Chemical Reactions and Equations

Correct Option: (i) (a) and (b)

Explanation:
In the given reaction:
• PbO is losing oxygen to form Pb. Loss of oxygen is reduction. So, PbO is getting reduced (d is correct).
• C is gaining oxygen to form CO₂. Gain of oxygen is oxidation. So, Carbon is getting oxidised (c is correct).
• Statement (a) says "Lead is getting reduced". This is incorrect because Pb in PbO is already in compound form; it's PbO that is reduced.
• Statement (b) says "CO₂ is getting oxidised". This is incorrect because CO₂ is the product of oxidation, not undergoing oxidation itself.
Thus, the incorrect statements are (a) and (b).

Correct Option: (d) displacement reaction.

Explanation:
This is a displacement reaction where aluminium (Al) displaces iron (Fe) from its compound Fe₂O₃ because Al is more reactive than Fe. It is also a highly exothermic reaction known as the thermite reaction, used for welding railway tracks.

Correct Option: (a) Hydrogen gas and iron chloride are produced.

Explanation:
Iron reacts with dilute HCl to form iron(II) chloride and hydrogen gas.
Chemical Equation: Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)
This is a common metal-acid reaction where a metal displaces hydrogen from the acid.

Balanced Chemical Equation: A chemical equation in which the number of atoms of each element is equal on both sides (reactants and products).

Why balancing is essential:
1. Law of Conservation of Mass: Matter cannot be created or destroyed. Balancing ensures the total mass of reactants equals the total mass of products.
2. Atom Conservation: The number of atoms of each element must remain unchanged.
3. Quantitative Calculations: Balanced equations are necessary for stoichiometry, to calculate reactant/product masses in chemical processes.
Example: Unbalanced: H₂ + O₂ → H₂O | Balanced: 2H₂ + O₂ → 2H₂O

(a) 3H₂(g) + N₂(g) → 2NH₃(g)
(b) 2H₂S(g) + 3O₂(g) → 2H₂O(l) + 2SO₂(g)
(c) 3BaCl₂(aq) + Al₂(SO₄)₃(aq) → 2AlCl₃(aq) + 3BaSO₄(s)↓
(d) 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)

Note: (c) is a double displacement and precipitation reaction. (d) is a highly vigorous reaction where alkali metals react with water.

Balanced Equations:
(a) 2HNO₃ + Ca(OH)₂ → Ca(NO₃)₂ + 2H₂O
(b) 2NaOH + H₂SO₄ → Na₂SO₄ + 2H₂O
(c) NaCl + AgNO₃ → AgCl↓ + NaNO₃ (Already balanced)
(d) BaCl₂ + H₂SO₄ → BaSO₄↓ + 2HCl

Tip: Always check the number of atoms of each element on both sides after balancing.

(a) Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)
(b) Zn(s) + 2AgNO₃(aq) → Zn(NO₃)₂(aq) + 2Ag(s)
(c) 2Al(s) + 3CuCl₂(aq) → 2AlCl₃(aq) + 3Cu(s)
(d) BaCl₂(aq) + K₂SO₄(aq) → BaSO₄(s)↓ + 2KCl(aq)

Reaction Types: (a) Combination, (b) & (c) Displacement, (d) Double Displacement (Precipitation).

(a) 2KBr(aq) + BaI₂(aq) → 2KI(aq) + BaBr₂(s) — Double Displacement
(b) ZnCO₃(s) → ZnO(s) + CO₂(g) — Decomposition (Thermal)
(c) H₂(g) + Cl₂(g) → 2HCl(g) — Combination
(d) Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g) — Displacement (Metal-Acid)

Note: In (a), the state of barium bromide should be (aq) if soluble, but it's generally considered a precipitate in such reactions.

Exothermic Reaction: A reaction that releases heat energy to the surroundings. The reaction mixture becomes warm.
Examples: Combustion (burning), respiration, neutralisation (acid + base), formation of slaked lime (CaO + H₂O).

Endothermic Reaction: A reaction that absorbs heat energy from the surroundings. The reaction mixture becomes cold.
Examples: Decomposition of calcium carbonate (CaCO₃ → CaO + CO₂), photosynthesis, dissolving ammonium chloride in water.

Respiration is a biochemical process in living cells where glucose reacts with oxygen to produce carbon dioxide, water, and energy.
Chemical Equation: C₆H₁₂O₆(aq) + 6O₂(aq) → 6CO₂(aq) + 6H₂O(l) + energy
Since energy (in the form of ATP and heat) is released during this process, it is classified as an exothermic reaction. This energy is essential for all life processes.

Combination Reaction: Two or more substances combine to form a single product.
Decomposition Reaction: A single compound breaks down into two or more simpler substances.

They are opposites because one builds up (combination) and the other breaks down (decomposition).
Example:
Combination: CaO(s) + H₂O(l) → Ca(OH)₂(aq) (Slaked lime formation)
Decomposition: 2HgO(s) → 2Hg(l) + O₂(g) (Mercuric oxide decomposition)

Decomposition by Heat (Thermal Decomposition):
2Pb(NO₃)₂(s) → 2PbO(s) + 4NO₂(g) + O₂(g)

Decomposition by Light (Photochemical Decomposition):
2AgCl(s) → 2Ag(s) + Cl₂(g) (used in black & white photography)

Decomposition by Electricity (Electrolysis):
2H₂O(l) → 2H₂(g) + O₂(g) (Electrolysis of water)

Displacement Reaction: A more reactive element displaces a less reactive element from its compound.
Example: Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

Double Displacement Reaction: Two compounds exchange ions to form two new compounds. Often results in a precipitate.
Example: Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s)↓ + 2NaCl(aq)

Key Difference: Displacement involves one element and one compound; Double displacement involves two compounds exchanging partners.

Reaction: 2AgNO₃(aq) + Cu(s) → Cu(NO₃)₂(aq) + 2Ag(s)

Copper is more reactive than silver and displaces silver from silver nitrate solution, leading to the formation of silver metal (which can be recovered) and copper nitrate solution.

Precipitation Reaction: A double displacement reaction in which an insoluble solid (called precipitate) separates from the solution.

Example 1:
AgNO₃(aq) + NaCl(aq) → AgCl(s)↓ + NaNO₃(aq)
(White precipitate of silver chloride)

Example 2:
Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s)↓ + 2KNO₃(aq)
(Yellow precipitate of lead iodide)

(a) Oxidation: Gain of oxygen (or loss of hydrogen).
Example 1: 2Cu + O₂ → 2CuO (Copper gains oxygen)
Example 2: CH₄ + 2O₂ → CO₂ + 2H₂O (Methane gains oxygen)

(b) Reduction: Loss of oxygen (or gain of hydrogen).
Example 1: CuO + H₂ → Cu + H₂O (Copper oxide loses oxygen)
Example 2: ZnO + C → Zn + CO (Zinc oxide loses oxygen)

Note: Oxidation and reduction always occur together (redox reactions).

Element X: Copper (Cu)
Black Compound: Copper(II) Oxide (CuO)

Explanation: Copper metal (shiny brown) on heating reacts with oxygen to form black copper oxide.
Reaction: 2Cu(s) + O₂(g) → 2CuO(s)

Paint is applied to iron articles to prevent rusting (corrosion).

Reason: Paint forms a protective coating that isolates the iron surface from moisture and air (oxygen), the two essential components required for rusting (formation of hydrated iron(III) oxide, Fe₂O₃.xH₂O).

Nitrogen gas is inert (unreactive) and prevents oxidation of oils and fats.

Reason: When oils/fats are exposed to air, they get oxidised (rancidity) which changes their smell and taste. Flushing with nitrogen creates an oxygen-free environment, slowing down oxidation and extending the shelf life of the food items.

(a) Corrosion: The process of gradual deterioration (eating away) of metals due to reaction with moisture, air, acids, etc.
Example: Rusting of iron: 4Fe + 3O₂ + 2xH₂O → 2Fe₂O₃.xH₂O (Hydrated ferric oxide - rust)

(b) Rancidity: The oxidation of oils and fats present in food, resulting in bad smell and taste.
Example: When butter is left open for a long time, it develops an unpleasant odour and flavour due to rancidity.