A Fundamental Approach To Ordinary Chemistry PDF
A Fundamental Approach To Ordinary Chemistry PDF
A Fundamental Approach To Ordinary Chemistry PDF
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Contents
1. Introduction to chemistry...........................................................................................
1.1. Definition of chemistry.......................................................................................
1.2. Applications of chemistry....................................................................................
1.3. A laboratory.....................................................................................................
1.4. Uses of a laboratory...........................................................................................
1.5. Precautions/ safety measures/rules and regulations taken in the laboratory........................
1.6. Apparatus used in the laboratory...........................................................................
1.7. Apparatus and their uses......................................................................................
1.7.1. The Bunsen burner.........................................................................................
1.7.2. Flames of the Bunsen burner.............................................................................
1.7.2.1. Luminous flame..........................................................................................
1.7.2.2. Characteristics of a luminous flame..................................................................
1.7.2.3. Non luminous flame....................................................................................
1.7.2.4. Characteristics of a non luminous flame............................................................
1.7.2.5. Parts of a non luminous flame........................................................................
1.7.2.6. Differences between luminous and non-luminous flame........................................
1.7.2.7. Steps followed when lighting a Bunsen burner....................................................
2. States of matter.......................................................................................................
2.1. Solids................................................................................................................
2.2. Liquids..............................................................................................................
2.3. Gases................................................................................................................
2.4. Change of state....................................................................................................
2.5. The Kinetic Particle Theory of Matter........................................................................
2.6. Diffusion............................................................................................................
2.6.1. Diffusion of gases..............................................................................................
2.6.2. Diffusion of liquids............................................................................................
2.6.3. Factors Affecting Rate of Diffusion........................................................................
2.6.3.1. Temperature..................................................................................................
2.6.3.2. Mass of particles............................................................................................
3. Chemical and physical changes...................................................................................
3.1. Differences between physical and chemical changes......................................................
4. Solutions, crystals, compounds and mixtures...................................................................
4.1. Solutions............................................................................................................
4.2. Solute................................................................................................................
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4.3. Solvent..............................................................................................................
4.4. Types of solutions.................................................................................................
4.4.1. A saturated solution...........................................................................................
4.4.2. A super saturated solution....................................................................................
4.4.3. A suspension....................................................................................................
4.5. Differences between solutions and suspensions.............................................................
4.6. Crystals.............................................................................................................
4.7. How to grow a large crystal of copper (II) sulphate........................................................
4.8. Water of crystallization..........................................................................................
5. Compounds and mixtures..........................................................................................
5.1. Compound..........................................................................................................
5.2. Mixture..............................................................................................................
5.3. Differences between mixtures and compounds.............................................................
5.4. Methods of separation of mixtures............................................................................
5.5. The methods of separation of mixtures include the following...........................................
5.5.1. Filtration.........................................................................................................
5.5.2. Crystallisation & Evaporation to Dryness................................................................
5.5.3. Evaporation to dryness........................................................................................
5.5.4. Decanting........................................................................................................
5.5.5. Using separating funnel......................................................................................
5.5.6. Distillation......................................................................................................
5.5.6.1. Simple Distillation..........................................................................................
5.5.6.2. Separation of a mixture of water and ethanol by distillation.......................................
5.5.6.3. Fractional distillation......................................................................................
5.5.7. Sublimation.....................................................................................................
5.5.8. Magnetic Attraction...........................................................................................
5.5.9. Chromatography...............................................................................................
6. Elements, compounds, atoms and symbols.....................................................................
6.1. An element.........................................................................................................
6.2. An atom.............................................................................................................
6.3. A molecule.........................................................................................................
6.4. Radicals.............................................................................................................
6.5. Chemical symbols................................................................................................
6.6. Metals and non-metals...........................................................................................
6.7. Properties of metals and non-metals..........................................................................
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7. Atomic structure and the periodic table..........................................................................
7.1. An atom.............................................................................................................
7.2. Composition of an atom.........................................................................................
7.3. Structure of an atom..............................................................................................
7.4. Atomic number....................................................................................................
7.5. Atomic mass.......................................................................................................
7.6. Electronic configuration.........................................................................................
7.7. Electronic configuration of ions................................................................................
7.8. Isotopes.............................................................................................................
8. The periodic table....................................................................................................
8.1. Chemical families.................................................................................................
8.2. Bonding.............................................................................................................
8.3. Ionic/Electrovalent Bonding....................................................................................
8.4. Covalent Bonding.................................................................................................
8.5. Dative bonding....................................................................................................
8.6. Metallic Bonding..................................................................................................
9. Valency.................................................................................................................
9.1. Elements and radicals with their valencies...................................................................
9.2. Radicals.............................................................................................................
10. Chemical formulae...............................................................................................
10.1. Writing chemical formulae...................................................................................
10.2. Calculating the number of atoms of elements in a compound........................................
11. Chemical equations...............................................................................................
11.1. Balanced chemical equations................................................................................
11.2. Balancing chemical equations...............................................................................
12. Types of chemical reactions.....................................................................................
12.1. Direct combination or direct synthesis....................................................................
12.2. Simple decomposition........................................................................................
12.3. Simple replacement...........................................................................................
12.4. Double replacement...........................................................................................
13. The atmosphere and combustion...............................................................................
13.1. Air is a mixture of gases......................................................................................
13.1.1. Oxygen...........................................................................................................
13.1.2. Uses of oxygen.................................................................................................
13.1.3. Nitrogen.........................................................................................................
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13.1.4. Carbon dioxide.................................................................................................
13.1.5. Noble gases.....................................................................................................
13.1.6. Uses of noble gases............................................................................................
13.1.7. Water vapour....................................................................................................
13.2. Hygroscopic, deliquescent and efflorescent substances................................................
13.2.1. Hygroscopic substances......................................................................................
13.2.2. Deliquescent substances......................................................................................
13.2.3. Efflorescent substances.......................................................................................
13.2.4. Water of crystallisation.......................................................................................
13.3. Drying agents...................................................................................................
13.4. Burning substances in air.....................................................................................
13.6. Rusting...........................................................................................................
13.7. Combustion.....................................................................................................
14. Methods of gas collection.......................................................................................
15. Oxygen..............................................................................................................
15.1. Oxides............................................................................................................
16. Oxidation and reduction.........................................................................................
17. Water and hydrogen............................................................................................ 101
17.1. Sources of water............................................................................................. 101
17.2. Properties of pure water.................................................................................... 101
17.3. Test for water................................................................................................. 102
17.4. Purification of water........................................................................................ 102
17.5. Reactions of metals with water............................................................................103
17.6. Uses of water................................................................................................. 105
17.7. Reactivity Series............................................................................................. 106
18. Hydrogen......................................................................................................... 109
19. Acids, bases and salts.......................................................................................... 113
19.1. Acids........................................................................................................... 113
19.1.1. Common Acids............................................................................................... 113
19.1.2. Laboratory acids: 3 common laboratory acids/Mineral acids.......................................113
19.1.3. Basicity of an acid........................................................................................... 113
19.1.4. Some Acids with Their Basicity..........................................................................114
19.1.5. Strong and Weak Acids..................................................................................... 114
19.1.6. Comparing Strong and Weak Acids with Concentrated and Dilute Acids........................115
19.1.7. Properties of Dilute Acids.................................................................................. 115
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19.1.8. Storage of Acids.............................................................................................. 116
19.1.9. Uses of Acids................................................................................................. 116
19.1.10. Acids and Hydrogen Ions...............................................................................116
19.2. Bases and Alkalis............................................................................................ 117
19.2.1. Properties of Alkalis......................................................................................... 117
19.2.2. Neutralisation reactions..................................................................................... 117
19.2.3. Preparation of bases......................................................................................... 118
19.2.4. Strong and Weak Bases..................................................................................... 118
19.2.5. Uses of Alkalis............................................................................................... 119
19.3. Indicators and pH............................................................................................ 119
19.3.1. pH............................................................................................................... 119
19.3.2. pH scale........................................................................................................ 119
19.4. Indicators...................................................................................................... 120
19.5. Measuring pH of a Solution...............................................................................120
19.6. Ionic Equations............................................................................................... 121
19.7. Salts............................................................................................................ 123
19.7.1. Types of salts................................................................................................. 124
19.7.2. Preparation of Salts.......................................................................................... 124
19.7.3. Soluble and Insoluble Salts................................................................................ 125
19.7.4. Preparation of Insoluble Salts............................................................................. 125
19.7.5. Preparation of Soluble Salts...............................................................................127
19.7.6. Action of heat on salts...................................................................................... 132
19.8. Solubility of salts............................................................................................ 134
19.8.1. Solubility curves............................................................................................. 135
19.8.2. Uses of solubility............................................................................................ 136
19.9. Determination of solubility of salts......................................................................137
20. Carbon and its compounds.................................................................................... 139
20.1. Oxides of carbon............................................................................................. 145
20.1.1. Carbon dioxide............................................................................................... 145
20.1.2. Carbon monoxide............................................................................................ 152
20.2. Carbonates and hydrogen carbonates....................................................................160
20.3. Calcium Oxide (quicklime)................................................................................ 164
20.4. Sodium carbonate (soda ash).............................................................................. 166
20.5. The Carbon cycle............................................................................................ 169
20.6. Hardness of water............................................................................................ 173
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21. Electrolysis....................................................................................................... 176
21.1. Laws of electrolysis......................................................................................... 191
21.2. Application of electrolysis.................................................................................196
22. Formulae, stoichiometry and the mole concept...........................................................206
22.1. Relative Atomic Mass....................................................................................... 206
22.2. Percentage Composition....................................................................................208
22.3. Calculating the Mass of an Element in a Compound.................................................208
22.4. Calculating the Mass of Water in a Compound........................................................209
22.5. Mole............................................................................................................ 209
22.6. Molar Mass................................................................................................... 210
22.7. Different Kinds of Chemical Formulae..................................................................211
22.8. Calculating the Empirical Formula from Percentage Composition................................212
22.9. From Empirical formula to Molecular Formula.......................................................212
22.10. Molar Volume of Gases..................................................................................... 216
22.11. Calculations using chemical equations..................................................................217
22.11.1. Constructing Chemical Equations.....................................................................217
22.11.2. Calculations from Equations........................................................................... 218
22.11.2.1. Reacting Masses....................................................................................... 218
22.11.2.2. Reacting Masses and Volumes......................................................................220
22.11.2.3. Calculations involving energy changes...........................................................222
22.12. Concentration of Solutions................................................................................. 222
22.13. Quantitative analysis........................................................................................ 225
22.14. Uses of Titrations in Analysis.............................................................................228
22.14.1. Identification of Acids and Alkalis....................................................................228
22.14.3. Determination of basicity of an acid..................................................................231
22.14.4. Determination of water of crystallisation in oxalic acid (COOH) 2. xH2O.....................232
22.15. Calculations of volume of solutions......................................................................235
22.17. Gay Lussacs law............................................................................................ 236
22.18. Gas laws....................................................................................................... 237
23. Qualitative analysis............................................................................................. 239
24. Sulphur and its compounds................................................................................... 248
24.1. Sulphur......................................................................................................... 248
24.2. Extraction of sulphur by the Fraschs process..........................................................248
24.3. Extraction of sulphur from natural gas..................................................................249
24.4. Uses of sulphur............................................................................................... 249
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24.5. Allotropes of sulphur........................................................................................ 249
24.5.1. Rhombic sulphur............................................................................................. 250
24.5.2. Monoclinic sulphur.......................................................................................... 250
24.5.3. Amorphous sulphur......................................................................................... 251
24.5.4. Plastic sulphur................................................................................................ 251
24.5.5. Colloidal sulphur............................................................................................. 251
24.6. Properties of sulphur........................................................................................ 252
24.6.1. Physical properties.......................................................................................... 252
24.6.2. Chemical properties of sulphur........................................................................... 252
a. Action of heat on sulphur (in absence of air).................................................................252
b. Combustion of sulphur (in a plentiful supply of air)........................................................252
c. Reaction with metals and non-metals..........................................................................252
d. Action of acids on sulphur....................................................................................... 253
24.7. Oxides of sulphur............................................................................................ 253
24.7.1. Sulphur dioxide.............................................................................................. 253
24.7.1.1. Laboratory preparation of sulphur dioxide..........................................................253
24.7.1.2. Properties of sulphur dioxide...........................................................................254
24.7.1.3. Uses of sulphur dioxide.................................................................................255
24.7.2. Sulphur trioxide.............................................................................................. 255
24.7.2.1. Preparation of sulphur trioxide.........................................................................255
24.8. Sulphuric acid................................................................................................ 256
24.8.1. Industrial manufacture of sulphuric acid by the contact process...................................256
24.8.2. Properties of sulphuric acid................................................................................ 257
24.8.2.1. Physical properties....................................................................................... 257
24.8.2.2. Chemical properties...................................................................................... 258
24.8.3. Uses of sulphuric acid...................................................................................... 259
24.9. Sulphates...................................................................................................... 260
24.10. Hydrogen sulphide.......................................................................................... 262
24.10.1. Laboratory preparation of hydrogen sulphide.......................................................262
24.10.2. Testing for hydrogen sulphide..........................................................................262
24.10.3. Properties of hydrogen sulphide.......................................................................263
24.10.3.1. Physical properties.................................................................................... 263
24.10.3.2. Chemical properties................................................................................... 263
25. Nitrogen and its compounds..................................................................................265
25.1. Nitrogen....................................................................................................... 265
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25.2. Laboratory preparation of nitrogen from air............................................................265
25.3. Test for nitrogen.............................................................................................. 267
25.4. Properties of nitrogen....................................................................................... 267
25.5. Uses of nitrogen.............................................................................................. 268
25.6. Nitrogen monoxide (nitrogen(II) oxide).................................................................269
25.6.1. Laboratory preparation of nitrogen monoxide.........................................................269
25.6.2. Tests for nitrogen monoxide............................................................................... 269
25.6.3. Properties of nitrogen monoxide..........................................................................270
25.7. Nitrogen dioxide............................................................................................. 270
25.7.1. Laboratory preparation of nitrogen dioxide.............................................................270
25.7.2. Properties of nitrogen dioxide............................................................................. 271
25.8. Ammonia...................................................................................................... 272
25.8.1. Laboratory preparation of ammonia......................................................................272
25.8.2. Industrial preparation of ammonia (Haber process)...................................................273
25.8.3. Tests for ammonia........................................................................................... 273
25.8.4. Properties of ammonia...................................................................................... 274
25.8.5. Solubility of ammonia in water...........................................................................274
25.8.6. Experiment to demonstrate the high solubility of ammonia gas in water.........................274
25.8.7. Action of ammonia on copper (II) oxide................................................................275
25.8.8. Combustion of ammonia................................................................................... 276
25.8.9. Reaction with hydrogen chloride.........................................................................277
25.8.10. Reaction with chlorine................................................................................... 277
25.8.11. Uses of ammonia......................................................................................... 278
25.8.12. Ammonia solution........................................................................................ 279
25.8.12.1. Preparation of ammonia solution...................................................................279
25.8.13. Ammonium salts.......................................................................................... 279
25.8.13.1. Nitrogenous fertilizers................................................................................ 279
25.8.13.2. Effect of heat on ammonium salts..................................................................280
25.8.13.3. Test for ammonium salts.............................................................................281
25.8.14. Reactions of ammonia solution and sodium hydroxide solution................................282
25.8.15. Nitric acid.................................................................................................. 284
25.8.15.1. Laboratory preparation of nitric acid..............................................................284
25.8.15.2. Industrial preparation of nitric acid................................................................285
25.8.15.3. Uses of nitric acid..................................................................................... 285
25.8.15.4. Properties of nitric acid............................................................................... 286
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25.8.15.5. Nitric acid acting as a strong acid..................................................................286
25.8.15.6. Nitric acid as an oxidizing agent...................................................................287
25.9. Nitrates......................................................................................................... 288
25.9.1. Action of heat on nitrates...................................................................................288
25.9.2. Test for nitrates............................................................................................... 290
26. Chlorine and its compounds..................................................................................298
26.1. Chlorine........................................................................................................ 298
26.1.1. Laboratory preparation of chlorine.......................................................................298
26.1.2. Industrial manufacture of chlorine.......................................................................299
26.1.3. Properties of chlorine....................................................................................... 300
26.1.4. Tests for chlorine............................................................................................. 306
26.1.5. Uses of chlorine.............................................................................................. 306
26.2. Hydrogen chloride........................................................................................... 306
26.2.1. Laboratory preparation of hydrogen chloride..........................................................306
26.2.2. Test for hydrogen chloride................................................................................. 307
26.2.3. Properties of hydrogen chloride...........................................................................307
26.3. Hydrochloric acid............................................................................................ 307
26.3.1. Preparation of hydrochloric acid..........................................................................307
26.3.2. Properties of hydrochloric acid........................................................................... 308
26.3.3. Uses of hydrochloric acid.................................................................................. 309
26.3.4. Properties of hydrogen chloride in methylbenzene....................................................309
26.4. Testing for soluble chloride................................................................................309
27. Extraction of metals............................................................................................ 313
27.1. Introduction................................................................................................... 313
27.2. Concentration of ores....................................................................................... 313
27.3. Sodium......................................................................................................... 314
27.3.1. Extraction of sodium........................................................................................ 314
27.3.2. Uses of sodium metal....................................................................................... 315
27.4. Copper......................................................................................................... 315
27.4.1. Extraction of copper......................................................................................... 316
27.4.2. Concentration of the ore.................................................................................... 316
27.4.3. Roasting and reduction..................................................................................... 316
27.4.4. Refining of the impure copper............................................................................ 317
27.4.5. Uses of copper................................................................................................ 318
27.5. Iron............................................................................................................. 318
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27.5.1. Extraction of iron............................................................................................ 318
27.5.2. Casting iron (pig-iron)...................................................................................... 319
27.5.3. Wrought iron.................................................................................................. 319
27.5.4. Steel............................................................................................................ 320
27.5.5. Recycling of metals......................................................................................... 321
27.5.6. Alloy........................................................................................................... 321
28. Organic chemistry.............................................................................................. 324
28.1. Introduction................................................................................................... 324
28.2. Hydrocarbons................................................................................................. 324
28.3. Homologous series.......................................................................................... 324
28.4. Functional groups............................................................................................ 324
28.5. Alkanes........................................................................................................ 325
28.5.1. General properties of alkanes..............................................................................330
28.5.1.1. Physical properties....................................................................................... 330
28.5.1.2. Chemical properties...................................................................................... 330
28.6. Petroleum (crude oil)........................................................................................ 333
28.6.1. Fractional distillation of petroleum.......................................................................333
28.6.2. Cracking....................................................................................................... 335
28.6.3. Bio gas......................................................................................................... 335
28.6.4. Disadvantages of bio gas production.....................................................................336
28.6.5. Alkenes........................................................................................................ 336
28.7. Ethene.......................................................................................................... 337
28.7.1. Laboratory preparation of ethene.........................................................................337
28.7.1.1. Physical properties....................................................................................... 338
28.7.1.2. Chemical properties...................................................................................... 338
28.7.2. Uses of ethene................................................................................................ 340
28.7.3. Alkynes........................................................................................................ 341
28.7.4. Ethyne (acetylene)........................................................................................... 342
28.7.4.1. Physical properties of ethyne...........................................................................342
28.7.4.2. Chemical properties of ethyne.........................................................................342
28.7.4.3. Uses of ethyne............................................................................................. 343
28.7.5. Alkanols (alcohol)........................................................................................... 344
28.8. Ethanol (ethyl alcohol)...................................................................................... 344
28.8.1. Manufacture of ethanol..................................................................................... 344
28.8.2. Properties of ethanol........................................................................................ 346
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28.8.2.1. Chemical properties...................................................................................... 346
28.8.2.2. Uses of ethanol............................................................................................ 347
28.9. Carboxylic acids and esters................................................................................ 347
28.9.1. Properties of carboxylic acids............................................................................. 347
28.10. Esters........................................................................................................... 348
28.10.1. Uses of esters.............................................................................................. 349
28.11. Soap............................................................................................................ 349
28.11.1. Manufacture of soap..................................................................................... 349
28.11.2. The cleansing action of soap........................................................................... 349
28.11.3. Soapless (synthetic detergents)........................................................................350
28.11.4. Advantages of synthetic detergents over soap......................................................350
28.11.5. Advantages of soap over synthetic detergents......................................................351
28.12. Polymerisation............................................................................................... 351
28.13. Addition polymerization.................................................................................... 351
28.14. Polyethene..................................................................................................... 351
28.15. Polypropene................................................................................................... 352
28.16. Polyvinyl chloride (P.V.C).................................................................................352
28.17. Synthetic rubber.............................................................................................. 352
28.18. Condensation polymerization.............................................................................353
28.19. Types of polymers........................................................................................... 353
28.20. Advantages of synthetic polymers over natural polymers...........................................355
28.21. Disadvantages of synthetic polymers....................................................................355
29. Energy changes.................................................................................................. 360
29.1. Introduction................................................................................................... 360
29.2. Enthalpy....................................................................................................... 360
29.3. Exothermic and endothermic reactions..................................................................360
29.4. Types of enthalpy changes................................................................................. 362
29.4.1. Determination of enthalpy (heat) of combustion of ethanol.........................................363
29.5. Determination of the enthalpy (heat) of solution o f sodium chloride.............................366
29.6. Enthalpy of neutralization.................................................................................. 367
29.7. Determination of heat of neutralization.................................................................367
29.8. Enthalpy of displacement..................................................................................369
29.8.1. Determination of the enthalpy (heat) of displacement of the reaction between copper(II)
sulphate solution and zinc.............................................................................................. 369
30. Rate of reaction and equilibrium............................................................................. 374
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30.1. Rate of reaction.............................................................................................. 374
30.1.1. Determination of rate of reaction.........................................................................374
30.1.2. Determination of rates of reaction by measuring the volume of the gas evolved with time
374
30.1.3. Effects of temperature on the rate of reaction..........................................................378
30.1.3.1. Investigation of the effect of temperature on the rate of reaction...............................378
30.1.4. Effect of a catalyst on the rate of reaction...............................................................379
30.1.4.1. Investigation of the effect of catalyst on the rate of reaction.....................................379
30.1.5. Effect of surface area on the rate of reaction...........................................................381
30.1.5.1. Investigation of the effect of surface area on the rate of reaction...............................381
30.1.6. Effect of light on the rate of reaction.....................................................................382
30.1.6.1. Investigation of the effect of light on the rate of reaction.........................................382
30.1.7. Effect of pressure on the rate of reaction................................................................383
31. Equilibrium...................................................................................................... 383
31.1. Factors affecting equilibrium..............................................................................383
31.2 Trial questions for paper 1 and 2..384
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1. Introduction to chemistry
1.3. A laboratory
This is a specialized and organized room where scientific experiments are conducted
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When your experiment is completed, turn off the water supply and gas supply and disconnect
any electrical connections and return all the materials and apparatus in their proper places
15
Reagent bottles A glass container to hold
liquid chemicals
16
Bunsen burner The function of a Bunsen
burner is to heat substances.
17
b. Flames of the Bunsen burner
The Bunsen burner produces two different flames depending on whether the air hole is open
or closed.
A flame is a combination of burning gases giving out heat and light.
The Bunsen burner produces two different flames namely
1. Luminous flame
2. Non-luminous flame
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The flame produces soot (smoke) when burning
It has four zones which have different colours
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Has four zones Has three zone
Is quiet Is noisy
Is not hot enough Is very hot
Forms soot No soot is formed
Produced when air holes are closed Produced when air holes are open
Produces a lot of light Produces little light
1.7.b.7. Steps followed when lighting a Bunsen burner
Fix the Bunsen burner properly
Connect the Bunsen burner to the gas tap
Close the air holes by turning the metal rings (collar)
Put on the gas tap and switch on for a few seconds
Light the gas by applying a match stick on top of the chimney
Lastly open the air holes by turning on the metal rings
2. States of matter
Matter is anything that occupies space and has weight.
The states of matter are solids, liquids and gases
2.1. Solids
The characteristics of solids include:
- Fixed volume
- Fixed shape
- Incompressible
- Do not flow
2.2. Liquids
The characteristics of liquids include:
- Fixed volume
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- No fixed shape
-Takes the shape of the container
- Incompressible
- Flow easily
Have definite volume but no definite shape and take up the shapes of their containers e.g.
water, milk
Liquids cannot be compressed by squeezing
When liquids are heated their particles move faster and finally turn into gas
The temperature at which a liquid changes into a gas is called the boiling point e.g. the
boiling point for pure water is 100oC.
2.3. Gases
The characteristics of gases include:
-No fixed volume
- No fixed shape
- Compressible
- Flow in all direction
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Solid
Freezing
Evaporation
Gases Condensatio
Liquids
n
Sublimation is the process of changing a substance from a solid to a gas state without passing
through the liquid state when heated. Substances which sublime include: iodine crystals,
ammonium chloride and iron (III) chloride
Note: water exists in all the three states. When in solid state, it is called ice, in liquid state, it
is called water, and in gaseous state, it is called water vapour/steam. But at room temperature
water is in liquid state.
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1. Matter can be coloured (e.g. sulphur is yellow) but particles are not.
2. Substances feel hot/cold but particles dont get hot/cold. The temperature is due to speed of
movement of particles. If hot, particles move fast.
3. Matter expands when heated but particles dont. They increase distance between particles
during expansion.
Changes of State
Melting
Melting is change from solid to liquid by absorbing heat to break force of attraction holding
particles together.
The temperature at which solid melts is melting point.
Freezing
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Freezing is the change of liquid to solid by cooling down of liquid.
Freezing point is the temperature at which liquid freezes.
B C
24
- There are spaces between particles of matter; the amount of space varies between each
states
- The particles are constantly move; each state moves in different speed
2.6. Diffusion
Diffusion is the spreading and mixing of particles in gases and liquids.
25
2.6.2. Diffusion of liquids
CuSO4 crystals placed in beaker of water, blue particles of the crystals is spread throughout the
water to form a uniformly blue solution.
26
2.7.
2.7.1.1. Temperature
The higher the temperature, the more particles of matter absorb energy making them move faster,
the higher the rate of diffusion; the lower the temperature, the slower the rate of diffusion
Greater mass, the slower it diffuses; smaller mass, the faster it diffuses
Cotton wool soaked in aqueous ammonia and another soaked in hydrochloric acid are placed on
opposite sides of the tube. Ammonium hydroxide (NH4OH) vapor and hydrogen chloride (HCl)
vapor diffuses in the tube and a compound is produced inside the tube closer to HCl soaked
cotton as the particles are heavier. The greater the mass, the slower particles diffuse and the
smaller the mass, the faster particles diffuse.
3. Chemical and physical changes
A chemical change is a process that occurs and a new substance is formed e.g.
Physical change is a change in which no new substance is formed e.g. freezing (water to ice)
Magnetization of iron
Melting of ice
No energy is evolved or absorbed during the Energy is absorbed of evolved during the
reaction reaction
Where M-Mass
I-Irreversible
S-Substance
E-energy
4.1. Solutions
A solution is a uniform mixture of two or more substances. When sugar is added to water and
stirred, the sugar dissolves in water.
In this process sugar is called a solute, water is called a solvent and a mixture of sugar and water
is called a solution.
4.2. Solute
4.3. Solvent
This is a solution which cannot dissolve any more solute at a given temperature in presence of
undissolved solute.
This is a solution which contains more solute than it can hold at a given temperature
4.4.3. A suspension
This is a liquid containing small particles of a solid which are spread throughout it and settle on
standing e.g. a solution of chalk in water.
A suspension contains solid particles which can be seen but a solution contains no solid particles.
In suspension the mixture can be separated by filtration while in solution the mixture cannot be
separated by filtration.
In suspension the solid particles settle on standing but in solutions no solid particles can settle on
standing.
4.6. Crystals
Crystals are formed when hot solutions cool. Is a hot saturated solution cools rapidly, crystals
formed are many and small but once cooled slowly crystals formed are few and big in size
Substances which form crystals are called crystalline substances i.e. sodium chloride and copper
(II) sulphate. Some solids do not form crystals e.g. charcoal and glass; and these are called non-
crystalline substances.
Choose one good crystal of copper (II) sulphate and tie a thin thread around it. Hang it in the
solution and place the beaker in a warm place for several days. The crystals grow large and the
solvent slowly evaporates.
4.8. Water of crystallization
This is a definite amount of water some substances chemically combine with when they form
crystals from their solutions in water
Some substances do not contain water of crystallization and are said to be anhydrous e.g.
sodium chloride (NaCl), potassium nitrate (KNO3) etc.
5.1. Compound
Water (H2O): this is a compound made up of hydrogen (H) and oxygen (O) as elements
Common salt sodium chloride (NaCl): this is a compound made up of sodium (Na) and
chlorine (Cl)
Glucose (C6 H12O6): this is a compound made up of carbon (C), hydrogen (H) and oxygen (O)
Iron (II) sulphide (FeS): this is a compound made up of iron (Fe) and sulphur (S)
5.2. Mixture
A mixture is a substance which consists of two or more elements or compounds not chemically
combined together.
Examples include: Salt and water, Salt and sand, Water and alcohol, Chalk and water, Air
Mixture Compound
Can be Separated by physical means e.g. Substances in it cant be separated by physical
filtration means
Physical Properties of mixtures e.g. colour and Physical properties of compounds are quite
density are the average of constituent different from those elements in them
substances
Energy is not usually given out or absorbed Energy is given out or absorbed when a
when a mixture is formed compound is made
Its Composition is variable, the substances can Its composition is not variable; the elements
be combined in any proportion by mass are combined in definite proportions by mass
Where S- Separated
P- Properties
E- Energy
C- Composition
There are two types of liquid mixtures i.e. immiscible and miscible liquid mixtures
Miscible liquids:
These are liquids which mix freely and form one layer
Immiscible liquids
These are liquids which do not mix easily and form more than one layer.
5.5.1. Filtration
Filtration separates insoluble solid from a liquid.
- Mixture is poured through a filter paper with tiny holes
- Large solid particles cannot pass through the pores and trapped in it as residue while tiny liquid
particles pass through as filtrate.
Common salt and sand are placed in a beaker and water is then added. The mixture is warmed
gently while stirring until the salt completely dissolves. Salt dissolves but sand does not dissolve.
The solution is then filtered. After filtering, the salt solution is obtained separate from sand. The
salt solution is therefore called a filtrate and sand, the residue.
The salt solution is then poured into an evaporating basin. The water evaporates when the salt
solution is heated. It evaporates completely, leaving salt crystals behind on the evaporating dish.
5.5.4. Decanting
This method can also be used to separate a mixture of sand and salt in water
Put the mixture of sand and salt in the beaker and add water. Stir the mixture. Salt will dissolve
but sand will not. Leave the beaker to stand for a few minutes for the sand to settle.
The sand will settle on the bottom of the salt solution. Pour off the salt solution carefully without
disturbing the mixture.
The sand will be left on the bottom of the beaker. Then evaporate the salt solution to dryness.
NB:
This method is not as good as filtration and should always be discouraged in the laboratory
Immiscible liquids are those which do not mix at all. Such liquids separate into distinct layers
according to their densities. Examples include:
Water and paraffin
To separate such liquids using a separating funnel, the following procedures are followed.
The mixture of paraffin/oil and water is poured into a separating funnel. Shake vigorously and
then allow to settle
Results
Water has a high density than paraffin and therefore water separates to the bottom and paraffin
goes on top of water making different layers
The tap is opened and the water layer runs out first
Close the tap and put another container where paraffin will be collected. Open the tap to collect
paraffin in a separate container
Separating funnel
5.5.6. Distillation
This is the process of heating a liquid to form vapour and then cooling it back to form a liquid
During distillation, the solution is heated so that its liquid component boils and escapes as a
vapour. The vapour is then cooled by running water and condensed into liquid called the
distillate.
To achieve even boiling and preventing too much bumping (frothing and bubbling) in the
flask, anti-bumping granules or boiling chips is added to the distilling flask containing the
mixture.
The distillation apparatus can also be used to determine the boiling point of a liquid. This is done
by use of a thermometer. The thermometer is passed through one holed stopper (cork) of the
flask and the temperature of the vapour is noted. This will be the boiling point of the liquid.
NB: The impurities which contain dissolved salts remain in the flask and therefore pure water is
not good for drinking because it lacks mineral salts.
Water and ethanol are two miscible liquids and they have different boiling points, they are
therefore separated by fractional distillation
The boiling point of pure water is 100oC and that of ethanol is 78oC
The mixture of ethanol and water is poured into a boiling flask fitted with a fractionating
column. Glass beads are placed in the fractionating column. The use of the glass beads in the
fractionating column is to increase the surface area for effective separation of vapours
The mixture is then heated up to 78oC; temperature reading on the thermometer becomes steady
for some time. Ethanol vapourises and ethanol vapour is condensed by cold water flowing in the
condenser forming a liquid. The liquid which is received now is pure ethanol and is then
collected in the receiving flask.
Heating is continued up to 100o when water evapourates and is collected in a different container
The use of the thermometer is to record temperature of the two vapours
The use of porcelain pieces which are placed in the distilling flask is to enable the mixture boil
gently
This is a process of separation of two or more liquids with different boiling points into different
fractions.
Fractional distillation is the process used to separate liquids which are miscible
Miscible liquids are those which mix freely in all proportions to form one uniform solution
The process can therefore be used to separate
Crude oil is separated into its constituents such as petrol, diesel and kerosene
Fractional distillation of liquefied air
5.5.7. Sublimation
This is the process where when a solid is heated; it changes to a gas directly without passing
through the liquid state.
Separation of mixtures by sublimation
Sublimation can be used to separate a mixture of two substances where one sublimes and the
other does not. It can therefore be used to separate a mixture of iodine and common salt (sodium
chloride)
When a mixture of iodine and common salt (sodium chloride) is heated, iodine changes to a gas
and common salt remains in solid form. Therefore the two substances can be separated from one
another.
Ammonium chloride
Anhydrous aluminum chloride
Benzoic acid
In hospitals, magnets are often used to remove iron splinters from a patients eyes.
Electromagnets are also used for removing scrap steel and iron at the junkyard. These scrap
metal can then be sent for recycling.
Magnetic substances are substances which can be attracted by a magnet e.g. iron. Therefore a
magnet can be used to separate iron from sulphur because iron is attracted by a magnet and
sulphur cannot be attracted by a magnet and is therefore left behind.
5.5.9. Chromatography
When a spot of ink is applied to the chromatography paper (usually filter paper), the dyes in the
ink are attracted to the surface of the paper. The chromatography paper is then immersed in
a solvent. The solvent level should not be above the ink spot.
As the solvent (usually water or ethanol) is soaked up by the paper, the solvent dissolves the
dyes.
A dye that is strongly attracted to the paper and not very soluble in the solvent will be left
behind. A dye that is weakly attracted to the paper and very soluble in the solvent will move up
with the solvent through the paper.
6. Elements, compounds, atoms and symbols
6.1. An element
An element is a substance which cannot be split up into two or more simpler particles by
chemical means. Examples of elements include
Copper
Sulphur
Carbon
Oxygen
Hydrogen
Iron
6.2. An atom
An atom is the smallest, indivisible particle of an element which can take part in a chemical
reaction.
6.3. A molecule
A molecule is the smallest indivisible particle of an element which can exist in free and separate
state. Molecules are formed when two or more atoms combine together e.g. water (H2O) is a
molecule made up of 2 atoms of hydrogen and 1 atom of oxygen
6.4. Radicals
A radical is a group of atoms which cannot exist on their own but exists in a compound.
Examples include:
Sulphate (SO4)
Carbonate (CO3)
Nitrate (NO3)
Sulphate radical cannot exist on its own but can exist in composition like in sulphuric acid
(H2SO4), calcium sulphate (CaSO4), sodium sulphate (Na2SO4)
A chemical symbol of an element is one or two letters which represent one atom of an element.
The letters used are the first letters of an element in English or Latin names of the element. The
first letter should be CAPITAL and the second letter should be small.
Element Symbol
Hydrogen H
Helium He
Lithium Li
Beryllium Be
Boron B
Carbon C
Nitrogen N
Oxygen O
Fluorine F
Neon Ne
Sodium Na
Magnesium Mg
Aluminum Al
Silicon Si
Phosphorous P
Sulphur S
Chlorine Cl
Argon Ar
Potassium K
Calcium Ca
Scandium Sc
Titanium Ti
Vanadium V
Chromium Cr
Manganese Mn
Iron Fe
Cobalt Co
Nickel Ni
Copper Cu
Zinc Zn
Lead Pb
Mercury Hg
Silver Ag
Gold Au
Xenon Xe
Iodine I
Barium Ba
Elements whose symbols were derived from their Latin names are summarized below
Element Latin names Symbol
Potassium Kalium K
Sodium Natrium Na
Iron Ferrum Fe
Copper Cupium Cu
Lead Plumbium Pb
Mercury Hydrogyrum Hg
Silver Argentum Ag
Metals: Copper (Cu), Iron (Fe), Sodium (Na), Potassium (K), Mercury (Hg), Silver (Ag), Gold
(Au) etc
Non-metals: Chlorine (Cl), Carbon (C), Oxygen (O), Hydrogen (H), Helium (He), Nitrogen (N),
Fluorine (F), Bromine (Br) etc..
A non-metal: is an element which forms negative ions by gaining electrons e.g. Cl-, Br- etc
6.7. Properties of metals and non-metals
Metals Non-metals
Good conductors of heat and electricity Poor conductors of heat and electricity
Have high density Have low density
Solids with high melting points Most are gases. Solids have low melting points
Ductile i.e. can be drawn into wires Not ductile
Malleable i.e. can be made into sheets Not malleable
Strong and tough i.e. have high tensile strength Not strong i.e. have low tensile strength
Are rustrous i.e can be polished Not rustrous i.e. cannot be polished
7. Atomic structure and the periodic table
7.1. An atom
An atom is the smallest indivisible particle of an element which can take part in a chemical
reaction.
An atom is mainly composed of three particles namely:- electrons, protons and neutrons
Electrons
Protons
Mass (1)
Neutrons
Have no charge.
Mass (1)
NB: Neutrons are symbolized by N. Both neutrons and protons are found inside the nucleus and
are together called nucleons.
The central part of an atom is called the nucleus. Protons and neutrons are located in the nucleus.
These make up nucleon number.
In a neutral atom the number of electrons is equal to the number of protons. The charge of an
electron is equal but opposite to that of a proton.
This is the number of protons in the nucleus of an atom. Since protons are equal to electrons in
an atom, atomic number can also be defined as the number of electrons found in an atom.
This is the number of protons plus the number of neutrons in the nucleus of an atom. It is the
sum of protons and neutrons in an atom.
Example
16
1. An atom of an element is represented by the symbol 8 X . State
23
2. Work out the following, 11 Na . State
40
Calcium 20 Ca
36
Argon 18 Ag
35
Chlorine 17 Cl
In an atom, electrons occupy levels or shells and move in these shells around the nucleus. The
number of shells in an atom depends on the number of electrons present. The maximum number
of electrons occupying the innermost shell is 2. The rest of the shells contain a maximum of 8
electrons.
To write electronic configuration we write as n:n:n.... where first n denotes the first shell, second
the second shell and so and so for.
Examples
Draw the structure of the following atoms and state their electronic configuration.
2:6 (Oxygen)
2:8:7 (Chlorine)
2:8:6 (Sulphur)
5.
40
a. Draw the electronic structure of element X represented by 20 X
b. State the number of neutrons and protons found inside the nucleus of X
An ion is an electrically charged particle. Ions are formed by gain or loss of electrons.
Metals form positively charged ions by losing electrons and non-metals form negatively charged
ions by gaining electrons.
The number of electrons gained or lost is equal to the valence of the atom e.g.
1. Sodium electronic configuration 2:8:1 loses 1 electron to form Na+. Na+ therefore has
electronic configuration of 2:8
2. Calcium electronic configuration 2:8:8:2 loses 2 electrons to form Ca2+. Ca2+ therefore
has electronic configuration of 2:8:8
3. Oxygen electronic configuration 2:6 gains 2 electrons to form O2-. O2- therefore has
electronic configuration of 2:8
4. Chlorine electronic configuration 2:8:7 gains 1 electron to form Cl-. Cl- therefore has
electronic configuration of 2:8:8
Exercise
The number of electrons, protons and neutrons in atoms W, X, Y and Z are shown in the table
below.
Atom Electrons Protons Neutrons
W 8 8 8
X 16 16 16
Y 13 13 14
Z A 3 4
a. Determine
i. The value of A
i. W
ii. W2-
iii. Y
iv. Y3+
c. State the two atoms that are isotopes of the same element
7.8. Isotopes
Protons 1 1 1
Electrons 1 1 1
Neutrons 0 1 2
Mass No. 1 2 3
Protons 17 17
Electrons 17 17
Neutrons 18 20
Mass No. 35 37
23 24
11 Na and 11 Na
24 25
12 Mg and 12 Mg
Examples
35 37
1. The relative abundances of 17 Cl and 17 Cl are 75% and 25% respectively by weight.
35 37
Relative atomic mass of chlorine = Mass due to 17 Cl + Mass due to 17 Cl
75 X 35 25 X 37
= 100 + 100
=35.5
Exercise
13 12 13
1. Carbon has two main isotopes C
6 and C .
6 C
6 has relative abundance of 1.11%
12
and C
6 has relative abundance 98.89%. Calculate the relative atomic mass of carbon.
23
2. Given the sodium atom, 11 Na
There over 103 elements so far discovered by scientists. The periodic table is a table in which all
the elements so far discovered are put. It is an arrangement of elements in order of their atomic
numbers.
Group
An element is put in a particular group depending on the number of electrons it has in its last or
outermost orbital e.g oxygen 2:6 has 6 electrons in its last orbital so its in group 6.
Sodium 2:8:1 has 1 electron in its outermost shell, so its put in group 1 etc.
Period
An element is put in a particular period depending on the number of orbitals it has e.g. oxygen 2:
6; Has 2 orbitals so it is period 2 and group 6.
Note
All elements in the same group have the same valency and similar chemical properties
Elements in group I III are metals
Elements in group I are alkaline metals while elements in group II are called alkaline earth
metals.
Elements in group VII are halogens while elements in group VIII are Noble gases or inert
gases.
The number of electrons in the outermost shell corresponds to the group number.
For elements in groups I-IV the valency is given by the group number
Hydrogen is put in groups I and VII because it behaves as both a group I and group VII element
Helium is put in group VIII because it has a fully filled outermost orbital and behaves like a
group VIII element.
8.1. Chemical families
On the left of the periodic table there are metals, and to the right lie the nonmetals.
In the middle are metalloids and these exhibit both metallic and nonmetallic properties.
Metals are malleable, ductile, and have luster; most of the elements on the periodic table are
metals. They oxidize (rust and tarnish) readily and form positive ions (cations).
Transition metals (also called the transition elements) are known for their ability to refract light
as a result of their unpaired electrons.
The actinides and lanthanides are collectively called the rare earth elements and are filling
the f orbitals.
They are rarely found in nature. Uranium is the last naturally occurring element; the rest are
man-made.
Non-metals do not conduct electricity well because they do not have free electrons.
All the elemental gases are included in the nonmetals. Notice that hydrogen is placed with the
metals because it has only one valence electron, but it is a nonmetal.
Specific families
They are all metals which are highly electropositive i.e. they have a very high tendency to lose
electrons than other elements in the period. They form ions with a single charge by loss of one
electron e.g. Na+, Li+, K+.
They are larger than any other elements in the period e.g. lithium (Li)
They have low densities, low melting points and low conductivities
They are soft and shinny and hence can be cut with a knife
The most reactive metal family, these must be stored under oil because they react violently with
water
They dissolve and create an alkaline, or basic, solution, hence their name
a) Lithium
Has valency 1
Reactions
i. With water
Lithium reacts slowly with cold water to form its hydroxide and liberates hydrogen
Gives a red flame colour when burnt in air i.e. burns with a red flame
b) Sodium (Na)
Also a member of group 1
Occurrence
Widely distributed in nature as sodium chloride NaCl, sodium nitrate NaNO3 or sodium
carbonate Na2 CO3 etc
Reactions
i. With water
Sodium attacks cold water rapidly, evolving hydrogen and forming sodium hydroxide
The sodium oxide formed absorbs water i.e. deliquescent to form sodium hydroxide solution
These are also reactive metals, but they dont explode in water; pastes of these are used in
batteries.
These are elements, which have two electrons each in their outermost orbital. Examples include
magnesium (Mg), calcium (Ca)
a. Magnesium (Mg)
Atomic number 12
Extracted electronically
Reactions
i. With air
Damp air puts a layer of an oxide on magnesium later forming a hydroxide and carbonate
Burns in air with a bluish white flame leaving behind an oxide and nitride
All dilute mineral acids i.e. HCl, H2SO4, HNO3 react with magnesium liberating hydrogen
NB: Hot concentrated sulphuric acid yields sulphur dioxide instead of hydrogen gas
When heated, magnesium reacts with halogens to produce salts called halides
b. Calcium (Ca)
Extraction
Mainly by electrolysis of its fused calcium chloride
Properties of calcium
Reactions
i. With air
Calcium hydroxide formed combines with carbon dioxide to form calcium carbonate
Explanation
The calcium oxide (CaO) formed is deliquescent i.e. absorbs moisture from the atmosphere so it
forms calcium hydroxide Ca(OH)2 (aq). The calcium hydroxide Ca(OH)2 (aq) formed later
absorbs carbon dioxide (CO2), also from the atmosphere forming calcium carbonate (CaCO3)
NB: The calcium hydroxide produced is sparingly soluble and tends to precipitate on the metal
and stops any further reaction from going on.
Halogens (7A)Known as the salt formers, they are used in modern lighting and always exist
as diatomic molecules in their elemental form.
They include fluorine (F2), chlorine (Cl2), bromine (Br2) and iodine (I2)
Properties
Chlorine
Poisonous gas
Denser than air
Atomic number 17
Atomic mass 35
Reactions
i. With water
Halogens react with water forming acids. Chlorine reacts with water forming two acids i.e.
hypochlorous acid (HOCl) and hydrochloric acid (HCl)
Halogens generally react with NaOH to produce a pale yellow solution of sodium chloride or
sodium bromide
Halogens are oxidizing agents i.e. they accept electrons, therefore they are electronegative
elements
Noble gases (8A)Known for their extremely slow reactivity, these were once thought to never
react; neon, one of the noble gases, is used to make bright signs.
Sometimes referred to as Group O elements
They include helium (He), neon (Ne), argon (Ar), Krypton (Kr), xenon (Xe)
i. Helium (He)
Electronic configuration 2
Very stable because all the three orbitals are full i.e. chemically satisfied
NB: their electron arrangement makes them very unreactive i.e. they do not lose or gain electrons
and this accounts for their low reactivity
8.2. Bonding
Elements always try to achieve the stable structure of the noble gases. In doing so, they combine
chemically forming bonds.
Types of bonds
E.g. NaCl made up of 2 ions; positively charged Na, negatively charged Cl.
8.3. Ionic/Electrovalent Bonding
Ionic bonding is the transfer of electrons from one atom to another to become/achieve an inert
gas configuration, forming ions.
Ionic bonds are formed between METALLIC and NON- METALLIC ATOMS ONLY.
- Metals lose electrons to form positive ions (cations)
- Non-metals gain electrons to form negative ions (anions)
The formation of ions is resulted from transfer of electrons from one atom to another atom(s).
The ions produced are of opposite charges and unlike charges attract, causing them to be held
together with a strong electrostatic force.
E.g. Formation of NaCl
Sodium atom loses an electron by transferring the electron to chlorine atom, making both stable.
The loss of electron forms a cation, Na+, and the gain of an electron forms anion, Cl-. The
opposite charges acquired by both ions attract each other, forming a strong ionic bond of NaCl.
E.g. Formation of MgF2
Sodium atom loses two electrons by transferring the electrons to fluorine atoms, one each,
making both stable. The loss of electron forms a cation, Mg2+, as it loses 2 electrons, and the gain
of electron forms anion, F-. The opposite charges acquired by both ions attract to each other,
forming a strong ionic bond of MgF2.
Deducing formula of ionic compounds
We can know the charge of elements by looking at groups of periodic table. Group I to group III
elements have a charge of +1, increasing to +3, going to the right. Group V to group VII
elements have a charge of -3, decreasing to -1, going to the right.
E.g. Aluminium sulfate
We have to balance the charges to make a stable bond
Ions present: Al3+ SO42-
SO42-
Al3+ SO42-
Total change: 6+ 6-
Properties
1. Ionic compounds are hard crystalline solids with flat sides and regular shapes because the ions
are arranged in straight rows in strong ionic bonds.
2. Ionic compounds have very high melting points and boiling points.
3. The strong forces holding ionic compounds prevent them to evaporate easily. Hence, ionic
compounds have no smell.
4. Solid ionic compounds dont conduct electricity but they do when they are aqueous or molten.
This is because in liquid/aqueous state the ions which conduct electricity are free to move. In
solids, these ions are fixed in place.
5. Ionic compounds are soluble in water but insoluble in organic compounds. This is because the
ions attract water molecules which disrupts the crystal structure, causing them to separate and go
into solution. Vice versa is when in organic solvent.
The hydrogen atom has one valency. To become stable, the hydrogen atom needs one more
electron, just like helium which has 2 valency electrons. When 2 hydrogen atoms join, they share
their electrons, on which, the share becomes 2 electrons, which is now a noble gas configuration,
being shared between these 2 atoms. Write the bond as H H single bond, which means they
share an electron pair (2 electrons).
E.g. Chlorine (Cl2) molecule
The chlorine (Cl) atom has 7 valency electrons and needs one electron, each, to form a noble gas
configuration between two Cl atoms. Hence they share an electron EACH to share 2 electrons
between the atoms. Hence, each Cl atom now has 8 valency electrons which is a noble gas
configuration.
E.g.O2 molecule
An O atom has 6 valency electrons and needs 2 electrons, each, to form a noble gas
configuration. Hence, EACH SHARE THE AMOUNT OF ELECTRONS EACH IS SHORT OF,
in this case 2 electrons, to form a stable molecule. The contribution hence now becomes 4
electrons and what left on each oxygen atom are 4 electrons. Combine each 4 electrons on
oxygen atom with the 4 electrons shared and hence get 8
E.g. H2O molecule
Apart from oxygen sharing between oxygen atoms, it can share electrons with other atoms.
Oxygen needs 2 electrons and when bonded with hydrogen, which need an atom each, they
combine to provide 2 electrons on both sides of oxygen bonded with hydrogen atoms. Each
hydrogen atom with oxygen atom form a single bond: O H.
E.g. CO2 molecule
Carbon needs 4, oxygen needs 2. Share two from oxygen part, WHICH HAS THE SMALLEST
NUMBER OF SHORT OF ELECTRONS, TO SHARE THE AMOUNT OF ELECTRONS
THAT EACH ATOM NEEDS, to form 4 shared atoms. Now oxygen is stable but carbon needs 2
more, which it can get from another oxygen atom. The atoms are now stable and since each bond
has 2 pairs of electrons, this is a double bond: C = O.
NB:
A pair of shared electrons between 2 atoms forms SINGLE BOND, X Y.
Two pairs of shared electrons between 2 atoms forms DOUBLE BOND, X = Y.
It is essentially another type of covalent bonding since it involves sharing of electrons. The
difference here is that the electrons to be shared are donated solely by one of the atoms in the
bond
This results from the existence of a lone pair of electrons, which are not directly concerned with
valency.
This kind of bonding exists in the ammonium molecule (NH4+). Coordinate bonding is
sometimes referred to as dative bonding.
8.6. Metallic Bonding
Metallic bonding is bonding within atoms of metals caused by attractive force between
positively charged metal ions and negatively charged free electrons. The atoms are packed
closely together in giant lattice structures.
A valency is the number of hydrogen atoms which combine with or displace one atom of an
element or one group of the radical e.g. one atom of oxygen combines with two atoms of
hydrogen to form water.
Therefore the valency of oxygen is 2. Hydrogen is regarded as the standard and its valency is 1.
A valency can also be defined as the number of electrons an element or radical must gain or lose
in order to attain a stable electronic configuration e.g. oxygen has electronic configuration 2:6,
therefore it needs two electrons to be a noble gas with electronic configuration of 2:8
Element Symbol
Hydrogen 1
Helium 0
Lithium 1
Beryllium 2
Boron 3
Carbon 4
Nitrogen 3
Oxygen 2
Fluorine 1
Neon 0
Sodium 1
Magnesium 2
Aluminum 3
Silicon 4
Phosphorous 3 and 5
Sulphur 2
Chlorine 1
Argon 1
Potassium 0
Calcium 1
Chromium 3
Manganese 2
Iron 2 and 3
Cobalt 2
Nickel 2
Copper 2
Zinc 2
Lead 2
Mercury 2
Silver 1
Bromine 1
Iodine 1
9.2. Radicals
Hydroxide OH 1
Bromide Br 1
Nitrate NO3 1
Chlorine Cl 1
Oxide O 2
Carbonate CO3 2
Sulphate SO4 2
Sulphite SO3 2
Phosphate PO4 3
10. Chemical formulae
This is a group of letters and numbers which represent the name of a compound. In writing
chemical formulae, we therefore use symbols of elements and their valences.
First write the symbols of the element or radical that makes up that compound e.g.
Na Cl Ca SO4 Al O
Write the valences on top right side of the symbols of elements or radicals
Rewrite the symbols again reversing the valences from top right side to the bottom right
side of symbols and radicals
Giving
Na1 Cl1 Ca2 (SO4) 2 Al2 O3
Note: if the valency of any element is 1, it should not be written in the final formula. Also if
the two valencies are similar, they should not be written.
The valency of elements combined in a radical should be written outside the brackets and affects
all the elements enclosed when counting the number of atoms e.g. Al2(SO4)3
For valencies in a formula which are multiples should be cancelled to their lowest possible
values
Examples
i. Sodium hydroxide
Sodium Hydroxide
Symbols Na OH
Magnesium Chloride
Symbols Mg Cl
Carbon Oxide
Symbols C O
Valencies on top C4 O2
Reversing valencies C4 O2
Cancelling valencies to the lowest possible values from the final formula: CO2
Ammonium Sulphate
Symbols NH4 SO4
Examples:
Calculate the number of atoms of the elements contained in the following compounds
1 atom of sodium
1 atom of chlorine
1 atom of calcium
1 atom of sulphur
4 atoms of oxygen
8 atoms of hydrogen
I atom of sulphur
4 atoms of oxygen
10 atoms of aluminium
15 atoms of oxygen
3 atoms of copper
3 atoms of sulphur
27 atoms of oxygen
30 atoms of hydrogen
1. The numeral behind the formula represents the number of molecules and is multiplied
through each element in the whole formula to get total number of atoms of each element
2. The numeral in front of each element is multiplied only through that element
3. The numeral outside the brackets should be multiplied only through those elements inside
the brackets
Exercise
Write the chemical formulae of the following compounds and calculate the number of atoms of
each element
Steps
Write the formula for the reactants on the left hand side and that for the products on the right
hand side and check the valencies for the elements forming the formula to confirm if they are
right
i. Sulphuric acid reacts with sodium hydroxide to form sodium sulphate and water
Write the states of matter for each element or compound in the equation:
H2SO4 (aq) + 2NaOH (aq) Na2SO4 (aq) + H2O (l)
A balanced equation is where the number of atoms on the left hand side is equal to the number of
atoms on the right hand side for each kind of element e.g. 2H2 (g) + O2 (g) H2O (l)
Write the formula for the reactants on the left hand side and that for products on the right
hand side. H2 (g) + O2 (g) H2O (l)
Check the valencies for the elements forming the compounds to confirm if formulae are
right
=H2O
Count the number of atoms of each element on the right hand side and on the left hand
side to see if they balance
LHS RHS
H=2 H=2
O=1 O=1
If they dont balance, look for a number which can be multiplied in the formula above to
make all elements balance on both sides of the equation.
LHS RHS
H=4 H=4
O=2 O=2
NB: Balanced
1. A numeral to balance the equation must be written behind the formula of the element and
affects the number of atoms of all elements in the formula e.g. 2H2O
2. Never fix the number in front of the element or formula to be balanced or in the middle of
the formula
3. Great care should be taken to ensure that all formulae are correctly written. Failure of the
equation to balance is an indicator that the equation may be wrong or some formulae are
wrongly written
4. Sometimes it is easier to balance an equation by using fractions and they multiplied by a
number to remove the fraction
Examples
LHS RHS
Na=1 Na=2
O=2 O=1
Not balanced
LHS RHS
Na=4 Na=4
O=2 O=1
Balanced
Or 2Na (g) + O2 (g) Na2O (l) multiplying by two through out to remove
the fraction
4Na (s) + O2 (g) 2Na2O (s)
LHS RHS
Fe=1 Fe=2
O=2 O=3
Not balanced
LHS LHS
Fe=4 Fe=4
O=6 O=6
Balanced
3
Or 2Fe (s) + 2 O2 (g) Fe2O3 (s)) multiplying by two through out to
LHS RHS
Zn =1 Zn =1
H=1 H=2
Cl =1 Cl =2
Not balanced
LHS RHS
Zn =1 Zn =1
H=2 H=2
Cl =2 Cl =2
Balanced
Note:
All the above steps are not required in examinations, but they are important to make you
understand. After you have got the basics then it is advised to balance the equation using
your head and directly write the balanced equation which is always required
Some equations are already balanced and therefore do not need to be balanced by
fractions
More examples:
Exercise
When two or more elements or compounds combine directly to form a more complex compound,
the process is called direct combination.
When a compound is broken down into two or more simpler substances, the type of reaction is
called decomposition e.g. during heating. It has a general formula of AB A + B.
Examples
When an element in a free state replaces another element in a compound, the reaction is called
simple replacement.
When an element in a compound replaces another element in another compound, this type of
reaction is called double displacement.
Examples
12
13. The atmosphere and combustion
Oxygen -21%
Nitrogen- 78%
13.1.1. Oxygen
photosynthesis
It is used as fuel; liquid oxygen is used to burn fuel in some air rockets
Oxyacetylene flame which is produced when acetylene burns in oxygen is used for cutting
metals
13.1.3. Nitrogen
It is used to dilute air in the atmosphere so that burning and rusting do not take place so fast.
Without nitrogen in the air, burning and rusting would be very fast.
This occupies a volume of 0.03% in the atmosphere. Carbon dioxide comes as a result
respiration and burning fossil fuels e.g. petrol, diesel etc...
It is removed from the atmosphere by green plants during the process of photosynthesis.
These include Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn)
They do not react with any substance under ordinary conditions because they are inert and have
full outermost shells.
Argon is used in some electric bulbs to stop the hot filament from darkening the glass of the bulb
In addition to the gases, the atmosphere also contains water vapour. This water vapour comes as
a result of evapouration from the oceans, seas, rivers etc The percentage composition of
vapour varies from 1% to 4% by volume.
Are substances which absorb water from the atmosphere e.g. sodium chloride, calcium chloride,
calcium oxide, concentrated sulphuric acid and anhydrous copper (II) sulphate
Once these substances are left exposed to the atmosphere, they absorb water from the
atmosphere.
These are substances which absorb water from the atmosphere forming solution. Examples
include calcium chloride, sodium hydroxide, phosphorous oxide, iron (III) chloride and sodium
nitrate.
Note:
All deliquescent substances are hygroscopic but not all hygroscopic substances are deliquescent
Deliquescent substances are hygroscopic substances which absorb water from the atmosphere
forming a solution
Examples include:
This is a definite amount of water which some substances chemically combine with, when they
form crystals from their solutions in water.
They are used to dry gases because they have a high affinity for water and therefore can absorb
water from moist gases
Drying agent Gases dried
Calcium oxide Ammonia
Concentrated sulphuric acid All gases except ammonia
Anhydrous calcium chloride All gases except ammonia
Phosphorous (V) oxide All gases except ammonia
Metals burn in air to form basic oxides which when dissolved in water form alkaline solutions
Sodium
This burns with a bright yellow flame forming a yellow solid of sodium peroxide and a little
sodium oxide
Magnesium
This burns with a brilliant flame forming a white smoke and ash i.e. an oxide and a nitride
Aluminium
This burns in air when heated strongly and becomes very hot
Zinc
Zinc burns in air with a green flame to form zinc oxide. It does not form a nitride.
Lead
This melts on heating to shiny beads and then forms lead oxide which is brown when hot and
turns yellow on cooling
Yellow on cooling
Copper
Non metals
Carbon
If oxygen is in short supply, carbon burns with little oxygen to form carbon monoxide
Phosphorous
This burns with a bright yellow flame and produces dense white fumes of phosphorous pentoxide
Sulphur
This burns in air with a bright blue flame forming cloudy fumes of sulphur dioxide
Put wet iron wool in a marked test tube and invert the test tube in a beaker of water. Note and
mark the length of the air column X
Leave the experiment to stand for a week. After a week, water will rise to a certain height in the
test tube. Note the new height of the air column Y.
The length of the air column used in the rusting of iron wool is X-Y.
This is equal to the volume of oxygen used up because oxygen is used up in rusting.
volume of oxygenused up
X 100
The percentage of oxygen used up = volume of total air
XY
= X X100
From the experiment, if the value of X is 10 and that of Y is 7.9, the percentage of oxygen used
up
107.9
= 10 X100
2.1
= 10 X100
=21%
Therefore the volume of oxygen used in rusting is 21% which is the composition of oxygen in
the atmosphere.
Burning is a chemical reaction in which a substance chemically combines with oxygen and
usually heat is produced.
When a candle burns, Carbon dioxide and water are produced.
An experiment to show that Carbon dioxide and water are produced when a candle burns
Pass the gases from the burning candle through a U-tube placed in cold water
Also pass the gases through a test tube containing lime water
After a short time, drops of a colourless liquid collect at the bottom of the U-tube which turn
anhydrous copper (II) sulphate from white to blue
The lime water in the test tube soon turns milky
The colourless liquid is water since it turns anhydrous copper (II) sulphate from white to blue
Lime water in the test tube turns milky due to the presence of Carbon dioxide produced.
Lime water
Therefore when a candle burns, Carbon dioxide and water are produced
13.5. Rusting
Rusting is a chemical change where by iron combines with oxygen under moist conditions to
form a brown substance called iron rust.
Rust is a brown coat formed when iron is left in damp air. Rust is chemically called hydrated iron
(III) oxide (Fe2O3.xH2O)
There are two conditions necessary for rusting to take place. These are:
Oxygen
Water
Place iron nails in the test tube and close the test tube with wet cotton wool. Leave the
experiment to stand for a week.
After a week, the iron nails were found to have turned brown i.e. had rusted.
Close the mouth of the test tube with another cotton wool. This reduces on the amount of damp
air reaching the calcium chloride such that it does not become damp so quickly.
Water is first boiled for a few minutes to drive off any air in it. The iron nails are then placed at
the bottom of the test tube. Oil or grease is added on top of water to form a layer on its surface.
The oil layer prevents any entry of air.
At the end of the week the iron nails did not rust due to absence of air. Therefore iron nails
cannot rust in water if it does not contain air.
Results
Explanation
Prevention of rusting
Iron can be prevented from rusting when either air or water is kept away from it. The methods of
preventing rusting include:
- Surface protection
- Sacrificial protection
- Use of stainless steel
Disadvantage If the layer is broken, air and water can reach metal to rust
Sacrificial Protection
Is to sacrifice more reactive metal to corrode with water and air by layering it over less reactive
metal e.g. iron covered by magnesium.
If the layer is broken, water & air reach underneath layer, overlying metal still protect it.
Applications:
1) Galvanised Iron is steel coated with zinc, usually used on roofs.
2) Protecting ships blocks of zinc are attached to hulls to corrode instead of steel which is the
ship metal.
3) Underground steel pipes these are attached to magnesium block using insulated copper
cables. Magnesium corrodes first than steel.
13.6. Combustion
Combustion is the burning of substances in air. It is divided into
Complete and incomplete combustion
Complete combustion occurs when a substance is completely burnt in excess oxygen
C(s) + O2 (g) CO2(g)
Incomplete combustion occurs when a substance burns partially in a limited amount of oxygen
e.g.
2C(s) + O2 (g) 2CO(g)
14. Methods of gas collection
This method is used to collect gases which are insoluble in water and denser than air e.g. sulphur
dioxide and hydrogen chloride
This method is used to collect gases which are soluble in water and less dense than air or lighter
than air e.g. ammonia
d. Gas syringe
Potassium chlorate is heated in presence of manganese (IV) oxide and oxygen gas is produced. It
is collected in the gas jar over water.
A catalyst: is a substance which alters the rate of a chemical reaction and remains chemically
unchanged at the end of the reaction.
b. Preparation of oxygen by decomposition of hydrogen peroxide in the presence of
manganese (IV) oxide
Hydrogen peroxide is added to manganese (IV) oxide in a round bottomed flask. Effervescence
of a colourless gas occurs and oxygen gas is produced according to the equation
No heat is applied.
By fractional distillation of liquefied air since nitrogen has a lower boiling point of 90K remains
as a liquid
Properties of oxygen
- As rocket fuel
- In steel making, to burn off impurities
- In oxy-acetylene cutting and welding
- In oxygen tanks for deep sea divers and mountain climbers to provide oxygen
- For respiration for most animals
- Used as oxygen tents in hospital to aid patients with respiratory problems
15.1. Oxides
An oxide is a compound formed when an element combines with oxygen e.g. magnesium oxide
(MgO), Calcium oxide (CaO) etc
Types of oxides
a. Basic oxides
These are oxides of metals that react with water to form alkalis. Examples of basic oxides
include calcium oxide (CaO), zinc oxide (ZnO), potassium oxide (K2O), sodium oxide (Na2O)
etc
These basic oxides react with water to form corresponding alkalis e.g.
There are oxides of non-metals which react with water to produce acids.
Examples of acidic oxides include Carbon dioxide (CO2), sulphur dioxide (SO2), sulphur
trioxide (SO3), nitrogen dioxide (NO2) etc
Examples of acids that can be produced from the reactions of these oxides with water are H2CO3,
H2SO3, H2SO4, HNO3 etc
Note:
An acid anhydride is an oxide of a non-metal which reacts with water to form an acid.
c. Amphoteric oxides
These are oxides of metals which show both acidic and basic properties e.g. ZnO, Al2O3 and PbO
(Lead oxide)
d. Neutral oxides
These are oxides which show neither basic nor acidic characters. They are usually the lower
oxides of non-metals e.g. water (H2O), dinitrogen oxide (N2O), carbon monoxide (CO), nitrogen
monoxide (NO) etc
Mixed oxides
These are oxides which react like a mixture of two or more simpler oxides e.g trilead tetraoxide
(Pb3O4), triiron tetraoxide (Fe3O4) and dinitrogen tetraoxide (N2O4)
Peroxides
These are oxides which produce twice as much oxygen as would be expected from the usual
valency of the element in the oxide e.g. sodium peroxide (Na2O2), hydrogen peroxide (H2O2)
etc
16. Oxidation and reduction
Oxidation
Reduction
Oxidizing agents
Or
Examples of oxidizing agents include: oxygen (O2), concentrated sulphuric acid (H2SO4),
potassium permanganate (KMnO4) manganese (IV) oxide (MnO2), nitric acid (HNO3), potassium
dichromate VII (K2CrO4), chlorine (Cl2) etc
Reducing agents
Examples of reducing agents include Hydrogen (H2), carbon (C), carbon monoxide (CO),
hydrogen sulphide (H2S), sulphur dioxide (SO2) and ammonia (NH3).
Redox reactions
A redox reaction is a reaction in which reduction and oxidation occur at the same time.
Therefore in such a reaction, one substance is reduced and another one oxidized.
17. Water and hydrogen
Water
Rain water: this is fairly purely than other kinds of water. It usually contains dissolved
gases like Carbon dioxide, oxygen and dust particles.
Well or spring water: this is good for drinking because it passes through soil and rock and
therefore some filtration has occurred. It also contains some dissolved gases.
River water: This contains a little dissolved matter e.g. salt and plenty of dirt and mud
Sea/lake water: This contains a lot of salt and therefore has a salty taste. Due to the
presence of salt in seas and lakes, such water is not good for drinking.
It is a colourless liquid
It is neutral to litmus
It turns anhydrous copper (II) sulphate from white powder to blue crystals
Impure water is purified by using the process of distillation. Impure water is placed in a
distillation flask and heated to 100oC. The steam comes out while solid impurities remain in the
flask. The steam is cooled and condensed by cold water in the condenser.
Sewage Fertilizers
Acidic gases
Oil
Hot water
Detergents
Insecticides
17.5. Reactions of metals with water
Potassium
Sodium React violently with cold water
Calcium
Magnesium
Aluminum React with steam
Zinc
Iron
Lead
Copper Do not react with cold water or steam
Mercury
Silver
Gold
Potassium
This is a bright silvery metal. It reacts explosively with cold water to produce a colourless
solution of potassium hydroxide and a colourless gas of hydrogen
2K(s) +2 H2O(l) 2KOH(aq) +H2(g)
Sodium
Calcium
This reacts quietly with cold water producing calcium hydroxide and hydrogen
Ca(s) +2 H2O(l) Ca(OH)2 (aq) +H2(g)
Magnesium
This metal reacts with cold water very slowly but when magnesium is hot, it catches fire with
steam and burns with a bright light flame to produce a white ash of magnesium oxide and
hydrogen
Mg(s) +H2O(g) MgO (s) +H2(g)
In cold water, it slowly forms magnesium hydroxide and hydrogen
Mg(s) +2 H2O(l) Mg(OH)2 (aq) +H2(g)
Reaction of magnesium with steam
Clean about 6cm of magnesium ribbon with sand paper. Place wet sand in a small test tube. Heat
is applied to both ends of the wet sand and magnesium ribbon.
The wet sand will produce steam on heating which will react with magnesium ribbon to form a
white ash of magnesium oxide and hydrogen gas is produced
Mg(s) +H2O(g) MgO (s) +H2(g)
Iron
This reacts slowly with steam and produces a black triiron tetraoxide (Fe3O4) and hydrogen gas
3Fe(s) +4H2O(g) Fe3O4 (s) +4H2(g)
3) Displacement Reactions
Displacement reaction is the displacement of ions of metal from compounds of metals lower in
reactivity series by metals higher in reactivity series.
Hydrogen is prepared by the action of dilute acids on metals e.g. action of dilute sulphuric acid
or hydrochloric acid on zinc, magnesium or iron in the presence of a little copper (II) sulphate to
speed up the reaction.
Set up:
Procedure
Dilute hydrochloric acid or dilute sulphuric acid is added through a thistle funnel to zinc granules
in a conical flask
The hydrogen produced is passed through anhydrous calcium chloride in a U-tube or
concentrated sulphuric acid in a wash bottle for drying.
The dry hydrogen gas is then collected by upward delivery because it is lighter than air.
If the reaction is too slow, copper (II) sulphate solution is added to the flask to speed up the
reaction hence acting as a catalyst.
Equation
Metals like sodium, potassium and calcium react with water to produce hydrogen.
Mg(s) +2H2O(l) MgO (aq) +H2(g)
Ca(s) +2 H2O(l) Ca(OH)2 (aq) +H2(g)
2K(s) +2 H2O(l) 2KOH(aq) +H2(g)
2Na(s) +2 H2O(l) 2NaOH(aq) +H2(g)
Test for hydrogen gas
It is tested by using a burning splint. When a burning splint is placed in a gas jar full of hydrogen
gas, it burns with a pop sound.
Properties of hydrogen
Physical properties
It does not support burning but burns readily in air with a faint blue flame forming water
2H2(g) + O2(g) 2H2O(l)
It reacts with non-metals e.g.
Hydrogen reacts with chlorine in presence of sun light to produce hydrogen chloride
H2(g) + Cl2(g) 2HCl(g)
It also combines with nitrogen to produce ammonia gas
3H2(g) + N2(g) 2NH3(g)
Hydrogen is a reducing agent i.e. it reduces oxides of copper, lead and iron to their metals
CuO(s) + H2(g) Cu(s) + H2O(l)
PbO(s) + H2(g) Pb(s) + H2O(l)
Fe2O3 (s) + 3H2(g) 2Fe(s) +3H2O(l)
Action of hydrogen on copper (II) oxide
When hydrogen is passed over heated copper (II) oxide, copper is formed. The black copper (II)
oxide turns brown. A colourless liquid (water) forms on the cooler parts of the test tube.
CuO(s) + H2(g) Cu(s) + H2O(l)
19.1. Acids
An acid is a substance which when dissolved in water produces hydrogen ions as the only
positively charged ions.
of ethanoic acids
- 10 mol/dm3 of ethanoic acid solution is a concentrated solution of weak acid
- 0.1 mol/dm3 of ethanoic acid solution is a dilute solution of weak acid
- 10 mol/dm3 of hydrochloric acid solution is a concentrated solution of strong acid
- 0.1 mol/dm3 of hydrochloric acid solution is a dilute solution of strong acid
Acids are covalent compounds and do not behave as acids in the absence of water as water reacts
with acids to produce hydrogen ions (H+) ions, responsible for its acidic properties.
E.g. Citric acid crystals dont react with metals and dont change colours of indicators; citric acid
in water reacts with metals and turns litmus red.
Hydrogen Ions
Hydrogen gas is formed by acids as hydrogen (H+(aq))ions are present in acid solutions
This means when a solid/gas acid dissolved in water, they produce hydrogen (H+(aq))ions in it
When metals react with acids, only the hydrogen ions react with metals, e.g.:
Chemical equation: 2Na(s) + 2HCl (aq) 2NaCl (aq) + H2 (g)
Ionic equation: 2Na(s) + 2H+ (aq) 2Na+ (aq) + H2 (g)
19.2. Bases and Alkalis
Bases are oxides or hydroxides of metals and react with an acid to from a salt and water only
Alkalis are bases(basic hydroxides) which are soluble in water
Laboratory Alkalis
All alkalis produce hydroxide ions (OH-) when dissolved in water. Hydroxide ions give the
properties of alkalis. They dont behave as acids in absence of water.
Alkalis are therefore substances that produce hydroxide ions, OH-(aq), in water.
19.3.1. pH
pH is the acidity or alkalinity of a substance. The pH scale ranges from 0 to 14. Solution with pH
7 are neutral, those with a pH less than 7 are acidic and those greater than 7 are alkaline. i.e
19.3.2. pH scale
Is used in measuring acidity and alkalinity in aqueous solutions
The PH scale is normally made up of pH values or numbers e.g. pH 7 for neutrality
Acidity ranges from 1 to 6 and alkalinity ranges from 8 to 14
Strength of an acid increases as the value of the numbers (pH) decreases i.e. (6<5<4<
1) represents increasing acidity
Strength of an alkali increases as the value of the numbers increases i.e. (8>9>10>14)
represents increasing alkalinity
19.4. Indicators
Indicators are substances that have different colours in acidic and alkaline solutions
Common indicators:
Litmus
Methyl orange
Phenolphthalein
The table shows the change of colours made by some indicators
Indicator Colour in acids colour changes at pH Colour in alkalis
Phenolphthalein Colourless 9 Pink
Methyl orange Red 4 Yellow
Litmus Red 7 Blue
Screened methyl orange Red 4 Green
Bromothymol blue Yellow 7 Blue
A hand-held pH probe is dipped into solution and a meter will show the pH digitally or by a
scale. Measures pH of water in lakes and streams accurately
3. pH sensor and computer
A probe is dipped into solution and will be sent to computer through interface used to measure
pH of solution. The pH reading is displayed on computer screen.
pH Around Us
- Substances in body involved in good digestion have different pH values
- Blood to heart and lungs contains carbon dioxide making blood slightly acidic
- Acids are used in food preservations (ethanoic acid to preserve vegetables; benzoic acid used in
fruit juices, jams and oyster sauce)
- pH affects plant growth some plants grow in acidic soil; some need alkaline soil
- When hair is cleaned with shampoo which is alkali to dissolve grease, hair can be damaged
unless its rinsed or acid conditioner is used to neutalise excess alkali
Displacement Reactions
E.g. Reactions between magnesium with zinc sulphate
Mg(s) + ZnSO4(aq) MgSO4(aq) + Zn(s)
Its ionic equation is written as:
Mg(s) + Zn2+(aq) + SO42-(aq) Mg2+(aq) + SO42-(aq) + Zn(s)
Since SO42-(aq) ions dont change, omit them, leaving:
Mg(s) + Zn2+(aq) Mg2+(aq) + Zn(s)
Neutralization
Neutralization is the reaction between acid and base to form salt and water only.
From ionic equation, it is known that the reaction only involves H+ ions from acids with OH- ions
from alkali to form water.
E.g. NaOH + H2SO4 forms Na2SO4 + H2O
H2SO4 (aq) + 2NaOH(aq) Na2SO4 (aq) + H2O(l)
Ionic equation is:
H+(aq) + OH-(aq) H2O(l)
Plants dont grow well in acidic soil. Quicklime (calcium hydroxide) is added to neutralise the
acidity of soil according to equation:
Acid (aq) + Ca(OH)2(aq) Ca(acid anion)(aq) + H2O(l)
19.7. Salts
A salt is a substance consisting of positive metallic ions and negative ions derived from an acid
OR:
A salt is a compound formed when the replaceable ionisable hydrogen of an acid is replaced by a
metal or an ammonium ion either wholly or partially
A normal salt: is formed when all the ionisable hydrogen in an acid has been replaced by a
metal or metallic radical e.g sodium chloride (NaCl), Sodium carbonate (Na2CO3)
An acid salt: is formed when only part of the ionisable hydrogen in an acid has been replaced by
a metal or metallic radical e.g sodium hydrogen sulphate (NaHSO4), sodium hydrogen carbonate
(NaHCO3) etc
Soluble Insoluble
All Nitrates -
All Sulphates Barium sulphate (BaSO4), Lead sulphate (PbSO4),
Calcium sulphate (CaSO4) is slightly soluble
All Chlorides Lead (II) chloride (PbCl2), silver chloride (AgCl)
Potassium, Sodium, Ammonium salts -
Potassium carbonate (K2CO3), sodium All Carbonates
carbonate (Na2CO3), ammonium
carbonate (NH4CO3)
Potassium oxide (K2O), sodium oxide All Oxides
(Na2O)
Oxides and hydroxides of potassium, All other oxides and hydroxides (those of calcium
sodium and ammonium and magnesium are slightly soluble)
E.g. Preparation of barium sulphate (BaSO4) from Barium chloride and dilute sulphuric
acid
Barium chloride (BaCl2) which contains the wanted barium ion, is reacted with dilute sulphuric
acid (H2SO4) which contains the wanted sulphate ion, to produce solid barium suphate (BaSO4)
and aqueous potassium chloride (KCl). BaSO4 is then separated from KCl by filtration, leaving a
colourless filtrate (KCl(aq)) & BaSO4 left on filter paper. The salt is then washed with water to
completely remove KCl & filter paper is squeezed with another filter paper to dry BaSO4.
Preparation of lead (II) sulphate
An insoluble salt of lead (II) sulphate is prepared by precipitation. This is done by mixing
solutions containing a sulphate and lead (II) ions
A soluble salt of lead is mixed with a soluble sulphate
When lead ions combine with the sulphate ions, lead (II) sulphate is formed
Pb2+(aq) + SO42-(aq) PbSO4(s)
Lead (II) nitrate solution is put in a beaker and the solution is heated. Dilute sulphuric acid is
then added. A white precipitate of lead (II) sulphate is formed. The precipitate is filtered off as a
residue and washed with distilled water. The crystals are then dried.
Pb(NO3)2( aq) + H2SO4 (aq) 2HNO3(aq) + PbSO4(s)
Ionically
Pb2+(aq)+ 2NO3-(aq) + 2H+(aq) + SO42-(aq) 2NO3-(aq) + 2H+(aq) + PbSO4(s)
Pb2+(aq) + SO42-(aq) PbSO4(s)
Preparation of copper (II) sulphate from sulphuric acid and copper (II) oxide
Dilute sulphuric acid is put in a beaker and copper (II) oxide is added little by little until in
excess. The excess oxide is the filtered off.
The filtrate which remains in the container is concentrated by heating. The solution is left to cool
for some time.
Crystals are formed and then filtered off and put on a filter paper for drying
CuO(s) + H2SO4(aq) CuSO4(aq) + H2O(l)
CuSO4(aq) + 5H2O(l) CuSO4.5H2O(s)
NB: Salts of lead (II) nitrate, zinc sulphate and magnesium sulphate can be prepared in the same
way.
Nitrates
Ammonium nitrate decomposes on heating to give dinitrogen oxide and water
NH4NO3(s) N2O(g) + H2O(l)
Metal nitrates decompose following the reactivity series
Calcium
Magnesium
Decompose on heating to form metal oxides, nitrogen dioxide and oxygen
Potassium and sodium nitrate melt into colourless liquids and then decompose to pale yellow
nitrites and oxygen
2KNO3(s) 2KNO2 (s) + O2 (g)
2NaNO3(s) 2NaNO2 (s) + O2 (g)
Mercury and silver nitrates decompose on heating to form their corresponding metals, nitrogen
dioxide and oxygen
Hg(NO3)2(s) Hg(s) + 2NO2 (g) + O2 (g)
Ag(NO3)2(s) Ag(s) + 2NO2 (g) + O2 (g)
The rest of the nitrates i.e. from calcium to copper nitrates decompose on heating to form metal
oxides, nitrogen dioxide and oxygen
2Ca (NO3)2(s) 2CaO(s) + 4NO2 (g) + O2 (g)
2Cu (NO3)2(s) 2CuO(s) + 4NO2 (g) + O2 (g)
2Pb (NO3)2(s) 2PbO(s) + 4NO2 (g) + O2 (g)
Sulphates and hydrates
Most sulphates are hydrated and when heated, they lose their water of crystallisation to form
anhydrous salts which are resistant to further heating and therefore do not decompose. Therefore
hydrated sulphates do not decompose on heating e.g.
MgSO4.7 H2O(s) MgSO4(s) + 7 H2O(l)
CuSO4.5 H2O(s) CuSO4(s) + 5H2O(l)
Blue White
Na2SO4.10H2O(s) Na2SO4(s) + 10H2O(l)
FeSO4.7 H2O(s) FeSO4(s) + 7 H2O(l)
On further heating, the anhydrous iron (II) sulphate formed decomposes to give sulphur dioxide,
sulphur trioxide and leaves a brown solid of iron (III) oxide
FeSO4(s) Fe2O3(s) + SO2(g)+ SO3(g)
When ammonium sulphate is heated it decomposes to give ammonia, sulphur trioxide and water
(NH4)2SO4(s) NH3 (g) + SO3(g) +H2O(l)
Example
a. Calculate the mass of potassium chlorate that can be crystallized by cooling the solution
from 70oC to 30oC
b. Use the solubility curve to determine the solubility of potassium chlorate at 50oC
Solution:
The solubility of potassium chlorate at 70oC is 160g/100g water and solubility at 30oC is
60g/100g water. Therefore the mass of potassium chlorate that can be crystallized by cooling the
solution from 70oC to 30oC = 160-60 =100g
The solubility of potassium chlorate at 50oC is 100g/100g water. This is obtained by
extrapolation of the line indicated by the dotted line on the graph
4.28
( X 100)
100g of water dissolve = 17.10 g
=25g
Therefore, the solubility of potassium nitrate is 25g/100g of water at 50oC
Exercise
1. 10g of a saturated sodium chloride solution was evaporated and 6g of solid sodium
chloride was left. Calculate
a. Solubility of sodium chloride (Ans = 150g)
b. The percentage of sodium chloride in a saturated solution (Ans = 60%)
2. 75g of a saturated solution contains 30g of a salt. Calculate its solubility(Ans =
66.67g/100g of water)
3. The solubility of X is 40g/100g of water. Calculate the mass of X that can be dissolved in
60g of water to give a saturated solution (Ans = 24g)
4. The table below shows the solubilities of salt P in water at different temperatures
Temperature /oC 10 20 30 40 50 60
Solubilities (g/100g of solvent) 18 20 24 30 38 50
a. Plot a graph of solubility of P
b. Use your graph to determine solubility of P at 25oC (Ans =22g/100g of water)
c. Calculate the mass of P that would dissolve in 45g of water at 25oC (Ans =9.9g)
20. Carbon and its compounds
Carbon
This is the element in group IV of the periodic table. It has atomic number 6 and atomic mass 12
Occurrence of carbon
Pure carbon is found in the form of diamond and impure carbon as graphite. Carbon is a
constituent of numerous naturally occurring substances such as coal, mineral oils, carbonates,
organic matter and in air as carbon dioxide.
Allotropes of carbon
Allotropy: is the existence of an element in more than one form, without change in physical
state. These different forms are called allotropes.
An allotrope is one of two or more distinct forms of an element. Carbon has three allotropes:
diamond, graphite and amorphous carbon. Others elements that show allotropy are:
1. Oxygen with two allotropes, that is, ordinary oxygen (O2) and ozone (O3).
2. Tin with two allotropes, that is, grey tin and white tin.
3. Sulphur has five allotropes, that is, amorphous sulphur, monoclinic sulphur, rhombic
sulphur, plastic sulphur and colloidal sulphur.
4. Phosphorus with two allotropes, that is white/yellow phosphorus and red phosphorus.
Graphite
Graphite is a black, soft, slippery, hexagonal crystalline substance. Its atoms are joined by strong
covalent bonds.
Structure of graphite
Graphite is a two dimensional layered structure. The carbon atoms within the layers are arranged
in hexagonal rings and each carbon atom is covalently bonded to three other carbon atoms.
For each carbon atom, three out of the four valence electrons localized during the formation of
the covalent bonds. The remaining electron is delocalized (mobile) over the whole layer. The
mobile electrons are free to move through the structure and therefore enable graphite to conduct
heat and electricity.
The hexagonal layers lie on top of one another and are joined by weak van der waals forces
which enable layers to slide over each other easily. That is why graphite is soft and can be used
as a lubricant.
Properties of graphite
1. It is a black material which feels greasy on touching
2. It is opaque and shiny
3. It has a density of 2.3g/cm3
4. It has hexagonal structures
5. It conducts electricity. This because it contains delocalised electrons. When an electric
field is applied, these electrons move freely conducting electric current.
6. Writes well on papers
7. Graphite is soft and slippery because the layers within the structure are held together by
weak van de waals forces which makes it possible for the layers to slide over each other
easily.
Uses of graphite
1. Used as protective coating for iron substances to prevent rusting.
2. Used as an electrode in electrolytic cells. Since graphite conducts electricity
3. As a lubricant for dynamos and electric motors.
4. For making pencil leads
Diamond
It is a colourless, transparent and sparkling crystalline substance. It is the hardest substance
known.
Structure of diamond
Diamond has a tetrahedral structure in which each carbon atom is joined by covalent bonds to
four other carbon atoms.
The valence electrons are all used in forming covalent bonds and therefore they are localized.
Diamond is a poor conductor of heat and electricity because it does not have free and mobile
electrons.
Uses of diamond
1. Used for manufacturing drilling and cutting hard substances such as glass and rock.
2. It is used to make jewelry e.g. necklaces and ear rings because of its high refractive index
giving it a shiny appearance.
Amorphous carbon
Amorphous carbon is black and has the lowest density among all the allotropes of carbon. It is a
fair conductor of electricity. It is a non-crystalline substance. Amorphous carbon exists in several
forms including wood charcoal, animal charcoal and lampblack. Coke and soot are other forms
of impure amorphous carbon. Animal charcoal is made by heating animal bones and remains in
a limited supply of air. Coke is made by heating coal in absence of air. Wood charcoal is formed
by burning in a limited supply of air. It can be used to remove poisonous gases such as ammonia,
sulphur dioxide and chlorine. It is also useful as fuel. Lamp black is made by burning oil in a
limited supply of air e.g. kerosene, petroleum and turpentine. It is used in making printers ink,
shoe polish, carbon paper and car tyres.
Exhaust fumes from cars contain carbon monoxide because of incomplete combustion of petrol
or diesel.
Charcoal continues to burn slowly with a yellow flame without any further heating. The amount
of charcoal gradually decreases and finally only a small amount of ash is left. The presence of
ash implies that wood charcoal is not pure carbon.
This reaction is used in extraction of the metals. Those metals higher in reactivity series than
carbon have a higher affinity for oxygen and will not give it up to carbon.
Experiment:
The gas is then passed through a bottle containing water or potassium hydrogen carbonate
solution to absorb any hydrochloric acid fumed.
The gas is collected by downward delivery in a gas-jar since the gas is denser than air.
Ionic equation
Ca2+(aq) + CO32- (aq) + 2H+(aq ) + 2Cl- (aq) Ca2+(aq) + 2Cl- (aq) + H2O(l) + CO2 (g)
NB: If the gas is not required dry it can be collected over water. This is possible because carbon
dioxide is only slightly soluble in water.
Dilute sulphuric acid is not used with calcium carbonate because the reaction produces calcium
sulphate which is sparingly soluble and thus forms a coating on the calcium carbonate which
stops further reaction.
The insoluble salt coats the carbonate preventing it from reacting with the acid.
Kipps apparatus
A continuous supply of carbon dioxide can be obtained from a Kipps apparatus using calcium
carbonate and dilute hydrochloric acid.
It is also obtained from the manufacture of cement. Cement is made by heating limestone
(calcium carbonate) with some sand and silicates to form impure calcium oxide.
Physical properties
1. It is a colourless gas
2. It has a faint sharp test
3. It has a very faint smell
4. It does not support burning
5. It is slightly soluble in water forming carbonic acid.
H2O(l) + CO2(g) H2CO3(aq)
7. It is denser than air. When carbon dioxide in a jar is poured into another jar containing a
lighted candle, the candle is extinguished. This shows that carbon dioxide is denser than
air. It displaces air from the jar containing a lighted candle hence starves the candle of
oxygen.
Chemical properties
Magnesium continues to burn in carbon dioxide because of its higher affinity for oxygen than
carbon. The heat from the burning magnesium decomposes carbon dioxide into carbon and
oxygen. The decomposition of carbon dioxide provides more oxygen which supports continued
burning of magnesium oxide.
The above test is used to distinguish carbon dioxide from any other gas. However, if excess
carbon dioxide is bubbled through the milky solution, the white precipitate dissolves to form a
colourless solution due to the formation of calcium hydrogen carbonate, which is soluble in
water.
With excess carbon dioxide, a white precipitate of sodium hydrogen carbonate is formed. The
precipitate is sparingly soluble in cold water.
When a jar of carbon dioxide is placed in a trough containing sodium hydroxide solution, the
solution quickly rises into the jar. This is because the gas is rapidly absorbed into the solution.
Carbon dioxide reacts with sodium hydroxide solution.
When solid sodium hydroxide is exposed to air, a colourless solution is formed and later a white
crystalline solid is formed. Sodium hydroxide is deliquescent and therefore absorbs water from
air to form a solution. The solution absorbs carbon dioxide from air forming a white crystalline
solid of sodium carbonate decahydrate.
1. Carbon dioxide is used in the manufacture of carbonated drinks because of its pleasant
taste in water.
2. Carbon dioxide is used as a refrigerating agent for perishable goods
3. Pieces of solid carbon dioxide are sometimes dropped into clouds to cool them to form
rain.
4. Carbon dioxide is used in fire extinguishers. Carbon dioxide being denser than air forms
a layer around the burning material. It covers the fire and starves it of oxygen hence the
fire is put out.
5. It is used during photosynthesis by green plants
6. It is used in the manufacture of sodium carbonate and sodium hydrogen carbonate
Exercise
1.
a.
i. Draw a labeled diagram of the set-up of the apparatus that can be used to prepare a dry sample of
carbon dioxide in the laboratory
ii. Write an equation that leads to the formation of carbon dioxide
iii. Write an ionic equation for the reaction leading to the formation of carbon dioxide
b. Carbon dioxide was passed through calcium hydroxide solution. Describe and explain the
reaction that took place.
c.
i. State what would be observed if burning magnesium ribbon was lowered into a jar of carbon
dioxide
ii. Write equation for the reaction that takes place
2.
a. Describe the structure of graphite
b. State two properties in which graphite differs from diamond
c. Graphite was heated in excess air and the gas given off passed through aqueous calcium
hydroxide for a long time
i. State what was observed
ii. Write equations for the reaction (s)
Charcoal is put in a combustion tube and heat is applied until red hot. Carbon dioxide is passed
over heated charcoal in a combustion tube. Carbon monoxide is produced.
The mixture of excess carbon dioxide and carbon monoxide is passed over concentrated
potassium hydroxide solution which absorbs carbon dioxide. Carbon monoxide is then collected
over water because it is insoluble in water.
Sodium methanoate is put in a flask and concentrated sulphuric acid is added drop wise through
a tap funnel.
Effervescence occurs and carbon monoxide is collected over water. In the flask, sodium
methanoate is first converted to methanoic acid which is later dehydrated with concentrated
sulphuric acid
Oxalic acid crystals are placed in a flask and concentrated sulphuric acid added through a thistle
funnel. The mixture is then warmed, effervescence occurs and a mixture of carbon dioxide and
carbon monoxide is produced. The mixture is then passed over concentrated potassium
hydroxide which absorbs carbon dioxide. Carbon monoxide is then collected over water.
1. It is a colourless gas.
2. It has no effect on litmus paper, that is, it is a neutral gas.
3. It burns in air with a blue flame forming carbon dioxide.
2CO(g) + O2(g) 2CO2(g)
This reaction also takes place in a charcoal burner when there is a sufficient supply of air.
At A, there is plentiful supply of oxygen and charcoal burns to form carbon dioxide.
At B, the rising carbon dioxide is reduced by red-hot charcoal to form carbon monoxide.
At the surface of the burner, the hot carbon monoxide burns in the air with a blue flame to form
carbon dioxide.
If the charcoal burner is in a poorly ventilated room with insufficient air, the reaction at the
surface fails to takes place. The poisonous carbon monoxide is released into the room. If
someone stays in such a room, he or she may die within a short while due to carbon monoxide
poisoning.
4. It is insoluble in water.
5. It is a reducing agent. It reduces some metallic oxides of copper, lead, zinc and iron, that
is, oxides of metals below carbon in activity series. The porcelain boat is heated strongly
and the excess carbon monoxide is lighted at the jet.
CuO(s) + CO(g) Cu(s) + CO2(g)
(black) (brown)
(red-brown) (grey)
(white) (grey)
Lead(II) oxide (yellow) is reduced to a grey solid. Carbon monoxide does not, however, reduce
the oxides of metals higher than carbon in the reactivity series. Such metals have a higher affinity
for oxygen than carbon monoxide.
6. It is a poisonous gas because it forms a fairly stable compound with haemoglobin which
reduces the oxygen-carrying capacity of blood.
Exercise
1)
a) Name the element present in pure charcoal
b) Explain why it is dangerous to use charcoal stove in a poorly ventilated room.
c) Write an equation for the reaction between charcoal and heated iron (III) oxide.
2) The figure below shows an experimental setup to investigate the effect of carbon monoxide
on oxides of metals.
a)
i) State the conditions for the reaction taking place in the combustion tube.
ii) Write the equation for the reaction taking place in the combustion tube.
b)
i) Name the gas X being burnt at the jet.
ii) Why is it necessary to burn gas X?
iii) Write equation for the combustion of gas X.
c) Name any other oxide that can be used instead of lead(II) oxide.
d) What would you expect to happen if lead (II) oxide was replaced with magnesium oxide?
Give a reason for your answer.
20.2. Carbonates and hydrogen carbonates
Carbonates
Carbonates are salts derived from carbonic acid (H2CO3). Aluminium carbonate does not exist.
Calcium
Magnesium
Zinc Insoluble in water
Iron Decompose on heating to form metal oxides and carbon dioxide
Lead
PbCO3(s) PbO(s) + CO2 (g)
Copper
Carbonates of potassium and sodium are not decomposed by heat. It is only lithium carbonate in
group I that decomposes on heating.
Carbonates of calcium, magnesium, zinc, iron, lead and copper are decomposed by heat to an
oxide and carbon dioxide.
When a white solid (powder) of lead (II) carbonate is heated strongly in a test-tube, a
colourless gas which turns lime-water milky is given off and a brown residue of lead (II) oxide
when hot and yellow when cold is formed.
When a green solid (powder) of copper (II) carbonate is heated, a black residue of copper(II)
oxide is formed.
Black residue
When a white solid (powder) of zinc carbonate is strongly heated, a yellow residue when hot
and white when cold is formed
Brown residue
White solids of magnesium carbonate and calcium carbonate decompose to white solids of
magnesium oxide and calcium oxide respectively.
Ammonium carbonate sublimes when heated. The cause of this sublimation is that ammonium
carbonate dissociates on heating to ammonia, water and carbon dioxide, which recombine on
cooling.
Add a dilute hydrochloric or sulphuric or nitric acid to the solution or solid to be tested.
Effervescence with liberation of a colourless gas that turns lime-water milky indicates the
presence of a carbonate (CO32-) or a hydrogen carbonate.
2H+ (aq) + CO32- (aq) H2O(l) + CO2 (g)
Hydrogen carbonates
These are salts derived from carbonic acid and are formed by partial replacement of hydrogen in
the acid by a metal. Therefore hydrogen carbonates are acidic salts. Common hydrogen
carbonates include sodium hydrogen carbonate (NaHCO 3) and calcium hydrogen carbonate
(Ca(HCO3)2)
Hydrogen carbonates are decomposed by heat to produce carbonates, carbon dioxide and water
When sodium hydroxide is exposed to air, it absorbs water forming a solution. The solution then
absorbs carbondioxide from the air and forms a crystalline solid of washing soda (sodium
carbonate decahydrate Na2CO3.10 H2O)
On further exposure, the hydrated sodium carbonate decahydrate loses its water of crystallisation
forming a white powder of sodium carbonate monohydrate
Na2CO3.10 H2O(s) Na2CO3.H2O(s) + 9H2O (l)
This white powder later absorbs carbon dioxide to form sodium hydrogen carbonate
Add magnesium sulphate or magnesium chloride solution to the test solution. A white precipitate
indicates the presence of a carbonate.
Hydrogen carbonates gives no precipitate but on heating, the magnesium hydrogen carbonate
decomposes to the insoluble magnesium carbonate (white precipitate).
Exercise
2. The figure below shows an experimental setup to investigate the effect of heat on lead (II)
carbonate.
(a) Write the equation for the reaction taking place in test-tube W.
(b) State what is observed in test-tube Q.
(c) What is observed in test-tube Q if lead (II) carbonate is replaced with sodium carbonate?
Give a reason for your answer.
3. (a) Write the equation for the reaction that would take place if
(i) Dilute hydrochloric acid is added to sodium hydrogen carbonate.
(ii) Sodium hydrogen carbonate is strongly heated.
(b) State what would be observed and write equation for the reaction that would take
place if magnesium sulphate solution is added to a solution containing
Calcium oxide is manufactured mainly from limestone, which is heated to very high
temperatures in a kiln. The limestone is mixed with coke or coal and it is fed into the kiln at the
top. Coke or coal burns and the heat decompose the limestone into the oxide and carbon dioxide.
The lime sinks to the bottom of the kiln and is removed; carbon dioxide is allowed to escape.
It can also be obtained by strongly heating sea shells. Sea shells contain calcium carbonate which
decomposes into the oxide and carbon dioxide.
Effervescence occurs and the gas (carbon dioxide) produced is passed through water to remove
traces of acid.
Then carbon dioxide is passed into a moderately concentrated solution of sodium hydroxide for
some time until finally a white precipitate of sodium hydrogen carbonate appears.
The white precipitate is filtered off and washed two or three times with cold water. The solid is
transferred into a dish and heated to a constant mass. Sodium carbonate is obtained as a fine
white powder.
The raw materials in this process are calcium carbonate (limestone) and sodium chloride in form
of brine. The ammonia dissolves in sodium chloride.
The mixture is reacted with carbondioxide down a large tower called the carbonator in which
there is an upwards flow of carbon dioxide (from decomposition of calcium carbonate) under
pressure. Sodium hydrogen carbonate is produced
Sodium hydrogen carbonate precipitates in the lower part of the tower in form of a wet sludge,
which is tapped off from the bottom of the tower.
After filtration and washing to remove ammonium compounds, sodium hydrogen carbonate is
heated to produce sodium carbonate.
2NaHCO3(s) Na2CO3(s) + CO2(g) + H2O(g)
Carbon dioxide is recycled for use. Ammonia is recovered from the ammonium chloride by
reacting ammonium chloride with calcium hydroxide, obtained by adding water to calcium oxide
(from decomposition of calcium carbonate). Ammonia is recycled for use.
Therefore, the end products of solvary process are calcium chloride and sodium carbonate.
Calcium chloride is used in extraction of sodium.
1. It is used for softening of water for domestic purpose. Calcium ions which are the
principal cause of hardness in water; are precipitated from water as calcium carbonate by
the addition of sodium carbonate.
Ca2+ (aq) + CO32- (aq) Ca CO3 (s)
Washing soda
Washing soda is sodium carbonate decahydrate (Na 2CO3.10H2O).When exposed to air, the
crystals lose mass and become coated with a fine white powder. Each molecule of washing soda
gives up, to the atmosphere, nine molecules of water of crystallization forming sodium carbonate
monohydrate (Na2CO3.H2O).
Washing soda is used for softening water by precipitating the calcium ions from solution as
calcium carbonate.
Exercise
1. (a) State what would be observed if sodium carbonate solution was added to
(i) Aqueous calcium hydroxide.
(ii) Dilute sulphuric acid.
(b) Write ionic equations for the reactions in (a) (i) and (ii).
2. A mixture containing copper (II) sulphate and copper (II) carbonate was shaken with
water and filtered.
(a) Identify the residue.
(b) To the residue was added dilute sulphuric acid.
(i) State what was observed.
(ii) Write the equation for the reaction.
1) Combustion: Carbon and its compounds burn in air to produce carbon dioxide e.g. burning
of coke, coal, wood, petrol, oils etc.
C(s) + O2 (g) CO2 (g)
2) Respiration: When sugars are oxidized in the body, carbondioxide is produced
1) Photosynthesis: Green plants absorb carbon dioxide from the atmosphere to make their own
food
2) Hardening of mortar: Mortar and white ash remain slaked lime which slowly absorbs
carbon dioxide is produced
3) Solution in water: Rain dissolves carbon dioxide to form a weak acid (carbonic acid) which
runs into rivers, lakes, seas and oceans
Exercise
1. (a) Zinc carbonate was strongly heated in a test-tube until no further change.
(i) State what was observed.
(ii) Write the equation for the reaction which took place.
(b) The residue formed in (a) above was added to dilute sulphuric acid and heated.
(c) Give two examples of other elements which show allotropy and name their allotropes.
3. (a) Name two common reagents used in the laboratory preparation of carbon dioxide.
(b) State what is observed when carbon dioxide is bubbled in fairly concentrated sodium
hydroxide solution for some time.
(d) Describe how you would show by a chemical test that graphite is made up carbon
atoms.
5. Carbon monoxide was passed over strongly heated copper (II) oxide.
(i) State what was observed.
(ii) Write the equation for the reaction.
(iii) Name any other oxide that shows similar reaction with carbon monoxide.
6. (a) Draw a well labeled diagram for preparation of sodium carbonate in the laboratory.
(b) (i) What is observed when washing soda (Na2CO3.10H2O) is exposed to atmosphere
for some time.
7. (a) Copper (II) carbonate was heated strongly until there was no further change.
(i) State what was observed.
(ii) Write an equation for the reaction.
(iii) Name one reagent which can be used to identify the gaseous product.
(b) Excess dilute sulphuric acid was added to the residue in (a) and the mixture warmed.
8. (a) (i) How can calcium oxide (quicklime) be obtained on large scale?
Diagram not required.
(b) (i) What would be observed when fresh calcium oxide is added to water in a beaker?
(ii) Write equation for the reaction that would occur.
Hardness of water is due to presence of calcium ions (Ca 2+) or magnesium ions (Mg2+) present in
water
1) Boiling
When temporary hard water is boiled, it becomes soft. This is because boiling decomposes
calcium hydrogen carbonate or magnesium hydrogen carbonate to calcium carbonate or
magnesium carbonate respectively thus removing the calcium (Ca2+) ions or magnesium (Mg2+)
ions from water making it soft
This removes temporary hardness from water as it reacts with hydrogen carbonates ions
dissolved in water
This precipitates calcium carbonate hence removing calcium ions from water making it soft
a. Leads to wastage of soap as it needs much soap before it forms lather. Initially soap is
used in removing calcium sulphate from water before a lather is formed
b. Produces scum which leaves dirty marks on clothes
Scum can also damage silk and nylon clothes.
Scum: is a solid precipitate formed when hard water reacts with soap. Scum is chemically
called calcium stearate (C17H35COOCa)
C17H35COONa(aq) + CaSO4(aq) (C17H35COO)2Ca(s) + Na2SO4(aq)
c. Leaves fur in kettles and pans in which it is boiled. This fur is a poor conductor of heat
and therefore wastes fuel or energy.
Fur: fur is solid calcium carbonate or magnesium carbonate formed inside kettles or pans
d. It forms boiler scales inside boilers; this is also a coat of calcium carbonate or magnesium
carbonate formed inside boilers. Both fur and boiler scales waste fuel as heat cant pass
through them easily.
21.
22. Electrolysis
Electrolysis is the decomposition of an electrolyte in aqueous solution or molten state by passing
an electric current through it.
Definitions
Conductors: These are substances that allow electricity to pass through them
In electrolytes, the conducting particles are called ions while in metals the conducting particles
are electrons.
Some substances do not conduct electricity in solid state e.g. solid sodium chloride or gaseous
state e.g. hydrogen chloride gas but conduct well in aqueous (solution) or molten form. This is
because in solid, the compounds consist of ions held together by strong forces of attraction but
ions separate in molten or solution form and can move freely.
Non-conductors/Insulators: These are substances that do not allow the passage of electricity
through them
Electrolytes are composed of ions. In the solid state, the ions are rigidly held in regular positions
and are not able to move freely. Melting the solid breaks the forces between the ions and
therefore the ions are free to move in a molten electrolyte. Dissolving a solid in water or any
other polar solvent, causes the breakdown of the lattice setting the ions free in aqueous state.
Types of electrolytes
Strong electrolytes
Weak electrolytes
Non electrolytes
Strong electrolyte: This is a compound which is completely ionised in dilute solution and in
molten state e.g. salts such as sodium chloride and mineral acids e.g. hydrochloric acid
Weak electrolyte: This is a compound which is only slightly ionized in dilute solution and in the
molten state. They have very few mobile ions and therefore slightly conduct an electric current
e.g. water, carbonic acid, ethanoic acid and ammonia solution.
A non-electrolyte is a solution or a molten compound which does not conduct electricity and
therefore cannot be decomposed by an electric current e.g. paraffin, sugar solution, ethanol etc.
Non-electrolytes exist only in the form of molecules and are incapable of ionisation. The
molecules have no charge and are therefore not able to carry an electric current.
Electrodes
These are poles of carbon (graphite) or metal where current enters and leaves the electrolyte. The
types of electrodes include;
Cathode: This is a negative electrode at which electrons enter the electrolyte or leave the
external circuit
Anode: This is the positive electrode at which the electrons leave/the electrolyte or enter the
external circuit.
NB: An electrode must be a good conductor of electricity and should not react with the
electrolyte.
Ions
Cation: This is a positively charged ion that will move to the cathode during electrolysis e.g. all
metallic ions e.g. Na+, NH4+, H+, Cu2+, Pb2+ etc.
Anion: This is a negatively charged ion that moves to the anode during electrolysis e.g. all non-
metal ions and radicals e.g. Cl-, SO42-, OH-, NO3-, Br- etc.
Faradays
This is the quantity of electricity carried by one mole of electrons and is equal to 96500
coulombs (C)
Exercise
(ii) X.
(iii) Y.
(ii) X?
This states that electrolytes consist of ions which are positively and negatively charged particles
that move to different electrodes during electrolysis.
In ionic compounds, these charges are held together by electrostatic forces but in solution or
molten state, these ions are free to move. The positive ions move to the cathode and the negative
to the anode.
When an electric current is applied to the electrolyte, the negatively charged ions called anions
move to the positively charged electrode called the anode. Once there, they lose electrons to
become atoms and are said to be discharged i.e. 2Xn-(aq or l) X2 (g) + ne
When two or more ions of similar charges are present under similar conditions in a solution e.g.
K+ and H+ or SO42- and OH-, one is preferentially selected for discharge. The selective discharge
depends on the following factors.
If Cu2+ and H+ ions are both present in solution, both migrate to the cathode but the H + being less
reactive than the Na+ is discharged first.
Cations Anions
K+ SO42-
Na+ NO3-
Decreasing Al3+ I-
Zn2+ OH-
Fe2+
Pb2+
H+
Cu2+
Hg2+
Ag+
E.g. when a solution of copper (II) sulphate is electrolysed using copper electrodes, copper ions
are discharged at the cathode but neither SO 42- nor OH- are discharged at the anode. Instead the
anode dissolves.
Cu2+(aq) + 2e Cu (s)
Cu (s) Cu2+(aq) + 2e
When copper (II) sulphate solution is electrolysed using the copper cathode and carbon anode,
the copper is discharged at the cathode and OH- is discharged at the anode.
Cu2+(aq) + 2e Cu (s)
E.g. Electrolysis of a solution of sodium chloride with mercury as a cathode and with platinum as
cathode. With platinum, the hydrogen ion is discharged in accordance with the order of the
activity series, sodium ion being higher in the series. The cathode product is hydrogen gas. If the
mercury cathode is used, there is a possibility of discharging sodium ion to form sodium
amalgam with mercury. This requires less energy than the discharge of hydrogen ions to form
hydrogen gas and so occurs in preference.
Ions present:
Reaction at cathode:
The hydrogen ions migrate to the cathode, gain electrons and become hydrogen gas.
The hydroxide ions and sulphate ions migrate to the anode. The hydroxide ions being less
reactive than sulphate ions are discharged and oxygen gas is formed.
Overall equation
4H+(aq) + 4OH-(aq) 2H2(g) + O2(g)+ 2H2O(l)
Note:
The acidity at the cathode decreases (pH increases) because the hydrogen ions are discharged as
hydrogen gas and therefore the concentration of hydrogen ions in solution decreases. At the
anode, the acidity increases (pH decreases). The discharge of hydroxide ions disturbs the ionic
equilibrium of water and therefore more water ionizes to restore it.
Therefore the excess hydrogen ions produced, with incoming sulphate ions, is equivalent to
increased concentration of sulphuric acid. This means that the total acidity at anode and cathode
together remains constant. This implies that the final change is that water is decomposed to
produce hydrogen and oxygen. That is why it is called electrolysis of water.
2. Two volumes of hydrogen are produced at the cathode and one volume of oxygen
produced at the anode i.e. Hydrogen : Oxygen =2:1
Overall equation
The bulb does not give out light while the lead (II) bromide is solid showing that no electric
current passes through the solid lead (II) bromide. As the lead (II) bromide melts, the bulb
gives out light. After a while, a brown colouration is observed at the anode and a shiny grey
solid (lead) is deposited at the cathode.
1. Write equations which represent the discharge of the following ions at the
(a) Cathode
(i) Ag+(aq)
(ii) Cu2+(aq)
(iii) Al3+(aq)
(iv) Na+(l)
(b) Anode
(i) Cl-(aq)
(ii) OH-(aq)
(iii) I-(aq)
4. The table below shows the observations made when an electric current was passed
through two substances Q and Z.
Substance Observation
(a) Which of the two substances would not conduct electricity in solid state? Explain.
(b) In what other state would you expect substance Q to conduct electricity?
(c) Name the particles that are responsible for conducting electricity in substance Q and
Z.
(d) Give the type of bonding that is present in substances Q and Z.
5.
a. Which ions would be discharged at the electrodes during electrolysis of a dilute
solution containing
(i) K+ and Mg2+?
b. Write the equation for the discharge of the ions in (a) (i) and (iv).
Reaction at cathode:
Copper (II) ions and hydrogen ions migrate to cathode. Copper (II) ions are discharged
because they are less reactive than hydrogen ions. Copper (II) ions gain electrons from the
cathode and copper is deposited. A brown layer of copper is deposited at the cathode and thus
the mass of the cathode increases.
Reaction at anode:
Both the sulphate and hydroxide ions migrate to the anode but none loses its electrons.
Instead the copper anode itself loses electrons and as it does so, it becomes copper (II) ions
which dissolves in solution. The anode electrode dissolves and its mass decreases.
Electrolysis of copper (II) sulphate solution (using copper cathode and platinum anode)
Ions present: From copper (II) sulphate: Cu2+ and SO42-
Copper (II) ions and hydrogen ions move to the cathode. Copper (II) ions being less reactive
than hydrogen ions are discharged. Copper (II) ions gain electrons and copper is deposited.
The blue colour of the electrolyte (copper (II) sulphate solution) fades as copper is deposited
because copper (II) ions are removed from the solution.
Sulphate ions and hydroxide ions move to the anode. Hydroxide ions being less reactive than
sulphate ions are discharged by giving up their electrons. Bubbles of a colourless gas
(oxygen) are formed at the anode.
The overall equation is obtained by adding the two equations after multiplying the first
equation by 2, to obtain the same number of electrons in both equations.
Copper (II) ions from copper (II) chloride and hydrogen ions from water migrate to the
cathode. Copper (II) ions are discharged because they are more concentrated than hydroxide
ions, thus chlorine gas (greenish yellow gas) is liberated.
However, if the copper (II) chloride solution is very dilute, some discharge of hydroxide ions
will also occur. As the copper (II) chloride solution is diluted, there will not be a point at
which chlorine ceases to be produced and oxygen replaces it. Instead, a mixture of the two
gases will come off, with the proportion of oxygen gradually increasing. The same case
arises in the electrolysis of sodium chloride solution and hydrochloric acid, because the same
anions are involved.
Conc. HCl H+, OH-, Cl- Platinum or carbon Carbon Hydrogen Chlorine Solution becomes alkaline
NaOH Na+, H+, OH- Platinum Platinum Hydrogen Oxygen Bubbles of colourless gas at the
cathode and anode
Conc. NaCl Na+, H+, OH-, Cl- Platinum or carbon carbon Hydrogen Chlorine Bubbles of colourless gas at the
cathode, greenish yellow gas at the
anode
NaOH is formed
Conc. NaCl Na+, H+, OH-, Cl- Mercury Carbon Sodium Chlorine Solution is diluted, grey metal at the
cathode, greenish yellow gas at the
anode
CuSO4 Cu2+, H +, OH-, Copper Carbon/Pt Copper Oxygen Blue colour fades
solution SO42-
Brown solid at the cathode
H2SO4 is produced
Dilute H+, OH-, SO42- Pt Pt Hydrogen Oxygen Total acidity at cathode and anode
H2SO4 reduces, Bubbles of colourless gas at
the cathode and anode
Conc. CuCl2 Cu2+, H+, OH-, Cl- Carbon Carbon Copper Chlorine Solution is diluted
Conc. CuCl2 Cu2+, H+, OH-, Cl- Copper Copper Copper Copper No change in pH
anode
dissolves Brown solid at the cathode and
Cathode mass increases
Exercise
(b) Name one other substance that can be used as electrode in the electrolysis of acidified
water.
2. An aqueous solution of silver nitrate solution was electrolysed using platinum electrodes.
(a) Write an equation for the reaction at
(i) Anode
(ii) Cathode
(b) Write the overall equation for the reactions taking place at anode and cathode.
(c) Write the equations for the reactions at anode and cathode is silver nitrate solution
was electrolysed using silver electrodes.
3. Dilute hydrochloric acid was electrolysed using carbon electrodes.
(a) Name the product(s) formed at the anode.
(b) Write the equation for the reaction at cathode.
4. An aqueous solution of zinc nitrate was electrolysed using platinum electrodes. Write an
equation for the reaction at
(i) Anode
(ii) Cathode.
5. A concentrated solution of copper(II) chloride was electrolysed using copper anode and
copper cathode. Write an equation for the reaction at
(i) Anode
(ii) Cathode
It states that the mass of a substance deposited at the electrodes is directly proportional to the
quantity of electricity passed. This can be illustrated by the graph below.
Mass of
substance
deposited
(g)
i.e. Q = I x t.
Therefore, the quantity of electricity can be found by measuring the current (I) in amperes and
the time (t) in seconds for which it flows. The unit for quantity of electricity is a coulomb (C).
One coulomb is equivalent to one ampere of current flowing for one second.
Exercise
It states that the mass of a substance deposited at the electrodes is inversely proportional to the
charge on its ion. For example, if the same quantity of electricity is separately passed through a
solution of silver ions and copper (II) ions, it is found that the number of moles of silver
deposited are twice the number of moles of copper deposited.
Faraday
This is the quantity of electricity required to deposit one mole of a substance from an ion with a
single charge. One mole of silver ions requires one faraday (1 mole of electrons) to discharge at
the cathode.
Ag+(aq) + e- Ag(s)
One mole of copper(II) ions requires 2 faradays (2 moles of electrons) to discharge at the
cathode.
Two moles of chloride ions require 2 faradays to discharge as one mole of chlorine molecules at
the anode.
From the above examples, one mole of electrons is equivalent to 1 faraday. 1 faraday is
equivalent to 96500 coulombs.
Example I: Quantity of electricity
A current of 4 amps was passed through a solution of magnesium sulphate solution for 5
minutes. Calculate the quantity of electricity
(ii) Used.
Solution:
= 2 x 96500 = 193000 C
(ii) Q = It
Q = 4 x 5 x 60 = 1200 C
When a current of 0.45 amps was passed through a solution of copper (II) sulphate for 1500
seconds, 0.222 g of copper were deposited. Calculate the relative atomic mass of copper.
Solution:
Q = It
675
675
A current of 10 amps is passed through molten magnesium chloride for 4 hours. How many
moles of magnesium metal are produced by this electrolysis?
Solution:
Q = It
But t = 4 x 60 x 60
Q = 10 x 4 x 60 x 60
Q = 144000 C
2 x 96500
Exercise
(ii) Na+
(iii) Fe3+
(iv) Pb2+
(v) OH-
(vi) H+
(vii) Br-
2. Molten aluminium chloride was electrolysed for 200 seconds using a current of 40A.
(a) Write the equation for the reaction at
(i) Cathode.
(ii) Anode.
3. Calculate the mass of aluminium deposited when a current of 960 A passes through a
solution of aluminium (III) oxide in fused cryolite for 800 seconds.
4. What volume of oxygen measured at s.t.p would be liberated in electrolysis of dilute
sulphuric acid by 96000 coulombs. (1 mole of gas occupies 22.4 litres at s.t.p.)
22.2. Application of electrolysis
(i) Electroplating
This is the process of coating a metal with another metal by the process of electrolysis.
Electroplating is done to protect metals from corrosion and to improve their appearance. The
metal to be plated is made the cathode in a suitable electrolyte containing ions of the plating
material. For example, during silver plating, the metal to be plated is made the cathode in a silver
salt solution as an electrolyte and pure silver is made the anode.
The silver salt solution contains positively charged silver ions which are attracted to the cathode
(metal to be plated). Once there, they gain electrons to form silver atoms.
Ag+(aq) + e+ Ag(s)
The anode (plate of pure silver) loses its electrons and forms silver ions which dissolve in the
solution to replace the ones moving to the cathode. The process continues until an adequate layer
of silver has been deposited on the metal being plated.
Other metals which can be used to coat other metals include chromium, nickel, copper and gold.
When iron is to be chromium plated, it is first electroplated with nickel to prevent corrosion and
then with chromium.
For successful electroplating, the material to be plated should be clean and the electric current,
temperature and concentration of the electrolyte should be exactly right. When a very low current
is used, electrolysis proceeds very slowly and a very smooth deposit can be obtained.
Exercise
1)
(a) Name the substance that can be used as the anode and cathode during
(i) Chromium plating of iron metal.
(b) Name the positive ions that must be present in the electrolyte used in a(i) and (ii).
(c) Write the equation for the reaction taking place at cathode in a(ii)
[valency of nickel (Ni) = 2]
(ii) Anodizing
Anodizing is the electrolytic process of coating objects made of aluminium with a very thin
oxide film to protect the metal from corrosion.
To ensure a very thin film of oxide, the oxidation is carried out by electrolysis, using the
aluminium object as the anode. The electrolyte is usually dilute sulphuric acid which gives
oxygen at the anode on electrolysis (refer to electrolysis of dilute sulphuric acid).
Under the correct conditions, oxygen reacts with the surface of the aluminium and coats it with a
thin invisible but protective coating of aluminium oxide.
Na+(aq) + e- Na(s)
The sodium formed dissolves in the mercury cathode to form a solution called sodium
amalgam. The amalgam is mixed with water producing sodium hydroxide solution, hydrogen
and pure mercury.
2NaHg(l) + 2H2O(l) 2NaOH(aq) + H2(g) + 2Hg(l)
Sodium amalgam
Sodium hydroxide is used in the manufacture of soap, in the textile industry (bleaching and
dyeing processes, in rayon manufacture and mercerising cotton to give it a silky sheen), in the
purification of bauxite for aluminium extraction and in the preparation of phenols and cresols
from coal-tar.
Exercise
1. Copper (II) sulphate solution was electrolysed using the set up shown in the figure below.
(ii) Write down all the ions present in aqueous copper (II) sulphate.
(i) Write an equation for the reaction that took place at A and B.
(ii) Comment on the colour of the solution after electrolysis.
2. The figure below shows electrolysis of dilute sulphuric acid using carbon electrodes. A
current of 11.0 amps was passed for 5 minutes and 20 seconds through the circuit.
(a) Name gas
(i) X.
(ii) Y.
(ii) Cathode.
(c) Calculate the volume of gas Y produced at room temperature.
(IF = 96500 C, 1 mole of gas at room temperature occupies 24 dm3)
(ii) Write an equation for the reaction that took place at each electrode.
(b) Calculate the mass of the product formed at the cathode when a current of 2 amps is
passed for 1 hour and 30 minutes. (Pb = 207, Br = 80)
5. A current of 0.25 amps was passed through copper(II) sulphate solution for 40 minutes.
Calculate the
(a) Quantity of electricity used.
(b) Quantity of electricity which deposits one mole of copper.
(c) Mass of copper deposited during the electrolysis.
(d) Moles of copper deposited. (Cu = 64, 1 F = 96500 C)
6. The circuit shown in the figure below was used in an experiment to study the effect of
electricity on lead (II) bromide.
(a) State what was observed.
(i) Before lead(II) bromide had melted.
(ii) X.
(ii) Cathode.
(ii) Cathode.
(ii) Y.
1
greater than hydrogen atom so 12 of carbon-12 atoms is equivalent to the mass of one
hydrogen atom.
Relative Atomic Mass - the average mass of one atom of the element (averaging isotopes) when
1
compared with 12 mass of a carbon-12 atom.
Note: The Relative Atomic Masses are already stated on the periodic table above each chemical
formula.
Relative Molecular Mass the average mass of one molecule of substance (averaging isotopes)
1
when compared with 12 mass of a carbon-12 atom.
Relative Formula Mass same as relative molecular mass but for ions or ionic compounds only
Relative Formula Mass total relative atomic masses of all atoms in a formula of ionic
compound
E.g. Relative formula mass of MgSO4
Mr = 24x1 + 32x1 + 4x16 = 120
Note: Relative molecular mass and relative formula mass have no units
Exercise
Calculate the relative formula masses or relative molecular masses of the following
compounds (H=1, Cl=35.5, Cu=64, S=32, O=16, Na=23, C=12, Fe=56, Zn=65, N=14,
Pb=207, Ag=108 )
a. Hydrogen chloride
b. Copper (II) sulphate
c. Sodium hydroxide
d. Sodium carbonate
e. Iron (II) sulphate
f. Copper (II) chloride
g. Zinc nitrate
h. Lead (II) carbonate
i. Silver chloride
j. Copper (II) sulphate pentahydrate
k. Sodium carbonate decahydrate
E.g.
23
% of Na = 58.5 x 100 = 39.32%
2( 56)
=( 160 x 100) %
112
=( 160 x 100) %
= 70%
Mr (Fe3O4) = 3(56) + 4(16) = 232
3 (56)
Percentage of Fe in Fe2O3 = ( 232 x 100) %
= 72%
Fe3O4 has more iron composition than that of Fe2O3.
2( 56)
=( 160 x 200g
= 140g
= 4.5g
23.5. Mole
Counting Particles
Unit for particles = mole
Symbol = mol
1 mol = 6 x 1023 atoms
Moles of Particles
Calculating the Number of Moles
Number of particles
n= 6 x 10
23
= 5 mol
e.g 2: Calculate the number of molecules in 0.25 mole of CO2. Hence, how many atoms are
present?
Number of particles
0.25mol = 6 x 10
23
Mass(m)
n = Relative formulaatomicmass ( Mr)
m = n x Mr
m = 0.4 x 56 = 22.4 g
Or
1 mole of iron weighs 56g
0.4mol of iron weigh 0.4 x 56 = 22.4 g
23.7. Different Kinds of Chemical Formulae
Ethene formula is C2H6
Molecular Formula shows the actual formula and kinds of atoms present, e.g. C2H6. It
expresses the composition of a compound showing the actual number of atoms in the compound.
Empirical Formula shows the simplest whole number ratio of the atoms present in a
compound, e.g. C2H6, ratio 1:3, therefore C1H3, simply CH3
Structural Formula shows how atoms are joined in the molecule. It can be represented by ball-
and-stick model or diagrammatically.
Ball-and-Stick Diagrammatic
a. Calculating the Empirical Formula of a Compound
Find the empirical formula of an oxide of magnesium consisting of 0.32g of oxygen and 0.48g of
magnesium. (Mg = 24, O = 32)
Solution
Elements present Mg O
Composition by mass 0.48 0.32
0.48 0.32
Number of moles 24 16
1 1
Mg1O1
Therefore, the empirical formula is MgO
23.8. Calculating the Empirical Formula from Percentage Composition
An oxide of sulphur consists of 40% sulphur and 60% oxygen (S = 32, O = 16)
Solution
Elements present S O
Percentage composition 40 60
40 60
Number of moles 32 16
1) 3
S1O3
Therefore, the empirical formula is SO3
Solution:
12 x 17.8
17.8g of carbon dioxide contains 44 g of carbon.
= 4.86g of carbon
2
1g of water contains 18 g of hydrogen.
2 x 9.27
9.27g of water contain 18 g of hydrogen
= 1.03g of hydrogen
Elements present: C H O
12 1 16
Simplest ratio: 4 : 10 : 1
Thus the simplest formula of Q is C4H10O.
Example
A compound Y contains 15.8% aluminium, 56.2% oxygen and 28% sulphur. (S=32, Al = 27 O
=16). (i) Calculate the empirical formula of Y.
Solution:
Elements present: A1 S O
27 32 16
1 1.5 6
Simplest ratio: 2 : 3 : 12
342 342
n = 1
Exercise
1) A compound contains 43.4% by mass of sodium, 11.3% carbon and 45.3% oxygen.
Calculate the simplest formula of the compound (Na = 23, C = 12, O = 16)
2) A compound contains 40% carbon, 6.67% hydrogen, the rest being oxygen. The relative
molecular mass of the compound is 180 (C=12, H =1, O = 16). Determine the empirical
formula of the compound and the molecular formula of the compound
3) An oxide of an element X was made of 50% X. Calculate the simplest formula of the
oxide (X = 32, O = 16)
4) A compound of molar mass 400 with 28% iron, 48% oxygen and the rest being sulphur
was dissolved in water. Calculate the empirical formula and molecular formula of the
compound .(Ans =Fe2(SO4)3)
5) When hydrated sodium carbonate crystals (Na2CO3.xH2O) were exposed to air for a long
time, there was loss of mass of 62.9%. What is the amount of the water of
crystallisation? (Na=23, C=12, O=16, H=1)
Na2CO3.xH2O(s) Na2CO3(s) + xH2O(l) .(Ans x=10)
6) A white crystalline salt (Z.xH2O) contains 51.2% of water of crystallisation. If the
formula mass of the crystals is 120, calculate the amount of water of crystallisation.(Ans
x=7)
One mole of all gases at room temperature and pressure (r.t.p.) = 24dm3
1dm3 = 1000cm3
Formulae:
Volume of a gas
Number of moles of a gas (n) = Molar volume
Example
What is the number of moles of 240cm3 of Cl2 at r.t.p.?
Solution
Since 24000cm3 of chlorine gas contain 1 mole
1
Then 1cm3 of chlorine gas contains 24000 moles
1
x 240
240cm3 of chlorine gas contain 24000 moles = 0.01 mol
Volume of a gas 1
x 240
Or Number of moles of a gas (n) = Molar volume = 24000 moles =
0.01 mol
Then multiply the ratio by no. of moles of Y to find the reacting mole of X.
1
Number of moles of X = 2 x 0.25 = 0.125 mole
To find the reacting mass of X, e.g. Y is given as 35g, we simply multiply the mole by the mass
of Y as they are always in ratio:
0.125 x 35 = 4.375 g
Example
Lead (II) nitrate reacts with potassium iodide according to the equation
Pb(NO3)2(aq) + 2KI (aq) PbI2(s) + 2KNO3 (aq)
Calculate the mass of lead (II) iodide that will be formed when 33.2g of potassium iodide reacts
with excess lead (II) nitrate (K=39, N=14, O=16, Pb=207, I=127)
Solution
Find the ratio first:
Number of moles of Lead ( II ) iodide 1
=
Number of moles of potassium iodide 2
1
Number of moles of PbI2 = 2 x 0.2 = 0.1 mole
Therefore 0.1 mole of PbI2 reacted with 0.2 mole of KI.
2. Calculate the loss in mass when 10g of calcium carbonate is heated to a constant mass
(Ca=40, C=12, O=16) (Ans =4.4g)
3. Calculate the mass of ammonium chloride that will just react completely with 14.8g of
calcium hydroxide (N=14, H=1, Cl=35.5) (Ans = 21.4g)
Then multiply the ratio by no. of moles of HCl to find the reacting mole of H2.
RFM of HCl = 1 + 35.5 = 35.5
14.6
Moles of HCl = 36.5 = 0.4moles
1 1
Number of moles of H2 = 2 x moles of HCl= 2 x 0.4moles = 0.2 moles
To find the volume of H2, simply multiply the mole by the molar volume at room temperature:
Volume of H2 = moles of H2 x Molar gas volume at room temperature
= 0.2 x 24 dm3= 4.8 dm3
But 1dm3 = 1000cm3
4.8dm3 x 1000 = 4800 cm3
4 800cm3 of gas is formed
Exercise
1 16
1g of sulphur absorb 2 X 32 kJ
1 16
x 16
16g of sulphur absorbs 2 X 32 kJ = 29kJ
Exercise
1. The formation of methanol from hydrogen and carbon dioxide is represented by the
following equation
2H2(g) + 2CO(g) CH3OH (l) H = -92kJmol-1
What would be the energy released when 3.2g of methanol is formed? (C=12, H=1,
O=16)
2. Methane burns in oxygen according to the equation
CH4(g) + 2O2(g) CO2(g) + 2H2O (l) H = -890kJmol-1
Calculate the volume of methane at s.t.p that will turn in excess oxygen to produce 2670kJ (1
mole of a gas occupies 22.4dm3 at s.t.p)
3. Calculate the heat produced when 48g of graphite is burnt in excess oxygen (C=12,
O=16, H=1, H = -390kJmol-1)
dm
Volume of solution( 3)
Concentration (C) = Number of moles ( mol )grams of a solute( g)
Molarity
Molarity: Is the number of moles of a substance contained in one litre of a solution
But 1 litre = 1000cm3
Therefore, molarity can also be defined as the number of moles of a substance contained in
1000cm3 of a solution
Units of molarity = moles per litre (moll-1) or moles per dm3 (moldm-3)
Molar solution
This is the solution containing one mole of a substance per litre of solution i.e. solutions which
are 1M.
Standard solution
This is a solution whose concentration is exactly known or is the solution which contains a
known mass and a known volume.
Calculating the Amount of Solute
Examples
1) Calculate the number of moles that are 20g of sodium hydroxide (Na=23, O=16, H=1)
Solution
RFM of NaOH =23+16+1 = 40
Mass(m) 20
Number of moles of NaOH, n = Relative formulaatomicmass (Mr ) = n = 40 =
0.5mol
2) 4g of sodium hydroxide was dissolved in water to make 200cm3 of a solution. Calculate
the molarity of the standard solution formed
Solution
RFM of NaOH =23+16+1 = 40
Mass(m) 4
Number of moles of NaOH, n = Relative formulaatomicmass (Mr ) = n = 40 =
0.1mol
200cm3 of solution contain 0.1mol of sodium hydroxide
0.1
1cm3 contains 200 mol of sodium hydroxide
0.1
3 x 1000
1000cm contains 200 mol of sodium hydroxide
=0.9mol
From
Mass(m)
Moles (n) = Relative formulaatomicmass (Mr )
m = 0.9 x 40
= 36g
4) How many grams of sodium sulphate crystals Na2SO4.10H2O would be required to make
500cm3 of 0.01M solution (Na=23, S=32, O=16, H=1)
Solution
1000cm3 contains 0.01 mol of Na2SO4.10H2O
0.01
3
1cm contains 1000 mol of Na2SO4.10H2O
0.01
500cm3 contains ( 1000 x 500) mol of Na2SO4.10H2O
=0.005mol
From
Mass(m)
Moles (n) = Relative formulaatomicmass (Mr )
Mass (m) = moles (n) x Relative formula mass (Mr)
m = 0.005 x 323
= 1.615g
Exercise
1) How many moles of sulphuric acid are contained in 250cm3 of 0.1M sulphuric acid? (Ans
= 0.0025M)
2) Calculate the mass of nitric acid (HNO3) required to make 200cm3 of 2M solution
(Ans=25.2g)
Volumetric Analysis
Is a measure of the concentrations of an acids/alkalis in solution
Examples
30.0 cm3 of 0.100 M NaOH reacted completely with 25.0 cm3 of H2SO4 in a titration. Calculate
the concentration of H2SO4 in mo mol/dm3 given that:
2NaOH(aq) + H2SO4(aq) Na2SO4(aq) + 2H2O(l)
Solution
Step 1: Find the reacting mole of NaOH
1000 cm3 of solution contain 0.1moles of sodium hydroxide
0.1
1cm3 contains 1000 moles
0.1
X 30
30cm3 contain 1000 moles = 0.0030 moles
Step 3: Find the ratio of number of moles of H2SO4 to number of moles of NaOH
2NaOH(aq) : H2SO4(aq) =2:1
1
X 0.030
0.030 moles of the NaOH reacts with 2 moles of H2SO4
=0.0015moles
0.0015
3 X 1000
1000cm contain 25 moles of H2SO4
= 0.06M
Molarity of H2SO4 = 0.06M
23.14. Uses of Titrations in Analysis
0.1
3
1cm contains 1000 moles
0.1
X 25.5
25.5cm3 contain 1000 moles = 0.00255 moles
Step 3: Find the ratio of number of moles of H2XO4 to number of moles of NaOH
2NaOH(aq) : H2XO4(aq) =2:1
1
X 0.0255
0.0255 moles of the NaOH reacts with 2 moles of H2XO4
=0.01275moles
Step 5: Find the concentration of H2XO4 in mol/dm3
0.01275
X 1000
1000cm3 contain 25 moles of H2XO4
= 0.51M
Molarity of H2XO4 = 0.51M
Example:
7.2g of an impure sample of hydrated sodium carbonate (Na 2CO3.10H2O) was dissolved in 250
cm3 of solution. 20 cm3 of this solution was required to completely react with 25 cm 3 of 0.1M
hydrochloric acid. Calculate:
a) The molarity of pure hydrated sodium carbonate
b) The mass of pure hydrated sodium carbonate per litre of solution
c) The percentage;
i. Purity
ii. Impurity of hydrated sodium carbonate
Solution
0.1
X 25
25 cm3 contain 1000 moles = 0.0025 moles
1
X 0.025
0.025 moles of the acid reacts with 2 moles of the Carbonate
=0.00125moles
0.00125
X 1000
1000cm3 contain 1cm3 contain 20 moles of sodium hydroxide
= 0.0625M
Molarity of sodium hydroxide = 0.0625M
ii. Molarity
concentrationgrams per litre
Molarity=
Molar mass (RMM /RFM )
17.785
250 cm3contains 1000 X 250 g of pure hydrated sodium carbonate
=62.07%
7.24.4688
X 100
ii. Percentage impurity = 7.2 = 37.93%
Solution
b.
i. Moles of the acid
1000 cm3 of acid contain 0.1moles
0.1
1cm3 contains 1000 moles
0.1
X 20
20 cm3 contain 1000 moles = 0.00250 moles
0.2
X 10
10 cm3 of solution contains 1000 moles of sodium hydroxide
0.1
1cm3 contains 1000 moles of sodium hydroxide
0.1
20 cm3 contain
X 20 moles of sodium hydroxide
1000
6.4
0.05=
Molar mass(RMM /RFM )
6.4
Molar mass ( RMM /RFM )= = 128
0.05
Note: Relative molecular mass and relative formula mass have no units
c. Value of x
(COOH)2. xH2O = 128
12 + 32 + X (2+16) = 128
90 + 18X = 128
12890
X = 18
X = 2.1
Number of moles of water of crystallisation =2. Hence (COOH) 2. 2H2O
Exercise
1. 25cm3 of sodium hydroxide reacted completely with 20cm3 of 0.1M hydrochloric acid.
Calculate the concentration of sodium hydroxide in
a. moles per litre (Ans = 0.08mol/litre)
b. Gram per litre (Ans = 3.2g/litre)
2. 25cm3 of a solution of sulphuric acid required 32cm3 of 0.1M sodium hydroxide for
neutralization. Calculate the molarity of the acid(Ans = 0.1M)
3. 20cm3 of sodium carbonate reacted completely with 25cm3 of 0.8M hydrochloric acid.
Calculate the concentration of sodium carbonate in
a. moles per litre (Ans = 0.5mol/litre)
b. Gram per litre (Ans = 53g/litre)
4. 20cm3 of an acid R.xH2O were dissolved in 1 litre of aqueous solution. 25cm3 of this
solution required 16cm3 of 0.5M sodium hydroxide solution. Calculate the relative
formula mass of the acid and determine x (Acid : alkali=1:2, R=89, H=1, O=16)Ans: x=2
5. 25cm3 of an acid HX was neutralized by 24 cm3 of a solution containing 5g of sodium
hydroxide per litre.
a. Calculate the molarity of the acid
b. If the acid solution contained 24g/l; calculate
i. The RFM of the acid
ii. The RAM of X (Ans X= 19)
6. 0.008g of a metallic oxide MO was dissolved in 80cm3 of 0.05M sulphuric acid. The
resultant solution which contained excess acid required 10 cm3 of a solution containing
16g of sodium hydroxide per litre for complete neutralisation (H = 1, O =16, S =32, Na =
23).
a. Write an equation for the reaction between
i. MO and sulphuric acid
ii. Sodium hydroxide and sulphuric acid
b. Calculate the number of moles of
i. Sodium hydroxide used
ii. Excess sulphuric acid
iii. Sulphuric acid which reacted with MO
c. Calculate the molar mass of MO and the atomic mass of M (Ans M = 24)
7.
a. If the acid solution contained 24g/l; calculate
i. The RFM of the acid
ii. The RAM of X (Ans X= 19)
=0.0025Moles
But 1 mole of acid reacts with 1 mole of base
Moles of NaOH = moles of HCl
0.2moles of NaOH are dissolved in 1000cm3
1000
1 moles of NaOH are dissolved in 0.2 cm3
1000 x 0.0025
0.0025Moles of NaOH are dissolved in 0.2 cm3
=12.5 cm3
Exercise
1. 25cm3of a 0.02M sodium hydroxide solution reacted with Vcm3 of an aqueous solution
containing 0.0025moles/cm3 of Z. calculate the volume V (2 moles of NaOH react with 1
mole of Z)
2. What volume of 0.1M hydrochloric acid would react with 25.0cm3 of sodium carbonate
(Na2CO3) solution containing 5.20g of anhydrous salt in 1dm3 of solution (Na=23, C=12,
O=16)
3. What volume of 0.05M sulphuric acid is required to neutralise completely 2.80g of
potassium hydroxide? (K=39, O=16, H=1)
4. 7.5g of compound U occupy, 5.6dm3 at s.t.p. Determine the molar mass of U.
1.9. Volumes of gases
According to Avogadros law, equal volumes of all gases at the same temperature and pressure
contain equal molecules
E.g. 2H2(g) + O2(g) 2H2O(g)
2moles 1mole 2mole
2Volumes 1Volume 2Volumes
Example
20cm3 of carbon monoxide is mixed with 30cm3 of oxygen and exploded. What is the
composition by volume of the resulting gas after cooling the mixture to the original temperature?
Solution
2CO(g) + O2(g) 2CO2 (g)
2moles 1mole 2mole
2Volumes 1Volume 2Volumes
20cm3 10cm3 20cm3
Therefore, 20cm3 of carbondioxide react with 10cm3 of oxygen and (30-10)cm3 of oxygen
remained un reacted. The carbon dioxide produced is equal to 20cm3. Thus the gaseous product
contained 20cm3 of carbon dioxide and 20cm3 of excess oxygen.
Charles law
This states that, the volume of a fixed mass of a gas is directly proportional to its absolute
temperature at a constant temperature
V
V T and T = a constant
P1 V 1 P2 V 2
T1 = T2
Exercise
1. The volume of a fixed mass of a gas is 200cm3 at 0oC and 760mmHg pressure. Calculate
the volume of the gas at 100 oC and 380mmHg pressure.
2. Calculate the volume of hydrogen measured at 25 oC when 88g of potassium react with
water at a pressure of 760mmHg (K=39, volume of the gas at s.t.p is 22400cm3)
24. Qualitative analysis
25. Sulphur and its compounds
25.1. Sulphur
Non metal
Yellow solid
Has 3 allotropes
In the outermost tube, superheated water at 170oC and at a pressure of 10 atmospheres, to keep it
in a liquid form, is sent down to the beds or deposits of sulphur. The sulphur melts and flows into
the reservoir at the base of the pump.
Hot compressed air under a pressure of about 15 atmospheres is sent down through the innermost
tube. It pushes the molten sulphur and water up through the middle tube and its collected in
containers. Water is evaporated off and almost 99% pure sulphur is obtained.
Extraction of sulphur
25.3. Extraction of
sulphur from natural gas
Natural gases obtained during the
refining of petroleum contain
hydrogen sulphide which is absorbed by special solvents. The gas is removed from the solvent
and a small portion of the gas is burnt in air to form sulphur dioxide.
The remaining portion of the gas is left to react with the sulphur dioxide to form sulphur and
water. The water is evaporated off.
Dissolve some powdered sulphur in carbon disulphide in a boiling tube. Place it in a beaker after
extinguishing all flames in the area around. Filter off the solution into another dry beaker and
place a clean filter paper on top of the beaker. Pierce some small holes in the filter paper and
place the set up near a window for a day to allow the carbon disulphide to evaporate. Large
rhombic crystals of sulphur will form.
25.5.2. Monoclinic sulphur
Monoclinic sulphur is also referred to as prismatic sulphur or -sulphur.
Place some powdered sulphur in an evaporating dish. Carefully heat it until it melts. Stir and
gradually add more sulphur until the crucible is full of molten sulphur. Stop heating and allow it
to cool. A crust will form on the surface of sulphur. Carefully pierce through the crust and
immediately pour off the liquid sulphur inside. Cut away the crust by cutting around the edge of
the crucible with a knife. Small needle shaped crystals will be seen inside the evaporating dish.
Transition temperature
This is the temperature at which a change from one form of sulphur to another form takes place.
It is 96oC. Rhombic sulphur is stable below 96oC. Above this temperature, it slowly changes to
the monoclinic form. Monoclinic stable is stable above 96oC and therefore below this
temperature it slowly changes to the rhombic form.
Rhombic sulphur consists of relatively large yellow, translucent, octahedral crystals with a
melting point of 114oC while monoclinic sulphur consists of needle shaped, pale yellow
transparent crystals with a melting point of 119oC
Rhombic sulphur has a density of 2.06g/cm3 while monoclinic sulphur has a density of 1.98
g/cm3.
Crystals of rhombic sulphur are stable below 96oC while monoclinic sulphur is stable above 96
o
C.
On further heating, the liquid becomes very dark, reddish brown in colour and less viscous again.
The chains break and become shorter which can flow more readily. Sulphur boils at 444oC and
forms a brown vapour. On cold surfaces, the vapour condenses directly into a yellow sublimate.
The mist is due to traces of sulphur trioxide formed simultaneously with sulphur dioxide.
S(s) + O2 (g) SO2 (g)
A hot copper foil glows in sulphur vapour forming a black solid, copper (I) sulphide.
Carbon combines directly with sulphur to form a liquid, carbon disulphide. Very high
temperatures are required for this reaction to occur.
Sulphur is oxidized by hot concentrated nitric acid to sulphuric acid. Bromine is added to speed
up the rate of reaction.
Hot concentrated sulphuric acid reacts with copper turnings giving off sulphur dioxide. The gas
is dried by passing it through a bottle containing concentrated sulphuric acid and collected in a
gas-jar by downward delivery since it is denser than air
Instead of using copper turnings, it is also possible to prepare the gas by using sodium sulphite
and dilute sulphuric acid
Na2SO3 (aq) + H2SO4 (aq) Na2SO4 (aq) + H2O (l) + SO2 (g)
Sulphur dioxide is very soluble in water and therefore heating reduces its solubility in water
formed in the flask. Since sulphur dioxide is very soluble in water, it cannot be collected over
water.
4. It is very soluble in water and its solubility in water can be shown by the fountain
experiment
5. It is a powerful reducing agent
a. It reduces iron (III) ions in the brown iron (III) sulphate solution to iron (II) ions
in the green iron (II) sulphate solution. Sulphur dioxide is oxidised to sulphuric
acid
2Fe3+(aq) + SO2(g) + H2O(l) 2Fe2+(aq) + 4H+(aq) + SO42-(aq)
b. Sulphur dioxide reduces concentrated nitric acid to nitrogen dioxide, the sulphur
dioxide being oxidized to sulphuric acid. Brown fumes are observed
3SO2 (g) + Cr2O72- (aq) + 2H+ (aq) 3SO42- (aq) + 2Cr3+ (aq) + H2O (l)
6. Sulphur dioxide acts as an oxidizing agent when it reacts with a more powerful reducing
agent than itself.
a. When sulphur dioxide is dissolved in water, it forms sulphurous acid which is a
bleaching agent. Sulphurous acid takes up oxygen from the dye to form sulphuric
acid. The removal of oxygen from a dye converts the dye to a colourless
compound. This is essentially a different reaction from that of other bleaching
agents, which oxidize the dye to a colourless compound.
It is prepared by passing a mixture of dry sulphur dioxide and dry air over heated vanadium (V)
oxide (or platinised asbestos) at a temperature of 450 500 o C under a pressure of 200
atmospheres. Sulphur trioxide is seen as dense white fumes which are solidified in a freezing
mixture of ice and sodium chloride.
Sulphur dioxide may contain some impurities such as arsenic compounds which may poison the
catalyst, that is, make the catalyst ineffective. Therefore sulphur dioxide is cleaned to remove
the impurities then is dried.
Then sulphur dioxide is mixed with air and passed along heated pipes containing pellets of
vanadium pentoxide (V2O5) (catalyst) at a temperature of 450 500 oC under a pressure of 200
atmospheres. Sulphur trioxide is formed.
Sulphur trioxide formed is dissolved in concentrated sulphuric acid to produce a fuming liquid
called oleum.
The oleum is diluted with a known amount of water to give concentrated sulphuric acid.
Note: Sulphur trioxide is not dissolved in water directly because the reaction is too exothermic
and the heat produced from the reaction vapourises the acid forming only tiny droplets of the
acid leading to a spray of sulphuric acid which would affect the workers in the factory.
Has a high affinity for water (hygroscopic) and that is why it is used as a drying agent
Note: Sulphuric acid has a high affinity for water. Never add water to the concentrated acid
because it can explode. It is therefore advisable to add the acid to water rather than water to
acid.
c) It neutralises alkalis like sodium hydroxide solution to form a salt and water only
3) As a dehydrating agent.
Sulphuric acid has a very high affinity for water and can remove it from substances including air,
that is, it is hygroscopic. It can be used as a drying agent for most gases.
a) When concentrated sulphuric acid is poured onto sugar (sucrose) in a beaker the sugar
turns yellow then brown and finally a black spongy mass of charcoal rises filling the
beaker. Steam is given off and the whole mass becomes very hot. The acid takes out the
elements of water from sugar leaving a black mass of carbon.
b) When concentrated sulphuric acid is added to blue crystals of copper (II) sulphate
(hydrated) and warmed, they change to a white solid of anhydrous copper (II) sulphate as
water of crystallisation is removed by concentrated sulphuric acid.
c) Other substances which are dehydrated by concentrated sulphuric acid include ethanol,
methanoic acid and oxalic aid
C2H5OH(l) C2H4 (g) + H2O(l)
Ethanol
HCOOH(l) CO(g) + H2O(l)
Methanoic acid
H2C2O4 (s) CO (g) + CO 2 (g) + H2O(l)
Oxalic acid
4) Action of concentrated sulphuric on nitrates and chlorides
Concentrated sulphuric acid displaces hydrochloric acid from metallic chlorides and nitric
acid from nitrates e.g. it reacts with sodium chloride when heated forming white fumes of
hydrogen chloride gas which dissolves in water forming hydrochloric acid.
2NaCl (s) + H2SO4 (aq) Na2SO4 (aq) + 2HCl (g) (With heating)
NaCl (s) + H2SO4 (aq) NaHSO4 (aq) + HCl (g) (Without heating)
25.9. Sulphates
Action of heat on sulphates (SO42-)
Most sulphates are hydrated and when heated, they lose their water of crystallisation to form
anhydrous salts which are resistant to further heating and therefore do not decompose. Therefore
hydrated sulphates do not decompose on heating e.g.
MgSO4.7 H2O(s) MgSO4(s) + 7 H2O(l)
Na2SO4.10H2O(s) Na2SO4(s) + 10H2O(l)
When a blue solid of hydrated copper (II) sulphate is heated, water vapour is given off as water of
crystallisation is lost, giving a white solid (residue). On further heating, it decomposes to form white
fumes of sulphur trioxide and a black residue of copper (II) oxide.
CuSO4.5 H2O(s) CuSO4(s) + 5H2O(l)
Blue White
CuSO4(s) CuO(s) + SO3 (g)
Overall equation:
On heating hydrated iron (II) sulphate (green), it loses its water of crystallisation
FeSO4.7 H2O(s) FeSO4(s) + 7 H2O(l)
On further heating, the anhydrous iron (II) sulphate formed decomposes to give white fumes of
sulphur trioxide together with sulphur dioxide and leaves a brown residue of iron (III) oxide
FeSO4(s) Fe2O3(s) + SO2(g)+ SO3(g)
When ammonium sulphate is heated it decomposes to give ammonia, sulphur trioxide and water
(NH4)2SO4(s) NH3 (g) + SO3(g) +H2O(l)
To the solution add dilute hydrochloric acid followed by barium chloride solution.
To the solution add dilute nitric acid and barium nitrate solution. A white precipitate is formed.
Ionic equation
Note: Carbonate ions (CO32-) and sulphite ions (SO32-) are precipitated as barium carbonate and
barium sulphite respectively, if carbonate and sulphite ions are present in solution.
The sulphite ions (SO42-) remain in solution because they do not react with dilute hydrochloric
acid acid or nitric acid.
Also lead (II) nitrate solution forms a white precipitate of lead (II) sulphate with a sulphate.
Add concentrated hydrochloric acid to iron (II) sulphide. Effervescence occurs and the hydrogen
sulphide formed is collected over warm water because it is quite soluble in cold water
To prepare hydrogen sulphide from sulphur, iron (II) sulphide is first prepared by heating the
mixture of iron and sulphur.
Fe(s) + S(s) FeS(s)
Hydrogen sulphide can be dried by using anhydrous calcium chloride. Concentrated sulphuric
acid is not used to dry the gas because it reacts with hydrogen sulphide forming a yellow
precipitate of sulphur.
Black
It reduces reducing agents and itself oxidized to sulphur which appears as a pale yellow
precipitate
a. Hydrogen sulphide reduces iron (III) chloride to iron (II) chloride. When
hydrogen sulphide is passed through iron (III) chloride solution (yellow solution)
a yellow precipitate of sulphur appears. On filtering, a green solution of iron (II)
chloride appears as the filtrate
Hydrogen sulphide burns with a blue flame forming water and sulphur dioxide.
When hydrogen sulphide is burnt in a limited supply of air, a yellow deposit of sulphur is
formed. The oxygen supply cannot oxidize the gas completely and therefore free sulphur is
deposited.
26.1. Nitrogen
Occurrence of nitrogen
Atomic number 7, atomic mass 14. It is one of the main elements needed for plant growth.
Nitrogen is the most abundant gas in the atmosphere, occupying about 78 per cent by
volume. It occurs in nature in a combined state as in minerals such as sodium nitrate. It is
found in living things in form of proteins.
Sodium hydroxide absorbs and removes carbondioxide from the air mixture and forms sodium
carbonate
It removes the oxygen by reacting it with the hot copper leading to formation of copper oxide
It can be dried by passing the gas through a U-tube containing glass beads wetted
with concentrated sulphuric acid to dry it and then collected in a syringe.
Nitrogen formed by this method is not pure. It contains several impurities, mainly the
noble gases as well as unreacted oxygen.
Commercially nitrogen is manufactured through fractional distillation of liquid air.
1. Nitrogen extinguishes a burning splint and the gas does not burn. This distinguishes it
from other gases that support burning like oxygen and dinitrogen oxide or any
combustible gas such as hydrogen, carbon monoxide, hydrogen sulphide.
2. Nitrogen has no smell. This distinguishes it from gases such as sulphur dioxide ammonia,
hydrogen chloride.
3. Nitrogen has no action on lime-water. This distinguishes it from carbon dioxide.
26.5. Properties of nitrogen
5. Nitrogen reacts only with the reactive metals (magnesium and calcium). When these
metals are heated strongly, they burn in nitrogen forming the corresponding nitride,
which is white in colour.
3Mg(s) + N2(g) Mg3N2(s)
The heat produced by the burning magnesium ribbon or calcium is strong enough to
break the triple bond in the nitrogen molecule forming free nitrogen atoms. The free
atoms are very reactive and combine with these metals to form a nitride. If a burning
wooden splint is placed in a jar of nitrogen, it gets extinguished. This is because the heat
it produces is not sufficient to break the tripple bond between the nitrogen atoms
The nitrides dissolve in water to form the corresponding hydroxide and ammonia.
6. In thunderstorms, a small amount of nitrogen reacts with the oxygen in the air to form
nitrogen monoxide and nitrogen dioxide.
N2(g) + O2(g) 2NO(g)
The electrical discharge in a thunderstorm provides sufficient energy for this reaction to
occur.
4. Because of its unreactive nature, nitrogen is used as an inert atmosphere for some
processes and chemical reactions. For example, empty oil tankers are filled with nitrogen
to prevent fires.
Remove the cover from a gas-jar of nitrogen monoxide. Brown fumes are at once
produced due to oxidation of the gas by oxygen of the air to nitrogen dioxide.
The experiment is set up as shown in figure 2.3. When lead(II) nitrate is heated, it makes a
cracking sound giving off a brown gas (nitrogen dioxide) and oxygen. Nitrogen dioxide is
liquefied in the freezing mixture and collects in the tube as green liquid. The oxygen passes
on as gas and escapes or it is collected over water.
Lead(II) nitrate is the most suitable to use because it crystallizes without water of
crystallization, which is found in crystals of most nitrates and which would interfere with the
preparation.
Nitrogen dioxide may also be prepared by the action of concentrated nitric acid on copper
turnings. The gas is collected by downward delivery in a gas-jar.
4. Nitrogen dioxide neutralizes alkalis forming a mixture of their corresponding nitrates and
nitrites. In this case nitrogen dioxide acts as an acid.
2NaOH(aq) + 2NO2(g) NaNO3(aq) + NaNO2(aq) + H2O(1)
6. Nitrogen dioxide oxidizes red hot metals and itself reduced to nitrogen.
4Cu(s) + 2NO2(g) 4CuO(s) + N2(g)
26.9. Ammonia
The usual drying agents such as concentrated sulphuric acid and anhydrous calcium chloride
are not used because ammonia reacts with them to form ammonium sulphate and tetraamine
calcium chloride respectively.
Ammonia is removed from the mixture of gases by cooling the mixture with a freezing
mixture. It is only ammonia that liquefies and can be removed from the mixture. The
unreacted nitrogen and hydrogen are recycled (figure 2.5).
Nitrogen used in this process is obtained by fractional distillation of liquid air and hydrogen
is obtained from natural gas or electrolysis of brine.
This is an experiment to demonstrate the high solubility of ammonia gas in water. A large round
thick walled flask is filled with ammonia gas. It is then fitted with two glass tubes C and D with
clips at one end (figure 2.6). The flask is inverted over a trough of water and the clip on tube D
opened to allow in a few drops of water and then closed. These are shaken with ammonia to
dissolve it. If the red litmus solution is added to water in the trough, water in the flask will turn to
blue indicating that it is an alkaline gas which dissolved in water. The clip on the tube C is
opened. Water runs up the tube and spreads at the end of the tube forming a fountain.
The few drops of water, which entered through tube D, dissolved all the ammonia gas in the flask
so that a partial vacuum was created in the flask. When the clip on tube C was opened,
atmospheric pressure pushed the water up the tube forming a fountain.
Here ammonia behaves as a reducing agent. A similar reaction takes place with the oxides of lead
and iron.
This experiment can also be used to demonstrate that ammonia contains nitrogen.
The figure 2.8 shows how ammonia is burnt. The role of the glass wool is to distribute oxygen
evenly throughout the gas vessel.
In presence of catalyst, ammonia is oxidised to nitrogen monoxide.
The figure 2.9 shows that set of the experiment. A hot platinum or copper wire which acts
as a catalyst is suspended in a beaker of concentrated ammonia and oxygen is bubbled
through the solution. The metal catalyst remains red-hot because the reaction is
exothermic. Brown fumes of nitrogen dioxide, which are formed due to oxidation of
nitrogen monoxide, are observed.
In excess ammonia, dense white fumes of ammonium chloride are formed. Hydrogen
chloride formed reacts with excess ammonia to form the white fumes, which later settle
to a white solid.
Ammonia
In the laboratory, ammonium salts are made by reacting the appropriate acid with ammonia.
For example, ammonium sulphate is made by neutralizing sulphuric acid with ammonia.
In industry, ammonium sulphate cannot be made using sulphuric acid, as the later is very
expensive. Instead, it is made by reacting ammonium carbonate with calcium sulphate.
Ammonium carbonate is first prepared by saturating ammonia solution with carbon dioxide.
Solid calcium sulphate is added and the mixture is stirred forming ammonium sulphate
solution and calcium carbonate.
(NH4)2CO3(aq) + CaSO4(s) (NH4)2SO4(aq) + CaCO3(s)
2. Ammonium sulphate decomposes on heating into ammonia and sulphuric acid. Although
the reaction is similar to that of ammonium chloride no sublimation occurs because
sulphuric acid is less volatile than ammonia. The ammonia gas escapes before sulphuric
acid volatiles such that the two cannot recombine.
(NH4)2SO4(s) 2NH3(g) + H2SO4(g)
Dinitrogen oxide is a colourless gas. It is fairly soluble in water and neutral to litmus. It is
denser than air and a glowing splint is relit when lowered into a gas-jar containing
dinitrogen oxide. The heat decomposes dinitrogen oxide into oxygen and nitrogen. It is
oxygen that relights the glowing splint.
Exercise
State what would be observed and write an ionic equation for the reaction that would take
place when aqueous ammonium chloride was
Ammonia solution precipitates metal hydroxides from solutions containing the metal ions.
When a few drops of ammonia solution are added to a solution of copper(II) ions, a blue
precipitate is formed.
A solution of zinc ions forms a white precipitate with a few drops of aqueous ammonia. The
precipitate dissolves in excess ammonia solution to form a colourless solution containing
complex tetraamine zinc ions.
Iron(II), iron(III), lead(II) and aluminum ions form precipitates of the hydroxides with
aqueous ammonia which are insoluble in excess ammonia solution.
(green)
(brown)
(white)
(white)
A solution of aluminium, Zinc and lead(II) ions reacts with sodium hydroxide solution to
form a white precipitate that is soluble in excess sodium hydroxide solution to form a
colourless solution.
A13+(aq) + 3OH-(aq) A1(OH)3(s)
(aluminate ion)
(zincate ion)
(plumbate ion)
Iron(II) and iron(III) ions in solution, react with sodium hydroxide solution to give a green
and brown precipitate respectively, insoluble in excess sodium hydroxide solution.
Magnesium and calcium ions in solution react with sodium hydroxide solution to give a
white precipitate insoluble in excess sodium hydroxide solution.
Brown fumes of nitrogen dioxide are produced during heating because of thermal
decomposition of nitric acid.
The experiment must be carried out in all-glass apparatus because nitric vapour attacks
rubber and cork.
Nitrogen monoxide is cooled and reacts with oxygen from excess air to produce brown fumes
of nitrogen dioxide.
Nitrogen dioxide together with excess air is dissolved in hot water to form nitric acid.
1. It is a strong acid.
2. It is a powerful oxidizing agent.
26.9.15.5. Nitric acid acting as a strong acid
Nitric acid is a very strong acid, being almost completely ionized in dilute solution with the
production of the hydrogen ion and the nitrate ion.
This ionization confers on it the usual acidic properties, modified to some extent by powerful
oxidizing action of the acid.
(b) It reacts with oxides and alkalis to form salt and water only.
CuO(s) + 2HNO3(aq) Cu(NO3)2(aq) + H2O(l)
Magnesium is the only metal that liberates hydrogen with nitric acid and only when the
acid is very dilute. Other metals are oxidised by the acid to the corresponding nitrates.
If the acid is 50% concentrated (equal volume of water as the volume of acid), nitrogen
monoxide is formed.
Lead reacts with nitric acid in a similar way. Aluminium and iron are made assive
because of the formation of the oxide layer, which forms a protective layer over the metal
and stops further reaction.
(ii) When a piece of red-hot charcoal is put into concentrated nitric acid, it continues to
burn and brown fumes are formed.
C(s) + 4HNO3(l) CO2(g) + 4NO2(g) + 2H2O(l)
(iii) When red phosphorus is gently heated with moderately dilute nitric acid, brown
fumes are formed.
P(s) + 5HNO3(aq) H3PO4(aq) + H2O(l) + 5NO2(g)
Lead(II) nitrate makes a cracking sound when heated. The sound is due to the fact that the air
inside the crystals splits them when it expands due to heating. A brown mixture of nitrogen
dioxide and oxygen is given off. Lead(II) oxide (residue) is brown when hot and yellow
when cold.
Most metallic nitrates decompose to a metal oxide, nitrogen dioxide (brown fumes) and
oxygen gas which relights a glowing splint.
(white) (white)
(white) (white)
(white)
(green) (black)
Zinc oxide is yellow when hot and white when cold. Zinc nitrate and copper(II) nitrate are
hydrated and when heated do not produce a cracking sound. They melt first and dissolve in
their water of crystallization forming a solution. The solution then evaporates and when most
of the water has evapourated, decomposition starts. Mercury(II) nitrate and silver nitrate
decompose to the metal, nitrogen dioxide and oxygen.
Exercise
When a green compound W was heated strongly, a brown gas was given off and a black
residue remained.
Nitric acid formed then oxidizes iron(II) to iron(III) and itself reduced to nitrogen
monoxide.
Fe2+(aq) Fe3+(aq) + e-
Nitrogen monoxide combines with the remaining iron(II) sulphate to form the dark brown
compound, nitroso-iron(II) sulphate.
The ring disappears if the solution is shaken. This is because when concentrated sulphuric
acid and water mix, a lot of heat is evolved which decomposes the compound.
Exercise
(b) Write equation for the reaction between nitric acid and ammonia.
(c) State one use of the product in (b).
2. (a) Describe the industrial preparation of nitric acid from ammonia. Your
description should include equations for the reactions that occur.
(b) Explain what happens when concentrated nitric acid is added to copper.
(c) Describe one chemical test that can be used to confirm the presence of a
nitrate.
(d) State what would be observed if concentrated nitric acid was heated with
3. (a) Draw a labeled diagram of the apparatus that can be used to prepare ammonia
in the laboratory.
(b) Describe an experiment that can be carried out to show that ammonia is a
(c) A copper coil was heated strongly and held over a concentrated solution of
(c) Name one other metal that reacts with nitrogen in a similar way to magnesium.
5. When compound x is heated with concentrated sulphuric acid, a gas which forms dense
white fumes with ammonia is liberated.
(a) Identify the anion in x.
(b) Write an ionic equation for the reaction between a solution of x and silver nitrate.
(c) State what would be observed if lead(II) nitrate solution was added to solution of x
and the mixture heated.
7. Dilute nitric acid reacts with copper to form a colourless gas, which on exposure to air
gives brown fumes soluble in water.
(a) Write the equation for the reaction between copper and nitric acid.
(b) Name the colourless gas.
(c) Explain how the brown fumes are formed.
(d) Write the equation to show the reaction between water and the brown fumes.
8. Excess lead(II) oxide was added to warm dilute nitric acid and the mixture was stirred.
After cooling, the mixture was filtered and a solution of ammonium hydroxide was added
to the filtrate.
(a) Write an equation for the reaction between lead(II) oxide and nitric acid.
(b) State what was observed when ammonium hydroxide solution was added to the
filtrate drop wise until in excess.
(c) Write an equation for the reaction in (b) above.
9. What would be observed if dilute sodium hydroxide solution was added drop wise until
in excess to a solution of
(i) Fe2+ salt.
(ii) Fe3+ salt.
10. (a) Describe how zinc sulphate crystals can be prepared from zinc in the
laboratory.
(b) A small amount of zinc sulphate was dissolved in dilute nitric acid and the
(i) State what would be observed when sodium hydroxide solution is added to
(c) (i) State what would be observed when aqueous ammonia is added to the
11. Study the figure 2.13 and answer questions that follow.
(a) Name
(i) gas X
(ii) liquid W.
(iii) one reagent that can be used to identify W.
(iii) Name another oxide that shows similar reaction with gas X.
12. Lead(II) nitrate was heated strongly in the apparatus shown in figure 2.14.
(a) Identify
(iii) Liquid Q.
(iv) Gas X.
(b) (i) State what was observed in the test-tube during the heating.
(ii) Write the equation for the reaction that took place.
(c) To the residue was added dilute nitric acid and the mixture warmed.
(i) Write the equation for the reaction.
(ii) State what was observed.
(d) To the resultant product in (c) was added sodium hydroxide solution drop wise until
in excess.
(i) State what was observed.
(ii) Write the equation(s) for the reaction(s) that took place.
13. Figure 2.15 shows the apparatus used for combustion of ammonia.
(a) Name
(i) gas Q
(ii) X
(b) Write the equation for the reaction that occurred in the test-tube.
(c) Name another substance that can be used instead of ammonium sulphate.
(d) State the role of
(i) The glass wool.
(ii) Calcium oxide.
(e) Explain why concentrated sulphuric acid is not used instead of calcium oxide.
(f) Write the equation for the combustion of ammonia.
(g) State one industrial use of ammonia.
14. Figure 2.16 shows an experimental setup for the laboratory preparation of nitrogen gas.
(iii) Write the equation for the reaction taking place in the combustion tube.
(d) Explain why nitrogen gas collected in this experiment is not pure.
27. Chlorine and its compounds
27.1. Chlorine
Chlorine is element number 17 in the periodic table of elements. It belongs to group VII, the
halogens. Chlorine comes from the Greek word chloros, meaning green.
The experiment is set up as shown in figure 3.1. Concentrated hydrochloric acid is poured
into a flask containing manganese(IV) oxide and the flask shaken well. The mixture is heated
and chlorine gas formed is passed through a bottle containing water to dissolve any fumes of
hydrogen chloride, which are produced from concentrated hydrochloric acid. It is then dried
by passing it through concentrated sulphuric acid and collected in a gas-jar by downward
delivery because it is denser than air.
The experiment need not be conducted in a fume-chamber, if the gas is collected over
brine.
Reaction at cathode:
Na+(aq) + e- Na(s)
Reaction at anode:
Sodium ion is discharged because it requires less energy than the discharge of hydrogen ions
in case a mercury cathode is used.
A tube containing the drying agent, anhydrous calcium chloride, is connected to the bottle to
prevent water from the atmosphere from entering the bottle as this would be absorbed by
iron(III) chloride which is very deliquescent. Excess chlorine, which is poisonous, escapes into
the fume chamber.
The formation of iron(III) chloride and not iron(II) chloride shows that chlorine is an oxidizing
agent. Iron(II) chloride is immediately oxidised by chlorine to iron(III) chloride.
Sodium chloride can be prepared in similar way.
Iron(II) chloride (white solid) is made in the same way, using dry hydrogen chloride
instead of chlorine.
Sodium, potassium, calcium and magnesium burn in chlorine forming white fumes of the
chloride which settle to a white solid.
Note: When iron(III) chloride crystals are dissolved in water they give a deep yellow
solution from which yellow crystals of hydrated iron(III) chloride may be obtained by
evaporation to the point of crystallization.
When chlorine gas is bubbled through cold aqueous alkalis, the hypochlorite and the
chloride of the metal are formed.
When chlorine is passed into hot concentrated alkalis, a mixture of the chlorate and
the chloride is formed.
Hypochlorous acid is a very reactive compound and readily gives up its oxygen to the
dye, to form a colourless compound, that is, the dye is oxidised to a colourless
compound.
Colourless
C1 Chlorine
Br Bromine
When chlorine gas is bubbled into a solution of potassium bromide in water, the
colourless solution immediately turns red due to formation of bromine water.
Chlorine displaces iodine from potassium iodine solution forming a dark brown solution
due to formation of iodine.
Bromine can displace iodine from iodides but cannot displace chlorine from chlorides.
Therefore on addition of few drops of bromine to a solution of potassium iodine in water,
the solution becomes brown due to the formation of iodine.
2KI(aq) + Br2(l) 2KBr(aq) + I2(aq)
Concentrated sulphuric acid is added to sodium chloride in the flask as shown in figure 3.6.
Effervescence occurs and misty fumes of a gas are formed. The gas is passed through a bottle
containing concentrated sulphuric acid to dry it and collected by upward displacement of air
since the gas is denser than air.
4. It reacts with alkalis and basic oxides producing salt and water only.
NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)
These ions are responsible for conducting electricity in the solution. The oxonium ion
liberates hydrogen with the more electropositive metals and carbon dioxide with carbonates
and hydrogencarbonates.
Thus these properties are not shown by hydrogen chloride in methylbenzene, that is, the
solution does not contain hydrogen or oxonium ions responsible for acidic characteristics and
the solution contains no ions which carry an electric current.
To the white precipitate add ammonia solution. The precipitate dissolves to form a colourless
solution.
Exercise
1. (a) Draw a well labeled diagram to show how a sample of dry hydrogen chloride can be
prepared in the laboratory.
(b) Dry hydrogen chloride gas was passed over heated iron filings. Write an equation for
the reaction that took place.
(c) The solid product in (b) was dissolved in water and aqueous sodium hydroxide added
to the resultant solution drop wise until in excess.
(d) Chlorine gas was passed through a solution of the product in (b)
(e) (i) Name one reagent that can be used to test for the anion formed in (d).
(ii) State what is observed when the reagent you have named is used.
2. (a) A mixture consists of sulphur and iron filings. Explain briefly how a sample of
sulphur can be obtained from the mixture.
(b) A sample of the mixture in (a) was heated in a porcelain dish.
3. State what would be observed and write ionic equation(s) for the reaction(s) that take
place when
(i) A solution of silver nitrate is added to potassium chloride solution.
(ii) A solution of barium chloride is added to sodium sulphate solution.
4. Chlorine can be prepared in the laboratory from hydrochloric acid.
(i) Name the other reagent used in the preparation of chlorine.
(ii) State the conditions for the reaction.
(iii) Write an equations for the reaction which takes place between hydrochloric acid and
the reagent you have named in (i).
5. During the preparation of chlorine in the laboratory, the gas may be passed through water
and concentrated sulphuric acid before collection.
(a) State the use of
(i) water
(ii) concentrated sulphuric acid
6. (a) Draw a labeled diagram of the apparatus you would use to prepare chlorine in the
laboratory, using potassium permanganate.
(b) State what is observed when
7. (a) With the aid of a well labeled diagram describe an experiment in the laboratory to
show that hydrogen chloride gas is very soluble in water.
(c) Hydrogen chloride was dried and passed over heated iron filings.
8. (a) (i) Draw a well labeled diagram to show the preparation of iron(III) chloride
using chlorine.
(b) (i) State what would be observed if aqueous ammonia was added to a solution
of iron(III) chloride.
10. When a compound M is heated with concentrated sulphuric acid, a gas that forms dense
white fumes with ammonia is liberated.
(a) Identify the anion in M
(b) (i) State what would be observed when a solution of M is added to silver nitrate
solution.
(ii) Write an ionic equation for the reaction which occurs in (i) above.
(c) State what would be observed when lead(II) nitrate solution is added to a solution of
M and the mixture warmed.
11. (a) The substance Y reacts with solid chloride to produce hydrogen chloride.
(i) Identify Y.
(ii) State the conditions for the reaction.
(iii) Write the equation for the reaction.
(b) (i) Name the substance that is formed when hydrogen chloride is dissolved in
water.
(ii) Explain why an aqueous solution of hydrogen chloride is an electrolyte
whereas the solution of the gas in methylbenzene is a non-electrolyte.
28. Extraction of metals
28.1. Introduction
The major factor determining the method used for extraction of metals from their ores is the
position of the metals in the electrochemical series. An ore is a naturally occurring substance
from which a metal can be extracted.
Very reactive metals, that is, those higher in the activity series occur mainly as chlorides.
They are extracted by electrolysis of their fused salts. Such metals include potassium,
sodium, calcium, magnesium and aluminium.
Metals in the middle of the series such as zinc, iron, lead and copper mainly occur as oxides,
carbonates and sulphides. They are extracted by reduction of the ore. Chemical reduction
involves the extraction of a metal from its ore by heating the ore with a strong reducing agent
such as coke. This method is used in extraction of iron. Thermal reduction involves roasting
(heating directly in air). It is applicable in the extraction of metals such as zinc, copper and
lead. At some stage in this process, the method might be intergrated with chemical reduction.
Metals lower in activity series, that is, mercury, silver and gold mainly occur as free metals in
the earths crust. They are mainly dug up in the pure form.
1. Magnetic separation
Many metals ores containing magnetic impurities are partially refined in this way. The
crushed ore is placed on a conveyer belt, which has a magnetic roller at one end. As the
ore passes over the magnetic roller, it is separated into two parts, one containing the
partially refined ore and the other containing all the magnetic impurities.
2. Froth flotation
Froth flotation is the process in which the ore is powdered, mixed with oil and water, and
air is brown through. The froth contains the ore which is skimmed off.
3. Hydrolic method
The rock containing the ore is blasted with a stream of water and earthly matter is washed
away, leaving the heavier ores.
4. Mechanical sorting
5. Solvent extraction
The ore is dissolved in the solvent. Some components of the ore dissolve in the solvent
while others do not.
28.3. Sodium
Sodium occurs as sodium carbonate (soda ash), sodium nitrate and sodium chloride (rock
salt).
Na+(l) + e+ Na(s)
28.4. Copper
The principal ores of copper are copper pyrites (CuFeS 2), cuprite (Cu2O), copper(I) sulphide
(Cu2S) and malachite (CuCO3.Cu(OH)2).
28.4.1. Extraction of copper
Copper pyrites is the ore usually used for the extraction of copper and there are three stages
involved.
Sulphur dioxide escapes from the top of the furnace. By adding silicon dioxide, SiO 2, and
heating in absence of air, the iron(II) oxide is converted into a slag of iron(II) silicate,
FeSiO3, which floats on top of the molten copper(I) sulphide and it is tapped off.
This copper produced is impure and is called blister copper because of the blistered
appearance on the copper surface caused by the escaping gases on cooling.
28.4.4. Refining of the impure copper
The impure copper is purified (refined) by electrolytic process using copper(II) sulphate
solution as the electrolyte. The cathode is pure copper and the impure copper is made the
anode as shown in figure 8.2.
During electrolysis, the copper atoms of the anode lose electrons to form copper(II) ions
which dissolve in the solution.
Then the copper(II) ions are attracted to the cathode where they gain electrons and become
copper atoms.
The overall effect is that copper gradually dissolves from the anode and is deposited on the
cathode. Copper from the cathode is removed by stripping.
Impurities which are higher than copper in the activity series, such as iron, also dissolve from
the anode but are not deposited on the cathode. They accumulate in solution in the
electrolyte. Impurities which are lower than copper in the activity series do not dissolve at
all. They fall to the bottom of the container as sludge. The elements which were present in
the original copper ore.
28.4.5. Uses of copper
1. Copper is used as a conductor of electric power in wires and cables.
2. It is used for making bronze which is used for manufacturing ball bearings.
3. It is used for making kettles for brewing beer.
4. Used for making roofing sheets because it is corrosion resistant.
28.5. Iron
The main iron ores are haematite (Fe2O3), magnetite (Fe3O4), iron disulphide (pyrites, FeS2)
and spathic iron ore (FeCO3).
Higher up the furnace, the source of oxygen is less and more coke combines with carbon
dioxide produced to form carbon monoxide.
Molten iron runs to the bottom of the furnace and is tapped off into moulds where it is
solidified. The moulds are called Pigs and therefore this impure form of iron is called pi-
iron. Limestone is decomposed by heat to calcium oxide and carbon doxide.
The iron contains impurities such as silicon dioxide (sand), which combine with calcium
oxide to form a molten slag that floats on top of the molten iron and it is tapped off.
Calcium oxide (from decomposition of limestone) reacts with these solid oxides, forming
calcium phosphate and calcium silicate, which float on the surface of the iron as slag.
Wrought iron is malleable. It is very tough and therefore can withstand some strain. It can be
used to make iron nails, iron sheets and agricultural implements.
28.5.4. Steel
Steel is an alloy of pure iron with a small percentage of carbon and other elements. Other
such as tungsten, chromium, nickel and manganese which are added to produce types of steel
(table 8.1) with different properties. Steel is hard, tough, strong and malleable.
Mild steel 99% iron, 0.5% carbon Car bodies, large structures
Stainless steel 74% iron, 18% chromium, Cutlery, kitchen sinks, surgical
8% nickel instruments
28.5.6. Alloy
An alloy is a metallic substance consisting of a mixture of two or more metals or a mixture of
a metal with a non-metal. Alloys have suitable properties when compared with pure metals.
Alloys are usually less malleable and ductile than pure metals. They also have low melting
points and electrical conductivity than pure metals. For example, solder has a lower melting
point than lead and tin. Because of its low melting point, solder can be used to join metals.
Examples of common alloys are given in table 8.2.
Exercise
1. (a) Extraction of metals is essentially a reduction process. Explain the statement using
extraction of iron as an example. Write the equation to illustrate your answer.
(b) State the conditions under which iron may react with
(i) Oxygen.
(ii) Water.
(iii) Chlorine.
(c) Write an equation for the reaction in (b) (ii) and (iii).
(ii) Name the elements which are used in making stainless steel.
(iv) Suggest a reason why the use of stainless is preferred to that of pure iron.
2. (a) Sodium metal is extracted by electrolysis of fused sodium chloride to which calcium
chloride has been added.
(i) Give a reason for the addition of calcium chloride.
(ii) Name a material that can be used as the cathode and another that can be used as
the anode.
(iii) Write equations for the reaction that take place at each electrode.
(v) Name one other element that can be extracted by similar method.
(b) Name a place in Uganda where a plant for the extraction of sodium could be
constructed. Give a reason for your answer.
(c) Describe what would be observed if a small piece of sodium metal was heated and
quickly plunged into a gas-jar of oxygen. Write an equation for the reaction that takes
place.
(i) Coke
(ii) Limestone
29. Organic chemistry
29.1. Introduction
Organic chemistry is a branch of chemistry dealing with compounds of carbon except oxides
of carbon, carbonates, hydrogencarbonates and carbides of metals. Carbon has the ability to
form bounds to itself. These bounds are very strong and can be single, double or triple bonds.
Owing to this fact, chains of varying sizes can be formed which contribute to a wide range of
stable compounds. These compounds are known as organic compounds. Carbon forms four
covalent bonds making it possible to have different groups attached to the chains of carbon
atoms. This will also lead to a wide diversity of compounds being formed.
29.2. Hydrocarbons
Hydrocarbons are compounds containing hydrogen and carbon atoms. They have a molecular
formula, CxHy, where x and y are whole numbers. They are classified into several types
according to their structures. The main classes of hydrocarbons are alkanes, alkenes and
alkynes.
i) All members conform to a general molecular formula. For example, in case of alkanes,
the general molecular formula is CnH2n+2 where n1.
ii) Each member differs in molecular formula from the next by CH 2, for example members
of the alkanes are CH4, C2H6, C3H8 and so on.
iii) All members show similar chemical reactions though they vary in vigour.
iv) The physical properties of members change gradually in the same direction along the
series.
H H
H C C H Ethane
H H
Structural formula
The structural formula shows the sequence and arrangement of atoms in a molecule. For
example, the structural formula of propane is shown below.
H H H
H C C C H or CH3CH2CH3
H H
Exercise 9.1
(a) Hexane
(b) Butane
Alkyl groups
Removal of one hydrogen atom from an alkane molecule leaves a monovalent group called
an alkyl group. Table 9.1 shows some groups and their molecular formulae.
Nomenclature of alkanes
The general rules of naming organic compounds were laid down by the International Union
of Pure and Applied Chemistry (I.U.P.A.C.)
(a) The first step is to choose the longest chain of carbon atoms which is called the parent
chain. In the structural formula below, the longest chain has eight carbon atoms and thus
it is taken as the parent chain and it is called octane.
CH3
4
3
2
CH3
It is possible at times to have longest chains of carbon atoms and the one which is taken
as a parent chain is that one which has a higher number of the chains. In the structural
formula below, the parent chain has three side chains and it is called heptanes.
(b) After identifying the parent chain, the carbon atoms are numbered from one end to the
other. The numbering should be in such a way that the carbon atoms carrying the side
chain gets the lowest number and the position of the side chain is indicated by the number
assigned to the carbon atom to which it is attached.
(c) If there are more than one side chain, then the numbering of the carbon atoms is done in
such a way that the sum of the numbers used to locate the side chains is lowest and this is
called the lowest sum rule. Consider the structural formula below.
From left to right, the sum of the numbers where the side chains are attached
= 2 + 3 + 6 = 11
From right to left, the sum of the numbers where the side chains are attached
= 2 + 5 + 6 = 13
Therefore, the numbering is from left to right and hence the name of the compound is
2,3,6-trimethylheptane.
(d) If there is more than one type of side chain, then the side chains are prefixed and should
be put in alphabetical order preceding the name of the parent chain.
In case of a particular side chain appearing twice of three times, then di, tri, tetra, pent
etc, are used. The locants (numbers used to locate a side chain) are written in increasing
order separated by commas and hyphens (-) separates the numbers from the prefix. Refer
to the example in (c), that is 2,3,6-trimethylheptane.
Exercise
(c) Ch
(d) Ch
(e) Ch
2. Write the structural formula of the following compounds.
(a) 2,2-dimethylpentane
(b) 3-ethylhexane
(c) 2-methylpropane
(ii) They are insoluble in water but soluble in non-polar solvents like
trachloromethane.
(iii) They are less dense than water. Their densities rise gradually with increasing
molecular mass.
During combustion, a great deal of heat is liberated. Owing to that fact, alkanes are used
as fuels for industrial and domestic purposes. For example, butane is used in gas cigarette
lighters. Methane is found in natural gas and bio gas. It is used in gas appliances. Butane
found in petrol is used to run petrol engines.
(ii) They undergo substitution reactions with the halogens, producing corresponding
compounds. A substitution reaction is a reaction in which one atom or group of atoms in a
molecule is replaced by another. For example, methane reacts with chlorine forming
chloromethane and hydrogen chloride, the reaction being catalysed by light (photo catalysis).
One hydrogen atom of the methane molecule has been replaced by a chlorine atom. In a
similar way, excess of chlorine may produce dichloromethane (CH2Cl2), trichloromethane
(CHCl3) and tetrachloromethane (CCl4).
Dichloromethane
Trichloromethane
tetrachloromethane
Exercise
When a mixture of ethane and chlorine was exposed to sunlight, the colour of chlorine
disappears.
(a) Write the equation for the reaction that takes place when ethane is mixed with a limited
amount of chlorine.
(b) Name the type of reaction.
(c) What is the role of sunlight in this reaction?
Isomerism
Isomerism is the occurrence of two or more compounds with the same molecular formula but
different structural formulae. Compounds which have the same molecular formula but
different structural formulae are called isomers. All alkanes with more than four carbon
atoms have more than one structure for a given molecular formula, that is, they exhibit
isomerism. The greater the number of carbon atoms in an alkane, the greater the number of
possible isomers.
The easiest way of finding the isomers is to draw the longest chain of carbon atoms first and
reduce it by one carbon atom at a time.
Compound Q has a molecular formula, C6H14. Give the structural formulae and names of
possible isomers of compound Q.
Fraction Uses
oil petrol
The molecules of petrol contain 5-9 carbon atoms and gases are mainly alkenes containing 2
to 4 carbon atoms.
(ii) The solid by-product is used as fertilizers since it contains a high nitrogen content.
29.6.4. Disadvantages of bio gas production
Some of the gases contained in bio gas are air pollutants. When bio gas is burnt, sulphur
dioxide is formed by oxidation of hydrogen sulphide.
Sulphur dioxide leads to the formation of acid rain which results in damage to plants and
aquatic organisms (refer to the effects of air pollution, chapter 12). Hydrogen sulphide reacts
with many metals. The tarnishing of silver objects is due to the reaction with hydrogen
sulphide to form silver sulphide, which is black. Paints which contain lead compounds are
also discoloured, due to the formation of black lead(II) sulphide. Hydrogen sulphide and
ammonia, which are contained in bio gas, cause eye irritation.
29.6.5. Alkenes
The alkenes are members of a homologous series of general molecular formula, CnH2n,
where 2. They are characterized by a carbon-carbon double bond as their functional group
and therefore alkenes are unsaturated compounds. An unsaturated compound is one in
which some atoms do not exert all their combining powers with other atoms. For example in
alkenes, not all the carbon atoms are bonded to four other atoms as illustrated below.
Nomenclature
In accordance with I.U.P.A.C. system, alkenes are named by dropping the ending ane from
the names of the corresponding alkenes and replacing it with the suffix ene for example
ethane (C2H4), propene (C2H6), butane (C4H8), pentene (C5H10) etc.
29.7. Ethene
When the reaction is complete, the junction between the flask and the bottle should be
disconnected to avoid the possibility of sodium hydroxide solution sucking back into hot
concentrated sulphuric acid.
Properties of ethene
Ethene as the first alkene may be used to indicate some of the physical and chemical
properties of alkenes.
29.7.1.1. Physical properties
(i) Ethene is a colourless, sweet smelling and non-poisonous gas.
2. Addition reactions
Alkenes are reactive compounds because the double bonds are readily converted to single bonds
by addition of other atoms. So ethene undergoes addition reactions. An addition reaction is a
reaction I which a molecule adds to an unsaturated molecule by breaking a double or a triple
bond.
1,2-dibromoethane
Addition of bromine across the double bond takes plae readily in the presence of an organic
solvent such as tetrachloro methane (CCl 4) or ether. The solvent dissolves the halogen to form a
solution such that when ethene is bubbled through the solution, the reaction takes place more
efficiently.
Since hydrobromic acid is a volatile liquid (can easily vaporize), hypobromous acid remains in
the solution to participate in the reaction with the alkene.
2-bromoethanol
Therefore, when ethene is bubbled through bromine water, the red colour of bromine water is
discharged.
This reaction is applied in changing double bonds in vegetable oils into single bonds. For
example in margarine production. The hardening of liquid vegetable oils into solid fats is called
addition hydrogenation. The oil is heated and mixed with a finely divided nickel catalyst, and
then hydrogen is blown through the mixture under pressure.
(c) Ethene decolourises acidified or alkaline potassium permanganate solution (from
purple to colourless). Also this reaction is a characteristic test for an unsaturated
compound.
Exercise
1. The compound with the molecular formula C10H22 can undergo the following reaction:
C10H22 C8H18 + W
(f) Write the equation for the reaction that took place in e(i) and (ii).
(g) Name the product formed in e(i)
2. Two hydrocarbon compounds are represented by the molecular formulae, C3H6 and C3H8
(a) To which hydrocarbon series does each of them belong?
(b) Give the name of each compound.
(c) Describe any chemical test that can be used to distinguish between the two
compounds.
(d) Which one of the compounds is unsaturated?
(e) The unsaturated compounds named in (d) was reacted with hydrogen under certain
conditions.
(i) State the conditions necessary for the reaction to take place.
(iii) State one industrial application of the type of reaction named in (ii).
29.7.3. Alkynes
Alkynes are a homologous series of unsaturated hydrocarbons of the general formula, C nH2n-
2, where n 2. They contain carbon-carbon triple bond ( C C ) as the functional group.
Nomenclature
The alkynes are named by dropping the ending ane of the corresponding alkane and
replacing it with the suffix yne. The table 9.3 shows the molecular formulae and names of
the first five alkynes.
Ethyne C2H2
Propyne C3H4
Butyne C4H6
Pentyne C5H8
Hexyne C6H10
2. Addition reactions
The triple bond in ethyne contributes a lot to its chemical properties. Ethyne undergoes
addition reactions with the substances that react with alkenes except that ethyne reacts more
slowly than alkenes. For example, it takes ethyne sometime to decolourize bromine unlike
ethene which does it almost instantaneously.
(i) Ethyne decolorizes bromine. The triple bond is first converted into a double bond, then
into a single bond.
1,2-dibromoethene
1,1,2,2-tetrabromoethane
(ii) A mixture of ethyne and hydrogen when passed over a nickel catalyst at about 150 oC
produces ethene.
HC CH + H2 CH2=CH2
CH2=CH2 + H2 CH3CH3
Exercise
Nomenclature
Alcohols are systematically named as alkanols, that is, the name is a particular member is
obtained by dropping the ending e of the corresponding alkane and replacing it with the
suffix o1. The table 9.4 shows the molecular formulae and name of the first five alkanols.
1 CH3OH Methanol
2 C2H5OH Ethanol
3 C3H7OH Propanol
4 C4H9OH Butanol
5 C5H11OH Pentanol
starch maltose
Yeast is then added at room temperature and one of its enzymes, maltase, catalyses the
hydrolysis of maltose to glucose.
glucose
Another enzyme present in yeast, zymase, catalyses the decomposition of glucose to ethanol
and carbon dioxide.
The resulting solution is crude ethanol which is converted to pure ethanol by fractional
distillation.
In Uganda, locally crude ethanol (known as Tonto) is obtained from bananas. Juice is
extracted from ripe bananas by squeezing them using spear grass leaves or banana leaves.
The juice is filtered to remove any solid impurities. The filtrate (juice) is then poured into a
locally made wooden container where it is mixed with ground roasted sorghum. The
container is covered and the mixture is allowed to ferment for two days. The resulting
solution is crude ethanol, locally known as Tonto.
Alternatively, millet is ground to flour and the flour is mixed with water to form a paste
which is covered in a container or buried in the ground for a few days so that it can ferment.
The fermented paste is removed and roasted to obtain malt. Miller grains are soaked in water
for sometime and allowed to germinate. It is then dried to give yeast. Yeast is added to malt
in appropriate proportion. A carefully determined amount of water is then added to the
mixture of yeast and malt to form a liquid mixture. The mixture is covered and stored in a
warm place for about 3 to 4 days so that an alcoholic drink locally called Malwa is formed.
Ethanol burns in air with a blue flame to give carbon dioxide and water. The reaction is
exothermic.
When excess concentrated sulphuric acid is added to ethanol and the mixture heated to
180oC, ethanol is dehydrated to ethene.
C2H5OH CH3COOH
ethanoic acid
C2H5OH CH3COOH
Exercise
Carboxylic acids are compounds of the homologous series of the general molecular formula
CnH2n+1COOH. In this series, CnH2n+1 represents the alkyl (R) groups. So carboxylic acids can
be written as RCOOH. The functional group of members in this series is the carboxyl group
(-COOH). They are sometimes called organic acids.
Nomenclature
Carboxylic acids are named as derivatives of alkanes by dropping the ending e of the
corresponding alkane and replacing it with the suffix oic and the functional group is always
at the end of the chain. The first three members in the series are methanoic acid (HCOOH),
ethanoic acid (CH3COOH) and propanoic cid (CH3CH2COOH).
29.9.1. Properties of carboxylic acids
1. Acidic character
When carboxylic acids dissolve in water, the solution formed turns blue litmus paper red.
When carboxylic acids dissolve in water, they are slightly ionized, that is they are weak
acids.
(i) Carboxylic acid react with strongly electropositive metals liberating hydrogen gas.
sodium ethanoate
(ii) Carboxylic acids react with bases to form salt and water.
zinc ethanoate
(iii) Carboxylic acid reacts with carbonates and hydrogencarbonates liberating carbon
dioxide.
2. Formation of esters
When a carboxylic acid reacts with an alcohol, a sweet smelling compound called ester is
formed.
Ethylethanoate (ester)
29.10. Esters
Esters have a formula, RCOOR1, where R and R1 are alkyl groups. Long-chain carboxylic
acids are very often referred to as fatty acids because their chief source is from esters in
animal fats and vegetable oils. Fats and oils are esters occurring naturally in plants and
animals. Esters produced from saturated fatty acids are usually solids at room temperature
and are called fats. Esters obtained from unsaturated fatty acids are usually liquids at room
temperature and are called oils.
When a mixture of sodium hydroxide or potassium hydroxide solution and an ester is heated,
the ester is hydrolysed to form sodium salt. The alkaline hydrolysis of any ester is called
saponification.
29.11. Soap
Soap is a sodium or potassium salt of a long-chain carboxylic acid, that is, potassium stearate
(C17H35COOK) and sodium stearate (C17H35COONa).
The soap is precipitated by addition of concentrated sodium chloride solution. This process is
called salting out of soap. Sodium chloride lowers the solubility of soap in the mixture and
causes the precipitation of soap which floats on top of the liquid. It is then removed and
compressed into a continuous block which is cut into bars.
Hard soap consists of sodium salts of carboxylic acids. Soft soap consists of potassium salts
for example potassium stearate. Liquid soap is a mixture of soft soap and coconut oil whereas
toilet soaps are produced by adding dyes and perfumes to the purified soap.
29.11.2. The cleansing action of soap
When soap is added to a cloth in water, the polar end containing COO -Na+ group dissolves in
water and non-polar organic end containing the hydrocarbon chain dissolves in non-polar
grease or oil deposits in a cloth. Agitation causes emulsification of grease or oil and then
dirty tiny droplets are carried away by clean water. If water is hard, soap first reacts with
dissolved calcium and magnesium ions to form insoluble salts which are precipitated as
scum.
Soapless detergents are more effective than soap in hard water since they do not form a scum.
This is because the calcium and magnesium salts of the hydrocarbon sulphonic acids of
which the detergents are composed, are soluble in water, so there is no precipitate (scum)
formed.
In the laboratory, soapless detergents can be prepared by boiling a vegetable oil with
concentrated sulphuric acid. The hydrogen sulphate compound formed is then neutralized by
adding sodim hydroxide solution. A precipitate forms and on evaporation, a white solid
soapless detergent is formed.
Soapless detergents can be manufactured in solid form (for example washing powders) or in
liquid form (for example washing-up liquids and shampoos). Washing powders contain a
number of other components. Phosphates are added to prevent scum formation. Sodium
perborate gives the washing powder a mild bleaching action. Sodium sulphate and silicate
help to keep the powder dry and free flowing. Some powders also contain enzymes to digest
organic dirt like food stains and blood. Common detergents include Omo, Noimi, Fab and
Axion.
29.11.4. Advantages of synthetic detergents over soap
(i) Synthetic detergents are more soluble in water than soap.
(ii) Synthetic detergents dont form scum with hard water unlike soap.
29.12. Polymerisation
Polymerization is the process by which many small molecules are combined to form a single
complex molecule. The small molecules that come together are called monomers and the
complex molecule formed is called a polymer. Polymers are long-chain molecules with
repeated units produced by the process of polymerization. There are two types of
polymerization, that is, addition polymerization and condensation polymerization.
29.14. Polyethene
It is a polymer formed by addition polymerization. Polyethene contains ethene as the
monomer. When many (n) ethene molecules combine, they form the polymer, polythene.
nCH2 = CH2 CH2 CH2 where n = about 600
There are two common types of polyethene depending on the conditions provided during
their manufacture.
Here the reaction involves heating ethene at about 200 oC and pressure of 1200
atmospheres in the presence of a small amount of oxygen. The polymer formed under
these conditions is soft because these conditions prevent close packing of chains. This
type of polyethene is used in making film and sheet material for plastic bags and
wrapping polyethene. It is also used as a film in solar heaters and driers instead of glass.
This is produced through the use of zieglar catalyst at a temperature of 60 oC and low
pressure of 1 atmoshpere. The polymer formed has fewer branched chains that are closely
packed. This is used for making moulds for rigid articles such as crates for milk and beer
bottles, toys, water pipes and electric cable pipes.
29.15. Polypropene
The monomer for this polymer is propene. Propene polymerises in presence of a Ziegler
catalyst to form a polymer, polypropene.
It is used to make beer bottle crates and ropes such as those used for drying clothes.
Polychloroethene is used to make water pipes, gramophone records and light fittings such as
sockets, plugs and bulb-holders.
29.17. Synthetic rubber
Synthetic rubber is made up of two monomers, that is, but-1,3-diene and phenylethene.
Therefore synthetic rubber is a co-polymer because it is made up of two different monomers.
Polymerization occurs in presence of a peroxide.
C6H5 C6H5
The elasticity of rubber is caused by the coiling of the rubber molecules. When rubber is
stretched, the molecules straighten out and when it is released, the molecules coil up again.
glucose starch
1. Natural polymers
These are polymers that exist in nature and are mainly manufactured by plants. Table 9.5
shows some of the natural polymers, their monomers and the type of polymerization they
undergo.
Natural rubber
Natural rubber is got from latex which slowly extrudes from the bark of rubber trees
when it is cut. It is coagulated by addition of ethanoic acid
Vulcanization of rubber
Natural rubber in its raw form is soft and sticky when warm and therefore it is unsuitable
for most of the intended uses. It can be made stronger, harder and more durable by
heating it with sulphur at about 140oC. The sulphur atoms are added across the carbon-
carbon double bonds to form cross-linkages between the polymeric chains of rubber. This
process is called vulcanization of rubber. Table 9.6 shows the properties of raw rubber
compared with vulcanized rubber. Vulcanized rubber is used for making toys, tyres, tyre
inner tubes and foot wear such as soles and gum boots.
2. Synthetic polymers
These are man-made polymers for example polyethene, nylon, terylene etc. All synthetic
polymers are plastic in nature hence are plastics. Plastics are man-made materials
composed of giant molecules based on carbon atoms. Plastics are classified into groups
according to the changes which occur on heating.
These are plastics which become soft and mouldable on heating without undergoing
any significant changes and on cooling they harden, for example polyethene, nylon
and polyvinyl chloride. There are no cross links between the chains in thermoplastics.
On heating, the chains move freely over each other thus the plastic melts.
These are plastics which decompose on heating and cannot be reshaped after
manufacture. They are rigid, hard and brittle, for example Bakelite. In thermosetting
plastics, there are strong cross links between the chains which gives a rigid structure.
Heating has no effect until a temperature high enough to break some of the cross links
is reached. Thus the plastic decomposes. There are cross-linking agents that convert
thermoplastics into thermosetting plastics, for example sulphur in vulcanization of
rubber.
3. Semi-synthetic polymers
Semi-synthetic polymers exist. Rayon is a semi-synthetic polymer because it is made
chemically from cellusoe in form of wood pulp. The wood pulp is treated with sodium
hydroxide solution and carbon disulphide. This converts the cellulose into a syrup
substance called viscose, which is then forced through small holes in a metal plate into a
bath of dilute sulphuric acid. The acid converts the viscose solution into glossy
transparent filaments which can be twisted to form rayon threads.
Exercise
1. Under certain conditions ethene undergoes a reaction that can be represented by the
following equation.
nCH2 = CH2 CH2 CH2
(a) Write the structural formula for the monomer of the polymer.
(b) Name the
(i) Monomer.
(ii) Polymer.
4. (a) In the manufacture of soap, oil or fat is heated with sodium hydroxide solution.
(i) Name the process of making soap.
(b) Sometimes when soap is used for washing clothes, a scum is formed.
(i) What is a scum?
(iii) Give the general name given to water which forms scum with soap.
(iv) Describe a chemical method by which the type of water you named in b(iii) can
be treated to avoid formation of scum. Write equations for the reactions that are
involved.
(a) (i) Name one soapless detergent that can be used instead of soap.
(ii) What is the advantage of using soapless detergents rather than soap?
(ii) Write equation for the reaction that takes place during fermentation.
(b) Write equation to show how ethanol can be converted to ehtene and indicate the
conditions for the reaction.
(c) (i) State what would be observed when ethene is reacted with bromine
7. (a) State the differences between fats and oils. Give one example of each.
(b) Briefly describe how soap can be prepared.
(c) State what would be observed if soap solution was shaken with a solution containing
magnesium hydrogencarbonate.
(e) State what would be observed if a solution of soapless detergent was used instead of
soap solution.
(b) Write the equation for the reaction that leads to the formation of ethanol.
(c) Briefly describe how ethanol produced can be concentrated.
11. (a) (i) What is a polymer?
(ii) Distinguish between a natural and artificial polymer. In each case give two
examples.
12. Ethanol can be converted to substances P and Q according to the reaction scheme shown
below.
C2H5OH C2H4 C2H6
step1 P step 2 Q
30.1. Introduction
Chemical changes are normally accompanied by energy changes. Energy is neither created
nor destroyed but it can be transformed from one form to another form. Chemical energy is
transformed into chemical energy in electrochemical cells. Fuels such as coal, store chemical
energy which is transformed into heat energy when it is burnt. During metabolism, the
chemical energy of food such as carbohydrates is converted to heat energy to keep the body
warm, to mechanical energy in muscles ans to electrical energy in the impulses within our
nerve fibres. The most common form of energy change in chemical reactions is the heat
change, and is our major concern.
30.2. Enthalpy
Enthalpy is the energy (heat) content of a substance which is stored in its bonds. Energy is
released when bonds are formed and to break bonds, energy must be supplied. The enthalpy
of a substance is denoted by H. Changes in enthalpy are denoted by H (delta H). Enthalpy
changes occur in a reaction when some old bonds in the reactants break and new bonds are
formed in the products.
Since H2 < H1, the H is negative. This can be illustrated using an energy level diagram
figure 10.1.
For example, when carbon reacts with oxygen, heat is evolved.
The chemical energy in carbon and oxygen is partly transferred to chemical energy in carbon
dioxide and partly evolved as heat. Thus carbon dioxide has less energy than the starting
materials, carbon and oxygen (figure 10.2). Therefore the value of enthalpy change is
negative.
An endothermic reaction is one during which heat is absorbed from the surroundings. When
an endothermic reaction occurs, the heat required for the reaction is taken from the reacting
materials and the temperature of the products falls below the initial temperature. Eventually,
the temperature of the products raises to room temperature again as heat is absorbed from the
surroundings. In this case, the heat content of the products is greater than that of the reactants
and the enthalpy change is positive.
H = H2 H1
Since H2> H1, then H is positive. Figure 10.3 illustrates an energy level diagram for an
endothermic reaction. The units for enthalpy change are kilo joules per mole (kJ mol-1)
For example, when hydrogen reacts with iodine, heat is absorbed from the surroundings.
Exercise
How much heat is given out when 20 g of carbon are completely burnt?
Solution:
12
12
= 655 kJ
A thin walled tin can is filled with a known volume of water. Ethanol is added to a specimen
bottle and a wick is filled in a cork through the bottles mouth. The specimen bottle with its
contents is weighed and its mass is recorded. The apparatus is set up as shown in figure 10.5.
The initial temperature of the water is recorded and the wick is lighted to heat the water. The
water is stired carefully with the thermometer. When the thermometer shows a convenient
temperature rise, the flame is blown off and the highest temperature reached is recorded. The
bottle and its contents is reweighed after cooling. A tin can is a good conductor of heat and
transmits most of the heat directly to the water.
Specimen results:
= temperature rise.
m = 1 x 100 = 100 g
= 40 24 = 16oC
46
= 0.044 mol.
0.044
= 152727 J or 152.7 kJ
The value obtained in this experiment is not very accurate because of heat losse to
surroundings.
Exercise
(ii) 4 g methanol?
(b) When 11.5 g of methane are burnt in excess oxygen, 640 kJ of heat are produced.
Calculate the
(ii) Volume of methane at room temperature that will burn to produce 1560kJ.
3. When 23.6 g of butane were burnt, the heat produced raised the temperature of 50 g of
water from 30oC to 40oC.
(a) Write an equation for the complete combustion of butane.
(b) Calculate the heat of combustion of butane.
2. Enthalpy of solution
The enthalpy of solution is the enthalpy change that occurs when one mole of a substance
is dissolved in sufficient amount of water such that no further heat change occurs on
dilution. When one mole of sulphuric acid dissolves in water, heat is evolved. The
reaction is exothermic and therefore the enthalpy of solution is negative. Some reactions
are endothermic. For example, when sodium chloride dissolves in water, heat is absorbed
and therefore the enthalpy of solution is positive.
30.5. Determination of the enthalpy (heat) of solution o f sodium chloride
A known volume of water is palced in a plastic cup and its temperature recorded. A known
mass of sodium chloride is added to the water. Carefully water is stirred with a stirrer and the
lowest temperature of the solution is recorded.
Specimen results:
Initial temperature = 24 oC
Calculations:
Heat evolved =
Assuming the specific heat capacity of the solution is 4.2 J g-1 oC-1
58.5
= 0.53 mol
This value is less than the accurate valuw, +4.97, because of the experimental errors.
Exercise
1. The heat of solution of sulphuric acid is -70kJmol -1. Calculate the mass of sulphuric acid
that will evolve 350 kJ of heat when sulphuric acid is dissolved in water.
2. When 231 g of ammonium nitrate were dissolved in water, 75 kJ of heat were absorbed.
Calculate the heat of solution of ammonium nitrate.
The enthalpy of neutralization of any strong acid and strong alkali is -57.3kJ mol -1 and is
constant. This is because the acids and alkalis and the salt produced in the reaction are
completely ionized in solution and the net reaction is the formation of water molecules. The
enthalpy of neutralization of a weak acid or alkali is less than 57.3 kJ mol -1 and is not constant.
This is because weak acids or alkalis are partly ionized in aqueous solution. Some heat is
absorbed for complete ionization of the acid or alkalis in order for neutralization to occur. The
enthalpy (heat) of ionization affects the overall enthrapy change.
Specimen Results:
Calculations:
2 2
= 24.5 oC
= 2658.6 J or 2.66 kJ
1000
= 0.05 mol
0.05
Heat of neutralization is -53.2kJ/mol
Exercise
1. 80 cm3 of 1 M nitric acid and 80 cm3 of 1 M sodium hydroxide, both at 25 oC were mixed
in a plastic beaker. The mixture was stirred and its maximum temperature was 31.34 oC.
(specific heat capacity of the solution = 4.2J/g/oC, density of the solution = 1 g/cm3)
(a) Write the ionic equation for the reaction which took place.
(b) Calculate the
(i) Number of moles contained in 80 cm3 of 1 M sodium hydroxide.
2. Aqueous hydrogen ions react with aqueous hydroxide ions according to the equation.
H+(aq) + OH-(aq) H2O(l) H = 57 kJ mol-1
Calculations:
Heat evolved = mc
1000
= 0.025 mol
0.025
= 29400 J or 29.4 kJ
1. The amount of ehat evolved when 16 g of copper was displaced from the solution by 2.4
g metal, Q, was 720 kJ. Calculate the heat of displacement. (Q = 24)
2. Iron reacts with copper(II) ions according to the equation.
Cu2+(aq) + Fe(s) Cu(s) + Fe2+(aq) H = -151 kJ mol-1
Calculate the mass of iron that will cause a heat change of -170 kJ.
3. (a) When methane burns in oxygen, heat is produced. Write an equation for the
combustion of methane in excess oxygen.
(b) The heat of combustion of methane is -890 kJ mol -1. calculate the volume of methane
gas at s.t.p that when burned in excess oxygen would raise the temperature of 178 g of
water by 10 oC.
4. (a) Bio gas contains mainly methane. Name two raw materials that can be used to
produce biogas.
(b) Methane burns in oxygen according to the equation:
Calculate the volume of methane at s.t.p. that will burn in excess oxygen to produce
2670kJ.
(c) When 0.382 g of ethanol was burnt, the heat evolved raised the temperature of 100 g
of water from 16.0 oC to 43.0 oC. Calculate the heat of combustion of ethanol.
(d) Name two products, other than water of incomplete combustion of ethanol.
7. (a) 50 cm3 of 2 M hydrochloric acid and 50 cm3 of 2 M sodium hydroxide, both at 22 oC,
were mixed in a plastic beaker. The mixture was stirred and its maximum temperature
was 35oC. (specific heat capacity of the solution = 4.2 J/g oC, density of the solution = 1
g/cm3)
(i) Write an ionic equation for the reaction which took place.
(ii) Calculate the heat of the reaction.
(b) 50 cm3 of 2 M ammonia solution was used instead of sodium hydroxide solution in
(a). State whether the heat of the reaction was greater than, smaller than or equal to the
value you have calculated in (a) (ii). Give a reason for your answer.
8. (a) Write an equation to show how ethanol can be prepared form glucose.
(b) State how a sample of ethanol obtained from the product of the reaction in (a) can be
purified.
(c) When 23 g of ethanol was completely burnt, 13600 J of heat was produced. Calculate
the molar heat of combustion of ethanol.
9. 7.5 g of methane, CH4 was completely burnt in air. Methane burns in air according to the
following equation:
CH4(g) + 2O2(g) CO2(g) + 2H2O(g) H = 890 kJ mol-1
Calculate the
10. When 6.4 g of zinc powder were added to 250 cm 3 of a 0.1 M copper(II) sulphate
solution in a plastic cup, 5.45 kJ of heat was liberated.
(a) Explain why a plastic cup was used instead of a metallic cup.
(b) Write an equation for the reaction between zinc powder and copper(II) sulphate.
(c) Calculate the
(i) Number of moles of zinc in 6.5 g of zinc powder.
(iii) Heat energy produced when 1 mole of zinc reacts with 1 mole of copper(II)
sulphate.
11. The formation of methanol from hydrogen and carbon monoxide is represented by the
equation.
2H2(g) + CO(g) CH3OH(g) H = -92 kJ/mol
What would be the energy released in kJ mol-1, when 3.2 g of methanol is formed?
12. An experiment was carried out to determine the molar heat of combustion of methanol. A
small lamp containing methanol was weighed and then lit. The heat produced by the
combustion of methanol was used to raise the temperature of 100 g of water in a metal
can (ignore the heat required to raise the temperature of the metal can). The spirit lamp
was weighed again after the experiment.
Results:
Mass of the spirit lamp + contents before heating = 36.17 g.
(a) What was the rise in temperature of water during the experiment?
(b) Calculate the amount of heat obtained by the water during the experiment.
(c) What mass of methanol was burnt during the experiment?
(d) Calculate the heat produced when
(i) 1 g of methanol was burnt.
13. The following pairs of compounds were reacted together and the maximium temperature
rise was recorded for each reaction.
A. 200 cm3 of 2 M sodium hydroxide and 200 cm3 of 2 M ethanoic acid.
B. 200 cm3 of 2 M ammonia solution and 200 cm3 of 2 M ethanoic acid.
C. 200 cm3 of 2 M sodium hydroxide and 200 cm3 of 2 M hydrochloric acid.
(a) State the pair which showed the
(i) Highest temperature.
The rate of a chemical reaction is the progress of the reaction in unit time. In other words, the
rate of a chemical reaction is the rate at which products are formed or the rate at which
reactants are used up in the reaction.
time in seconds
The determination of the rate of this reaction can be done by either measuring the volume of
hydrogen evolved with time or by measuring the time a given length of magnesium ribbon
takes to dissolve in varying concentrations of the acid.
To determine the rate of reaction at a given time, say t 1, the tangent to the curve is drawn at
that time as shown in figure 11.3. The gradient of the tangent is the rate of reaction at that
time, that is y/x. the units are cm3/s.
Exercise
In determination of the rate of reaction, 10 g of calcium carbonate were reacted with dilute
hydrochloric acid. The mass of the flask and its contents was weighed with time.
(a) Write the equation for the reaction that took place.
(b) Sketch a graph of mass of flask and its contents against time.
Factors which affect the rate of a chemical reaction are concentration, temperature, surface
are (particle size), pressure, catalyst and light. You are required to perform various
experiments to investigate the effect of these factors on the rate of chemical reactions.
The rate of the reaction depends on the frequency with which reacting particles collide,
which frequency depends on the concentration of the reactants. The higher the
concentration, the higher the frequency of collision and therefore the higher the rate of
the chemical reaction.
Experiment:
Make a mark with blue or black ink on a piece of paper. Place 50 cm 3 of 0.05 M sodium
thiosulphate solution into a beaker. Add 10 cm 3 of 1 M hydrochloric acid to the sodium
thiosulphate and at the same time start the stop clock. Gently shake the mixture to mix
the solution well and place the beaker on the paper over the mark. Watch the mark
through the solution from above the beaker. Stop the clock when the mark just
disappears. Vary the concentration of the thiosulphate solution by taking 40, 30, 20 and
10 cm3 each time by adding distilled water. Tabulate your results including 1/time. Plot
graphys of volume of sodium thiodulphate solution against 1/time (time -1) and against
time. The rate of reaction is proportional to the reciprocal of time (time -1). Your graphs
should appear as shown in figure 11.4a and 11.4b.
The mark disappears becaue the reaction between hydrochloric acid and sodium thiosulphate
forms a precipitate of sulphur which renders the mixture opaque.
Figure 11.4a shows that the higher the volume of the sodium thiosilphate, the less the time
taken to form a precipitate. Figure 11.4b shows that the rate of the reaction increases with
increase in volume of sodium thiosulphate solution.
Exercise
(c) On the same axis, sketch a graph of volume of hydrogen evolved against time when the
following are reacted with the same concentration and volume of sulphuric acid.
(i) 10 g of magnesium.
(ii) 50 g of magnesium.
Exercise
(ii) 40 oC.
Exercise
(b) On the same axis, sketch a graph of volume of oxygen evolved against time when
Solid react much more rapidly when powdered than when in large lumps. This is because
reactions with solids take place at the surface. Powdered solids present a large surface
area over which the reaction occurs than solids in lump form.
Exercise
1. Calcium carbonate lumps were mixed with dilute nitric acid in a conical flask. The mass
of the flask and its contents was weighed with time.
(a) Write the equation for the reaction that took place.
(b) (i) Sketch a graph of mass of flask and its contents against time. Label the graph A.
(ii) On the same axis, sketch the graph that would be obtained when powdered
calcium carbonate is used instead of calcium carbonate lumps. Label the graph B.
2. (a) Which one of the following reaction mixtures will produce hydrogen more quickly at
room temperature?
(i) Magnesium ribbon + dilute sulphuric acid.
(c) Suggest two other methods by which the rate of this reaction can be altered.
Add 1 cm3 of sodium chloride solution to two test-tubes. To each test-tube add a few drops of
silver nitrate solution. Immediately, a white precipitate forms. Put one test-tube in a dark cup
board and the other in sunlight for about 4 minutes. Record your observations.
Sodium chloride solution forms a white precipitate with silver nitrate solution according to
the equation.
In presence of light, the precipitate darkens because of the decomposition of silver chloride
to solver and chlorine. In absence of light, the precipitate remains white.
The effects of light on hydrogen peroxide and concentrated nitric acid explain why they are
stored in dark-glass bottles.
31.1.7. Effect of pressure on the rate of reaction
A change in pressure only affects reactions which occur in the gas phase. When pressure
of a gaseous mixture is increased, the gases are compressed. This brings the reacting
particles together and thus increases the frequency at which the reacting particles together
and thus increases the frequency at which the reacting particles collide hence increased
rate of reaction.
32. Equilibrium
Equilibrium is the point in a reversible reaction when the rate at which the reactants are
forming the products is equal to the rate at which the products are dissociating to the
reactants. Therefore, at equilibrium, both the products and the reactants are present. A
reversible reaction is one which proceeds in both directions, that is, forward and backward.
The factors that affect equilibrium are temperature, pressure, concentration and catalyst. The
effect of these factors on equilibrium was first investigated by Louise Le Chatelier who came
up with a principle known as Le chateliers principle. The principle states that when a
chemical equilibrium reaction is distributed externally by a change in one of the factors upon
which it depends, the equilibrium shifts in a direction so as to reduce the effects of the
change.
(a) Temperature
(b) Pressure
In a gaseous system, an increase in pressure leads to a decrease in the volume of the gases
involved and the reverse is true. Let us again consider the Haber process.
One volume of nitrogen combines with three volumes of hydrogen to produce two
volumes of ammonia. The forward reaction occurs with a decrease in volume from four
to two volumes. If additional pressure is applied to the system, the equilibrium shifts in
the direction of a reduction in volume, that is, the forward reaction is favoured and more
ammonia is produced. If pressure of the system is decreased, the equilibrium shifts in the
direction of an increase in volume, that is, the backward reaction occurs and more of the
reactants are produced. Gaseous equilibrium reactions which are not accompanied by a
change in volume, are not affected by pressure changes e.g
Exercise
(c) Concentration
If extra oxygen is pumped into the reaction vessel, the equilibrium shifts in the direction
that results in a decrease in oxygen concentration, that is, the forward reaction occurs and
more nitrogen monoxide is produced. If the concentration of one of the reactants
decreases instead of increasing, the equilibrium will shift to cancel this decrease and the
backward reaction will occur to restore the balance. If more of nitrogen monoxide is
added to the equilibrium mixture, the backward reaction will occur producing more
reactants, that is, the equilibrium shifts to the left in order to offset the effect of the
increase in concentration of nitrogen monoxide. If there is a decrease in the concentration
of nitrogen monoxide, the forward reaction is favoured and the equilibrium shifts to the
right producing more nitrogen monoxide.
(d) Catalyst
Catalysts do not have any effect on the position of the equilibrium. In an equilibrium
reaction, a catalyst increases the rate of both the forward and backward reactions, that is,
a catalyst enables an equilibrium to be attained much more quickly than when there is no
catalyst.
This idea of equilibrium is applied in some industrial process. In the Haber process,
ammonia is synthesized from nitrogen and hydrogen according to the equation.
Ammonia is produced with a decrease in volume and therefore high pressure will
increase the yield of ammonia. The reaction is exothermic therefore low temperature will
favour the production of ammonia. However, by lowering the temperature, the rate of the
reaction is reduced. The presence of a catalyst will give a sufficient reaction rate despite
the relatively low temperature. In general, a maximum yield of ammonia is obtained by
using the following conditions.
Aluminium oxide is added to make the catalyst more porous hence promoting its
effectiveness.
Exercise
1. The graph below shows the effect of temperature on the rate the reaction between
calcium carbonate of the same mass and excess 2 M hydrochloric acid.
(a) If curve B is for the reaction at 40 oC, which curve shows the reaction taking place at
(i) 20 oC
(ii) 60 oC
(b) Explain why the curves eventually end at the same level.
(c) State one other method that can be used to measure the rate of the reaction between
calcium carbonate and hydrochloric acid.
2. A certain mass of zinc powder was reacted with dilute hydrochloric acid at room
temperature.
(a) (i) Write an equation for the reaction.
(ii) Draw a graph to show how the volume of the gaseous product varied with time.
(ii) Using the same mass of zinc granules instead of the zinc powder?
(c) (i) Give a brief explanation of the cause of difference in the graphs G1 and G2.
(ii) Name one other factor that can cause similar results as in b(i) above.
4. (a) 12 g of large pieces of calcium carbonate were reacted with 50 cm 3 of 2 M
hydrochloric acid at room temperature. The decrease in mass was measured at regular
intervals.
(i) Write an equation for the reaction.
(b) State what would be observed if the same mass of calcium carbonate powder was
used instead of the large pieces. Give a reason for your answer.
(c) State what would be observed if the same mass of large pieces of calcium carbonate
was used at 40 oC. Give a reason for your answer.
5. The figure 11.11 shows the set up of the experiment used to study the rate of evolution of
a gas when 1.0g of powdered calcium carbonate was reacted with 50cm3 of 2 M
hydrochloric acid at 25 oC.
(a) Sketch a graph to show the variation of the volume of the gas evolved in the reaction
with time. Describe the shape of the graph.
(b) On the same diagram in (a) above sketch a graph to show the results obtained when
(i) 1.0 g of powdered calcium carbonate was reacted with 100 cm 3 of 1 M
hydrochloric acid at 25 oC.
(ii) 1.0 g of powdered calcium carbonate was reacted with 50 cm3 of 2 M
hydrochloric acid at 25 oC.
(c) Explain the shapes of the graphs you have sketched in (b) (i) and (ii) above.
(d) 1.0 g of powdered calcium carbonate was reacted with 20 cm 3 of 2 M hydrochloric
acid. Which one of the reactants was in excess?
6. When a certain volume of 0.1 M hydrochloric acid was reacted at room temperature with
excess of iron filings, a gas was produced.
(a) Draw a labeled diagram to show how the rate of reaction was determined.
(b) Write equation for the reaction that took place.
(c) Draw a sketch graph of the volume of the gas evolved against time.
(d) State how the rate of reaction would change if the reactions was carried out at a
temperature above room temperature.
TRIAL QUESTIONS FOR PAPER 1
- Choose the correct answer from the alternatives A, B, C and D and write it
on the answer sheet provided.
Molar gas volume at s.t.p = 22.4 dm3 and at room temperature is 24 dm3.
1 Faraday = 96500
1. How many electrons are contained in the ion of atom with atomic number
15?
A. 10 B. 15 C. 18 D. 13
96 96
C. 2 x 22.4 x 96 dm3 D. 2 x 17 x 96 dm3
6.5 6.5
10. Which one of the following salt will not give oxygen when heated?
A. Potassium chlorate (V) B. Copper (II) nitrate
11. Substance P was separately treated in two test tubes using sodium
hydroxide and ammonia solution. In both cases it formed precipitate
soluble in excess. The cation in P is
A, Mg2+ B. Zn2+ C. Al3+ D. Pb2+
3 100
15. Strontium carbonate (SrCO3) relative formular mass = 148 reacts with
hydrochloric acid according to the equation.
2 x 0.4 2 x 16.6
16. The mass in grams of potassium hydroxide used to prepare 500 cm3 of
0.2M solution for volumetric analysis is (K = 39 O = 16 H = 1)
A. 56 B. 5.6 C. 1.20 D. 11.2
17. Which of the following pairs of oxide will react with both dilute sulphuric
acid and sodium hydroxide solution?
A. CuO and ZnO B. PbO and MgO
C. Al2O3 and CuO D. ZnO and Al2O3
18. Which one of the following salt dissolves in water to give an alkaline
solution?
A. 14 x 100 B. 54 x 100
149 149
C. 18 x 100 D. 42 x 100
149 149
21. Magnesium was burnt in air. When was water added to the cold products,
a gas was given off. The gas was
24. A hydro carbon Z when completely burnt in excess oxygen produce 220g
of carbon dioxide and 45 g of water. The empirical fomular of z in
A. CH B.. CH2 C CH3 D. C2H3
25. During sewage treatment process the stage at which sludge is formed is
A. filtration B. sedimentation
31. 30 cm3 of 2.6M hydrochloric acid completely reacts with 5.0 g of impure
Zinc carbonate. The percentage purity of zinc carbonate in the mixture is
(Zn = 65, C = 12, O = 16, H = 1, Cl = 35.5)
32. A bottle of sulphuric acid is labelled 10M H2SO4. This means in every one
litre of solution there are
33. A beaker containing lime water is left exposed to the atmosphere for some
time. The white solid that formed at the bottom of the beaker is
A. Calcium oxide B. Calcium hydroxide
34. The gas collected when chlorine water is exposed to sun light is
A. Chlorine B. Oxygen C. Hydrogen chloride D. Hydrogen
35. Which one of the following salts is prepared by double decomposition?
A. Sodium sulphate B. Lead (II) nitrate
37. 50 cm3 of a mixture of Ammonia and Carbon dioxide was passed over
heated copper (II) Oxide, and nitrogen formed occupied 20 cm3 under the
same condition as a mixture. The percentage of ammonia in the mixture is
A. 20% B. 80% C. 40% D. 60%
C. Na+ + e Na (s)
39. the mass of copper deposited from a solution of Copper (II) chloride when
a current of 1.2A is passed for 3000Sec (Cu = 63.5)
41. In each of the questions 41 45 there is assertion (left and reason (right).
The correct answer will be as follows:
A. if both assertion and reason are true statements, and reason is correct
explanation of assertion.
B. if both assertion and reason are true statements but reason is not correct
explanation of the assertion.
Instruction summarised
Assertion Reason
A True True and reason is a correct explanation
B True True but reason not a correct explanation.
C True Incorrect
D Incorrect True
INSTRUCTION SUMMARISED
A B C D
1,2,3 only 1,3, only 2,4 only 4 only Correct
Correct Correct Correct
46. Which one of the following is responsible for forming fur in kettles used
1. Calcium sulphate
2. Calcium carbonate
3. Magnesium sulphate
4. Magnesium carbonate.
47. When sulphur dioxide is passed through sodium hydroxide solution for a
long time, which one of the following product(s) are formed?
1. Sodium sulphate 2. Sodium sulphite
49. Which of the following solutions contains the same number of hydrogen
ions?
1. 1 l of 1M H2SO4 2. 1 l of 2M H Ol
2. Temperature
3. Surface area
4. Pressure.
NameIndex No
P545/1
Paper 1
1 hours.
Instructions:
The paper consists of fifty compulsory objective type of questions. For each of the
questions 1 50, four alternatives lettered A D are given. Indicate the letter which
corresponds to the best alternative.
A. Water
B. Ethanol
D. Ethanoic acid.
2. Which of the following would give off carbon dioxide acid rapidest? 2 molar
3. Ammonia gas was passed over a copper foil, the foil glowed red hot a
colourless gas which turned reddish brown on exposure to air. The role of
copper toil was to.
B. sunlight
C. hydrogen peroxide
D. bromine
8. Which of the following is the reason why zinc is not a typical transition metal?
It (does not)
A. vertical pillars of pure calcium carbonate growing from the floor and
root a cave.
11. 25.00cm3 of phosphoric acid (H3PO4) completely reacted with 18.7 cm3 of
0.5M sodium
25 x 0.5
3 x 18.7 x 0.5 2 x 18.7 x 0.5
18 2 25 3 25
A. B. C. D.
3 x 2 x 25 x 0.5
18.7
12. The atomic mass of chlorine is 35.5. The mass of the isotopes of chlorine is
35 and 37: Which of the following is correct?
35 37
A. Cl is more abundant than Cl
37 35
B. Cl is heavier than Cl
37
C. Cl is radioactive
35
D. Cl is lighter and passes through the mass spectrometer fast.
C. Ammonia D. Ethene
H H
H H
polyethene when n= 50 is
21. Calcium carbonate reacts with hydrochloric acid according to the equation.
Manufacture of ;
23. 20.00cm3 of 0.25 molar solution of an acid Hn X completely reacts with 35.00
cm3 of 0.5M sodium hydroxide.
The stoichiometry of the reaction between the acid and alkali is.
A. 1 : 3.5 B. 4 : 1 C. 1 : 1 D. 1 : 4
24. The formula of the ion formed when excess ammonia reaction with copper (II)
sulphate is.
A. decomposition. B. cracking
C. decarboxylation D. dehydration
The volume of the gaseous products formed when 50cm 3 of nitrogen is mixed
with 120cm3 of hydrogen is
A. 170cm3 B. 90cm3
C. 80cm3 D. 70cm3
30. Which of the following pairs of metals would make the most efficient chemical
cell?
31. Which of the following salts dissolves in water to make a solution of pH less
than 7 is
32. 5.70g of hydrated salt X was strongly heat and weighed to constant mass
2.12g of the anhydrous salt remained. The number of molecules of water of
crystallization is
34. The mass of nitrogen formed when 240 coulombs of electricity is used in
electrolysis of molten magnesium nitride [Mg=24, N=14, 1F= 96,500
coulombs]
A. 240 x 28 B. 240 x 74
96500 x 6 96500 x 2
C. 240 x 14 D. 6 x 240 x 28
6 x 96,500 2 x 96,500
A. HNO3(1) B. H2SO4(aq)
C. NaCl(aq) D. HCl(1)
36. Which of the following processes take place naturally?
A. +4 B. +6 C. +2 D. -6
38. Which of the following gases is given off in Crater Lake; (L. Nyamunuuka) W.
Uganda?
39. The bleaching action of chlorine is different from sulphur (IV) oxide in that
chlorine (forms)
A. C2 H 2 B. C4H10
C. C6H12 D. C2 H 6
1000x 10 1000 x 10
C. 20 x 1 x1 D. 1000 x 0.1
1000 x 10 2 20
A B C D
1, 2, 3 1, 3 2, 4 4 only
correct correct correct correct
41. The reagent(s) most suitable to distinguish between Pb 2+ and Al3+ is (are)
1. Potassium iodide
2. Sodium carbonate
4. Ammonium hydroxide.
42. Which of the following elements forms more than one oxide?
1. Nitrogen
2. Sulphur
3. Carbon
4. Calcium.
1. Carbon dioxide
3. Ammonia
4. Methane.
1. C2 H 2
2. C3 H 3
3. N2
4. C10H16
Summarized instructions.
A B C D
46.
47.
NameIndex No
P545/1
Paper 1
1 hours.
Instructions:
A. Sublimation B. Chromatography
C. Filtration D. Crystallization
A. 18 9
B. 11 17
C. 9 19
D. 20 9
1000 10 1000 10
C. 20 x 1 x D 1000 x 0.1
1000 10 20
11. The equation for the reaction between carbon dioxide and
lime water is
A. 2: 8: 3 B. 2: 8: 4 C. 2: 8: 7 D. 2: 8:
18: 8: 5
A. 240 x 64 B. 64 x 96,500
2 x 96,500 2 x 240
C. 240 x 2 x 64 D. 2 x 96,500
96,500 240 x 64
A. +1 B. +3 C. -1 D. -3.
W top of group E
X F
Y G
Z H
A. W+E B. Z+E C. Z + H D. W+
H
20.64 2.52
2.52 20.64
H H H H
A. Br C C Br B. Br C C H
H H Br H
H Br
C Br C = C Br D. Br C = C Br
H Br
37. The formula of the iron formed when excess ammonia reacts
with copper (II) sulphate solution is
A. [Cu(NH4)4]2- B. [CuNH4)3]2+
C. [Cu(NH3)4]2- D. [Cu(NH3)4]2+
C. Decomposition D. Synthesis
39. Which of the following metals would form the best pair for
construction of the chemical cell?
CH3 H CH3 H
C C C C
H H H H
SUMMARIZED INSTRUCTIONS.
A B C D
1,2,3 correct 1,3 correct 2,4 correct 4 only correct
1. Ethene
2. Chlorine
3. Sulphur dioxide
4. Hydrogen sulphide
42. When sodium hydroxide was added to separate test tubes
containing cations, there was no observable change. Which
of the following were the cations?
1. Ca2+
2 Mg2+
3. Al3+
4. NH4+
3. Carbon dioxide
4. Calcium oxide
Choose:
SUMMERISED INSTRUCTIONS.
Assertio A B C D
n
True True True False
Reason True and True not False True
correct correct
explanation
explanation
46. In the contact process sulphur BECAUSEthe oxide is lost in
mists dioxide is dissolved in 98% when dissolved
in water.
Water milky
green to brown.
during saponification.
50. A solution of hydrogen chloride BECAUSE Methyl
benzene is a in methyl benzene liberates
non-ionising solvent.
Cu 2 Mg 2 Ca 2 Fe 2 Zn 2 Cu 2 K NH +4
A. B.
14
6 C
2. Isotopic carbon twelve ( ) is used to
C. Catalyst D. Proximity
6. An ore is a (an)
H 3C C CH3
H
The name of the compound is
A. butane B. propane
A. Iron B. Bromine
C. Platinum D. Rhodium
A. 50cm3 B. 70cm3
C. 40cm3 D. 30cm3
11. The mass of copper deposited when 240 coulombs of electricity is used in
electrolysis of copper sulphate is (Cu=63.5, IF- 96,500 coulombs).
63.5x240 63.5 x 240 240 x63.5 x 2
96500 2 x96500 96500
A. B. C. D.
2 x96500
240 x63.5
12. A glass rod was dipped in a solution X and held in a stream of hydrogen
chloride gas. The drop on the end of the glass rod turned milky. The cation
present in the solution X is;
4
A. Ca2+ B. NH C. Ag+ D. Ba2+
13. The volume of 0.25 molar hydrochloric acid required to react with 20.00cm3
of 0.1 molar sodium carbonate is given by;
20 x0.1 20 x 0.25 2 x 20 x 0.25 2 x 20 x0.1
2 x0.25 2 x 0.1 0.1 0.25
A. B. C. D.
A. XY3 B. X3 Y2 C. XY D. X 3Y
15. Which of the following gases does not reduce copper (II) oxide to copper?
A. Pb(OH)2 B. Zn(OH)2
C. Fe(OH)3 D. Cu(OH)2
17. Hydrogen peroxide decomposes in the presence of the Fe3+ the results of
the experiment are known in the sketch Y. To obtain the sketch X one has
to.
heat
19. Carbon burns in air according to the equation (C(s) + O2(g) CO2(g)
H 390KJ. The heat produced when 48g of carbon completely burns in air
is;
[C=12, ]
20. Part of the reactivity series is given below Cu, Mg, C, K, Na, Al. The
elements are not in the correct order. Which of the following equations
represents a possible reaction?
heat
A. 2Na2O(s) + C 4Na(s) + CO2(g)
heat
B. 2MgO(s) + C 2Mg(s) + CO2(g)
heat
C. 2 Al + 3CuO Al2O3(s) +3Cu(s)
21. Which of the following acids will yield the least volume of carbon dioxide
when reacted with a fixed mass of marble?
A. Polyethene B. Rubber
C. Cellulose D. Bakelite
23. The diagram below is used to illustrate the solubility of a gas. The delivery
tube on the left is to.
A. Equalize pressure
B. prevent suck back
A. Carbon B. Nitrogen
C. Oxygen D. Fluorine
26. Which of the following is carried out during the extraction of sulphur using
the frash method?
C. super heated water and hot air are sent to the sulphur deposits
D. carbon disulphide and hot water are sent to the sulphur beds.
C. H2 S D. CH4
C. Gamma ( ) D. neutron ( )
A. explosives
B. restore breath
35. When 0.1 Foradays were passed through a solution containing a metal ion
X .
5.96g of the metal was formed. The charge on the ion X is 1 Foraday =
96,500 coulombs, X= 179.
A. + B. 2+ C. 3+ D. 3-
36. Which of the following will form a green precipitate when sodium carbonate
solution is added?
37. A student soaked a white material in sodium chlorate I solution (Jik). The
material became crumpled and turned cream (dirty white) instead of
becoming white
40. Sulphuric acid reacts with sodium hydroxide according to the equation.
Summarized instructions:
A B C D
1, 2, 3 correct 1, 3 correct 2, 4 correct 4 only correct
1. polyethene 2. bakelite
3. rubber 4. decane
42. Which of the following is (are) radioactive?
266 3
109 Une 1 H
1. 2.
18 15
8 O 7 N
3. 4.
1. CH2 2. C3 H6
3. C5H10 4. C4 H8
44. Which of the following would form a yellow precipitate when potassium
iodide solution is added?
1. Pb2+ 2. Al3+
3. Ag+ 4. Cu2+
Choose;
Summarized instructions:
A B C D
Assertion True True True False
Reason True, correct True not False true
explanation explanation
47. Aluminium reacts explosively with Because Two moles of aluminium react with 3
copper (II) oxide moles of copper
(II) oxide
48. Sodium vapour is used in street Because It burns with a yellow flame to
lights light the dark streets.
49. Fluoro carbons are pollutants Because They deplete the ozone layer which
prevents ultra violet radition direct
contact.
Signature
--------------------------
545/1
Chemistry
Paper 1
Time: 1 hours
INSTRUCTIONS TO CANDIDATES.
This paper is made up of 50 objective type questions. Answer all questions. You are
required to write the correct answer, A, B, C or D in the box provided on the right hand
side of each question.
Do not use pencil.
3. Phosphoric acid H3PO4 dissociates into two hydrogen ions, the 3rd hydrogen ion is not
ionisable. From this, conclusion can be made that phosphoric acid is;
4. A gasless dense than air, soluble in water and alkaline in nature is best collected by;
5. Which one of the following would cause a basic change in sodium hydroxide when
bubbled in excess?
6. Which one of the following is the reason why sodium extracted by Downs process is
kept under nitrogen?
7. Which one of the following is used to prepare a sample of soap in the laboratory?
8. In the manufacture of sugar from sugar cane, the syrup is passed through a sulphuriser.
The role of the sulpuriser is to;
9. Which one of the following is observed when two drops of phenolphthalein indicator is
added to dilute hydrochloric acid>?
10. Which one of the following is observed at the anode when molten lead (II) bromide is
electrolysed?
12. Which one of the following is observed when at the anode when copper (II) chromate is
electrolysed?(Cu2+ is blue,CrO42- is yellow)
13. The loss in mass when Iron (II) sulphate -7-water is strongly heated is heated is due to;
14. Which one of the following is the reason why tree are planted in cities ?
C. provide CO2 for photo synthesis D. provide shelter for people and animals.
16. Which one of the following salt would have PH greater than 7?
17. In industrial areas, wall painted with lead compounds blacken. The ions responsible for
tarnishing the walls is;
A. S2- B. Cl- C. C4- D. O2-
18. The mass in grams of OH ions in 0.25M NaOH solution is (H=1, O=6)
0.25 4
0.25 17 0.25
0.25 17 17
A. 17 B. C. D. 4
19. The minimum volume of 1M HCl(aq) required to produce 0.25g of Hydrogen with excess
Magnesium is
22. The Fountain experiment is used to determine the solubility of gases except
23. Iron (III) chloride reacts with Sodium hydroxide according to the equation
The mass in grams of iron (II) hydroxide precipitated when 6cm3 of 0.1M sodium
hydroxide reacts with iron (II) chloride is;
[Fe(OH)3 =107]
24. 20.00cm3 of 0.1M sodium carbonate reacted with 25.00cm3 of XM phosphoric acid. The
morality of the acid is;
27. Which one of the following is used to determine the formular of water?
A. K B. L C. N D. M
29. Which one of the following is the reason why sulphur IV oxide is used in fruit juices?
It is (has)
32. When 40g of the oxide of element X was reduced,32.00g of the element was obtained.
The simplest formula of the oxide is; (X=65, O=16)
33. Which one of the following contain the same number of particles as 8.00g of oxygen?
37. An atom Y has 20 nuetrons and mass number 37,the electronic configuration of X is
38. Which one of the following hydrocarbons contains 20% by mass of hydrogen.
39. Which one of the following would bring about the change Fe2+ Fe3+ ?
In each of the following questions 41-45 one or more of the statement(s) is ( are) correct
Indicate the letter which corresponds to the best alternative according to the instructions.
Choose;
D If statements 4 is correct
A B C D
42. Which one of the following forms a black residue on strong heating?
44. The property of sulphur which is used in the Frasch process is (are);
In each of the questions 46-50, there is a statement on the left hand side (assertion) and
reason on the right hand side. Read the statements carefully then, choose the letter
which corresponds to the best alternative. The instructions are as in the table below;
A B C D
electrolysis
Atoms
48. Concentrated sulphuric acid is not used to BECAUSE the acid has low affinity for
water
dry ammonia
to hydrogen
Name.Centre/Index No/
Signature..
545/1
Chemistry
Paper 1
July/August, 2010
1 hours
Instructions
This paper consists of 50 objective type questions
Answer the questions by writing the correct alternative in the box on the right hand side
of the paper
1. A dilute aqueous solution og potassium iodide was electrolysed. The product at
the cathode is likely to be
A. iodine B. potassium
C. oxygen D. hydrogen
11. Concentrated nitric acid and concentrated sulphuric acid may be separated by their
action on
A) copper (II) oxide
B) litmus paper
C) iorn (II) sulpahte
D) sodium hydroxide solution
12. The element R reacts with the element T to form a compound R 2T3. The ion formed
by T is
A)T2+
B) T3+
C) T3-
D) T2-
13. When 2.02g of potassium nitrate was heated strongly 1.70g of a solid remained.
2KNO3 (s) 2KNO2 (s) + O2 (g)
What volume of oxygen measured at room temperature was evolved?
A) carbon dioxide
B) calcium oxide
C) steam
D) copper (II) oxide
14. Which one of the following substances does not react with red hot carbon?
A)carbon dioxide
B) calcium oxide
C) steam
D) copper (II) oxide
15. The yield of ammonia in the reaction;
N2 (g) + 3H2 (g) 2NH3 (g) + heat
A) raising the temperature
B) increasing the pressure
C) employing a suitable catalyst
D) adding an inert gas
16. Which one of the following is NOT a property of ethene
A) it turns potassium manganate (VII) colourless
B) it is an unsaturated hydrocarbon
C) it is a saturated hydrocarbon
D) it decolourises bromine water
17. Which one of the following combinations would produce oxygen at the fastest rate?
A) 100cm3 of 2mH2O2 heated at 30C
B) a mixture of 1.0g of MnO2 and 100cm3 of 2mH2O2 at room temperature
C) 100cm3 of 1MH2O2 heated at 30C
D) a mixture of 0.5g of MnO2 and 100cm3 of 2MH2O2 heated at 30C
18. Which one of the following is not an application of electrolysis?
A) synthesis of elements
B) purification of metals
C) anodizing aluminium
D) electroplating
19. The atomic number of magnesium is 12 and its relative atomic mass is 24. Which
one of the following represents the magnesium ion, Mg 2+
Protons Neutrons Electrons
A) 12 12 12
B) 12 10 12
C) 10 12 10
D) 12 12 10
20. When steam was passed over heated magnesium, the mass of the dry residue was
0.4g more
than that of the magnesium. How many moles of steam molecules were decomposed?
(H=1 O=16)
A) 0.40
B) 0.25
C) 0.025
D) 0.022
21. Most carbonates decompose readily on heating to form
A) an oxide and carbon dioxide
B) a metal and carbon dioxide
C) an oxide and oxygen
D) an oxide and carbon
22. A piece of metal foil was wrapped around an iron nail which was then put in a test
tube containing water. The foil that would best protect the iron nail form corrosion is;
A) lead
B) zinc
C) copper
D) tin
23. In which one of the following pairs would the substances be likely to react together if
their powders were mixed and heated?
A) carbon and lead (II) oxide
B) zinc and aluminium oxide
C) copper and lead (II) oxide
D) iron and magnesium oxide
24. Which one of the following is the best reason why zinc is able to displace copper
from solutions of copper salts?
A) zinc is more electronegative than copper
B) zinc loses electrons more easily than copper
C) zinc is a stronger oxidizing agent than copper
D) zinc has few electrons than copper
25. Methanol burns in oxygen according to the following equation
CH3OH + 3/2 O2 (g) CO2 (g) + 2H2O H=-730kJ
The heat given off when 6.4g of methanol burns in oxygen is
A) 32 X 730 kJ
6.4
B) 6.4 x 730 kJ
32
C) 6.4 X 730 kJ
32 16
D) 32 X 730 kJ
6.4 16
26. Which one of the following salts would give the highest volume of carbon dioxide
when reacted with
Name..Centre/Index No
Signature.
545/1
Chemistry
Paper 1
August 2011
1 hours
CHEMISTRY
PAPER 1
TIME: 1 hours
Instructions
Answer the questions by writing the correct alternative in the box on the right
hand side of the paper.
1. Fused calcium chloride when exposed to air changes from solids to liquids.
This is because the salt is:
A. Deliquescent
A
B. hygroscopic
C efflorescent
D. hydrated
2. Which of the following compounds is decomposed by heat leaving a black
residue?
D. zinc carbonate
A. hydrogen
B. nitrogen
C
C. hydrogen chloride
D. nitrogen dioxide
A. an electron
B. a neutron C
C. a proton
D. a neutron
5. Which one of the following oxides will react with magnesium most readily
when both are heated together?
A. zinc oxide B
B. copper (II) oxide
C. iron (II) oxide
D. lead (II) oxide
6. Sodium sulphite reacts with dilute hydrochloric acid according to the
equation
Na2SO3 (aq) + 2HCl (aq) 2NaCl (aq) + SO2 (g) + H2O (l)
The maximum volume in litres of the sulphur dioxide gas formed,
measured at room temperature when 6.3g of sodium sulphite reacts with
dilute hydrochloric acid is
A. 6.3 x 24 B. 126 x 24
126 6.3
7. Which one of the following statements best defines the term electrolysis?
A. conduction of electricity through a solution
B. decomposition of an electrolyte by electricity
C. ionization of an electrolyte B
D. migration of ions to the anode
11.Which one of the following reacts with water with evolution of heat?
A. calcium oxide
B. zinc oxide
C. magnesium oxide
D. lead (II) oxide
B. 25 x 0.5
2 x 20
C. 20 x 2 x 0.05
25 x 2
D. 20 x 2
25 x 0.05
13. The nuclear composition of our atoms W, X, Y and Z are shown in the
table below
Atom W X Y Z
Number of protons + neutrons 12 23 14 24
Number of protons 6 12 8 12
18. Which one of the following compounds does not produce an oxide when
heated moderately?
A. sodium hydroxide
B. copper (II) nitrate
C. calcium hydroxide
D. lead (II) nitrate
19. Which one of the following nitrates does not give off oxygen when
heated?
A. NaNO3
B. NH4NO3
C. Mg(NO3)2
D. Zn(NO3)2\
20. Which one of the following statements is correct about the electrolysis
of dilute sulphuric acid?
A. Hydrogen is liberated from the anode
B. acidity decreases at the cathode
C. sulphate ions are discharged ate the anode
D. the total amount of acid decreases
23. Hydrogen gas burns in chlorine in chlorine to form white misty fumes.
The misty fumes are;
A. smoke from the burning hydrogen
B. droplets of hydrochloric
C. chlorine water formed
D. unburnt hydrogen and chlorine
25. What mass of pure sodium hydroxide is required to make 5 litres of 0.9M
solution?
A. 0.9 x 40 x 5000
1000
B. 0.9 x 40 x 100
5000
D. 40 x 100 x 5000
0.9
27. Sulphur dioxide reacts with oxygen to give sulphur trioxide according to
the following equation.
2SO2(g) + O2(g) 2SO3(g) H =-98 kJ/ mol
28. A colourles solution ,Z, reacted with aqueous sodium hydroxide to give
a white precipitate soluble in excess alkali. When Z was treated with
aqueous potassium iodide, a yellow precipitate was observed.Z
contains
A. aluminium ions
B. lead ions
C. zinc ions
D. calcium ions
29. Which one of the following elements reacts with chlorine to form a
covalent compound?
A. calcium
B. hydrogen
C. copper
D. zinc
30. Which one of the following is the most ractive towards magnessium?
A. flourine
B. chlorine
C. bromine
D. iodine
31. 100cm3 of hydrochloric acid dissolves 3g of magnessium ribbon . The
molarity of the acid is
(Mg= 24)
A. 0.025
B. 0.25
C. 2.5
D. 25
34. Alumimium and sulpur are in group 3 and 6 of the periodic table
respectively. The formula of alumium sulphide is
A. AlS
B. AlS3
C. Al2S3
D. Al3S2
35. In the petroleum industry the hydrocarbons of higher molecular mass
are converted to hydrocarbons of lower molecular mass by
A. distillation
B. cracking
C. evaporation
D. cooling
36. Which of the following sets of elements are in the same group of the
periodic table? (atomic numbers; ( C=6, O=8, S=16,Al=13)
A. oxygen and carbon
B. oxygen and sulphur
C. oxygen and aluminium
D. carbon and sulphur
37. Which of the follwing cations NOT form a carbonate when reacted with
aqueous sodium carbonate.
A. Al3+(aq)
B. Fe2+ (aq)
C. Ca2+ (aq)
D. Mg2+(aq)
39. Which of the following halogens has the highest oxidizing power?
A. iodine
B. fluorine
C. chlorine
D. bromine
40. Which of the following reagents would produce a visible reaction when
added to aqueous sodium chloride?
A. aqueous barium chloride
B. concentrated hydrochloric acid
C. aqueous silver nitrate
D. dilute sulphuric acid
A B C D
1, 2, 3 1, 3 2,4 4
Only correct Only correct Only correct Only correct
41. Which of the following solutions contains the same mass of sodium
hydroxide?
1. 200cm3 of 2M sodium hydroxide
2. 400cm3 of 1M sodium hydroxide
3. 100cm3 of 4M sodium hydroxide
4. 400cm3 of 2M sodium hydroxide
43. Ethene
1. burns with a smoky flame
2. is an alkene
3. is a hydrocarbon
A
4. is an alkane
INSTRUCTIONS SUMMARIZED
A B C D
Assertio True True True False
n
Reason True correct True not correct false
explanation explanation
48. Sodium chloride does not because the ions in the solid are not
A
conduct electricity in the solid free to move.
state
because nitrogen has a lower
49. During the industrial molecular mass than oxygen.
preparation of nitrogen and
oxygen from liquid air, nitrogen A
comes off first
50 Magnesium can displace because the ions of both magnesium
copper from aqueous and copper carry a positive
copper (II) sulphate charge of two
B
Name..Centre/Index No
Signature.
545/1 Chemistry
Paper 1
April 2014
1 hours
Instructions
This paper consists of 50 objective type questions
Answer the questions by writing the correct alternative in the answer sheet
provide below.
1 11 21 31 41
2 12 22 32 42
3 13 23 33 43
4 14 224 34 44
5 15 25 35 45
6 16 26 36 46
7 17 27 37 47
8 18 28 38 48
9 19 29 39 49
10 20 30 40 50
1. An alloy of solder consists of
A. Zinc and Lead
B. Copper and Lead
C. Copper and Tin
D. Tin and Lead
6. In the Daniell chemical cell the reaction at the positive terminal (electrode)
is
A. Cu2+(aq)+ 2e Cu (s)
B. Cu (s) Cu2+(aq)+ 2e
C. Zn2+(aq)+ 2e Zn (s)
D. Zn (s) Zn2+(aq)+ 2e
11. Astatine belongs to group VII of the Periodic Table. Which of the following
is true about astatine? It
A. forms an ion with a positive.
B. forms a diatomic molecule.
C. is a bleaching agent.
D. is used in the treatment of water.
12. Calculate the molar heat of neutralization given off; 50cm3 of 1M sulphuric
acid was mixed with 50cm3 of 2M sodium hydroxide solution and the
temperature of the solution changed by 130C.
(Specific heat capacity of solution 4.2 Jg-1/0C, density of solution is 1gcm-1)
15. Carbon monoxide was passed over 2.32g of an oxide of iron the residue
remaining weighed 1.68g. The formula of the oxide is
[Fe =56, O = 16]
A. FeO B. Fe2O3
C. Fe3O4 D. Fe4O3
16. X, Y and Z are elements in the same short period of the periodic Table.
The oxide of X is amphoteric, the oxide of Y is basic, and the oxide of Z is
acidic. Arranged in order of increasing atomic number they are:
A. YXZ B. XYZ C. ZYX D. YZX
19. Sodium hydrogen carbonate reacts with sulphuric acid according to the
equation
2NaHCO3(s) + H2SO4 (aq) Na2SO4 (aq) + CO2 (g) +2H2O (l)
The volume of 0.4M sulphuric acid required to dissolve 4.2g of the sodium
hydrogen carbonate is
[Na=23,H=1,C=12,S=32,O=16]
A. 84 x 2 x 4.2 cm3 B. 1000 x4.2 x 0.4 cm3
1000 x 0.4 84 x 2
The table below shows the number of electrons, neutrons and protons in
particles P, Q, R, S, T and U.
Study the table below and answer questions 20 22
Particle Electron neutron protons
s s
P 19 20 19
Q 18 22 18
R 19 22 19
S 10 8 8
T 18 14 20
U 2 2 2
20. Which atoms of elements belong to the same group of the Periodic Table?
A. P and R B. Q and T
C. Q and U D. Q and R
25. Gas X was passed over heated copper (II) oxide. A brown substance and
water vapour were formed. Gas X could be
A. carbon monoxide
B. ammonia
C. sulphur (IV) oxide
D. hydrogen chloride
26. Sodium thiosulphate reacts with dilute hydrochloric acid according to the
equation
S2O32-(aq) + 2H+(aq) S(s)+ SO2(aq) + H2O(l)
A graph of the reciprocal of time (1/t) was plotted against volume of
thiosulphate that was used.
27. Magnesium was burnt in air. When was water added to the cold products,
a gas was given off. The gas was
A. hydrogen B. nitrogen
C. oxygen D. ammonia
28. A hydrocarbon Z when completely burnt in excess oxygen produced 220g
of carbon dioxide and 45 g of water. The empirical formula of Z is
A. CH B. C2 H
C. C5H2.5 D. C10H5
29. A beaker containing lime water is left exposed to the atmosphere for some
time. The white solid that formed at the bottom of the beaker is
A. calcium oxide B. calcium hydroxide
C. calcium carbonate D. calcium hydrogen carbonate.
30. The gas collected when chlorine water is exposed to sun light is
A. chlorine B. oxygen
C. hydrogen chloride D. hydrogen
31. A filter paper was soaked in a saturated solution of sodium sulphate and
placed on a slide as shown below. A crystal of copper (II) chromate was
placed onto the filter paper and a current was allowed to flow for some
time. Which of the following was observed?
32. 6.44g of hydrated salt X was strongly heat and weighed to constant mass
2.84g of the anhydrous salt remained. The number of molecules of water
of crystallization is
[X = 142, H2O = 18]
A. 3.6 B 0.2
C. 10.0 D. 6.0
33. 50 cm3 of a mixture of ammonia and carbon dioxide was passed over
heated copper (II) Oxide, and nitrogen formed occupied 20 cm3 under the
same condition as a mixture. The percentage of ammonia in the mixture
was
A. 20 B. 80
C. 40 D. 60
38. The bleaching action of chlorine is different from sulphur (IV) oxide in that
chlorine (forms)
A. choric (I) acid donates oxygen to the colored substance
B. chloric (I)acid removes oxygen from the colored substance
C. substances bleached turn brown on exposure.
D. hydrogen chloride adds hydrogen to the dye.
In each of the following questions 41 to 45, one or more answers given may
be31 correct. Read each question carefully and then choose.
A. If 1, 2, 3 only are correct
B. If 1, 3 only are correct
C. If 2, 4 only are correct
D. If only 4 is correct
INSTRUCTIONS SUMMARIZED
A B C D
1, 2, 3 1, 3 2,4 4
Only correct Only correct Only correct Only correct
42. Which of the following is/are property (ies) of concentrated sulphuric acid?
It
1. is reduced to sulphur (IV) oxide
2. gives off hydrogen when reacted with magnesium
3. is hygroscopic
4. is deliquescent
43. The gas given off when bananas ripen decolorizes acidified potassium
manganate (VII). The gas is probably
1. hydrogen sulphide
2. ethane
3. sulphur (IV) oxide
4. ethene
INSTRUCTIONS SUMMARIZED
Assertion Reason
A. True True and correct explanation
B. True True not correct explanation
C. True False
D. False True
48. Oxygen is liberated at the anode because The hydroxide ions are
when brine is electrolyzed with lower in the electrochemical
carbon electrode series than chloride ions
END.
Name..Centre/Index No/..
Signature.
545/1 Chemistry
Paper 1
August 2013
1 hours
Instructions
This paper consists of 50 objective type questions
Answer the questions by writing the correct alternative in the box on the right hand side
of the paper.
FOR EXAMINERS USE ONLY
1. Which of the following pairs of compounds is used to prepare lead (II) carbonate?
A. Lead (II) nitrate and sodium carbonate
B. Lead (II) nitrate and sodium hydrogen carbonate
C. Lead and carbonic acid
D. Lead (II) hydroxide and carbonic acid.
3. Which of the following form a precipitate when mixed with aqueous sodium
hydroxide?
A. NO2 B. CO2 C. NO2 D. Cl2
7. Which of the following pairs of compounds can be used to prepare nitric acid?
A. calcium nitrate and concentrated sulphuric acid
B. sodium nitrate and concentrated sulphuric acid
C. lead (II) nitrate hydroxide and concentrated sulphuric acid
D. barium nitrate and concentrated sulphuric acid.
11 when
In each of the following questions 41 to 45, one or more answers given may
be31 correct. Read each question carefully and then choose.
A. If 1, 2, 3 only are correct
B. If 1, 3 only are correct
C. If 2, 4 only are correct
D. If only 4 is correct
INSTRUCTIONS SUMMARIZED
A B C D
1, 2, 3 1, 3 2,4 4
Only correct Only correct Only correct Only correct
42. Which of the following is/are property (ies) of concentrated sulphuric acid?
It
1. is reduced to sulphur (IV) oxide
2. gives off hydrogen when reacted with magnesium
3. is hygroscopic
4. is deliquescent
43. The gas given off when bananas ripen decolorizes acidified potassium
manganate (VII). The gas is probably
1. hydrogen sulphide
2. ethane
3. sulphur (IV) oxide
4. ethene
INSTRUCTIONS SUMMARIZED
Assertion Reason
A. True True and correct explanation
B. True True not correct explanation
C. True False
D. False True
Signature.
545/1
Chemistry
Paper 1
August 2013
1 hours
PAPER 1
1hour 30 minutes
Instructions
Write the correct alternative in the box on the right hand side of the paper.
4. Which of the following anions forms a white precipitate with barium nitrate
solution followed by dilute nitric acid?
A. SO42- B. CO32-
C. Cl- D. NO3-
11. The atomic numbers of the elements X, Y, W and Z are 9, 11, 16 and12
respectively. Which one of the following pairs of the elements will form a
covalent compound?
C. X and W B. X and Y
C. Z and W. D. Y and W
12. The elements that can be extracted from their oxides by chemical
reduction using carbon (coke) are.
A. Al and Zn B. Zn and Fe
C. Mg and Cu. D. Ca and Cu
18. 20cm3 of 0.1M sodium carbonate solution react with 10cm3 of dilute
hydrochloric acid according to the equation below
Na2CO3(aq) + 2HCl (aq) 2 NaCl (aq) + CO2 (g) +H2O (l)
The molarity of the acid is
A. 0.1M B. 0.4M
C. 0.8M D. 0.2M
22. Which of the following reagents can be used to oxidize iron (II) ions?
A. sulphur dioxide B. hydrogen sulphide
C. chlorine. D. hydrogen chloride
25. A solid N dissolves in water to form a colourless gas that fumes with
hydrogen chloride gas. The solid M is likely to be
A. Na2O2 B. Mg3N2
C. Na2NO3 D. Mg(NO3) 2
C. 22.4x10 D. 10x85
85 22.4
35. 0.98g of gas M at stp occupies 8.4 dm3. The relative molecular mass of gas
M is (1mole of a gas occupies 22.4dm3 at stp)
A. 0.98 x 22.4 B. 0.98 x 8.4
8.4 22.4
36. Which of the following anions when in solution will react with lead (II)
nitrate to form a white precipitate which dissolves on heating and reappears
on cooling?
A. iodide ions B. chloride ions
C. sulphate ions D. carbonate ions
37. Which of the following salt does NOT produce ammonia on heating?
A. ammonium chloride B. ammonium carbonate
In each of the following questions 41 to 45, one or more answers given may
be correct. Read each question carefully and then choose.
A. If 1, 2, 3 only are correct
B. If 1, 3 only are correct
C. If 2, 4 only are correct
D. If only 4 is correct
INSTRUCTIONS SUMMARIZED
A B C D
1, 2, 3 1, 3 2,4 4
Only correct Only correct Only correct Only correct
41. Which of the following element(s) react directly with oxygen to form an
acidic oxide?
1. magnesium 2. silver
3. copper 4. gold
42. When dry hydrogen gas is passed over heated copper (II) oxide
1. copper (II) oxide is reduced. 2. Hydrogen is reduced.
3. hydrogen is oxidised 4. Copper is oxidised
44. Which of the following contains the same number of particles as 2.4 dm3 of
argon at room temperature?
[Mg = 24, Ca= 40, C= 12, O = 16, Al = 27, molar volume =24dm3 at room
temperature]
1. 2.4g of magnesium 2. 4.0g of calcium
3. 4.4g of carbon dioxide 4. 3.4g of aluminium
INSTRUCTIONS SUMMARIZED
Assertion Reason
A. True True and correct explanation
B. True True not correct explanation
C. True False
D. False True
46. Hydrogen gas is collected by because it can explode when mixed
upward delivery. with air.
47. the reactivity of alkali metals because The atomic radius of alkali
decreases down the group metals increases down the
group.
48. Hydrogen chloride gas can be because hydrogen chloride gas and
used instead of ammonia gas ammonia gas have the
in the fountain experiment. same solubility in water.
SENIOR FOUR
CHEMISTRY PAPER II
Time: 2 hours.
Instructions:
1 2 3 4 5 6 7 8 9 10 11 12 13 14 Total
SECTION A (Answer all questions in this section)
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(b) Determine the loss in mass (2 mks)
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4. State what is observed and in each case write ionic equation for the
reaction.
Observation (1
mk)
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Ionic equation (1 )
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Observation (1
mk)
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6. (a)(i) Define allotropy (1
mk)
__________________________________________________________________
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(b) Give one use of each allotrope named in a(ii) and in each case
give a property that make them suitable for the use. (2 mks)
__________________________________________________________________
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7.
The set-up above represents electrolysis of dilute sulphuric acid.
__________________________________________________________________
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8. Study the diagram below and answer the questions that follow.
HCl
Salt
Y
NaCl
Gas X _______________________________________________
Salt Y _______________________________________________
(b) Write equation for the reaction leading to the formation of gas
X (1
mk)
____________________________________________________________
_
__________________________________________________________________
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__________________________________________________________________
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I
VIII
II III IV V VI VII
U S R
P Q T
(i) Q and S (1
)
__________________________________________________________________
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(ii) T and R (1
)
__________________________________________________________________
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(b)(i) Arrange the elements above in order of increasing size
beginning with smallest. (1
mk)
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(a) Calculate
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SECTION B. (Answer any two questions)
12. (a) Starting Zinc granules, describe how zinc sulphate crystals
can be prepared in laboratory. (6
mks)
(b) Ammonia solution was added to zinc sulphate solution dropwise
until in excess.
(i) State what was observed. (1
mk)
(ii) Write ionic equation for the reaction that took place. (1 )
(c) (i) Name the reagent that can be used to identify sulphate ions
in zinc sulphate solution. (1
mk)
O = 16)
(b) The table below shows the solubilities of Salt A and Salt B at
different temperatures.
Temperature 0C 0 10 20 30 40 50 60
Solubility Salt A 13 20 32 45 63 85 110
g per 100g of water
Salt B 32.5 34 35 36 37 38 39
(i) Plot the graph of solubility against temperature for Salt A and Salt
B using the same axes. (4
mks)
(b)(i) Briefly describe how molten iron can be obtained from the ore
named in a(ii) (no diagram required). Include all the necessary
equations, uses of coke and calcium carbonate.
END.
Signature_______________________________
SENIOR FOUR
CHEMISTRY PAPER II
Time: 2 hours.
Instructions:
1 2 3 4 5 6 7 8 9 10 11 12 13 14 Total
1. Sulphur (IV) oxide can be prepared and collected in the laboratory by the
action of sulphuric acid on copper.
_______________________________________________________________________
________________________________________________________________________
(ii) What method is used to collect sulphur (IV) oxide . Explain your answer
(1mark)
____________________________________________________________________________
____________________________________________________________________________
(c) State what would be observed if sulphur (IV) oxide were bubbled into an
aqueous solution of Iron (III) oxide.
(1mark)
____________________________________________________________________________
____________________________________________________________________________
____________________________________________________________________________
(ii) Calculate the volume of carbon dioxide that would completely react with
the sodium hydroxide.
(2marks)
(i) cathode____________________________________________________________
_______
(ii)
anode_____________________________________________________________________
(iii) Write the equation for reaction that took place at the anode
(1mark)
(b) The pH of the solution at the cathode was tested using the universal
indicator.
____________________________________________________________________________
__________________________________________________________________________
____________________________________________________________________________
____________________________________________________________________________
____________________________________________________________________________
____________________________________________________________________________
___________________________________________________________________________
8. The table below shows the melting points and electrical conductivities of
substances A, B and C. Use this information to answer the questions that
follow
Electrical conductivity
Substance Melting Aqueous Molten state Solid state
s point oC solution
(a)Name the particles that are responsible for the electrical conductivity of
(1mark)
(b)Explain why substance C conducts electricity in the molten state but not
in the solid
state
(2marks)
__________________________________________________________________________
(i)B__________________________________because_______________________________
__
___________________________________________________________________________
(ii)C__________________________________because_______________________________
_
___________________________________________________________________________
9.(a) What is a fertilizer?
(2marks)
(b) Ammonium nitrate, potassium nitrate and calcium hydroxide are usually
added to soil. (i) Name one of the above salts that would be added to an
alkaline soil and explain your answer?
(2mark)
____________________________________________________________________________
____________________________________________________________________________
____________________________________________________________________________
____________________________________________________________________________
10. The table below shows the products formed when nitrates P, Q and R
are heated.
Nitrate Products
P Metal oxide, nitrogen (IV) oxide and
oxygen
Q Metal, nitrogen (IV) oxide and oxygen
R Metal nitrite and oxygen
(a)Arrange the metals in order of reactivity stating with the most reactive
(1mark)
_______________________________________________________________________
(b)Concentrated sulphuric acid was added to R and the test tube gently
heated.
____________________________________________________________________________
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______________________________________________________________________________
________________________________________________________________________
SECTION B
(c) When zinc is dipped in copper (II) sulphate, the solution turns from blue
to colourless and a brown deposit is formed.
(3marks)
(d) When sulphur (IV) oxide is bubbled in barium nitrate solution a white
precipitate is observed.
(3marks)
12.(a) Name and write the formulae of one ore from which sodium metal is
extracted (2mark)
(c) What differences, if any, are there in the products when a current is
passed through molten sodium chloride and molten sodium metal. Explain
your answer. (5marks)
14. The diagram bellow shows the extraction of sulphur by the Fraschs
process
Tank D
bed deposit
A B C
Signature_______________________________
SENIOR FOUR
CHEMISTRY PAPER II
Time: 2 hours.
Instructions:
1 2 3 4 5 6 7 8 9 10 11 12 13 14 Total
Cu =64 O=16 C=12 Na= 23 H=1 N=14 K=19
___________________________________________________________________________
Chloride
Cl-
Sulphate
SO42-(aq)
Carbonate
CO32-(aq)
4. 3.1g of a carbonate XCO3 was heated to constant mass and 2.0g of the
metal oxide was formed.
(a) Write the equations for reactions.
(1.5marks)
______________________________________________________________________________
(b) Calculate:
(ii)The volume of carbon dioxide (at stp) produced when the carbonate is
completely decomposed (1.5
marks)
5. Concentrated sulphuric acid was added to copper turnings and the mixture
heated. A white solid and a colourless gas were observed.
(a) Name the
(i) white solid___________________________________________________(1mark)
(iii)reagent used to identify the colourless gas and state what is observed
when the gas is reacted with the reagent
(2marks)
______________________________________________________________________________
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_____________________________________________________________________________
7. Part of the Periodic Table is shown in the figure below. The letters do not
represent the actual symbols of the elements.
(a)(i) State an element that can form an ion with a charge of 2- (1mark)
(ii) State the type of bond that exists in the compound formed when it reacts
with G.
(1mark)
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______________________________________________________________________________
(b)State and explain what would be observed when the universal indicator is
added to the following solutions.
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(ii)
(b)(i) State what would be observed when carbon dioxide is bubbled into
calcium hydroxide solution for a long time
(1.5marks)
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(ii)Explain why the precipitate is not observed with all salts assuming the
soap was sodium stearate
(2marks)
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SECTION B
11. (a)(i) Describe how you would prepare pure crystals of lead (II) nitrate in
the laboratory. (ii) Write the equation for the reaction.
(6marks)
(b) (i) State what would be observed when lead (II) nitrate crystals are
strongly heated.
(ii) Write the equation for the reaction.
(4marks)
(c) 0.15M hydrochloric acid was added to 25cm3 of 0.1M lead (II) nitrate.
(i) State what was observed
(ii) Write an ionic equation for the reaction
(2marks)
(iii) Calculate the volume of 0.15M hydrochloric acid needed to completely
react with the lead (II) nitrate solution.
(2marks)
(b) When a jar of hydrogen sulphide is inverted over a jar containing moist
sulphur (IV) oxide a yellow solid is observed.
(3marks)
(c) Concentrated sulphuric acid can not be used to dry ammonia gas
(3marks)
(d) When chlorine gas was passed through a solution containing potassium
bromide, a red colouration was observed. (3marks)
13.(a) (i)Describe and explain how a dry sample of hydrogen gas may be
prepared in the laboratory. (N.B No diagram is required) (4marks)
(ii) Write the equation for the reaction
(2marks)
(b) Dry hydrogen were passed over heated 8.0g copper (II) oxide.
(i) State what was observed.
(ii) Explain the observation
(3marks)
(iii) Write the equation for the reaction (2marks)
(iv) Calculate the volume of hydrogen at room temperature that reacted
completely with copper (II) oxide (2marks)
(b) The table below shows the solubilities of Salt A and Salt B at different
temperatures.
Temperature 0C 0 10 20 30 40 50 60
Solubility Salt A 13 20 32 45 63 85 110
g per 100g of Salt B 32.5 34 35 36 37 38 39
water
(i)Plot the graph of solubility against temperature for Salt A and Salt B using
the same axes. (4
marks) (ii)A saturated solution of
0 0
Salt A was cooled from 45 C to 25 C. Determine the mass of Salt A
deposited. (2 marks)
(iii)If salt A is potassium nitrate calculate the numbers of moles of potassium
nitrate obtained in b(ii) above. (2
marks)
Signature........................................................................
SENIOR FOUR
END OF TERM 1 EXAMINATION 2011
CHEMISTRY PAPER II APRIL
Time: 2 hours.
Instructions:
Section A consists of 10 structured questions. Attempt all questions. Answers to
these questions must be written in the spaces provided.
Section B consists of 4 semi-structured questions. Attempt any two any
questions in this section. Answers to these questions must be written in the
answer booklets provided.
In both sections all working must be clearly shown.
2.(a) Burning magnesium was lowered into a gas jar of sulphur (IV)
oxide gas.
(i) State what was observed. (1marks)
------------------------------------------------------------------------------------------------------------
------------------------------------------------------------------------------------------------------------
------------------------------------------------------------------------------------------------
(ii) Write equation for the reaction that took place. (1marks)
------------------------------------------------------------------------------------------------------------
----------------------------------------------------------------------------------------------------
(iii) State the property of sulphur (IV) oxide shown in the reaction. (marks)
------------------------------------------------------------------------------------------------------------
----------------------------------------------------------------------------------------------------
(b) Write equation for one other reaction in which sulphur (IV) oxide shows
property you have stated in a (iii). (1marks)
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----------------------------------------------------------------------------------------------------
3. The diagram below shows the set up of the apparatus that was used to
investigate the effect of hot manganese (IV) oxide on hydrogen chloride
gas.
(a) Write equation for the reaction in that took place in the glass tube
(1marks)
------------------------------------------------------------------------------------------------------------
----------------------------------------------------------------------------------------------------(b)
State
(i) the role of manganese (IV) oxide in the reaction in (a) (mark)
--------------------------------------------------------------------------------------------------- ----
(ii) what was observed in the beaker containing potassium iodide solution.
(1marks)
------------------------------------------------------------------------------------------------------------
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------------------------------------------------------------------------------------------------
(c) Write an ionic equation for the reaction to illustrate your answer in (b)(ii).
(1marks)
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4.(a)(i) Define the term salt. (1mark)
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------------------------------------------------------------------------------------------------------
(ii) Write the name and formula of an acid salt. (1mark)
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(b) When 2.4g of crystals of salt W were heated, 1.5g of anhydrous salt was
obtained. Determine the number of moles of water of crystallisation in one
mole of crystalline W.
(The molar mass of W =158, H =1, O= 16) (3marks)
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5. 25.0cm3 of 0.1M sodium carbonate solution reacted completely with a
solution containing 5.0g of sulphuric acid in 250cm3 of solution.
Calculate
(i) the molarity of sulphuric acid in moles per dm3. (3marks)
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(ii) the volume of acid used. (2marks)
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6. The elements carbon, magnesium and aluminium can combine with
oxygen to form oxides.
(a) State the class of the oxides of
(i) Carbon ( mark)
--------------------------------------------------------------------------------------------------------
(ii) Magnesium ( mark)
--------------------------------------------------------------------------------------------------------
(iii) Aluminium ( mark)
--------------------------------------------------------------------------------------------------------
(b) Write equation for the reaction that would take place in each case if the
oxides of carbon, magnesium were separately treated with water.
(i) oxide of carbon (1 mark)
------------------------------------------------------------------------------------------------------------
----------------------------------------------------------------------------------------------------
(ii) oxide of magnesium (1 mark)
------------------------------------------------------------------------------------------------------------
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(c) Name one other element that can combine with oxygen to form an oxide
in the same class as that of aluminium. (mark)
--------------------------------------------------------------------------------------------------------
7.(a) During the laboratory preparation of hydrogen gas, dilute sulphuric acid
was added to zinc metal followed by a little amount of cooper (II)
sulphate solution.
(i) Write equation for the reaction leading to the formation of hydrogen gas.
(1 mark)
------------------------------------------------------------------------------------------------------------
----------------------------------------------------------------------------------------------------
(ii) State the role of copper (II) sulphate. ( mark)
--------------------------------------------------------------------------------------------------------
(b) State the conditions under which hydrogen may be displaced from water
by:
(i) sodium metal ( mark)
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(ii) iron filings ( mark)
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(c) Write the equation for the reaction in (b)(ii) (1 mark)
------------------------------------------------------------------------------------------------------
8.(a) A mixture of magnesium powder and lead (II) oxide was heated strongly
until no further change.
(i) State what was observed. (1 mark)
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(ii) Write equation for the reaction that took place. (1 mark)
------------------------------------------------------------------------------------------------------------
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(b) The experiment in (a) was repeated using a mixture of copper turnings
and magnesium oxide.
State what was observed, ( mark)
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(b) Write equation for the reaction leading to the formation of the product
collected at Y (1mark)
------------------------------------------------------------------------------------------------------------
----------------------------------------------------------------------------------------------------(c)
Name the substance used as the anode. (1mark)
--------------------------------------------------------------------------------------------------------
(d) State;
(i) one industrial use of sodium hydroxide. ( mark)
(ii) what is observed if chlorine gas is bubbled through a cold dilute solution
of sodium hydroxide? (1mark)
--------------------------------------------------------------------------------------------------------
------------------------------------------------------------------------------------- ( mark)
(ii) how ammonia is collected and give a reason for your answer. (1mark)
------------------------------------------------------------------------------------------------------------
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(iii) why ammonia is not dried using fused calcium chloride or concentrated
sulphuric acid. (no equation required). (1 marks)
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(c) When X g of ammonium chloride was used in preparation of ammonia
as shown in the equation in (a0, 4.0g of pure and dry calcium chloride
was obtained. Determine the value of X
(H =1, N= 14, Cl = 35.5, Ca = 40)
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(b) The analysis in (a) was repeated on a day that had rained.
(i) State how the volume of carbon dioxide would compare with that
obtained in (a) ( marks)
(ii) Explain your answer in (b) (i) (1 marks)
(c) Name one reagent that can be used to remove carbon dioxide from a
sample of air and write the equation that would take place when the
reagent is used. (2marks)
(c) Write the structure of the compound formed when carbon combines with
chlorine. (1mark)
13(a) Outline how a reasonably pure and dry sample of copper (II) sulphate
crystals can be prepared in the laboratory, starting from copper (II) oxide.
Write equation(s) to illustrate your answer. (7marks)
(b) State what would be observed, and write equation for the reaction that
would take place if;
(i) a crystal of copper (II) sulphate was dropped in concentrated sulphuric
acid and the mixture warmed. (2marks)
(ii) to an aqueous solution of copper (II) sulphate was added dilute
sodium hydroxide solution drop-wise until the alkali was in excess.
(2marks)
(c) Lead (II) ions can react with sulphate ions according to the following
equation.
Pb 2+(aq) + SO4(aq) PbSO4(s)
(i) State what would be observed if lead (II) nitrate was added to aqueous
copper (II) sulphate solution. (1mark)
(ii) Excess aqueous lead (II) nitrate solution was added to a solution
containing 2.50g anhydrous copper (II) sulphate. The mixture was
stirred, filtered and the residue dried and weighed.
Calculate the mass of the dry residue that was obtained. (3marks)
(Pb = 207, Cu = 64, S = 32, O = 16)
14. Iron (III) oxide (haematite) is one of the common ores of iron from which
iron can be extracted in the blast furnace.
(a) Name
(i) one common ore of iron other than iron (III) oxide. (1mark)
(ii) one major impurity that can be found in the ore you have named in (a)(i)
(1mark)
(b) Outline the reactions which occur in the blast furnaces during the
extraction of iron from iron (III) oxide ore. (7marks)
(d) Most common compounds of iron are either those of iron (II) or iron (III).
Write the formula of one compound of
(i) iron (II) (mark)
(iii) iron (III) (mark)
(e) Name one reagent that can be used to distinguish between iron (II) and
iron (III) compounds and in each case state the observations that would be
made if the reagent you have named was used. (2marks)
Name.............................................................. Index /No..................................
Signature........................................................................
SENIOR FOUR
MID TERM EXAMINATION 2011
CHEMISTRY PAPER II
Time: 2 hours.
Instructions:
Section A consists of 10 structured questions. Attempt all questions.
Section B consists of 4 semi-structured questions. Attempt any two any
questions n this section.
1 2 3 4 5 6 7 8 9 10 11 12 13 14 Total
Relative atomic masses; O=16, C=12, Na= 23, H=1, 0=16, S=32, P=31
Molar volume at room temperature 24dm3
SECTION A (Answer all questions in this section)
2 The following table shows the atomic numbers and mass numbers of
elements A, B and C.
(b) Briefly describe a test you would carry out to identify gas X.
(2marks)
..................................................................................................................
..................................................................................................................
..................................................................................................................
(ii). Write an ionic equation for the reaction in (a) (ii). (1mark)
..................................................................................................................
.................................................................................................................
8. A compound X of formula mass 90 contains 26.67% carbon and 2.22%
hydrogen by mass, the rest being oxygen.
(a). Calculate
(i). the empirical formula of X (3marks)
..................................................................................................................
..................................................................................................................
..................................................................................................................
..................................................................................................................
..................................................................................................................
..................................................................................................................
..................................................................................................................
..................................................................................................................
..................................................................................................................
(ii). the molecular formula of X (1mark)
..................................................................................................................
..................................................................................................................
..................................................................................................................
.................................................................................................................
10. Study the diagram below and answer the questions that follow.
Concentrated
Hydrochloric
acid
Salt Y
..
(ii) Write equation for reaction in b (i) above (1mark)
SECTION B
Answer two questions only in this section. Start each question on a fresh page.
(b) (i) Briefly describe how a pure sample of hydrated copper (II)
sulphate is prepared in the laboratory. (7marks)
(i) Write the equation for the reaction. (1mark)
(c) Copper (II) sulphate-5- water was heated strongly until no further
change.
(i) State what was observed (1mark)
(ii) Write the equation for the reaction (1mark)
(c) A strip of zinc metal was dipped into aqueous copper (II)
sulphate solution.
(i) State what was observed (1mark)
(ii) Write an ionic equation for the reaction. (1mark)
N. P. K. FERTILIZER
(b) State the industrial processes by which the ammonia and nitric
acid used for making the N.P.K. fertilizer are produced (2marks)
Signature........................................................................
SENIOR FOUR
BEGINNING OF TERM EXAMINATION 2012
CHEMISTRY PAPER II
Time: 2 hours.
Instructions:
Section A consists of 10 structured questions. Attempt all questions.
Section B consists of 4 semi-structured questions. Attempt any two any
questions n this section.
1 2 3 4 5 6 7 8 9 10 11 12 13 14 Total
.................................................................................................................
2. Name one pair of substance in each case, which when mixed together
can be separated by
(a) Use of a separating funnel (1 marks)
(b) Fractional crystallization (1 mark)
(c) Filtration (1 mark)
(d) Chromatography (1 mark)
(e) Fractional distillation (1 mark)
3.(a) Excess carbon monoxide gas was passed over a heated sample of an
oxide of iron as shown in the diagram below. Study the diagram and the
data below to answer the questions that follow.
4. The table below shows the products formed when nitrates P, Q and R are
heated.
Nitrate Products
P Metal oxide, nitrogen (IV) oxide and oxygen
Q Metal, nitrogen (IV) oxide and oxygen
R Metal nitrite and oxygen
(a) Arrange the metals in order of reactivity stating with the most reactive.
(1mark)
(b) Concentrated sulphuric acid was added to R and the test tube gently
heated.
(i) State what was observed (1mark)
.
(iii) Write the equation for the reaction (1marks)
(c) Nitrogen compounds are pollutants. State two examples (1mark)
5. Magnesium burnt in air to form magnesium nitride.
(a) Write the equation for the reaction.
(1marks)
(b) Calculate the volume of air necessary to completely react with 6.0g of
magnesium if the percentage of nitrogen in air is 78 by volume.
(3marks)
..
..
6. Sulphur (IV) oxide can be prepared and collected in the laboratory by the
action of sulphuric acid on copper.
(a) (i) State the conditions for the reaction (1mark)
(ii) Write the equation for the reaction (1marks)
(b) Name one reagent that can be used for
(i) identifying sulphur (IV) oxide and ethene (1 mark)
..
(ii) distinguishing between sulphur (IV) oxide and ethene (1 mark)
..
(iii) State what would be observed if the reagent you have named in (a) (ii)
was treated with each ( mark @)
sulphur (IV) oxide .
and ethene.
(b) Write equation for the reaction leading to the formation of the product
collected at Y (1marks
(c) Name the substance used as the anode ( mark)
(d) State
(i) One industrial use of sodium hydroxide ( mark)
(ii) Write equation for the reaction when chlorine gas is bubbled through a
cold dilute solution of sodium hydroxide. (1 marks)
I VIII
II III IV V VI VII
U S R
P Q T
(a) Write the formula of a compound formed between pairs of elements and in
each case state the type bond formed.
SECTION B
Attempt only two questions
(b) When sodium hydroxide is heated with ammonium chloride a gas that
turns litmus blue is given off. (3marks)
(c) When sodium hydroxide is added drop-wise until in excess to zinc
sulphate solution, a white precipitate is formed. The precipitate
dissolves in excess sodium hydroxide solution to form a colourless
solution. (4marks)
12(a)(i) Describe how you would obtain a sample of sugar crystals from sugar cane
(7marks)
(ii) State two uses of sugar in the world of the sick (2marks)
(c) Outline how a pure dry sample of sodium chloride can be prepared in the
laboratory. (6marks)
(c). An aqueous solution of lead (II) nitrate was added to a solution of sodium
chloride and the resultant product was heated.
(i) State what was observed. (1mark)
(ii) Write an ionic equation for the reaction that took place. (1marks)
(d) Iron (III) chloride dissolves in water according to the following equation:
FeCl3(s) + 3H2O (l) Fe(OH)3(s) +3 HCl (aq)
(i) State what would be observed when a solution of sodium hydrogen
carbonate is added to iron (III) chloride solution. (1mark)
(ii) Write an ionic equation for the reaction that took place. (1mark)
14(a). Hydrogen can react with nitrogen in the presence of a catalyst to produce
ammonia on an industrial scale.
(i). State the sources of hydrogen and nitrogen which are used in the reaction
(1
marks)
(ii). Name the catalyst used in the reaction. (1mark)
(iii). Write equation for the reaction leading to the formation of ammonia. (1marks)
(b). In order for maximum yield of ammonia to be achieved during the industrial
preparation of ammonia as stated in (a) the reaction is carried out at a low
temperature.
(i). Give a reason for carrying out the reaction at low temperature. (1mark)
(ii). State one condition for the reaction other than low temperature and use of the
catalyst. (1mark)
(iii). Briefly describe and explain how the condition you have stated in (b)(ii) affects
the reaction. (2marks)
(c). The ammonia obtained by the reaction in (a) above can be oxidized in the
presence of a catalyst to manufacture nitric acid.
(i). Name the catalyst which is used in the reaction. (1mark)
(ii). State two conditions for the oxidation of ammonia during the manufacture of
nitric acid other than the catalyst which you have named in (c) (i).
(iii). Using equation only show how nitric acid can be produced by the catalytic
oxidation of ammonia. (3marks)
(d). Name one ammonium compound that is commonly used in agricultural industry
and state the purpose for which the compound is used in agriculture. (1mark)
Signature........................................................................
SENIOR FOUR
MID TERM EXAMINATION 2012
CHEMISTRY PAPER II
Time: 2 hours.
Instructions:
Section A consists of 10 structured questions. Attempt all questions.
Section B consists of 4 semi-structured questions. Attempt any two any
questions n this section.
1 2 3 4 5 6 7 8 9 10 11 12 13 14 Total
Relative atomic masses; O=16, C=12, Na= 23, H=1, 0=16, S=32, P=31
Molar volume at room temperature 24dm3
Molar volume at standard temperature and pressure 22.4dm3
1(a) A teabag was each suspended separately in a beaker containing of hot water
and another containing cold water and allowed to stand for 10 minutes
2. The experiment below was set up and left to stand for 5days
(b) Galvanization is one method used to prevent rusting. Explain why galvanized iron
does not rust. (1 marks)
3. The graphs below show the boiling curve of pure water and an aqueous
solution of sodium chloride A
Temperature (0C)
Time in minutes
............................................................................................................................
............................................................................................................................
............................................................................................................................
(b) Which boiling curve represents an aqueous solution of sodium chloride? Explain
your answer (2marks)
............................................................................................................................
............................................................................................................................
............................................................................................................................
(c) State one other factor that affects the boiling point of a liquid (1mark)
...............................................................................................................................
(d) Boiling points are used to determine the criteria of purity of a liquid.
State one other method .................................................................. (1mark)
4. Study the diagram below and answer the questions that follow
5. 3.1g of a carbonate XCO3 was heated to constant mass and 2.0g of the metal
oxide was formed.
(a) Write equation for the reactions. (1 marks)
.
(b) Calculate:
(i) the volume of carbon dioxide at stp produced when the carbonate is completely
decomposed (molar volume at stp is 22.4dm3) (2marks)
.................................................................................................................................
.................................................................................................................................
.................................................................................................................................
.................................................................................................................................
.................................................................................................................................
.................................................................................................................................
.................................................................................................................................
.................................................................................................................................
(ii) the atomic mass of X (0 = 16, C=12)
.................................................................................................................................
.................................................................................................................................
.................................................................................................................................
.................................................................................................................................
.................................................................................................................................
.................................................................................................................................
.................................................................................................................................
.................................................................................................................................
.
6.(a) (i) Name a substance which when reacted with sodium peroxide can produce
oxygen (mark)
.
(ii) Write an equation for the reaction leading to the formation of oxygen.
(1 marks)
(b) Write an equation for the reaction that takes place when
(i) Hydrogen is burnt in excess oxygen (1 marks)
..
(ii) Sodium is burnt in limited supply of oxygen (1 marks)
..
8. State what is observed and in each case write ionic equation for the reaction.
(a) Chlorine gas is bubbled through a solution of Iron (II) chloride.
Observation (1mark)
Ionic equation (1 marks)
(b) When lead (II) nitrate solution is added to a solution of sodium iodide
Observation (1mark)
Ionic equation (1 marks)
(b). Calculate
(i) volume of ammonia evolved at stp
(N =14, Cl=35.5, H = 1) (2marks)
(ii) number of moles of sodium hydroxide that was in excess (1mark)
10. The following chart below shows the contact process that is used to manufacture
sulphuric acid.
(ii) Write equation for the reaction taking place in Chamber A. (1mark)
...
(c) State what is observed when concentrated sulphuric acid is added to hydrated
(i) copper (II) sulphate. (1mark)
(ii) potassium chloride (mark)
SECTION B
Attempt any two questions in this section. Start each question on a fresh page.
11(a)(i) Describe how you would prepare pure crystals of lead (II) nitrate in the
laboratory.
(ii) Write the equation for the reaction. (7marks)
(c) 0.15M hydrochloric acid was added to 25cm 3 of 0.1M lead (II) nitrate.
(i) State what was observed
(ii) Write an ionic equation for the reaction (2marks)
(iii) Calculate the volume of 0.15M hydrochloric acid needed to completely react
with the lead (II) nitrate solution. (2marks)
12. Explain the following observations. Write equations to explain your answer
where applicable.
(a) Sodium conducts electricity both the solid and liquid state whereas sodium
chloride does not conduct electricity in the solid state but conducts electricity in
the molten state. (4marks)
(b) A solution of hydrogen chloride gas in methyl benzene has no effect on dry
litmus paper but an aqueous solution of hydrogen chloride gas turns moist
litmus paper red. (3marks)
(c) When a jar of hydrogen sulphide is inverted over a jar containing moist sulphur
(IV) oxide a yellow solid is observed. (3marks)
(b) Draw a labeled diagram of the set up of the apparatus to show that hydrogen
chloride gas very soluble in water (3marks)
(b). State what would be observed and write equation for the reaction that would
take place if aqueous hydrogen chloride was reacted with solid calcium
carbonate. (3marks)
a. Describe an experiment you would carry out in the laboratory to determine the
molecular formula of magnesium oxide.
Signature........................................................................
SENIOR FOUR
END OF TERM 1 EXAMINATION
APRIL 2014
CHEMISTRY PAPER II
Time: 2 hours.
Instructions:
Section A consists of 10 structured questions. Attempt all questions.
Section B consists of 4 semi-structured questions. Attempt any two any questions n this
section.
1 2 3 4 5 6 7 8 9 10 11 12 13 14 Total
---------------------------------------------------------------------------------------------------------------------
-------------- ---------------------------------------------------------------------------------------------------
--------------------------------- ---------------------------------------------------------------------------------
---------------------------------------------------(i) State two property of steel which is the reason
for its being used widely than iron
(2mark)
------------------------------------------------------------------------------------------------------------
------------------------ ------------------------------------------------------------------------------------------
------------------------------------------(b) State the constituents of each of the following
alloys
(i) Duralumin
(1mark)
------------------------------------------------------------------------------------------------------------
--- (ii) Solder
(1mark)
------------------------------------------------------------------------------------------------------------
---
4. Iron (II) sulphate 7 water and copper (II) sulphate 5 water are prepared in
the laboratory by action of sulphuric acid on iron and copper respectively.
(a) Name the conditions for the reaction of sulphuric acid and
(i) Iron
(mark)
. ---------------------------------------------------------------------------------------------------
---(ii) Copper
(mark)
---------------------------------------------------------------------------------------------------
---
(b). Write equation to show the reaction in which copper reacts with sulphuric
acid under the conditions stated to form copper (II) sulphate.
(1marks)
---------------------------------------------------------------------------------------------------
--- ---------------------------------------------------------------------------------------------------
---
(c) Write equations to show the effect of heat on the
(i) Iron (II) sulphate 7 water (1marks)
---------------------------------------------------------------------------------------------------
--- --------------------------------------------------------------------------------------------------
(ii) Copper (II) sulphate 5 water
(1mark)
---------------------------------------------------------------------------------------------------
---
---------------------------------------------------------------------------------------------------
--- --------------------------------------------------------------------------------------------------
(b)(i) From the observation in (a) what elements are present in a candle wax.
Explain your answer
(2marks)
---------------------------------------------------------------------------------------------------
--- ---------------------------------------------------------------------------------------------------
--- --------------------------------------------------------------------------------------------------
--------------------------------------------------------------------------------------------------
(ii) Write an equation to represent the combustion of a candle (1
marks)
---------------------------------------------------------------------------------------------------
---
(ii) oxygen (
marks)
(a) Write equation to show the reaction in which manganese (IV) oxide together with
the substance that you have named in (a) produce
(i) Chlorine (1 marks)
----------------------------------------------------------------------------------------------------------------
----------------------------------------------------------------------------------------------------------------
----
(ii) Oxygen
(1marks)
---------------------------------------------------------------------------------------------------------------
---------------------------------------------------------------------------------------------------------------
----
(b) Write an equation for the oxidation of iron by oxygen (1mark)
-----------------------------------------------------------------------------------------------------------------
----------------------------------------------------------------------------------------------------------------
7(a). 20.0cm3 of a solution containing 6.3g per dm3 of a dibasic acid H2X. nH2O
required exactly 20.15cm3 of a 0.1M sodium hydroxide solution for complete
neutralization. (i). Calculate the
concentration of the acid in moles per dm3 of the solution. (3marks)
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-- ------------------------------------------------------------------------------------------------------------
-- ------------------------------------------------------------------------------------------------------------
-- ------------------------------------------------------------------------------------------------------------
-- ------------------------------------------------------------------------------------------------------------
-- ------------------------------------------------------------------------------------------------------------
-- ------------------------------------------------------------------------------------------------------------
-- ------------------------------------------------------------------------------------------------------------
--- ------------------------------------------------------------------------------------------------------------
---
------------------------------------------------------------------------------------------------------------
--- ------------------------------------------------------------------------------------------------------------
--- ------------------------------------------------------------------------------------------------------------
---
10. Study the table below and answer the following questions.
(b) Which substance(s) would dissolve in water and can be separated from the
solution by (marks@)
(i) fractional distillation
----------------------------------------------------------------------------------
(ii) evaporation of the water
----------------------------------------------------------------------------
11(a) Describe how dry crystals of copper (II) nitrate can be prepared in the laboratory
starting from copper (II) oxide
(4marks)
(b) State what would be observed when an aqueous solution of copper (II) nitrate is
added
(i) sodium hydrogen carbonate solution (
mark)
(ii) zinc powder (1
mark)
(c)(i) Write ionic equations for the reactions that took place in (b)(i) and (b)(ii)
(2mark)
(ii) Explain the observation in(b)(i)
(2marks)
(b) Copper (II) nitrate crystals were heated strongly until no further change.
State what was observed and write the equation.
(3marks)
(d). State
(i) how you would distinguish between hand water and soft water. (1
marks)
(ii). one use of hard water. (
mark)
(e). Name one compound which when dissolved in water will cause.
(i). permanent hardness of water
(mark)
(ii). temporary hardness of water (
mark)
13. Explain each of the following observations. Write equations where appropriate
(a) Sodium metal conducts electricity in the solid state whereas sodium chloride
conducts electricity in the molten state.
(5marks)
(b) Hard water wastes soap.
(5marks)
(d) When a mixture of magnesium powder and copper (II) oxide are heated together
a mixture of a white and brown solids are formed.
(5marks)
14. In recent years water pollution has become a serious problem; some of the
main sources of pollution being sewage and use of detergents
(b) The flow diagram below shows what takes place in the sewage works.
Water Chlorination
bodies
(3marks)
TOTAL 5
2(a) (i) Allotropy is the existence of an element in two or more forms due to different 1
arrangement of atoms in the same physical state. (1)
(c) 1
2C(s)+O2(g) 2CO(g)
TOTAL 5
3(a)(i) Copper; electrons()
(ii) aqueous copper (II) sulphate; ions ()
(b) (i) at electrode A (+ve) the electrode dissolves ()and decreases in mass()
at electrode B (-ve) a brown coating ()and increases in mass()
(ii) Cu(s) Cu2+ (aq) + 2e (1)
(ii) purification of metals ()
electroplating ()
extraction of metals
TOTAL 5
4(a) (i) Iron; dilute ()sulphuric acid
ii) Copper; hot concentrated() sulphuric acid
TOTAL 5
5(a) A white precipitate () colourless liquid on the watch glass, () the level of
calcium hydroxide rose ()
TOTAL 5
6(a) (i) hydrochloric acid ()
(ii) hydrogen peroxide ()
TOTAL 5
3
7(a) (i) In 20.15 cm of sodium hydroxide there are (20.15 x 0.1)/1000 (1) 5
= 2.015 x 10-3
Reaction ratio acid to base 1: 2
In 20.00 cm3 of acid there are (2.015 x 10-3)/2 = 1.0075 x 10-3 (1)
In 1000cm3 of acid there are (1.0075 x 10-3 x1000)/20 (1)
= 0.05
TOTAL 5
8(a) (i) Ammonium chloride ()and calcium hydroxide ()
(ii) heat()
(iii) Ca(OH)2(s) + 2NH4Cl(s) CaCl 2(s) + 2NH3(g) + 2H2O(l) (1)
TOTAL 5
10(a o Add small spatula ends full of copper (II) oxide powder to dilute nitric acid () in 6
) a beaker until the oxide can dissolve no more.()
o Warm the mixture to increase the rate of reaction.
o Filter () the mixture to remove the unreacted copper (II) oxide powder.
o Place the filtrate in an evaporating basin and heat () until the solution is
saturated. ()
o Allow the solution to cool () and crystallise. ()
o Filter off the crystals () wash () in little water () and dry () between filter
papers.
o CuO(s) +2HNO3(aq) Cu(NO3)2(aq) + H2O(l) (1)
(b) 2
(i) Cu2+(aq) + 2HCO3-(aq) CuCO3(s) + CO2(g) +H2O(l) (1)
2+
(ii) Cu (aq) + Zn(s) Cu (s) + Zn2+ (1)
(iii) The green precipitate is copper (II) carbonate ()that is insoluble()
(c) 3
The blue ()crystals turn to black, ()a brown gas given off(),( a colourless liquid
given off)( three observations)
2Cu(NO3)2(s) 2CuO (s) + 4NO2 (g) + O2(g) (1)
(d) 3
TOTAL 15
12(a (i) No of grams of a solute that dissolves in 100g of solvent to form a saturated 2
) solution (until no more solute can dissolve) at that temperature. (2)
(ii) Factors : temperature(1) nature of solvent (1)
Explanation: as temperature rises solubility off most solutes increase(1) 4
(iii) the higher the temperature the higher the solubility (1)
Covalent solutes dissolve in organic solvents
Ionic solutes dissolve in polar solvents(1)
(b) 70
3
60
50
40
Solubility(grams/100gra
solubility g/100g solvent ms of solvent
30 Exponential
(Solubility(grams/100gra
ms of solvent )
20
10
0
0 100 200
Temperature
(c)
(d)
Solubilit
y at 150C 4.75g (1)
0
At 75 C 32.5 g (1)
(ii) mass that crystallises out 32.5 4.75 = 27.75g (1)
(i) add soap solution(); soft water lathers easily(), hard water does not lather
easily()
(ii) brewing beer()
TOTAL 15
13(a Sodium metal has delocalized () electrons () that are responsible for carrying the 3
) current.
In the solid state sodium chloride have ions () that are held together by strong
electrostatic forces, ()the ions are not free to move to carry the current.
In the molten state the forces of attraction between the ions are weakened (), and
the ions are mobile () and free to can carry the current.
Hard water contains calcium (1) or magnesium ions. On addition of soap, scum ()is
(b) formed which is an insoluble salt ()of magnesium or calcium stearate as follows: 4
2+ -
Ca (aq) + 2St (aq) Ca St2(s) (1)
It is only when all the calcium and magnesium ions have been removed() from water
that a lather will form()
Sulphuric acid is an acid ()ammonia is alkaline ()and the two react forming a salt() 3
(c) H2SO4(aq) + 2NH3(g) (NH4)2SO4(aq) (1)
Magnesium is higher in the reactivity series than copper(1), it reduces (1), copper
(d) (II) oxide to copper (),; the brown solid and itself is oxidised (1), to magnesium
oxide (),; the white solid. 5
Mg(s) + CuO(s) MgO (s) + Cu(s) (1)
TOTAL 15
14(a Waste water refuse containing human excreta and other waste (1), from factories, 3
) gutters, toilets, ,bathrooms and kitchens (mentions at least two sources) (1) that
enter sewers() and flow to sewage works ()
(i) Sedimentation; involves the clustering together of suspended particles into big
(b) lumps () which sink and settle at the bottom. Potassium aluminium sulphate 1
(Alum) (1)is added
(ii) Anaerobic digester; there is no oxygen here () appropriate temperature () is 1
and right pH() are required
(iii) Aeration tank; To enable the an aerobic micro organism grows here () and 1
breaks down any remaining organic material()
(iv) Chlorination; To kill germs(1)
1
(iv) detergents contain phosphates (1) that promote the growth of algae that
deprive oxygen (1)from aquatic life
TOTAL 15
Name.............................................................. Index /No..................................
Signature........................................................................
SENIOR FOUR
MID TERM EXAMINATION
MARCH 2014
CHEMISTRY PAPER II
Time: 2 hours.
Instructions:
Section A consists of 10 structured questions. Attempt all questions.
Section B consists of 4 semi-structured questions. Attempt any two any questions n this
section.
1 2 3 4 5 6 7 8 9 10 11 12 13 14 Total
Attempt all questions
1 Air is a mixture of gases
(a) State
(i) two reasons why air is regarded as a mixture and not a compound.
(1mark)
.
(ii) the method by which the major components of air are separated industrially.
(1mark)
.
(iii) Give a reason for your answer in (a)(ii)
(marks)
.
(b) Write equation to show the reaction that take place between the most abundant
components of air with magnesium. (1
marks)
..
(c) Name one major pollutant of the atmosphere and state its effect
(1mark)
Pollutant.
Effect
2(a)(i) State what is observed when sodium carbonate solution is added to copper (II)
sulphate solution.
(mark)
(b)(i) State what was observed when dilute nitric acid was added to to the product of
the raction in (b)(i)
(1marks)
3. The table below shows the physical properties of elements W,X,Y and Z.
Subastance Melting point 0C boiling point 0C Solubility in water Density at 250C
W -117 78.5 very soluble 0.8
X -78 -33 very soluble 0.77 x 10-3
Y -23 77 insoluble 0.16
Z -219 -183 slightly soluble 1.33 x 10-3
(a) (i)Identify the substance that would dissolve in water and could be separated from
solution by fractional distillation. .
(1mark)
(ii) Give a reason for you answer.
(1mark)
(b) Which substances is a liquid at room temperature and when mixed with
water forms two separate layers?
(1mark)
(b) Write down the 2 possible structural formulae of P and give their systematic
names
(2marks)
...
...
..
..
5(a) Name the reagents that are neded to prepare sodium hydrogen carbonate and
write the equation for the reaction.
(1mark)
Reagents..
.......
Equation
(1marks)
...
...
(b) Iron (III) chloride dissolves in water according to the following equation :
FeCl3(s)+3H2O(l) Fe(OH)3(s)+3HCl(aq)
(i) State what is observed when a solution of iron (III) chloride is added to sodium
hydrogen carbonate solution.
(1mark)
(ii) Explain your answer.
(1mark)
.
..
6. Part of the Periodic Table is shown below, with some elements labelled Q to Z.
The letters are not the usual symbols of these elements.
I II II IV V VI VII VII
Q W Z
X Y
R T
(ii) W and X
(1mark)
(c) Arrange the elements above in order of increasing size beginning with smallest.
(1mark)
..
(b) Write the equation of the following half equations and in each case state whether
the reaction is reduction or oxidation reaction.
(i) The conversion of hydrogen ions (H+) to hydrogen molecules (H2)
(1mark)
............................................................................................................................
...........................................................................................................................
(ii) The conversion of iron (II) ions (Fe2+) to iron (III) ions (Fe3+)
(1mark)
............................................................................................................................
...........................................................................................................................
(iii) The conversion of chlorine molecules to chloride ions (Cl --)
(1mark)
......................................................................................................................
...... ......................................................................................................................
....
.8. Name one reagent that can be used to distinguish between each of the following
pairs of ions. In each case state what would be observed when the reagent was
treated separately with each pairs of ions.
(a). chloride and iodide ions.
(i). reagent
(1mark)
.......
.....
(ii). Observation
(1mark)
Chloride
iodide
(ii). Observation
(1mark)
Hydrogen
carbonate.
Carbonate
(ii). Observation
(1mark)
Sulphate
Sulphite
9. The diagram below shows the set up of apparatus that was used to investigate
the effect of hot platinum on a mixture of sulphur dioxide and oxygen.
(a)(i) State what was observed. (1mark)
..
(ii). Give a reason (s) for your observation in (a) (i).
(1mark)
..
(iii). Write equation for the reaction that took place. (1mark
)
(b) (i).State the purpose of using the hydrogen peroxide mixed with manganese (IV)
oxide.
(1mark)
10. 1.00g of pure ammonium chloride was boiled with 20.0cm 3 of a solution of
sodium hydroxide until evolution of ammonia had ceased. If the resulting
solution required 11.0 cm3 of 0.1M hydrochloric acid for neutralization.
(a). Write equation for the reaction that took place between ammonium chloride and
sodium hydroxide (1 marks)
(b). Calculate
(i) Volume of ammonia evolved at stp
(N =14, Cl=35.5, H = 1) (2marks)
SECTION B
Attempt any two questions in this section. Start each question on a fresh page.
11. Soap is prepared in laboratory by reacting vegetable oil or fat and an alkali
(a) State
(i) the process by which soap is prepared.
(1mark) (ii) the differences between fats and oils
(2marks)
(iii) one source each of a fat and oil
(1mark )
(c) Starting from the named oil or fat; describe how soap can be prepared in
laboratory.
(5marks)
(d) Explain the following; illustrate your answer using equations.
When soap was added to aqueous calcium hydrogen carbonate solution; a white
precipitate was formed. But when the soap solution was added to aqueous
calcium hydrogen carbonate solution that had been boiled no precipitate way
formed.
(6marks)
12(a)(i) With the aid of equations, outline how a dry sample of hydrogen chloride gas
can be prepared in the laboratory staring from sodium chloride.
N.B. Diagram not required
(5marks)
(ii) State two uses of hydrochloric acid.
(1mark)
(4marks)
(ii) Name the reagent used to dehydrate ethanol
(1mark) (iii) State the conditions for the reaction
(1mark)
(iv) Write the equation for the dehydration of ethanol.
(1mark)
14(a). Hydrogen can react with nitrogen in the presence of a catalyst to produce
ammonia on an industrial scale.
(i). State the sources of hydrogen and nitrogen which are used in the reaction
(1marks)
(ii). Name the catalyst used in the reaction.
(1mark)
(iii). Write equation for the reaction leading to the formation of ammonia. (1
marks)
(b). In order for maximum yield of ammonia to be achieved during the industrial
preparation of ammonia as stated in (a) the reaction is carried out at a low
temperature.
(i). Give a reason for carrying out the reaction at low temperature.
(1mark)
(ii). State one condition for the reaction other than low temperature and use of the
catalyst
(1mark)
(iii). Briefly describe and explain how the condition you have stated in (b)(ii) affects
the reaction.
(2marks)
(c). The ammonia obtained by the reaction in (a) above can be oxidized in the
presence of a catalyst to manufacture nitric acid.
(i). Name the catalyst which is used in the reaction.
(1mark)
(ii). State two conditions for the oxidation of ammonia during the manufacture of
nitric acid other than the catalyst which you have named in (c) (i).
(1mark)
(iii). Using equation only show how nitric acid can be produced by the catalytic
oxidation of ammonia.
(3marks)
(d). Name one ammonium compound that is commonly used in agricultural industry
and state the purpose for which the compound is used in agriculture.
(2mark)
(b) 1
3Mg(s) + N2(g) Mg3N2(s) (1)
(c) 1
Pollutant; carbon dioxide ()
Effect ; global warming () or EQUIVALENT
TOTAL 5
2.(a) (i) A green precipitate() is formed
(b) (i) the green solid dissolves () with effervescence () to form a green/blue
solution ()
TOTAL 5
3(a) (i) W (1)
(iii) Difference boiling points(1)
TOTAL 5
3
4.(a) (i) 24,000 cm of a gas contain 1mole 2
96 cm3 of a gas contain (1 x 96)/ 24,000 = 0.004 moles ()
0.004 moles are contained in 0.224 g of alkene ()
1 mole will contain 0.224/0.004 () = 56 ()
(iii) CnH2n = 56
14n = 56 n=4
C4H8 ()
1
(b)
TOTAL 5
5(a) Reagents ; sodium hydroxide solution () and carbon dioxide () 1
(ii) Iron (III) chloride undergoes hydrolysis forming hydrochloric acid () which
decomposes () the hydrogen carbonate into carbon dioxide; the colourless
1
gas()
TOTAL 5
6.(a) (i) the most reactive metal is R () 1
(ii) the most reactive non-metal Y ()
(iii) A monoatomic gaseous element Z ()
(b) 2
Z, W, Q, Y, X, T, R (1)
(c) 1
TOTAL 4
7(a) (i) oxidation is loss of electrons ()
(i) The conversion of hydrogen ions (H+) to hydrogen molecules (H2)
(b) 2H+(aq) + 2e H2(g) (1) reduction ()
(ii (ii) The conversion of iron (II) ions (Fe2+) to iron (III) ions (Fe3+)
Fe2+(aq) Fe3+(aq)+e (1) oxidation ()
TOTAL 5
9.(a) chloride and iodide ions. 2
(i) reagent; lead (II) nitrate solution (1)
(ii) Observation
Chloride white precipitate()
Iodide yellow precipitate()
(b) o Vegetable oil and dilute sodium hydroxide (1)are boiled()together in a beaker
using a low Bunsen flame()
o The mixture is stirred continuously () 5
o The mixture is shaken with water to see whether it forms lather. ()
o When the mixture begins to form lather, a saturated solution of sodium chloride (1)
is added to precipitate () the soap.(separate the soap from the reaction mixture)
o The soap is removed from the reaction mixture (), additives added and shaped.
TOTAL 15
12(a) o Concentrated sulphuric acid (1) is placed in a funnel since it is a liquid.
o Solid () sodium chloride is placed in a flask (and the mixture is warmed gently.)
o The gas is passed through a wash bottle containing concentrated sulphuric acid ()
to dry it ()
o The gas is denser than air () and is collected by down delivery. ()
o H2SO4 (aq) + NaCl(s) NaHSO4(s) + HCl(g) (1)
TOTAL 15
14(a) (i) Nitrogen is obtained from air () by fractional distillation of liquid air. Hydrogen 4
is obtained from water gas(1)
(ii) Catalyst used is finely ground iron (1)
TOTAL 15
545/2
Chemistry
Paper 2
June/July 2015
2 hours.
INSTRUCTIONS TO CANDIDATES:
Section A consists of 10 structured questions. Attempt all questions in this
section. Answers to these questions must be written in the spaces provided.
Section B consists of 4 semi-structured questions. Attempt any two questions
from this section. Answers to the questions must be written in the answer
booklets provided.
In both sections, all working must be shown clearly.
1 2 3 4 5 6 7 8 9 10 11 12 13 14 Total
2(a) Complete the equations below, showing the effect of heat. (4marks)
(i) NaNO3(s)
(ii) AgNO3 (s)
(iii) Ca(NO3)2 (s)
(b) The equation below shows the reaction that should take place if potassium
carbonate was heated.
The basic strength of the oxides of the elements shown in the Periodic
Table increase in the order
Na2O CaOMgO Al2O3 SiO2 P2O3 SO2
(a) State the trend in the metallic character of the elements
(i) Across the Periodic Table. (1mark)
.
(ii) Down the group (1mark)
(b)(i) Identify the oxide which is amphoteric. (1mark)
(ii) Write equation for the reaction of the oxide you have identified in (b)(i) with
nitric acid. (1mark)
5.(a) Write equation for the reaction that would take place if, over strongly
heated iron filings was passed a current of
(i) dry hydrogen chloride (1mark)
(iii) dry chlorine gas (1mark)
..
(b)(i) State what would be done in order to convert the product in reaction (a)(i)
in its aqueous solution to the product in reaction (a)(ii) (mark)
(iii) Write equation to illustrate your answer in (b)(i) (1mark)
..
(iii) Name one reagent that would be used to confirm that the reaction in (a)(ii)
had taken place. (mark)
State what would be observed if the reagent you have named in (b)(iii) was
treated with the product of the reaction in (b)(ii) (1mark)
6. The diagram below shows the set up of apparatus that was used to make
an electrochemical cell for comparing the reactivities of copper and silver.
(a) Draw an arrow on the diagram to show the direction in which electrons are
flowing.(mark)
(b)(i) Name the substance X(mark)
(iii) Identify on e substance that was used to make the electrolyte containing
silver ions.
(1mark)
(c)write equations for the reaction taking place at
(i) The copper electrode
(1mark)
.
(ii) The electrode labeled X
(1mark)
.
7.(a) Name one reagent that can be used to distinguish between;
(i) Zn2+(aq) andPb2+(aq) (mark)
.
(ii) SO42-(aq) andCO32-(aq)
(1mark)
.
(b) State what would be observed if the reagent you have named in (a) was used to
separately treat (1mark)
2+ 2+
(i) Zn (aq) and Pb (aq)
(2mark)
(b) A solution was made by transferring exactly 20.0 cm 3 of 1M sulphuric acid into a
250cm3 volumetric flask, then carefully adding distilled water to bring the volume of
thenew solution to the mark. Determine the molar concentration of the new solution
with respect to hydrogen ions. (3marks)
10. When 40.0 cm3 of a 2M hydrochloric acid was mixed with 40.0 cm3 of a
2M sodium hydroxide both at initial temperature of 25.80C the temperature
of
the solution rose to t0C. Determine t
(specific heat capacity of water =4.2Jg-1 0C-1, density of water=1g cm-1 and
enthalpy of neutralization of hydrochloric acid by sodium hydroxide
=56.5kJ mol-1)
(4marks)
(b)(i) State two reactions which can be used to show that water contains hydrogen.
(2marks)
(iii) Write the equation to illustrate your statements in (b)(i)
(3marks)
(iv) Write an ionic equation for the reaction leading to the production of hydrogen
starting from zinc metal. (1marks)
(d) State what would be observed and write equation for the reaction that would take
place if
(i) Dry hydrogen was passes over heated copper (II) oxide (3marks)
(iii) Burning sodium was lowered into a gas jar containing pure dry hydrogen.
(2marks)
12(a) Describe how a sample of dry chlorine can be prepared in the laboratory starting
from potassium manganate (VII). (No diagram is required, but youranswer should
include conditions, equation, the drying agent and method of collection).
(5marks)
(b) Moist blue litmus paper was introduced into a test tube containing dry chlorine
State
(i) what was observed. (1mark)
(ii) the industrial application of the reaction in (b) (i)
(1mark)
(c) Write the equations for the reaction that would take place if chlorine was bubbled
through
(i) a dilute solution of potassium hydroxide. (1marks)
(ii) aqueous sodium bromide. (1marks)
(d) A colourless gas, T bubbles out of aqueous solution of chlorine that was left
exposed to sunlight.
(i) Identify T
(1mark)
(ii) Explain how T is produced and write equation to illustrate your explanation
(2marks)
(e) Write an equation to show the reaction that takes place when phosphorous
is
treated with chlorine.
(1mark)
13(a)(i) State how sulphur dioxide which is used for the production of sulphuric
acid
is obtained. (
mark)
(ii) Write the equation for the reaction that leads to the formation of sulphur
dioxide described in (a) (i) (1mark)
(iii) Name one reagent that can be used to confirm the presence of sulphur
dioxide and state what would be observed if sulphur dioxide was treated
with the reagent.
(2marks)
(iv) Give a reason for your observation in (a) (iii)
(1mark)
(d) Ethene can undergo a reaction leading tom the formation of polyethene.
(i) Name the reaction that leads to the formation of polyethene from ethene
and
statewhat the reaction you have named means in general terms. (1marks)
(ii) Write the equation for the reaction of ethene to produce polyethene
(1mark)
(1marks)
(f) Some other polymers exist, which are not synthetically produced.
(i) state what such polymers are called. (
mark)
(ii) Give one example of such polymers that you have stated in (f)(i) and give
one
use of it. (1mark)
14. Dilute nitric acid reacts with marble chips at room temperature according to
the equation
CaCO3(s) + 2HNO3 (aq) Ca(NO3)2(aq)+CO2(g)+H2O(l)
The of this reaction can be affected by a number of factors
(b) Exlain in each case the effect of the action taken in (a) on the rate of reaction
as stated by you (4marks)
(c) The table below shows time in minutes for the complete reaction when
various
volumes of 2M nitric acid were added separately to equal mass of marble
chips contained in six different beakers.
Volume of 2M nitric acid cm3 35.0 20.0 10.0 6.0 3.5 2.5
Time of completion of reaction 0 10 20 30 40 50
(min)
(i) Plot a graph of volume of nitric acid (vertical) against time (horizontal axis)
(4marks)
(ii) Using your graph, describe the effect of concentration of nitric acid on the
rate of reaction.
(4marks)
(iii) State how the graph you have plotted in (c)(i) can be use to determine the
rate of reaction at a given time
(2marks)
END
Total 5
2(a) (i) 2NaNO3(s) 2NaNO2(s )+ O2(g) (1)
(ii) 2AgNO3(s)2Ag2(s)+ O2(g) + 2NO2(g) (1)2Ca(NO3)2(s)2CaO(s)+ O2(g)
+ 4NO2(g) (1)
(b) (i) The reaction does not take place or The reaction is hypothetical(1)
(ii) Potassium carbonate is stable to heat. ()
Total 6
3(a) (i) Oxygen jar(1)
(ii) Nitrogen contains a triple bond that requires more heat to break. ()It
takes more time to break compared to oxygen double bond ()
(c) The universal indicator turned blue in each of the jars. (1)
Total 6
4(a) (i) Metallic character decreases across the period (1)
(ii) Metallic character increases across the period (1)
Total 4
5(a) (i) Fe(s)+2HCl(aq) FeCl2(aq)+H2 (g)(1)
(ii) 2Fe(s)+3Cl2(aq) 2FeCl3(aq)(1)
(b) In 20cm3of sulphuric acid there are (20 x 1)/1000 =2.0 x 10 -2()
In 250cm3 of diluted solution thee are 2.0 x 10-2()
In 1000 cm3 of diluted solution thee are [ 2.0 x 10-2 x 1000]/250() =0.08()
No of moles of hydrogen ions 0.08 x2 = 0.16()
+
The new solution is 0.16M H ions()
Total 4
9(a) (i) X =17 2: 8: 7(1)
(ii) Y = 20 2: 8:8:2 (1)
13.5 = t -25.8()
t =39.3()
Total 15
14(a) (i) Marble chips ground into a powder ; the rate of reaction increases() 1
(ii) Nitric acid diluted; the rate of reaction decreases ()
(b) Ground marble chips have a large () surface area (). This provides a greater
chance ()for more particles to participate in the reaction()
Dilute nitric acid; there are fewer hydrogen ions () per unit volume of the acid
solution. The average distance between the hydrogen ions and carbonate ions in 4
the marble chips increases (). The number () and frequency of collisions
()between the hydrogen and carbonate ions decrease resulting in a slower rate of
reaction
(c) Axes (1) scale (1) shape (1) accurate plot (1) 4
INSTRUCTIONS TO CANDIDATES:
Section A consists of 10 structured questions. Attempt all questions in this
section. Answers to these questions must be written in the spaces provided.
Section B consists of 4 semi-structured questions. Attempt any two questions
from this section. Answers to the questions must be written in the answer
booklets provided.
In both sections, all working must be shown clearly.
1 2 3 4 5 6 7 8 9 10 11 12 13 14 Total
(N =14, Cl=35.5, H = 1)
(ii) cathode. (1mark)
3.(a) A hydrocarbon Y contains 85.7% carbon. Calculate the simplest formula.
(2marks)
......................................................................................................................
.....................................................................................................................
.....................................................................................................................
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..................................................................................................................
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.....................................................................................................................
.................................................................................................................
.....................................................................................................................
.....................................................................................................................
.....................................................................................................................
..................................................................................................................
....................................................................................................................
.....................................................................................................................
B
(ii) Write an equation for the reaction (1marks)
(b) State what is observed in the test tube containing bromine water. (1mark)
5(a) Write equation for the reaction that would take place if, strongly heated
iron filings was passed a current of
(i) dry hydrogen chloride (1mark)
(ii)dry chlorine gas (1mark)
(b)(i) State what would be done in order to convert the product in reaction (a)(i)
in its aqueous solution to the product in reaction (a)(ii) (mark)
(iv) Write equation to illustrate your answer in (b)(i) (1mark)
(iv) Name one reagent that would be used to confirm that the reaction in (a)(ii)
had taken place. (mark)
State what would be observed if the reagent you have named in (b)(iii) was
treated with the product of the reaction in (b)(ii) (1mark)
6. The following chart below shows the contact process that is used to
manufacture sulphuric acid.
(ii) Name gas X and briefly describe how the gas can be identified in the
laboratory (1mark)
(b) Carbon dioxide was bubbled into the solution in the beaker.
(i) State what was observed ( mark)
(a) Name the particles that are responsible for the electrical conductivity of
Substance A ...................................................................................... (1 mark)
.................................................................................................................................
.................................................................................................................................
SECTION B
Answertwoquestions only in this section. Start each question on a fresh page.
(d) When a mixture of magnesium powder and copper (II) oxide are heated
together a mixture of a white and brown solid are formed. (5marks)
(b). State what would be observed and in each case, write an equation for
the reaction. (5marks)
(i) sodium burnt in excess oxygen.
(ii) phosphorous in excess oxygen
MAKING GUIDE 2013
Silicon (IV) oxide is acidic (); will not dissolve in the acidic (). Zinc oxide is
(b) amphoteric (), it acts as a base and dissolves in the acid() 2
TOTAL 5
2 (a) (i) X 2.8.2 (1)
(ii)
(i) anode
(b)
2H-(l) H2(g) + 2e (1)
(ii) cathode
X2+ (l) +2e X(l) (1)
TOTAL 5
3 % of hydrogen = 100 - 85.7 = 14.3 ()
Element C H
% 85.7 14.3
(ii) (CH2)n = 56
14n = 56
n = 56/14 = 4 ()
The formula is C4H8 ()
TOTAL 5
4(a) (i)A is sulphuric acid () B is ethene ()
(b) (i) The bromine water turns from brown ()to colourless ()
(ii) CH2 =CH2 + Br2CH2BrCH2Br(1)
TOTAL 5
5(a) (iii) Fe(s)+2HCl(aq) FeCl2(aq)+H2 (g) (1)
(iv) 2Fe(s)+3Cl2(aq) 2FeCl3(aq)(1)
TOTAL 6
6(a) X is air ()
Y is 98% sulphuric acid ()
TOTAL 5
8(a) (i) bubbles() of a colourless gas()
(ii) The gas is hydrogen (1) a burning splint ()is introduced into the test tube, the
gas explodes() or pops
TOTAL 5
9.(a) An indicator is substance that changes colour according to the hydrogen ion
concentration or hydroxide concentration. (1)
TOTAL 15
12(a Sodium metal has delocalized electrons (1)that are responsible for carrying the 5
) current.
In the solid state sodium ions () and chloride ions () are firmly held by strong
electrostatic forces (1) therefore t are not mobile () hence inability to conduct
electricity.
In the molten state heat () breaks () the strong electrostatic forces allowing
the ions to move () and hence conduction of electricity.
Hard water contains calcium () and magnesium () ions. On addition of soap, scum
(b) (1) is formed which is an insoluble salt (1)of magnesium or calcium stearateas
follows:
Ca2+(aq) + 2St-(aq) Ca St2(s)(1) 5
It is only when all the calcium and magnesium ions have been removed() from
water that a lather will form() thus wastage of soap
(c) Sulphuric acid is an acid,(1) ammonia is alkaline,(1) and the two react () forming a
salt.(1)
H2SO4(aq) + 2NH3(g) (NH4)2SO4(aq)(1)
5
(d) Magnesium is higher in the reactivity series than copper() it reduces (1)
copper(II)oxide to copper(),; the brown solid and itself is oxidised(1) to
magnesium oxide(),; the white solid.
Equation: CuO(g) + Mg(s) Cu (s) + MgO(s) (1) 5
TOTAL 15
13(a (i) platinum-rhodium (1). 5
0
) (ii) temperature 900 C () , excess oxygen (), slight high pressure
(iii) 4NH3(g) + 5 O2(g) 4NO(g)+ 6H2O(l) (1)
(i) Observations: a brown gas was given off (). The copper turnings dissolved
(b) forming a blue solution ()
4
(ii) Explanation: The acid is reduces () to nitrogen (IV) oxide,()
the brown gas and copper is oxidized() to copper (II) nitrate (), the blue
solution.
TOTAL 15
14(a (i) hydrogen peroxide; room temperature ()and a catalyst() 4
)
(ii) (ii) potassium nitrate; heat()
(iii) Equation:
2H2O2(aq) 2H2O (l)) + O2(l) (1)
2KNO3(s) 2KNO2(s) + O2(l) (1)
(i) The rate of formation of oxygen is the increase in volume of oxygen per unit
time. (2)
(b)
(ii) adding a catalyst e.g. manganese (IV) oxide (1)
Increasing the concentration of hydrogen peroxide (1) 5
Exposing the mixture to sunlight. (1)
(iii) photosynthesis()
Sodium first melts and then burns with bright yellow flame ()forming a yellow
solid()
2Na(s) + O2(g) Na2O2(g) (1)
Phosphorous burns with yellow flame () and forms dense white fumes() 5
P4(s) + 5O2(g) 2P2O5(s) (1)
TOTAL 15
Name..Centre/Index No
Signature.
545/2
Chemistry
Paper 2
August 2015
Instructions:
This paper consists of two sections A and B
Attempt all questions in section A
Answers to section A must be written in the spaces provide only
Attempt two questions in section B
Marks
(b) Manganese (IV) oxide was added to hydrogen peroxide and the mixture
was then exposed to sunlight
(i) State the rate of gas bubbling in this mixture would differ from that in (a)
(1mark)
..
(ii) Give reasons for your answer in (b) (i) (1 mark)
c) Other than the use of Manganese (IV) oxide, suggest one other thing
that could be done to enhance change in the rate of gas bubbling by the
hydrogen peroxide. (1 marks)
14. a) Write an equation for the ionization of sodium carbonate Na2CO3 in water
b) Water was added to 250cm3 of a 0.4M sodium carbonate solution to make
1 litre of a diluted solution. Calculate the concentration of sodium ions in
moles per litre of the diluted solution. (3 marks)
5. a) Excess carbon monoxide gas was passed over a heated sample of an
oxide of iron as shown in the diagram below. Study the diagram and the
data below to answer the questions that follow.
(v) Write an equation for the reaction which took place in the dish (1 marks)
(b) Rusting is a destructive process in which iron is converted into hydrated
iron (iii) oxide.
(v) State two conditions necessary for rusting to occur
(1
mark)
One method to protect iron from rusting (1 mark)
7. The elements: Aluminium, sulphur and calcium can each combine with
oxygen to form oxides.
a) State the class of the oxide of
(i) Aluminium ( mark)
.
(ii) Sulphur ( mark)
.
(iii) Calcium ( mark)
(b) (i) State what would be observed if a few drops of water were added to the
oxide of calcium. (1 marks)
(ii) Write the equation that takes place when water is added to the oxide of
calcium. (1 mark)
(c) Name one other oxide in the same class as the oxide of aluminium.
(1
marks)
8. The diagram below shows sketches of curves that were obtained in three
experiments to study the rates of reactions of dilute sulphuric acid and iron
metal.
mass of
flask and its
contents
Y Z
X
M
Time (minutes)
10. 1.9g of magnesium chloride were dissolved in distilled water. Silver nitrate
solution was added until in excess.
(a) Write equation for the reaction that took place. (1 marks)
(b) Calculate the mass of silver nitrate that was used for complete reaction.
(RFM of MgCl2 = 95, N = 14.0, O = 16.0, Ag = 108.0) (2 marks)
SECTION B
11. Explain the following observation. (Where necessary illustrate your answer
with equations)
(a) When heated, a mixture of ethanol and concentrated sulphuric acid
produces a gas that decolorizes bromine water. (4 marks)
(b) Sodium chloride in the solid state does not conduct electricity but it will
when molten and is decomposed by electric current. (6 marks)
12. Draw a labeled diagram for the set up of apparatus that can be used to
electrolyze lead (II) bromide. (3 marks)
(b) Describe the reactions that take place during the electrolysis of copper (ii)
sulphate solution using
(i) Graphite electrodes (6 marks)
(ii) Copper electrodes (4 marks)
(c) State one industrial application of electrolysis other than the purification of
copper. (1 mark)
(b) The solubility of a salt X at various temperatures are shown in the table
below.
(d) State
(i) How you would distinguish between soft water and hard water (1 marks)
(ii) One use of hard water (1 marks)
(e) Name one compound which when dissolved in water will cause;
(i) Permanent hardness of water
(iii) Temporary hardness of water
14( a)( i) Name the raw material from which sulphuric acid can be manufactured
by the contact process (1 mark)
(ii) With the help of equations, outline the reactions which take place during
the contact process. (5
marks)
(b) Explain why fuming sulphuric acid has no effect on litmus paper whereas
dilute sulphuric acid readily turns blue litmus paper red. (3 marks)
(c) State what would be observed and write equation for the reaction that
would take place when concentrated sulphuric acid was added to;
(i) Iron (ii) sulphate -7- water (3 marks)
(ii) Potassium chloride (3 marks)
END
Name..Centre/Index No
Signature.
545/2
Chemistry
Paper 2
August 2011
CHEMISTRY
PAPER 2
2 HOURS
Instructions:
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When 9.28 of the salt was heated, 4.09g of the residue was
obtained. Calculate the number of moles of water of crystallization in
the hydrated salt
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2. (a) Write equation for the reaction that would take place if burning
magnesium was lowered into a gas jar of;
i) Carbon dioxide (1
marks)
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ii) Nitrogen (1
marks)
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ii) Write equation for the reaction that would take place if a few
drops of water were added to the product in (a) (ii) (1
marks)
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3. Impure alcohol can be prepared from a solution of glucose,
C6H12O6 mixed with yeast.
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(1 mark)
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b) State
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ii) One method that can be used to test for the purity of the
alcohol(1mark)
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i) Q (1 mark)
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ii) R (1 mark)
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b) State
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ii) The type of bond in the compound that can be formed when Q
reacts with R. (1
marks)
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6. a) State on use of sodium hydroxide in
i) A chemistry laboratory (1
marks)
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ii) An industry (1
marks)
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ii) Write an ionic equation for the reaction that took place.
(1 marks)
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(
mark)
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ii) State what would be observed if the reagent you have named
in (c) (i) was treated with carbonate ion
( mark)
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a) State
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b) Write equation for the reaction that takes place when ethene is
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i) Hydrochloric acid (
mark)
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b) Describe briefly how the products in (a) (i) and (ii) can be
distinguished.
(1
marks)
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(1 marks)
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10. a) Copper (II) carbonate was heated strongly until there was
no further change.
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b) When heated, zinc nitrate decomposes to give nitrogen dioxide
and oxygen according to the following equation
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State the condition under which the electrolysis was carried out
( mark)
calcium hydroxide by nitric acid and state the units for your answer.
(4 marks)
(Specify heat capacity of the solution and its density are 4.2Jg -1oC-
1
and 1.0gcm-3 respectively)
13. a) i) Name the raw material for the manufacture of sulphuric acid
by the contact process. (
mark)
ii) Outline with aid of equations the reactions which take place
during the contact process. (6
marks)
(Your answer should include equations for the reaction that takes
place; but No diagram is required)
(5 marks)
d) Dry ammonia was passed over hot lead (II) oxide. Write the
equation for the reaction which took place.
(1 marks)