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Chemical Reactions and Equations

In our daily lives, we observe many changes where the original substance loses its nature and identity. For instance, milk turning sour in summer, iron rusting in humid air, grapes fermenting, food being cooked or digested, and even the process of respiration are all examples of chemical changes. Whenever a chemical change happens, we can say that a chemical reaction has taken place.

But how do we know for sure that a chemical reaction has occurred? Certain observable changes help us identify them.

Indicators of a Chemical Reaction: A chemical reaction can often be identified by one or more of the following observations:

Change in state: The physical state of the substance might change (e.g., solid to liquid or gas).
Change in colour: There might be a distinct change in the colour of the substances involved. For example, when lead nitrate solution is mixed with potassium iodide solution, a yellow precipitate is formed.
Evolution of a gas: Bubbles of gas may be seen forming. For instance, when zinc granules react with dilute sulphuric acid, hydrogen gas is evolved.
Change in temperature: The temperature of the mixture might increase (exothermic reaction) or decrease (endothermic reaction). The reaction between zinc and dilute acid, for example, causes the flask to become warm.

A classic example is burning a magnesium ribbon. The ribbon, a solid metal, burns with a dazzling white flame and turns into a white powder called magnesium oxide. This involves a change in state and the release of energy as light and heat.

CHEMICAL EQUATIONS

Describing a chemical reaction in a full sentence can be lengthy. A shorter, more precise way is to use a word-equation.

For the reaction of burning magnesium in air, the word-equation is: $\underset{\text { (Reactants) }}{\text { Magnesium }+ \text { Oxygen }} \rightarrow \underset{\text { (Product) }}{\text { Magnesium oxide }}$

Reactants are the substances that undergo a chemical change. They are written on the left-hand side (LHS).
Products are the new substances formed during the reaction. They are written on the right-hand side (RHS).
An arrow ( $\rightarrow$ ) points from the reactants to the products, indicating the direction of the reaction.

Writing a Chemical Equation

To make the representation even more concise and useful, we use chemical formulae instead of words. This symbolic representation of a chemical reaction is called a chemical equation.

The word-equation for burning magnesium can be written as: $\mathrm{Mg}+\mathrm{O}_{2} \rightarrow \mathrm{MgO}$

This type of equation, where the number of atoms of each element is not equal on both sides, is called a skeletal chemical equation. It is an unbalanced equation.

Balanced Chemical Equations

According to the law of conservation of mass, mass can neither be created nor destroyed in a chemical reaction. This means the total mass of the reactants must equal the total mass of the products. For this to be true, the number of atoms of each element must remain the same before and after the reaction.

Therefore, every chemical equation must be balanced.

Let's look at the reaction of zinc with sulphuric acid: $\mathrm{Zn}+\mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow \mathrm{ZnSO}_{4}+\mathrm{H}_{2}$

If we count the atoms of each element on both sides:

Zn: 1 on LHS, 1 on RHS (Balanced)
H: 2 on LHS, 2 on RHS (Balanced)
S: 1 on LHS, 1 on RHS (Balanced)
O: 4 on LHS, 4 on RHS (Balanced)

Since the number of atoms of each element is the same on both sides, this is a balanced chemical equation.

Steps to Balance a Chemical Equation (Hit-and-Trial Method)

Let's balance the skeletal equation for the reaction of iron with steam: $\mathrm{Fe}+\mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{Fe}_{3} \mathrm{O}_{4}+\mathrm{H}_{2}$

List the number of atoms of each element on both the reactant (LHS) and product (RHS) sides.
- Fe: 1 on LHS, 3 on RHS
- H: 2 on LHS, 2 on RHS
- O: 1 on LHS, 4 on RHS
Start with the compound that has the most atoms. In this case, it's $\mathrm{Fe}_{3} \mathrm{O}_{4}$ . Within this compound, start with the element that has the most atoms, which is oxygen (O).
- There are 4 oxygen atoms on the RHS and only 1 on the LHS. To balance oxygen, place a coefficient '4' in front of $\mathrm{H}_{2} \mathrm{O}$ on the LHS.
- Do not change the chemical formula (e.g., you cannot change $\mathrm{H}_{2} \mathrm{O}$ to $\mathrm{H}_{2} \mathrm{O}_{4}$ ). $\mathrm{Fe}+4 \mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{Fe}_{3} \mathrm{O}_{4}+\mathrm{H}_{2}$
Balance the other elements. Now, the number of hydrogen atoms on the LHS is $4 \times 2 = 8$ . On the RHS, there are only 2. To balance hydrogen, place a coefficient '4' in front of $\mathrm{H}_{2}$ on the RHS. $\mathrm{Fe}+4 \mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{Fe}_{3} \mathrm{O}_{4}+4 \mathrm{H}_{2}$
Balance the remaining element. Finally, balance the iron (Fe) atoms. There is 1 Fe atom on the LHS and 3 on the RHS. Place a coefficient '3' in front of Fe on the LHS. $3 \mathrm{Fe}+4 \mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{Fe}_{3} \mathrm{O}_{4}+4 \mathrm{H}_{2}$
Check the final equation. Count the atoms of each element on both sides to verify that the equation is balanced.
- Fe: 3 on LHS, 3 on RHS
- H: 8 on LHS, 8 on RHS
- O: 4 on LHS, 4 on RHS The equation is now balanced.

Making Equations More Informative

To provide more information, we can include the physical states of the reactants and products.

(s) for solid
(l) for liquid
(g) for gas
(aq) for aqueous (dissolved in water)

The balanced equation for iron and steam becomes: $3 \mathrm{Fe}(\mathrm{~s})+4 \mathrm{H}_{2} \mathrm{O}(\mathrm{~g}) \rightarrow \mathrm{Fe}_{3} \mathrm{O}_{4}(\mathrm{~s})+4 \mathrm{H}_{2}(\mathrm{~g})$ [!note] The symbol (g) is used for $\mathrm{H}_{2} \mathrm{O}$ to show it is in the form of steam, not liquid water.

Reaction conditions like temperature, pressure, or catalysts can be written above or below the arrow.

$\mathrm{CO}(\mathrm{~g})+2 \mathrm{H}_{2}(\mathrm{~g}) \xrightarrow{340 \mathrm{~atm}} \mathrm{CH}_{3} \mathrm{OH}(\mathrm{l})$
$6 \mathrm{CO}_{2}(\mathrm{aq})+12 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \xrightarrow[\text { Chlorophyll }]{\text { Sunlight }} \mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}(\mathrm{aq})+6 \mathrm{O}_{2}(\mathrm{aq})+6 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})$

TYPES OF CHEMICAL REACTIONS

Chemical reactions involve the breaking of old chemical bonds and the making of new ones to form new substances. Reactions can be classified into several types based on the changes occurring.

Combination Reaction

A combination reaction is a reaction in which two or more reactants combine to form a single product.

Example

When calcium oxide (quick lime) is mixed with water, it forms calcium hydroxide (slaked lime).

\underset{\text { (Quick lime) }}{\mathrm{CaO}(\mathrm{~s})}+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \underset{\text { (Slaked lime) }}{\mathrm{Ca}(\mathrm{OH})_{2}(\mathrm{aq})}

This reaction releases a large amount of heat, making the mixture warm.

Other Examples of Combination Reactions:

Burning of coal: $\mathrm{C}(\mathrm{~s})+\mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{CO}_{2}(\mathrm{~g})$
Formation of water: $2 \mathrm{H}_{2}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g}) \rightarrow 2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})$

Reactions that release heat along with the formation of products are called exothermic chemical reactions. The reaction of quick lime with water is an exothermic reaction.

Examples of Exothermic Reactions:

Burning of natural gas (methane): $\mathrm{CH}_{4}(\mathrm{~g})+2 \mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{CO}_{2}(\mathrm{~g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{~g})$
Respiration: During digestion, carbohydrates are broken down into glucose. This glucose reacts with oxygen in our cells to provide energy. $\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}(\mathrm{aq})+6 \mathrm{O}_{2}(\mathrm{aq}) \rightarrow 6 \mathrm{CO}_{2}(\mathrm{aq})+6 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})+\text { energy }$
Decomposition of vegetable matter into compost.

Decomposition Reaction

A decomposition reaction is the opposite of a combination reaction. It is a reaction in which a single compound breaks down into two or more simpler substances. These reactions require an input of energy in the form of heat, light, or electricity.

1. Thermal Decomposition (using heat)

When a decomposition reaction is carried out by heating, it is called thermal decomposition.

Example

Heating ferrous sulphate crystals (green) causes them to lose water and then decompose into ferric oxide (a brown solid), sulphur dioxide, and sulphur trioxide gas.

2 \mathrm{FeSO}_{4}(\mathrm{~s}) \xrightarrow{\text { Heat }} \mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{~s})+\mathrm{SO}_{2}(\mathrm{~g})+\mathrm{SO}_{3}(\mathrm{~g})

Another important example is the decomposition of limestone (calcium carbonate) into quick lime (calcium oxide) and carbon dioxide, which is used in manufacturing cement. $\underset{(\text { Limestone })}{\mathrm{CaCO}_{3}(\mathrm{~s})} \quad \xrightarrow{\text { Heat }} \underset{\text { (Quick lime) }}{\mathrm{CaO}(\mathrm{~s})}+\mathrm{CO}_{2}(\mathrm{~g})$

2. Electrolytic Decomposition (using electricity)

This type of decomposition occurs when an electric current is passed through a substance.

Example

The electrolysis of water breaks it down into hydrogen gas and oxygen gas.

2H_2O(l) \xrightarrow{\text{Electrolysis}} 2H_2(g) + O_2(g)

The volume of hydrogen gas collected is double the volume of oxygen gas because a water molecule (

H_2O

) contains two atoms of hydrogen for every one atom of oxygen.

3. Photolytic Decomposition (using light)

This decomposition is caused by light energy.

Example

When white silver chloride is exposed to sunlight, it decomposes into grey silver and chlorine gas.

2 \mathrm{AgCl}(\mathrm{~s}) \xrightarrow{\text { Sunlight }} 2 \mathrm{Ag}(\mathrm{~s})+\mathrm{Cl}_{2}(\mathrm{~g})

Similarly, silver bromide also decomposes in sunlight. These reactions are the basis for black and white photography.

2 \mathrm{AgBr}(\mathrm{~s}) \xrightarrow{\text { Sunlight }} 2 \mathrm{Ag}(\mathrm{~s})+\mathrm{Br}_{2}(\mathrm{~g})

Reactions that absorb energy are known as endothermic reactions. All decomposition reactions are endothermic because they require energy to break bonds.

Displacement Reaction

A displacement reaction is a reaction in which a more reactive element displaces (removes) a less reactive element from its salt solution.

Example

If an iron nail is placed in a blue copper sulphate solution, the iron, being more reactive than copper, displaces the copper.

\mathrm{Fe}(\mathrm{~s})+\mathrm{CuSO}_{4}(\mathrm{aq}) \rightarrow \mathrm{FeSO}_{4}(\mathrm{aq})+\mathrm{Cu}(\mathrm{~s})

Observations:

The blue colour of the copper sulphate solution fades and turns light green due to the formation of iron sulphate.
A reddish-brown coating of copper metal is deposited on the iron nail.

Other Examples of Displacement Reactions:

Zinc displaces copper from copper sulphate: $\mathrm{Zn}(\mathrm{~s})+\mathrm{CuSO}_{4}(\mathrm{aq}) \rightarrow \mathrm{ZnSO}_{4}(\mathrm{aq})+\mathrm{Cu}(\mathrm{~s})$
Lead displaces copper from copper chloride: $\mathrm{Pb}(\mathrm{~s})+\mathrm{CuCl}_{2}(\mathrm{aq}) \rightarrow \mathrm{PbCl}_{2}(\mathrm{aq})+\mathrm{Cu}(\mathrm{~s})$

Double Displacement Reaction

A double displacement reaction is a reaction in which there is an exchange of ions between the reactants. These reactions usually occur in solution and often result in the formation of a precipitate.

Example

When a solution of sodium sulphate is mixed with a solution of barium chloride, an exchange of ions (

\mathrm{Na}^{+}

with

\mathrm{Ba}^{2+}

and

\mathrm{SO}_{4}^{2-}

with

\mathrm{Cl}^{-}

) occurs.

\mathrm{Na}_{2} \mathrm{SO}_{4}(\mathrm{aq})+\mathrm{BaCl}_{2}(\mathrm{aq}) \rightarrow \mathrm{BaSO}_{4}(\mathrm{~s})+2 \mathrm{NaCl}(\mathrm{aq})

A white, insoluble substance called barium sulphate is formed. This insoluble substance is known as a precipitate. Any reaction that produces a precipitate is called a precipitation reaction.

Oxidation and Reduction

Oxidation is defined as the gain of oxygen or the loss of hydrogen by a substance. Reduction is defined as the loss of oxygen or the gain of hydrogen by a substance.

These two processes always occur simultaneously in a reaction. Such reactions are called oxidation-reduction reactions or redox reactions.

Example

When copper powder is heated in air, it gains oxygen and forms black copper(II) oxide. Here, copper is oxidised.

2 \mathrm{Cu}+\mathrm{O}_{2} \xrightarrow{\text { Heat }} 2 \mathrm{CuO}

If hydrogen gas is passed over this heated copper(II) oxide, the reverse happens. The copper(II) oxide loses oxygen and becomes copper again. The hydrogen gains this oxygen to form water. $\mathrm{CuO}+\mathrm{H}_{2} \xrightarrow{\text { Heat }} \mathrm{Cu}+\mathrm{H}_{2} \mathrm{O}$ In this reaction:

Copper oxide ( $\mathrm{CuO}$ ) is reduced to copper ( $\mathrm{Cu}$ ) because it loses oxygen.
Hydrogen ( $\mathrm{H}_{2}$ ) is oxidised to water ( $\mathrm{H}_{2} \mathrm{O}$ ) because it gains oxygen.

Other Examples of Redox Reactions:

$\mathrm{ZnO}+\mathrm{C} \rightarrow \mathrm{Zn}+\mathrm{CO}$ $ZnO + C \to Zn + CO$
- $\mathrm{ZnO}$ is reduced to $\mathrm{Zn}$ .
- $\mathrm{C}$ is oxidised to $\mathrm{CO}$ .
$\mathrm{MnO}_{2}+4 \mathrm{HCl} \rightarrow \mathrm{MnCl}_{2}+2 \mathrm{H}_{2} \mathrm{O}+\mathrm{Cl}_{2}$ $MnO_{2} + 4 HCl \to MnCl_{2} + 2 H_{2} O + Cl_{2}$
- $\mathrm{MnO}_{2}$ is reduced to $\mathrm{MnCl}_{2}$ .
- $\mathrm{HCl}$ is oxidised to $\mathrm{Cl}_{2}$ (as it loses hydrogen).

Chemical Reactions and EquationsClass 10 Chemistry Notes

Chemical Reactions and Equations

CHEMICAL EQUATIONS

Writing a Chemical Equation

Balanced Chemical Equations

TYPES OF CHEMICAL REACTIONS

Combination Reaction

Decomposition Reaction

Displacement Reaction

Double Displacement Reaction

Oxidation and Reduction

EFFECTS OF OXIDATION REACTIONS IN EVERYDAY LIFE

Corrosion

Rancidity