NCERT Explained: Class X Science, Chapter 1 - Chemical Reactions and Equations

Introduction to the Topic

Have you ever wondered why milk left at room temperature during summer turns sour, or why a shiny iron nail rusts when left exposed to humid air? What actually happens when food is digested in our stomachs, or when we breathe in oxygen and breathe out carbon dioxide? All these everyday phenomena are powered by a fundamental process in chemistry: chemical reactions.

In Class X Science, Chapter 1, titled Chemical Reactions and Equations, we embark on a journey to understand how substances interact, transform, and create entirely new substances with unique chemical properties. Whenever a chemical change occurs, we say that a chemical reaction has taken place. In this comprehensive guide, we will break down the fundamental concepts from the latest NCERT syllabus into simple, digestible explanations, clear step-by-step examples, and practical real-world applications to help you ace your examinations and build a solid foundation in chemistry.

Key Concepts Explained

1. What is a Chemical Reaction?

A chemical reaction is a process in which one or more substances—known as reactants—undergo a chemical transformation to form one or more different substances known as products. During a chemical reaction, existing chemical bonds between atoms are broken, and new chemical bonds are formed.

How can we determine whether a chemical reaction has taken place? According to the NCERT curriculum, a chemical reaction is usually accompanied by one or more of the following observable changes:

• Change in state: For instance, hydrogen gas and oxygen gas combine to form liquid water.
• Change in color: When an iron nail is dipped in a blue copper sulphate solution, the solution gradually turns pale green.
• Evolution of a gas: When zinc granules react with dilute sulphuric acid, hydrogen gas is evolved with effervescence.
• Change in temperature: Quicklime reacting with water releases a large amount of heat energy, making the container hot.

2. Chemical Equations: Representing Chemical Reactions

Describing a chemical reaction in full sentences can be long and tedious. Chemistry uses chemical equations as a concise shorthand notation. A chemical equation uses chemical formulas and symbols to represent a reaction.

Word Equations vs. Chemical Equations

Consider the reaction where magnesium ribbon burns in oxygen to form magnesium oxide. In word form, this is written as:

Magnesium + Oxygen → Magnesium Oxide

Here, Magnesium and Oxygen are the reactants (written on the left-hand side, or LHS), and Magnesium Oxide is the product (written on the right-hand side, or RHS). The arrow points from reactants to products, indicating the direction of the reaction.

Using chemical formulas, this reaction becomes:

Mg + O₂ → MgO

Noting Physical States

To make a chemical equation more informative, the physical states of the reactants and products are specified alongside their chemical formulas:

• (s) for solid
• (l) for liquid
• (g) for gas
• (aq) for aqueous (a solution in water)

For example: 2Mg(s) + O₂(g) → 2MgO(s)

3. Balancing Chemical Equations

If you count the number of atoms of each element on both sides of the simple equation Mg + O₂ → MgO, you will notice that there are two oxygen atoms on the LHS but only one on the RHS. Such an equation is called a skeletal chemical equation because it is unbalanced.

Why Must We Balance Chemical Equations?

The necessity of balancing equations comes directly from the Law of Conservation of Mass formulated by Antoine Lavoisier. This fundamental law states that mass can neither be created nor destroyed in a chemical reaction. Consequently, the total mass of elements present in the products must equal the total mass of elements present in the reactants. In simple terms, the number of atoms of each element must remain the same before and after a chemical reaction.

Step-by-Step Method to Balance Equations (Hit-and-Trial Method)

Let us balance the skeletal equation for the reaction of iron with steam:

Fe + H₂O → Fe₃O₄ + H₂

• Step 1: Count the atoms of each element on both sides.
  • Reactants (LHS): Fe = 1, H = 2, O = 1
  • Products (RHS): Fe = 3, H = 2, O = 4
• Step 2: Balance the element with the maximum number of atoms first. Here, oxygen has 4 atoms on the RHS and 1 on the LHS. To balance oxygen, multiply H₂O on the LHS by 4:
  Fe + 4H₂O → Fe₃O₄ + H₂
• Step 3: Balance the remaining elements (Hydrogen and Iron).
  • Hydrogen now has 8 atoms on LHS (4 × 2) and 2 on RHS. Multiply H₂ on RHS by 4:
    Fe + 4H₂O → Fe₃O₄ + 4H₂
  • Iron has 3 atoms on RHS and 1 on LHS. Multiply Fe on LHS by 3:
    3Fe + 4H₂O → Fe₃O₄ + 4H₂
• Step 4: Check the atom count on both sides.
  • LHS: Fe = 3, H = 8, O = 4
  • RHS: Fe = 3, H = 8, O = 4

The equation is now fully balanced!

4. Types of Chemical Reactions

NCERT Class X Chapter 1 categorizes chemical reactions into five main types based on the nature of chemical changes taking place:

A. Combination Reactions

A reaction in which two or more substances (elements or compounds) combine to form a single new substance is called a combination reaction.

General Formula: A + B → AB

Example 1: Reaction of quicklime (calcium oxide) with water to form slaked lime (calcium hydroxide):

CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat

Application Note: A solution of slaked lime is used for whitewashing walls. Calcium hydroxide reacts slowly with carbon dioxide in the air to form a thin, shiny layer of calcium carbonate (CaCO₃) on the walls after two to three days.

Exothermic vs Endothermic Reactions: Reactions like quicklime slaking that release heat energy along with products are called exothermic chemical reactions. Burning of natural gas (methane) and respiration are other classic examples of exothermic reactions.

B. Decomposition Reactions

A reaction in which a single compound breaks down into two or more simpler substances is known as a decomposition reaction. These reactions require energy input in the form of heat, light, or electricity, making them endothermic reactions.

General Formula: AB → A + B

Depending on the energy source used, decomposition reactions are classified into three types:

• Thermal Decomposition (Heat): Decomposition carried out by heating. For example, heating green ferrous sulphate crystals causes them to lose water and decompose into ferric oxide, sulphur dioxide, and sulphur trioxide:
  2FeSO₄(s) –[Heat]→ Fe₂O₃(s) + SO₂(g) + SO₃(g)
  Another key example is the decomposition of lead nitrate, producing brown fumes of nitrogen dioxide (NO₂):
  2Pb(NO₃)₂(s) –[Heat]→ 2PbO(s) + 4NO₂(g) + O₂(g)
• Electrolytic Decomposition (Electricity): Decomposition carried out by passing electric current. For example, electrolysis of water decomposes it into hydrogen and oxygen gases in a 2:1 volume ratio:
  2H₂O(l) –[Electric Current]→ 2H₂(g) + O₂(g)
• Photolytic Decomposition (Light): Decomposition brought about by sunlight. White silver chloride turns grey in sunlight due to its decomposition into silver and chlorine:
  2AgCl(s) –[Sunlight]→ 2Ag(s) + Cl₂(g)
  Similarly, silver bromide decomposes in sunlight. These reactions are heavily utilized in black and white photography.

C. Displacement Reactions

A chemical reaction in which a more reactive element displaces a less reactive element from its aqueous salt solution is called a displacement reaction.

General Formula: A + BC → AC + B

Example: When shiny iron nails are placed in a blue copper sulphate solution, iron (being more reactive than copper) displaces copper. The blue color fades to light green due to the formation of ferrous sulphate, and reddish-brown copper deposits on the iron nail:

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

Other examples include zinc and lead displacing copper from copper solution because zinc and lead are chemically more reactive than copper:

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

D. Double Displacement Reactions

Reactions in which there is an exchange of ions between two reactant compounds to form two new compounds are called double displacement reactions.

General Formula: AB + CD → AD + CB

Often, these reactions produce an insoluble solid called a precipitate. Thus, they are also referred to as precipitation reactions.

Example: Mixing sodium sulphate solution with barium chloride solution forms a white insoluble precipitate of barium sulphate immediately:

Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s)↓ + 2NaCl(aq)

E. Oxidation and Reduction (Redox) Reactions

Chemical reactions always involve electron transfer, which is simplified at the Class X level in terms of oxygen and hydrogen exchange:

• Oxidation: The gain of oxygen or the loss of hydrogen by a substance during a reaction.
• Reduction: The loss of oxygen or the gain of hydrogen by a substance during a reaction.
• Redox Reaction: A reaction in which one reactant gets oxidized while the other gets reduced simultaneously.

Example 1: Heating black copper oxide with hydrogen gas:

CuO + H₂ –[Heat]→ Cu + H₂O

• CuO loses oxygen to become Cu → CuO is reduced.
• H₂ gains oxygen to become H₂O → H₂ is oxidized.
• CuO acts as the oxidizing agent (supplies oxygen), while H₂ acts as the reducing agent.

Example 2: Reaction of zinc oxide with carbon:

ZnO + C → Zn + CO

Here, ZnO is reduced to Zn, and C is oxidized to carbon monoxide (CO).

5. Everyday Effects of Oxidation Reactions

Oxidation isn't confined to chemistry laboratories; it constantly affects materials around us in daily life. Two major consequences detailed in NCERT Chapter 1 are corrosion and rancidity.

A. Corrosion

When a metal is attacked by environmental substances such as moisture, acids, air, or oxygen, it degrades over time. This process is called corrosion.

• Rusting of Iron: Iron objects acquire a flaky, reddish-brown coating when exposed to moist air for prolonged periods.
• Black coating on Silver: Silver ornaments turn dark due to reaction with hydrogen sulphide gas present in the air, forming silver sulphide.
• Green coating on Copper: Copper vessels acquire a greenish layer of basic copper carbonate when exposed to moist carbon dioxide over time.

Impact and Prevention: Corrosion causes serious damage to bridges, iron railings, ships, and automobile bodies, leading to massive economic loss annually. It can be prevented through painting, greasing, oiling, galvanization (coating iron with a protective layer of zinc), electroplating, or alloying.

B. Rancidity

When fats and oils present in food items are oxidized by atmospheric oxygen, they undergo degradation. Their taste, aroma, and flavor change, rendering them unpalatable. This phenomenon is known as rancidity.

Methods to Prevent Rancidity:

• Adding antioxidants (substances that retard oxidation) to foods containing fats and oils.
• Flushing packaged food products (like potato chips bags) with unreactive gas such as nitrogen to prevent the food from coming into contact with oxygen.
• Storing food in air-tight containers to slow down oxidation.
• Keeping food in refrigerators at low temperatures.

Summary & Key Takeaways

To summarize Class X Science Chapter 1, keep these core concepts fresh in your mind for quick revision:

• Chemical Reaction: A process leading to the chemical transformation of one set of chemical substances into another, accompanied by changes in state, color, temperature, or gas evolution.
• Law of Conservation of Mass: Total mass of reactants must equal total mass of products. Therefore, all chemical equations must be balanced so that the number of atoms of each element is identical on both sides.
• Combination Reaction: Two or more reactants combine to form a single product (e.g., quicklime + water → slaked lime).
• Decomposition Reaction: A single reactant breaks down into multiple simpler products via heat (thermal), electricity (electrolytic), or light (photolytic).
• Displacement Reaction: A more reactive element displaces a less reactive element from its aqueous salt solution.
• Double Displacement Reaction: Mutual exchange of ions between two ionic compounds, often yielding an insoluble precipitate.
• Redox Reaction: Simultaneous occurrence of oxidation (gain of O / loss of H) and reduction (loss of O / gain of H).
• Corrosion: Degradation of metals due to air, water, and chemical attack (e.g., rusting of iron).
• Rancidity: Oxidation of fats/oils in food altering taste and smell; prevented by nitrogen flushing and airtight storage.