Chapter 1: Chemical Reactions and Equations (Detailed Explanation)

1. Introduction to Chemical Reactions

A chemical reaction occurs when substances (called reactants) undergo a change to form new substances (called products). During this process, the atoms in the reactants are rearranged to form the products.

Features of Chemical Reactions:

• Change in color: A change in color indicates that a chemical reaction has occurred. For example, the reaction of iron with oxygen leads to the formation of rust, which is reddish-brown. 
• Evolution of gas: Gas may be released during a chemical reaction. For example, when hydrochloric acid reacts with zinc, hydrogen gas is released. 
• Formation of a precipitate: A solid substance that forms when two solutions react is called a precipitate. For example, when silver nitrate reacts with sodium chloride, silver chloride (a white precipitate) forms. 
• Change in temperature: A chemical reaction may release or absorb energy in the form of heat. For example, the reaction of calcium oxide with water is exothermic (releases heat).

2. Types of Chemical Reactions

A. Combination Reaction

Definition: A combination reaction occurs when two or more reactants combine to form a single product.

General Form: 

            A + B → AB

Example 1:

Reaction of hydrogen with oxygen:

           2H₂(g) + O₂(g) → 2H₂O(l)

Hydrogen and oxygen combine to form water. This is a simple combination reaction.

Example 2:

Formation of calcium oxide:

            Ca(s) + O₂(g) → CaO(s)

Calcium reacts with oxygen to form calcium oxide.

B. Decomposition Reaction

Definition: A decomposition reaction is the opposite of a combination reaction. It occurs when a single reactant breaks down into two or more simpler products.

• General Form:
                AB → A + B

Example 1:

Decomposition of mercury(II) oxide:

        2HgO(s) → 2Hg(l) + O₂(g)

When mercury(II) oxide is heated, it decomposes into mercury and oxygen.

Example 2:

Decomposition of calcium carbonate:

      CaCO3(s) → CaO(s) + CO2(g)

When calcium carbonate is heated, it decomposes to form calcium oxide (quicklime) and carbon dioxide.

C. Displacement Reaction

Definition: In a displacement reaction, one element displaces another from a compound.

Types of Displacement Reactions:

(i) Single Displacement Reaction: One element replaces another element in a compound.

Example:

• Zinc displaces copper from copper sulfate:

          Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)

Zinc metal displaces copper from copper sulfate solution, forming zinc sulfate and copper metal.

(ii) Double Displacement Reaction: Two compounds exchange their ions to form two new compounds.

Example:

• Reaction between sodium chloride and silver nitrate:

         NaCl(aq) + AgNO₃(aq) → NaNO₃(aq) + AgCl(s)

Here, sodium chloride reacts with silver nitrate to form sodium nitrate and a white precipitate of silver chloride.

D. Redox Reaction (Reduction-Oxidation Reaction)

Definition: A redox reaction involves the transfer of electrons between two substances. One substance gets oxidized (loses electrons) and another gets reduced (gains electrons).

Example:

• Rusting of iron:

       4Fe + 3O₂ → 2Fe₂O₃​

In this reaction, iron loses electrons (oxidation) and oxygen gains electrons (reduction), forming iron (III) oxide, which is rust.

E. Neutralization Reaction

Definition: In a neutralization reaction, an acid reacts with a base to form a salt and water.

Example:

• Reaction of hydrochloric acid with sodium hydroxide:

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

Hydrochloric acid reacts with sodium hydroxide to form sodium chloride (salt) and water.

3. Balancing Chemical Equations

Balancing chemical equations ensures that the number of atoms of each element is the same on both sides of the equation, which follows the law of conservation of mass.

Steps to Balance Equations:

1. Write the unbalanced equation: Write down the chemical equation as given.
2. Count the number of atoms: Count how many atoms of each element are present on both sides of the equation.
3. Balance the atoms: Add coefficients (the numbers in front of compounds) to balance the atoms of each element on both sides.
4. Check and simplify: Ensure all atoms are balanced and that coefficients are in the smallest whole number ratio.

Example of Balancing:

• Unbalanced equation:

             H2 + O2 → H₂O

• Balanced equation:

             2H₂ + O₂ → 2H₂O

Here, we added coefficients to balance hydrogen and oxygen atoms on both sides.

4. Importance of Reaction Conditions

• Effect of Temperature:

  • High temperature generally speeds up reactions. For example, heating a substance increases the energy of particles, making them collide more often and with greater energy.

• Effect of Pressure:

  • In reactions involving gases, increasing pressure can increase the rate of reaction or shift equilibrium reactions.

• Effect of Concentration:

  • A higher concentration of reactants increases the frequency of collisions between particles, thus increasing the reaction rate.

• Catalysts:

  • A catalyst is a substance that increases the rate of a reaction without being consumed in the process. It lowers the activation energy required to start the reaction.
  • Example: The enzyme in our body acts as a catalyst in many biological reactions, like digestion.

5. Examples of Chemical Reactions

(i) Magnesium and Oxygen Reaction (Combination):

           2Mg(s) + O2(g) → 2MgO(s)

Magnesium burns in oxygen to form magnesium oxide. This is a combination reaction.

(ii) Calcium Carbonate Decomposition (Decomposition):

          CaCO3(s) → CaO(s) + CO2(g)

Calcium carbonate decomposes when heated to form calcium oxide and carbon dioxide.

Summary of Chapter 1: Chemical Reactions and Equations

• Chemical reactions involve the rearranging of atoms and lead to the formation of new substances.
• Types of chemical reactions include combination, decomposition, displacement (single and double), redox, and neutralization reactions.
• Balancing chemical equations ensures that mass is conserved in a reaction.
• Reaction conditions like temperature, pressure, concentration, and catalysts can affect the rate of a reaction.