Chemical Reactions and Equations

The Law

Law of Conservation of Mass: In any chemical reaction, the total mass of reactants equals the total mass of products. Atoms cannot be created or destroyed — only rearranged. This makes balancing compulsory.

Step 1 — Write the skeleton equation

Al + O₂ → Al₂O₃
Count atoms: Left: Al=1, O=2 | Right: Al=2, O=3 — Not balanced.

Step 2 — Balance Al

Right side has 2 Al → put coefficient 2 on left Al:
2Al + O₂ → Al₂O₃
Al: 2=2 ✓ | O: 2 ≠ 3 ✗

Step 3 — Balance O (LCM method)

Left O₂ gives O in multiples of 2; right Al₂O₃ gives O in multiples of 3.
LCM(2,3) = 6. Need 3O₂ on left and 2Al₂O₃ on right.
This also means Al on left must be 4Al:

Verification

Al: Left = 4 | Right = 2×2 = 4 ✓
O: Left = 3×2 = 6 | Right = 2×3 = 6 ✓

Classification

This is a chemical change — the browning is caused by the enzymatic oxidation of phenolic compounds (like catechol) in the apple when exposed to atmospheric oxygen (O₂), forming melanin pigments.

Four Pieces of Evidence

• New substance formed: Melanin (brown pigment) is a completely different compound from the original phenols.
• Irreversible change: The brown apple cannot be turned back to its original white state.
• Change in properties: Taste becomes slightly bitter; texture softens at the cut surface.
• Absorption of O₂: The process requires oxygen (removing O₂ by vacuum or coating with lemon juice slows it).

Simplified Equation

Balanced Equation

Product Name

Calcium hydroxide — also called Slaked lime. It is an alkali used to neutralise soil acidity.

Classification

Combination Reaction (two substances → one product) AND Exothermic Reaction (releases heat energy to surroundings).

Why Does It Get Hot?

In an exothermic reaction, the chemical bonds formed in the product (Ca–OH bonds in Ca(OH)₂) release more energy than is required to break the reactant bonds. This excess energy is released as heat, raising the temperature of the surroundings.

Observation 1 — Loss of Water

The light green crystals gradually lose their water of crystallisation. The crystals first shrink and become a white/pale powder (anhydrous FeSO₄).

Observation 2 — Colour Change

On further heating, the white FeSO₄ powder turns reddish-brown as Fe₂O₃ (ferric oxide) is formed. This is a clear sign of a new substance being produced.

Observation 3 — Pungent Smell

A sharp, acrid smell is noticed — this is due to SO₂ and SO₃ gases being evolved. A damp litmus paper turns red (acidic nature confirmed).

Equation

1. Thermal Decomposition

Energy Source: Heat (Δ)

Application: Industrial production of quicklime (CaO) used in cement manufacturing and for treating acidic soil.

2. Electrolytic Decomposition

Energy Source: Electric current

Application: Hydrogen fuel production; extraction of metals like aluminium and sodium from their molten salts (electrolytic refining).

3. Photolytic Decomposition

Energy Source: Light (hν — photons)

Application: Black & white photography; photochromic lenses (darken in sunlight due to AgCl decomposition, become clear indoors).

Fe in CuSO₄ — Reaction Occurs ✓

Iron is more reactive than copper, so Fe displaces Cu from its salt solution.

Observation: Blue solution gradually turns light green (FeSO₄ is pale green). Reddish-brown copper metal deposits on the iron surface.

Cu in ZnSO₄ — No Reaction ✗

Copper is LESS reactive than zinc. A less reactive metal cannot displace a more reactive one from its solution.

Observation: No change in the colour of solution or appearance of copper strip. The solution remains colourless.

Zn in FeSO₄ — Reaction Occurs ✓

Zinc is more reactive than iron, so Zn displaces Fe from its salt solution.

Observation: Pale green solution of FeSO₄ becomes colourless (ZnSO₄ solution). Grey iron deposits on the zinc strip.

Balancing the Thermite Equation

Skeleton: Fe₂O₃ + Al → Al₂O₃ + Fe

Fe: Left=2, Right=1 → put 2Fe on right. Al: Left=1, Right=2 → put 2Al on left. O: Left=3, Right=3 ✓

Check: Fe 2=2 ✓ | Al 2=2 ✓ | O 3=3 ✓

Why Al Displaces Fe

Aluminium is higher in the activity series than iron. More reactive metals have a greater tendency to lose electrons and form oxides. Al therefore has a stronger drive to form Al₂O₃ than Fe does to remain as Fe₂O₃, so it displaces iron.

Why Liquid Iron?

The thermite reaction is highly exothermic — it releases enormous heat, reaching temperatures of approximately 2500°C. Since the melting point of iron is only 1538°C, the iron produced is immediately in the liquid (molten) state.

Industrial Application

Railway Track Welding (Thermite Welding): The liquid iron flows into the gap between two rail sections and solidifies, creating a seamless, strong weld. No external power source is needed — the reaction is self-sustaining once ignited.

Molecular Equation

Precipitate: Barium sulphate (BaSO₄) — white solid, insoluble in water and dilute acids.

Net Ionic Equation

The spectator ions Na⁺ and Cl⁻ are not shown as they don't participate in the reaction.

Classification

Double Displacement Reaction — Ba²⁺ and Na⁺ exchange their anions (Cl⁻ and SO₄²⁻). Also classified as a Precipitation Reaction since an insoluble product (precipitate) forms.

OIL RIG Framework

OIL: Oxidation Is Loss of electrons
RIG: Reduction Is Gain of electrons

What is Oxidised?

Hydrogen (H₂) is oxidised.
H₂: oxidation state = 0 → in H₂O: +1. Increase in oxidation state = loss of electrons = Oxidation.

What is Reduced?

Copper (Cu in CuO) is reduced.
Cu in CuO: +2 → Cu metal: 0. Decrease in oxidation state = gain of electrons = Reduction.

Oxidising and Reducing Agents

Oxidising Agent: CuO — it oxidises H₂ (and gets reduced itself).
Reducing Agent: H₂ — it reduces CuO (and gets oxidised itself).

Colour Change

Black copper(II) oxide (CuO) → reddish-brown metallic copper (Cu).
Colour change: Black → Reddish-brown

Oxidation State Changes

Mn: In MnO₂ → +4 | In MnCl₂ → +2 | Change: +4 → +2 (decrease by 2)

Cl: In HCl → −1 | In Cl₂ → 0 | Change: −1 → 0 (increase by 1)

Oxidised and Reduced

Cl (in HCl) is Oxidised: oxidation state increases from −1 to 0.

Mn (in MnO₂) is Reduced: oxidation state decreases from +4 to +2.

Oxidising and Reducing Agents

Oxidising Agent: MnO₂ — it accepts electrons from HCl (gets reduced from +4 to +2).

Reducing Agent: HCl — it donates electrons (Cl gets oxidised from −1 to 0).

Observation

Black MnO₂ dissolves; yellowish-green Cl₂ gas evolves with a characteristic pungent smell. The solution becomes pale pink/colourless (MnCl₂ is very pale pink).

Three Reasons for Faster Corrosion Near Pool

• Higher humidity: Water vapour acts as a medium (electrolyte) for the electrochemical oxidation of Fe. Fe + O₂ + H₂O → rust is significantly accelerated.
• Chlorine ions (Cl⁻): Pool water contains chlorine. Cl⁻ ions break down the passive oxide layer on iron, exposing fresh metal to further attack.
• Temperature variations: Outdoor temperature fluctuations cause expansion/contraction, creating micro-cracks in any protective layer, exposing more iron surface.

Two Prevention Methods

1. Galvanisation (Zinc Coating): A layer of zinc (Zn) is applied to iron. Zn is more reactive — it acts as a sacrificial anode, oxidising preferentially and protecting iron even if the coating is scratched: Zn → Zn²⁺ + 2e⁻ (instead of Fe).

2. Electroplating with Chromium/Nickel: An inert, hard metal coat prevents O₂ and H₂O from contacting iron at all. No contact = no electrochemical cell = no rust.

Chemistry of Rancidity

Fats and oils contain unsaturated fatty acid chains (C=C double bonds). Atmospheric O₂ attacks these double bonds in a process called auto-oxidation:

This is an oxidation reaction producing hydroperoxides and then breakdown products that smell/taste rancid.

How Nitrogen Prevents Rancidity

Nitrogen (N₂) is a chemically inert gas — it does not react with fats or oils. By replacing air with N₂ in the packet, all oxygen is excluded. Without O₂, auto-oxidation cannot begin, and the fats remain fresh indefinitely (until the packet is opened).

Three Other Prevention Methods

• Antioxidants (BHA, BHT, Vit E): These molecules donate hydrogen atoms to free radicals, breaking the oxidation chain reaction before it damages fats.
• Refrigeration: Low temperature slows the rate of oxidation reaction (reaction rate decreases with temperature).
• Airtight packaging: Vacuum-sealed containers reduce O₂ availability, preventing contact between fat and oxygen.

The Apparent Contradiction Explained

The second student's argument incorrectly assumes only the magnesium's mass should be considered. The Law of Conservation of Mass applies to the total mass of ALL reactants — not just one of them.

Balanced Equation

Here, Mg combines with O₂ from the air. Mass of MgO = Mass of Mg + Mass of O₂ consumed.

Numerical Example

If 24g of Mg reacts with 16g of O₂ → 40g of MgO is formed.
Total Reactant mass = 24 + 16 = 40g
Total Product mass = 40g ✓

The product is heavier than Mg alone because it incorporates the mass of oxygen from air.

Balanced Equation

Gases at Each Electrode

Cathode (–ve): Hydrogen gas (H₂) — H⁺ ions gain electrons here (reduction).
Anode (+ve): Oxygen gas (O₂) — OH⁻ ions lose electrons here (oxidation).

Volume Ratio

H₂ : O₂ = 2 : 1
From the equation, 2 moles H₂ and 1 mole O₂ are produced. By Avogadro's law, equal moles of gases occupy equal volumes → ratio is 2:1.

Why Add Dilute H₂SO₄ or NaOH?

Pure water is a very poor conductor of electricity (very few ions). Adding dilute H₂SO₄ or NaOH provides ions (H⁺, SO₄²⁻ or Na⁺, OH⁻) that carry the current through the solution, making electrolysis possible.

Testing the Gases

Testing H₂: Bring a burning splint near the tube — H₂ burns with a blue flame and a characteristic "pop" sound.

Testing O₂: Bring a glowing splint near the tube — O₂ relights the glowing splint (supports combustion).

2Na + Cl₂ → 2NaCl

Combination + Redox.
Two substances → one product (Combination). Also: Na: 0 → +1 (oxidised); Cl: 0 → −1 (reduced) — Redox occurs simultaneously.

2KClO₃ →^{Δ, MnO₂} 2KCl + 3O₂

Thermal Decomposition + Redox.
One compound → simpler products (Decomposition). Cl: +5 → −1 (reduced); O: −2 → 0 in O₂ (oxidised) — Redox also occurs.

Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag

Displacement + Redox.
Cu more reactive than Ag → Cu displaces Ag from solution (Displacement). Cu: 0 → +2 (oxidised); Ag: +1 → 0 (reduced) — Redox simultaneously.

NaOH + HCl → NaCl + H₂O

Double Displacement + Neutralisation.
Na⁺ (from NaOH) and H⁺ (from HCl) exchange their anions OH⁻ and Cl⁻ (Double Displacement). Acid + Base → Salt + Water defines it as Neutralisation. No change in oxidation states → NOT a redox reaction.