Chapter 1: Chemical Reactions and Equations | CBSE Class 10 | NCERT Notes

CHEMICAL REACTIONS & EQUATIONS

Chemical reaction is a process in which one or more reactants are converted to one or more products.

During a chemical reaction, chemical change occurs.

The substances that undergo chemical change are called reactants. The new substance formed is called product.

EXAMPLES FOR CHEMICAL REACTIONS

• Burning of a clean magnesium ribbon with a dazzling white flame to form white powder (magnesium oxide). It is due to the reaction of magnesium with oxygen in air.

• Take lead nitrate solution in a test tube. Add potassium iodide solution. A yellow precipitate of lead iodide appears at the bottom.
• Take few zinc granules in a conical flask or test tube. Add dilute HCl or H₂SO₄. Bubbles are observed around zinc granules due to release of hydrogen. Conical flask becomes hot.

The following observations helps to determine whether a chemical reaction has taken place.

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

CHEMICAL EQUATIONS

The description of a chemical reaction can be written in a shorter form. The simplest way is word-equation. E.g.

Magnesium + Oxygen → Magnesium oxide

                                                        (Reactants)                      (Product)

The reactants are written on left-hand side (LHS) with a plus sign (+) between them. Products are written on the right-hand side (RHS) with a plus sign (+) between them.

The arrowhead points towards the products, and shows the direction of the reaction.

WRITING A CHEMICAL EQUATION

Chemical equations can be simplified by using chemical formulae. A chemical equation represents a chemical reaction. E.g.

Mg + O₂ → MgO (skeletal chemical equation)

If the number of atoms of each element is same on LHS & RHS, the equation is balanced. If not, it is unbalanced. It is called a skeletal chemical equation.

BALANCED CHEMICAL EQUATIONS

According to the law of conservation of mass, mass can neither be created nor destroyed in a chemical reaction. i.e., total mass of the elements present in the products is equal to the total mass of the elements in the reactants.

Number of atoms of each element remains the same before and after a chemical reaction. So, skeletal chemical equation must be balanced.

E.g. Word-equation of a chemical reaction is given:

Zinc + Sulphuric acid → Zinc sulphate + Hydrogen

It is represented by the following chemical equation:

Zn + H₂SO₄ → ZnSO₄ + H₂

Element	Number of atoms in reactants (LHS)	Number of atoms in products (RHS)
Zn	1	1
H	2	2
S	1	1
O	4	4

Thus it is a balanced chemical equation.

STEPS OF BALANCING A CHEMICAL EQUATION

Balancing a chemical equation using least whole number coefficient is called hit-and-trial method.

The steps are given below:

Fe + H₂O → Fe₃O₄ + H₂

Step I: Draw boxes around each formula. Do not change anything inside the boxes.

Fe + H₂O → Fe₃O₄ + H₂

Step II: List the number of atoms of different elements.

Element	Number of atoms in reactants (LHS)	Number of atoms in products (RHS)
Fe	1	3
H	2	2
O	1	4

Step III: Select the compound (reactant or product) having maximum number of atoms (Fe₃O₄). In that, select the element having maximum number of atoms (oxygen).

Atoms of oxygen	In reactants	In products
(i)   Initial (ii) To balance	1 (in H2O) 1 x 4	4 (in Fe3O4) 4

Fe + 4 H₂O → Fe₃O₄ + H₂

Step IV: Balance the number of hydrogen atoms.

Atoms of hydrogen	In reactants	In products
(i)   Initial (ii) To balance	8 (in 4 H2O) 8	2 (in H2) 2 x 4

Fe + 4 H₂O → Fe₃O₄ + 4 H₂ (partly balanced)

Step V: Balance the number of iron atoms.

Atoms of iron	In reactants	In products
(i) Initial (ii) To balance	1 (in Fe) 1×3	3 (in Fe3O4) 3

3 Fe + 4 H₂O → Fe₃O₄ + 4 H₂

Step VI: Count atoms of each element on both sides of the equation to check the correctness.

3Fe + 4H₂O → Fe₃O₄ + 4H₂ (balanced equation)

Step VII: If necessary, physical states such as gaseous (g), liquid (l), aqueous (aq) and solid (s) states are included in a chemical equation. Aqueous (aq) means the reactant or product is present as a solution in water.

3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

H₂O (g) indicates that water is used in the form of steam.

Sometimes, reaction conditions (temperature, pressure, catalyst, etc.) are indicated above and/or below the arrow in the equation. E.g.

TYPES OF CHEMICAL REACTIONS

TYPES OF CHEMICAL REACTIONS

In a chemical reaction, atoms of one element do not change into atoms of another element. Nor do atoms disappear from the mixture or appear from elsewhere.

Chemical reactions involve breaking and making of bonds between atoms.

1. COMBINATION REACTIONS

These are the reactions in which a product is formed by combining two or more reactants.

E.g. Take some calcium oxide (quick lime) in a beaker. Slowly add water to this. Beaker becomes hot. Calcium oxide reacts vigorously with water to produce slaked lime (calcium hydroxide) releasing heat.

CaO (s)     +     H₂O (l)    →     Ca(OH)₂(aq)   +   Heat

(Quick lime)                                  (Slaked lime)

 

Formation of slaked lime

Slaked lime solution is used for whitewashing. Calcium hydroxide reacts slowly with CO₂ in air to form a thin layer of calcium carbonate (CaCO₃) after 2-3 days of whitewashing giving a shiny finish to the walls. Marble is also CaCO₃.

Other examples of combination reactions:

(i) Burning of coal:

C(s) + O₂(g) → CO₂(g)

(ii) Formation of water from H₂ and O₂:

2H₂(g) + O₂(g) → 2H₂O(l)

Reactions in which heat is released along with formation of products are called exothermic chemical reactions.

Other examples of exothermic reactions:

(i)       Burning of natural gas:

CH₄(g) + 2O₂ (g) → CO₂ (g) + 2H₂O (g)

(ii)     Respiration:

C₆H₁₂O₆(aq) + 6O₂(aq) → 6CO₂(aq) + 6H₂O(l) + energy

(Glucose)

(iii)    Decomposition of vegetable matter into compost.

(iv)    Burning magnesium to form magnesium oxide.

2. DECOMPOSITION REACTIONS

These are the reactions in which single reactant breaks down to give simpler products.

Examples:

• Take 2 g ferrous sulphate crystals in a dry boiling tube. Heat it over flame. Green colour of ferrous sulphate is changed with characteristic odour of burning sulphur. When heated, ferrous sulphate (FeSO₄. 7H₂O) lose water. It then decomposes to ferric oxide (Fe₂O₃), sulphur dioxide (SO₂) and sulphur trioxide (SO₃).

• Decomposition of CaCO₃ to calcium oxide (CaO) & CO₂ on heating is used in industries. CaO has many uses (e.g. manufacture of cement). When a decomposition reaction is carried out by heating, it is called thermal decomposition.

Other examples of thermal decomposition:

Example 1

Take 2 g lead nitrate powder in a boiling tube. 

Hold the boiling tube and heat it over a flame. 

Brown fumes of nitrogen dioxide (NO₂) are emitted.

Example 2

At the base of a plastic mug, drill two holes and fit rubber stoppers. Insert carbon electrodes in rubber stoppers. Connect electrodes to a 6 volt battery.

Fill the mug with water such that the electrodes are immersed. Add few drops of dil. H₂SO₄ to the water.

Take two test tubes filled with water and invert them over the electrodes. Switch on the current. Bubbles are formed at electrodes by displacing water in test tubes.

At cathode (-ve electrode), hydrogen gas is collected. At anode (+ve electrode), oxygen is collected.

At cathode, double amount of gas is collected as compared to anode because during the break down of water, 2H molecule is released with 1 oxygen molecule.

When we bring a burning candle to the gas at cathode, it burns immediately. But gas at anode does not burn.

Example 3

Take 2 g silver chloride in a china dish. 

Place this in sunlight for some time. 

Its white colour turns grey. This is due to the decomposition of silver chloride into silver and chlorine by light.

2AgCl (s)  → 2Ag (s) + Cl₂ (g)

Silver bromide also behaves in the same way.

2AgBr (s)  → 2Ag (s) + Br₂ (g)

Decomposition reactions of silver chloride and silver bromide are used in black & white photography.

Decomposition reactions require energy (heat, light or electricity) to break down the reactants.

Reactions in which energy is absorbed are called endothermic reactions. E.g.

Take 2 g barium hydroxide in a test tube. Mix 1 g of ammonium chloride. The following reaction occurs.

Ba(OH)₂ +NH₄Cl → BaCl₂ + NH₄OH

It is an endothermic reaction because it absorbs heat from environment. So bottom of the test tube becomes cool.

3. DISPLACEMENT REACTIONS

These are the reactions in which a more reactive element displaces a less reactive element from its compound.

Examples:

Take 3 iron nails cleaned by rubbing with sand paper.

Take two test tubes A and B. Add 10 mL copper sulphate (CuSO₄) solution in each test tube.

Immerse two iron nails in CuSO₄ in test tube B. Keep one iron nail aside for comparison.

After 20 minutes, take out the iron nails. They become brownish in colour. Blue colour of CuSO₄ solution fades.

This is due to the following chemical reaction:

Here, iron has been displaced by copper from the CuSO₄ solution.

Other examples of displacement reactions:

Zinc and lead are more reactive elements than copper. They displace copper from its compounds.

4. DOUBLE DISPLACEMENT REACTIONS

These are the reactions in which there is an exchange of ions between the reactants.

Examples:

Take 3 mL sodium sulphate solution in a test tube. In another test tube, take 3 mL barium chloride solution. Mix the two solutions.

A white substance (BaSO₄) is formed by the reaction of SO₄^2– and Ba²⁺. This water insoluble substance is called precipitate. Any reaction that produces a precipitate is called precipitation reaction.

The other product (sodium chloride) remains in solution.

Reaction between lead nitrate & potassium iodide to form a yellow precipitate of lead iodide.

Pb(NO₃)₂(aq)+ 2KI(aq) → PbI₂(s) + 2KNO₃(aq)

5. OXIDATION AND REDUCTION

If a substance gains oxygen during a reaction, it is said to be oxidised. Such reaction is called oxidation.

If a substance loses oxygen during a reaction, it is said to be reduced. Such reaction is called reduction.

Examples:

Heat a china dish containing about 1 g copper powder.

The surface of copper powder becomes coated with black copper(II) oxide. This is because oxygen is added to copper and copper oxide is formed.

2Cu + O₂   → 2CuO

If hydrogen gas is passed over the heated CuO, the black coating turns brown as the reverse reaction takes place and copper is obtained.

CuO + H₂   →  Cu + H₂O

During this reaction, copper(II) oxide loses oxygen and is reduced. The hydrogen gains oxygen and is oxidised.

When one reactant gets oxidised and the other gets reduced in a reaction, it is called oxidation-reduction reaction (redox reaction).

Other examples of redox reactions:

Carbon is oxidised to CO and ZnO is reduced to Zn.

ZnO + C → + Zn + CO

HCl is oxidised to Cl₂ whereas MnO₂ is reduced to MnCl₂.

MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂

If a substance gains oxygen or loses hydrogen during a reaction, it is oxidised. If a substance loses oxygen or gains hydrogen during a reaction, it is reduced.

A magnesium ribbon burns in air (oxygen) to form magnesium oxide. Here magnesium is oxidised.

EFFECTS OF OXIDATION REACTIONS IN EVERYDAY LIFE

1. CORROSION

It is a process by which a metal is attacked by substances such as moisture, acids, etc. E.g. black coating on silver, green coating on copper.

Iron articles get coated with a reddish-brown powder when left for some time. This is called rusting of iron.

Corrosion causes damage to car bodies, bridges, iron railings, ships and to all objects made of metals.

2. RANCIDITY

When fats and oils are oxidised, they become rancid and their smell and taste change.

Oxidation and rancidity can be prevented by:

• Adding antioxidants (substances which prevent oxidation) to foods containing fats and oil.
• Keeping food in air tight containers.
• Flushing bags of chips with gas such as nitrogen.