Chemical Reactions and Equations — Study Notes

1. Introduction

Chemistry is the branch of science that deals with the study of matter, its properties, and the changes it undergoes. When a substance undergoes a change that results in the formation of one or more new substances, we say that a chemical reaction has taken place. Chemical reactions are the fundamental events that drive all transformations in nature, from the rusting of iron gates to the digestion of food in our bodies and the photosynthesis carried out by green plants. Understanding chemical reactions helps us explain how new substances are formed, how energy is released or absorbed, and how matter is conserved throughout these transformations.

A chemical equation is the shorthand representation of a chemical reaction using symbols and formulas of the substances involved. The substances that undergo reaction are called reactants, and they are written on the left-hand side, while the new substances formed are called products, written on the right-hand side. An arrow separates the two sides and indicates the direction in which the reaction proceeds. A chemical equation not only tells us which substances react but also, when balanced, tells us the exact quantitative relationships between the reactants and products.

In this chapter, we will explore the different types of chemical reactions, learn how to write and balance chemical equations, understand the concepts of oxidation and reduction, and study real-life applications such as corrosion and rancidity. By the end of this chapter, you will be able to represent any simple chemical change symbolically and predict the type of reaction taking place.

2. Writing a Chemical Equation

A chemical equation is written by placing the chemical formulas of the reactants on the left and the products on the right, separated by an arrow. For example, when magnesium burns in air to form magnesium oxide, the word equation is written as:

Magnesium + Oxygen → Magnesium oxide

The symbolic equation for the same reaction is written as:

$$\text{Mg} + \text{O}_2 \rightarrow \text{MgO}$$

Balancing a Chemical Equation

According to the law of conservation of mass, matter can neither be created nor destroyed in a chemical reaction. Therefore, the total number of atoms of each element must remain the same on both sides of the equation. An equation in which the number of atoms of each element is equal on both sides is called a balanced chemical equation.

To balance the magnesium oxide equation, we count the atoms on each side. On the left we have one Mg atom and two O atoms (from O2), while on the right we have one Mg and one O. To balance oxygen, we put a coefficient 2 before MgO. Now there are two Mg atoms on the right, so we put a coefficient 2 before Mg on the left.

$$\boxed{2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}}$$

Steps for Balancing

1. Write the skeletal equation with correct formulas of all reactants and products.
2. Count the number of atoms of each element on both sides.
3. Balance the atoms one element at a time by adding coefficients (never change the subscripts in a formula).
4. Check that all elements are balanced and reduce coefficients to the smallest whole-number ratio.

3. Types of Chemical Reactions

Combination Reaction

When two or more substances combine to form a single product, the reaction is called a combination reaction. For example, quicklime (calcium oxide) reacts with water to form slaked lime with the release of a large amount of heat:

$$\text{CaO} + \text{H}_2\text{O} \rightarrow \text{Ca(OH)}_2 + \text{Heat}$$

Decomposition Reaction

A reaction in which a single compound breaks down into two or more simpler substances is called a decomposition reaction. It is the opposite of a combination reaction and usually requires energy in the form of heat, light, or electricity. For example, when ferrous sulphate is heated, it decomposes with a change of colour:

$$2\text{FeSO}_4 \xrightarrow{\text{Heat}} \text{Fe}_2\text{O}_3 + \text{SO}_2 + \text{SO}_3$$

When electricity is passed through water, it decomposes into hydrogen and oxygen gas. This is called electrolysis of water.

$$2\text{H}_2\text{O} \xrightarrow{\text{Electric current}} 2\text{H}_2 + \text{O}_2$$

When silver chloride is kept in sunlight, it turns grey because it decomposes into silver and chlorine. This is a photochemical (light-driven) decomposition reaction:

$$2\text{AgCl} \xrightarrow{\text{Sunlight}} 2\text{Ag} + \text{Cl}_2$$

Displacement Reaction

In a displacement reaction, a more reactive element displaces a less reactive element from its compound. For example, when iron nails are dipped in copper sulphate solution, iron displaces copper from the solution:

$$\text{Fe} + \text{CuSO}_4 \rightarrow \text{FeSO}_4 + \text{Cu}$$

Double Displacement Reaction

In a double displacement reaction, two compounds exchange their ions to form two new compounds. For example, when sodium sulphate solution is mixed with barium chloride solution, a white precipitate of barium sulphate is formed:

$$\text{Na}_2\text{SO}_4 + \text{BaCl}_2 \rightarrow \text{BaSO}_4 \downarrow + 2\text{NaCl}$$

Oxidation and Reduction

Oxidation is the gain of oxygen or loss of hydrogen (and loss of electrons), while reduction is the loss of oxygen or gain of hydrogen (and gain of electrons). Reactions in which oxidation and reduction occur simultaneously are called redox reactions. The substance that gives oxygen (or loses hydrogen) is called the oxidising agent, and the substance that removes oxygen (or donates hydrogen) is the reducing agent. For example:

$$\text{CuO} + \text{H}_2 \xrightarrow{\text{Heat}} \text{Cu} + \text{H}_2\text{O}$$

Here copper oxide is reduced to copper (loses oxygen) while hydrogen is oxidised to water (gains oxygen). CuO is the oxidising agent and H2 is the reducing agent.

Exothermic and Endothermic Reactions

Reactions that release heat energy are called exothermic reactions, for example respiration and burning of fuels. Reactions that absorb heat energy are called endothermic reactions, for example decomposition of calcium carbonate (limestone):

$$\text{CaCO}_3 \xrightarrow{\text{Heat}} \text{CaO} + \text{CO}_2$$

4. Effects of Oxidation in Daily Life

Corrosion

The slow destruction of a metal due to the action of air, moisture, and chemicals present in the environment is called corrosion. For example, iron rusts to form reddish-brown iron oxide, silver tarnishes to form black silver sulphide, and copper develops a green coating of basic copper carbonate. Rusting of iron requires both oxygen and water (moisture) to be present.

Rancidity

When oils and fats in food are oxidised, their smell and taste change, and the food becomes unfit for consumption. This is called rancidity. It can be prevented by storing food in airtight containers, keeping it in a refrigerator, adding antioxidants such as vitamin C and vitamin E, or packing it in the presence of nitrogen gas instead of air.

Quick Revision Tables

Table 1: Types of Chemical Reactions

Table 2: Important Reactions and their Observations

Mind Map

Important Diagrams (SVG)

Diagram 1: Combination of Calcium Oxide with Water

Diagram 2: Displacement of Copper by Iron

Common Mistakes

1. Changing the subscripts of a formula while balancing an equation instead of adding coefficients. Changing subscripts changes the identity of the compound.
2. Writing oxygen as O instead of O2 and hydrogen as H instead of H2, forgetting that these gases exist as diatomic molecules.
3. Forgetting to balance water molecules in decomposition reactions, or omitting the state symbols such as (s), (l), (g), and (aq).
4. Confusing displacement with double displacement reactions; remember double displacement involves an exchange of ions between two compounds.
5. Saying "oxidation is always addition of oxygen"; oxidation also includes loss of hydrogen and loss of electrons.
6. Forgetting that rusting of iron requires both oxygen and moisture together, not just oxygen.
7. Marking the reaction arrow or the heat symbol incorrectly, or omitting the arrow above the equals sign in decomposition reactions.

Exam Tips

1. Always balance the chemical equation first before attempting any mole-based or mass-based calculation.
2. Memorise the common diatomic molecules: H2, O2, N2, Cl2, Br2, I2, and F2, and use them correctly in equations.
3. For precipitate questions, remember the colour of the precipitate: BaSO4 is white, PbI2 is yellow, AgCl is white, and Fe(OH)3 is reddish-brown.
4. Identify oxidation-reduction by tracking the change in oxidation states: increase means oxidation, decrease means reduction.
5. Practise writing word equations from symbolic equations and vice versa, as questions often appear in both forms.
6. Note the conditions over the arrow: heat (Δ), light (sunlight), catalyst, and electric current, as these are frequently tested.
7. Learn the two daily-life applications well: corrosion (rusting) and rancidity, including at least two methods of prevention for each.

Conclusion

Chemical reactions and equations form the bedrock of the entire study of chemistry. A balanced equation accurately represents the law of conservation of mass and gives us quantitative information about the substances involved. The classification of reactions into combination, decomposition, displacement, double displacement, and redox reactions helps us organise and predict chemical behaviour. Oxidation and reduction are inseparable processes that occur everywhere in nature, from the corrosion of metals to the burning of fuels and the functioning of our bodies. By mastering these fundamental ideas, you equip yourself to understand not only chemistry but also many processes in biology, physics, and everyday life. Practise balancing equations regularly and relate every concept to a real-life example to make your learning permanent and exam-ready.