Chapter 1: Chemical Reactions and Equations

Key Concepts

• Chemical reactions involve changes in the nature and identity of substances, signifying chemical changes.
• Indicators of a chemical reaction include change in state, color, evolution of gas, and temperature changes.

Activities

Activity 1.1

• Objective: Burn a magnesium ribbon to observe the formation of magnesium oxide.
• Procedure: Clean magnesium ribbon, burn it, and collect the ash.
• Observation: White powder formation indicates a chemical reaction.

Activity 1.2

• Objective: Observe hydrogen gas formation by reacting zinc with dilute sulfuric acid.

Activity 1.3

• Objective: Study temperature changes when zinc reacts with dilute hydrochloric acid and lead nitrate with potassium iodide.

Chemical Equations

1.1 Writing Chemical Equations

• Chemical reactions can be represented by word equations or chemical equations.
• Example: [ \text{Mg} + \text{O}2 \rightarrow \text{MgO} ]

1.1.1 Balanced Chemical Equations

• The number of atoms for each element must be equal on both sides of a reaction equation, as per the law of conservation of mass.
• Balancing involves adjusting coefficients (not subscripts).

Types of Chemical Reactions

1.2.1 Combination Reaction

• Two or more reactants combine to form one product.
• Example: [ \text{CaO} + \text{H}_2\text{O} \rightarrow \text{Ca(OH)}_2 ]

1.2.2 Decomposition Reaction

• A single compound breaks down into two or more simpler substances.
• Example: [ \text{CaCO}_3 \xrightarrow{\text{Heat}} \text{CaO} + \text{CO}_2 ]

1.2.3 Displacement Reaction

• One element displaces another in a compound.
• Example: [ \text{Fe} + \text{CuSO}_4 \rightarrow \text{FeSO}_4 + \text{Cu} ]

1.2.4 Double Displacement Reaction

• Exchange of ions between two compounds.
• Example: [ \text{Na}_2\text{SO}_4 + \text{BaCl}_2 \rightarrow \text{BaSO}_4 + 2\text{NaCl} ]

1.2.5 Oxidation and Reduction

• Oxidation: Gain of oxygen or loss of hydrogen.
• Reduction: Loss of oxygen or gain of hydrogen.
• Example: [ \text{CuO} + \text{H}_2 \rightarrow \text{Cu} + \text{H}_2\text{O} ]

Everyday Effects

1.3.1 Corrosion

• Metals react with environmental substances leading to deterioration (e.g., rusting of iron).

1.3.2 Rancidity

• Oxidation of fats and oils causes bad odor and taste; prevented by antioxidants or nitrogen flushing.

Summary

• Chemical equations represent the transformation of reactants to products, showing balanced atoms and potentially physical states.
• Reactions include combination, decomposition, displacement, and double displacement.
• Redox reactions involve oxidation and reduction processes with associated real-world implications like corrosion and rancidity.

Exercises

• Balance given chemical equations.
• Identify types of reactions.
• Explain concepts such as exothermic/endothermic reactions, corrosion, and rancidity with examples.