The Complete General Science Notes (Chemistry) for Railway Exams

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 The Complete General Science Notes (Chemistry) for Railway Exams

General Science - Chemistry

S.NO Content Page Number

1 Chemical reaction and equation 2

2 Matter 9
3 Atoms and molecules 16

4 Carbon and its compounds 21

5 Periodic Classification of Elements 25

6 Coal and petroleum 32

7 Acids Bases and Salts 34

8 Metals and Nonmetals 39
9 Important chemicals and their uses 44

10 Important Chemistry one liners 45

CHEMICAL REACTIONS AND EQUATIONS

CHEMICAL REACTION & EQUATIONS
 A complete chemical equation represents the reactants, products and their physical
State symbolically.
 Following observations helps us to determine whether a chemical reaction has taken place
 Change in state
 Change in colour
 Evolution of a gas
 Change in temperature.
 Some of the example of chemical reactions in everyday life
 Photosynthesis
 Aerobic Cellular Respiration

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 The Complete General Science Notes (Chemistry) for Railway Exams
 Combustion of wood
 Rusting of iron
 Metathesis
 Digestion
 Cooking an egg
 Souring of milk
 Rotting bananas
 Exothermic reactions are reactions or processes that release energy, usually in the form of heat or light
 Reactions in which energy is absorbed are known as endothermic reactions.
BALANCED CHEMICAL EQUATIONS
 Mass can neither be created nor destroyed in a chemical reaction. That is, the total mass of the elements
present in the products of a chemical reaction has to be equal to the total mass of the elements present in the
reactants
 The number of atoms of each element remains the same, before and after a chemical reaction
Some of the example of balanced equations
Zn + H2SO4-ZnSO4 + H2
3Fe + 4H2O Fe3O4 + 4H2

TYPES OF CHEMICAL REACTIONS

1. COMBINATION REACTION
 In a combination reaction two or more substances combine to form a new single Substance.
Example of combination reaction
 Calcium oxide reacts vigorously with water to produce slaked lime(calcium hydroxide) Releasing a large
amount of heat

CaO (s) + H2O (l) -->Ca (OH)2 + Heat

 A solution of slaked lime produced by the above reaction is used for white washing wall

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 The Complete General Science Notes (Chemistry) for Railway Exams
NOTE: Calcium hydroxide reacts slowly with the carbon dioxide in air to form a thin layer of calcium
carbonate on the walls. Calcium carbonate is formed after two to three days of whitewashing and gives a
shiny finish to the walls. It is interesting to note that the chemical formula for marble is also CaCO3.

2. DECOMPOSITION REACTION
 Decomposition reactions are opposite to combination reactions. In a decomposition reaction, a single
substance decomposes to give two or more substances
 In this reaction, you can observe that a single reactant breaks down to give simpler products. This is a
decomposition reaction.

2FeSO4(s) Heat---->Fe2O3(s) + SO2(g) + SO3(g)

 Decomposition of Silver bromide into silver and chlorine by light.

 Silver bromide used in black and white photography

 Decomposition of calcium carbonate to calcium oxide and carbon dioxide on heating is an important
decomposition reaction used in various industries. Calcium oxide is called lime or quick lime. It has many
uses – one is in the manufacture of cement. When a decomposition reaction is carried out by heating, it
is called thermal decomposition

3. DISPLACEMENT RECTION
 It is a reaction between an element and a compound. When they react, one of the elements of the compound-
reactant is replaced by the element-reactant to form a new compound and an element.
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 The Complete General Science Notes (Chemistry) for Railway Exams
Fe+ CuSO4 --> FeSO4 + Cu
 In this reaction, iron has displaced or removed another element copper from copper sulphate solution. This
reaction is known as displacement reaction
4. DOUBLE DISPLACEMENT REACTIONS
 When two compounds react, if their ions are interchanged, then the reaction is called double displacement
reaction. The ion of one compound is replaced by the ion of the another compound
Na2SO4 (aq) + BaCl2 (aq) → BaSO4 (s) + 2NaCl(aq)
PRECIPITATION REACTIONS
 When aqueous solutions of two compounds are mixed, if they react to form an insoluble compound and a
soluble compound, then it is called precipitation reaction. Because the insoluble compound, formed as one of
the products, is a precipitate and hence the reaction is so called.
 Precipitation reactions produce insoluble salts.
NEUTRALISATION REACTION
 When an acid and a base react together to form salt and water known as neutralization reaction
HCl + NaOH → H2O + NaCl

OXIDATION AND REDUCTION

OXIDATION
 The chemical reaction which involves addition of oxygen or removal of hydrogen or loss of electrons is
called oxidation.
2 Mg + O2 →2 MgO (addition of oxygen)
CaH2 → Ca + H2 (removal of hydrogen)
Fe2+→Fe3+ + e− (loss of electron)
REDUCTION
 The chemical reaction which involves addition of hydrogen or removal of oxygen or gain of electrons is
called reduction.
2 Na + H2 → 2 NaH (addition of hydrogen)
CuO + H2 → Cu + H2 O (removal of oxygen)
Fe3++ e−→Fe2+ (gain of electron)
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 The Complete General Science Notes (Chemistry) for Railway Exams
REDOX REACTIONS
 Generally, the oxidation and reduction occurs in the same reaction (simultaneously). If one reactant gets
oxidized, the other gets reduced. Such reactions are called oxidation-reduction reactions or Redox reactions.
2 PbO + C → 2 Pb + CO2
Zn + CuSO4 → Cu + ZnSO4

Oxidation Reduction
Addition of oxygen Removal of oxygen
Removal of hydrogen Addition of hydrogen
Loss of electron Gain of electron

OXIDATION AND REDUCTION AGENTS

 Substance that loses oxygen or gains hydrogen is known as an oxidizing agent
 Substance that loses hydrogen or gains oxygen is known as a reducing agent
 Compounds with oxygen atom are called oxidizing agent and compounds with hydrogen atom are called
reducing agent
 Some compounds can act as either oxidizing agents or reducing agents. One example is hydrogen gas, which
acts as an oxidizing agent when it combines with metals and As a reducing agent when it reacts with non-
metals.
Oxidation reactions in daily life
 The shining surface of metals tarnishes due to the formation of respective metal oxides on their surfaces. This
is called corrosion.
 The freshly cut surfaces of vegetables and fruits turn brown after some time because of the oxidation of
compounds present in them
RANCIDITY
 When oils and fats or foods containing oils and fats are exposed to air, they get oxidized due to which the
food becomes stale and gives a bad taste or smell. This is called Rancidity.
Following ways to preventing rancidity
 Adding antioxidants
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 The Complete General Science Notes (Chemistry) for Railway Exams
 Refrigerating
 Storing food in airtight containers with nitrogen gas

COMMON NAME AND FORMULA OF CHEMICAL COMPOUNDS

Chemical Compounds Chemical formula Common names
Calcium oxide Cao Quick lime
Calcium hydroxide Ca(OH)2 Slaked lime
Calcium carbonate CaCO3 Limestone
Trichloro Methane CHCl3 Chloroform
Calcium Oxychloride CaOCl2 Bleaching powder
sodium hydrogencarbonate NaHCO3 Baking soda
Sodium carbonate Na2CO3 Washing soda
Calcium sulphate CaSO4 .1/2H2O Plaster of paris
hemihydrate
calcium sulfate dihydrate CaSO .2H2O Gypsum
Acetic acid CH3COOH Vinegar
Silicon Oxide SiO2 Sand

Methane CH4 Marsh Gas

Nitrous oxide N2O Laughing Gas

Deuterium Oxide D2O Heavy water

Solid Carbondixide CO2 Dry ice

Calcium Carbonate CaCo3 Chalk

Sulphuric Acid H2SO4 Oil of vitriol

Zinc sulphate ZnSO4 White Vitriol

Copper sulphate CuSO4.5H2O Blue Vitriol

Sodium hydroxide NaOH Caustic Soda

Potassium carbonate K2CO3 Potash Ash

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Mercurous chloride Hg2Cl2 Calomel
Sucrose C12H22O11 Sugar
Silver nitrate AgNO3 Lunar caustic
Ethyl Alcohol C2H6O Alcohol
Hydrochloric acid HCl Muriatic acid

CHEMICAL COMPOUNDS AND FORMULA

Chemical Compounds Chemical formula
Sodium chloride NaCl
Zinc sulphate ZnSO4
Glucose C6H12O6
Ferric oxide Fe2O3
Ferrous sulphate FeSO4
Lead oxide PbO
Lead nitrate Pb(NO3)2
silver chloride AgCl
Silver bromide AgBr
Sodium sulphate Na2SO4

CHEMICAL BONDING
Attraction between atoms, ions or molecules that enables the formation of chemical compounds is called
chemical bonding
TYPES OF CHEMICAL BONDING
1. Ionic bond
Chemical bond formed between two atoms due to transfer of electron from one atom to the other atom
2. Covalent bond
A covalent bond is a chemical bond that involves the sharing of electron between two atoms
3. Metallic bond
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 The Complete General Science Notes (Chemistry) for Railway Exams
Metallic bond is the force of attraction between metal ions to a number of electrons within its sphere of
influence.

MATTER
 Matter is made up of small particles
 The matter around us exists in three states— solid, liquid and gas.
 The forces of attraction between the particles are maximum in solids, intermediate in liquids and minimum in
gases
 The spaces in between the constituent particles and kinetic energy of the particles are minimum in the case of
solids, intermediate in liquids and maximum in gases
 Particles of matter are continuously moving, that is, they possess what we call the kinetic energy. As the
temperature rises, particles move faster. So, we can say that with increase in temperature the kinetic energy
of the particles also increases
 The states of matter are inter-convertible. The state of matter can be changed by changing temperature or
pressure.
DIFFUSION
 The mixing of a substance with another substance due to the motion of its particles is called diffusion. It is
one of the properties of material. The diffusion of one substance to another substance goes on until a uniform
mixture is formed. Diffusion takes place in gases, liquids and solids. Diffusion increases on increasing the
temperature of the diffusing substance.
STATES OF MATTER
 Matter around us exists in three different states– solid, liquid and gas. These states of matter arise due to the
variation in the characteristics of the particles of matter
1. THE SOLID STATE
 Solid have a definite shape, distinct boundaries and fixed volumes, that is, have negligible compressibility.
Solids have a tendency to maintain their shape when subjected to outside force. Solids may break under force
but it is difficult to change their shape, so they are rigid.
2. THE LIQUID STATE

 Liquids have no fixed shape but have a fixed volume. They take up the shape of the container in which they
are kept. Liquids flow and change shape, so they are not rigid but can be called fluid
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 The rate of diffusion of liquids is higher than that of solids
 Particles move freely and have greater space between each other as compared to particles in the solid state
3. THE GASEOUSSTATE
 Gases are highly compressible as compared to solids and liquids
 Gases have lower density than other states of matters
 The liquefied petroleum gas (LPG) cylinder that we get in our home for cooking or the oxygen supplied to
hospitals in cylinders is compressed gas
 The oxygen supplied to hospitals in cylinders is compressed gas.
 Compressed natural gas (CNG) is used as fuel these days in vehicles.
 The rate of diffusion of gas is higher than that of solids and liquids

 We come to know of what is being cooked in the kitchen without even entering there, the smell of hot cooked
food reaches us in seconds because rate of diffusion of gas is higher than that of solids and liquids.

MATTERS CHANGE ITS STATE?

Water can exist in three states of matter–
• Solid, as ice,
• Liquid, as the familiar water, and
• Gas, as water vapour.

1. EFFECT OF CHANGE OF TEMPARATURE

Increasing the temperature of solids, the kinetic energy of the particles increases.Due to the increase in kinetic
energy, the particles start vibrating with greater speed. The energy supplied by heat overcomes the forces of attraction
between the particles. The particles leave their fixed positions and start moving more freely. A stage is reached when
the solid melts and is converted to a liquid. The minimum temperature at which a solid melts to become a liquid at
the atmospheric pressure is called its melting point
 The melting point of ice is 273.15 K
Supply heat energy to water, particles start moving even faster. At a certain temperature, a point is reached when the
particles have enough energy to break free from the forces of attraction of each other. At this temperature the liquid

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starts changing into gas. The temperature at which a liquid starts boiling at the atmospheric pressure is known as its
boiling point

 State of matter can be changed into another state by changing the temperature

LATENT HEAT
The heat energy required to convert a solid into a liquid or vapour, or a liquid into a vapour, without change of
temperature known as latent heat

2. EFFECT OF CHANGE OF PRESSURE

 Increasing or decreasing the pressure can change the state of matter
 Pressure and temperature determine the state of a substance, whether it will be solid, liquid or gas
 Gases can be liquefied by applying pressure and lowering temperature and liquid also convert to solid by
applying the pressure and lowering the temperature

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 Atmosphere (atm) is a unit of measuring pressure exerted by a gas

 The unit of pressure is Pascal (1 atmosphere = 1.01 * 105 Pa)
Solid carbon dioxide
 It is stored under high pressure.
 Solid CO2 gets converted directly to gaseous state on decrease of pressure to 1 atmosphere* without coming
into liquid state. This is the reason that solid carbon dioxide is also known as dry ice
Sublimation
Sublimation is the change of solid state directly to gaseous state without going through liquid state.
Evaporation
Evaporation is a surface phenomenon. Particles from the surface gain enough energy to overcome the forces
of attraction present in the liquid and change into the vapour state.
Rate of evaporation depends upon the surface area exposed to the atmosphere, the temperature, the humidity and the
wind speed.

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Some measurable quantities and their units

Quantity Unit
Temperature Kelvin
Length Metre
Mass Kilogram
Weight Newton
Volume Cubic Metre
Density kilogram per cubic metre
Pressure Pascal

IS MATTER AROUND US PURE

 Depending upon the chemical composition, matter is classified into elements, compounds and mixtures
 A mixture contains more than one substance mixed in any proportion
 Air is a mixture of nitrogen, oxygen, carbon dioxide, water vapour and other gases. Soil is a mixture of clay,
sand and various salts. Milk, ice cream, rock salt, tea, smoke, wood, sea water, blood, tooth paste and paint
are some other examples of mixtures. Alloys are mixtures of metals.
 Mixtures can be separated into pure substances using appropriate separation techniques
TYPES OF MIXTURES
1. Homogeneous mixture

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2. Heterogeneous mixture
Homogeneous mixture
1. A mixture in which the components cannot be seen separately is called a homogeneous mixture.
2. It has a uniform composition and every part of the mixture has the same properties
3. Tap water, milk, air, ice cream, sugar syrup, ink, steel, bronze and salt solutions are homogeneous mixtures
Heterogeneous mixture
1. A mixture in which the components can be seen separately is called a heterogeneous mixture.
2. It does not have a uniform composition and properties.
3. Soil, a mixture of iodine and common salt, a mixture of sugar and sand, a mixture of oil and water, a mixture
of sulphur and iron filings and a mixture of milk and cereals are heterogeneous mixture.
SOLUTIONS
 A solution is a homogeneous mixture of two or more substances. You come across various types of solutions
in your daily life. Lemonade, soda water etc.
 We can also have solid solutions (alloys) and gaseous solutions (air)
 The particles of a solution are smaller than 1 nm (10-9 metre) in diameter. So, they cannot be seen by naked
eyes
 Because of very small particle size, they do not scatter a beam of light passing through the solution. So, the
path of light is not visible in a solution
 Solution has a solvent and a solute as its components.
SUSPENSION
 Suspension is a heterogeneous mixture
 The particles of a suspension scatter a beam of light passing through it and make its path visible.
 The particles of a suspension can be seen by the naked eye.
COLLOIDS
 A colloid is a heterogeneous mixture.

 The size of particles of a colloid is too small to be individually seen by naked eyes.
 Colloids are big enough to scatter a beam of light passing through it and make its path visible.
EXAMPLES OF COLLOIDS

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Dispersed phase Dispersing Medium Type Example
Liquid Gas Aerosol Fog, clouds, mist
Solid Gas Aerosol Smoke, automobile exhaust
Gas Liquid Foam Shaving cream
Liquid Liquid Emulsion Milk, face cream
Gas Solid Foam Foam, rubber, sponge, pumice
Solid Liquid Sol Milk of magnesia, mud
Liquid Solid Gel Jelly, cheese, butter
Solid Solid Solid Sol Coloured gemstone, milky glass

SEPARATING THE COMPONENTS OF A MIXTURE

 Separate the volatile component (solvent) from its non-volatile solute by the method of evaporation.
Applications:
1. Ink is a mixture of a dye in water
 Centrifugation is the process by which fine insoluble solids from a solid- liquid mixture can be separated in
a machine called a centrifuge. A centrifuge rotates at a very high speed. On being rotated by centrifugal
force, the heavier solid particles move down and the lighter liquid remains at the top.
Applications:
1. Used in diagnostic laboratories for blood and urine tests.
2. Used in dairies and home to separate butter from cream.
3. Used in washing machines to squeeze out water from wet clothes
 Separation of components of a mixture containing two miscible liquids that boil without decomposition and
have sufficient difference in their boiling points this method is called distillation
Applications:
1. Salt water turned to fresh water using distillation process
 The crystallization method is used to purify solids
Applications:
1. Purification of salt that we get from sea water.
2. Separation of crystals of alum from impure samples.
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 Chromatography is a separation technique. It is used to separate different components of a mixture based on
their different solubilities in the same solvent
Applications
1. To separate colours in a dye
2. To separate pigments from natural colours
3. To separate drugs from blood.
HOMOGENEOUS & HETEROGENEOUS MIXTURE
Homogeneous mixture Heterogeneous mixture
Consists of single phase Consists of two or more phases

Has the same uniform appearance and Has different non uniform appearance and
composition composition
Components are unrecognizable Components are recognizable

Examples: Air, saline solution and Example: Sand, oil and water
bitumen

ATOMS AND MOLECULES

 Antoine L. Lavoisier laid the foundation of chemical sciences by establishing two important laws of chemical
combination.
LAW OF CONSERVATION OF MASS
During a chemical reaction, the sum of the masses of the reactants and products remains unchanged. This is
known as the Law of Conservation of Mass

LAW OF CONSTANT PROPORTIONS

This law was stated by Proust as “In a chemical substance the elements are always present in definite
proportions by mass”. This Law known as the Law of Definite Proportions or Law of definite proportions.
 Dalton’s atomic theory provided an explanation for the law of conservation of mass and the law of definite
proportions.
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According to Dalton’s atomic theory
 All matter is made of very tiny particles called atoms, which participate in chemical reactions
 Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction
 Atoms of a given element are identical in mass and chemical properties.
 Atoms of different elements have different masses and chemical properties

ATOMS
 An atom is the smallest particle of an element that can take part in a chemical reaction
 Atomic radius is measured in nanometers.
1/109 m = 1 nm
1 m = 109 nm
 Hydrogen atom is smallest atom of all. Atomic radius of hydrogen atom is
0.037 x 10-9
Radii Example
10–10 Atom of hydrogen
10–9 Molecule of water
10–8 Molecule of hemoglobin
10–4 Grain of sand
10–3 Ant
10–1 Apple
ATOMIC MASS
 Atomic mass is defined as the mass of a single atom of a chemical element
 One atomic mass unit is a mass unit equal to exactly one-twelfth (1/12th) the mass of one atom of carbon-12.
The relative atomic masses of all elements have been found with respect to an atom of carbon-12
ATOMIC MASS OF SOME ELEMENTS
Element Atomic mass
Hydrogen 1

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Carbon 12
Nitrogen 14
Oxygen 16
Sodium 23
Magnesium 24
Sulphur 32
Chlorine 35.5
Calcium 40

Avogadro constant
 The Avogadro constant 6.022 × 1023 is defined as the number of atoms in exactly 12 g of carbon-12.

THE STRUCTURE OF AN ATOM

 J.J. Thomson was the first one to propose a Model for the structure of an atom.
 Thomson proposed that:
(i) An atom consists of a positively charged sphere and the electrons are embedded in it.
(ii) The negative and positive charges are equal in magnitude. So, the atom as a whole is electrically neutral
 Rutherford’s model of the atom proposed that a very tiny nucleus is present inside the atom and electrons
revolve around this nucleus. The stability of the atom could not be explained by this model
 Neils Bohr’s model of the atom was more successful. He proposed that electrons are distributed in different
shells with discrete energy around the nucleus. If the atomic shells are complete, then the atom will be stable
and less reactive.
 Electron was discovered by JJ Thomson
 Proton was discovered by Rutherford
RUTHERFORD’S ATOMIC MODEL
According to this model:
1. The atom contains large empty space.
2. There is a positively charged mass at the centre of the atom, known as nucleus.

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3. The size of the nucleus of an atom is very small compared to the size of an atom.
4. The electrons revolve around the nucleus in close circular paths called orbits.
5. An atom as a whole is electrically neutral, i.e., the number of protons and electrons in an atom are equal.
BOHR’S MODEL OF AN ATOM
 In order to overcome the objections raised against Rutherford’s model of the atom, Neils Bohr put forward the
following postulates about the model of an atom:
1. Only certain special orbits known as discrete orbits of electrons, are allowed inside the atom.
2. While revolving in discrete orbits the electrons do not radiate energy.
 These orbits or shells are called energy levels
NEUTRONS
 J. Chadwick discovered the neutron
 Neutrons are present in the nucleus of all atoms, except hydrogen
 Mass of an atom equal to sum of the masses of protons and neutrons present in the nucleus
ELECTRONS DISTRIBUTED IN DIFFERENT ORBITS (SHELLS)
 Distribution of electrons into different orbits of an atom was suggested by Bohr and Bury.
 Maximum number of electrons present in a shell is given by the formula 2n2
Where n=1,2,3,4,….
 These orbits or shells are represented by the letters K,L,M,N,…
 The maximum number of electrons that can be accommodated in the outermost orbit is 8.
 Electrons are not accommodated in a given shell, unless the inner shells are filled. That is, the shells are filled
in a step-wise manner.

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VALENCE ELECTRONS
 Electrons present in the outermost shell of an atom are known as the valence electrons
 The elements with same number of electrons in the valence shell show similar properties and those with
different number of valence electrons show different chemical properties
 Elements, which have 1 or 2 or 3 valence electrons (except Hydrogen), are metals.
 Elements with 4 to 7 electrons in their valence shell are non-metals.
VALANCY
 Valency of an element is the combining capacity of the element with other elements and is equal to the
number of electrons that take part in a chemical reaction
 Valency of the elements having valence electrons 1, 2, 3, 4 is 1, 2, 3, 4 respectively
 Valency of an element with 5, 6 and 7 valence electrons is 3, 2 and 1 (8–valence electrons) respectively.
Because 8 is the number of electrons required by an element to attain stable electronic configuration
 Elements having completely filled outermost shell show Zero valency
ATOMIC NUMBER
 Atomic number of an element is the same as the number of protons in the nucleus of its atom.
MASS NUMBER
 Mass number of an atom is equal to the number of protons and neutrons in a nucleus
ISOTOPES
 Two or more forms of an element having the same atomic number, but different mass number are called
Isotopes (17Cl35, 17Cl37).
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 Applications
1. An isotope of uranium is used as a fuel in nuclear reactors.
2. An isotope of cobalt is used in the treatment of cancer.
3. An isotope of iodine is used in the treatment of goitre.
ISOBARS
 Atoms of different elements having the same mass number, but different atomic numbers are called Isobars
(18Ar40, 20Ca40).
ISOTONES
 Atoms of different elements having the same number of neutrons, but different atomic number and different
mass number are called Isotones ( 6C13, 7N14 ).

CARBON AND ITS COMPOUNDS

 All living structures are carbon based.
 Carbon is found both in free state as well as combined state in nature
 Earth’s crust has only 0.02% carbon in the form of minerals like carbonates, hydro carbonates,coal and
petroleum and the atmosphere has 0.03% of carbon dioxide. In spite of this small amount of carbon available
in nature
 Both diamond and graphite are formed by carbon atoms they are allotrope of carbon
 The gas/kerosene stove used at home has inlets for air so that a sufficiently oxygen-rich mixture is burnt to
give a clean blue flame.
 If bottoms of cooking vessels getting blackened, it means that the air holes are blocked and fuel is getting
wasted
 Cooking Gas mainly consist of Butane
 Ethanol is used as a fuel in cars along with petrol
 Ethyl alcohol is used as an antiseptic to sterilize wounds and syringes in hospitals
 Methane popularly known as marsh gas. Natural consists of over 90 percent methane and some amount of
propane and butane
 Paddy field is biggest source of methane gas
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 Bio gas consists of 55to 70 percent methane and 30 to 45 percent carbon
 Ethylene used for
1. Preparation of mustard gas
2. Preservation and artificial ripening of green fruits
3. Manufacturing of PVC pipes
COVALENT BOND
 Carbon always have a covalent bond
 The bond formed by sharing of electrons between two atoms are known as covalent bond
 The boiling and melting points of the carbon compounds is low
 Most carbon compounds are poor conductors of electricity because they form covalent bond so it does not
give rise to free electrons. All electrons are used in making the covalent bond
 Graphite is a good conductor of heat and electricity because it has free electrons
Melting & boiling points compounds of carbon
Compound Formula Melting point (K) Boiling point (K)
Acetic acid CH3COOH 290 391
Chloroform CHCl3 209 334
Ethanol CH3CH2OH 156 351
Methane CH4 90 111
Allotropes of carbon
 Allotropy is a property by which an element can exist in more than one form that are physically different and
chemically similar. The different forms of that element are called its allotropes
 The element carbon occurs in different forms in nature with widely varying physical properties. Both diamond
and graphite are formed by carbon atoms, the difference lies in the manner in which the carbon atoms are
bonded to one another
 Carbon exists in different allotropic forms and based on their physical nature they are classified as below.
Crystalline forms of Carbon
1. Diamond
2. Graphite

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3. Fullerene
Amorphous forms of carbon
1. Charcoal
2. coke
3. Lamp black
4. Gas carbon
SOME FUCTIONAL GROUPS OF CARBON COMPOUNDS
Hetero atom Class of compounds Formula of functional group
Cl/Br Halo- (Chloro/bromo) —Cl, —Br (substitutes for
Alkane hydrogen atom)

Oxygen 1. Alcohol —OH

2. Aldehyde

3. Ketone

4. Carboxylic acid

FORMULA OF SATURATED COMPOUNDS OF CARBON AND HYDROGEN’S

No of C atoms Name Formula

1 Methane CH4
2 Ethane C2H6
3 Propane C3H8
4 Butane C4H10
5 Pentane C5H12

6 Hexane C6H14

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7 Heptane C7H16
8 Octane C8H18
9 Nonane C9H20
10 Decane C10H22

ETHANOL
 Ethanol is commonly known as alcohol. All alcoholic beverages and some cough syrups contain ethanol. Its
molecular formula is C2H5OH
 Ethanol and ethanoic acid are carbon compounds of importance in our daily lives
 Ethanol is a liquid at room temperature. Ethanol is commonly called alcohol and is the active ingredient of all
alcoholic drinks
 Ethanol is a colourless liquid, having a pleasant smell and a burning taste.
 Ethanol is used as an anti-freeze in automobile radiators
 Ethanol is used in medical wipes, as an antiseptic
 Ethanol is a good solvent, it is also used in medicines such as tincture iodine, cough syrups, and many tonics
 Ethanol is used for effectively killing microorganisms like bacteria, fungi, etc., by including it in many hand
sanitizers.
ETHANOIC ACID
 Ethanoic acid or acetic acid is one of the most important members of the carboxylic acid family. Its molecular
formula is C2H4O2.
 Ethanoic acid is commonly called acetic acid and belongs to a group of acids called carboxylic acids
 5-8% solution of acetic acid in water is called vinegar and is used widely as a preservative in pickles.
 Ethanoic acid is used in printing on fabrics
 The melting point of pure ethanoic acid is 290k and hence it often freezes during winter. They look like
glaciers, so it is called glacial acetic acid
SOAPS & DETERGENTS
 Soap is a sodium or potassium salt of long chain carboxylic acid
 Soap is effective only in soft water

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 Detergent is ammonium or sulphonate salt of long chain of carboxylic acid
 Detergent are effective both soft and hard water
ORGANIC COMPOUNDS IN DAILY LIFE
 Organic compounds are inseparable in human life
 Various classes of organic compounds and their uses in our daily life as follows:
Hydrocarbons
1. Fuels like LPG, Petrol, Kerosene.
2. Raw materials for various important synthetic materials.
3. Polymeric materials like tyre, plastic containers.
Alcohols
1. As a solvent and an antiseptic agent.
2. Raw materials for various important synthetic materials.
Aldehydes
1. Formaldehyde as a disinfectant.
2. Raw materials for synthetic materials.
Ketones
1. As a solvent.
2. Stain Remover.

PERIODIC CLASSIFICATION OF ELEMENTS

 In 1800, there were only 31 known elements. By 1865, their number became 63. Now 118 elements have been
discovered.
 Presently, 118 elements are known. All these have different Properties. Out of these 118, only 94 are naturally
occurring.
 All the elements are unique in their nature and property. To categorize these elements according to their
properties, scientists started to look for a way.
 Scientists made several attempts to classify elements according to their properties Such as Newlands Law of
Octaves, Dobereiner triads Law and Mendeleev

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DOBEREINER’STRIADS LAW
 He tried to arrange the elements with similar properties into groups. He identified some groups having three
elements each. So he called these groups ‘triads’ Dobereiner showed that when the three elements in a triad
were written in the order of increasing atomic masses. The atomic mass of the middle element was roughly
the average of the atomic masses of the other two elements
 Example: In the triad group (1), arithmetic mean of atomic masses of 1st and 3rd elements, (6.9 + 39.1)/2 =
23. So the atomic mass of Na (middle element) is 23.
Limitations
 Dobereiner could identify only three triads from the elements known at that time and all elements could not be
classified in the form of triads. „
 The law was not applicable to elements having very low and very high atomic mass.
Newlands Law of Octaves
 In 1866, John Newlands arranged 56 known elements in the increasing order of their atomic mass.
 He started with the element having the lowest atomic mass (hydrogen) and ended at thorium which was the
56th element. He found that every eighth element had properties similar to that of the first. This arrangement
was known as 'law of octaves'
 Law of Octaves was applicable only upto calcium, as after calcium every eighth element did not possess
properties similar to that of the first.
 Newlands’ Law of Octaves worked well with lighter elements only
 Newlands’ table was restricted to only 56 elements and did not leave any room for new elements
Mendeleev periodic table
 At the time of Mendeleev started his work, 63 elements were known. He examined the relationship between
the atomic masses of the elements and their physical and chemical properties
 He observed that most of the elements got a place in a Periodic Table and were arranged in the order of their
increasing atomic masses
 Mendeleev’s Periodic Table contains vertical columns called ‘groups’ and horizontal rows called ‘periods’
 It has eight vertical columns called ‘groups’ and seven horizontal rows called ‘period’.
Limitations
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 The increasing order of atomic mass was not strictly followed throughout.
Eg. Co & Ni, Te & I.
 No place for isotopes in the periodic table
 No proper position could be given to the element hydrogen. Non-metallic hydrogen was placed along with
metals like lithium (Li), sodium (Na) and potassium (K).
Modern Periodic Table
 Elements are arranged in order of increasing atomic number
 Modern Periodic Law can be stated as follows “The chemical and physical properties of the elements are the
periodic functions of their atomic numbers”. Based on the modern periodic law, the modern periodic table is
derived
 Modern Periodic Table has 18 vertical columns known as groups and 7 horizontal rows known as ‘periods’.
 Elements present in any one group have the same number of valence electrons
 The valence of an element is determined by the number of valence electrons present in the outermost shell of
its atom
 Metals are found on the left-hand side of the Periodic Table
 Non-metals are found on the right-hand side of the Periodic Table
 Modern Periodic Table, a zig-zag line separates metals from non-metals. The borderline elements boron,
silicon, germanium, arsenic, antimony, tellurium and polonium are intermediate in properties and are called
metalloids
 Halogens are located on the 17th group on the periodic table
 Noble gases are located on the 18th group on the periodic table
 Based on the physical and chemical properties of elements, they are grouped into various families.
Groups in modern periodic table
Group 1 Alkali metals
Group 2 Alkaline earth metals
Group 3 o 12 Transition metals
Group 13 Boron Family
Group 14 Carbon Family

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Group 15 Nitrogen Family
Group 16 Oxygen Family (or) Chalcogen Family
Group 17 Halogens
Group 18 Noble gases

Modern Periodic Table

Position of hydrogen in the periodic table
 Hydrogen is the lightest, smallest and first element of the periodic table. Its electronic configuration (1s1) is the simplest
of all the elements.
 It occupies a unique position in the periodic table. It behaves like alkali metals as well as halogens in its properties
 In the periodic table, it is placed at the top of the alkali metals.
1. Hydrogen can lose its only one electron to form a hydrogen ion (H+) like alkali metals.
2. It can also gain one electron to form the hydride ion (H- ) like halogens.
3. Alkali metals are solids while hydrogen is a gas
 The position of hydrogen in the modern periodic table is still under debate as the properties of hydrogen are unique.
Position of Noble gases in the periodic table
 The elements Helium, Neon, Argon, Krypton, Xenon and Radon of group 18 in the periodic table are called as Noble
gases or Rare gases. They are monoatomic gases and do not react with other substances easily, due to completely filled
subshells. Hence they are called as inert gases. They are found in very small quantities and hence they are called as rare
gases.
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ATOMIC NUMBERS
Atomic number Symbol Elements
1 H Hydrogen
2 He Helium
3 Li Lithium
4 Be Beryllium
5 B Boron
6 C Carbon
7 N Nitrogen
8 O Oxygen
9 F Fluorine
10 Ne Neon
11 Na Sodium
12 Mg Magnesium
13 Al Aluminum
14 Si Silicon
15 P Phosphorus
16 S Sulfur
17 Cl Chlorine
18 Ar Argon
19 K Potassium
20 Ca Calcium
21 Sc Scandium
22 Ti Titanium
23 V Vanadium
24 Cr Chromium
25 Mn Manganese
26 Fe Iron
27 Co Cobalt
28 Ni Nickel
29 Cu Copper
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30 Zn Zinc
31 Ga Gallium
32 Ge Germanium
33 As Arsenic
34 Se Selenium
35 Br Bromine
36 Kr Krypton
37 Rb Rubidium
38 Sr Strontium
39 Y Yttrium
40 Zr Zirconium
41 Nb Niobium
42 Mo Molybdenum
43 Tc Technetium
44 Ru Ruthenium
45 Rh Rhodium
46 Pd Palladium
47 Ag Silver
48 Cd Cadmium
49 In Indium
50 Sn Tin
51 Sb Antimony
52 Te Tellurium
53 I Iodine
54 Xe Xenon
55 Cs Cesium
56 Ba Barium
57 La Lanthanum
58 Ce Cerium
59 Pr Praseodymium
60 Nd Neodymium

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61 Pm Promethium
62 Sm Samarium
63 Eu Europium
64 Gd Gadolinium
65 Tb Terbium
66 Dy Dysprosium
67 Ho Holmium
68 Er Erbium
69 Tm Thulium
70 Yb Ytterbium
71 Lu Lutetium
72 Hf Hafnium
73 Ta Tantalum
74 W Tungsten
75 Re Rhenium
76 Os Osmium
77 Ir Iridium
78 Pt Platinum
79 Au Gold
80 Hg Mercury
81 Tl Thallium
82 Pb Lead
83 Bi Bismuth
84 Po Polonium
85 At Astatine
86 Rn Radon
87 Fr Francium
88 Ra Radium
89 Ac Actinium
90 Th Thorium
91 Pa Protactinium

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92 U Uranium
93 Np Neptunium
94 Pu Plutonium
95 Am Americium
96 Cm Curium
97 Bk Berkelium
98 Cf Californium
99 Es Einsteinium
100 Fm Fermium
101 Md Mendelevium
102 No Nobelium
103 Lr Lawrencium
104 Rf Rutherfordium
105 Db Dubnium
106 Sg Seaborgium
107 Bh Bohrium
108 Hs Hassium
109 Mt Meitnerium
110 Ds Darmstadtium
111 Rg Roentgenium
112 Cn Copernicium
113 Nh Nihonium
114 Fl Flerovium
115 Mc Moscovium
116 Lv Livermorium
117 Ts Tennessine
118 Og Oganesson

COAL AND PETROLEUM

 Coal, petroleum and natural gas are fossil fuels.

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 Fossil fuels were formed from the dead remains of living organisms millions of years ago.
COAL
 It is as hard as stone and is black in colour
 Coal is processed in industry get some useful products such as coke, coal tar and coal gas
 Coke is a tough, porous and black substance. It is an almost pure form of carbon. Coke is used in the
manufacture of steel and in the extraction of many metals.
 Coal Tar is a black, thick liquid with an unpleasant smell. Products obtained from coal tar are used as starting
materials for manufacturing various substances used in everyday life and in industry, like synthetic dyes,
drugs, explosives, perfumes, plastics, paints, photographic materials, Roofing materials.
 Coal gas is obtained during the processing of coal to get coke
 Different types of coals are peat, lignite, bituminous, and anthracite
 Anthracite is one of variety of coal contains the highest percentage of carbon
 Lignite coal is called brown coal, is the lowest grade coal with the least concentration of carbon
Petroleum and Natural Gas
 Petrol and diesel are obtained from a natural resource called petroleum. Due to its great commercial
importance, petroleum is also called black gold.
 Petroleum is a mixture of Hydrocarbon
 Petroleum is a dark oily liquid. It has an unpleasant odour. It is a mixture of various constituents such as
petroleum gas, petrol, diesel, lubricating oil, paraffin wax, etc.
 The process of separating the various constituents/ fractions of petroleum is known as refining
 Natural gas is a very important fossil fuel because it is easy to transport through pipes.
 Natural gas is stored under high pressure as compressed natural gas (CNG). CNG is used for power
generation. It is a cleaner fuel.
 Many useful substances are obtained from petroleum and natural gas. These are termed as ‘Petrochemicals’
Constituents of Petroleum and their Uses below
Constituents of Petroleum Uses
LPG Fuel for home and industry
Petrol Motor fuel, aviation fuel
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Lubricating oil Lubrication
Paraffin wax Ointments, candles, Vaseline
Bitumen Paints, road surfacing
Kerosene Fuel for stoves, lamps and jet aircrafts

ACIDS, BASES AND SALTS

 Acids are sour in taste and change the colour of blue litmus to red, whereas, bases are bitter and change the
colour of the red litmus to blue
 An acid and a base neutralize each other and form a salt and water. A salt may be acidic, basic or neutral in
nature.
 Special types of substances are used to test whether a substance is acidic or basic. These substances are known
as indicators. The indicators change their colour when added to a solution containing an acidic or a basic
substance. Turmeric, litmus, China rose petal are some of the naturally occurring indicators
LITMUS
 The most commonly used natural indicator is litmus. It is extracted from lichens .It has a
mauve (purple) colour in distilled water. When added to an acidic solution, it turns red and when added to a
basic solution, it turns blue. It is available in the form of a solution, or in the form of strips of paper, known as
litmus paper. Generally, it is available as red and blue litmus paper
ACIDS
 Acidic nature of a substance is due to the formation of H+ ions in solution
 When an acid reacts with a metal, hydrogen gas is evolved and a corresponding salt is formed
Acid + Metal → Salt + Hydrogen gas
 Some metals do not react with acid and liberate hydrogen gas. Example: Ag, Cu.
 When an acid reacts with a metal carbonate or metal hydrogen carbonate, it gives the corresponding salt, carbon
dioxide gas and water
Na2 CO3 + 2HCl → 2NaCl + H2 O + CO2
NaHCO3 + HCl → NaCl + H2 O + CO2
 Acidic solutions in water conduct electricity because they produce hydrogen ions
 Acid is a molecule or ion which is capable of donating proton
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 An acid is a substance which can accept the electron
 Some naturally occurring acids
Natural source Acid
Vinegar Acetic acid
Orange Citric acid
Tamarind Tartaric acid
Tomato Oxalic acid
Sour milk (Curd) Lactic acid
Lemon Citric acid
Ant sting Formic acid
Nettle sting Formic acid
Apple Malic acid

Note
 The atmosphere of Venus is made up of thick white and yellowish clouds of sulphuric acid
 The accidental touch of Nettle leaves creates a pain and burning sensation, which is due to inject of Methanoic
acid into the skin of the person
BASE
 Basic nature of a substance is due to the formation of OH- ions in solution
 Bases react with metals to form salt with the liberation of hydrogen gas.
Zn + 2 NaOH → Na2 ZnO2 + H2 ↑
 Bases react with acids to form salt and water. The reaction between a base and an acid is known as
Neutralisation reaction
KOH + HCl → KCl + H2O
 Basic solution in water conduct electricity because they produce hydroxide ions
 Base is a molecule or ion which is capable of accepting proton
 An base is a substance which can produce the electron
HOW STRONG ARE ACIDS AND BASE SOLUTIONS

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 A scale for measuring hydrogen ion concentration in a solution is called pH scale. The ‘p’ in pH stands for
‘potenz’ in German meaning power. pH scale is a set of numbers from 0 to 14 which is used to indicate
whether a solution is acidic, basic or neutral
 The pH of a neutral solution is 7. Values less than 7 on the pH scale represent an acidic solution. As the pH
value increases from 7 to 14, it represents an increase in OH– ion concentration in the solution, that is,
increase in the strength of alkali
1 Acids have pH less than 7
2 Bases have pH greater than 7
3 A neutral solution has pH equal to 7
 Strength of acids and bases depends upon the number of H+ ions and OH– ions produced, respectively. If we
take hydrochloric acid and acetic acid of the same concentration, say one molar, then these produce different
amounts of hydrogen ions. Acids that give rise to more H+ ions are said to be strong acids, and acids that give
less H+ ions are said to be weak acids.
Substances pH values
Human blood 7.35-7.45
Pure water 7
Lemon juice 2.2
Gastric juice 1.2
Milk of magnesia 10
Human urine 6
Beers 4.5
Wines 2.8-3.8
Black coffee 5.2
Milk 6.5 – 6.7
Normal rain 5.6 - 6
Acid rain 4.2-4.4

IMPORTANCE OF PH IN EVERYDAY LIFE DAY LIFE

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 Our body works within the pH range of 7.0 to 7.8. Living organisms can survive only in a narrow range of pH
change
 When pH of rain water is less than 5.6, it is called acid rain. When acid rain flows into the rivers, it lowers the
pH of the river water. The survival of aquatic life in such rivers becomes difficult.
 Tooth decay starts when the pH of the mouth is lower than 5.5. Tooth enamel, made up of calcium
hydroxyapatite (a crystalline form of calcium phosphate) is the hardest substance in the body. It does not
dissolve in water, but is corroded when the pH in the mouth is below 5.5. Using toothpastes, which are
generally basic, for cleaning the teeth can neutralise the excess acid and prevent tooth decay
 It is very interesting to note that our stomach produces hydrochloric acid. It helps in the digestion of food
without harming the stomach. During indigestion the stomach produces too much acid and this causes pain
and irritation. To get rid of this pain, people use bases called antacids. These antacids neutralize the excess
acid. Magnesium hydroxide (Milk of magnesia), a mild base, is often used for this purpose
USES OF ACIDS „
 Sulphuric acid is called King of Chemicals because it is used in the preparation of many other compounds.
It is used in car batteries also.
 Hydrochloric acid is used as a cleansing agent in toilets.
 Carbonic acid is used in aerated drinks. „
 Tartaric acid is a constituent of baking powder
 Citric acid is used in the preparation of effervescent salts and as a food preservative. „
 Nitric acid is used in the manufacture of fertilizers, dyes, paints and drugs. „
 Oxalic acid is used to clean iron and manganese deposits from quartz crystals. It is also used as bleach for
wood and removing black stains. „
USES OF BASES
 Sodium hydroxide is used in the manufacture of soap.
 Magnesium hydroxide is used as a medicine for stomach disorder.
 Ammonium hydroxide is used to remove grease stains from cloths.
 Calcium hydroxide is used in white washing of building.

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SALTS
 Salt is the product of reaction between acids and bases.
 Salts of a strong acid and a strong base are neutral with pH value of 7. On the other hand, salts of a strong acid
and weak base are acidic with pH value less than 7 and those of a strong base and weak acid are basic in
nature, with pH value more than 7.
 Most of the salts are soluble in water. For example, chloride salts of potassium and sodium are soluble in
water. But, silver chloride is insoluble in water
 Salt is hygroscopic in nature.
USES OF SALTS
COMMON SALT (SODIUM CHLORIDE - NaCl)
 It is used in our daily food and used as a preservative.
BLEACHING POWDER (CaOCl2 )
 For bleaching cotton and linen in the textile industry, for bleaching wood pulp in paper factories and for
bleaching washed clothes in laundry.
 Oxidizing agent in many chemical industries.
 To make drinking water free from germs.
BAKING SODA (NaHCO3 )
 The baking soda is commonly used in the kitchen for making tasty crispy pakoras, etc. Sometimes it is added
for faster cooking
 Baking soda is also an ingredient in antacids. Being alkaline, it neutralizes excess acid in the stomach and
provides relief.
 It is also used in soda-acid fire extinguishers
WASHING SODA (Na2CO3.10H2O)
 Sodium carbonate (washing soda) is used in glass, soap and paper industries.
 It is used in the manufacture of sodium compounds such as borax.
 Sodium carbonate can be used as a cleaning agent for domestic purposes.
 It is used for removing permanent hardness of water.
PLASTER OF PARIS ( CaSO4 .1/2 H2O)
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 Plaster of Paris, the substance which doctors use as plaster for supporting Structured bones in the right
position.
 Plaster of Paris is used for making toys, materials for decoration and for making surfaces smooth
GYPSUM (CaSO4 .2H2O)
 Manufacture of wallboard, cement, plaster of Paris, soil conditioning, a hardening retarder in Portland cement

METALS & NON-METALS

 Metals are lustrous whereas non-metals have no lustre. Metals are malleable and ductile. Non-metals do not
have these properties.
 Metals are good conductors of heat and electricity but non-metals are poor Conductors.
 On burning, metals react with oxygen to produce metal oxides which are basic in nature. Non-metals react
with oxygen to produce non- metallic oxides which are acidic in nature.
 Some metals react with water to produce metal hydroxides and hydrogen gas. Generally, non-metals do not
react with water.
 Metals react with acids and produce metal salts and hydrogen gas. Generally, non-metals do not react with
acids.
 Some metals react with bases to produce hydrogen gas.

METALS
 Metals, in their pure state, have a shining surface. This property is Called metallic lustre.
 Metals can be beaten into thin sheets. This property is called malleability. Gold and silver are most
malleability metal.
 Ability of metals to be drawn into thin wires is called ductility. Gold is the most ductile metal. You will be
surprised to know that a wire of about 2 km length can be drawn from one gram of gold.
 Metals are good conductors of heat and have high melting points. The best conductors of heat are silver and
copper. Lead and mercury are comparatively poor conductors of heats
 Metals are good conductors of electricity. The best conductors of electricity is silver
 When an acid reacts with a metal, hydrogen gas is evolved and a corresponding salt is formed

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Acid + Metal → Salt + Hydrogen gas
 All metals except mercury exist as solids at room temperature
 Gallium and cesium have very low melting points. These two metals will melt if you keep them on your
palm
 Alkali metals (lithium, sodium, potassium) are so soft that they can be cut with a knife. They have low
densities and low melting points
 Concentrated Acid: It has relatively large amount of acid dissolved in a solvent.
 Dilute Acid: It has relatively smaller amount of acid dissolved in solvent.
 Almost all metals combine with oxygen to form metal oxides.
Metal + Oxygen → Metal oxide
 All metals do not react with oxygen at the same rate. Different metals show different relativities towards
oxygen. Metals such as potassium and sodium react so vigorously.
 Metals such as potassium and sodium react so vigorously that they catch fire if kept in the open. Hence, to
protect them and to prevent accidental fires, they are kept immersed in kerosene oil.
 Anodising is a process of forming a thick oxide layer of aluminium. Aluminium develops a thin oxide layer
when exposed to air. This aluminium oxide coat makes it resistant to further corrosion
 Silver and gold do not react with oxygen even at high temperatures
When Metals react with Water?
 Metals react with water and produce a metal oxide and hydrogen gas. Metal oxides that are soluble in water
dissolve in it to further form metal hydroxide. But all metals do not react with water.
Metal + Water → Metal oxide + Hydrogen
Metal oxide + Water → Metal hydroxide
 Metals like potassium and sodium react violently with cold water. In case of sodium and potassium, the
reaction is so violent and exothermic that the evolved hydrogen immediately catches fire.
 The reaction of calcium with water is less violent. The heat evolved is not sufficient for the hydrogen to
catch fire. Calcium starts floating because the bubbles of hydrogen gas formed stick to the surface of the
metal.
 Magnesium does not react with cold water. It reacts with hot water to form magnesium hydroxide and
hydrogen. It also starts floating due to the bubbles of hydrogen gas sticking to its surface.
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 Metals like aluminium, iron and zinc do not react either with cold or hot water. But they react with steam to
form the metal oxide and hydrogen.
 Metals such as lead, copper, silver and gold do not react with water at all
When Metals react with Acids?
 Metals react with acids to give a salt and hydrogen gas.
Metal + Acid → Salt + Hydrogen
 Hydrogen gas is not evolved when a metal reacts with nitric acid. It is because HNO3 is a strong oxidising
agent. It oxidises the H2 produced to water and itself gets reduced to any of the nitrogen oxides (N2O, NO,
NO2 ). But magnesium (Mg) and manganese (Mn) react with very dilute HNO3 to evolve H2 gas.
 Aqua regia is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in
the ratio of 3:1. Aqua regia is a highly corrosive, fuming liquid. It is one of the few reagents that is able to
dissolve gold and platinum. It is used for cleaning and refining gold.

The Reactivity Series

 The reactivity series is a list of metals arranged in the order of their decreasing activities

Symbol Metal
Sk Potassium

Na Sodium

Ca Calcium

Mg Magnesium

Al Aluminum

Zn Zinc

Fe Iron

Pb Lead

H Hydrogen

Cu Copper

Hg Mercury

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 The Complete General Science Notes (Chemistry) for Railway Exams
Ag Silver

Au Gold

LIST OF METAL AND THEIR ORES

METALS ORES
Sodium Trona
Borax
Common salt
Aluminum Bauxite
Potassium Nitrate
Carnalite
Magnesium Magnesite
Dolomite
Epsom salt
Silver Ruby silver
Horn silver
Mercury Cinnabar
Tin Cassiterite
Lead Galena
Gold Calaverite
Silvenites
Calcium Dolomite
Gypsum
Fluorspar
Asbestos
Iron Haemethite
Magnetite
Bismuth Bismuthate

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NON-METALS
 Examples of non-metals are carbon, sulphur, iodine, oxygen, hydrogen, etc.
 Non-metals are either solids or gases except bromine which is a liquid at room temperature
 Iodine is a non-metal but it is lustrous
 Carbon is a non-metal that can exist in different forms. Each form is called an allotrope.
 Diamond, an allotrope of carbon, is the hardest natural substance known and has a very high melting and
boiling point.Graphite, another allotrope of carbon, is a conductor of electricity
 Non-metals produce acidic oxides when dissolve in water
CORROSION
 When a metal is attacked by substances around it such as moisture, acids, etc., it is said to corrode and this
process is called corrosion. The black coating on silver and the green coating on copper are other examples
of corrosion
PREVENTION OF CORROSION
 Rusting of iron can be prevented by painting, oiling, greasing, galvanizing , chrome plating, anodizing or
making alloys
 Galvanization is a method of protecting steel and iron from rusting by coating them with a thin layer of zinc.
ALLOYING
 Alloying is a very good method of improving the properties of a metal.
 Alloy is a homogeneous mixture of two or more metals, or a metal and a Non-metal.
 Iron is the most widely used metal. But it is never used in its pure state. This is because pure iron is very soft
and stretches easily when hot. But, if it is mixed with a small amount of carbon, it becomes hard and strong.
When iron is mixed with nickel and chromium, we get stainless steel, which is hard and does not rust.
 Pure gold is very soft. It is, not suitable for making jewelry. It is alloyed with either silver or copper to make it
hard.
IMPORTANT ALLOYS
Alloy Combinations
Solder Lead and Tin
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 The Complete General Science Notes (Chemistry) for Railway Exams
Brass Copper and zinc
Stainless steel Iron, Chromium and Nickel
Bronze Copper and Tin
Invar Iron and Nickel
Constantan Copper and Nickel
Gun metal Copper ,tin and zinc
Sterling silver Silver and copper
German silver Copper , zinc and Nickel

 An amalgam is an alloy of mercury.

 Electrical conductivity and melting point of an alloy is less than that of pure metals.

 Some alloys have lower melting point than pure metals (Example: Solder is an alloy of lead and tin which has
lower melting point than each of the metals).

 Solder is used for welding electrical wires together.

 Alloys do not get corroded or get corroded to very less extent

IMPORTANT CHEMICAL AND ITS USES

Chemical name Common name Uses
Aluminium Used in Heat resistant
clothing, Cookware and
manufacturing of aircraft
Acetic Acid Vinegar Cooking, baking and pickling
Acetylsalicylic Acid Aspirin Medical
Argon Used in incandescent lighting
equipment’s such as Bulbs,
CFLs
Ammonium Phosphate Fertilizer Used as a fertilizer in

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 The Complete General Science Notes (Chemistry) for Railway Exams
Agricultural
Aluminium potassium Alum Used in Water Purification
Sulphate ,Some types of Toothpastes
and Pickling Agent
Ammonium Nitrate Fertilizers and Explosives
Bismuth Fire detection systems and
bullets
Calcium Carbonate LimeStone Marble, Limestone and
Precipitated Chalk
Calcium oxide Quicklime Cement Production
Carbon Graphite, Fossil Fuels, Clay,
Charcoal and Diamond
Copper Manufacturing of Electrical
Wires & cables
Glycerin Making of Skin Products
Ethanol Antiseptic, Rocket Fuels, Fuel
cells and Engine Fuel
Helium Treating Asthma and Barcode
Reading
Lithium Portable Battery and Making
of Optical devices
Mercury Quicksilver Barometers and Thermometer
Sodium Nitrate Gunpowder making and
treating of dentine
hypersensitivity
Sulphuric acid Vitriol Electrolyte and Industrial
Cleaning agent
Zinc Galvanizing

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 The Complete General Science Notes (Chemistry) for Railway Exams

CHEMISTRY ONE LINERS

 Graphite is used a lubricant in heavy machines
 Aspirin is obtained from latex tree
 Ionic compounds conduct electricity when dissolved in water and They are soluble in water and are also
crystalline solids
 Father of modern chemistry is Lavoisier
 Cathode rays consists beam of electrons
 Nucleus of an atom consists of protons and neutrons
 Proton was discovered by Rutherford
 A swimmer finds it easier to swim in sea water than plain water because of sea water has higher density
 An electric iron has heating element made of Nichrome
 Heaviest naturally occurring element of periodic table uranium
 Pungent smell of garlic is due to asulphur compound
 White phosphorous is stored under water because it is dangerously reactive in air
 Mercury is known as quick silver
 Red phosphorous present at the tip of the match stick
 Magnesium burns with dazzling white flame
 Sodium benzoate Is used as food preservative
 Potassium is used for the manufacturing of fertilizers
 Fluorine is the most electronegative element in of the periodic table
 Francium is the most electropositive element in of the periodic table
 Bhopal gas Tragedy of 1984 is related to Methyl Isocyanate
 A powerful eye irritate present in smog is Peroxyacetyl nitrate
 Plastic is type of polymer
 Platinum is known as white gold
 Petroleum is a mixture of Hydrocarbon
 Acetyl salicylic acid commonly used as a pain killer
 Iron is commonly used for making an electromagnet
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 Halite commonly known as rock salt. Halite is the source of common salt
 Xenon is known as a stranger gas
 Rayon is known as a artificial silk
 Reinforced glass is used in bullet proof screens
 When quick lime is added to water heat is liberated
 Nail polish remover contains Acetone
 Zeolite is suitable for water purification
 Silicon used in the manufacture of high voltage insulators
 Chemical name of Green vitriol is Iron sulphate
 Sodium silicate is chemical name of quartz
 Camphor can easily be purified by the process of sublimation
 The National Chemical Laboratory is located in Pune
 Quick silver is another name of mercury
 Natural rubber is heated with sulphur in vulcanization process
 Titanium dioxideis the chemical name for marble
 Deep blue colour is imported to glass by the presenceof Cobalt oxide
 Anthracite is one of variety of coal contains the highest percentage of carbon
 Henri Becquerel discovered the radioactivity
 Cooking oil can be converted into vegetable ghee by the process of Hydrogenation
 Silver iodide is used to produce artificial rain
 Lightest element in the universe is hydrogen
 Germanium and silicon is most commonly used in semiconductors
 Silver nitrate is commonly used in voting ink .It is first used in india 1962 in mysore
 Hydrogen is the lowest density element and Osmium the highest density element
 Silver bromide is commonly used chemical in photography
 Tungsten has highest Melting and boiling point
 Radon is the heaviest gas
 Hydrogen peroxide is used to restore the colour of old oil paintings
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 Ethylene Glycol is used in car radiators as it increases the freezing temperature
 Age of fossils and archeological excavation is determined by radioactive carbon (C-14)
 Non-stick utensil is made up of Teflon
 Gelatinused to prevent the melting of ice
 Ferric chloride is used to stop bleedingbecause it is a strong coagulant
 Barium is the responsible for green colour in fireworks
 Liquid hydrogen is used as a rocket fuel
 Fluorescent tube contains helium gas and neon gas
 Copper is the first metal used by man
 Titanium is called strategic metal
 Lithium is the lightest metal. It weighs about half as much as water
 Antacids drugs are used to productive relief burning sensation in stomach
 Backlites used in electrical insulator, switches, handles of cook wares
 Periodic table
Group 13 Boron family

Group 14 Carbon family

Group 15 Nitrogen family

Group 16 Chalcogen family

Group 17 Halogen family

Group 18 Group 18

 Element common to all acids is Hydrogen

 Balloons are filled withHelium
 Most abundant metal in earth’s crust is Aluminium
 Carbon occur in the nature in purest form is Diamond
 Gelatineis used to avoid melting of ice

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 Tooth enamel is made up of Calcium Phosphate
 Calcium Phosphate acid is used in soft drinks
 In the absence of air and under high temperature and pressure the dead organisms are converted into
petroleum and natural gas
 Nuclear fuel in the sun is Helium
 Metal constituent of chlorophyll is Magnesium
 Carbon dioxide is responsible for the swelling of bread
 Kerosene is a mixture of Aliphatic hydrocarbons
 Most of the explosions in mines occurs due to mixing of Methane with air
 Titanium is known as Metal of Future
 Impurity present in ore is Gangue
 Paper is chemically Cellulose
 Xenon is also known as Stranger Gas
 Butane is used in cigarette lighters
 Metals are lustrous because they have free electrons
 Noble gases are Colourless and Odourless
 Petroleum is found in Sedimentary Rocks
 Lead pencil contains Graphite
 Platinum is called white Gold
 Nickle is used for the synthesis of Vanaspati Ghee
 Ammonia (NH3) is synthesized through Haber’s process
 Ozone is allotrope of oxygen
 Cesium used in photoelectric cells which is used to convert sunlight into electricity
 Calcium hydride is used to prepare fire proof and water proof clothes
 During the process of rusting the weight of iron Increases increased due to the weight of oxygen which has
combined with the iron
 Fuse wire is made up of Lead and Tin
 Gases used by sea divers for breathing are Oxygen and Helium

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 Ozone blackens silver’s shine
 Egg shell is made up of Calcium Carbonate
 Acid rain is caused when the air is polluted by Nitrous Oxide &Sulphur dioxide Gases
 Alum is used as a Water Purifier
 Electro negative elements are non-metal
 Lignite known as the brown goal
 Black lung disease occurs in people working in coal mines
 Lead pollutes big cities air.Sources of lead is emissions from motor vehicles and industrial sources
 Crook Glass is used to make sun glass
 Acetylsalicylic Acid commonly known as Aspirin
 Barium Hydroxide is known Baryta water
 Benzoic acid is one of the most common preservatives used in food processing industry
 Deuterium is Isotope of hydrogen
 Efficiency of the catalyst depends on its molecular state
 Mine explosions are mostly caused by mixing of Air and Methane
 Natural rubber is a polymer derived from Isoprene
 Iron Pyrite is known as Fools Gold
 Ozone is diamagnetic in nature
 Oxides of metals are alkaline
 Paraffin wax is Saturated hydrocarbon
 Mercury Vapour& Argon is filled inside a Tube light
 Vinegar is an aqueous solution of Acetic acid
 Bee Sting contains a Methanoic Acid
 A Photoelectric cell contains selenium metal
 Zinc Phosphide is used a Rat Poison
 Hydrogen was the first element to be produced after Big Bang
 The nature of saliva is acidic.
 Steel contains 0.1–2 percent carbon

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 Commonly used medicine for typhoid is chloromycetin.
 The chemical that is used in making artificial rain is silver nitrate.
 Aqua regia is a mixture of HCI and NHO3
 Hematite is ore of iron
 Acid is used to write on glass- Hydrochloric acid
 Calcium and magnesium ion cause hardness of water.
 Pencil “lead” is made up of – Graphite
 The hardest substance available on earth is Diamond
 Lavioser was the first person to classify elements into metals and non-metal.
 Copper and its alloys are natural antimicrobial material
 Dead organisms are transformed into petroleum and natural gas in absence of air
 Alum is commonly used in water purification
 Ozone is Allotrope of Oxygen
 Deuterium is Isotope of Hydrogen
 Biogas chiefly contains Methane
 Carbon dioxide is responsible for the swelling of bread
 Chemical name of Picric Acid is Tri Nitro Phenol
 Egg shell is made up of Calcium Carbonate (CaCO3)
 Uses of isotopes
Iron 59 Anemia
Iodine 131 Goitre
Cobalt 60 Cancer
Carbon11 Brain scan

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