1. Chemical Equations

Chemical Reactions and Equations: Balanced and unbalanced chemical equations and balancing of chemical equations.Consider the following situations of daily life and think what happens when , In all the above situations, the nature and the identity of the initial substance have somewhat changed. We have already learnt about physical and chemical changes of matter in our previous classes. Whenever a chemical change occurs, we can say that a chemical reaction has taken place

  • Milk is left at room temperature during summers.
  • An iron tawa/pan/nail is left exposed to humid atmosphere.
  • Grapes get fermented.
  • Food is cooked.
  • Food gets digested in our body.
  • We respire.

Chemical Reaction: The transformation of chemical substance into another chemical substance is known as Chemical Reaction. For example: Rusting of iron, the setting of milk into curd, digestion of food, respiration, etc.

ExampleThe burning of magnesium in the air to form magnesium oxide is an example of a chemical reaction.

2Mg(s) + O2(g) → 2MgO(s)

Before burning in air, the magnesium ribbon is cleaned by rubbing with sandpaper . This is done to remove the protective layer of basic magnesium carbonate from the surface of the magnesium ribbon. Reactant: Substances which take part in a chemical reaction are called reactants. Example: Mg and O2.

Product: New substance formed after a chemical reaction is called a product.

Example: MgO.

Chemical Reaction's Characteristics

1. Change in temperature: The chemical reaction between quick lime water to form slaked lime is characterized by a change in temperature (which is a rise in temperature).

2. Change in state of substance: The combustion reaction of candle wax is characterised by a change in state from solid to liquid and gas (because the wax is a solid, water formed by the combustion of wax is a liquid at room temperature.

3. Formation of precipitate: The chemical reaction between sulphuric acid and barium chloride solution is characterised by the formation of a white precipitate of barium sulphate.

BaCl2(aq) + H2SO4(aq) ----→ BaSO4(s) (ppt) + 2HCl(aq)

4. Evolution of gas: The chemical reaction between zinc and dilute sulphuric acid is characterised by the evolution of hydrogen gas.

 Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)

Chemical Equations - The chemical equation of the reaction is the representation of a chemical change in terms of symbols and formulae of the reactants and products.

Example: Potassium Hydrochloric Potassium Manganese Water Chlorine permanganate acid → chloride chloride

(a) Balanced Chemical Equation: A balanced chemical equation has the number of atoms of each element equal on both sides.

Example: Zn + H2SO4 → ZnSO4 + H2

In this equation, numbers of zinc, hydrogen and sulphate are equal on both sides, so it is a Balanced Chemical Equation.

(b) Unbalanced Chemical Equation: If the number of atoms of each element in reactants is not equal to the number of atoms of each element present in the product, then the chemical equation is called Unbalanced Chemical Equation.

Example: Fe + H2O → Fe3O4 + H2

Balancing a Chemical Equation: To balance the given or any chemical equation, follow these steps:

Fe + H2O → Fe3O4 + H2

Table as shown here.

Fe + 4 × H2O → Fe3O4 + H2To balance the oxygen, one needs to multiply the oxygen on the LHS by 4, so that, the number of oxygen atoms becomes equal on both sides.

Now, the number of hydrogen atoms becomes 8 on the LHS, which is more than that on the RHS. To balance it, one needs to multiply the hydrogen on the RHS by 4.

Fe + 4 × H2O → Fe3O4 + 4 × H2

After that, the number of oxygen and hydrogen atoms becomes equal on both sides. The number of iron is one on the LHS, while it is three on the RHS. To balance it, multiply the iron on the LHS by 3.

3 × Fe + 4 × H2O → Fe3O4 + 4 × H2

Now the number of atoms of each element becomes equal on both sides. Thus, this equation becomes a balanced equation.

After balancing, the above equation can be written as follows:

3Fe + 4H2O → Fe3O4 + 4H2.

1. Chemical Equations




CHEMICAL REACTIONS:- The process in which two or more substance combine with each other to form new substances with new properties is called chemical reaction.

There are two parts of a chemical reactions :-

(i) Reactants:-   The substances which take part in a chemical reaction are known as reactants.

(ii) Products:-  The new substances formed during a chemical reaction are known as products.

There are 5 ways to tell if a chemical reaction has occurred.

  • Change in state.
  • Change  in colour.
  • Change in temperature.
  • Evolution of a gas.
  • Formation of precipitate.

  Chemical reaction in everyday life:-

  • Digestion of food.
  • Respiration.
  • Rusting of iron.
  • Formation of curd.
  • Burning of magnesium ribbon.

Chemical Equations:-  A chemical equation is a written representation of a chemical reaction.

The representation of chemical reaction using symbols and formulae of the substances is called chemical equation.

A   +   B             C   +     D

Reactants                   Products   

n this equation, A and B are called reactants and C and D are called the products. The arrow shows the direction of the chemical reaction. The necessary condition such as temperature, pressure or any catalyst should be written on arrow between reactants and products.

E.g. Magnesium is burnt in air to form magnesium oxide.

(i) Word equation for above reaction would be -

 Magnesium + oxygen                Magnesium oxide

   ( Reactants )                                      ( Product )

Skeletal equation for above reaction would be - 

Mg +    O2                   MgO


  • LAW OF CONSERVATION OF MASS :-  Mass can neither be created nor be destroyed in a chemical reaction.
  • So number of elements involved in chemical reaction should remain same at reactant and products side.

For Example ,

Zn   +   H2SO4                  ZnSO4  +         H2

(Zinc)    ( Sulphuric Acid)        (Zinc Sulphate)   ( hydrogen) 

Let us check the number of atoms of different elements on both sides of the arrow .


As the number of atoms of each element is same on both sides of arrow. This is   a balanced chemical equation.

 Let us take another example :-   

Fe     +    H2O     Fe3O4   +   H2

STEP 1 :-   Write a chemical equation.

Fe   +    H2O  →   Fe3O4  +  H2

 STEP 2:-  List the number of atoms of different elements present in the unbalanced equation.


















STEP 3 :_ Select the element which has the maximum number of atoms . now equalize the number of atoms by putting coefficient in front of it.

Fe   +      4  H2O           Fe3O4   +     H2

STEP 4 :-  Fe and H atoms are still not balanced choose any elements now to balance. To equalize the number of H atoms,

Fe   +     4 H2O         Fe3O4   +     4 H2

STEP 5 :- Now, take Fe and equalize the number of Fe atoms.

3 Fe   + 4 H2O     Fe3O 4     +    4 H2

Now all the atoms of elements are equal on both sides.

STEP 6 :-  To make the chemical equation more information ,write the physical states of reactants and products.

Solid state = (s)

Liquid state = (l)

Gaseous State = (g)

Aqueous state = (aq)

3 Fe (S)    +   4H2O(g)          Fe3O4S4H2(g) 

STEP 7:-  Write necessary conditions of temperature pressure or catalyst on above or below arrow.

For Example:-          


2. Types of Chemical Reactions

Types of Chemical Reactions: Combination Reaction, Decomposition Reaction, Displacement Reaction, Double Displacement Reaction, Neutralization Reactions, Exothermic – Endothermic Reactions and Oxidation-Reduction Reactions.

A. Combination or Synthesis Reactions:

The reactions in which two or more chemicals combine to generate a single new compound.

Types of Combination reactions:

I. Combination of two elements to form a compound

Burning of hydrogen in air or oxygen to produce water.

H²(hydrogen) + O²( oxygen)    H²O(water) + O²( oxygen)

   II. Combination Reactions involving an Element and a Compound

Burning of carbon monoxide in oxygen to form carbon dioxide.

2CO (carbon monoxide) + O2(Oxygen) 2CO2(Carbon dioxide)

III. Combination Reactions involving Two Compounds

Combination of ammonia and hydrogen chloride to produce ammonium chloride.

NH³( Ammonia) + HCl(Hydrogen chloride) → NH⁴Cl³( Ammonium chloride)

B . Decomposition Reaction: Reactions in which one compound decomposes in two or more compounds or elements are known as Decomposition Reaction. A decomposition reaction is just the opposite of combination reaction.
A general decomposition reaction can be represented as follows :
A + B
When calcium carbonate is heated, it decomposes into calcium oxide and carbon dioxide.

CaCO3(s) heat−→− CaO(s) + CO2(g)
Calcium carbonate Calcium oxide + Carbon dioxide

When ferric hydroxide is heated, it decomposes into ferric oxide and water
2Fe(OH)3(s)  Fe2O3(s) + 3H2O(l)

Thermal Decomposition: The decomposition of a substance on heating is known as Thermal Decomposition.
2Pb(NO3)2(s) heat−→− 2PbO(s) + 4NO2(g) + O2(g)

Electrolytic Decomposition: Reactions in which compounds decompose into simpler compounds because of passing of electricity, are known as Electrolytic Decomposition. This is also known as Electrolysis.
Example: When electricity is passed in water, it decomposes into hydrogen and oxygen.
2H2O(l) Undefined control sequence \xrightarrow 2H2(g) + O2(g)

Photolysis or Photo Decomposition Reaction: Reactions in which a compound decomposes because of sunlight are known as Photolysis or Photo Decomposition Reaction.
Example: When silver chloride is put in sunlight, it decomposes into silver metal and chlorine gas.

2AgCl(s) (white) Sunlight−→−−−−− 2Ag(s) (grey) + Cl2(g)

Photographic paper has a coat of silver chloride, which turns into grey when exposed to sunlight. It happens because silver chloride is colourless while silver is a grey metal.

C. Displacement Reaction: The chemical reactions in which a more reactive element displaces a less reactive element from a compound is known as Displacement Reactions. Displacement reactions are also known as Substitution Reaction or Single Displacement/ replacement reactions.
A general displacement reaction can be represented by using a chemical equation as follows :
A + BC
AC + B
Displacement reaction takes place only when ‘A’ is more reactive than B. If ‘B’ is more reactive than ‘A’, then ‘A’ will not displace ‘C’ from ‘BC’ and reaction will not be taking place.
When zinc reacts with hydrochloric acid, it gives hydrogen gas and zinc chloride.

Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

When zinc reacts with copper sulphate, it forms zinc sulphate and copper metal.
Zn(s) + CuSO4(aq) ZnSO4(aq) + Cu(s)

D. Double Displacement Reaction: Reactions in which ions are exchanged between two reactants forming new compounds are called Double Displacement Reactions.
When the solution of barium chloride reacts with the solution of sodium sulphate, white precipitate of barium sulphate is formed along with sodium chloride.

BaCl2(aq) + Na2SO4(aq) BaSO4(s) (Precipitate) + 2NaCl(aq)

When sodium hydroxide (a base) reacts with hydrochloric acid, sodium chloride and water are formed.
NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)

Note: Double Displacement Reaction, in which precipitate is formed, is also known as precipitation reaction. Neutralisation reactions are also examples of double displacement reaction.

Precipitation Reaction: The reaction in which precipitate is formed by the mixing of the aqueous solution of two salts is called Precipitation Reaction.

Neutralization Reaction: The reaction in which an acid reacts with a base to form salt and water by an exchange of ions is called Neutralization Reaction.


E. Oxidation and Reduction Reactions:
Oxidation: Addition of oxygen or non-metallic element or removal of hydrogen or metallic element from a compound is known as Oxidation.
Elements or compounds in which oxygen or non-metallic element is added or hydrogen or metallic element is removed are called to be Oxidized.

Reduction: Addition of hydrogen or metallic element or removal of oxygen or non-metallic element from a compound is called Reduction.
The compound or element which goes under reduction in called to be Reduced.
Oxidation and Reduction take place together. For example

Oxidizing agent:

  • The substance which gives oxygen for oxidation is called an Oxidizing agent.
  • The substance which removes hydrogen is also called an Oxidizing agent.

Reducing agent:

  • The substance which gives hydrogen for reduction is called a Reducing agent.
  • The substance which removes oxygen is also called a Reducing agent.

The reaction in which oxidation and reduction both take place simultaneously is called Redox reaction.
When copper oxide is heated with hydrogen, then copper metal and hydrogen are formed.

CuO + H2  Cu + H2O
(i) In this reaction, CuO is changing into Cu. Oxygen is being removed from copper oxide. Removal of oxygen from a substance is called Reduction, so copper oxide is being reduced to copper.

(ii) In this reaction, H2 is changing to H2O. Oxygen is being added to hydrogen. Addition of oxygen to a substance is called Oxidation, so hydrogen is being oxidised to water.

  • The substance which gets oxidised is the reducing agent.
  • The substance which gets reduced is the oxidizing agent

F. Exothermic and Endothermic Reactions:
Exothermic Reaction: Reaction which produces energy is called Exothermic Reaction. Most of the decomposition reactions are exothermic.
Respiration is a decomposition reaction in which energy is released.

When quick lime (CaO) is added to water, it releases energy.

Endothermic Reaction: A chemical reaction in which heat energy is absorbed is called Endothermic Reaction.
Example: Decomposition of calcium carbonate.

2. Types of Chemical Reactions


There are five type of chemical reactions.

COMBINATION REACTION:-   The reaction in which two or more reaction combine to form a single product is called combination reaction They are represented by equation of the following term. 

 A  + B                   AB      

Reactant           product

For Example :-  

Burning of coal  

 C(s)      +      o2(g)    →      Co2 (g)   

(carbon)     (oxygen)     (carbon dioxide)

Formation of water

H2 (g)    +        O2 (g)    →    2H2 O (l)           

 (Hydrogen)      (oxygen)          (Water)   

Formation of slaked line

 CaO (s)     +      H2 0(L)           Ca(OH)2 (aq)

(calcium oxide)   (water)   (calcium Hydroxide)/ slaked lime

Formation of slaked lime by the reaction of calcium oxide with water.

DECOMPOSITION REACTION :-   The reaction I which single compound breaks down to simpler product is called decomposition reaction. They are represented by equation of the following term .

  AB              A+B

  (i) Thermat decomposition :-   when decomposition is carried out by heating.                           

For example     

(i) 2 Fe SO4 (s)            CaO(s) +CO2 (g)

(Ferrous sulphate )        (Ferric Oxide)                               

 Green colour                 Red brown colour

(ii)  CaCO3 (s)                        CaO(s)            +        CO2 (g)

(calcium carbonate)       ( calcium oxide )   (Carbon dioxide )  Limestone quick lime

(iii) 2Pb (N03)2   (s) + 2PbO(s)     4NO2 (g)  +   O2(g)    

(Lead nitrate )     (lead oxide)  (Nitrogen dioxide) oxygen   

                                                         Brown fumes

 Heating of lead nitrate and emission of nitrogen  dioxide

 ELECTRONIC DECOMPOSITION :- When decomposition is carried out by passing electricity .

 For Example:-  (i) Electrolysis of water.

 Electric current                                                                                        

 2H2 0              2H2       +    O2

(Water)            (Hydrogen)    (Oxygen)

Electrolysis of water is done as follows:-

Electrolysis of water

When decomposition is carried out in presence of sunlight .

For Example:-

This is why silver chloride turns grey in sunlight because of the decomposition of silver chloride into silvers and chloride by light.

This reaction is used in black and write Photography.

DISPLACEMENT REACTION :-  This reaction in which more reactive element displace less reative element from its salt solution is called displacement reaction . they aare represented by equation of the following term.

A    +    BC          AC   +  B    

For example :- 

(i) Fe (s)  +  CuSO4 (aq)           FeSO4  (aq)   +   Cu(s)

(Iron )    (copper Sulphate)    ( iron Sulphate)      ( Copper)

Fe is more reactive than therefore iron (Fe) has displaced copper (Cu) from copper sulphate solution.

Iron nails dipped in copper sulphate solution.

(ii)  Zn(s) + CuS04 (aq)            ZnSO4 (aq)    +  Cu (s)

(Zinc)  (copper sulphate)     ( Zinc  sulphate )   copper

Zinc is more reactive than copper , therefore it displace copper from copper sulphate solution.

(iii) Pb (s)   + CuCl2 (aq)            PbCl2 (aq)    + Cu(s)

(Lead)    (copper chloride)        ( lead chloride)    copper

Lead is more reactive elements than copper, therefore it displaces copper from copper chloride solution.

DOUBLE DISPLACEMENT REACTIVE :- The reaction in which the reactant ions exchange to form new products is called double displacement reaction . they are represented by equation of the following term.

AB    +    CD          AD    +        CB

For Example :-

Na2 SO4 (aq)   +  BaCl2 (aq)        BaSO4 (s)    +  2NaCl (aq)

(sodium sulphate) (barium chloride) (barium sulphate) (sodium chloride)

White precipitate of BaSO4  is formed .the insoluble substance formed is known as precipitate . any reation that produces  a precipitate can be called a precipitation reaction.


OXIDATION    →    

(i)  The addition of oxygen to reactant

(ii)   the removal of hydrogen from  reactant

 For Example  :-

  2 Cu +         O2              2CuO

 (Copper)  (oxygen)       (copper oxide )

                                          Black substance

Oxidation  of copper to copper oxide

Reduction  : -      (i) the  addition of hydrogen to reactant

(ii) the removal of oxygen from a reactant.

 Redox Reactions :-  The reaction on written one substance gets oxidisied and other get reduced is known as redox reaction

  For example :-    

 CuO  +    H2         Cu   +   H2 O

In this reaction CuO is reduced to Cu and H2  is oxidized to H2 O . So oxidation and reduction taking place together, therefore it is a redox reaction.

 ZnO +  C           Zn    +   C O

    Here, C is oxidized to CO because oxygen is being added and ZnO IS REDUCED TO Zn because O is being removed.

   NOTE :→  

*  If a substance gains oxygen during a reaction , it is said to be oxidized.

*   If a substance loses oxygen during a reactant , it is said to be reduced.

ENDETHERMIC REACTION:-    Reaction in which energy is absorded are known as endothermic reaction.

For example :-  

 CaCO3   (S)              CaO (s)   +   CO2  (g)

FXOTHERMIC REACTION :-    Reaction in which heat is released along with formation of products.

For Example, 

CH(q)   +  2 O2  (g)         CO2 (g) +2H2 O(g) + Heat

3. Effects of Oxidation Reactions in Everyday Life

Effects of Oxidation Reactions in Everyday life: Corrosion and Rancidity.

Corrosion: The process of slow conversion of metals into their undesirable compounds due to their reaction with oxygen, water, acids, gases etc. present in the atmosphere is called Corrosion.

Example: Rusting of iron.

Corrosion of Copper: Copper objects lose their lustre and shine after some time because the surface of these objects acquires a green coating of basic copper carbonate, CuCO3.Cu(OH)2 when exposed to air.

Corrosion of Silver Metal: The surface of silver metal gets tarnished (becomes dull) on exposure to air, due to the formation of a coating of black silver sulphide(Ag2S) on its surface by the action of H2S gas present in the air.

Rancidity: The taste and odour of food materials containing fat and oil changes when they are left exposed to air for a long time. This is called Rancidity. It is caused due to the oxidation of fat and oil present in food materials.

Methods to prevent rancidity:

By adding anti-oxidant.

Vacuum packing.

Replacing air by nitrogen.

Refrigeration of foodstuff.

A. Chemical Reaction: During the chemical reactions, the chemical composition of substances changes or make new substances are formed known as chemical reaction.

B. Chemical Equation: Chemical reactions can be written in chemical equation form which should always be balanced are know as chemical equation.

C. Types of Chemical Reactions:

Combination reaction: A single product is formed from two or more than two reactants.

2Mg + O2 → 2MgO

Decomposition reaction: A single reactant breaks down two or more then two products.

Thermal decomposition: 

2Pb(NO2)2 → 2PbO + 4NO2 + O2


2H20 → 2H2 + O2

Photochemical reaction: 

2AgBr → 2Ag + Br2

Displacement reaction: One element is displaced by another element.this format call di

Zn + CuSO4 → ZnSO4 + Cu

Double displacement reaction: Exchange of ions between reactants.

AgNO3 + NaCl → AgCl + NaNO3

Redox reaction: Both oxidation and reduction take place simultaneously.

CuO + H2 → Cu + H2O

Exothermic reaction: A chemical reaction in which heat energy is evolved.

C + O2 → CO2 (g) + heat

Endothermic reaction: A chemical reaction in which heat energy is absorbed.

ZnCO3 + Heat → ZnO + CO2

5. Reduction: Reaction that shows the loss of oxygen or gain of hydrogen.

ZnO + C → Zn + CO

ZnO is reduced to Zn—reduction. C is oxidized to CO—Oxidation.

Effects of Oxidation Reactions in Our Daily Life:

Corrosion: It is an undesirable change that occurs in metals when they are attacked by moisture, air, acids and bases.

Example, Corrosion (rusting) of Iron: Fe2O3. nH2O (Hydrated iron oxide)

Rancidity: Undesirable change that takes place in oil containing food items due to the oxidation of fatty acids.

Preventive methods of rancidity: Adding antioxidants to the food materials, storing food in the airtight container, flushing out air with nitrogen gas and refrigeration

3. Effects of Oxidation Reactions in Everyday Life


1 COROSION :-    When a metal is exposd to moisture ,air, acid eyc. For some time , a layer of hydrated oxide, is formed which e weaknes the metal and hence metal is said to be corroded.

Example of corrosion are:-

  1. Rusting of iron .
  2. Black coating on silver
  3. Green coating on copper.

• Rusting of iron :-    when iron is exposed to oxygen in the presence of  moisture, reddish brown power is formed.

This is process is knon as rusting of iron.

• Method to prevent corrosion are:-  

  1. Galvanization
  2.  Electroplating
  3. By putting paints

2) RANCIDITY:-    The oxidation of fats and oils when exposed to air is known as rancidity . due to rancidity, bad smell and bad taste of food occurs

Method of prevent rancidity are :-

  1. By adding antioxidants.
  2. Refrigeration. 
  3. Replacing air by nitrogen.
  4. Keeping food in air tight containers.

 Chips manufactures fill bag of chips with nitrogen because it is non reactive gas and it prevent the chips from getting oxidized.

1. The Chemical Properties of Acids and Bases

Indicators: Indicators are substances which indicate the acidic or basic nature of the solution by the colour change.

Types of Indicator:  Some common types of indicators are:

1. Natural Indicators: Indicators obtained from natural sources are called Natural Indicators. Litmus, turmeric, red cabbage, China rose, etc., are some common natural indicators used widely to show the acidic or basic character of substances.

Turmeric: Turmeric is another natural indicator. Turmeric is yellow in colour. Turmeric solution or paper turns reddish brown with base. Turmeric does not change colour with acid.

Red Cabbage: The juice of red cabbage is originally purple in colour.

2. Olfactory Indicator: Substances which change their smell when mixed with acid or base are known as Olfactory Indicators. For example; Onion, vanilla etc.

Onion: Paste or juice of onion loses its smell when added with base. It does not change its smell with acid.

Vanilla: The smell of vanilla vanishes with base, but its smell does not vanish with an acid.

3. Synthetic Indicator: Indicators that are synthesized in the laboratory are known as Synthetic Indicators. For example; Phenolphthalein, methyl orange, etc.

Phenolphthalein is a colour less liquid. It remains colour less with acid but turns into pink with a base.

Methyl orange is originally orange in colour. It turns into the red with acid and turns into yellow with base.

Acids, Bases and Salts

Classification of matter

On the basis of

a) composition –  elements, compounds and mixtures

b) state – solids, liquids and gases

c) solubility – suspensions, colloids and solutions

Types of mixtures – homogeneous and heterogeneous

Types of compounds – covalent and ionic

an Acid and a Base?

Ionisable and non-ionisable compounds

An ionisable compound when dissolved in water or in its molten state, dissociates into ions almost entirely. Example: NaCl, HCl, KOH, etc.

A non-ionisable compound does not dissociate into ions when dissolved in water or in its molten state. Example: glucose, acetone, etc.

Arrhenius theory of acids and bases

Acids: Acids are sour in taste, turn blue litmus red, and dissolve in water to release H+ ions.

Properties of Acids:

Acids have a sour taste.

Turns blue litmus red.

Acid solution conducts electricity.

Release H+ ions in aqueous solution.

Types of Acids: Acids are divided into two types on the basis of their occurrence i.e., Natural acids and Mineral acids.

(i) Natural Acids: Acids which are obtained from natural sources are called Natural Acids or Organic Acids.


Methanoic acid (HCOOH)

Acetic acid (CH3COOH)

Oxalic acid (C2H2O4) etc.

(ii) Mineral Acids: Acids that are prepared from minerals are known as Mineral Acids Example; Inorganic acids, man-made acids or synthetic acid are also known as Mineral Acids.


Nitric acid (HNO3)

Carbonic acid (H2CO3)

Phosphoric acid (H3PO4) etc.

Chemical Properties of Acid:

(i) Reaction of acids with metal: Acids give hydrogen gas along with respective salt when they react with a metal.

Metal + Acid → Salt + Hydrogen

Test For Hydrogen Gas: The gas evolved after reaction of acid with metal can be tested by bringing a lighted candle near it. If the gas bums with a pop sound, then it confirms the evolution of hydrogen gas.

(ii) Reaction of acids with metal carbonate: Acids give carbon dioxide gas and respective salts along with water when they react with metal carbonates.

Metal carbonate + Acid → Salt + Carbon dioxide + Water

 (iii) Reaction of acid with hydrogen carbonates (bicarbonates): Acids give carbon dioxide gas, respective salt and water when they react with metal hydrogen carbonate.


Test For Evolution of Carbon Dioxide Gas: Carbon dioxide turns lime water milky when passed through it. This is the characteristic test for carbon dioxide gas.

The gas evolved because of reaction of the acid with metal carbonate or metal hydrogen carbonate turns lime water milky. This shows that the gas is carbon dioxide gas.


  • Strong Acids

An acid which is completely ionised in water and produces (H+) is called Strong Acid.

Examples: Hydrochloric acid (HCl), Sulphuric acid (H2SO4), Nitric acid (HNO3)

  • Weak Acids

An acid which is partially ionised in water and thus produces a small amount of hydrogen ions (H+) is called a Weak Acid.

Example: Acetic acid (CH3COOH), Carbonic acid (H2CO3) , When a concentrated solution of acid is diluted by mixing water, then the concentration of Hydrogen ions (H+) or hydronium ion (H3O–) per unit volume decreases.

Bases: Bases are bitter in taste, have soapy touch, turn red litmus blue and give hydroxide ions (OH–) in aqueous solution.

Examples: Sodium hydroxide (caustic soda) – NaOH

Calcium hydroxide – Ca(OH)2

Potassium hydroxide (caustic potash) – (KOH)

Properties of Bases:

1.Have a bitter taste.

2.Turns red litmus blue.

3.Conducts electricity in solution.

4.Release OH– ions in Aqueous Solution

Types of bases: Bases can be divided in two types – Water soluble and Water-insoluble.

The hydroxide of alkali and alkaline earth metals are soluble in water. These are also known as alkali.

For example; sodium hydroxide, magnesium hydroxide, calcium hydroxide, etc.

Chemical properties of bases:

(i) Reaction of Base with Metals: When alkali (base) reacts with metal, it produces salt and hydrogen gas.

Alkali + Metal → Salt + Hydrogen

(ii) Reaction of Base with Oxides of Non-metals: Non-metal oxides are acidic in nature. For example; carbon dioxide is a non-metal oxide. When carbon dioxide is dissolved in water it produces carbonic acid.

Therefore, when a base reacts with non-metal oxide, both neutralize each other resulting respective salt and water.

Base + Non-metal oxide → Salt + Water

(iii) Neutralisation Reaction: An acid neutralizes a base when they react with each other and respective salt and water are formed.

Acid + Base → Salt + Water

(iv) Reaction of Acid with Metal Oxides: Metal oxides are basic in nature. Thus, when an acid reacts with a metal oxide both neutralize each other. In this reaction, the respective salt and water are formed.

Acid + Metal Oxide → Salt + Water

Common in all bases: A base dissociates hydroxide ion in water, which is responsible for the basic behaviour of a compound.

Example: When sodium hydroxide is dissolved in water, it dissociates hydroxide ion and sodium ion.

Neutralisation Reaction: When an acid reacts with a base, the hydrogen ion of acid combines with the hydroxide ion of base and forms water

Dilution of Acid and Base: The hydrogen ion in an acid and hydroxide ion in a base, per unit volume, shows the concentration of acid or base.By mixing of acid to water, the concentration of hydrogen ion per unit volume decreases. Similarly, by addition of base to water, the concentration of hydroxide ion per unit volume decreases. This process of addition of acid or base to water is called Dilution and the acid or base is called Diluted.The dilution of acid or base is exothermic. Thus, acid or base is always added to water and water is never added to acid or base. If water is added to a concentrated acid or base, a lot of heat is generated, which may cause splashing out of acid or base and may cause severe damage as concentrated acid and base are highly corrosive



1.Hydrochloric acid (HCl)

2.Sulphuric acid  (H2SO4)

3.Nitric acid (HNO3)


1.Sodium hydroxide (NaOH)

2.Potassium hydroxide (KOH)

3.Calcium hydroxide (Ca(OH)2)

Bronsted Lowry theory

1.A Bronsted acid is an H+ (aq) ion donor.

2.A Bronsted base is an H+ (aq) ion acceptor.


In the reaction: HCl (aq) + NH3 (aq) → NH+4(aq) + Cl− (aq)

HCl – Bronsted acid and Cl− : its conjugate acid

NH3 – Bronsted base and NH+4 : its conjugate acid

Physical test

Given are two possible physical tests to identify an acid or a base.

a. Taste

An acid tastes sour whereas a base tastes bitter.

The method of taste is not advised as an acid or a base could be contaminated or corrosive.

b. Effect on indicators by acids and bases

An indicator is a chemical substance which shows a change in its physical properties, mainly colour or odour when brought in contact with an acid or a base.

Below mentioned are commonly used indicators and the different colours they exhibit:

a) Litmus

In a neutral solution – purple

In acidic solution – red

In basic solution – blue

b) Methyl orange

In a neutral solution – orange

In acidic solution – red

In basic solution – yellow

c) Phenolphthalein

In a neutral solution – colorless

In acidic solution – remains colorless

In basic solution – pink

Acid-Base Reactions & Reactions of acids and bases

a) Reaction of acids and bases with metals

Acid + active metal →  salt + hydrogen + heat

2HCl + Mg → MgCl2 + H2 (↑)

2NaOH + Zn → Na2ZnO2 + H2 (↑)

b) Reaction of acids with metal carbonates and bicarbonates

Acid + metal carbonate or bicarbonate →  salt + water + carbon dioxide.

2HCl + CaCO3 → CaCl2 + H2O + CO2

H2SO4 + Mg (HCO3)2 → MgSO4 + 2H2O + 2CO2

c) Neutralisation reaction

1. Reaction of metal oxides and hydroxides with acids

Metal oxides or metal hydroxides are basic in nature.

Acid + base → salt + water + heat

 H2SO4 + MgO → MgSO4 + H2O

2HCl + Mg (OH) 2 → MgCl2 + 2H2O

2. Reaction of non-metal oxides with bases

Non-metal oxides are acidic in nature

Base + Nonmetal oxide →  salt + water + heat

2NaOH + CO2→ Na2CO3 + H2O


Bases undergo neutralisation reaction with acids.

They are comprised of metal oxides, metal hydroxides, metal carbonates and metal bicarbonates.

Most of them are insoluble in water.

1. The Chemical Properties of Acids and Bases



These are the substances which have sour taste

  • They turn blue litmus solution  red
  • They give H+ ions in aqueous solution

Type  of acids are as follows

  1. Strong Acid :  Hcl , H2SO4
  2. Weak Acid : CH3COOH, HCOOH
  3. Concentrated Acid : which have more amount of acid and less amount of water.
  4. Dilute Acid : which have more amount of water and less amount of acid


* These are the substances which have bitter taste and soapy in touch .

  • They turn red litmus solution blue
  • They give OH- ions in aqueous solution.

Type of bases are as follows :-

  1. Strong Bases :-  NaOH, KOH
  3. ALKALI :- These are the bases which are soluble in water (like NaOH,KOH)

SALTS :-  A salts is a substance produced from the reaction of an acid and a base.


INDICATORS :- The substances that change their colour/ smell when they are added to acidic or alkaline solutions.

Types of Indicators.

  1. Natural Indicators :- Litmus , turmeric
  2. Synthetic Indicators :-  Phenolphthalein, Mythyl
  3. Olfactory Indicators:- Onion, vanilla essence.


(i)  Reaction of Metal with :-     

(a)  ACIDS 

Acid   +  Metal                 Salt +    Hydrogen  gas                      

For Example,

2Hcl           +          Zn                    ZnCl2      +    H2

(Hydrochloric acid) (zinc)       (zinc chloride ) (hydrogen)

(b)  Bases  Base + metal        salt + hydrogen gas

For example,

2NsoH         +          Zn           Na3ZnO2        +        H2  

(sodium hydroxide )  ( zinc)  ( sodium zincate)  ( hydrogen gas )

 Hydrogen gas released can be tested by bringing burning candle near gas bubbles , it burst with pop sound.

(ii)    REACTION OF METAL CARBONATES / Metal Hydrogen carbonates with 

(a)  Acids :- 

(b) Acid   +   metal carbonate/ metal hydrogen carbonate     salt + carbon dioxide + water.

  For example :-

  1. 2 Hcl + Na2CO3                  2NaCl +  CO2    +   H2O
  2. HCl +  NaHCO3                  NaCl +  CO2  +H2O
  3. Ca(OH)2   +   CO2     →         CaCO3 H2O

                ( Line Water)                   ( White precipitate )

On passing CO Through lime water , lime water turns milky and in this way CO2 can be tested.

And when excess CO2 is passed , milkiness disappers and the following reaction take place.

Passing carbon dioxide gas through calcium hydroxide.

CaCo3    +    CO2     +     H2O        Ca (HCO3)

(iii)     Reaction of acids and bases with each other  

Acid     +       base               Salt   +    H2O

Neutralisation Reaction :-   The reaction between an aid and a base to give a salt and water is known as neutralization reaction.

  For Example :-    

HCl (aq)   + NaOH (aq)    →  NaCl(aq) + H2O(l)

(iV)  Reaction of Metallic Oxides with Acids  :-

Metallic oxide +  acid          salt   +  water

 Metallic oxides are basic in nature

For Example :- 

CaO  + 2HCl       CaCl2   +    H2O

(v) Reaction of Non- metallic Oxide with base :-

Non- metallic oxide + base    →        salt   +    H2

Non- metallic oxide are acidic in nature.

For Example :-

CO2   +   Ca(OH)2       CaCO3  +    H2O

1. Physical Properties of Metals and Non-Metals

Metals: Physical properties of metals, chemical properties of metals and non-metal oxide.

Metals are the elements that conduct heat and electricity and are malleable and ductile. Examples are Iron (Fe), Aluminium (Al), Silver (Ag), Copper (Cu), Gold (Au), Platinum (Pt), Lead (Pb), Potassium (K), Sodium (Na), Calcium (Ca) and Magnesium (Mg) etc.

Metals are the elements which form positive ions by losing electrons. Thus, metals are known as Electropositive Elements.

How Do Metals and Nonmetals React

Metals lose valence electron(s) and form cations.
Non-metals gain those electrons in their valence shell and form anions.
The cation and the anion are attracted to each other by strong electrostatic force, thus forming an ionic bond.
For example: In calcium chloride, the ionic bond is formed by opposite charged calcium and chloride ions.
Calcium atom loses 2 electrons and attains the electronic configuration of the nearest noble gas (Ar). By doing so, it gains a net charge of +2.

The two Chlorine atoms take one electron each, thus gaining a charge of -1 (each) and attain the electronic configuration of the nearest noble gas (Ar).

Physical Properties of Metals

  • Hardness: Most of the metals are hard, except alkali metals, such as sodium, potassium, lithium, etc. are very soft metals. These can be cut by using a knife.
  • Conduction: Metals are a good conductor of heat and electricity. This is the cause that electric wires are made of metals like copper and aluminium.
  • Melting and Boiling Point: :Metals have generally high melting and boiling points. (Except sodium and potassium metals which have low melting and boiling point.)
  • Strength: Most of the metals are strong and have high tensile strength. Because of this, big structures are made using metals, such as copper (Cu) and iron (Fe). (Except Sodium (Na) and potassium (K) which are soft metals).
  • State: Metals are solid at room temperature except for mercury (Hg).
  • Malleability: Metals are malleable. This means metals can be beaten into a thin sheet. Because of this property, iron is used in making big ships.
  • Ductility: Metals are ductile. This means metals can be drawn into thin wire. Because of this property, a wire is made of metals.
  • Density: Most of the metals have a high density.
  • Colour: Most of the metals are grey in colour. But gold and copper are exceptions.

 Non-Metals: Physical Properties of non-metals, chemical properties of non-metals, non¬metal oxides, Reaction of metal and Non-metal, Ionic bonds and formation of an ionic bond. Non-metals are the elements that do not conduct electricity and are neither malleable nor ductile.
Examples: Carbon (C), Sulphur (S), Phosphorous (P), Silicon (Si), Hydrogen (H), Oxygen (O), Nitrogen (N), Chlorine (Cl), Bromine (Br), Neon (Ne) and Argon (Ar) etc.
Non-metals are the elements which form negative ions by gaining an electron. Thus, non¬metals are also known as Electronegative Elements.

Physical properties of non-metals

  • Hardness: Non-metals are not hard rather they are generally soft. But the diamond is an exception; it is the hardest naturally occurring substance.
  • State: Non-metals may be solid, liquid or gas.
  • Lustre: Non-metals have a dull appearance. Diamond and iodine are exceptions.
  • Sonority: Non-metals are not sonorous, i.e., they do not produce a typical sound on being hit.
  • Conduction: Non-metals are a bad conductor of heat and electricity. Graphite which is allotrope of carbon is a good conductor of electricity and is an exception.
  • Malleability and ductility: Non-metals are brittle.
  • Melting and boiling point: Non-metals have generally low melting and boiling points.
  • Density: Most of the non-metals have low density.
  • Colour: Non-metals are in many colours.

Carbon in the form of graphite is non-metal which conduct electricity.

Iodine is non-metal which is lustrous having a shining surface.

Carbon in the form of diamond is a non-metal which is extremely hard.

Diamond is a non-metal which has a very high melting point and boiling point.

1. Physical Properties of Metals and Non-Metals



 Physical properties of metals and non- metals.

 1     LUSTRE  -

  Metals have shining surface.

    Non-metals do not have shinig surface

 * Except iodine   

 2 Hardness -

    Metals are generally hard.

  • Except sodium lithium and potassium which are soft and can be cut with knife.

 Non – metals are generally soft.

  • Except diamond a form of carbon which is the hardest natural substan

3 Malleability - 

 Metals can be beaten into thin sheets gold and silver are the most malleable metals.

Non- metals are non-mallerable.

4 Ductility -

 Metals can be drawn into thin wires.

 Non- metals can not be drawn into thin wires. They are non- ductile.

5 Conductor of heat and electricity -

Metals are good conductor of heat and electricity of heat. Lead and mercury are poor conductor of heat.

Metals are good conductor of heat and electricity.

Non- metals are poor conductor of heat EXCEPT graphite.

6 state -

 The metals exist as solids. EXCEPT mercury

The non- metals exist as solids or gaseous.EXEPT  bromine.

7 density -

Metals have high density and high melting point.

Except sodium and potassium.

8 Oxides -

 Metallic oxides are basic in nature.

Non metallic oxides are acidic in nature.

9 Sonorous -

Metals produce a sound on strinking a hard surface.

Non-metals are not sonorous.

1. Bonding in Carbon

Bonding in Carbon: The Covalent bond, Electron dot structure, Physical properties of organic compounds, Allotropes of Carbon.

Hard Water

Hard water contains salts of calcium and magnesium, principally as bicarbonates, chlorides, and sulphates. When soap is added to hard water, calcium and magnesium ions of hard water react with soap forming insoluble curdy white precipitates of calcium and magnesium salts of fatty acids.

2C17H35COONa+MgCl2 → (C17H35COO)2Mg+2NaCl

2C17H35COONa+CaCl2 → (C17H35COO)2Ca+2NaCl

These precipitates stick to the fabric being washed and hence, interfere with the cleaning ability of the soap. Therefore, a lot of soap is wasted if the water is hard.

Covalent Bonds

Difficulty of Carbon to Form a Stable Ion

To achieve the electronic configuration of the nearest noble gas, He, if the carbon atom loses four of its valence electrons, a huge amount of energy is involved. C4+ ion hence formed will be highly unstable due to the presence of six protons and two electrons.

If the carbon atom gains four electrons to achieve the nearest electronic configuration of the noble gas, Ne, C4− ion will be formed. But again, a huge amount of energy is required. Moreover, in C4+ ion it is difficult for 6 protons to hold 10 electrons. Hence, to satisfy its tetravalency, carbon shares all four of its valence electrons and forms covalent bonds.

Ionic Bond

Ionic bonding involves the transfer of valence electron/s, primarily between a metal and a nonmetal. The electrostatic attractions between the oppositely charged ions hold the compound together.

Ionic compounds:

Are usually crystalline solids (made of ions)

Have high melting and boiling points

Conduct electricity when melted

Are mostly soluble in water and polar solvents

Covalent Bond

A covalent bond is formed when pairs of electrons are shared between two atoms. It is primarily formed between two same nonmetallic atoms or between nonmetallic atoms with similar electronegativity.

Lewis Dot Structure

Lewis structures are also known as Lewis dot structures or electron dot structures.

These are basically diagrams with the element’s symbol in the centre. The dots around it represent the valence electrons of the element.

Lewis structures of elements with atomic number 5-8

Covalent Bonding in H2, N2 and O2

Formation of a single bond in a hydrogen molecule:

Each hydrogen atom has a single electron in the valence shell. It requires one more to acquire the nearest noble gas configuration (He).

Therefore, both the atoms share one electron each and form a single bond.


 Formation of a double bond in an oxygen molecule:

Each oxygen atom has six electrons in the valence shell (2, 6). It requires two electrons to acquire the nearest noble gas configuration (Ne).

Therefore, both the atoms share two electrons each and form a double bond.

Formation of a triple bond in a nitrogen molecule:

Each nitrogen atom has five electrons in the valence shell (2, 5). It requires three electrons to acquire the nearest noble gas configuration (Ne).

Therefore, both atoms share three electrons each and form a triple bond.

Single, Double and Triple Bonds and Their Strengths

A single bond is formed between two atoms when two electrons are shared between them

A double bond is formed between two atoms when four electrons are shared between them, i.e., one pair of electrons from each participating atom. It is depicted by double lines between the two atoms.

A triple bond is formed between two atoms when six electrons are shared between them, i.e., two pairs of electrons from each participating atom. It is depicted by triple lines between the two atoms.

Bond strength:

  The bond strength of a bond is determined by the amount of energy required to break a bond.

This is to signify that the energy required to break three bonds is higher than that for two bonds or a single bond.

Bond length:

Bond length is determined by the distance between nuclei of the two atoms in a bond.

The order of bond length for multiple bonds is: Triple bond<double bond<single bond

The distance between the nuclei of two atoms is least when they are triple bonded.

Covalent Bonding of N, O with H and Polarity

In ammonia (NH3), the three hydrogen atoms share one electron each with the nitrogen atom and form three covalent bonds.

1.Ammonia has one lone pair.

2.This causes the N atom to acquire a slight negative charge and H atom a slight positive charge

3.All three N-H covalent bonds are polar in nature.N atom is more electronegative than

the H atom. Thus, the shared pair of electrons lies more towards N atom.


In water (H2O), the two hydrogen atoms share one electron each with the oxygen atom and form two covalent bonds.

 Covalent Bonding in Carbon

A methane molecule (CH4) is formed when four electrons of carbon are shared with four hydrogen atoms as shown below.


Diamond has a regular tetrahedral geometry. This is because each carbon is connected to four neighbouring carbon atoms via single covalent bonds, resulting in a single unit of a crystal. These crystal units lie in different planes and are connected to each other,  resulting in a rigid three-dimensional cubic pattern of the diamond.


1.Has a high density of 3.5g/cc.

2.Has a very high refractive index of 2.5.

3.Is a good conductor of heat.

4.Is a poor conductor of electricity.


In graphite, each carbon atom is bonded covalently to three other carbon atoms, leaving each carbon atom with one free valency. This arrangement results in hexagonal rings in a single plane and such rings are stacked over each other through weak Van der Waals forces.


1.Has a density of 2.25 g/cc.

2.Has a soft and slippery feel.

3.Is a good conductor of electricity.


C60, also known as Buckminsterfullerene, is the very popular and stable form of the known fullerenes. It is the most common naturally occurring fullerene and can be found in small quantities in soot. It consists of 60 carbon atoms arranged in 12 pentagons and 20 hexagons, like in a soccer ball.

1. Bonding in Carbon



Carbon:-  (i)Most carbon compounds are poor conductors of electricity. Therefore, the bonding in these componds does not give rise to any ions

  1. They have low melting and boiling points as compared to ionic compounds
  2. Forces of attraction between the molecules are not very strong.
  3. The atomic number of carbonis
  4. It has four elements I its outermost shell and needs to gain or lose four electrons to attain noble gas configuration.

If carbon were to gain lose electrons:-

  • It  could gain four electrons forming C4-  anion. But it would be difficult for the nucleus with six protons to hold on to ten electrons.
  • It could lose four electrons forming C4+ cation. But it would require a large amount of energy to remove electrons. Carbon overcomes this problem by sharing its valences electrons with other of carbon or with atoms of other elements.

 Bonding in carbon

 Carbon form covalent bounds.  

  • Covalent bound formation involves sharing of electrons between bonding atoms which may be either same or different.
  • The number of lectrons contributed by an atom for sharing is known as its covalency.

       For example 

  1. Molecule of hydrogen  

  2. Molecule of oxygen

  1. Molecule of Nitrogen

  1. Structure of methane

Methane CH4  is widely used as a fixed and is a major component of bio-gas and compressed natural gas ( CNG) .

  1. These compounds are molecula in nature  i.e. , they exist as single molecules)
  2. These are insoluble in water and soluble in benzene , kerosence and petrol etc.
  3. These compounds are poor conductor of electricity.

Allotropes of carbon

  The property due to which an element exists intwo or more forms, which differ in their physical and chemical properties is known as ‘Allotropes’ and the various forms are called “Allotropes”.

  • carbon exists in two allotropic  form
  • (i) crystalline                 (ii)  amorphous.

The crystalline forms are diamond and graphite the amorphous forms are coat , characol etc.

  • In diamound , each carbon is bonded to for other carbon atoms forming rigid 3-D dtructure. Diamond is the nardest substance. * in graphite each carbon is bonded to three other carbon atom. Graphite structure is formed by the hexagonal arrays. Graphite is smooth and slippery . it is very good conductor of electricity.
  • Fullerenes form another class of carbon allotropes. The first one to be identified was C-60, which has carbon atoms arranged in the shape of a football.

1. Early Classification of Elements

Dobereiner’s Triads

Dobereiner arranged a group of three elements with similar properties in the order of increasing atomic masses and called it a triad. He showed that the atomic mass of the middle element is approximately the arithmetic mean of the other two. But, Dobereiner could identify only the following three triads from the elements known at that time.

Newlands’ Law of Octaves

Assumptions and Limitations:

1. The law was applicable for elements with atomic masses up to 40.

2. Properties of new elements discovered did not fit into the law of octaves.

3. In a few cases, Newlands placed two elements in the same slot to fit elements in the table.

4. He also grouped unlike elements under the same slot.

1. Early Classification of Elements

Chapter 5

Periodic Classification of Elements

Early classification of elements.

  • Classification means identifying similar species and grouping them together.
  • Lavoisier divided lements into two main types known as metals and non- metals.

Doberiner’s law of triads :- 

Doberiner tried to arrange the elements with similar properties. He showed that when the three elements in a triad were written in the order of increasing atomic masses ; the atomic masses ; the atomic mass of the middle element was roughly the average of the atomic masses of other two elements.

i.e atomic masses Li, Na and K are 7,23, and 39 respectively , tus the mean of 1st and 3rd elements is 23 and the atomic mass of middle element is 23.

Limitation :-  He could identify only a few such triads and so the law could not gain importance.

For example, Fe, Co, Ni, all the three elements have nearly equal atomic mass and thus does not follow this law.

Newland’s law of octaves :-

 He found that every eighth element had properties similar to that of the first. HP compared this to the octaves found in music . he called it ‘law of octaves’.

For Example. The properties of lithium and sodium were found to be the same sodium is the eighth  element after lithium.

 Limitation :-

(i) law of octaves was applicable only upto calcium as after calcium every eigth element did not possess properties similar to that of first.

(II)    According  to him , only 56 elements exist in nature and no more elements would be discovered in the future. But later on several new element  were discovered wose properties did not fit into law of octaves. 

(III)    In order to fit new elements into his table newland  adjust two elements in the same column, but put some unlike elements under the same column. 

(IV)    Thus, newland’s  classification was not accepted.

2. Common Things in All Acids and Bases

Hydronium ion

Hydronium ion is formed when a hydrogen ion accepts a lone pair of electrons from the oxygen atom of a water molecule, forming a coordinate covalent bond.


Dilution is the process of reducing the concentration of a solution by adding more solvent (usually water) to it.

It is a highly exothermic process. Strong acid or base: When all molecules of a given amount of an acid or a base dissociate completely in water to furnish their respective ions, H+(aq) for acid and OH−(aq) for base).

Weak acid or base: When only a few of the molecules of a given amount of an acid or a base dissociate in water to furnish their respective ions, H+(aq) for acid and OH−(aq) for base).

Dilute acid: contains less number of H+(aq) ions per unit volume.

Concentrated acid: contains more number of H+(aq) ions per unit volume

2. Common Things in All Acids and Bases

Common Things in All Acids and Bases:-

All acids have H+ ios in common

All bases have OH- ions in common

 Acids produce H+ ions which  ar responsible for their acidic properties.

Acids solution in water conducts electricity.

An Acid or a Base in water solution:-

  • Acids produce H+Ions in presence of water.
  • H+ ions cannot exist alone , they exist as H3O+ ions.

H+   +   H2O      H3O+

HCl  +  H2O     H3O+ +  Cl-

Bases when dissolved in water gives  Oh-  ions.

For Example:- 


Bases soluble in water are called alkali.

 It is recommended that the acid / base should be added to water and not water. Is added to acid/ base, the heat generated may cause the mixture to splash out and cause burns and the glass

conatiner may also break due to exessing  heating.

  Mixing an acid or base with water results in decrease of concentration of ions per unit volume. Such a process is called dilution.

3. Strongness about Acids or Bases

Universal indicator

A universal indicator has a pH range from 0 to 14 that indicates the acidity or alkalinity of a solution.

A neutral solution has pH=7


In pure water, [H+]=[OH]=10−7 mol/L. Hence, the pH of pure water is 7.
The pH scale ranges from 0 to 14.
If pH < 7 
acidic solution
If pH > 7
basic solution

Importance of pH in everyday life

1. pH sensitivity of plants and animals

Plants and animals are sensitive to pH. Crucial life processes such as digestion of food, functions of enzymes and hormones happen at a certain pH value.

2. pH of a soil

The pH of a soil optimal for the growth of plants or crops is 6.5 to 7.0.

3. pH in the digestive system

The process of digestion happens at a specific pH in our stomach which is 1.5 to 4.
The pH of the interaction of enzymes, while food is being digested, is influenced by HCl in our stomach.

4. pH in tooth decay

Tooth decay happens when the teeth are exposed to an acidic environment of pH 5.5 and below.

5. pH of self-defence by animals and plants

Acidic substances are used by animals and plants as a self-defence mechanism. For example, bee and plants like nettle secrete a highly acidic substance for self-defence. These secreted acidic substances have a specific pH.

3. Strongness about Acids or Bases

Strongness  about  Acids or Bases 

Strength of acid and base can be determined by universal indicator.

concentrations of H+ ions in the solution.

pH scale :-  A scale for measuring hydrogen ion concentration in a solution called PH scale.

pH = 7                           Neutral Solution

pH Less than 7             Acidic  Solution

pH more than 7           Basic solution

pH of some common substances shown on a pH paper.

Importance of  pH in everyday life :-

1 ,  Plants and animals are Ph sensitive.

  • Our body works within the pH range of  7- 7.8 .
  • When pH of rain water is less than 5.6, it is called acid rain.
  • When acid rain flows into the rivers ,it lower the pH of river water and makes the survival  of aquatic life difficult.

2, pH of the soil

  • Plants require a specific pH range for their healthy growth.

3,  pH in our digestive system.

  • Our stomach produce hydrogenchloric acid which helps in digestion without harming the stomach.
  • During indigestion, stomach produces more acid and cause pain and irritation.
  • To get rid of pain ,prople use mild bases called antacid to neutralize the excess acid. Magnesium hydroxide ( Milk of magnesia) is an antacid.

  4,  Ph changes as the cause of tooth decay.

  • Tooth decay starts when pH of mouth is lower than 5.5
  • Tooth enamel is made up of calcium phosphate ( hardest substance in body) .it does not dissolve in water but carodes when pHis low than 5.5
  • Using basic toothpaste , tooth decay can be prevented    

5, self defence by animals and plants through chemical warfare.

  • Bee sting leaves an acid which cause pain and irritation. Bakig soda (mild base) gives relief by rubbing it on stung area.
  • Stinging hair of nettle leaves inject  methanoic  acid causing  buring or pain rubbing this with leaf of dock plant give relief.

4. About Salts


A salt is a combination of an anion of an acid and a cation of a base.

Examples – KCl, NaNO3 ,CaSO4, etc.

Salts are usually prepared by the neutralisation reaction of an acid and a base.

Common salt

Sodium Chloride (NaCl) is referred to as common salt because it’s used all over the world for cooking. Sodium chloride (NaCl) is also known as Common or Table Salt. It is formed after the reaction between sodium hydroxide and hydrochloric acid. It is a neutral salt. The pH value of sodium chloride is about 7. Sodium chloride is used to enhance the taste of food. Sodium chloride is used in the manufacturing of many chemicals.

Family of salts

Salts having the same cation or anion belong to the same family. For example, NaCl, KCl, LiCl.

pH of salts

A salt of a strong acid and a strong base will be neutral in nature. pH = 7 (approx.).

A salt of a weak acid and a strong base will be basic in nature. pH > 7.

A salt of a strong acid and a weak base will be acidic in nature. pH < 7.

The pH of a salt of a weak acid and a weak base is determined by conducting a pH test.

Preparation of Sodium hydroxide

Chemical formula – NaOH

Also known as – caustic soda

Preparation (Chlor-alkali process):

Electrolysis of brine (solution of common salt, NaCl) is carried out.

At anode: Cl2 is released

At cathode: H2 is released

Sodium hydroxide remains in the solution.

Bleaching powder

Chemical formula – Ca(OCl)Cl or CaOCl2

Preparation – Ca(OH)2(aq)+Cl2(g)→CaOCl2(aq)+H2O(l)

On interaction with water – bleaching powder releases chlorine which is responsible for bleaching action.

Baking soda

Chemical name – Sodium hydrogen carbonate

Chemical formula – NaHCO3

Preparation (Solvay process):

a. Limestone is heated: CaCO3→CaO+CO2

b. CO2 is passed through a concentrated solution of sodium chloride and ammonia:



1. Textile industry

2. Paper industry

3. Disinfectant

Washing soda

Chemical name – Sodium hydrogen carbonate

Chemical formula – NaHCO3

Preparation (Solvay process) – 

a. Limestone is heated: CaCO3 → CaO + CO2

b. CO2 is passed through a concentrated solution of sodium chloride and ammonia:

NaCl(aq) + NH3(g) + CO2(g) + H2O(l) → NaHCO3(aq) + NH4Cl(aq)


1. In glass, soap and paper industries

2. Softening of water

3. Domestic cleaner

Crystals of salts

Certain salts form crystals by combining with a definite proportion of water. The water that combines with the salt is called water of crystallisation.

Plaster of paris

Gypsum, CaSO4.2H2O (s) on heating at 100°C (373K) gives CaSO4. ½ H2O and 3/2 H2O

CaSO4. ½ H2O is plaster of paris.

CaSO4. ½ H2O means two formula units of CaSO4 share one molecule of water.

Uses – cast for healing fractures.

4. About Salts

About  salts

  1. Strong acid  +  strong base       neutral salt :  pH = 7
  2. Salt of strong   +  weak base      Acidic salt :    pH <7
  3. Salt of strong base + weak acid  Basic salt  :  pH  >7

Chemical from common salt

1;  sodium hydroxide  ( NaOH)

  When electricity is passed through an aqueous solution of NaCl (called brine) it decomposes to form NaOH.

This process is called chlor- alkali process.

2NaCl  +   2H2O              2NaOH + Cl2    + H2

 At cathode  :     H2    gas

 At anode :    Cl2 gas

Near cathode :   NaOH solution is formed


  1. Cl2 :   Water treatment , PVC , pesticides.
  2. H2 :      Fuels,  margarine.
  3. Hcl :       Medicines, cleaning steels.
  4.  NaOH :   De- greasing metals, soaps and paper making.

(2) Bleaching powder  ( CaOCl2)

   Bleaching powder is produced by the action of chlorine on dry slaked lime   ( Ca(OH)2).

 Cl2  +  Ca(OH)2        CaOCl2  +   H2O

 USES :   * Bleaching cotton and linen in textile industry

  • Bleaching washed clothes in laundry .
  • Oxidising in chemical industries.
  •  Make drinking water free from germs

(3)   Baking Soda ( NaHCO3)

 The chemical name of the compound is sodium hydrogen carbonate.

NaCl+H2O+CO2+NH3      NH4Cl  +  NaHCO3

                                     (Ammonium Chloride) (Sodium hydrogen carbonate)

  • IT is a mild non- corrosive base.
  • When it is heated during cooking.

2 NaHCO3      Na2CO3  +   H2O  +  CO2


  1. For making baking powder, which is a mixture of baking soda ( sodium hydrogen carbonate ) and tartaric acid. Following reaction takes place.

 NaHCO3   +  H+         CO2  H2O + Sodium salt of acid

  1. Sodium hydrogen carbonate is an ingredient in antacids.
  2. Used in soda-acid , fire extinguisher

 (4)  Washing Soda  :-   ( Na2 CO3IOH2O)

  Recrystallisation of sodium carbonate gives washing sod. It is a basic salt.

  Na2CO3 + IOH2O     →    Na2CO3. IOH2O


  1. In glass , soap and paper industries.
  2. Manufacture of borax
  3. Cleaning agent for domestic purpose.
  4. For removing permanent hardness of water.

 (5)  Plaster of paris CaSO4 .2H2 O. 

On heating gypsum ( CaSo4.2H2O) at 373k, ut loses water molecules and becomes plaster of paris (POP) .

It is a white powder and on mixing with waater it changes to gypsum.

 Removing water of crystallization.


It is a fixed number of water molecules present in one formula unit of a salt.

For Example; 

  • CuSO4.5H2O   has 5 water molecules.
  • Na2CO3.10H2  has 10 water molecules.
  •  CaSO4.2H2O  has 2 water molecules.

2. Chemical Properties of Metals

Chemical Properties of Metals

1. Reaction with oxygen: Most of the metals form respective metal oxides when reacting with oxygen.

Metal + Oxygen → Metal Oxide


Reaction of Potassium with Oxygen: Potassium metal forms potassium oxide when reacts with oxygen.

Reaction of Sodium with Oxygen: Sodium metal forms sodium oxide when reacts with oxygen.

 Reaction of Copper metal with Oxygen: Copper does not react with oxygen at room temperature but when burnt in air, it gives oxide.

Silver, gold and platinum do not combine with the oxygen of air even at high temperature. They are the least reactive.

2. Reaction of metals with water: Metals form respective hydroxide and hydrogen gas when reacting with water.

Metal + Water → Metal hydroxide + Hydrogen

Reaction of Sodium metal with Water: Sodium metal forms sodium hydroxide and liberates hydrogen gas along with lot of heat when reacting with water.

Reaction of Iron with Water: Reaction of iron with cold water is very slow and comes into notice after a long time. Iron forms rust (iron oxide) when reacts with moisture present in the atmosphere. Iron oxide and hydrogen gas are formed by passing of steam over iron metal.

3. Reaction of metals with dilute acid: Metals form respective salts when reacting with dilute acid. Reaction of Sodium metal with dilute hydrochloric acid: Sodium metal gives sodium chloride and hydrogen gas when react with dilute hydrochloric acid.

Hydrogen gas is not when metal is treated with nitric acid (HNO3):

Nitric acid is strong oxidising agent and it oxidises the hydrogen gas (H2) liberated into water (H2O) and itself get reduced to some oxide of nitrogen like nitrous oxide (N2O)3 nitric oxide (NO) and nitrogen dioxide (NO2).

Copper, gold, silver are known as noble metals. These do not react with water or dilute acids.

The order of reactivity of metal towards dilute hydrochloric acid or sulphuric acid is in the order;

K > Na > Ca > Mg > Al > Zn > Fe > Cu > Hg > Ag

Metal Oxides

Chemical Properties: Metal oxides are basic in nature. The aqueous solution of metal oxides turns red litmus blue.

Reaction of Metal oxides with Water: Most of the metal oxides are insoluble in water. Alkali metal oxides are soluble in water. Alkali metal oxides give strong base when dissolved in water.

Reaction of Sodium oxide with Water: Sodium oxide gives sodium hydroxide when reacts with water.

Reaction of Zinc oxide and Aluminium oxide: Aluminum oxide and zinc oxide are insoluble in water. Aluminium oxide and zinc oxide are amphoteric in nature. An amphoteric substance shows both acidic and basic characters. It reacts with base like acid and reacts with an acid like a base.

When zinc oxide reacts with sodium hydroxide, it behaves like an acid. In this reaction, sodium zincate and water are formed

Reactivity Series of Metals: The order of intensity or reactivity of metal is known as Reactivity Series. Reactivity of elements decreases on moving from top to bottom in the given reactivity series. In the reactivity series, copper, gold, and silver are at the bottom and hence, least reactive. These metals are known as Noble metals.

Reactivity of some metals are given in descending order :

K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu


2. Chemical Properties of Metals

Chemical properties of Metals

(i)  Reaction with air :-

All metals combine with oxygen to form metal oxide .

Metal + O2           Metal oxide

For example,

2Cu  +  O2          2CUO    copperoxide (black)

 4Al + 3O2        2Al2O3     Aliminium oxide

  • Sodium and potassium react so vigorously that they catch fire in open so they are kept immersed in kerosence
  • Surfaces of Mg , Al, Zn pb are covered with a thin layer of oxide whish prevent them from further oxidation. Anodizing is a process of forming a tick oxide layer of aluminium.
  • Iron does not burn on heating but iron filling burn vigorously.
  •  Copper does not burn but the hot metal is coated with a black coloured layer of copper (ii) oxide
  • Silver and gold do not react with oxygen even at high temperatures.

Amphoteric oxide    Metal oxides which react with both acids as well as bases to produce salts and water are called amphoteric oxides.

Example ;   

Al2O3   +    6HCl      2AlCl3 +   H2O

 Al2O3  +  2NaOH    2NaAlO2    +   H2O

                                (SODIUM ALUMINATE )

Most metal oxides are insoluble in water but some of these dissolve in water to form alkalis. Sodium oxide and potassium oxide dissolved in water to produce alkali.

Na2O(S) +    H2O(l)        2 NaOH (aq)

K2O(S)    +   H2O(l)         2 KOH (aq)

(ii)  Reactions of metals with water:-

 Metal+ water        Metal oxide  +  hydrogen

Metal oxide  +  water      metal hydroxide

Metals like potassium and sodium react violenty with cold water.

Na  +   2 H2O          NaOH + H2  + heat energy.

The reaction of calcium of water is less violent

Ca + 2H2O    →   Ca(oH)2 + H2

Magnesium react with hot water to form magnesium hydroxide and hydrogen.

Mg +  2H2O     →  Mg(OH)   H2

Metals like aluminium iron and zinc do not react with cold or hot water. But they react with steam to form metal oxide and hydrogen.

2 Al + 3 H20         Al2O3  +   3H2

3 Fe  + 4 H2O     →  Fe3O4   +   4H2

Metal such as lead ,copper, silver and gold do not react with water at all.

(ii) Reaction of metals with acids.

Metal +  Dilute acid     salt + hydrogen

Copper and silver do not react wit dil acids.

 For example 

 Fe +  2HCl          FeCl2  +  H2

Mg  +  2HCl          MgCl2  +   H2

Zn  +   2HCl           ZnCl2  +  H2

(iii)   Reaction of metals with solution of other Metal salts ; 

Metal A+  Salts solution B      Salt solution A + Metal B   

Reaction of metals with salt solutions.

More reactive metals can displace less reactive metals from their compounds in solution form.

Fe +  CuSO4          FeSO4  +    Cu       

Fe displaces Cu because Fe is more reactive metals than Cu .


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

3. Metals and Non-Metals Reactions

Ionic Compounds

The electrostatic attractions between the opposite charged ions hold the compound together.
Example: MgCl2, CaO, MgO, NaCl etc.

Properties of Ionic Compound

Ionic compounds

  1. Are usually crystalline solids (made of ions).
  2. Have high melting and boiling points.
  3. Conduct electricity when in aqueous solution and when melted.

Physical Nature

Ionic solids usually exist in regular, well-defined crystal structures.

Electric Conduction of Ionic Compounds

Ionic compounds conduct electricity in the molten or aqueous state when ions become free and act as charge carriers.
In solid form, ions are strongly held by electrostatic forces of attractions and are not free to move; hence do not conduct electricity.

For example, ionic compounds such as NaCl does not conduct electricity when solid but when dissolved in water or in a molten state, it will conduct electricity.

Salt solution conduct electricity

Melting and Boiling Points of Ionic Compounds

In ionic compounds, the strong electrostatic forces between ions require a high amount of energy to break. Thus, the melting point and boiling point of an ionic compound are usually very high.

Solubility of Ionic Compounds

Most ionic compounds are soluble in water due to the separation of ions by water. This occurs due to the polar nature of water.
For example, NaCl is a 3-D salt crystal composed of Na+ and Cl ions bound together through electrostatic forces of attractions. When a crystal of NaCl comes into contact with water, the partial positively charged ends of water molecules interact with the Cl ions, while the negatively charged end of the water molecules interacts with the Na+ ions. This ion-dipole interaction between ions and water molecules assist in the breaking of the strong electrostatic forces of attractions within the crystal and ultimately in the solubility of the crystal.

.Properties of Ionic compound

  • Ionic compounds are solid. Ionic bond has a greater force of attraction because of which ions attract each other strongly. This makes ionic compounds solid.
  • Ionic compounds are brittle.
  • Ionic compounds have high melting and boiling points because force of attraction between ions of ionic compounds is very strong.
  • Ionic compounds generally dissolve in water.
  • Ionic compounds are generally insoluble in organic solvents; like kerosene, petrol, etc.
  • Ionic compounds do not conduct electricity in the solid state.

3. Metals and Non-Metals Reactions


  • Atoms of the metals lose electron from their valence shell to form cation.
  • Atoms of the non- metals gain electrons in the valence to form anion.

Formation of sodium chloride.

Formation of magnesium chloride

Properties of ionic compound

  1. PHYSICAL NATURE  :-  They are solid and hard (because of the strong force of attraction between the positive and negative ions) . They are brittle.
  2. Melting and boiling point:_  They have high melting and boiling pont.
  3. Solubility :- soluble in water and insoluble in solvents such as kerosene , petrol etc.
  4. Conductor of electricity :- Ionic compound conduct electricity in molter (ions move to the opposite electrodes  when electricity is passed )
  • They do not conduct electricity in solid state as movements of ions is not possible in solid They conduct electricity in molten state.

4. Occurence of Metals

Occurrence of Metals

Most of the elements, especially metals occur in nature in the combined state with other elements. All these compounds of metals are known as minerals. But out of them, only a few are viable sources of that metal. Such sources are called ores.

Au, Pt – exist in the native or free state.

Extraction of Metals

Metals of high reactivity – Na, K, Mg, Al.

Metals of medium reactivity – Fe, Zn, Pb, Sn.

Metals of low reactivity – Cu, Ag, Hg


Converts sulphide ores into oxides on heating strongly in the presence of excess air.

It also removes volatile impurities.



Converts carbonate and hydrated ores into oxides on heating strongly in the presence of limited air. It also removes volatile impurities.




Extracting Metals Low in Reactivity Series

By self-reduction- when the sulphide ores of less electropositive metals like Hg, Pb, Cu etc., are heated in air, a part of the ore gets converted to oxide which then reacts with the remaining sulphide ore to give the crude metal and sulphur dioxide. In this process, no external reducing agent is used.

1. 2HgS(Cinnabar)+3O2(g)+heat→2HgO(crude metal)+2SO2(g)


2. Cu2S(Copperpyrite)+3O2(g)+heat→2Cu2O(s)+2SO2(g)

2Cu2O(s)+Cu2S(s)+heat→6Cu(crude metal)+SO2(g)

Extracting Metals in the Middle of Reactivity Series

Smelting – it involves heating the roasted or calcined ore (metal oxide) to a high temperature with a suitable reducing agent. The crude metal is obtained in its molten state.


Aluminothermic reaction – also known as the Goldschmidt reaction is a highly exothermic reaction in which metal oxides usually of Fe and Cr are heated to a high temperature with aluminium.



Extraction of Metals Towards the Top of the Reactivity Series

Electrolytic reduction:

1. Down’s process: Molten NaCl is electrolysed in a special apparatus.

At the cathode (reduction):


Metal is deposited.

At the anode (oxidation):


Chlorine gas is liberated.

2. Hall’s process: Mixture of molten alumina and a fluoride solvent usually cryolite, (Na3AlF6) is electrolysed.

At the cathode (reduction):

2Al3++6e–→ 2Al(s)

Metal is deposited.

At the anode (oxidation):

6O2– → 3O2(g)+12e–

Oxygen gas is liberated.

Enrichment of Ores

It means the removal of impurities or gangue from ore, through various physical and chemical processes. The technique used for a particular ore depends on the difference in the properties of the ore and the gangue.

Refining of Metals

Refining of metals – removing impurities or gangue from crude metal. It is the last step in metallurgy and is based on the difference between the properties of metal and the gangue.

Electrolytic Refining

Metals like copper, zinc, nickel, silver, tin, gold etc., are refined electrolytically.

Anode: impure or crude metal

Cathode: a thin strip of pure metal

Electrolyte: aqueous solution of metal salt

From anode (oxidation): metal ions are released into the solution

At cathode (reduction): the equivalent amount of metal from solution is deposited

Impurities deposit at the bottom of the anode.

4. Occurence of Metals


MINERALS :- The elements which occur naturally in the earth’s crust are called minerals.

ORES:- Minerals that contain very high percentage of particular metal andthe metal can be profitably extracted from it, such minerals are called ores.

  • Metals at the bottom of the activity series are least reactive they are often found in free state . For Example –  Ag,Au, Cu.
  • Metals at the ttop of the acitivity series (k,Na, Co, Mg, and Al) are so reactive that they never found in free state.
  • Metals in the middle of the activity series ( Zn, Fe,Pb etc are moderately reactive . they occur as sulphates ,oxides or carbonates.
  • They ore of many metals are oxide because oxygen is very reactive and is abundant on the earth.
  • Steps involved in the extraction of pure metals from ores.

 Step involved in the extraction of metals from ores.


 1 ENCRICHMENT OF ORES    Ores are usually contaminated with large amounts of impurities such as soil, sand etc called gangue these impuriestes are removed from the ore prior to the extraction of mertal.

 2  Extraction of metals   Metals low in the activity series are very anreactive the oxides of these metals are reduced to metals by heating.

 For example           

 2HgO         2Hg  +  O2

Mercury oxide is reduced to mercury on heating

The metals in the middle of the activity series ( Zn, Fe, Pb, Cu ) are moderately active. The metal sulphides  and carbontes  are converted into metal oxide. The sulphate ores are converted into oxides by heating strongly in the presence of excess air this process is known as roasting .

2 ZnS + 3O2 →      2ZnO + CO2

  • The carbonate ores are changed into oxides by heating strongly in limited air this process is called calcination .

 Zn CO3          ZnO + CO2

  • Then metal oxides are reduced to corresponding metals by using reducing agent like carbon.

 ZnO  +  C       Zn   +  CO

  • This reaction of iron (iii) oxide ( Fe2O3) with aluminium is used to join railway tracks or cracked machine parts . this reaction is known as thermit reaction.

 Fe2O3    +  2Al      2 Fe +  Al2O3 + heat

 Metals high up in the reactivity series are very reactive. The metalo are obtained by electrolytic reduction. The metals are deposit at the cathode and chlore is deposited at anode.

At cathode     Na+    +  e-         Na
At anode        2Cl-       Cl2   +  2e-

Refining of metals

 The most widely used method for refining impure metals is electronic refining.

Electrolytic refining of copper

Electrolytic refining :-   Metals  ( Cu, Zn , Ag, Au etc ) are refined electrolytically . the impure metal is made the anode and athin strip of pure metal is made the cathode . a solution a metal salt is used as an electrolyte.

Electrolytic refining of copper.

Anode :  Impure copper

Cathode: Strip of pure copper

The insoluble impurities settle at the bottom of the anode and is called anode mud.  

5. Corrosion



Alloys are homogeneous mixtures of metal with other metals or nonmetals. Alloy formation enhances the desirable properties of the material, such as hardness, tensile strength and resistance to corrosion.

Examples of a few alloys:

Brass: copper and zinc

Bronze: copper and tin

Solder: lead and tin

Amalgam: mercury and other metal


Gradual deterioration of material usually a metal by the action of moisture, air or chemicals in the surrounding environment.


4Fe(s)+3O2(from air)+xH2O(moisture)→2Fe2O3. xH2O(rust)

Corrosion of copper:

Cu(s)+H2O(moisture)+CO2(from air)→CuCO3.Cu(OH)2(green)

Corrosion of silver:

Ag(s)+H2S(from air)→Ag2S(black)+H2(g)

Prevention of Corrosion


1. Coating with paints or oil or grease: Application of paint or oil or grease on metal surfaces keep out air and moisture.

2. Alloying: Alloyed metal is more resistant to corrosion. Example: stainless steel.

3. Galvanization: This is a process of coating molten zinc on iron articles. Zinc forms a protective layer and prevents corrosion.

4. Electroplating: It is a method of coating one metal with another by the use of electric current. This method not only lends protection but also enhances the metallic appearance.

5. Corrosion


The surface of metals is corroded when they are exposed to moist air for a long period of time. This is called corrosion

For example 

  • Silver becomes black when exposed to air and form a coating of silver sulphide.
  • Copper react with  moist Co2 and form a green coat of copper carbonate.
  • Iron acquires a coating of brown floky substance called rust.

 Prevention of corrosion. 

The rusting of iron can be prevented by painting , oiling, greasing, galvanizing, chrome plating anodizing and making alloys.

  • Galvanisation:-  It is a method of protecting steel and iron from ressting by coating them a thin layer of zinc.
  • Alloying is a very good method of improving the properties of a metal an alloy is a homogenous mixture of two or more metals or a metal and a non- metal
  • Iron      it is mixed with small amount of carbon
  • Steel    iron + nickel and chromium.
  • Brass   copper  + tin
  • Solder   →  lead  +  tin ( used for welding electric wire together )

2. Nature of Carbon Compounds

Physical Properties of Organic Compounds

Most of the organic compounds have low boiling and melting point, due to the weak force of attraction (i.e., the inter-molecular force of attraction) between these molecules. Most carbon compounds are poor conductors of electricity, due to the absence of free electrons and free ions.

Chains, Branches and Rings

Saturated and Unsaturated Hydrocarbons

Saturated hydrocarbons: These hydrocarbons have all carbon-carbon single bonds. These are known as alkanes. General formula = CnH2n+2 where n = 1, 2, 3, 4.…..

Unsaturated hydrocarbons: These hydrocarbons have at least one carbon-carbon double or triple bond.

Hydrocarbons with at least one carbon-carbon double bond are called alkenes. General formula = CnH2n where n = 2, 3, 4…..

Hydrocarbons with at least one carbon-carbon triple bond are called alkynes. General formula = CnH2n−2 where n = 2, 3, 4…..

Chains, Rings and Branches

Carbon chains may be in the form of straight chains, branched chains or rings.

In cyclic compounds,

 Structural Isomers

The compounds with the same molecular formula and different physical or chemical properties are known as isomers and the phenomenon is known as isomerism.

 The isomers that differ in the structural arrangement of atoms in their molecules are called structural isomers and the phenomenon is known as structural isomerism.


Benzene is the simplest organic, aromatic hydrocarbon.

Physical properties: colourless liquid, pungent odour, flammable, volatile.


Cyclic in nature with chemical formula, C6H6, i.e., each carbon atom in benzene is arranged in a six-membered ring and is bonded to only one hydrogen atom.

It includes 3-double bonds which are separated by a single bond.

 Hence, this arrangement is recognized to have conjugated double bonds and two stable resonance structures exist for the ring.

Functional Groups and Nomenclature

Functional Groups

An atom or a group of atoms which when present in a compound gives specific physical and chemical properties to it regardless of the length and nature of the carbon chain is called a functional group.

Classification of Functional Groups

Main Functional Groups:

(i) Hydroxyl group (-OH): All organic compounds containing -OH group are known as alcohols. For example, Methanol (CH3OH), Ethanol (CH3−CH2−OH), etc.

(ii) Aldehyde group (-CHO): All organic compounds containing -CHO group are known as aldehydes. For example, Methanal (HCHO), Ethanal (CH3CHO), etc.

(iii) Ketone group (-C=O): All organic compounds containing (-C=O) group flanked by two alkyl groups are known as ketones. For example, Propanone (CH3COCH3), Butanone (CH3COCH2CH3), etc.

(iv) Carboxyl group (-COOH): All organic acids contain a carboxyl group (-COOH). Hence, they are also called carboxylic acids.

For example, Ethanoic acid (CH3COOH), Propanoic acid (CH3CH2COOH), etc.

(v) Halogen group (F, CI, Br, I): The alkanes in which one or more than one hydrogen atom is substituted by- X (F, CI, Br or I) are known as haloalkanes. For example, Chloromethane (CH3Cl), Bromomethane (CH3Br), etc.

Homologous Series

Homologous series constitutes organic compounds with the same general formula, similar chemical characteristics but different physical properties. The adjacent members differ in their molecular formula by −CH2.

Physical Properties

The members of any particular family have almost identical chemical properties due to the same functional group. Their physical properties such as melting point, boiling point, density, etc., show a regular gradation with the increase in the molecular mass. 

2. Nature of Carbon Compounds

Nature of carbon compounds.

Catenation:-     The property of elements to form long chains or rings by self linking of their own atoms through covalentbonds is called catenation. These compounds may have long, branched chains of carbon atoms may be linked by single, double or triple bound. The extent of catenation depends upon the strength of the bounds between the atoms involved in catenation.

Saturated and unsaturated compounds

  • Compounds of carbon which are linked by only single bonds between carbon atoms are called saturated compounds.

For example,    C   -    C  

Carbon atom linked together with single bound.

But the valencies of each carbon atom remain unsatisfied , so each carbon atom is bonded to 3 hydrogen atoms:-

atom is bonded to 3 hydrogen atoms:-

 C2H6 is called methane.

Electron dot structure of ethane

Compounds of carbon having double or triple bounds between their carbon atoms are called unsaturated compounds.

For example,      

C 2 H 2  is called ethyne

.To satisfy the valency , carbon form double bond.

Strength of compound

The bonds that carbon forms with other elements are very strong, so these    compounds because very stable. Carbon form strong bonds is due to its small size. Nucleus hold shaired pair of electrons strougly.

Chains , branches and rings,

Straight chain compounds: the compounds which conatin straight chain of carbon atom e.g butane (c4H10 ) petane ( C5H12) etc.

Branched chain compounds:   Those compounds which are branched.

E.g.    ISO- butane  ( C4H10), isopentane ( C5H12)

Ring compounds:-  they  are also known as closd chain compounds.

Cyclic compounds are called ring coumpounds.

  E.g.  cyclohexane (C6H12)

           Cyclopropane ( C3H6)    etc. 

Hydrocarbons :-  All those compounds which contain only carbon and hydrogen are called hydrocarbon. The saturated hydrocarbon which contain single bond are called alkanes. The unsaturated hydro carbons which contain one or more double bonds are called alkanes those containing one or more triple bounds are called alkynes.

Functional group:-  The atoms or group of atoms which determine the properties of a copmpond is known as functional group.   

E.g.  –Cl  ,  -Br  ( choro/ bromo alkane

             -OH (alcohol) 

Homologous series :-  A series of compounds in which the same functional group substitutes hydrogen ina carbon chain is called a homologous series.   E.g   ,     CH3C1    and   C2H5C1     these differ by a     CH2 unit. 

 Nomenclature :-   For naming organic compounds based on their structures, are followed by UPSC rules.

IUPAC name of an organic compounds consists of 3 parts –

  1. Prefix :- in case functional group is present , it is indicated in the name of the compound with either as a prefix or as a suffix.
  2. Word root :- A word root indicates the nature of basic carbon skeleton.
  3. Suffix :- while adding the suffix to the word rrot, the terminal ‘e’ of carbon chain is removed . if the carbon chain is unsaturated , then final ‘are’ is substituted by ‘en and yne’ respectively for double and triple bonds.

3. Chemical Properties of Carbon Compounds

Chemical Properties

Combustion Reactions

Combustion means burning of carbon or carbon-containing compounds in the presence of air or oxygen to produce carbon dioxide, heat and light.

Flame Characteristics

Saturated hydrocarbons give clean flame while unsaturated hydrocarbons give smoky flame. In the presence of limited oxygen, even saturated hydrocarbons give smoky flame.



The reactions in which two molecules react to form a single product having all the atoms of the combining molecules are called addition reactions.
The hydrogenation reaction is an example of the addition reaction. In this reaction, hydrogen is added to a double bond or a triple bond in the presence of a catalyst like nickel, palladium or platinum.


The reaction in which an atom or group of atoms in a molecule is replaced or substituted by different atoms or group of atoms is called substitution reaction. In alkanes, hydrogen atoms are replaced by other elements.

CH4+Cl2+Sunlight → CH3Cl+HCl

3. Chemical Properties of Carbon Compounds


(i)  Combustion :- carbon burns in oxygen to give carbon dioxide along with the release of heat and heat .

 C  +   O2         CO2 +  heat and light.

CH4  +  O2       CO2  +  H2O + heat and light

CH3CH2OH + O2   CO2 + H2O + heat and light

Saturated hydrocarbon will give a clean flame unsaturated carbon compounds will give a yellow flame with lots of balck soke. The gas slove used at home has inlets for air so that a sufficiently oxygen-rich mixture is burnt to give a clean blue flame. Fuel such as  coal  and  petroleum  have some amount of nitrogen and sulphur in them. The combustion results in the formation of oxides of sulphur and nitrogen which are major pollutants in the environment.

•  Coal Coal and petroleum has been formed from biomass. Coal is the remains of trees, forms that lived millions of year ago . oil and gas are the remains of millions of tiny plants and animals that lived in the sea.

(ii) Oxidation :-   The substance which are used for oxidation are known as oxidizing agent.

E.g  alkaline kMnO4  , acidifed  k2 Cr2O7 .


(iii) Addition Reaction  unsatured hydrocarbons ( alkenes and alkynes) add hydrogen in the presence of catalysts to give saturated hydrocarbons. 

Catalyst are substance that cause a reaction to occur or proced ata different rate without the reaction itself being affected

  (IV) Substitution Reaction 

Saturated hydrocarbons give substitution reaction e.g methane in presence of sunlight undergo chlorination.

CH4  +  Cl2        CH3Cl +  Hcl (in the presence of sunlight )

4. Carbon Compounds- Ethanol and Ethanoic Acid

Ethanol and Ethanoic Acid


(i) Ethanol, C2H5OH is a colourless liquid having a pleasant smell.

(ii)   It boils at 351 K.

(iii)  It is miscible with water in all proportions.


1. As a solvent in the manufacture of paints, dyes, medicines, soaps and synthetic rubber.

2. As a solvent to prepare the tincture of iodine.

How Do Alcohols Affect Human Beings?

(i)     It causes addiction, damages the liver if taken in excess.

(ii)    High consumption of ethanol may even cause death.

Reactions of Ethanol with Sodium

Ethanol reacts with sodium to produce hydrogen gas and sodium ethoxide. This reaction supports the acidic character of ethanol.

  2C2H5OH+2Na → 2C2H5ONa+H2(↑)

Elimination Reaction

An elimination reaction is a type of reaction in which two substituents are removed from a molecule. These reactions play an important role in the preparation of alkenes.

Dehydration Reaction

Ethanol reacts with concentrated sulphuric acid at 443 K to produce ethylene. This reaction is known as dehydration of ethanol because, in this reaction, a water molecule is removed from the ethanol molecule.


(reaction taking place in presence of Conc.H2SO4)

Ethanoic Acid or Acetic Acid

(i)        Molecular formula: CH3COOH

(ii)       It dissolves in water, alcohol and ether.


SaponificationWhen a carboxylic acid is refluxed with alcohol in the presence of a small quantity of conc.H2SO4, a sweet-smelling ester is formed. This reaction of ester formation is called esterification.

 A soap is a sodium or potassium salt of long-chain carboxylic acids (fatty acid). The soap molecule is generally represented as RCOONa, where R = non-ionic hydrocarbon group and  −COO−Na+ ionic group. 

Reaction of Ethanoic Acid with Metals and Bases

Ethanoic acid (Acetic acid) reacts with metals like sodium, zinc and magnesium to liberate hydrogen gas.


It reacts with a solution of sodium hydroxide to form sodium ethanoate and water.


Carboxylic acids react with carbonates and bicarbonates with the evolution of CO2 gas. For example, when ethanoic acid (acetic acid) reacts with sodium carbonate and sodium bicarbonate, CO2 gas is evolved.



Friendly Carbon

Why Carbon Can Form so Many Compounds

Catenation occurs most readily with carbon due to its small size, electronic configuration and unique strength of carbon-carbon bonds. Tetravalency, catenation and tendency to form multiple bonds with other atoms account for the formation of innumerable carbon compounds.


Catenation is the self-linking property of an element by which an atom forms covalent bonds with the other atoms of the same element to form straight or branched chains and rings of different sizes. It is shown by carbon, sulphur and silicon.


In its native state, sulphur show catenation up to 8 atoms in the form of S8 molecule. It

4. Carbon Compounds- Ethanol and Ethanoic Acid

Carbon compounds ethanol and ethanoio acid

 Alcohol :-  compounds containing –OH group attached to a carbon atom are known as alcohol.

E.g, Ethanol ( C2H5OH ) : Commonly known as alcohol.

 Properties of ethanol

Ethanol is a liquid at room temperature and is soluble n a water . in take of ever a small quantity of pure ethanol ( called absolute alcohol ) can be lethal.

Reactions of ethanol

2Na  +  2CH3CH2OH       2CH3CH2ONa+  +  H2 

                                                 Sodium ethoxide

Ethanol reacts with sodium to liberate H2 gas.

Reaction with cone. H2SO4

Heating ethanol at 443 K with excess concentrated sulphuric acid results in the dehydration of ethane to give ethane.

                              HO+ conc

Alcohol as a fuel :- alcohol (ethanol) is added to petrol up to  20% and the mixture is called gasol.

Harmful effects of drinking  Alcohol :

  • If the alcohol is used for drinking purpose conatins some methyl alcohol ( CH3OH) as impurity then it may cause series poisoning and loss of eye sight.
  • It damage liven if taken regularly in large quantities.
  • Dyes are also added to colour the alcohol blue so that it can be identified easily.
  • This is called denatured alcohol.

Ethanoic Acid :-  Ethanoic acid, commercially known as acetic acid belongs to a group of a acids called carboxylic acid. 

Reactions of ethanoic  acid :

  1. Esterification Reaction :- ethanoic acid reacts with absolute ethanol in the presence of an acid catalyst to give an ester.

Ester are sweet – smelling substances.

These are used in making perfums.

•  Saponification :-   ester is converted back to alcohol and sodium salt of carboxylic acid . this reactionis known as saponification because it is used in the preparation of soap.

  1. Reaction with a base

  NaOH  +  CH3COONa  +  H2O

  1. Reaction with carbonates and bicarbonates

Ethanoic acid reacts with carbonates and hydrogencarbonates to give rise toa salt, carbon dioxide and water . the salt produced is commonly called sodium acetate.

   2CH3COOH  +  Na2CO3         2CH3COONa +  H2O + CO2

  CH3COOH +  NaHCO3          CH3COONa + H2O +  CO2

5. Soaps and Detergents

Soaps and Detergents

Cleansing Action of Soap

When soap is added to water, the soap molecules uniquely orient themselves to form spherical shape micelles.

 The agitation or scrubbing of the fabric helps the micelles to carry the oil or dirt particles and detach them from the fibres of the fabric.

5. Soaps and Detergents


Soap :-   soaps are sodium or potassium molecules is towards the oil droplet while the ionic- end faces outsides. The soap micelle thus helps in pulling out the dirt in water and clothes can be cleaned. salts of long chain acid carboxylic acid. Structures of soap molecules called micelle , where one end of the

Micelles :-  soaps re the molecules in which the two ends have differing properties, one I hydrophilic, it interacts with water , while the other end is hydrophobic i.e it interacts with hydrocarbons. These molecules have a unique orientation that keeps the hydrocarbon portion out of water . this is called

Detergent :- they are ammonium or sulphate salts of long chain carboxylic acids. The charged ends of these compounds do not form insoluble precipitates with the calcium and magnesium oins  in hard water. Thus, they remain effective in hard water. Detergents are used to make shampoos and products for cleaning soaps. 

2. Mendeleev's Periodic Table

Mendeleev’s Periodic Table and Law

The physical and chemical properties of elements are periodic functions of their atomic weights.

Features of Mendeleev’s Periodic Table

● Twelve horizontal rows, which were condensed to 7, known as periods.

● Eight vertical columns known as groups.

● Groups I to VII subdivided into A and B subgroups.

● Group VIII doesn’t have any subgroups and contains three elements in each row.

● Elements in the same group exhibit similar properties.

Achievements of Mendeleev’s Periodic Table

1. A systematic study of elements: Elements with similar properties were grouped together, that made the study of their chemical and physical properties easier.

2. Correction of atomic masses: Placement of elements in Mendeleev’s periodic table helped in correcting the atomic masses of certain elements. For example, the atomic mass of beryllium was corrected from 13.5 to 9. Similarly, atomic masses of indium, gold, platinum etc., were also corrected.

Limitations of Mendeleev’s Periodic Table

1. Position of hydrogen: Hydrogen resembles both, the alkali metals (IA) and the halogens (VIIA) in properties, so, Mendeleev could not justify its position.

2. Position of isotopes: Atomic weight of isotopes differ, but, they were not placed in different positions in Mendeleev’s periodic table.

3. Anomalous pairs of elements: Cobalt (Co) has higher atomic weights but was placed before Nickel (Ni) in the periodic table.

4. Placement of like elements in different groups: Platinum (Pt) and Gold (Au) have similar properties but were placed in different groups.

5. Cause of periodicity: He could not explain the cause of periodicity among the elements.

Periodicity of Properties: The repetition of properties of elements after certain regular intervals is known as Periodicity of Properties.

Merits of Mendeleev’s Periodic Table

1.Mendeleev’s left vacant places in his table which provided an idea for the discovery of new elements. Example: Eka-boron, Eka-aluminium and Eka-silicon.

2.Mendeleev’s periodic table was predicted properties of several undiscovered elements on the basis of their position in Mendeleev’s periodic table.

3.It is useful in correcting the doubtful atomic masses of some elements.

4.Noble gases could accommodate in the Mendeleev’s periodic table without disturbing the periodic table after discovery.

Limitations of Mendeleev’s Periodic Table

(a) No fixed position for hydrogen: No correct position of the hydrogen atom was in Mendeleev’s periodic table.

Example: Position of hydrogen with alkali metals and halogens (17th group).

(b) No place for isotopes: Position of isotopes were not decided.

Example: Cl-35 and Cl-37.

(c) No regular trend in atomic mass: Position of some elements with lower atomic masses before with higher atomic mass.

Example: Ni-58.7 before Co-58.9.

Mendeleev’s original periodic table is reproduced in the table below

2. Mendeleev's Periodic Table

 Mendeleev’s Periodic Table

Mendeleev arranged 63 elements known at that time in the periodic table. According to Mendeleev “ the properties of the elements are a periodic function of their atomic masses the table consists of eight vertical coloumn called groups and horizontal rows called periods. 

Achievements :- 

(I) The arrangement of elements in group and periods made the study of elements quite systematic in the sence that if properties of one element in a particular group are known, those of the other can be easily .

(II) many gaps were left in  this table for undiscovered elements however, properties of these elements could be predicted in advance from their expected position. This helped in the discovery of these elements the lements silicon , gallium and germanium were ddiscovered in this manner.

(III) Mendeleev   corrected the atomic masses of certain elements with the help of their excepted position and properties. 

(Iv) When  inert  gasses were discovered they could be placed in a new group without disturbing the existing order.

 Limitations :-     

  1. He could not assign a correct position of hydrogen in his periodic table 1 as the properties of hydrogen resembles both with alkali metals as well as with halogens.
  2. The atomic masses do not increases a regular manner in going from one elements to the next , so it was not possible to predict how many elements could be discovered between two elements.
  3. The isotopes of some element will be given different position if atomic number is taken as basis which will disturb the symmetry of the periodic table.

3. Modern Periodic Table

The Modern Periodic Table: In 1913, Henry Moseley showed that the atomic number of an element is a more fundamental property than its atomic mass.

Modern Period Law: The physical and chemical properties of elements are the periodic function of their atomic number.

Modern periodic table is based on atomic number of elements. Atomic number (Z) is equal to the number of protons present in the nucleus of an atom of an element. Modern periodic table contains 18 vertical column known as group and seven horizontal rows known as periods. On moving from left to right in a period, the number of valence electrons increases from 1 to 8 in the elements present. On moving from left to right in a period, number of shell remains same. All the elements of a group of the periodic table have the same number of valence electrons.

Trends in Modern Periodic Table: Valency, Atomic size, metallic and non-metallic characters, and Electronegativity.

(i) Valency: The valency of an element is determined by the number of valence electrons present in the outermost shell of its atom (i.e. the combining capacity of an element is known as its valency).

In Period: On moving from left to right in a period, the valency first increases from 1 to 4 and then decreases to zero (0).

In Groups: On moving from top to bottom in a group, the valency remains same because the number of valence electrons remains the same.

Example: Valency of first group elements = 1 Valency of second group elements = 2.

(ii) Atomic size: Atomic size refers to radius of an atom. It is a distance between the centre of the nucleus and the outermost shell of an isolated atom.

In Period : On moving from left to right in a period, atomic size decreases because nuclear charge increases.

Example: Size of second period elements: Li > Be > B > C > N > O > F

Point to know: The atomic size of noble gases in corresponding period is largest

due to presence of fully filled electronic configuration (i.e. complete octet).

In Group: Atomic size increases down the group because new shells are being

added in spite of the increase in nuclear charge.

Example ; Atomic size of first group element : Li < Na < K < Rb < Cs < Fr

Atomic size of 17th group elements : F < Cl < Br < I

(iii) Metallic character: It is the tendency of an atom to lose electrons. In Period: Along the period from left to right, metallic characters decreases because a tendency to lose electron decreases due to the increase in nuclear charge. Example: Metallic character of second period elements: Li > Be > B > C >> N > O > F

In Group: Metallic character, when moving from top to bottom increases because the atomic size and tendency to lose electrons increases.

Example: First group element : Li < Na < K < Rb < Cs

17th group elements: F < Cl < Br < I

(iv) Non-metallic character: It is tendency of an atom to gain electrons.

In Period: Along the period from left to right, non-metallic character increases because tendency to gain electrons increases due to increase in nucleus charge. Example ; Non-metallic character of 2nd period elements : Li < Be < B < C < N < O < F In Group: On moving from top to bottom in a group, non-metallic character decreases because atomic size increases and tendency to gain electrons decreases. Ex. Non-metallic character of 17th period element: F > Cl > Br > I

(v) Chemical Reactivity

In metals: Chemical reactivity of metals increases down the group because tendency to lose electrons increases. Example ; Li < Na < K < Rb < Cs (1st group) In non-metals: Chemical reactivity of non-metals decreases down the group because tendency to gain electrons decreases. Example: F > Cl > Br > I (17th group)

Period: The horizontal rows in the Modern Periodic Table and Mendeleev’s Periodic Table are called periods.

There are 18 groups and 7 (seven) periods in the Modern Periodic Table.

Atomic size: The atomic size may be visualised as the distance between the centre of the nucleus and the outermost shell of an isolated atom.

The trend of atomic size (radius) in moving down a group: Ongoing down in a group of the Periodic Table, the atomic size increases because a new shell of electrons is added to the atoms at every step. There is an increase in distance between the outermost shell electrons and the nucleus of the atom.

The trend of atomic size (radius) in moving from left to right in a period: On moving from left to right along a period, the size of atoms decreases because on moving from left to right, the atomic number of elements increases which means that the number of protons and electrons in the atoms increases. Due to the large positive charge on the nucleus, the electrons are pulled in more closely to the nucleus and the size of the atom decreases.

Characteristics of triads of J.W. Dobereiner.

  • Elements of a triad show similar chemical properties.
  • These elements of a triad show specific trends in their physical properties.
  • The atomic mass of the middle element was roughly the average of the atomic masses of the other two elements.

Example: Atomic mass of Na is 23 in the triad Li, Na and K. This atomic mass is the average of the atomic masses of Li and K which have atomic masses 7 and 39 respectively.

Triads as formed by Dobereiner.

1st Triad

Li – Lithium

Na – Sodium

K – Potassium

2nd Triad

Ca – Calcium

Sr – Strontium

Ba – Barium

3rd Triad

Cl – Chlorine

Br – Bromine

I – Iodine

Mendeleev’s Periodic Law: It states that “the properties of elements are the periodic functions of their atomic masses.” It means the properties of the elements depend on their atomic masses and the elements are given a position in the periodic table on the basis of their increasing atomic masses.

Merits of Mendeleev’s Periodic Table

(i) Mendeleev left a number of gaps in his table to accommodate the new elements which would be discovered later on. So Mendeleev boldly predicted the existence of some more elements. He even predicted the properties of some of these elements and named them as Eka-boron, Eka-aluminium and Eka-silicon respectively. Later on the elements were discovered, for example, gallium replaced Eka-aluminium and it showed properties similar to that of aluminium.

(ii) He gave the proper position to the noble gases which were discovered later on, without disturbing the existing order of elements. He placed them in a new group.

Limitations of Mendeleev’s classification:

  • The position of isotopes could not be explained because isotopes have the same chemical properties but different atomic masses. If the elements are arranged according to atomic masses, the isotopes should be placed in different groups of the Periodic Table.
  • The atomic masses do not increase in a regular manner in going from one element to the next.
  • He could not assign a correct position to hydrogen in his table because hydrogen has some properties similar to alkali metals and some properties similar to halogens.

Modem Periodic Law: This law was proposed by Henry Moseley, a scientist in 1913. According to this Law, “Properties of elements are the periodic function of their atomic number.” It means that the properties of elements depend on their atomic number and the elements are given positions in the periodic table on the basis of their increasing atomic number. As atomic number determines the distribution of electrons in the orbits, and electrons of the outermost orbit determine the properties of an element.

Groups and periods in the Modem (long form) Periodic Table: There are 18 groups (vertical columns) and 7 periods (horizontal lines) in the Modern (or long form) Periodic Table. The number of the period is equal to the number of shells in the atoms of the elements belonging to that period.

Trends in Mendeleev’s Periodic Table

The properties of elements are periodic functions of their atomic mass.

It has 8 groups.

  • No place could be assigned to isotopes of an element.
  • There were three gaps left by Mendeleev in his Periodic Table.
  • No fixed position was given to hydrogen in this Periodic Table.
  • No distinction was made between metals and non-metals.
  • Transition elements are placed together in Group VIII.
  • Inert gases were not known at the time of Mendeleev.

Trends in Modem Periodic Table

(i) Valency: Elements belonging to the same group have the same number of valence electrons and thus the same valency. Valency in a particular period from left to right first increases as positive valency and then decreases as negative valency.

Example: In elements of 2nd period:

Li has 1+ valency, then Be2+, B3+, C4+ covalency, N3- valency, then O2- and F(-) valency.

(ii) The atomic size or atomic radius increases: as we move down in a group and it decreases as we move from left to right in a period. Atomic size increases down a group due to the increase in the number of shells. Atomic size decreases along a period due to an increase in the nuclear charge which tends to pull the electrons closer to the nucleus and reduces the size of the atom.

Non-metallic characters increase from left to right in a period due to increase in the electronegativity and these characters decrease from top to bottom in a group due to the decrease in the electronegativity of atoms while going down in a group.

1. Need for classification of elements:

Increase in the discovery of different elements made it difficult to organise all that was known about the elements. To study a large number of elements with ease, various attempts were made. The attempts resulted in the classification of elements into metals and non-metals.

2. Dobereiner’s triads:

Johann Wolfgang Dobereiner, a German chemist, classified the known elements in groups of three elements on the basis of similarities in their properties. These groups were called triads.

(i) Characteristics of Triads:

Properties of elements in each triad were similar.

Atomic mass of the middle element was roughly the average of the atomic masses of the other two elements.

3. Newlands’ Law of Octaves:

John Newlands’, an English scientist, arranged the known elements in the order of increasing atomic masses and called it the ‘Law of Octaves’. It is known as ‘Newlands’ Law of Octaves’.

(i) Characteristics of Newlands’ Law of Octaves:

It contained the elements from hydrogen to thorium.

Properties of every eighth element were similar to that of the first element.

(ii) Table showing Newlands’ Octaves:

(iii) Limitations of Newlands’ law of Octaves:

The law was applicable to elements up to calcium (Ca).

It contained only 56 elements.

In order to fit elements into the table, Newlands’ adjusted two elements like cobalt and nickel in die the same slot and also put some unlike elements under the same note.

4. Mendeleev’s Periodic Table: Dmitri Ivanovich – 5 ’ Mendeleev, a Russian demist, was the most important contributor to the early development of a periodic table of elements wherein the elements were arranged on the basis of their atomic mass and chemical properties.

Characteristics of Mendeleev’s Periodic Table:

  • Mendeleev arranged all the 63 known elements in increasing order of their atomic masses.
  • Mendeleev’s Periodic Law: The properties of elements are the periodic function of their atomic masses.

Achievements of Mendeleev’s Periodic Table:

Mendeleev adjusted few elements with a slightly greater atomic mass before the elements with slightly lower atomic mass, so that elements with similar properties could be grouped together. For example, aluminium appeared before silicon, cobalt appeared before nickel.

Mendeleev left some gaps in his periodic table.

He predicted the existence of some elements that had not been discovered at that time. His predictions were quite true as elements like scandium, gallium and germanium were discovered later.

The gases like helium, neon and argon, which were discovered later, were placed in a new group without disturbing the existing order.


No fixed positions were given to hydrogen in the Mendeleev’s periodic table.

Positions of Isotopes of all elements was not certain according to Mendeleev’s periodic table.

Atomic masses did not increase in a regular manner in going from one element to the next.

5. Modem Periodic Table: Henry Moseley, gave a new ! property of elements, ‘atomic number’ and this was I adopted as the basis of Modem Periodic Table.

(i) Modem Periodic Law: Properties of elements are a periodic function of i their atomic number.

(ii) The position of elements in Modem Periodic Table:

The modem periodic table consists of 18 groups and 7 periods.

  • Elements present in any one group have the same number of valence electrons. Also, the number of shells increases as we go down the group.
  • Elements present in any one period, contain the same number of shells. Also, with increase in atomic number by one unit on moving from left to right, the valence shell electrons increases by one unit.
  • Each period marks a new electronic shell getting filled.

Trends in the Modern Periodic Table:

Valency:  Valency of an element is determined by the number of valence electrons present in the outermost shell of its atom.

Valency of elements in a particular group is same.

Atomic Size: Atomic size refers to the radius of an atom.In a period, atomic size and radii decreases from left to right.In a group, atomic size and radii increases from top to bottom.

3. Modern Periodic Table

Modern Periodic Table

This law was given by Henry Moseley in 1913 . it states properties of the elements are the periodic function of their atomic numbers.

Periodicity may be defined as the repetition of the similar properties the elements placed in a group and separated placed in a group and seprated by certain definite gap of atomic numbers the cause of periodicity is the resemblance in properties of the elements is the repetition of the same valence shell electronic configuration

  • Mosely proposed this modern periodic table according to which “ the physical and chemical properties of elements are periodic function of their atomic number and the 7 horizontal rows are called periods and 18 vertical columns known as groups.
  • The elements belonging  to a particular group make a family and usually named after the first member .
  • In a group all elements contain the same number if valence electrons.
  • In a period , all elements contain the same number of shells , but as we move from left to right , the number of valence shell electrons increases by one unit.
  • The maximum number of electrons that can be accommodated in a shell can be calculated by the formula 2n2 , where n is the formula of the given shell from the nucleus.

For example

K shell        → 2(1)2 =  2, hence the first periodic has 2 elements .  

L shell        → 2(2)2  = 8, hence the second period has 8 elements.     

The third, fourth , fifth, sixth, and seventh periods have 8,18, 18, 32, and 32 respectively.

(i)    Trends in modern periodic table  

Some trends were observed of the elements in moving down the group ( from top to bottom of the table ) and across a period ( from left to right ) are as follows:-

Valency :-  the valancy of an elements is determined by the number if valence electrons present in the outermost shell of its atom (i.e 8 electrons in valence shell, in some special cases it is 2 electrons.

(ii) Atomic size :- the term atomic size refers to the radius of an atom the atomic radius of hydron atom is 37 pm ( picometre , 1=10-12m)

  • The atomic radius decreases on moving from left to right along a period. This is due to an increase  nuclear charge which tends to pull the electrons closer to the nucleus and reduces the size of the atom.
  •  In a group atomic decreases  from top  to buttom due to increase in number of shells.

(iii)   Metallic and non- metallic properties

  • The metals like Na and Mg are towards the left hand side of the periodic table. The non- metals like sulphur and chlorine are formed on the right hand side. In middle , silicon, is classified as a semi- metal or metalloid ( which exhibits some properties of both metals and non-  metals )
  • In a eriod from left to right metallic Character  increases.
  •  In a group , metallic character increases from top to bottom while non- metallic character decrease.

(iV)  Electronegativity :-

The relative tendency of an atom to attract the shared electrons pair of electrons towards itself is called electronegativity.

In a period left to right the value to right the value of electro negativity increases while in a group from top to buttom, the value of electronegativity decreases.