Chemical Reactions and Equations

Everyday Examples of Chemical Changes

• Milk left at room temp becomes sour (curdling).
• Iron exposed to moisture rusts.
• Fermentation of grapes produces alcohol.
• Cooking changes food’s texture and taste.
• Digestion converts food into nutrients.
• Respiration oxidizes food, releasing heat and energy.

Chemical Changes

• Defined as changes where the chemical nature of initial substances is altered.
• Occur due to chemical reactions.

Laboratory Examples of Chemical Reactions

1. Magnesium and Oxygen
   • Magnesium ribbon burns in oxygen, producing white ash (magnesium oxide).
2. Potassium Iodide and Lead Nitrate
   • Mixing leads to formation of yellow precipitate (lead iodide).
3. Zinc and Hydrochloric Acid
   • Zinc reacts with dilute HCl or H2SO4, releasing hydrogen gas and heat.

Observations of Chemical Reactions

• Change in state
• Change of color
• Evolution of gas
• Change in temperature
• Change in chemical nature

Chemical Equations

• Definition: Representation of a chemical reaction using symbols and formulas.
• Example:
  • Magnesium + Oxygen → Magnesium Oxide
  • Written as: Mg + O2 → MgO
• Components:
  • Reactants: Substances before reaction (left side).
  • Products: Substances after reaction (right side).
  • Arrow indicates direction of reaction.

Balancing Chemical Equations

• Law of Conservation of Mass: Mass of reactants = mass of products.
• Steps to Balance:
  1. Enclose formulas in boxes.
  2. Count atoms of each element in reactants and products.
  3. Adjust coefficients to balance the atoms.
  4. Check for balance on both sides.
  5. Include physical states if necessary.

Example of Balancing

• Unbalanced: Fe + H2O → Fe3O4 + H2
• Balanced: 3Fe + 4H2O → Fe3O4 + 4H2

Types of Chemical Reactions

1. Combination Reaction

   • Two or more reactants combine to form a single product.
   • Example: C + O2 → CO2
   • Example: H2 + O2 → 2H2O

2. Exothermic Reactions

   • Heat released during the reaction.
   • Example: Combustion of natural gas.

3. Decomposition Reaction

   • A single compound breaks down into simpler substances.
   • Example: Heating calcium carbonate → calcium oxide + carbon dioxide.
   • Types:
     • Thermal Decomposition
     • Electrolytic Decomposition: Break down by electricity (e.g., electrolysis of water).
     • Photolytic Decomposition: Break down by light (e.g., silver chloride to silver and chlorine).

4. Displacement Reaction

   • One element displaces another in a compound.
   • Example: Zn + CuSO4 → ZnSO4 + Cu

5. Double Displacement Reaction

   • Two compounds exchange ions to form new compounds.
   • Example: Na2SO4 + BaCl2 → BaSO4 (precipitate) + NaCl

Oxidation and Reduction

• Oxidation: Gain of oxygen (e.g., copper + oxygen → copper oxide).
• Reduction: Loss of oxygen (e.g., copper oxide + hydrogen → copper + water).
• Redox Reactions: Reactions involving both oxidation and reduction.

Effects of Oxidation

1. Corrosion
   • Deterioration of metals (e.g., rusting of iron).
2. Rancidity
   • Oxidation of fats and oils leads to off smells and tastes.
   • Preventive measures: Add antioxidants, store in tight containers, use nitrogen gas in packaging.