Chemistry Notes Olevel
Chemistry Notes Olevel
Chemistry Notes Olevel
Contents
1. Introduction to chemistry .................................................................................................................... 15
1.1. Definition of chemistry ............................................................................................................... 15
1.2. Applications of chemistry ........................................................................................................... 15
1.3. A laboratory ................................................................................................................................ 15
1.4. Uses of a laboratory .................................................................................................................... 15
1.5. Precautions/ safety measures/rules and regulations taken in the laboratory ............................... 15
1.6. Apparatus used in the laboratory ................................................................................................ 16
1.7. Apparatus and their uses ............................................................................................................. 19
1.7.1. The Bunsen burner .................................................................................................................. 19
1.7.2. Flames of the Bunsen burner .................................................................................................. 19
1.7.2.1. Luminous flame .................................................................................................................. 20
1.7.2.2. Characteristics of a luminous flame .................................................................................... 20
1.7.2.3. Non luminous flame ............................................................................................................ 20
1.7.2.4. Characteristics of a non luminous flame ............................................................................. 20
1.7.2.5. Parts of a non luminous flame ............................................................................................. 21
1.7.2.6. Differences between luminous and non-luminous flame .................................................... 21
1.7.2.7. Steps followed when lighting a Bunsen burner................................................................... 21
2. States of matter ................................................................................................................................... 21
2.1. Solids............................................................................................................................................... 21
2.2. Liquids ............................................................................................................................................ 22
2.3. Gases ............................................................................................................................................... 22
2.4. Change of state ................................................................................................................................ 23
2.5. The Kinetic Particle Theory of Matter ............................................................................................ 26
2.6. Diffusion ......................................................................................................................................... 27
2.6.1. Diffusion of gases ....................................................................................................................... 27
2.6.2. Diffusion of liquids ..................................................................................................................... 27
2.6.3. Factors Affecting Rate of Diffusion............................................................................................ 28
2.6.3.1. Temperature ............................................................................................................................ 28
2.6.3.2. Mass of particles ..................................................................................................................... 28
3. Chemical and physical changes .......................................................................................................... 29
3.1. Differences between physical and chemical changes ..................................................................... 29
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4. Solutions, crystals, compounds and mixtures ..................................................................................... 30
4.1. Solutions ......................................................................................................................................... 30
4.2. Solute .............................................................................................................................................. 30
4.3. Solvent ............................................................................................................................................ 30
4.4. Types of solutions ........................................................................................................................... 30
4.4.1. A saturated solution .................................................................................................................... 30
4.4.2. A super saturated solution ........................................................................................................... 30
4.4.3. A suspension ............................................................................................................................... 30
4.5. Differences between solutions and suspensions ............................................................................. 30
4.6. Crystals ........................................................................................................................................... 31
4.7. How to grow a large crystal of copper (II) sulphate ....................................................................... 31
4.8. Water of crystallization ................................................................................................................... 31
5. Compounds and mixtures.................................................................................................................... 32
5.1. Compound ....................................................................................................................................... 32
5.2. Mixture............................................................................................................................................ 32
5.3. Differences between mixtures and compounds............................................................................... 32
5.4. Methods of separation of mixtures.................................................................................................. 33
5.5. The methods of separation of mixtures include the following ........................................................ 33
5.5.1. Filtration...................................................................................................................................... 33
5.5.2. Crystallisation & Evaporation to Dryness................................................................................... 34
5.5.3. Evaporation to dryness ................................................................................................................ 34
5.5.4. Decanting .................................................................................................................................... 35
5.5.5. Using separating funnel .............................................................................................................. 36
5.5.6. Distillation................................................................................................................................... 37
5.5.6.1. Simple Distillation .................................................................................................................. 37
5.5.6.2. Separation of a mixture of water and ethanol by distillation .................................................. 38
5.5.6.3. Fractional distillation .............................................................................................................. 38
5.5.7. Sublimation ................................................................................................................................. 39
5.5.8. Magnetic Attraction .................................................................................................................... 40
5.5.9. Chromatography ......................................................................................................................... 42
6. Elements, compounds, atoms and symbols ......................................................................................... 43
6.1. An element ...................................................................................................................................... 43
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6.2. An atom........................................................................................................................................... 43
6.3. A molecule ...................................................................................................................................... 43
6.4. Radicals ........................................................................................................................................... 43
6.5. Chemical symbols ........................................................................................................................... 44
6.6. Metals and non-metals .................................................................................................................... 45
6.7. Properties of metals and non-metals ............................................................................................... 46
7. Atomic structure and the periodic table .............................................................................................. 47
7.1. An atom........................................................................................................................................... 47
7.2. Composition of an atom .................................................................................................................. 47
7.3. Structure of an atom ........................................................................................................................ 48
7.4. Atomic number ............................................................................................................................... 48
7.5. Atomic mass.................................................................................................................................... 49
7.6. Electronic configuration .................................................................................................................. 49
7.7. Electronic configuration of ions ...................................................................................................... 52
7.8. Isotopes ........................................................................................................................................... 53
8. The periodic table ............................................................................................................................... 54
8.1. Chemical families ........................................................................................................................... 56
8.2. Bonding ........................................................................................................................................... 63
8.3. Ionic/Electrovalent Bonding ........................................................................................................... 63
8.4. Covalent Bonding ........................................................................................................................... 65
8.5. Dative bonding ................................................................................................................................ 68
8.6. Metallic Bonding ............................................................................................................................ 68
9. Valency ............................................................................................................................................... 70
9.1. Elements and radicals with their valencies ..................................................................................... 70
9.2. Radicals ........................................................................................................................................... 71
10. Chemical formulae .......................................................................................................................... 72
10.1. Writing chemical formulae ......................................................................................................... 72
10.2. Calculating the number of atoms of elements in a compound .................................................... 73
11. Chemical equations ......................................................................................................................... 76
11.1. Balanced chemical equations ...................................................................................................... 76
11.2. Balancing chemical equations ..................................................................................................... 76
12. Types of chemical reactions ............................................................................................................ 81
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12.1. Direct combination or direct synthesis........................................................................................ 81
12.2. Simple decomposition ................................................................................................................. 81
12.3. Simple replacement ..................................................................................................................... 81
12.4. Double replacement .................................................................................................................... 82
13. The atmosphere and combustion..................................................................................................... 83
13.1. Air is a mixture of gases ............................................................................................................. 83
13.1.1. Oxygen ........................................................................................................................................ 83
13.1.2. Uses of oxygen ............................................................................................................................ 83
13.1.3. Nitrogen ...................................................................................................................................... 83
13.1.4. Carbon dioxide ............................................................................................................................ 84
13.1.5. Noble gases ................................................................................................................................. 84
13.1.6. Uses of noble gases ..................................................................................................................... 84
13.1.7. Water vapour ............................................................................................................................... 84
13.2. Hygroscopic, deliquescent and efflorescent substances .............................................................. 84
13.2.1. Hygroscopic substances .............................................................................................................. 84
13.2.2. Deliquescent substances.............................................................................................................. 85
13.2.3. Efflorescent substances ............................................................................................................... 85
13.2.4. Water of crystallisation ............................................................................................................... 85
13.3. Drying agents .............................................................................................................................. 85
13.4. Burning substances in air ............................................................................................................ 86
13.6. Rusting ........................................................................................................................................ 89
13.7. Combustion ................................................................................................................................. 93
14. Methods of gas collection ............................................................................................................... 94
15. Oxygen ............................................................................................................................................ 96
15.1. Oxides ......................................................................................................................................... 98
16. Oxidation and reduction ................................................................................................................ 100
17. Water and hydrogen ...................................................................................................................... 102
17.1. Sources of water ........................................................................................................................ 102
17.2. Properties of pure water ............................................................................................................ 103
17.3. Test for water ............................................................................................................................ 103
17.4. Purification of water ................................................................................................................. 103
17.5. Reactions of metals with water ................................................................................................. 104
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17.6. Uses of water............................................................................................................................. 106
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
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................................................................................................................................... 119
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
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19.7.4. Preparation of Insoluble Salts ................................................................................................... 125
19.7.5. Preparation of Soluble Salts ...................................................................................................... 127
19.7.6. Action of heat on salts............................................................................................................... 131
19.8. Solubility of salts ...................................................................................................................... 133
19.8.1. Solubility curves ....................................................................................................................... 134
19.8.2. Uses of solubility ...................................................................................................................... 134
19.9. Determination of solubility of salts ........................................................................................... 135
20. Carbon and its compounds ............................................................................................................ 138
20.1. Oxides of carbon ....................................................................................................................... 144
20.1.1. Carbon dioxide .......................................................................................................................... 144
20.1.2. Carbon monoxide ...................................................................................................................... 151
20.2. Carbonates and hydrogen carbonates ........................................................................................ 158
20.3. Calcium Oxide (quicklime) ....................................................................................................... 162
20.4. Sodium carbonate (soda ash) .................................................................................................... 164
20.5. The Carbon cycle ...................................................................................................................... 167
20.6. Hardness of water ..................................................................................................................... 171
21. Electrolysis.................................................................................................................................... 174
21.1. Laws of electrolysis .................................................................................................................. 190
21.2. Application of electrolysis ........................................................................................................ 195
22. Formulae, stoichiometry and the mole concept ............................................................................ 204
22.1. Relative Atomic Mass ............................................................................................................... 204
22.2. Percentage Composition ........................................................................................................... 206
22.3. Calculating the Mass of an Element in a Compound ................................................................ 206
22.4. Calculating the Mass of Water in a Compound ........................................................................ 207
22.5. Mole .......................................................................................................................................... 207
22.6. Molar Mass ............................................................................................................................... 208
22.7. Different Kinds of Chemical Formulae..................................................................................... 209
22.8. Calculating the Empirical Formula from Percentage Composition .......................................... 210
22.9. From Empirical formula to Molecular Formula........................................................................ 210
22.10. Molar Volume of Gases ............................................................................................................ 214
22.11. Calculations using chemical equations ..................................................................................... 215
22.11.1. Constructing Chemical Equations ......................................................................................... 215
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22.11.2. Calculations from Equations ................................................................................................. 216
22.11.2.1. Reacting Masses................................................................................................................ 216
22.11.2.2. Reacting Masses and Volumes ......................................................................................... 218
22.11.2.3. Calculations involving energy changes ............................................................................. 220
22.12. Concentration of Solutions........................................................................................................ 220
22.13. Quantitative analysis ................................................................................................................. 223
22.14. Uses of Titrations in Analysis ................................................................................................... 226
22.14.1. Identification of Acids and Alkalis ....................................................................................... 226
22.14.3. Determination of basicity of an acid ..................................................................................... 229
22.14.4. Determination of water of crystallisation in oxalic acid (COOH)2. xH2O ............................ 230
22.15. Calculations of volume of solutions.......................................................................................... 233
22.17. Gay Lussac‘s law ...................................................................................................................... 234
22.18. Gas laws .................................................................................................................................... 235
23. Qualitative analysis ....................................................................................................................... 237
24. Sulphur and its compounds ........................................................................................................... 246
24.1. Sulphur ...................................................................................................................................... 246
24.2. Extraction of sulphur by the Frasch‘s process .......................................................................... 246
24.3. Extraction of sulphur from natural gas ..................................................................................... 247
24.4. Uses of sulphur ......................................................................................................................... 247
24.5. Allotropes of sulphur ................................................................................................................ 247
24.5.1. Rhombic sulphur ....................................................................................................................... 247
24.5.2. Monoclinic sulphur ................................................................................................................... 248
24.5.3. Amorphous sulphur ................................................................................................................... 249
24.5.4. Plastic sulphur ........................................................................................................................... 249
24.5.5. Colloidal sulphur ....................................................................................................................... 249
24.6. Properties of sulphur ................................................................................................................. 249
24.6.1. Physical properties .................................................................................................................... 249
24.6.2. Chemical properties of sulphur ................................................................................................. 250
a. Action of heat on sulphur (in absence of air) .................................................................................... 250
b. Combustion of sulphur (in a plentiful supply of air)......................................................................... 250
c. Reaction with metals and non-metals ............................................................................................... 250
d. Action of acids on sulphur ................................................................................................................ 250
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24.7. Oxides of sulphur ...................................................................................................................... 251
24.7.1. Sulphur dioxide ......................................................................................................................... 251
24.7.1.1. Laboratory preparation of sulphur dioxide ........................................................................... 251
24.7.1.2. Properties of sulphur dioxide ................................................................................................ 252
24.7.1.3. Uses of sulphur dioxide......................................................................................................... 253
24.7.2. Sulphur trioxide ........................................................................................................................ 253
24.7.2.1. Preparation of sulphur trioxide ............................................................................................. 253
24.8. Sulphuric acid ........................................................................................................................... 254
24.8.1. Industrial manufacture of sulphuric acid by the contact process .............................................. 254
24.8.2. Properties of sulphuric acid....................................................................................................... 255
24.8.2.1. Physical properties ................................................................................................................ 255
24.8.2.2. Chemical properties .............................................................................................................. 255
24.8.3. Uses of sulphuric acid ............................................................................................................... 257
24.9. Sulphates ................................................................................................................................... 257
24.10. Hydrogen sulphide .................................................................................................................... 259
24.10.1. Laboratory preparation of hydrogen sulphide ....................................................................... 259
24.10.2. Testing for hydrogen sulphide .............................................................................................. 260
24.10.3. Properties of hydrogen sulphide ........................................................................................... 260
24.10.3.1. Physical properties ............................................................................................................ 260
24.10.3.2. Chemical properties .......................................................................................................... 260
25. Nitrogen and its compounds ......................................................................................................... 262
25.1. Nitrogen .................................................................................................................................... 262
25.2. Laboratory preparation of nitrogen from air ............................................................................. 262
25.3. Test for nitrogen ........................................................................................................................ 264
25.4. Properties of nitrogen ................................................................................................................ 264
25.5. Uses of nitrogen ........................................................................................................................ 265
25.6. Nitrogen monoxide (nitrogen(II) oxide) ................................................................................... 266
25.6.1. Laboratory preparation of nitrogen monoxide .......................................................................... 266
25.6.2. Tests for nitrogen monoxide ..................................................................................................... 266
25.6.3. Properties of nitrogen monoxide ............................................................................................... 267
25.7. Nitrogen dioxide ....................................................................................................................... 267
25.7.1. Laboratory preparation of nitrogen dioxide .............................................................................. 267
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25.7.2. Properties of nitrogen dioxide ................................................................................................... 268
25.8. Ammonia................................................................................................................................... 269
25.8.1. Laboratory preparation of ammonia.......................................................................................... 269
25.8.2. Industrial preparation of ammonia (Haber process).................................................................. 270
25.8.3. Tests for ammonia..................................................................................................................... 270
25.8.4. Properties of ammonia .............................................................................................................. 271
25.8.5. Solubility of ammonia in water ................................................................................................. 271
25.8.6. Experiment to demonstrate the high solubility of ammonia gas in water ................................. 271
25.8.7. Action of ammonia on copper (II) oxide .................................................................................. 272
25.8.8. Combustion of ammonia ........................................................................................................... 273
25.8.9. Reaction with hydrogen chloride .............................................................................................. 274
25.8.10. Reaction with chlorine .......................................................................................................... 274
25.8.11. Uses of ammonia................................................................................................................... 275
25.8.12. Ammonia solution ................................................................................................................. 276
25.8.12.1. Preparation of ammonia solution ...................................................................................... 276
25.8.13. Ammonium salts ................................................................................................................... 276
25.8.13.1. Nitrogenous fertilizers....................................................................................................... 276
25.8.13.2. Effect of heat on ammonium salts..................................................................................... 277
25.8.13.3. Test for ammonium salts ................................................................................................... 278
25.8.14. Reactions of ammonia solution and sodium hydroxide solution .......................................... 279
25.8.15. Nitric acid.............................................................................................................................. 281
25.8.15.1. Laboratory preparation of nitric acid ................................................................................ 281
25.8.15.2. Industrial preparation of nitric acid ................................................................................... 282
25.8.15.3. Uses of nitric acid ............................................................................................................. 282
25.8.15.4. Properties of nitric acid ..................................................................................................... 283
25.8.15.5. Nitric acid acting as a strong acid ..................................................................................... 283
25.8.15.6. Nitric acid as an oxidizing agent ....................................................................................... 284
25.9. Nitrates ...................................................................................................................................... 285
25.9.1. Action of heat on nitrates .......................................................................................................... 285
25.9.2. Test for nitrates ......................................................................................................................... 287
26. Chlorine and its compounds.......................................................................................................... 295
26.1. Chlorine..................................................................................................................................... 295
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26.1.1. Laboratory preparation of chlorine ........................................................................................... 295
26.1.2. Industrial manufacture of chlorine ............................................................................................ 296
26.1.3. Properties of chlorine ................................................................................................................ 297
26.1.4. Tests for chlorine ...................................................................................................................... 303
26.1.5. Uses of chlorine ........................................................................................................................ 303
26.2. Hydrogen chloride .................................................................................................................... 303
26.2.1. Laboratory preparation of hydrogen chloride ........................................................................... 303
26.2.2. Test for hydrogen chloride ........................................................................................................ 304
26.2.3. Properties of hydrogen chloride ................................................................................................ 304
26.3. Hydrochloric acid...................................................................................................................... 304
26.3.1. Preparation of hydrochloric acid ............................................................................................... 304
26.3.2. Properties of hydrochloric acid ................................................................................................. 305
26.3.3. Uses of hydrochloric acid ......................................................................................................... 306
26.3.4. Properties of hydrogen chloride in methylbenzene ................................................................... 306
26.4. Testing for soluble chloride ...................................................................................................... 306
27. Extraction of metals ...................................................................................................................... 311
27.1. Introduction ............................................................................................................................... 311
27.2. Concentration of ores ................................................................................................................ 311
27.3. Sodium ...................................................................................................................................... 312
27.3.1. Extraction of sodium ................................................................................................................. 312
27.3.2. Uses of sodium metal ................................................................................................................ 313
27.4. Copper ....................................................................................................................................... 313
27.4.1. Extraction of copper .................................................................................................................. 314
27.4.2. Concentration of the ore............................................................................................................ 314
27.4.3. Roasting and reduction.............................................................................................................. 314
27.4.4. Refining of the impure copper .................................................................................................. 315
27.4.5. Uses of copper........................................................................................................................... 316
27.5. Iron ............................................................................................................................................ 316
27.5.1. Extraction of iron ...................................................................................................................... 316
27.5.2. Casting iron (pig-iron) .............................................................................................................. 317
27.5.3. Wrought iron ............................................................................................................................. 317
27.5.4. Steel........................................................................................................................................... 318
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27.5.5. Recycling of metals................................................................................................................... 319
27.5.6. Alloy ......................................................................................................................................... 319
28. Organic chemistry ......................................................................................................................... 322
28.1. Introduction ............................................................................................................................... 322
28.2. Hydrocarbons ............................................................................................................................ 322
28.3. Homologous series .................................................................................................................... 322
28.4. Functional groups...................................................................................................................... 322
28.5. Alkanes ..................................................................................................................................... 323
28.5.1. General properties of alkanes .................................................................................................... 328
28.5.1.1. Physical properties ................................................................................................................ 328
28.5.1.2. Chemical properties .............................................................................................................. 328
28.6. Petroleum (crude oil) ................................................................................................................ 331
28.6.1. Fractional distillation of petroleum ........................................................................................... 331
28.6.2. Cracking .................................................................................................................................... 333
28.6.3. Bio gas ...................................................................................................................................... 333
28.6.4. Disadvantages of bio gas production ........................................................................................ 334
28.6.5. Alkenes ..................................................................................................................................... 334
28.7. Ethene ....................................................................................................................................... 335
28.7.1. Laboratory preparation of ethene .............................................................................................. 335
28.7.1.1. Physical properties ................................................................................................................ 336
28.7.1.2. Chemical properties .............................................................................................................. 336
28.7.2. Uses of ethene ........................................................................................................................... 338
28.7.3. Alkynes ..................................................................................................................................... 339
28.7.4. Ethyne (acetylene) .................................................................................................................... 339
28.7.4.1. Physical properties of ethyne ................................................................................................ 339
28.7.4.2. Chemical properties of ethyne .............................................................................................. 340
28.7.4.3. Uses of ethyne ....................................................................................................................... 341
28.7.5. Alkanols (alcohol) ..................................................................................................................... 341
28.8. Ethanol (ethyl alcohol) .............................................................................................................. 342
28.8.1. Manufacture of ethanol ............................................................................................................. 342
28.8.2. Properties of ethanol ................................................................................................................. 343
28.8.2.1. Chemical properties .............................................................................................................. 343
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28.8.2.2. Uses of ethanol ...................................................................................................................... 344
28.9. Carboxylic acids and esters ....................................................................................................... 344
28.9.1. Properties of carboxylic acids ................................................................................................... 345
28.10. Esters ......................................................................................................................................... 346
28.10.1. Uses of esters ........................................................................................................................ 346
28.11. Soap........................................................................................................................................... 346
28.11.1. Manufacture of soap.............................................................................................................. 346
28.11.2. The cleansing action of soap ................................................................................................. 347
28.11.3. Soapless (synthetic detergents) ............................................................................................. 347
28.11.4. Advantages of synthetic detergents over soap ...................................................................... 348
28.11.5. Advantages of soap over synthetic detergents ...................................................................... 348
28.12. Polymerisation .......................................................................................................................... 348
28.13. Addition polymerization ........................................................................................................... 348
28.14. Polyethene ................................................................................................................................. 349
28.15. Polypropene .............................................................................................................................. 349
28.16. Polyvinyl chloride (P.V.C)........................................................................................................ 349
28.17. Synthetic rubber ........................................................................................................................ 350
28.18. Condensation polymerization ................................................................................................... 350
28.19. Types of polymers..................................................................................................................... 350
28.20. Advantages of synthetic polymers over natural polymers ........................................................ 352
28.21. Disadvantages of synthetic polymers ........................................................................................ 353
29. Energy changes ............................................................................................................................. 357
29.1. Introduction ............................................................................................................................... 357
29.2. Enthalpy .................................................................................................................................... 357
29.3. Exothermic and endothermic reactions ..................................................................................... 357
29.4. Types of enthalpy changes ........................................................................................................ 359
29.4.1. Determination of enthalpy (heat) of combustion of ethanol ..................................................... 360
29.5. Determination of the enthalpy (heat) of solution o f sodium chloride ...................................... 363
29.6. Enthalpy of neutralization ......................................................................................................... 364
29.7. Determination of heat of neutralization .................................................................................... 364
29.8. Enthalpy of displacement .......................................................................................................... 366
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29.8.1. Determination of the enthalpy (heat) of displacement of the reaction between copper(II)
sulphate solution and zinc ......................................................................................................................... 366
30. Rate of reaction and equilibrium ................................................................................................... 371
30.1. Rate of reaction ......................................................................................................................... 371
30.1.1. Determination of rate of reaction .............................................................................................. 371
30.1.2. Determination of rates of reaction by measuring the volume of the gas evolved with time ..... 371
30.1.3. Effects of temperature on the rate of reaction ........................................................................... 375
30.1.3.1. Investigation of the effect of temperature on the rate of reaction ......................................... 375
30.1.4. Effect of a catalyst on the rate of reaction................................................................................. 376
30.1.4.1. Investigation of the effect of catalyst on the rate of reaction ................................................ 376
30.1.5. Effect of surface area on the rate of reaction ............................................................................ 378
30.1.5.1. Investigation of the effect of surface area on the rate of reaction ......................................... 378
30.1.6. Effect of light on the rate of reaction ........................................................................................ 379
30.1.6.1. Investigation of the effect of light on the rate of reaction ..................................................... 379
30.1.7. Effect of pressure on the rate of reaction .................................................................................. 380
31. Equilibrium ................................................................................................................................... 380
31.1. Factors affecting equilibrium .................................................................................................... 380
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1. Introduction to chemistry
1.3. A laboratory
This is a specialized and organized room where scientific experiments are conducted
15
Place broken glasses and solids into bin always located in the laboratory
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
16
Litmus paper Litmus paper contains a
chemical that turns blue in
basic solutions and reddish-
pink in acidic solutions
Pipette For transferring or measuring
out small quantities of liquid
17
liquids)
Mortar and pestle Used for pounding and
grinding solid substances to
make them into smaller
particles for ease of handling
Beaker A beaker is a container used
to hold, mix, and pour liquid
or powdered chemicals.
18
Tripod stand A tripod is a three legged
frame for supporting a few
apparatus during heating
19
2. Non-luminous flame
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1.7.b.5. Parts of a non luminous flame
Inner zone: this is a zone of cool un-burnt gas
Middle zone: this is blue or green in colour, here some of the gas is burnt but not all because
there is no enough air
Pale blue/ purple zone: in this zone burning of the gas is complete
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
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- Fixed shape
- Incompressible
- Do not flow
2.2. Liquids
The characteristics of liquids include:
- Fixed volume
- 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
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- Flow in all direction
Solid
Freezing
Gases Liquids
Condensation
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.
23
Particles in solid: Particles in liquid: Particles in gas:
- Are packed close together in - Are packed closely but not - Are far apart and in random
orderly arrangement orderly arranged arrangement
- Have little empty space - Have little empty space - Are free to move anywhere
between them between them but more than in the container
- Can vibrate but cannot move in solids
freely about their fixed - Are not held fixed but free to
position move throughout liquid
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.
24
From the graph:
A-B: the temperature of solid increases to melting point.
B-C: the temperature remains constant as heat is absorbed to break forces of attraction instead
for raising temperature. Solid and liquid are present.
C-D: liquid heats as heat energy increases temperature.
D-E: liquid temperature rises to boiling point.
E: heat energy is absorbed by particles to break the attractive forces so that they move freely and
far apart as gas particles. That‘s why the temperature remain constant
Freezing
Freezing is the change of liquid to solid by cooling down of liquid.
Freezing point is the temperature at which liquid freezes.
25
A
B C
26
2.6. Diffusion
Diffusion is the spreading and mixing of particles in gases and liquids.
27
2.6.3. Factors Affecting Rate of Diffusion
2.6.3.1. 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
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.
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3. Chemical and physical changes
A chemical change is a process that occurs and a new substance is formed e.g.
rusting, explosion of hydrogen in air,
burning of magnesium to ash,
Burning of a paper to ash, etc…
Physical change is a change in which no new substance is formed e.g. freezing (water to ice)
Evaporation of water (water to steam)
Magnetization of iron
Sublimation of solid iodine
Melting of ice
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4. Solutions, crystals, compounds and mixtures
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
A solute is a dissolved substance e.g. sugar, salt.
4.3. Solvent
A solvent: is a substance that dissolves a solute e.g. water, ethanol, petrol.
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.
30
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
A crystal is a solid that has solidified into regular fixed shape.
Crystals have regular shapes, flat surfaces and sharp edges
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.
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When a hydrated substance is heated, it loses its water of crystallization
CuSO4. 5H2O (s) CuSO4(s) + 5H2O (g)
Blue solid white solid
5.1. Compound
A compound is a substance which consists of two or more elements chemically combined
together
Examples of compounds include:
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
32
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
The word SPEC is used to recall these differences
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.
33
5.5.2. Crystallisation & Evaporation to Dryness
Crystallisation – separation of dissolved solid from a solution as well-formed crystals
Evaporation to Dryness – separation of dissolved solid from a solution as crystals of salt by
evaporating all the liquid off.
Why crystallisation occurs?
- Solubility of most solutes decrease as temperature decrease, when solution cools, solution can‘t
hold more solute (saturated) so the extra solute separates as pure crystals.
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.
This is called evapouration to dryness.
34
Using water bath carry out evapouration to dryness
To avoid spitting off (jumping off) of salt crystals formed, the evapourating basin is placed on a
water bath or sand bath. The steam produced from the water bath will heat up the salt solution
until crystals are formed.
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
This method can also be used to separate chalk from water
35
Diagram showing decanting
36
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
Distillation helps in separating of substances
(miscible liquid) and also in purification of liquids
37
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.
38
Crude oil is separated into its constituents such as petrol, diesel and kerosene
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
39
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.
Other substances which sublime other than iodine include:
Ammonium chloride
Anhydrous aluminum chloride
Iron (III) chloride
Benzoic acid
40
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.
41
5.5.9. Chromatography
- To separate liquid components in a mixture
This is a process of separating different coloured substances using a
porous paper. The coloured substances are moved over the paper at
different rates by a moving solvent.
The components of ink are separated by chromatography
Chromatography shows that ink consists of many coloured
substances.
This can be done as follows
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.
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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)
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6.5. Chemical symbols
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.
Common elements and their symbols include
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
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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 Argentums Ag
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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
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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.
47
7.3. Structure of an atom
The central part of an atom is called the nucleus. Protons and neutrons are located in the nucleus.
These make up nucleon number.
48
7.5. Atomic mass
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.
Atomic mass = protons + neutrons
Consider the following element
―Z‖ is the atomic mass and ―A‖ is the atomic number
The number of protons =number of electrons = b = atomic number
The number of neutrons = atomic mass – atomic number i.e. Z-A
Example
1. An atom of an element is represented by the symbol . State
a. Its atomic number = 8
b. The mass number = 16
c. The number of neutrons = 8 (i.e. 16 -8 =8)
2. Work out the following, . State
a. Its atomic number = 11
b. The mass number = 23
c. The number of neutrons = 12 (i.e. 23 - 11 = 12)
Examples of electronic structures of some elements:
Oxygen
Calcium
Argon
Chlorine
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Examples
Draw the structure of the following atoms and state their electronic configuration.
1. Oxygen (O): atomic number = 8
2:6 (Oxygen)
2. Carbon (C): atomic number = 6
2:4 (Carbon)
3. Chlorine (Cl): atomic number =17
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2:8:7 (Chlorine)
4. Sulphur (S): atomic number = 16
2:8:6 (Sulphur)
5.
a. Draw the electronic structure of element X represented by
b. State the number of neutrons and protons found inside the nucleus of X
6. Fill in the following table
Atom Atomic No. No. of electrons Electronic
configuration
Hydrogen (H) 1
Helium (He) 2
Lithium (Li) 3
Boron (B) 4
Nitrogen (N) 7
Neon (Ne) 10
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Phosphorous (P) 15
Chlorine (Cl) 17
Potassium (K) 19
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ii. W2-
iii. Y
iv. Y3+
c. State the two atoms that are isotopes of the same element
d. Write the formula of the compound formed between X and Y
7.8. Isotopes
These are atoms of the same element with
Same atomic number
Different atomic masses due to difference in number of neutrons
Isotopes of an element have the same chemical properties.
E.g. Hydrogen has three isotopes
Isotopes Hydrogen Deuterium Tritium
Protons 1 1 1
Electrons 1 1 1
Neutrons 0 1 2
Mass No. 1 2 3
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Calculating relative atomic masses
Examples
1. The relative abundances of and are 75% and 25% respectively by weight.
Calculate the relative atomic masses of chlorine
Relative atomic mass of chlorine = Mass due to + Mass due to
= +
=35.5
Exercise
1. Carbon has two main isotopes and . has relative abundance of 1.11% and
has relative abundance 98.89%. Calculate the relative atomic mass of carbon.
2. Given the sodium atom,
a. Give the examples of isotopes of sodium
b. Write down the electronic configuration of sodium
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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 V-VIII are non-metals
Elements in group IV are metalloids except Carbon which is a non-metal
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
For elements in groups V-VII the valency is given by 8-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.
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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).
They are excellent conductors of both heat and electricity.
The metals can be broken down into several groups.
Transition metals (also called the transition elements) are known for their ability to refract light
as a result of their unpaired electrons.
They also have several possible oxidation states.
Ionic solutions of these metals are usually colored, so these metals are often used in pigments.
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.
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Specific families
Alkali metals (1A) Group one—
They have one electron each in their outermost shell.
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 have a valency of one.
They are strong reducing agents.
They form compounds by either metallic or ionic bonding.
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 atomic number 3 and electronic configuration 2:1
Has valency 1
Silvery white in colour
Extracted by electrolysis of lithium chloride
Its compounds are mainly deliquescent
Not very typical of group I elements, but resembles them
Harder than K and Na
Reactions
i. With water
Lithium reacts slowly with cold water to form its hydroxide and liberates hydrogen
Lithium + Water Lithium hydroxide + Hydrogen
2Li (s) + 2H2O(l) 2LiOH (aq) + H2 (aq)
ii. With air
Gives a red flame colour when burnt in air i.e. burns with a red flame
4Li (s) + O 2 (g) 2Li2O(s)
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iii. With chlorine
58
Alkaline earth metals (Group 2A elements)
These are also reactive metals, but they don‘t 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
Silvery white metal
Electronic configuration 2:8:2
Extracted electronically
Reactions
i. With air
Dry air does not attack magnesium
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
2Mg(s) + O2 (g) → 2MgO (s) (Magnesium oxide)
3Mg(s) + N2 (g) → Mg3 N2 (s) (Magnesium Nitride)
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When heated, magnesium reacts with halogens to produce salts called halides
Mg(s) + Cl2 (g) → MgCl2 (s)
b. Calcium (Ca)
Widespread in the earth‘s crust as CaCO3, CaSO4, CaF2, Ca3(PO4)2 etc…
Extraction
Mainly by electrolysis of its fused calcium chloride
Properties of calcium
Silvery white in colour
Soft enough to cut with a knife
Reactions
i. With air
On exposure to air at room temperature, calcium tarnishes in color as it froms an oxide
Calcium + oxygen → calcium oxide
2Ca(s) + O2 (g) → 2CaO (s) (Calcium oxide)
Calcium oxide later combines with water to form calcium hydroxide
Calcium oxide + water → calcium hydroxide
CaO (s) + H2O (aq) → Ca(OH)2 (aq)
Calcium hydroxide formed combines with carbon dioxide to form calcium carbonate
Ca(OH)2 (aq) + CO2 (g) → CaCO3 (s) + H2O (l)
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)
ii. With water
Calcium reacts with water to form calcium hydroxide liberating hydrogen
Ca (s) + 2H2O (aq) → Ca(OH)2 (aq) + H2 (g)
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.
60
They include fluorine (F2), chlorine (Cl2), bromine (Br2) and iodine (I2)
Have seven electrons in their outermost shell
Properties
Chlorine
Greenish yellow gas
Poisonous gas
Denser than air
Bleaches damp litmus paper
Atomic number 17
Electronic configuration 2:8:7
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)
Cl2 (g) + H2O(l) → HCl (aq) + HOCl (aq)
ii. With dilute sodium hydroxide
Halogens generally react with NaOH to produce a pale yellow solution of sodium chloride or
sodium bromide
NaOH(aq) + Cl2(g) → NaCl(aq) + H2O(l)
iii. Displacement reactions
Written in order Cl, Br, I each halogen displaces those on its right from their solutions of simple
salts i.e. chlorine displaces bromine from its own solution and bromine forms a brown solution
2NaBr (aq) + Cl2(g) → 2NaCl (aq) + Br2(aq)
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.
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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
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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
They are prominently two types of bonds
The electrovalent bond and covalent bonds
Elements are made of atoms
Atom is smallest unit of an element, having properties of that element.
Molecule is group of two or more atoms chemically joined together, e.g. chlorine molecule has 2
chlorine atoms
Chemical formula shows the number and kinds of atoms in a molecule, e.g. chlorine molecule
has formula Cl2, where Cl is chlorine symbol and the subscript number (2) shows that there are 2
atoms in a chlorine gas molecule.
Compounds
Compound is substance containing two or more elements chemically combined together e.g.
Magnesium is an element; oxygen is an element – they can only be burnt to form magnesium
oxide compound.
Composition of compounds
Ions or molecules make up compounds
Ions are atoms having electrical charge
E.g. NaCl made up of 2 ions; positively charged Na, negatively charged Cl.
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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.
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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 don‘t 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.
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In covalent bond, TRY TO SUBTITUTE THE SHORT OF ELECTRONS OF TWO/MORE
ATOMS BETWEEN EACH OTHER TO FORM THE 2 OR 8 VALENCE ELECTRONS.
THE SHARED ELECTRONS APPEAR IN PAIRS!
E.g. H2 molecule
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
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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.
Three pairs of shared electrons between 2 atoms forms TRIPLE BOND, X ≡ Y.
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8.5. Dative bonding
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.
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1. Metals can be bent (ductile) and can be stretched (malleable) because the layers of atoms in
metals slide over each other when force is applied but will not break due to attractive force
between electrons and metal ions.
2. Metals conduct electricity as it has free electrons which carry current.
3. Metals conduct heat as it has free electrons which gains energy when heated and moves faster
to collide with metal atoms, releasing heat in collisions.
4. Metals have high melting and boiling points because the bond between metals is very strong.
Hence very high heat energy needed to break the bonds.
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9. Valency
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
70
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
Radical Symbol Valency
Hydroxide OH 1
Bromide Br 1
Nitrate NO3 1
Chlorine Cl 1
Hydrogen carbonate HCO3 1
Hydrogen sulphate HSO4 1
Oxide O 2
Carbonate CO3 2
Sulphate SO4 2
Sulphite SO3 2
Phosphate PO4 3
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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.
Sodium Chloride Calcium Sulphate Aluminium Oxide
Na Cl Ca SO4 Al O
Write the valences on top right side of the symbols of elements or radicals
Na1 Cl1 Ca2 SO42 Al3 O2
Rewrite the symbols again reversing the valences from top right side to the bottom right
side of symbols and radicals
Na1 Cl1 Ca2 SO42 Al3 O2
Giving
Na1 Cl1 Ca2 (SO4) 2 Al2 O3
Na1Cl1 Ca2 (SO4)2 Al2O3
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
Na1 Cl1 Ca2/2 (SO4) 2/2 Al2 O3
Formula: NaCl CaSO4 Al2O3
Examples
i. Sodium hydroxide
Sodium Hydroxide
72
Symbols Na OH
Valencies on top Na1 OH1
Reversing valencies Na1 OH1
Neglecting 1 in final formula: NaOH
73
CaSO4 = one mole of calcium sulphate contains
1 atom of calcium
1 atom of sulphur
4 atoms of oxygen
(NH4)2SO4 = one mole of ammonium sulphate contains
2 atom of nitrogen
8 atoms of hydrogen
I atom of sulphur
4 atoms of oxygen
5Al2O3 = 5 moles of aluminium oxide contains
10 atoms of aluminium
15 atoms of oxygen
Exercise
Write the chemical formulae of the following compounds and calculate the number of atoms of
each element
Chemical name Common name Formula
Calcium hydroxide (solid) Slaked lime Ca(OH)2
Calcium oxide Lime, quick lime CaO
74
Calcium hydroxide solution Lime water Ca(OH)2
Potassium hydroxide Caustic potash KOH
Sodium hydrogen carbonate Baking soda NaHCO3
Calcium carbonate Chalk, limestone or CaCO3
marble
Sodium hydroxide Caustic soda NaOH
Iron (III) oxide (hydrated) Iron rust Fe2O3.xH2O
Potassium nitrate Salt patre KNO3
Sodium carbonate hydrated Washing soda Na2CO3.10H2O
75
11. Chemical equations
Writing chemical equations
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
H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + H2O(l)
ii. Hydrogen reacts with oxygen to form water
H2(g) + O2(aq) 2H2O(l)
NB: The following symbols represent
76
Count the number of atoms of each element on the right hand side and on the left hand
side to see if they balance
H2 (g) + O2 (g) → H2O (l)
LHS RHS
H=2 H=2
O=1 O=1
If they don‘t balance, look for a number which can be multiplied in the formula above to
make all elements balance on both sides of the equation.
2H2 (g) + O2 (g) → 2H2O (l)
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
77
Na=1 Na=2
O=2 O=1
Not balanced
4Na (s) + O2 (g) → 2Na2O (s)
LHS RHS
Na=4 Na=4
O=2 O=1
Balanced
Or 2Na (g) +½ O2 (g) → Na2O (l) multiplying by two through out to remove
the fraction
4Na (s) + O2 (g) → 2Na2O (s)
ii. Fe (s) + O2 (g) → Fe2O3 (s)
LHS RHS
Fe=1 Fe=2
O=2 O=3
Not balanced
iii. 4Fe (s) + 3O2 (g) → 2Fe2O3 (s)
LHS LHS
Fe=4 Fe=4
O=6 O=6
Balanced
Or 2Fe (s) + O2 (g) → Fe2O3 (s)) multiplying by two through out to remove
the fraction
4Fe (s) + 3O2 (g) → 2Fe2O3 (s)
iv. Zn (s) + HCl (aq) → ZnCl2 (aq) + H2 (g)
LHS RHS
78
Zn =1 Zn =1
H=1 H=2
Cl =1 Cl =2
Not balanced
Zn (s) + 2HCl (aq) → ZnCl2 (aq) + H2 (g)
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:
i. KClO3 (s) → KCl (s) + O2(g)
2KClO3 (s) → 2KCl (s) + 3O2(g)
ii. Mg (s) + HCl(aq)→ MgCl2(aq) + H2(g)
Mg (s) + 2HCl(aq)→ MgCl2(aq) + H2(g)
iii. CaCO3 (s) → CaO (s) + CO2(g)
CaCO3 (s) → CaO (s) + CO2(g) Directly balanced
iv. NaOH (aq) + HCl(aq) →NaCl (s) + H2O(l)
NaOH (aq) + HCl(aq) →NaCl (s) + H2O(l) Directly balanced
Exercise
79
c) 4Na (s) + O2 (g) → 2Na2O (s)
d) NaNO3 (s) → NaNO2 (s) + O2 (g)
e) Pb(NO3)2(s) → PbO (s) + NO2 (g) +O2 (g)
80
12. Types of chemical reactions
There are four types of chemical reactions
81
12.4. Double replacement
When an element in a compound replaces another element in another compound, this type of
reaction is called double displacement.
It has a general formula of AB + CD → AD + BC
Examples
ZnSO4(s) + Na2CO3 (s) → ZnCO3(s) + Na2SO4 (s)
2KOH (aq) + H2SO4 (aq) → K2SO4 (aq) + 2H2O (l)
NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l)
NH4OH (aq) + HCl (aq) → NH4Cl (aq) + H2O (l)
82
13. The atmosphere and combustion
The composition of the atmosphere
13.1.1. Oxygen
It occupies 21% of the atmospheric air
It is added to the atmosphere during the process of
photosynthesis
thermal decomposition of nitrates
it is removed from the atmosphere during respiration, rusting and combustion
13.1.3. Nitrogen
This gas occupies the largest volume of 78%
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.
83
13.1.4. Carbon dioxide
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.
CO2(g) + H2O (l) → C6 H12O6(aq) + O2 (g)
84
13.2.2. Deliquescent substances
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
85
13.4. Burning substances in air
The part of air used for burning is oxygen
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
2Na(s) + O2 (g) → Na2O2 (s) (sodium peroxide)
4Na(s) + O2 (g) → 2Na2O (s) (sodium oxide)
When sodium oxide is dissolved in water, sodium hydroxide is formed
Na2O (s) + H2O (l) → 2NaOH (aq) (sodium hydroxide)
When sodium peroxide is dissolved in water effervescence/bubbles of a colourless gas (oxygen)
are observed and a colourless solution (sodium hydroxide solution) are observed
2Na2O2(s) + 2H2O (l) → 4NaOH (aq) + O2 (g)
Magnesium
This burns with a brilliant flame forming a white smoke and ash i.e. an oxide and a nitride
2Mg(s) + O2 (g) → 2MgO (s) (Magnesium oxide)
3Mg(s) + N2 (g) → Mg3 N2 (s) (Magnesium Nitride)
Aluminium
This burns in air when heated strongly and becomes very hot
It forms aluminium oxide and a little nitride
4Al(s) +3 O2 (g) → 2 Al2 O3 (s) (Aluminium oxide)
2Al(s) + N2 (g) → 2AlN (s) (Aluminium nitride)
Zinc
Zinc burns in air with a green flame to form zinc oxide. It does not form a nitride.
2Zn(s) + O2 (g) → 2ZnO (s) (Zinc oxide)
Lead
This melts on heating to shiny beads and then forms lead oxide which is brown when hot and
turns yellow on cooling
2Pb(s) + O2 (g) → 2PbO (s) (Lead II oxide)
Brown when hot
Yellow on cooling
86
Copper
This burns with a green flame to form a black oxide
2Cu(s) + O2 (g) → 2CuO (s) (copper II oxide)
Non metals
Carbon
This burns in a plentiful supply of air to form Carbon dioxide
C(s) + O2 (g) → CO2 (g)
If oxygen is in short supply, carbon burns with little oxygen to form carbon monoxide
2C(s) + O2 (g) → 2CO (g)
Phosphorous
This burns with a bright yellow flame and produces dense white fumes of phosphorous pentoxide
P4 (s) + 5O2 (g) → P2 O5 (s)
Or 4P (s) + 5O2 (g) → P2 O5 (s)
Sulphur
This burns in air with a bright blue flame forming cloudy fumes of sulphur dioxide
S(s) + O2 (g) → SO2 (g)
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
87
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.
= X100
From the experiment, if the value of X is 10 and that of Y is 7.9, the percentage of oxygen used
up
= X100
= X100
=21%
Therefore the volume of oxygen used in rusting is 21% which is the composition of oxygen in
the atmosphere.
1.8. Burning candle
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
88
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.
Ca(OH)2 (aq) + CO2 (g) → CaCO3 (s) + H2O(l)
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)
4Fe(s) +3O2 (g) → 2Fe2O3 (s)
Fe2O3 (s) + xH2O (l) →Fe2O3.xH2O (s)
Conditions necessary for rusting
There are two conditions necessary for rusting to take place. These are:
Oxygen
Water
For rusting to take place both of these must be present.
An experiment to show that iron rusts when left exposed to damp air
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.
89
An iron nail rusts in presence of water and air/damp air.
Clean iron nails or iron fillings are placed at the bottom of a test tube. Cotton wool is pushed half
way the tube and anhydrous calcium chloride placed on top of the cotton wool, to absorb all the
moisture from air, leaving it dry.
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.
90
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
1. Rusting took place
2. Rusting did not take place
3. Rusting did not take place
91
4. Rusting did not take place
Explanation
A- For A all conditions for rusting are present
B- For B oxygen gas is absent
C- For C both water and oxygen gas are absent
D- For D water is is absent
Therefore rusting takes place in presence of water and oxygen gas.
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
92
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)
93
14. Methods of gas collection
The method of gas collection depends on the properties of the gas.
It depends on whether the gas is;
Soluble or insoluble in water
Lighter or denser than air
There are three methods of gas collection:
a. Over water method/downward displacement of water
The method is used to collect gases which are insoluble or slightly soluble in water e.g. carbon
monoxide , oxygen, nitrogen and hydrogen.
b. Downward delivery or upward displacement of air:
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
94
d. Gas syringe
95
15. Oxygen
Laboratory preparation of oxygen
Potassium chlorate is heated in presence of manganese (IV) oxide and oxygen gas is produced. It
is collected in the gas jar over water.
2KClO3 (s) → 2KCl (s) + 3O2 (g)
Manganese (IV) oxide acts as a catalyst.
A catalyst: is a substance which alters the rate of a chemical reaction and remains chemically
unchanged at the end of the reaction.
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
96
2H2O2 (l) → 2H2O (l) + O2 (g)
The gas is then collected over water.
No heat is applied.
Manganese (IV) oxide acts as a catalyst.
By fractional distillation of liquefied air since nitrogen has a lower boiling point of 90K remains
as a liquid
Properties of oxygen
97
Has a boiling point of -183oC
It relight a glowing wooden splint
Chemical properties
- 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.
Na2O (s) + H2O (l) → 2NaOH (aq) (Sodium hydroxide)
K2O (s) + H2O (l) → 2KOH (aq) (Potassium hydroxide)
98
b. Acidic oxides/Acid anhydrides
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…
CO2 (g) + H2O (l) → H2CO3 (aq)
SO2 (g) + H2O (l) → H2SO3 (aq)
SO3 (g) + H2O (l) → H2SO4 (aq)
2NO2 (g) + H2O (l) → 2HNO3 (aq)
Note:
Acidic oxides are also called acid anhydrides.
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…
99
16. Oxidation and reduction
Oxidation
Reduction
Oxidizing agents
100
Reducing agents
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.
101
17. Water and hydrogen
Water
102
17.2. Properties of pure water
It is a colourless liquid
It has a flat taste
It is neutral to litmus
It is freezes at 0oC and boils at 100oC and pressure of 1 atmosphere
It has a maximum density of 1g/cm3 at 4oC
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
This reacts vigorously with cold water to produce a colourless solution of sodium hydroxide and
a colourless gas of hydrogen
2Na(s) +2 H2O(l) → 2NaOH(aq) +H2(g)
During the reaction of sodium with water, a hissing sound is produced. However, sodium is less
reactive than potassium
Calcium
This reacts quietly with cold water producing calcium hydroxide and hydrogen
104
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) +2H2O(g) → MgO (aq) +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 (aq) +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)
105
17.6. Uses of water
Acts as a solvent
Raw material in the manufacture of ammonia and ethanol
Used in the production of steam to drive turbines
For heating and cooling in industries
106
2) Reaction of Metals with Dilute Hydrochloric Acid
Potassium, sodium, calcium, magnesium, zinc and iron reacts with dilute hydrochloric acid to
form:
M(s) + 2HCl(aq) → MCl2(aq) + H2(g)
Metal + Acid → Metal Chloride + Hydrogen
Lead reacts with warm hydrochloric acid slowly
Copper and gold have no reaction with dilute hydrochloric acid
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.
107
E.g. Copper (II) oxide reacts with carbon and can be reduced by bunsen burner flame
temperature
CuO(s) + C(s) → Cu(s) + CO2(g)
For iron (III) oxide to be reduced, it needs very high temperature.
108
18. Hydrogen
Laboratory preparation of dry hydrogen gas
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
109
NB: Dilute nitric acid is not used in preparation of hydrogen because it is an oxidizing agent and
would therefore oxidize hydrogen produced to water.
Metals like sodium, potassium and calcium react with water to produce hydrogen.
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
110
It is neutral to litmus paper i.e. has no effect on litmus paper
It is the lightest gas (much less dense than air) and diffuses fast
It is not very soluble in water
It is the first element in the periodic table
Has an atomic number of one
Exists as a diatomic element
Chemical 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(l)
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)
Fe3O2 (s) + 2H2(g) → 3Fe(s) +2 H2O(l)
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)
111
Hydrogen also reacts with alkenes to give alkanes (hydrogenation)
CH2=CH2(g) + H2(g) → CH3CH3(g)
It combines with highly electropositive metals to form hydrides
2Na(s) + H2=(g) → 2NaH (s)
NB: Hydrogen chloride (HCl), Hydrogen sulphide (H2S ), ammonia (NH3) and water (H2O ) are
called hydrides.
A hydride is a compound of an element with hydrogen only.
Uses of hydrogen
112
19. Acids, bases and salts
19.1. Acids
An acid is a substance which when dissolved in water produces hydrogen ions as the only
positively charged ions
113
19.1.4. Some Acids with Their Basicity
114
Weak acids react slowly with metals than strong acids – hydrogen gas bubbles are produced
slowly.
19.1.6. Comparing Strong and Weak Acids with Concentrated and Dilute Acids
Concentration Strength
Is the amount of solute (acids or alkalis) dissolved in 1 dm3 of a Is how much ions can be
solution disassociated into from
acid or alkali
It can be diluted by adding more water to solution or concentrated by The strength cannot be
adding more solute to solution changed
Comparing 10 mol/dm3 and 0.1 mol/dm3 of hydrochloric acids and 10 mol/dm3 and 0.1 mol/dm3
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
115
NaHCO3(s or aq) + HCl(aq) → NaCl(aq) + CO2(g) + H2O (l)
5) - Acids react with metal oxides and hydroxides
Metal oxides & hydroxides react slowly with warm dilute acid to form salt and water
Cu(OH)2(s) + H2SO4(aq) → CuSO4(aq) + 2H2O (l)
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 don‘t react with metals and don‘t 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
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- However when dissolved in organic solutions, they don‘t show acidic properties
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)
All alkalis produce hydroxide ions (OH-) when dissolved in water. Hydroxide ions give the
properties of alkalis. They don‘t behave as acids in absence of water.
Alkalis are therefore substances that produce hydroxide ions, OH-(aq), in water.
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Energy is given out during the reaction. Common neutralisation reactions include:
NaOH(aq) + HCl(aq)→NaCl(aq) + H2O(l)
2NaOH(aq) + H2SO4 (aq)→Na2SO4 (aq) +2H2O(l)
KOH(aq) + HCl(aq)→KCl(aq) + H2O(l)
Na2CO3 (aq) + H2SO4 (aq)→ Na2SO4 (aq) +H2O(l) + CO2(g)
5) Alkalis react with ammonium compounds
They react with heated solid ammonium compounds to produce ammonia gas
(NH4)2SO4(s) + Ca(OH)2(aq) → CaSO4(aq) + 2NH3(g) + 2H2O(l)
6) Alkalis react with solutions of metal ions
Barium sulphate, BaSO4(aq), contains Ba2+(aq) ions
Ca(OH)2(aq) + BaSO4(aq) → Ba(OH)2(s) + CaSO4(aq)
The solid formed is a precipitate – the reaction is called a precipitation reaction
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Weak base - Base that partially ionises in water. The remaining molecules remain unchanged as
a base. Their reactions are reversible. E.g. NH3
NH3(g) + H2O(l) ⇌ NH4+(aq) + OH-(aq)
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.
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
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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 it‘s rinsed or acid conditioner is used to neutalise excess alkali
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19.6. Ionic Equations
Ionic equation is equation involving ions in aqueous solution, showing formation and changes
of ions during the reaction
Rule to make ionic equations:
- Only formulae of ions that change is included; ions don‘t change = omitted
- Only aqueous solutions are written as ions; liquids, solids and gases written in full
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FeO(s) + 2H+(aq) + SO42-(aq) → Fe2+(aq) + SO42-(aq) + H2O(g)
Note: FeO is written in full as it‘s solid (although it‘s an ionic compound)
Since SO42-(aq) ions don‘t change, omit SO42- ions, leaving:
FeO(s) + 2H+(aq) → Fe2+(aq) + H2O(g)
E.g. Reaction between calcium carbonate and hydrochloric acid
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)
Its ionic equation is:
CaCO3(s) + 2H+(aq) + 2Cl-(aq) → Ca2+(aq) + 2Cl-(aq) + CO2(g) + H2O(l)
Since 2 Cl-(aq) ions don‘t change, omit Cl- ions, leaving:
CaCO3(s) + 2H+(aq) → Ca2+(aq) + CO2(g) + H2O(l)
Reactions Producing Precipitate
E.g. Reaction between calcium hydroxide and barium sulphate
Ca(OH)2(aq) + BaSO4(aq) → Ba(OH)2(s) + CaSO4(aq)
Its ionic equation is written as:
Ca2+(aq) + 2OH-(aq) + Ba2+(aq) + SO42-(aq) → Ba(OH)2(s) + Ca2+(aq) + SO42-(aq)
Since Ca2+(aq) and SO42-(aq) ions don‘t change, omit them, leaving:
Ba2+(aq) + 2OH-(aq) → Ba(OH)2(s)
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 don‘t 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
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H2SO4 (aq) + 2NaOH(aq) → Na2SO4 (aq) + H2O(l)
Ionic equation is:
H+(aq) + OH-(aq)→ H2O(l)
Plants don‘t 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
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Examples of salts and their corresponding salts
Acid Salt Name of the salt
HCl NaCl Sodium chloride
HCl KCl Potassium chloride
HCl NH4Cl Ammonium chloride
HNO3 AgNO3 Silver nitrate
H2SO4 CuSO4 Copper sulphate
H2SO4 MgSO4 Magnesium sulphate
H2SO4 Na2SO4 Sodium sulphate
H2SO4 NaHSO4 Sodium hydrogen sulphate
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…
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E.g. chlorides of aluminum and iron can be prepared from their elements directly
Al(s) + 3Cl2(g) → 2AlCl3(s)
Fe(s) + 3Cl2(g) → 2FeCl3(s)
NB: The gas is passed over the heated metal in each case
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
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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.
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Pb2+(aq)+ 2NO3-(aq) + 2H+(aq) + 2Cl-(aq)→ 2NO3-(aq) + 2H+(aq) + PbCl2(s)
Pb2+(aq) + 2Cl-(aq)→ PbCl2(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.
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Filter to remove excess zinc and other insoluble impurities. The filtrate is a colourless solution
containing zinc sulphate
Allow the filtrate to crystallize by heating it so that water can freely evaporate
Colourless needle like crystals will be obtained after evaporation
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
ZnSO4(aq) +7 H2O(g) → ZnSO4.7 H2O(s)
NB: Other salts such as magnesium sulphate, copper (II) sulphate and iron (II) sulphate can be
prepared in the same way. But with copper (II) sulphate, concentrated sulphuric acid is used
instead of dilute sulphuric acid with copper metal.
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NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(g)
The process can also be used to prepare salts of potassium and ammonium
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Zinc is added to dilute sulphuric acid until in excess to ensure no more sulphuric acid is present.
Then the mixture is filtered off to separate zinc from zinc sulphate. The filtrate which contains
zinc sulphate is then placed in evaporating dish to evaporate most of water then it‘s cooled after
zinc sulphate crystals are formed. The crystals are then filtered and squeezed between filter
papers to dry.
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K2CO3(s) + H2SO4 (aq) → K2SO4 (aq) + CO2 (g) + H2O (l)
The same process is used as reaction of acid with metal, just that carbon dioxide is produced.
Carbon dioxide can be tested by bubbling it into limewater which will turn limewater colourless
to milky.
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
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Potassium Decompose to pale yellow metal nitrites and oxygen
Sodium
2MNO3(s) → 2MNO2 (s) + O2 (g)
Calcium
Magnesium
Aluminum 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.
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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)
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Others are sparingly soluble in water at room temperature e.g. lead chloride, silver chloride,
barium sulphate, lead sulphate, calcium carbonate etc…
The solubility of these sparingly soluble salts increases as temperature increases.
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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
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Calculate the mass of potassium nitrate dissolved in 100g of water at 50oC, using the results
obtained. This will be the solubility of potassium nitrate at 50oC
Sample results
Mass of dish =14.32g
Mass of dish + potassium nitrate solution =35.70g
Mass of dish + potassium nitrate solid =18.60g
Temperature of the saturated solution = 50oC
Therefore
Mass of water =35.70 – 18.60 =17.10g
17.10g of water dissolved 4.28g of potassium nitrate =18.60-14.32 = 4.28g
17.10g of water dissolved 4.28g of potassium nitrate
18.60g of water will dissolve =
=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)
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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.
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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.
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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.
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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.
141
C(s) + O2(g)→ CO2(g)
Exhaust fumes from cars contain carbon monoxide because of incomplete combustion of petrol
or diesel.
142
Some powdered wood charcoal is placed in a deflagrating spoon and heated over a Bunsen
burner flame until it glows red-hot. The spoon is immediately transferred into a jar of oxygen
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.
143
2ZnO(s) + C(s) → 2Zn(s) + CO2 (g)
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:
144
Dilute hydrochloric acid from a tap funnel is added to calcium carbonate in a flat-bottomed flask
Effervescence occurs and a colourless gas, which is carbon dioxide, is formed according to the
equation.
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.
145
Lead (II) carbonate is also not used because when it reacts with dilute hydrochloric acid or
sulphuric acid, the reaction soon slows down and then stops. This is due to the formation of lead
(II) chloride or lead (II) sulphate, both of which are insoluble salts.
The insoluble salt coats the carbonate preventing it from reacting with the acid.
Kipp’s apparatus
A continuous supply of carbon dioxide can be obtained from a Kipp‘s 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.
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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
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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.
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(c) Reaction with alkalis
Sodium hydroxide solution readily absorbs carbon dioxide to produce sodium carbonate.
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.
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Uses of carbon dioxide
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
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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.
Apparatus
151
Preparation of carbon monoxide from sodium methanoate (HCOONa)
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.
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2KOH(aq) + CO2(g) → K2CO3(aq) + H2O(l)
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.
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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.
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6. It is a poisonous gas because it forms a fairly stable compound with haemoglobin which
reduces the oxygen-carrying capacity of blood.
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Does not burn Burns with a blue flame
No action on oxides Reduces metallic oxides
Reacts with heated charcoal to form carbon No action on charcoal
monoxide
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.
156
157
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.
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Brown residue when hot
Yellow residue on cooling
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.
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Carbonates of lithium, potassium, sodium and ammonium are soluble in water. The other
carbonates are insoluble in water.
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 (NaHCO3) 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
160
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.
161
(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.
162
It can also be obtained by strongly heating sea shells. Sea shells contain calcium carbonate which
decomposes into the oxide and carbon dioxide.
163
20.4. Sodium carbonate (soda ash)
Laboratory preparation of sodium carbonate
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.
164
Soda ash is obtained at Lake Magadi in Kenya by the solvary process.
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.
165
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
166
Such an action, that is, the giving up of water of crystallization to the atmosphere is termed as
efflorescence.
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
167
C6H12O6 (l) + 6O2 (g) → 6CO2 (g) + 6H2O (l)
3) Thermal decomposition of calcium carbonates: Carbon dioxide passes into air when
limestone or chalk is heated.
CaCO3 (s) → CaO (g) + CO2 (l)
4) Fermentation: in the manufacture of ethanol during fermentation carbon dioxide is
produced as a bi-product.
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
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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.
169
(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?
170
(ii) Write equation for the reaction that would occur.
Hardness of water is due to presence of calcium ions (Ca2+) or magnesium ions (Mg2+) present in
water
1) Boiling
171
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
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2) Allows for use of cheap lead pipes for water supply. Does not cause lead poisoning
3) Necessary for formation of animal shells. Some aquatic animals need calcium to form their
shells e.g. water snails.
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 can‘t pass
through them easily.
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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
174
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.
175
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
176
Y Doesn‘t conduct Doesn‘t conduct Not soluble
(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
177
The positively charged ions called cations move to the negatively charged electrode called
cathode where they gain electrons and become atoms which are then said to be discharged i.e.
Mn+(aq or l) + ne→ M (s)
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+
178
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 SO42- 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
179
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
180
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
181
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
182
(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
183
Molten Q Conducts an electric current and a brown
substance is deposited at the cathode.
(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).
184
Ions present:
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)
185
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.
186
Chloride ions from copper (II) chloride and hydroxide ions from water move to the anode but
chloride 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
187
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
2-
solution SO4
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
188
(a) Write an equation for the reaction that took place at the
(i) Anode
(ii) Cathode
(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
189
22.1. Laws of electrolysis
The laws of electrolysis were stated by Faraday. According to his laws, the amount of substance
produced during electrolysis depends on the
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.
190
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.
191
2Cl-(aq) → Cl2(g) + 2e-
From the above examples, one mole of electrons is equivalent to 1 faraday. 1 faraday is
equivalent to 96500 coulombs.
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
192
Q = 0.45 x 1500 = 675 C
= 2 x 96500 = 193000 C.
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
193
So 2 x 96500 C produces 1 mol of magnesium.
2 x 96500
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.
194
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.)
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.
195
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
196
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.
197
(iv) Extraction of metals
Reactive metals such as aluminium and sodium are extracted by electrolysis of the fused
electrolytes
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.
Sodium amalgam
Sodium hydroxide is used in the manufacture of soap, rayon, paper and in the purification of
bauxite for aluminium extraction.
198
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.
199
(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.
(ii) Write an equation for the reaction that took place at each electrode.
200
(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.
201
7. Copper(II) sulphate solution was electrolysed using carbon electrodes.
(a) State what was observed at the
(i) Anode.
(ii) Cathode.
8. The figure below shows an arrangement of the apparatus used for the purification of
copper.
(a) Name the substance used as
(i) Anode.
(ii) Cathode.
(ii) Y.
202
9. An aqueous solution of potassium iodide was electrolysed in a U-tube using carbon
electrodes. Iodide was formed at the anode and it dissolved to form a yellowish brown
solution around the electrode. At the cathode, bubbles of a colourless gas were seen to
evolve. The solution near the cathode had a pH of about 11.
(a) Explain with an equation how the change from iodide ions to iodine took place at the
anode.
(b) What was the gas evolved at the cathode?
(c) A solution of potassium iodide in water is neutral (pH = 7). Explain why the pH
increased to 11 near the cathode during electrolysis.
203
23. Formulae, stoichiometry and the mole concept
atom.
Relative Atomic Mass - the average mass of one atom of the element (averaging isotopes) when
compared with mass of a carbon-12 atom.
Ar =
Note: The Relative Atomic Masses are already stated on the periodic table above each chemical
formula.
204
Relative Molecular Mass – the average mass of one molecule of substance (averaging isotopes)
when compared with mass of a carbon-12 atom.
In short: Mr = =
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
205
23.2. Percentage Composition
E.g.
% of Na = x 100 = 39.32%
=( x 100) %
=( x 100) %
=( x 100) %
= 70%
Mr (Fe3O4) = 3(56) + 4(16) = 232
= 72%
Fe3O4 has more iron composition than that of Fe2O3.
206
e.g. Determine the mass of iron in 200g of Fe2O3.
Mr(Fe2O3)= 2(56) + 3(16) = 160
=( x 200g
= 140g
=( x 12.5g
= 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
n=
n=
= 5 mol
207
e.g 2: Calculate the number of molecules in 0.25 mole of CO2. Hence, how many atoms are
present?
0.25mol =
n =
208
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
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
Number of moles
209
Divide by the smallest number
1 1
Mg1O1
Therefore, the empirical formula is MgO
Solution
Elements present S O
Percentage composition 40 60
Number of moles
1) 3
S1O3
Therefore, the empirical formula is SO3
210
Solution:
= 4.86g of carbon
= 1.03g of hydrogen
Elements present: C H O
12 1 16
211
Simplest ratio: 4 : 10 : 1
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
212
[(27 x 2) + (32 x 3) + (16 x 12)]n = 342
342n = 342
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)
213
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:
Example
What is the number of moles of 240cm3 of Cl2 at r.t.p.?
Solution
Since 24000cm3 of chlorine gas contain 1 mole
Then 1cm3 of chlorine gas contains moles
214
Molar Volume and Molar Mass
Gases have same volume but not necessarily same mass
Example: 1 mole of hydrogen gas (H2) has a mass of 2g,
1 mole of Carbon Dioxide gas (CO2) has a mass of 44g
Example
Find the volume of 7g of N2 at r.t.p.
Solution
Step 1: Find the number of moles from the mass of nitrogen
Molar mass of N2 =2 x 14 =28g
215
From above, we know that H2O is short 1 oxygen atom. Therefore we multiply the product by 2
first. Note: all atoms in molecules are automatically multiplied by 2.
O2 + H2 → 2H2O
O H H
O H H
H
H
O
O
Now we can cancel off oxygen atoms. However, hydrogen atoms on the reactant side are short of
2 atoms. Therefore, we multiply the hydrogen molecule by 2 so that the short is balanced. The
equation is fully balanced when we are able to cancel off all atoms of that element on both sides.
O2 + 2H2 → 2H2O
O H H
O H H
H H
H H
O
O
Then multiply the ratio by no. of moles of Y to find the reacting mole of X.
216
Number of moles of X = 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:
217
1. Ammonium chloride reacts with calcium hydroxide according to the equation
Ca(OH)2(s) + 2NH4Cl (aq) → CaCl2(s) + 2NH3(g) +2H2O (l)
If 14.8g of calcium hydroxide was reacted completely with ammonium chloride, what mass of
ammonia gas will be evolved? (H=1, N=14, O=16, Ca=40) (Ans = 6.8g of ammonia)
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.
218
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
219
23.11.2.3. Calculations involving energy changes
Carbon reacts with sulphur according to the equation
C(s) + 2S(s) → CS2(s) ∆H = +116kJmol-1
The amount of heat absorbed when 16g of sulphur reacts with excess carbon is (C=12, S=32)
Solution
2moles of sulphur absorb 116kJ
1g of sulphur absorb kJ
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)
220
Concentration (C) =
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.
221
1cm3 contains mol of sodium hydroxide
=0.9mol
From
Moles (n) =
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
1cm3 contains mol of Na2SO4.10H2O
222
500cm3 contains ( x 500) mol of Na2SO4.10H2O
=0.005mol
From
Moles (n) =
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
223
Detecting the End Point
End point is the point at which neutralisation of acid and alkali is complete
- Sharp indicators (phenolphthalein and methyl orange) are used to detect end points effectively
- Litmus and universal indicators are not used as the changes at the end point are not sharp
224
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
Step 3: Find the ratio of number of moles of H2SO4 to number of moles of NaOH
2NaOH(aq) : H2SO4(aq) =2:1
=0.0015moles
225
= 0.06M
Molarity of H2SO4 = 0.06M
Step 3: Find the ratio of number of moles of H2XO4 to number of moles of NaOH
2NaOH(aq) : H2XO4(aq) =2:1
=0.01275moles
226
Step 5: Find the concentration of H2XO4 in mol/dm3
= 0.51M
Molarity of H2XO4 = 0.51M
Example:
7.2g of an impure sample of hydrated sodium carbonate (Na2CO3.10H2O) was dissolved in 250
cm3 of solution. 20 cm3 of this solution was required to completely react with 25 cm3 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
227
1cm3 contains moles
=0.00125moles
= 0.0625M
Molarity of sodium hydroxide = 0.0625M
ii. Molarity
228
1 cm3 contains g of pure hydrated sodium carbonate
Percentage purity =
=62.07%
a.
b.
i. Moles of the acid
1000 cm3 of acid contain 0.1moles
229
1cm3 contains moles
containing 6.4g of the oxalic acid (COOH)2. xH2O solution. 2moles of sodium hydroxide react with 1
mole of oxalic acid(C=12, O=16, H=1)
Calculate:-
a. The moles of sodium hydroxide that reacted
b. The molecular mass of (COOH)2.xH2O
c. The value of X (number) of moles of water of crystallisation in (COOH)2.xH2O
Solution
a. Moles of sodium hydroxide
1000 cm3 contain 0.1moles of sodium hydroxide
230
20 cm3 contain moles of sodium hydroxide
= 128
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
X =
X = 2.1
Number of moles of water of crystallisation =2. Hence (COOH)2. 2H2O
Exercise
231
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)
232
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
=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)
233
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 Avogadro‘s 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.
234
2Volumes 1Volume 2Volumes
10cm3 5cm3 10cm3
So the volume of excess oxygen =20-5 = 15cm3
Volumes of carbon dioxide =10cm3
Total volume = product + excess oxygen
Total volume = (15 x 10) = 25cm3
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 α T and = a constant
235
From Charles‘ law =
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)
236
24. Qualitative analysis
Blue
Cu 2+ Colour of solid
white
White
green yellow
Cu 2+ Blue
Fe 2+
Cu 2+ 2+
Fe 3+
Zn
green Pb2+
Cu 2+ Blue Al3+
Fe 2+
green Cu 2+
Cu 2+
green Yellow
Fe 2+ Cu 2+
Fe 2+ Fe 3+
Cu 2+
Green Fe 2+ 1
Powdery, or crystalline solid
NH4+ Pb 2+ Al 3+
Zn 2+
Colourless solution
2
237
Add water to the solid
Solid dissolves to form
NH4+
2+
Pb 2+
Cu Fe 2+ Fe 3+ Al 3+
Zn 2+
Blue solution Green Yellow Colourless solution
solution solution
238
Heat solid gently and then strongly
Brown gas that turns wet blue litmus red
239
Add aqueous ammonia drop wise until excess
Cu 2+
8
240
Add aqueous ammonia drop wise until excess
Cu2+ confirmed
9
Colourless solution
10
241
White ppt
11
242
Add potassium Iodide solution drop wise
until excess
Pb 2+
Cu 2+ gives slightly similar results but check if you
have observed a pale blue ppt with sodium
hydroxide solution 13
White ppt
SO42-
confirmed 14
243
Silver nitrate solution followed
by dilute nitric acid solution
White ppt
Cl -
confirmed 15
White ppt
244
Lead nitrate solution followed by
dilute nitric acid solution
White ppt
SO42- Cl-
17
Filtration
White residue
Colourless filtrate
18
245
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 it‘s collected in
containers. Water is evaporated off and almost 99% pure sulphur is obtained.
Extraction of sulphur
246
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.
247
Preparation of rhombic or octahedral sulphur (alpha sulphur)
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.
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
248
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.
249
25.6.2. Chemical properties of sulphur
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.
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.
250
S(s) + 2H2SO4(s) → 3SO2(g) + 2H2O(l)
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)
251
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
SO2 (g) + 2HNO3 (l) → H2SO4 (l) + 2NO2 (g)
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.
3SO2 (g) + 2H2O (l) + Dye → H2SO4 (aq) + (Dye + 2H (colourless)
252
b. Sulphur dioxide reacts with hydrogen sulphide to form a yellow deposit of
sulphur. Sulphur dioxide oxidizes hydrogen sulphide to water by supplying
oxygen and itself reduced to sulphur
2H2S (g) + SO2 (g) → 2H2O (l) + 3S (s)
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.
253
2 SO2 (g) + O2 (g) ⇌ 2 SO3 (g)
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.
The oleum is diluted with a known amount of water to give concentrated sulphuric acid.
254
25.8.1. Properties of sulphuric acid
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
2) As an oxidizing agent
255
When concentrated and hot, it acts as an oxidizing agent to both metals and non metals and it is
reduced to sulphur dioxide. It accepts electrons or supplies oxygen in its reaction.
a) It oxidizes copper to copper (II) sulphate. Reaction with zinc and iron produces similar
results.
Cu(s) + 2H2SO4(l) → CuSO4 (aq) + 2H2O(l) + SO2(g)
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.
C12H22O11(s) → 12C (s) + 11H2O(l)
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.
CuSO4 .5H2O(s) → CuSO4 (s) + 5H2O(l)
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
256
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.
257
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 purpose of adding dilute nitric acid or hydrochloric acid is to remove the carbonate ion and
sulphite ions if they are present in solution.
258
2H+(aq) + SO32-(aq) → H2O(l) + SO2(g)
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.
259
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
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
H2S(g) + 2FeCl3(aq) → 2FeCl2(aq) + 2HCl(aq) + S(s)
b. Hydrogen sulphide reacts with concentrated nitric acid to form brown fumes of
nitrogen dioxide and a yellow deposit of sulphur. Also the mixture becomes hot.
The hydrogen sulphide reduces the nitric acid to nitrogen dioxide.
2HNO3(l) + H2S(g) →2H2O(l) + 2NO2(g) + S(s)
260
2. Combustion of hydrogen sulphide
a. With plentiful supply of air
Hydrogen sulphide burns with a blue flame forming water and sulphur dioxide.
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26. Nitrogen and its compounds
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.
262
c. Direct combination of chlorine and ammonia gives nitrogen and ammonium
chloride
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
263
Note:
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.
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)
264
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.
2. It is used in food packaging, for example in crisp packets, to keep the food fresh and in
this case to prevent the crisps being compressed.
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.
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26.7. Nitrogen monoxide (nitrogen(II) oxide)
Nitrogen dioxide (the brown fumes) is produced partly by the action of the acid upon the
copper and partly by the oxidation of the main product (nitrogen monoxide) by the oxygen of
the air in the flask.
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.
266
2. Action on iron(II) sulphate solution
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
267
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)
5. Nitrogen dioxide does not burn, but supports combustion of substances whose flames are
hot enough to decompose it and so liberate free oxygen with which the substance may
combine. It supports the combustion of carbon, sulphur, phosphorus and magnesium.
2C(s) + 2NO2(g) → 2CO2(g) + N2(g)
6. Nitrogen dioxide oxidizes red hot metals and itself reduced to nitrogen.
4Cu(s) + 2NO2(g) → 4CuO(s) + N2(g)
268
2NO2(g) → 2NO(g) + O2(g)
26.9. Ammonia
269
Ammonium sulphate may be used instead of ammonium chloride.
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.
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26.9.4. Properties of ammonia
It is a colourless gas with a choking smell. It is less dense than air and thus collected by
upward delivery. It is an alkaline gas and therefore turns red litmus blue. It is the only
common alkaline gas.
The solution is only weakly alkaline because of the reversible nature of this reaction, which
results in a relatively low concentration of hydroxide ions. Ammonia gas dissolved in water is
usually known as aqueous ammonia.
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
271
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.
272
26.9.8. Combustion of ammonia
Ammonia is a good reducing agent, which means that it can be easily oxidised. Ammonia burns
with a green/yellow flame, in an atmosphere of air slightly enriched by oxygen forming nitrogen
and water.
The figure 2.8 shows how ammonia is burnt. The role of the glass wool is to distribute oxygen
evenly throughout the gas vessel.
273
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.
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26.9.11. Uses of ammonia
1. Ammonia solution is used in laundry work to remove temporary hardness.
2. Ammonia is used to manufacture ammonium sulphate and ammonium nitrate used as
fertilizers.
3. It is used in manufacture of nitric acid.
4. It is used in production of nylon.
5. It can be used as a refrigerant because it evapourates readily, removing heat from the
surrounding as it does so. It can be easily liquefied by compression.
Exercise
Ammonia
275
26.9.12. Ammonia solution
276
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.
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)
277
3. Ammonium nitrate is decomposed to nitrogen(I) oxide (dinitrogen oxide) and water.
NH4NO3(s) → N2O(g) + 2H2O(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.
Caution: Ammonium nitrate should not be heated in the laboratory because it explodes
on strong heating.
Exercise
State what would be observed and write an ionic equation for the reaction that would take
place when aqueous ammonium chloride was
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26.9.14. Reactions of ammonia solution and sodium hydroxide solution
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.
When excess aqueous ammonia is added to the blue precipitate, the precipitate dissolves to
form a deep blue solution containing complex tetraamine copper(II) ions.
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)
279
(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.
(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.
280
Mg2+(aq) + 2OH-(aq) →Mg(OH)2(s)
When a mixture of potassium nitrate and concentrated sulphuric acid is heated gently,
potassium nitrate gradually dissolves and effervescence occurs givning off nitric acid which
is condensed in another flask placed in a sink and cooled by tap water as shown in figure
2.11.
Brown fumes of nitrogen dioxide are produced during heating because of thermal
decomposition of nitric acid.
281
The experiment must be carried out in all-glass apparatus because nitric vapour attacks
rubber and cork.
Nitric acid is manufactured by the catalytic oxidation of ammonia. Ammonia and excess air
are passed over a heated platinum catalyst at about 800oC, forming nitrogen monoxide. The
reaction is exothermic.
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.
282
26.9.15.4. Properties of nitric acid
1. It is a strong acid.
2. It is a powerful oxidizing agent.
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.
283
26.9.15.6. Nitric acid as an oxidizing agent
(a) When concentrated nitric acid is added to a green solution of iron(II) sulphate and
warmed, it oxidizes it to a yellow or brown solution of iron(III) sulphate.
Fe2+(aq) → Fe3+(aq) +e-
(b) Concentrated nitric acid reacts with copper giving off nitrogen dioxide.
4HNO3(l) + Cu(s) → Cu(NO3)2(aq) + 2H2O(l) + 2NO2(g)
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)
26.10. Nitrates
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)
285
2Mg(NO3)2(s) →2MgO(s) + 4NO2(g) + O2(g)
(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.
286
26.10.2. Test for nitrates
The formula of the brown ring is FeSO4.NO. Concentrated sulphuric acid reacts with
nitrate ions to give nitric acid.
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.
287
FeSO4.NO(aq) →FeSO4(aq) + NO(g)
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
288
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.
289
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.
290
(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.
291
(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.
292
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.
293
(d) Explain why nitrogen gas collected in this experiment is not pure.
294
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.
295
b) By oxidation of concentrated hydrochloric acid with potassium permanganate solid
potassium permanganate is placed in a flask and concentrated hydrochloric acid is
dropped on to it from a tap funnel as shown in figure 3.2 Green/yellow gas is produced
which is collected over brine.
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)
296
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.
297
4. Reaction with hydrogen sulphide
Chlorine reacts with hydrogen sulphide forming a yellow particles of sulphur and white
fumes of hydrogen chloride.
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.
298
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.
299
When a thin strip of dutch metal (alloy of copper and zinc, mainly copper) is placed in a
gas-jar full of chlorine, it burns spontaneously with a green flame (due to the copper) to
form copper(II) chloride and a little of zinc chloride.
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.
300
Cl2(g) + H2O(l) → HCl(aq) + HOCl(aq)
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
301
11. Displacement reactions of chlorine
Chlorine is higher in the reactivity series than bromine and iodine and therefore can
displace them from solutions of their salts in water. This is because chlorine is more
reactive than bromine and iodine due to the fact the incoming electron is more strongly
attracted into the outer energy level of the smaller atom. The attraction force on the
electron will be greater for chlorine than for bromine and iodine, since the outer energy
level of chlorine is closer to the nucleus. As one goes down the group, the extra election
is further away from the nucleus. It will, therefore, be attracted less strongly thus the
reactivity of the halogens decreases down the group.
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.
302
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.
303
Sodium sulphate is not formed because it requires higher temperatures than is provided in
this experiment. Sodium chloride is the one most commonly used because it is cheap and
readily available.
304
27.3.2. Properties of hydrochloric acid
1. It turns blue litmus paper red.
2. It reacts with metals producing hydrogen gas. It reacts with metals above hydrogen in the
activity series.
4. It reacts with alkalis and basic oxides producing salt and water only.
NaOH(aq) + HCl(aq) →NaCl(aq) + H2O(l)
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27.3.3. Uses of hydrochloric acid
1. It is used in removal (de-scaling) of rust from iron before it is galvanized. It is also used
in cleaning metals before they are electroplated.
2. It is used in manufacture of plastics such as polychloroethene.
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.
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Nitric acid prevents the precipitation of other insoluble silver salts such as silver carbonate.
The only common insoluble chlorides are lead(II) chloride and silver chloride. Lead (II)
chloride is soluble in hot water.
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.
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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.
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(i) Bubbled in lead(II) nitrate solution.
(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.
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(ii) Explain why an aqueous solution of hydrogen chloride is an electrolyte
whereas the solution of the gas in methylbenzene is a non-electrolyte.
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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 earth‘s 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
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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).
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and collected under dry nitrogen. Nitrogen is inert under ordinary conditions and therefore
hardly reacts with sodium. This process is called Down’s process.
Na+(l) + e+ → Na(s)
28.4. Copper
The principal ores of copper are copper pyrites (CuFeS2), cuprite (Cu2O), copper(I) sulphide
(Cu2S) and malachite (CuCO3.Cu(OH)2).
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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, SiO2, 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.
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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.
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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.
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The wastes gases, mainly nitrogen and oxides of carbon, escape from the top of the furnace.
They are used in a heat exchange process of heat incoming air and so help to reduce the
energy costs of the process. Slag is used in making road foundations, phosphorus fertilizers
and cement.
317
4P(s) + 5O2(g) →2P2O5(s)
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
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28.5.5. Recycling of metals
With increasing use of metals, natural deposits of metal ores will eventually run out. The
metal ores will last longer if metals are recycled. For example, tin can be removed from scrap
food cans to be reused in making new food cans.
Metal ores can also be made to last longer by use of alternative materials. For example,
instead of using iron, plastic materials can be used in making dustbins. Many parts of cars are
now made of plastic materials. This is also advantageous in that plastic materials do not
corrode and can be moulded into much more complex shapes and textures than the metals
which they have replaced.
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.
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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.
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3. Iron is extracted by a reduction process in the blast furnace.
(a) Name two raw materials used besides the iron ore.
(b) Write the equation leading to the production of iron from its ore.
(c) State one use of iron.
(d) State what is observed when iron nail, is dropped in beaker of copper(II) sulphate
solution.
(i) Coke
(ii) Limestone
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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 n≥1.
ii) Each member differs in molecular formula from the next by CH2, 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.
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29.5. Alkanes
Alkanes are hydrocarbons with the general molecular formula, CnH2n+2, where n ≥ 1, for
example methane (CH4), ethane (C2H6), propane (C3H8), butane (C4H10), pentane (C5H12),
hexane (C6H14), heptanes (C7H16), octane (C8H18), nonane (C9H20), decane (C9H22) et.
Alkanes are saturated hydrocarbons. This means that the molecules of alkanes consist of
carbon and hydrogen atoms and single covalent bonds only. In other words, all the atoms
exert their usual combining power with other atoms. That is each carbon atom is bonded to
four other atoms. This is illustrated below.
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
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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
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׀
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.
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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
326
(b) Ch
(c) Ch
(d) Ch
(e) Ch
327
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.
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(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.
329
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.
330
Exercise
Compound Q has a molecular formula, C6H14. Give the structural formulae and names of
possible isomers of compound Q.
331
Straight-chain alkanes obtained from the fractional distillation of crude oil, are purified and
mixed with an aqueous solution of yeast. The yeast acts on the alkanes and converts them to
proteins, which are separated out and dried. The end-product is a powder which is used as
animal feed.
Fraction Uses
332
29.6.2. Cracking
Fractional distillation of crude oil yields only about 20 percent of petrol. With the demand for
petrol increasing, it has become necessary to devise a new process of obtaining it, that is, by
cracking of gas oil. Cracking is the process of breaking down the long chain hydrocarbons
into shorter-chain molecules. Large molecules of oils can be broken down into smaller
molecules of petrol and gases by cracking.
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.
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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
334
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.
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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 (CCl4) 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.
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Bromine dissolves partially in water to form a solution called bromine water which contains
hypobromous acid and hydrobromic acid.
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.
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(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, C 3H6 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).
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29.7.3. Alkynes
Alkynes are a homologous series of unsaturated hydrocarbons of the general formula, CnH2n-
2, where n ≥ 2. They contain carbon-carbon triple bond (– C ≡ C –) as the functional group.
They are unsaturated compounds.
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
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29.7.4.2. Chemical properties of ethyne
1. Combustion
Ethyne burns in air with a yellow sooty flame forming carbon dioxide and steam. The reaction is
exothermic. The soot is due to unburnt carbon because of its high content in ethyne.
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
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29.7.4.3. Uses of ethyne
(i) Used in manufacture of polyvinyl chloride plastic which has a wide variety of uses.
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
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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.
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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.
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It is sometimes called elimination reaction because a molecule of water is eliminated from
the alcohol to form an alkene.
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.
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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).
(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)
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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.
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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.
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.
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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.
(ii) Synthetic detergents don‘t 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.
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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.
There are two common types of polyethene depending on the conditions provided during
their manufacture.
Here the reaction involves heating ethene at about 200oC 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 60oC 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.
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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.
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Polymer Monomer Type of polymerization
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.
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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 vulvanisation 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.
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3. They are usually stronger than the corresponding natural material.
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.
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(ii) What is the purpose of adding saturated sodium chloride solution?
(b) Sometimes when soap is used for washing clothes, a scum is formed.
(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.
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(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
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(a) Name substances P and Q.
(b) write the structural formula of P
(c) Name the reagent used in step 1.
(d) State the products for complete combustion of Q in excess air.
(e) Name the catalyst used in step 2.
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30. Energy changes
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.
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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.
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Hydrogen iodide has more energy than the starting materials, hydrogen and iodine (figure
10.4). Therefore the value of enthalpy change is positive.
Exercise
How much heat is given out when 20 g of carbon are completely burnt?
Solution:
12
12
= 655 kJ
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30.4.1. Determination of enthalpy (heat) of combustion of ethanol
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 bottle‘s 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:
Calculations:
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Where m = mass of the water.
∆θ = temperature rise.
m = 1 x 100 = 100 g
∆θ = 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
361
1. The following reaction takes place when methanol is burnt in oxygen.
2CH3OH(l) + 3O2(g) → 2CO2(g) + 4H2O(g) ∆H = -1452 Kj mol-1
(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.
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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 = ∆θ
Assuming the specific heat capacity of the solution is 4.2 J g-1 oC-1
58.5
= 0.53 mol
0.53
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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:
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Temperature of 1 M hydrochloric acid, θ2 = 24 oC
Calculations:
2 2
= 24.5 oC
= 2658.6 J or 2.66 kJ
1000
= 0.05 mol
0.05
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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.
(ii) Heat evolved when 80 cm3 of 1 M sodium hydroxide react with 80 cm3 of 1 M
nitric acid.
2. Aqueous hydrogen ions react with aqueous hydroxide ions according to the equation.
H+(aq) + OH-(aq) – H2O(l) ∆H = 57 kJ mol-1
Specimen Results:
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Intial temperature = 25.0 oC
Calculations:
1000
= 0.025 mol
0.025
= 29400 J or 29.4 kJ
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Exercise 10.5
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.
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(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 cm3 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:
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Mass of the spirit lamp + content after heating = 34.07 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.
370
31. Rate of reaction and equilibrium
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.
371
A graph of volume of hydrogen evolved against time is plotted. A typical graph has the form
of figure 11.2.
To determine the rate of reaction at a given time, say t1, 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.
372
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 cm3 of 0.05 M sodium
thiosulphate solution into a beaker. Add 10 cm3 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
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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
374
(i) 0.5 M sulphuric acid.
(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.
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Figure 11.5a shows that the higher the temperature the less the time taken to form a
precipitate. Figure 11.5b shows that the rate of the reaction increases with increase in
temperature.
Exercise
(ii) 40 oC.
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manganese(IV) oxide. When the graphs of volume of oxygen against time are plotted
using the same axes, they appear as shown in figure 11.6.
Exercise
(b) On the same axis, sketch a graph of volume of oxygen evolved against time when
377
31.1.5. Effect of surface area on the rate of reaction
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.
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(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.
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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 chatelier’s 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
380
The forward reaction is exothermic and therefore the backward reaction is endothermic.
If heat is supplied, the equilibrium shifts in the direction which requires more heat, that
is, the backward reaction which uses up the excess heat occurs. However, if the
equilibrium vessel is cooled, the equilibrium shifts to the right, producing more ammonia.
(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
381
What will be the effect on the concentration of hydrogen iodide and ammonia in the
equilibrium mixture of
(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,
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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.
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(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?
3. 12.0 g of clean magnesium ribbon were added to 50 cm3 of 1 M sulphuric acid. The
volume of the gas evolved was measured at fixed time interval until the reaction stopped.
(a) Write the equation of the reaction that took place.
(b) (i) sketch a graph of volume of the gas (on y-axis) against time (on-axis). Label the
graph G1.
(ii) On the same axis sketch the graph that would be obtained if 12.0 g of magnesium
powder were used instead of magnesium ribbon. Label this graph G2.
(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.
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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 50cm 3 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 cm3 of 1 M
hydrochloric acid at 25 oC.
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(ii) 1.0 g of powdered calcium carbonate was reacted with 50 cm 3 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 cm3 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.
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