Chapter 1: Chemical Reactions and Equations

Chemical Reactions and Changes

Daily Life Situations Involving Chemical Changes

• Milk left at room temperature during summers – Spoils due to bacterial action (fermentation).
• Iron tawa/pan/nail exposed to humid air – Rusts due to oxidation.
• Grapes fermenting – Sugar converts into alcohol (fermentation).
• Food cooking – New substances form due to heating (chemical change).
• Food digestion – Breakdown of food into simpler substances by enzymes.
• Respiration – Glucose breaks down to release energy (oxidation).

All these indicate chemical changes where the identity of the substance changes.

Activity 1.1: Burning of Magnesium Ribbon

Observation: Burns with a dazzling white flame. Forms a white powder (Magnesium Oxide).

Reaction: 2Mg + O₂ → 2MgO

Conclusion: Change in state: Magnesium (solid) to Magnesium Oxide (powder). Change in color: Silver magnesium to white powder. Evolution of heat and light: Exothermic reaction.

Activity 1.2: Reaction of Zinc with Dilute Acid

Observation: Bubbles form (hydrogen gas). Flask becomes warm (exothermic reaction).

Reaction: Zn + H₂SO₄ → ZnSO₄ + H₂↑

Conclusion: Gas evolution: Hydrogen gas. Temperature change: Flask becomes warm.

Activity 1.3: Precipitation Reaction

Observation: Yellow precipitate forms.

Reaction: Pb(NO₃)₂ + 2KI → PbI₂↓ + 2KNO₃

Conclusion: Change in color: Yellow precipitate of lead iodide. Formation of an insoluble substance (precipitate).

Indicators of a Chemical Reaction

• Change in state
• Change in color
• Evolution of a gas
• Change in temperature

These observations help us identify a chemical reaction.

Chemical Equations

Definition of a Chemical Equation

A chemical equation is a symbolic representation of a chemical reaction. It expresses reactants and products with an arrow (→) indicating the reaction direction.

Example: Burning of Magnesium Ribbon (Activity 1.1)

• Word Equation: Magnesium + Oxygen → Magnesium Oxide
• Reactants: Magnesium (Mg) and Oxygen (O₂).
• Product: Magnesium Oxide (MgO).

Parts of a Chemical Equation

• Reactants: Substances that undergo change (written on LHS).
• Products: New substances formed (written on RHS).
• Arrow (→): Indicates the direction of the reaction.
• Plus Sign (+): Separates multiple reactants or products.

Types of Chemical Equations

• Word Equation: Represents a reaction using words. Example: Magnesium + Oxygen → Magnesium Oxide
• Symbolic Equation: Uses chemical symbols and formulas. Example: Mg + O₂ → MgO

Writing a Chemical Equation

1. Representation of Chemical Equations: A chemical equation is a concise way to represent a chemical reaction. Instead of writing names, chemical symbols and formulas are used.
2. Example: Burning of Magnesium
   Word Equation: Magnesium + Oxygen → Magnesium Oxide
   Chemical Equation: Mg + O₂ → MgO
3. Skeletal Chemical Equation: In equation Mg + O₂ → MgO, the number of atoms on LHS and RHS are:
   LHS: Mg = 1, O = 2
   RHS: Mg = 1, O = 1
   Unequal oxygen atoms → The equation is unbalanced. An unbalanced equation is called a skeletal equation because the mass is not equal on both sides.
4. Importance of Balancing Equations: A chemical reaction follows the Law of Conservation of Mass (mass remains constant). A balanced chemical equation ensures the same number of atoms on both sides. Balancing equations is the next step in correctly representing chemical reactions.

Balanced Chemical Equations

Law of Conservation of Mass

Mass can neither be created nor destroyed in a chemical reaction. The total mass of reactants = total mass of products. The number of atoms of each element remains the same before and after the reaction.

Steps to Balance a Chemical Equation

Example: Balancing Fe + H₂O → Fe₃O₄ + H₂

1. Draw Boxes: Do not change anything inside the chemical formulas.
2. Count Atoms of Each Element:

Element	Reactants (LHS)	Products (RHS)
Fe	1	3
H	2	2
O	1	4

3. Balance Oxygen Atoms: Place 4 before H₂O to balance oxygen. Fe + 4H₂O → Fe₃O₄ + H₂
4. Balance Hydrogen Atoms: Place 4 before H₂ on RHS. Fe + 4H₂O → Fe₃O₄ + 4H₂
5. Balance Iron (Fe) Atoms: Place 3 before Fe on LHS. 3Fe + 4H₂O → Fe₃O₄ + 4H₂
6. Verify Balanced Equation:

Element	Reactants (LHS)	Products (RHS)
Fe	3	3
H	8	8
O	4	4

✅ The equation is balanced.

Writing Physical States in Chemical Equations

• (s) → Solid
• (l) → Liquid
• (g) → Gas
• (aq) → Aqueous (dissolved in water)

Final Balanced Equation with Physical States:

3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

Water is in gaseous (steam) form in this reaction.

Reaction Conditions in Equations

Temperature, pressure, or catalysts may be mentioned above/below the reaction arrow.

Example:

CO(g) + 2H₂(g) ^{340 atm}→ CH₃OH(l)

Methanol is produced under 340 atm pressure.

Example (Photosynthesis):

6CO₂(aq) + 12H₂O(l) ^{Sunlight, Chlorophyll}→ C₆H₁₂O₆(aq) + 6O₂(aq) + 6H₂O(l)

Types of Chemical Reactions

1.2.1 Combination Reaction

A combination reaction is a reaction in which two or more reactants (elements or compounds) combine to form a single product. These reactions are often exothermic, meaning they release heat.

Example: Reaction of Calcium Oxide with Water

Reaction: CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat

Quick lime (CaO) reacts vigorously with water to form slaked lime (Ca(OH)₂) and releases heat.

Other Examples of Combination Reactions:

• Burning of Coal: C(s) + O₂(g) → CO₂(g)
• Formation of Water: 2H₂(g) + O₂(g) → 2H₂O(l)

Exothermic Reactions

Definition: Reactions that release heat along with product formation.

Examples:

• Burning of Natural Gas: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g) + Heat
• Respiration (Energy Release in Cells): C₆H₁₂O₆(aq) + 6O₂(aq) → 6CO₂(aq) + 6H₂O(l) + Energy

The breakdown of glucose in our body cells to release energy is an exothermic process. Decomposition of Vegetable Matter into Compost is also an exothermic reaction.

Activity 1.1 (Burning of Magnesium Ribbon)

Type of Reaction: Combination Reaction

Reason: Magnesium (Mg) reacts with oxygen (O₂) to form a single product, magnesium oxide (MgO).

Exothermic Nature: Heat is released during the reaction.

1.2.2 Decomposition Reaction

A decomposition reaction is a reaction in which a single reactant breaks down into two or more simpler products. These reactions require energy in the form of heat, light, or electricity, making them endothermic reactions (energy is absorbed).

Types of Decomposition Reactions

1. Thermal Decomposition (Heat-Based Decomposition)

Definition: When a decomposition reaction is carried out by heating, it is called thermal decomposition.

Examples:

• Decomposition of Ferrous Sulphate (FeSO₄): On heating, green ferrous sulphate crystals decompose into ferric oxide (Fe₂O₃), sulphur dioxide (SO₂), and sulphur trioxide (SO₃).
  Reaction: 2FeSO₄(s) ^Heat → Fe₂O₃(s) + SO₂(g) + SO₃(g)
  Observation: Green crystals turn brown, and a characteristic smell of burning sulphur is noticed.
• Decomposition of Calcium Carbonate (CaCO₃): On heating, calcium carbonate (limestone) decomposes into calcium oxide (quick lime) and carbon dioxide gas.
  Reaction: CaCO₃(s)^Heat → CaO(s) + CO₂(g)
  Importance: Quick lime is used in cement production.
• Decomposition of Lead Nitrate (Pb(NO₃)₂): On heating, lead nitrate decomposes to form lead oxide (PbO), nitrogen dioxide (NO₂), and oxygen (O₂).
  Reaction: 2Pb(NO₃)₂(s) ^Heat → 2PbO(s) + 4NO₂(g) + O₂(g)
  Observation: Brown fumes of nitrogen dioxide are seen.

2. Photochemical Decomposition (Light-Based Decomposition)

Definition: When decomposition occurs due to light energy, it is called photochemical decomposition.

Examples:

• Decomposition of Silver Chloride (AgCl): In sunlight, white silver chloride decomposes into grey silver metal and chlorine gas.
  Reaction: 2AgCl(s) ^Sunlight → 2Ag(s) + Cl₂(g)
  Observation: The white colour of AgCl turns grey.
• Decomposition of Silver Bromide (AgBr):
  Reaction: 2AgBr(s)^Sunlight → 2Ag(s) + Br₂(g)
  Use: This reaction is used in black and white photography.

3. Electrolytic Decomposition (Electricity-Based Decomposition)

Definition: When decomposition occurs due to electric current, it is called electrolytic decomposition.

Example: Electrolysis of Water (H₂O): When electric current is passed through water, it decomposes into hydrogen gas (H₂) and oxygen gas (O₂).
Reaction: 2H₂O(l) ^Electricity→ 2H₂(g) + O₂(g)
Observation: Hydrogen collects at the cathode. Oxygen collects at the anode. Twice the volume of hydrogen is collected compared to oxygen.

Endothermic vs. Exothermic Reactions

• Endothermic Reactions: Reactions that absorb energy (heat, light, or electricity ) to break down reactants. Examples: Decomposition of calcium carbonate, silver chloride, and water.
• Exothermic Reactions: Reactions that release energy (heat) during product formation. Example: Combustion of fuels, respiration, and decomposition of organic matter into compost.

Experiment on Endothermic Reaction (Barium Hydroxide + Ammonium Chloride)

Ba(OH)₂ + 2NH₄Cl → BaCl₂ + 2NH₃ + H₂O

Heat is absorbed, and the test tube feels cold. This confirms it is an endothermic reaction.

Summary of Decomposition Reactions

Type	Energy Required	Examples
Thermal Decomposition	Heat	FeSO₄, CaCO₃, Pb(NO₃)₂
Photochemical Decomposition	Light	AgCl, AgBr
Electrolytic Decomposition	Electricity	H₂O

Displacement Reaction

Activity 1.9: Reaction Between Iron and Copper Sulphate Solution

The iron nail turns brownish in color. The blue color of copper sulphate solution fades.

Iron (Fe) is more reactive than copper (Cu). It displaces copper from copper sulphate (CuSO₄) solution. The iron nail gets coated with copper, giving it a brownish appearance. The solution changes from blue to light green due to the formation of iron sulphate (FeSO₄).

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

This type of reaction is called a displacement reaction, where a more reactive metal replaces a less reactive metal in a compound.

Other Examples of Displacement Reactions:

• Zinc displacing Copper: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
• Lead displacing Copper: Pb(s) + CuCl₂(aq) → PbCl₂(aq) + Cu(s)

Conclusion:

More reactive metals displace less reactive metals from their compounds. Displacement reactions are common in metal reactivity and are used in various industrial processes.

Double Displacement Reaction

Activity 1.10: Reaction Between Sodium Sulphate and Barium Chloride Solutions

A white insoluble substance (precipitate) is formed. This precipitate is barium sulphate (BaSO₄), which is insoluble in water.

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

Explanation: Sodium sulphate (Na₂SO₄) and barium chloride (BaCl₂) are both soluble in water. When mixed, barium ions (Ba²⁺) react with sulphate ions (SO₄²⁻) to form barium sulphate (BaSO₄), which is insoluble and precipitates out. Sodium chloride (NaCl) remains dissolved in water.

Concept of Double Displacement Reaction:

Definition: A reaction in which two compounds exchange their ions to form new products. This reaction is also called a precipitation reaction because it forms an insoluble solid (precipitate).

Recall Activity 1.2: Reaction Between Lead(II) Nitrate and Potassium Iodide Solutions

1. What was the colour of the precipitate formed? The precipitate was yellow. The compound formed was lead(II) iodide (PbI₂).
2. Balanced Chemical Equation: Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)
3. Is this also a double displacement reaction? Yes, this is a double displacement reaction because Pb²⁺ and I⁻ ions exchange places, forming lead(II) iodide precipitate and soluble potassium nitrate (KNO₃).

Key Takeaways:

Double displacement reactions involve the exchange of ions between two compounds. Precipitation reactions are a type of double displacement reaction where an insoluble solid (precipitate) is formed.

Examples:

• BaSO₄ (white precipitate) from BaCl₂ and Na₂SO₄.
• PbI₂ (yellow precipitate) from Pb(NO₃)₂ and KI.

Oxidation and Reduction (Redox Reactions)

Activity 1.11: Oxidation of Copper

When copper powder is heated, it turns black. This happens because copper reacts with oxygen to form copper(II) oxide (CuO).

2Cu + O₂ ^Heat→ 2CuO

Copper (Cu) gains oxygen and forms Copper(II) oxide (CuO) → Oxidation.

Reduction of Copper Oxide with Hydrogen

CuO + H₂ → ^Heat→ Cu + H₂O

Copper(II) oxide (CuO) loses oxygen → Reduced. Hydrogen (H₂) gains oxygen to form water (H₂O) → Oxidised.

Definition of Oxidation and Reduction

• Oxidation: A substance gains oxygen or loses hydrogen.
• Reduction: A substance loses oxygen or gains hydrogen.

Redox Reaction: A reaction in which one reactant gets oxidised while the other gets reduced.

Other Examples of Redox Reactions

• Example 1: ZnO + C → Zn + CO
  Carbon (C) gains oxygen → Oxidised to CO. Zinc oxide (ZnO) loses oxygen → Reduced to Zn.
• Example 2: MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂
  Hydrochloric acid (HCl) is oxidised to chlorine (Cl₂). Manganese dioxide (MnO₂) is reduced to manganese chloride (MnCl₂).

Oxidation of Magnesium in Activity 1.1

When magnesium ribbon burns in air, it forms magnesium oxide (MgO).

2Mg + O₂ → 2MgO

Magnesium gains oxygen → Oxidised to MgO.

Key Takeaways

• Oxidation: Gain of oxygen / Loss of hydrogen.
• Reduction: Loss of oxygen / Gain of hydrogen.
• Redox Reaction: Both oxidation and reduction happen simultaneously.

Examples:

• Copper + Oxygen → Copper Oxide (Oxidation).
• Copper Oxide + Hydrogen → Copper + Water (Reduction).
• Zinc Oxide + Carbon → Zinc + Carbon Monoxide (Redox Reaction).
• Manganese Dioxide + HCl → Manganese Chloride + Water + Chlorine (Redox Reaction).

Corrosion

What is Corrosion?

Corrosion is the process in which metals are attacked by substances in their environment, such as moisture, acids, and air. This reaction damages the metal surface over time.

Examples of Corrosion

• Rusting of Iron: Iron (Fe) reacts with oxygen (O₂) and water (H₂O) to form hydrated iron oxide (Fe₂O₃·xH₂O), known as rust.
  Rusting Equation: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃
• Tarnishing of Silver: Silver (Ag) reacts with hydrogen sulfide (H₂S) in the air to form a black coating of silver sulfide (Ag₂S).
  Equation: 2Ag + H₂S → Ag₂S + H₂
• Green Coating on Copper: Copper (Cu) reacts with oxygen (O₂), water (H₂O), and carbon dioxide (CO₂) to form basic copper carbonate (CuCO₃·Cu(OH)₂), a green coating.
  Equation: 2Cu + O₂ + H₂O + CO₂ → CuCO₃·Cu(OH)₂

Effects of Corrosion

• Weakens metals and reduces their lifespan.
• Damages iron objects like bridges, railings, ships, and car bodies.
• Causes economic loss, as a lot of money is spent on maintenance and replacement.

Prevention of Corrosion (Will be covered in Chapter 3)

• Painting, oiling, or greasing metal surfaces.
• Galvanization (coating with zinc) for iron and steel.
• Electroplating or alloying (e.g., stainless steel) to prevent rusting.

Rancidity

What is Rancidity?

Rancidity occurs when fats and oils are oxidized, leading to unpleasant taste and smell. This happens when food is exposed to air (oxygen) for a long time.

Causes of Rancidity

• Oxidation of fats and oils by oxygen in the air.
• Exposure to heat, light, or moisture speeds up the process.

Prevention of Rancidity

• Use of Antioxidants: Antioxidants prevent oxidation and delay rancidity. Examples: BHA (Butylated hydroxyanisole), BHT (Butylated hydroxytoluene), Vitamin C, and Vitamin E.
• Airtight Packaging: Storing food in airtight containers reduces exposure to oxygen.
• Refrigeration: Keeping food at low temperatures slows down oxidation.
• Flushing with Nitrogen Gas: Manufacturers flush bags of chips with nitrogen to prevent oxidation and keep them fresh for longer.
• Using Dark or Opaque Containers: Protects food from light, which can speed up oxidation.