Chemistry Topic 7,8,9

Indicator equilibrium

HIn ⇌ H⁺ + In⁻

Indicator colour change

Occurs when pH = pKa

Choosing indicators

Strong acid + strong base → pH ≈ 7; Weak acid + strong base → pH 7–11; Strong acid + weak base → pH 3–7; Choose indicator whose range covers equivalence point

Titration curve features

Starting pH shows acid/base; Equivalence point = steep rise midpoint; Buffer region before equivalence; Half‑equivalence

Conductometric titration pattern

Conductivity decreases → minimum at equivalence → increases after

Oxidation vs reduction

Oxidation = loss of electrons; Reduction = gain of electrons

Oxidising vs reducing agents

Oxidising agent gets reduced; Reducing agent gets oxidised

Oxidation number rules

Element = 0; Ion = charge; F = –1; H = +1 (hydrides = –1); O = –2 (peroxides = –1); Sum = charge or 0

Balancing redox (half‑equations)

Write oxidation + reduction halves; Balance atoms except O & H; Balance O with H₂O; Balance H with H⁺; Balance charge with e⁻; Equalise electrons; Add and cancel species

Galvanic cell basics

Spontaneous; Anode = oxidation (negative); Cathode = reduction (positive); Electrons flow anode → cathode; Salt bridge maintains charge balance

Cell diagram example

Zn(s) | Zn²⁺ || Cu²⁺ | Cu(s)

E°cell formula

E°cell = E°cathode – E°anode

Standard hydrogen electrode

E° = 0.00 V; 1 M H⁺; 1 atm H₂; Pt electrode

Electrolytic cell basics

Non‑spontaneous; Requires external power; Anode = positive; Cathode = negative

Electrolysis products

Molten salts discharge ions directly; Aqueous solutions compete with water; Example

Faraday’s laws

q = n(e⁻) × F; q = I × t