Chemical Reactions & Equations

What is a Chemical Reaction?

New substances with different properties are formed. Atoms rearrange — they are never created or destroyed.

🎬 Signs of a Chemical Reaction

📋 Writing Chemical Equations

Word equation: Magnesium + Oxygen → Magnesium oxide

Symbolic (skeleton):

Mg + O₂ → MgO

Balanced equation:

2Mg + O₂ → 2MgO

Balancing ensures atoms of each element are equal on both sides — Law of Conservation of Mass.

📌 State Symbols

Added in brackets to show physical state:

(s) = solid (l) = liquid
(g) = gas (aq) = aqueous

Example:

2H₂(g) + O₂(g) → 2H₂O(l)

Gas evolved: ↑ Precipitate: ↓

⚡ Energy in Reactions

Exothermic — Heat/light released
Burning, respiration, neutralisation

Endothermic — Energy absorbed
Photosynthesis, decomposition of CaCO₃

C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + Energy

Sign	Example	Observation
Change in colour	2FeSO₄ → Fe₂O₃ + SO₂ + SO₃	Green → reddish brown
Gas evolved (↑)	Zn + H₂SO₄ → ZnSO₄ + H₂↑	Bubbles, pop sound with flame
Precipitate (↓)	Pb(NO₃)₂ + 2KI → PbI₂↓ + 2KNO₃	Yellow solid forms
Change in temp	CaO + H₂O → Ca(OH)₂	Gets very hot (exothermic)
Change in state	2Ag + S → Ag₂S	Silver tarnishes (solid product)

🔗 Combination Reaction

Two or more reactants combine → single product. A + B → AB

🎬 Magnesium Burns in Oxygen

Magnesium + Oxygen

2Mg(s) + O₂(g) → 2MgO(s)

Exothermic

• Bright white flame
• White powder (MgO) formed
• Used in fireworks, photography
• MgO dissolves in water → Mg(OH)₂

Slaking of Lime

CaO(s) + H₂O(l) → Ca(OH)₂(aq)

Exothermic

• Large amount of heat released
• CaO = quicklime; Ca(OH)₂ = slaked lime
• Used in construction, water treatment
• Turns red litmus blue (basic)

Carbon Burns

C(s) + O₂(g) → CO₂(g)

Exothermic

Incomplete combustion:

2C + O₂ → 2CO (toxic!)

CO is colourless, odourless — causes carbon monoxide poisoning.

Formation of SO₃

2SO₂(g) + O₂(g) → 2SO₃(g)

Industrial process (Contact process for H₂SO₄).

SO₃ + H₂O → H₂SO₄

Sulphuric acid — the "king of chemicals".

💥 Decomposition Reaction

Single compound breaks into two or more simpler substances. AB → A + B

🎬 Electrolysis of Water

Thermal Decomposition

Heat causes breakdown:

CaCO₃(s) → CaO(s) + CO₂(g)

Endothermic

2Fe(OH)₃(s) → Fe₂O₃(s) + 3H₂O(g)

2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂

Brown NO₂ gas; yellow PbO formed.

Electrolytic Decomposition

Electric current breaks compounds:

2H₂O(l) → 2H₂(g) + O₂(g)

Endothermic

• H₂ at cathode (−): 2 vol
• O₂ at anode (+): 1 vol
• Ratio H₂:O₂ = 2:1 by volume

Photodecomposition

Light energy causes decomposition:

2AgCl(s) → 2Ag(s) + Cl₂(g)

2AgBr(s) → 2Ag(s) + Br₂(g)

Endothermic

• Used in black-and-white photography
• AgCl: white → grey in sunlight
• Sun cream: similar principle

Heating FeSO₄

2FeSO₄ → Fe₂O₃ + SO₂↑ + SO₃↑

Endothermic

Colour change: Green → Brown
Smell: Burning sulphur (pungent)
This is used as a test to identify FeSO₄.

↔️ Displacement Reactions

More reactive element displaces less reactive one from its compound.

🎬 Iron Nail in Copper Sulphate Solution

Single Displacement

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

Blue solution → green
Iron nail coated with red-brown copper

Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂↑

Zinc displaces hydrogen — H₂ gas burns with pop sound.

Double Displacement

Exchange of ions between two compounds:

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

White precipitate of BaSO₄

Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂↓ + 2KNO₃(aq)

Yellow precipitate of lead iodide.

Reactivity Series

K > Na > Ca > Mg > Al > Zn
> Fe > Ni > Sn > Pb > H
> Cu > Hg > Ag > Au

Higher element displaces lower.
Metals above H displace H from acids.

Neutralisation

Acid + Base → Salt + Water

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

Exothermic

Used in antacid tablets (Mg(OH)₂ neutralises excess HCl in stomach).

🔥 Oxidation & Reduction (Redox)

Oxidation and reduction always occur simultaneously in a reaction.

🎬 CuO reduced by Hydrogen

Definitions

Oxidation:
• Gain of oxygen
• Loss of hydrogen
• Loss of electrons (OIL)

Reduction:
• Loss of oxygen
• Gain of hydrogen
• Gain of electrons (RIG)

CuO + H₂ Reaction

CuO(s) + H₂(g) → Cu(s) + H₂O(l)

CuO is reduced (loses O) → Cu
H₂ is oxidised (gains O) → H₂O

Colour change: Black → Red-brown
H₂ = reducing agent; CuO = oxidising agent

Corrosion

Slow oxidation of metals in moist air:

4Fe(s) + 3O₂(g) + xH₂O → 2Fe₂O₃·xH₂O

Iron → Rust (reddish-brown)

Prevention: Painting, galvanisation (Zn coating), alloying, electroplating

Rancidity

Oxidation of fats/oils in food → unpleasant smell and taste.

Fats + O₂ → Rancid compounds

Prevention:
• Antioxidants (BHA, BHT)
• Nitrogen flushing in packets
• Refrigeration
• Vacuum packaging

🔴 Corrosion of Iron — Visual

Fe (Iron)

Fresh metal

→

…wait…

After days

→

Fe₂O₃·xH₂O

Rust (reddish-brown)

⚖️ Balancing Chemical Equations

Law of Conservation of Mass — atoms are neither created nor destroyed.

🎬 Atom Count Visualiser

✏️ Practice: Balance these equations

1. Hydrogen + Oxygen → Water

H₂ + O₂ → H₂O

2. Iron + Water → Iron oxide + Hydrogen

Fe + H₂O → Fe₃O₄ + H₂

3. Aluminium + Copper chloride → Aluminium chloride + Copper

Al + CuCl₂ → AlCl₃ + Cu

Steps to Balance

1. Write the skeletal equation
2. List atoms of each element on both sides
3. Add coefficients (not subscripts)
4. Start with the most complex formula
5. Balance H and O last
6. Cross-check all atoms

Common Mistakes

❌ Changing subscripts (H₂O → H₃O)
❌ Forgetting diatomic molecules (O₂, H₂, N₂)
❌ Not updating all elements after adding coefficient

✅ Only change coefficients before formulae
✅ Diatomic: H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂