Chapter 1: Chemical Reactions and Equations

1. Introduction to Chemical Reactions

A chemical reaction is a process in which one or more substances (reactants) are converted into one or more new substances (products) with different properties.

Characteristics of a Chemical Reaction

• Change in state: Solid, liquid, or gas formation.

• Change in color: Example – Rusting of iron (Iron turns reddish-brown).

• Evolution of gas: Example – Reaction of zinc with hydrochloric acid produces hydrogen gas.

• Change in temperature:

  • Exothermic reaction – Heat is released (e.g., combustion of fuels).

  • Endothermic reaction – Heat is absorbed (e.g., decomposition of calcium carbonate).

• Formation of precipitate: Example – Mixing silver nitrate and sodium chloride forms a white precipitate of silver chloride.

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2. Chemical Equations

A chemical equation is a representation of a chemical reaction using symbols and formulas.

Example of a Chemical Equation

Magnesium burns in oxygen to form magnesium oxide:

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

Balanced and Unbalanced Chemical Equations

A balanced chemical equation has an equal number of atoms of each element on both sides.

Example (Balanced Equation for the combustion of hydrogen):

$$2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$$

Law of Conservation of Mass

Mass is neither created nor destroyed in a chemical reaction.
Thus, the number of atoms of each element must be equal on both sides of a balanced equation.

Steps to Balance a Chemical Equation

1. Identify reactants and products.

2. Count the number of atoms of each element.

3. Use coefficients to balance atoms.

4. Ensure total atoms are equal on both sides.

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3. Types of Chemical Reactions

1. Combination Reaction

When two or more reactants combine to form a single product.

Example:

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

This is an exothermic reaction as heat is released.

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2. Decomposition Reaction

A single compound breaks down into two or more simpler substances.

Types of Decomposition Reactions:

• Thermal Decomposition (heat required)

  $$\text{CaCO}_3 \xrightarrow{\text{heat}} \text{CaO} + \text{CO}$$

• Electrolytic Decomposition (electricity required)

  $$2\text{H}_2\text{O} \xrightarrow{\text{electricity}} 2\text{H}_2 + \text{O}_2$$

• Photolytic Decomposition (light required)

  $$2\text{AgBr} \xrightarrow{\text{light}} 2\text{Ag} + \text{Br}_2$$

(Used in black and white photography.)

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3. Displacement Reaction

A more reactive element displaces a less reactive element from its compound.

Example:

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

(Iron displaces copper from copper sulfate solution.)

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4. Double Displacement Reaction

Exchange of ions between two compounds to form new compounds.

Example:

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

(BaSO₄ is a white precipitate.)

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5. Oxidation and Reduction (Redox Reactions)

• Oxidation: Addition of oxygen or removal of hydrogen.

• Reduction: Removal of oxygen or addition of hydrogen.

• Redox Reaction: Both oxidation and reduction occur together.

Example:

$$\text{CuO} + \text{H}_2 \rightarrow \text{Cu} + \text{H}_2\text{O}$$

(CuO is reduced to Cu, and H₂ is oxidized to H₂O.)

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4. Corrosion and Rancidity

Corrosion

• Slow destruction of metals by reaction with air, water, or chemicals.

• Example: Rusting of iron.

  $$4\text{Fe} + 3\text{O}_2 + 6\text{H}_2\text{O} \rightarrow 4\text{Fe(OH)}_3$$

• Prevention:

  • Painting

  • Galvanization (coating with zinc)

  • Oiling and greasing

  • Alloying (mixing metals like stainless steel)

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Rancidity

• Spoiling of food due to oxidation of fats and oils.

• Prevention:

  • Antioxidants (e.g., vitamin C, vitamin E)

  • Refrigeration

  • Airtight storage

  • Vacuum packing

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5. Important Questions for Practice

1. Short Answer Questions:

1. Define a chemical reaction. Give an example.

2. What is a balanced chemical equation? Why is it necessary to balance a chemical equation?

3. Write the balanced equation for the decomposition of lead nitrate.

4. What is corrosion? How can it be prevented?

5. Explain the difference between oxidation and reduction with examples.

2. Long Answer Questions:

1. Describe different types of chemical reactions with examples.

2. Explain how you can balance the following equation step by step:

   $$\text{Fe} + \text{H}_2\text{O} \rightarrow \text{Fe}_3\text{O}_4 + \text{H}_2$$

3. What is rancidity? How can it be prevented in food items?

4. Describe an experiment to show a double displacement reaction.

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Conclusion

• Chemical reactions play a vital role in our daily life, from rusting of iron to cooking food.

• The law of conservation of mass governs all chemical equations, ensuring balance.

• Different types of reactions help us understand various chemical changes and their applications.

• Oxidation and reduction reactions are crucial for many biological and industrial processes.