AS Chapter 10 · Periodicity

Chemical Periodicity

The Periodic Table arranges elements in order of increasing atomic number (proton number). A period is a horizontal row; a group is a vertical column. Elements in the same group have the same number of outer (valence) electrons and show similar chemical properties.

1.1 Atomic Radius

The atomic radius is half the distance between the nuclei of two adjacent, identical atoms that are bonded together (covalent radius) or in contact (van der Waals radius).

Trends Across Period 3 (Na → Ar)

↓ Atomic radius decreases across a period. Nuclear charge increases but electrons are added to the same shell → greater attraction pulls electrons closer.

↑ Atomic radius increases down a group. Extra electron shells are added; shielding effect increases → outer electrons are further from the nucleus.

1.2 First Ionisation Energy

The first ionisation energy (IE₁) is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous unipositive ions under standard conditions.

X(g) → X⁺(g) + e⁻ ΔH = IE₁

Trend Across Period 3

General increase Na → Ar: increasing nuclear charge, same shielding → electrons held more tightly.

Exceptions:

Al < Mg — Al's outer electron is in a 3p orbital (higher energy, easier to remove) vs Mg's 3s
S < P — S has a paired electron in one 3p orbital; electron-electron repulsion makes it easier to remove

Trend Down a Group

IE₁ decreases down a group: increasing atomic radius and shielding → outer electrons are less strongly attracted to the nucleus.

1.3 Electronegativity

Electronegativity (Pauling scale) is the ability of an atom to attract the shared pair of electrons in a covalent bond towards itself.

Increases across a period (left → right): more protons, smaller atomic radius
Decreases down a group: larger atomic radius, more shielding
Most electronegative element: F (4.0)

1.4 Melting Points of Period 3 Elements

Element	Structure	Melting Point / °C	Explanation
Na	Giant metallic lattice	98	Weak metallic bonding (1 delocalised e⁻)
Mg	Giant metallic lattice	650	Stronger metallic bonding (2 delocalised e⁻)
Al	Giant metallic lattice	660	Strongest metallic bonding (3 delocalised e⁻)
Si	Giant covalent (diamond-type)	1414	Many strong Si–Si covalent bonds; high energy to break
P	Simple molecular (P₄)	44	Weak van der Waals forces between P₄ molecules
S	Simple molecular (S₈)	113	Weak van der Waals forces; S₈ larger than P₄ → slightly stronger
Cl	Simple molecular (Cl₂)	−101	Very weak van der Waals; small molecule
Ar	Monatomic	−189	Extremely weak van der Waals forces

1.5 Reactions of Period 3 Elements with Oxygen

Period 3 Oxides — Acid/Base Character

Oxides of metals (Na, Mg, Al) are generally basic. The oxide of the metalloid Si is weakly acidic. Oxides of non-metals (P, S, Cl) are acidic.

Basic ──────────────────────── Acidic

Na₂O → MgO → Al₂O₃ → SiO₂ → P₄O₁₀ → SO₃ → Cl₂O₇

Formation Equations

4Na(s) + O₂(g) → 2Na₂O(s) 2Mg(s) + O₂(g) → 2MgO(s) 4Al(s) + 3O₂(g) → 2Al₂O₃(s) Si(s) + O₂(g) → SiO₂(s) P₄(s) + 5O₂(g) → P₄O₁₀(s) S(s) + O₂(g) → SO₂(g) 2SO₂(g) + O₂(g) → 2SO₃(g) (catalyst, high T)

1.6 Reactions of Oxides with Water

Na₂O(s) + H₂O(l) → 2NaOH(aq) (strongly alkaline, pH ≈ 13) MgO(s) + H₂O(l) → Mg(OH)₂(aq) (slightly alkaline, pH ≈ 9–10) Al₂O₃ — does not dissolve in water (amphoteric oxide) SiO₂ — does not dissolve in water (weakly acidic oxide) P₄O₁₀(s) + 6H₂O(l) → 4H₃PO₄(aq) (acidic, pH ≈ 2) SO₂(g) + H₂O(l) → H₂SO₃(aq) (acidic, pH ≈ 2–3) SO₃(g) + H₂O(l) → H₂SO₄(aq) (strongly acidic, pH ≈ 1)

1.7 Amphoteric Character of Al₂O₃

Al₂O₃ is amphoteric — it reacts with both acids and bases:

Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l) (reacts as base) Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) → 2NaAl(OH)₄(aq) (reacts as acid)

1.8 Period 3 Chlorides and Their Reactions with Water

Chloride	Bonding	Reaction with H₂O	pH
NaCl	Ionic	Dissolves; neutral	7
MgCl₂	Ionic	Dissolves; slightly acidic (Mg²⁺ ions hydrolyse slightly)	6–7
AlCl₃	Covalent (polar)	[Al(H₂O)₆]³⁺ → [Al(OH)(H₂O)₅]²⁺ + H⁺ acidic	3
SiCl₄	Covalent	SiCl₄ + 2H₂O → SiO₂ + 4HCl steamy fumes	1
PCl₃	Covalent	PCl₃ + 3H₂O → H₃PO₃ + 3HCl	1–2
PCl₅	Covalent	PCl₅ + 4H₂O → H₃PO₄ + 5HCl	1

The ionic chlorides (NaCl, MgCl₂) dissolve without reacting with water. Covalent chlorides (AlCl₃, SiCl₄, PCl₃, PCl₅) hydrolyse — they react with water, producing acidic solutions and steamy fumes of HCl.

Key Summary — Period 3 Trends

Atomic radius: decreases →, increases ↓
IE₁: generally increases → (with exceptions at Al and S); decreases ↓
Electronegativity: increases →, decreases ↓
Metallic oxides → basic; non-metal oxides → acidic
Ionic chlorides → neutral/slightly acidic; covalent chlorides → acidic (hydrolyse)

AS A2 Chapter 11 · Group 2

Group 2 – The Alkaline Earth Metals

BeZ=4

MgZ=12

CaZ=20

SrZ=38

BaZ=56

Group 2 elements all have the outer electron configuration ns². They form +2 ions by losing both outer electrons. All are silver-white reactive metals.

2.1 Physical Trends Down Group 2

Property	Trend	Reason
Atomic radius	Increases	Extra electron shells added; shielding increases
Ionic radius (M²⁺)	Increases	Same reason as atomic radius
First IE	Decreases	Outer electrons further away; more shielding
Melting point	Generally decreases*	Weaker metallic bonding; larger, fewer electrons per unit volume
Reactivity	Increases	Easier to lose outer electrons; lower IE

2.2 Reactions with Water

Be(s) + H₂O → no reaction (Be is protected by an oxide layer) Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g) (very slow, cold water) Mg(s) + H₂O(g) → MgO(s) + H₂(g) (reacts with steam) Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g) (vigorous; fizzing) Sr(s) + 2H₂O(l) → Sr(OH)₂(aq) + H₂(g) (more vigorous) Ba(s) + 2H₂O(l) → Ba(OH)₂(aq) + H₂(g) (very vigorous)

Trend

Reactivity with water increases down the group: Mg barely reacts with cold water; Ba reacts vigorously. This is due to the decreasing ionisation energy.

2.3 Reactions with Oxygen

2Mg(s) + O₂(g) → 2MgO(s) 2Ca(s) + O₂(g) → 2CaO(s) 2Sr(s) + O₂(g) → 2SrO(s) 2Ba(s) + O₂(g) → 2BaO(s) (also forms BaO₂ at high pressure)

2.4 Reactions with Dilute Acids

Mg(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂(g) Ca(s) + 2HCl(aq) → CaCl₂(aq) + H₂(g)

2.5 Thermal Stability of Group 2 Carbonates and Nitrates

Polarising power is the ability of a cation to distort (polarise) the electron cloud of a nearby anion. It increases with higher charge density — i.e. smaller ionic radius and/or higher charge.

Explaining Thermal Stability via Polarising Power

When a Group 2 cation sits next to a carbonate or nitrate anion in the lattice, it attracts electron density away from the anion — distorting and weakening the C–O or N–O bonds. This makes decomposition easier.

Polarising Power Down Group 2

Cation	Ionic Radius / pm	Charge Density	Polarising Power
Mg²⁺	72	Highest	Strongest — greatly distorts CO₃²⁻ / NO₃⁻
Ca²⁺	100	High	Strong
Sr²⁺	126	Moderate	Moderate
Ba²⁺	142	Lowest	Weakest — barely distorts CO₃²⁻ / NO₃⁻

As cation size increases down the group, charge density and polarising power decrease → the anion is less distorted → more thermal energy is required to decompose it → thermal stability increases down the group.

Carbonates — Thermal Decomposition

MgCO₃(s) → MgO(s) + CO₂(g) (decomposes at ~300 °C) CaCO₃(s) → CaO(s) + CO₂(g) (decomposes at ~840 °C) SrCO₃(s) → SrO(s) + CO₂(g) (decomposes at ~1100 °C) BaCO₃(s) → BaO(s) + CO₂(g) (decomposes at ~1360 °C)

Mechanism — Why Mg²⁺ Makes MgCO₃ Least Stable

The small, highly charged Mg²⁺ ion exerts a strong electric field on the adjacent CO₃²⁻ ion. This polarises the CO₃²⁻ electron cloud, weakening the C–O bonds and destabilising the carbonate lattice. Only a small amount of thermal energy (~300 °C) is then needed to break it apart, releasing CO₂.

Ba²⁺ is much larger and has a much lower charge density. It barely distorts the CO₃²⁻ ion, so the C–O bonds remain strong and ~1360 °C is required for decomposition.

Key link: smaller cation → greater polarising power → greater distortion of CO₃²⁻ → weaker C–O bonds → lower decomposition temperature → less thermally stable.

Nitrates — Thermal Decomposition

2Mg(NO₃)₂(s) → 2MgO(s) + 4NO₂(g) + O₂(g) 2Ca(NO₃)₂(s) → 2CaO(s) + 4NO₂(g) + O₂(g) 2Ba(NO₃)₂(s) → 2BaO(s) + 4NO₂(g) + O₂(g)

Mechanism — Polarising Power and Nitrate Stability

The same principle applies to nitrates. The NO₃⁻ ion contains N–O bonds that are weakened when polarised by a nearby cation. Mg²⁺, with its high charge density, strongly distorts the NO₃⁻ electron cloud, significantly weakening the N–O bonds. This means Mg(NO₃)₂ decomposes at a relatively low temperature.

Moving down the group, the increasing cation size reduces polarising power, so the N–O bonds in the nitrate are less distorted and the thermal stability of the nitrate increases.

All Group 2 nitrates decompose to the metal oxide, NO₂ (brown gas), and O₂ — unlike Group 1 nitrates (except LiNO₃) which decompose only to the nitrite.

2.6 Solubility Trends

The solubility of an ionic compound depends on the balance between two competing energetic processes:
Lattice energy (LE) — energy released when gaseous ions form the solid lattice (always exothermic). Larger LE = harder to dissolve.
Hydration energy (ΔH_hyd) — energy released when gaseous ions are surrounded by water molecules (always exothermic). Larger ΔH_hyd = easier to dissolve.

ΔH_solution = ΔH_hyd(cation) + ΔH_hyd(anion) − Lattice Energy
If ΔH_solution is negative (or small and positive), the compound is soluble.

Key Principle — Which Changes Faster?

Both lattice energy and hydration energy decrease as the cation gets larger (lower charge density). The crucial question is: which decreases faster?

If LE decreases faster than ΔH_hyd → ΔH_solution becomes more negative → solubility increases down the group
If ΔH_hyd decreases faster than LE → ΔH_solution becomes more positive → solubility decreases down the group

Hydroxides — Solubility Increases ↓

Mg(OH)₂ — slightly soluble (sparingly)
Ca(OH)₂ — slightly soluble (limewater)
Sr(OH)₂ — soluble
Ba(OH)₂ — soluble

Sulfates — Solubility Decreases ↓

MgSO₄ — very soluble (Epsom salt)
CaSO₄ — slightly soluble (gypsum)
SrSO₄ — insoluble
BaSO₄ — insoluble (used in barium meal)

Hydroxides — Why Solubility Increases Down the Group

Lattice Energy vs Hydration Energy for M(OH)₂

The OH⁻ ion is small. In M(OH)₂ lattices, the lattice energy is dominated by the interaction between M²⁺ and the small OH⁻. As the cation grows larger down the group, the lattice energy decreases relatively quickly because the interionic distance increases significantly.

The hydration energy of M²⁺ also decreases down the group, but more slowly than the lattice energy (because the large OH⁻ hydration energy stays roughly constant and partially compensates).

Net effect: LE decreases faster than ΔH_hyd → ΔH_solution becomes less endothermic (or more exothermic) → hydroxides become more soluble down the group.

Hydroxide	Lattice Energy trend	ΔH_hyd trend	Net effect on solubility
Mg(OH)₂ → Ba(OH)₂	Decreases steeply ↓	Decreases gently ↓	Solubility increases ↑

Sulfates — Why Solubility Decreases Down the Group

Lattice Energy vs Hydration Energy for MSO₄

The SO₄²⁻ ion is large. In MSO₄ lattices, because SO₄²⁻ is large, the interionic distance is already large even for Mg²⁺. As the cation grows down the group, the lattice energy decreases only slowly — adding a larger cation to an already large anion makes little proportional difference to interionic distance.

However, the hydration energy of M²⁺ decreases steeply as the cation becomes larger and its charge density drops.

Net effect: ΔH_hyd decreases faster than LE → ΔH_solution becomes more endothermic → sulfates become less soluble down the group.

Sulfate	Lattice Energy trend	ΔH_hyd trend	Net effect on solubility
MgSO₄ → BaSO₄	Decreases gently ↓	Decreases steeply ↓	Solubility decreases ↓

Summary Comparison — The Key Insight

Compound	Anion size	What decreases faster down the group?	Solubility trend
M(OH)₂	Small (OH⁻)	Lattice energy	Increases ↑ (Mg → Ba)
MSO₄	Large (SO₄²⁻)	Hydration energy of M²⁺	Decreases ↓ (Mg → Ba)

The size of the anion is the deciding factor: a small anion (OH⁻) gives lattices whose LE changes dramatically with cation size; a large anion (SO₄²⁻) gives lattices that are relatively insensitive to cation size, so hydration energy dominates the trend.

2.7 Important Uses

Compound	Use
CaO (quicklime)	Neutralise acidic soils; making slaked lime
Ca(OH)₂ (slaked lime)	Neutralise soil acidity; water treatment; whitewash
CaCO₃ (limestone)	Building material; manufacture of glass; CaO production
MgO	Refractory lining in furnaces; antacids
Mg(OH)₂	Milk of magnesia (antacid)
BaSO₄	Barium meal X-ray contrast agent (insoluble — not absorbed)
MgSO₄·7H₂O	Epsom salts; laxative

2.8 Test for Sulfate Ions (SO₄²⁻)

Qualitative Test

Add dilute HCl (to remove CO₃²⁻ interference), then add BaCl₂(aq):

Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s) (white precipitate, insoluble in HCl)

Remember: sulfate solubility decreases down the group (Mg soluble → Ba insoluble); hydroxide solubility increases down the group (Mg insoluble → Ba soluble). This is a common exam question!

AS A2 Chapter 12 · Group 17

Group 17 – The Halogens

FZ=9

ClZ=17

BrZ=35

IZ=53

AtZ=85

Group 17 elements all have the outer electron configuration ns²np⁵. They need one more electron to complete their outer shell, forming X⁻ ions. They are strong oxidising agents.

3.1 Physical Properties

Halogen	State (25 °C)	Colour	Boiling Point
F₂	Gas	Pale yellow	−188 °C
Cl₂	Gas	Yellow-green	−34 °C
Br₂	Liquid	Red-brown	59 °C
I₂	Solid	Shiny grey-black (purple vapour)	184 °C

Boiling points increase down the group: larger molecules have more electrons → stronger van der Waals (London dispersion) forces.

3.2 Oxidising Power — Displacement Reactions

Trend: Oxidising Power Decreases Down Group 17

A halogen higher in the group can oxidise (displace) the halide ion of a halogen lower in the group.

Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq) (solution turns orange-brown) Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq) (solution turns brown; blue-black with starch) Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(aq) (solution turns brown) Cl₂ displaces Br⁻ and I⁻; Br₂ displaces I⁻ only; I₂ displaces neither

Why does oxidising power decrease?

The ability to gain an electron (E + e⁻ → E⁻) is related to electron affinity. Down the group, the atom is larger with more electron shells → the incoming electron is less strongly attracted to the nucleus. Also, bond dissociation enthalpy of X–X decreases (less energy needed to break the bond, which actually favours reactivity), but the dominant factor is the lower electron affinity.

3.3 Reactions with Sodium Hydroxide (NaOH)

Cold, dilute NaOH — Disproportionation

Cl₂(g) + 2NaOH(aq) → NaCl(aq) + NaOCl(aq) + H₂O(l) (cold, dilute) Cl₂(g) + 2OH⁻(aq) → Cl⁻(aq) + OCl⁻(aq) + H₂O(l) (ionic equation)

This is disproportionation: Cl is simultaneously oxidised (0 → +1 in OCl⁻) and reduced (0 → −1 in Cl⁻).

Hot, concentrated NaOH — Further Disproportionation

3Cl₂(g) + 6NaOH(aq) → 5NaCl(aq) + NaClO₃(aq) + 3H₂O(l) (hot, concentrated)

Cl is oxidised to ClO₃⁻ (chlorate(V), +5) and reduced to Cl⁻ (−1).

3.4 Reaction with Water — Chlorine

Cl₂(g) + H₂O(l) ⇌ HCl(aq) + HClO(aq) HClO = hypochlorous acid (chloric(I) acid) — weak acid

HClO is a weak acid but a powerful oxidising agent and bleach. It kills bacteria by oxidising cell components. This explains why chlorine is used in water purification.

3.5 Hydrogen Halides (HX)

Preparation from NaCl + Concentrated H₂SO₄

NaCl(s) + H₂SO₄(l) → NaHSO₄(s) + HCl(g) (cold) NaCl(s) + H₂SO₄(l) → Na₂SO₄(s) + 2HCl(g) (hot)

This works for HCl because Cl⁻ cannot reduce H₂SO₄. However, for HBr and HI, the concentrated H₂SO₄ is itself an oxidant:

NaBr(s) + H₂SO₄(l) → NaHSO₄(s) + HBr(g) initially, but then: 2HBr(g) + H₂SO₄(l) → Br₂(g) + SO₂(g) + 2H₂O(l) (Br⁻ oxidised to Br₂) 8HI(g) + H₂SO₄(l) → 4I₂(g) + H₂S(g) + 4H₂O(l) (I⁻ reduces H₂SO₄ further to H₂S)

To make HBr or HI, use phosphoric acid (H₃PO₄) instead of H₂SO₄ — it is not an oxidant.

3.6 Reducing Power of Halide Ions

Trend: Reducing Power Increases Down Group

F⁻ < Cl⁻ < Br⁻ < I⁻

Larger ions are more easily oxidised (electrons more easily removed).

Halide	With conc. H₂SO₄	Observation
Cl⁻	No reduction of H₂SO₄	Steamy HCl fumes
Br⁻	Reduces H₂SO₄ to SO₂	Steamy HBr + brown Br₂ fumes + SO₂ smell
I⁻	Reduces H₂SO₄ to S, H₂S	HI + violet I₂ vapour + yellow S + rotten egg (H₂S) smell

3.7 Tests for Halide Ions

Silver Nitrate Test

Acidify with dilute HNO₃ (to remove CO₃²⁻ and SO₄²⁻ interference), then add AgNO₃(aq):

Ag⁺(aq) + Cl⁻(aq) → AgCl(s) (white ppt; soluble in dilute NH₃) Ag⁺(aq) + Br⁻(aq) → AgBr(s) (cream ppt; soluble in conc. NH₃ only) Ag⁺(aq) + I⁻(aq) → AgI(s) (yellow ppt; insoluble in NH₃)

3.8 Uses of Chlorine and Its Compounds

Water purification — Cl₂ kills bacteria; but may form carcinogenic chloroorganic compounds
Bleaching — NaOCl in household bleach; CaOCl₂ (bleaching powder)
PVC manufacture — CH₂=CH₂ + Cl₂ → CH₂ClCH₂Cl → HCl + CH₂=CHCl (then polymerised)
Solvents, pesticides — chlorinated hydrocarbons

Chlorine gas is toxic — it was used as a chemical weapon in WWI. Always handle in a fume cupboard.

A2 Chapter 13 · Nitrogen & Sulfur

Nitrogen and Sulfur Chemistry

4.1 Nitrogen — Oxides and Their Properties

Oxide	Name	Oxidation State of N	Properties
N₂O	Dinitrogen oxide (nitrous oxide)	+1	Colourless gas; anaesthetic ("laughing gas")
NO	Nitrogen monoxide (nitric oxide)	+2	Colourless gas; free radical; toxic; reacts with O₂ → NO₂
NO₂	Nitrogen dioxide	+4	Brown/reddish gas; toxic; acidic
N₂O₃	Dinitrogen trioxide	+3	Blue liquid; mixed anhydride of HNO₂
N₂O₄	Dinitrogen tetraoxide	+4	Colourless gas; dimer of NO₂
N₂O₅	Dinitrogen pentoxide	+5	White solid; anhydride of HNO₃

4.2 The Haber Process — Industrial Synthesis of Ammonia

The Haber process synthesises ammonia from nitrogen and hydrogen using a catalyst.

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ mol⁻¹

Optimum Conditions

Temperature: 450 °C (compromise — low T favours product yield but rate too slow; high T increases rate but shifts equilibrium left)
Pressure: 200 atm (high pressure favours forward reaction: 4 mol gas → 2 mol gas; but very high pressure is expensive and dangerous)
Catalyst: Iron (Fe) with promoters K₂O and Al₂O₃ (increases rate without affecting equilibrium position)
Removal of product: NH₃ is liquefied and removed continuously to shift equilibrium to the right (Le Chatelier's principle)

Source of Raw Materials

N₂ — from fractional distillation of liquid air
H₂ — from steam reforming of natural gas:

CH₄(g) + H₂O(g) ⇌ CO(g) + 3H₂(g) (Ni catalyst, 700–1000 °C) CO(g) + H₂O(g) ⇌ CO₂(g) + H₂(g) (water-gas shift reaction)

4.3 Reactions of Ammonia

As a Base

NH₃(g) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq) (weak base; Kₐ small) NH₃(g) + HCl(g) → NH₄Cl(s) (white smoke) NH₃(g) + HNO₃(g) → NH₄NO₃(s)

As a Ligand

Cu²⁺(aq) + 4NH₃(aq) → [Cu(NH₃)₄]²⁺(aq) (deep blue complex)

Catalytic Oxidation (Ostwald Process)

4NH₃(g) + 5O₂(g) → 4NO(g) + 6H₂O(g) (Pt/Rh catalyst, ~900 °C) 2NO(g) + O₂(g) → 2NO₂(g) 3NO₂(g) + H₂O(l) → 2HNO₃(aq) + NO(g) (NO recycled)

4.4 Nitric Acid — Properties and Uses

Nitric acid (HNO₃) is a strong acid and a powerful oxidising agent. Concentrated HNO₃ dissolves most metals (except Au, Pt, Al — passivation).

Cu(s) + 4HNO₃(conc.) → Cu(NO₃)₂(aq) + 2NO₂(g) + 2H₂O(l) (brown gas NO₂) 3Cu(s) + 8HNO₃(dilute) → 3Cu(NO₃)₂(aq) + 2NO(g) + 4H₂O(l) (colourless NO)

4.5 Sulfur — Allotropes

Rhombic sulfur (α-S₈): stable at room temperature; S₈ crown-shaped rings; yellow crystals
Monoclinic sulfur (β-S₈): stable 95.6–119 °C; also S₈ rings; needle-shaped crystals
Transition temperature: 95.6 °C (rhombic ⇌ monoclinic)

4.6 Oxides of Sulfur

S(s) + O₂(g) → SO₂(g) (burning sulfur; S oxidation state: 0 → +4) 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = −196 kJ mol⁻¹ (V₂O₅ catalyst, 450 °C)

4.7 The Contact Process — Industrial Synthesis of H₂SO₄

Steps

Burn sulfur (or roast sulfide ores): S + O₂ → SO₂
Catalytic oxidation: 2SO₂ + O₂ ⇌ 2SO₃ (V₂O₅ catalyst, 450 °C, 1–2 atm)
Absorb SO₃ in concentrated H₂SO₄ → oleum: SO₃ + H₂SO₄ → H₂S₂O₇
Dilute oleum: H₂S₂O₇ + H₂O → 2H₂SO₄

Note: SO₃ cannot be absorbed directly in water because a stable acid mist forms that is difficult to collect.

SO₃(g) + H₂O(l) → H₂SO₄(l) (theoretical, but forms mist in practice) SO₃(g) + H₂SO₄(l) → H₂S₂O₇(l) (oleum / pyrosulfuric acid) H₂S₂O₇(l) + H₂O(l) → 2H₂SO₄(l)

4.8 Properties of Sulfuric Acid

As a Strong Acid (dilute H₂SO₄)

H₂SO₄(aq) → 2H⁺(aq) + SO₄²⁻(aq) Mg(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂(g) CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l) Na₂CO₃(s) + H₂SO₄(aq) → Na₂SO₄(aq) + H₂O(l) + CO₂(g)

As an Oxidising Agent (concentrated H₂SO₄)

Cu(s) + 2H₂SO₄(conc.) → CuSO₄(s) + SO₂(g) + 2H₂O(l) C(s) + 2H₂SO₄(conc.) → CO₂(g) + 2SO₂(g) + 2H₂O(g)

As a Dehydrating Agent (concentrated H₂SO₄)

C₁₂H₂₂O₁₁(s) → 12C(s) + 11H₂O(l) (charring of sucrose) HCOOH(l) → CO(g) + H₂O(l) (dehydration of formic acid)

4.9 Acid Rain and Environmental Impact

Formation of Acid Rain

S(in fuels) + O₂(g) → SO₂(g) 2SO₂(g) + O₂(g) → 2SO₃(g) (in atmosphere) SO₃(g) + H₂O(l) → H₂SO₄(aq) (acid rain, pH < 5.6) N₂(g) + O₂(g) → 2NO(g) (lightning / car engines, high T) 2NO(g) + O₂(g) → 2NO₂(g) 3NO₂(g) + H₂O(l) → 2HNO₃(aq) + NO(g)

Effects: Kills aquatic life; damages forests; corrodes limestone buildings (CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂); human respiratory damage.

Solutions: Catalytic converters in cars; flue gas desulfurisation (CaO + SO₂ → CaSO₃); switching to low-sulfur fuels.

4.10 NO₂ as a Homogeneous Catalyst for SO₂ Oxidation

In the atmosphere, the direct oxidation of SO₂ to SO₃ by O₂ is extremely slow without a catalyst. Nitrogen dioxide (NO₂), which is also a pollutant produced by vehicle engines and lightning, acts as a homogeneous catalyst — it is in the same gas phase as the reactants and dramatically accelerates SO₂ oxidation, contributing significantly to acid rain formation.

The NO₂ Catalytic Cycle — Two Steps

NO₂ catalyses the oxidation of SO₂ to SO₃ through a two-step cycle in which NO₂ is consumed in step 1 and regenerated in step 2. It is not used up overall — the hallmark of a catalyst.

Oxidation of SO₂ by NO₂
NO₂(g) + SO₂(g) → NO(g) + SO₃(g)
NO₂ acts as the oxidising agent, donating its oxygen atom to SO₂ to form SO₃. NO₂ is reduced to NO (oxidation state: N goes from +4 → +2). SO₂ is oxidised to SO₃ (S: +4 → +6).

↓

Regeneration of NO₂ from NO
2NO(g) + O₂(g) → 2NO₂(g)
The NO produced in step 1 is rapidly re-oxidised by atmospheric O₂ back to NO₂. This regenerates the catalyst, allowing the cycle to repeat indefinitely.

↓

∑

Overall reaction (NO₂ cancels out)
2SO₂(g) + O₂(g) → 2SO₃(g)
Adding both steps together, NO₂ does not appear in the net equation — confirming its role as a catalyst. The O₂ from the atmosphere is the ultimate oxidant.

Why This Mechanism Works — Oxidation States

Species	Oxidation State of N or S	Change
NO₂ → NO (Step 1)	N: +4 → +2	N is reduced (gains electrons from SO₂)
SO₂ → SO₃ (Step 1)	S: +4 → +6	S is oxidised (loses electrons to NO₂)
NO → NO₂ (Step 2)	N: +2 → +4	N is re-oxidised by atmospheric O₂

Key Features of Homogeneous Catalysis Here

NO₂ and SO₂ are both gases — catalyst and reactant are in the same phase (homogeneous catalysis)
NO₂ provides an alternative reaction pathway with a lower activation energy than the direct O₂ + SO₂ reaction
NO₂ is regenerated at the end of the cycle — its concentration is unchanged overall
The variable oxidation state of nitrogen (+2 in NO ↔ +4 in NO₂) is what makes this catalytic cycle possible — a classic feature of transition-element-like chemistry (though N is a non-metal)

Connection to the Lead Chamber Process (Historical)

This NO₂ catalytic cycle is essentially the same chemistry exploited in the historical Lead Chamber Process for making sulfuric acid (before the Contact Process replaced it). In that industrial process, NO₂ was deliberately added to react with SO₂ and O₂ in large lead-lined chambers filled with steam:

NO₂(g) + SO₂(g) → NO(g) + SO₃(g) 2NO(g) + O₂(g) → 2NO₂(g) (NO recycled in the chamber) SO₃(g) + H₂O(g) → H₂SO₄(l) (dilute acid collected on chamber walls)

Understanding the atmospheric mechanism therefore also illuminates early industrial chemistry.

Exam tip: When asked why NO₂ is described as a homogeneous catalyst for SO₂ oxidation in the atmosphere, state that: (1) both NO₂ and SO₂ are in the same gas phase, (2) NO₂ is consumed in step 1 but regenerated in step 2, and (3) the overall equation shows no change in NO₂ — so it lowers Eₐ without being used up.

The NO₂ catalytic cycle is a key reason why air pollution from car exhaust (which produces both SO₂ from fuel sulfur and NO/NO₂ from engine combustion) accelerates acid rain formation far more effectively than either pollutant would alone.

4.11 Nitrogen Cycle — Brief Overview

Nitrogen fixation: N₂ → NH₃ (bacteria, lightning, Haber process)
Nitrification: NH₄⁺ → NO₂⁻ → NO₃⁻ (soil bacteria)
Denitrification: NO₃⁻ → N₂ (anaerobic bacteria)
Assimilation: plants take up NO₃⁻/NH₄⁺ to make proteins

A2 Chapter 14 · Transition Elements

Chemistry of Transition Elements

ScZ=21

TiZ=22

VZ=23

CrZ=24

MnZ=25

FeZ=26

CoZ=27

NiZ=28

CuZ=29

ZnZ=30

5.1 Definition and Properties

A transition element is a d-block element that forms at least one stable ion with an incompletely filled d sub-shell.

Note: Sc and Zn are NOT transition elements by this definition. Sc³⁺ has configuration [Ar] with empty 3d; Zn²⁺ has configuration [Ar]3d¹⁰ with a full 3d sub-shell.

Characteristic Properties of Transition Elements

Variable (multiple) oxidation states
Form coloured ions in solution
Act as catalysts (homogeneous and heterogeneous)
Form complex ions with ligands
Exhibit paramagnetism (unpaired d electrons)

5.2 Electronic Configurations

Element	Electron Configuration	Note
Sc	[Ar] 3d¹ 4s²
Ti	[Ar] 3d² 4s²
V	[Ar] 3d³ 4s²
Cr	[Ar] 3d⁵ 4s¹	Half-filled d is extra stable
Mn	[Ar] 3d⁵ 4s²
Fe	[Ar] 3d⁶ 4s²
Co	[Ar] 3d⁷ 4s²
Ni	[Ar] 3d⁸ 4s²
Cu	[Ar] 3d¹⁰ 4s¹	Fully filled d is extra stable
Zn	[Ar] 3d¹⁰ 4s²

When writing ion configurations, remove 4s electrons first: Fe²⁺ = [Ar] 3d⁶; Fe³⁺ = [Ar] 3d⁵

5.3 Variable Oxidation States

Element	Common Oxidation States	Examples
V	+2, +3, +4, +5	VO²⁺ (+4), VO₄³⁻ (+5)
Cr	+2, +3, +6	Cr³⁺ (green), Cr₂O₇²⁻ (+6)
Mn	+2, +4, +7	Mn²⁺ (pale pink), MnO₂ (+4), MnO₄⁻ (+7)
Fe	+2, +3	Fe²⁺ (pale green), Fe³⁺ (yellow/brown)
Co	+2, +3	Co²⁺ (pink), [Co(NH₃)₆]³⁺ (yellow)
Cu	+1, +2	Cu⁺ (colourless), Cu²⁺ (blue)

Variable oxidation states arise because the 3d and 4s electrons have similar energies, so different numbers can be involved in bonding.

5.4 Complex Ions and Ligands

A complex ion consists of a central metal ion surrounded by ligands — molecules or ions that donate a lone pair of electrons to the metal (coordinate/dative bonds).

Types of Ligands

Type	Example	Donor atoms
Monodentate	H₂O, NH₃, Cl⁻, CN⁻, OH⁻	1
Bidentate	en (ethane-1,2-diamine), C₂O₄²⁻ (oxalate)	2
Hexadentate	EDTA⁴⁻	6

Common Complex Ions — Shapes and Colours

Complex	Coordination Number	Shape	Colour
[Fe(H₂O)₆]²⁺	6	Octahedral	Pale green
[Fe(H₂O)₆]³⁺	6	Octahedral	Yellow/brown
[Cu(H₂O)₆]²⁺	6	Octahedral	Pale blue
[Cu(NH₃)₄(H₂O)₂]²⁺	6	Octahedral	Deep blue
[CuCl₄]²⁻	4	Tetrahedral	Yellow
[Co(H₂O)₆]²⁺	6	Octahedral	Pink
[CoCl₄]²⁻	4	Tetrahedral	Blue
[Ag(NH₃)₂]⁺	2	Linear	Colourless
[Ni(H₂O)₆]²⁺	6	Octahedral	Green
[MnO₄]⁻	4	Tetrahedral	Deep purple
[Cr(H₂O)₆]³⁺	6	Octahedral	Violet/green

5.5 Why Transition Metal Complexes Are Coloured

d-d Electron Transitions

In a transition metal complex, ligands split the 3d orbitals into two energy levels (crystal field splitting). An electron can absorb a photon of visible light to jump from the lower to the higher d level. The colour we see is the complementary colour of the wavelength absorbed.

The colour depends on: the metal ion, its oxidation state, the ligands present, and the coordination geometry.

Zn²⁺ and Sc³⁺ form colourless ions because Zn²⁺ (3d¹⁰) has a full d sub-shell and Sc³⁺ has an empty 3d — no d-d transitions are possible.

5.6 Ligand Substitution Reactions

Cu²⁺ + NH₃

[Cu(H₂O)₆]²⁺(aq) + 4NH₃(aq) → [Cu(NH₃)₄(H₂O)₂]²⁺(aq) + 4H₂O(l) pale blue → deep blue (inky blue)

Co²⁺ + Cl⁻

[Co(H₂O)₆]²⁺(aq) + 4Cl⁻(aq) → [CoCl₄]²⁻(aq) + 6H₂O(l) pink → blue (equilibrium: can reverse by adding water)

Ag⁺ + NH₃ (Tollens' reagent)

Ag⁺(aq) + 2NH₃(aq) → [Ag(NH₃)₂]⁺(aq)

5.7 Reactions of Aqueous Transition Metal Ions with NaOH and NH₃

Ion	+ NaOH(aq)	+ excess NaOH	+ NH₃(aq)	+ excess NH₃
[Fe(H₂O)₆]²⁺	Fe(OH)₂ green ppt	No change	Fe(OH)₂ green ppt	No change
[Fe(H₂O)₆]³⁺	Fe(OH)₃ rust-brown ppt	No change	Fe(OH)₃ rust-brown ppt	No change
[Cu(H₂O)₆]²⁺	Cu(OH)₂ pale blue ppt	No change	Cu(OH)₂ pale blue ppt	Deep blue [Cu(NH₃)₄(H₂O)₂]²⁺
[Cr(H₂O)₆]³⁺	Cr(OH)₃ grey-green ppt	[Cr(OH)₄]⁻ dark green soln (amphoteric)	Cr(OH)₃ grey-green ppt	[Cr(NH₃)₆]³⁺ purple soln
[Co(H₂O)₆]²⁺	Co(OH)₂ blue ppt → pink	No change	Co(OH)₂ blue/pink ppt	[Co(NH₃)₆]²⁺ straw/yellow soln
[Mn(H₂O)₆]²⁺	Mn(OH)₂ cream/buff ppt	No change	Mn(OH)₂ cream ppt	No change

[Fe(H₂O)₆]²⁺(aq) + 2OH⁻(aq) → Fe(OH)₂(s) + 6H₂O(l) (green ppt) [Fe(H₂O)₆]³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s) + 6H₂O(l) (rust-brown ppt) [Cu(H₂O)₆]²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s) + 6H₂O(l) (pale blue ppt) [Cr(H₂O)₆]³⁺(aq) + 3OH⁻(aq) → Cr(OH)₃(s) + 6H₂O(l) (grey-green ppt) Cr(OH)₃(s) + OH⁻(aq) → [Cr(OH)₄]⁻(aq) (excess NaOH; amphoteric)

5.8 Transition Metals as Catalysts

Process	Catalyst	Type
Haber process (NH₃)	Fe	Heterogeneous
Contact process (SO₃)	V₂O₅	Heterogeneous
Hydrogenation of alkenes	Ni	Heterogeneous
Catalytic converter	Pt, Pd, Rh	Heterogeneous
Decomposition of H₂O₂	MnO₂	Heterogeneous
Reaction of I⁻ and S₂O₈²⁻	Fe²⁺/Fe³⁺	Homogeneous
Wacker process (ethanol)	PdCl₂/CuCl₂	Homogeneous

Why Are Transition Metals Good Catalysts?

They have variable oxidation states → can participate in redox reactions as both oxidant and reductant in a catalytic cycle
They can adsorb reactants on their surface (heterogeneous catalysis), weakening bonds and lowering activation energy
Their partially filled d orbitals allow the formation of intermediate species

5.9 Chromium Chemistry — Chromate/Dichromate

2CrO₄²⁻(aq) + 2H⁺(aq) ⇌ Cr₂O₇²⁻(aq) + H₂O(l) yellow chromate(VI) ⇌ orange dichromate(VI) Add acid → Cr₂O₇²⁻ formed; add alkali → CrO₄²⁻ formed

Cr₂O₇²⁻ is a strong oxidising agent in acidic solution:

Cr₂O₇²⁻(aq) + 14H⁺(aq) + 6e⁻ → 2Cr³⁺(aq) + 7H₂O(l) orange → green

5.10 Manganese Chemistry

MnO₄⁻(aq) + 8H⁺(aq) + 5e⁻ → Mn²⁺(aq) + 4H₂O(l) purple → very pale pink (effectively colourless)

KMnO₄ is a powerful oxidising agent in acidic solution. It is self-indicating in redox titrations.

2MnO₄⁻(aq) + 5C₂O₄²⁻(aq) + 16H⁺(aq) → 2Mn²⁺(aq) + 10CO₂(g) + 8H₂O(l) KMnO₄ titration of Fe²⁺ or C₂O₄²⁻

MnO₄⁻(aq) + 4H⁺(aq) + 3e⁻ → MnO₂(s) + 2H₂O(l) (neutral/weakly acidic) MnO₄⁻(aq) + e⁻ → MnO₄²⁻(aq) (strongly alkaline)

5.11 Copper Chemistry

2Cu⁺(aq) → Cu(s) + Cu²⁺(aq) (disproportionation; Cu⁺ unstable in water) Cu²⁺(aq) + 2e⁻ → Cu(s)

Cu₂O (brick-red solid) is evidence for the existence of Cu⁺. Cu⁺ is stable in the solid state and in insoluble compounds, e.g. CuCl, Cu₂O, CuI.

5.12 Isomerism in Transition Metal Complexes

Complex ions can exhibit different types of isomerism — compounds with the same molecular formula but different structural arrangements. This is particularly rich in transition metal chemistry due to the variety of geometries and ligands possible.

Overview — Types of Isomerism in Complex Ions

Structural isomerism — different connectivity between atoms: ionisation isomerism, linkage isomerism, coordination isomerism
Stereoisomerism — same bonding connectivity, different arrangement in space: geometric (cis/trans) isomerism, optical isomerism

Geometric (cis/trans) Isomerism

In geometric isomerism, the same ligands can occupy different positions around the central metal ion due to restricted rotation about metal–ligand bonds. It occurs in square planar and octahedral complexes that contain two or more different types of ligand.

Square Planar Complexes — e.g. [Pt(NH₃)₂Cl₂]

Platinum(II) forms square planar complexes with coordination number 4. When two different pairs of ligands are present, two distinct geometric isomers arise:

cis-[Pt(NH₃)₂Cl₂] — two Cl ligands on the SAME side (adjacent positions, 90° apart) trans-[Pt(NH₃)₂Cl₂] — two Cl ligands on OPPOSITE sides (180° apart) cis isomer = cisplatin (anti-cancer drug); trans isomer = transplatin (inactive)

Cisplatin — A Medically Important Geometric Isomer

cis-[Pt(NH₃)₂Cl₂] (cisplatin) is a widely used anti-cancer drug. It enters the cell nucleus and cross-links guanine bases on adjacent DNA strands, distorting the double helix and triggering cancer cell death. The trans isomer (transplatin) cannot achieve this specific binding geometry — it is pharmacologically inactive. This is one of the most striking demonstrations in medicine of how spatial arrangement alone determines biological function.

Octahedral Complexes — e.g. [Co(NH₃)₄Cl₂]⁺

In octahedral [MA₄B₂] complexes, two geometric isomers exist:

cis-[Co(NH₃)₄Cl₂]⁺ — two Cl ligands adjacent (90° apart) (violet) trans-[Co(NH₃)₄Cl₂]⁺ — two Cl ligands directly opposite (180° apart) (green)

Complex Formula	Geometry	Geometric Isomers?	Type
[MA₆]	Octahedral	No	All ligands identical
[MA₅B]	Octahedral	No	Only one arrangement possible for B
[MA₄B₂]	Square planar	Yes	cis / trans
[MA₄B₂]	Octahedral	Yes	cis / trans
[MA₃B₃]	Octahedral	Yes	fac / mer
[MA₂B₂]	Tetrahedral	No	All positions equivalent in tetrahedral geometry

Facial (fac) and Meridional (mer) Isomerism

For octahedral [MA₃B₃] complexes, two geometric isomers exist based on whether the three identical ligands occupy one triangular face or span a meridional plane:

fac-[Co(NH₃)₃Cl₃] — three Cl on one triangular face; all Cl–Co–Cl angles = 90° mer-[Co(NH₃)₃Cl₃] — three Cl span a meridian; one Cl–Co–Cl angle = 180°

Optical Isomerism

Optical isomers (enantiomers) are pairs of molecules that are non-superimposable mirror images of each other. They rotate plane-polarised light in opposite directions: the (+) or d form rotates light clockwise; the (−) or l form rotates it anticlockwise. A complex is optically active when it has no plane of symmetry (it is chiral).

Optical isomerism is especially prevalent in octahedral complexes containing bidentate ligands such as en (ethane-1,2-diamine) or oxalate (C₂O₄²⁻), which wrap around the metal and create a chiral environment:

[Co(en)₃]³⁺ — tris(en)cobalt(III): exists as Δ (delta) and Λ (lambda) enantiomers [Cr(C₂O₄)₃]³⁻ — tris(oxalato)chromate(III): also optically active cis-[Co(en)₂Cl₂]⁺ — IS optically active (no plane of symmetry) trans-[Co(en)₂Cl₂]⁺ — NOT optically active (plane of symmetry through the two Cl and the metal)

Rules for Optical Activity in Metal Complexes

A complex is optically active only if it has no plane of symmetry (is chiral)
Octahedral [M(bidentate)₃] complexes are always chiral — both Δ and Λ enantiomers exist
cis-[M(bidentate)₂X₂] is chiral; trans-[M(bidentate)₂X₂] is NOT chiral
Tetrahedral [MABCD] with 4 different ligands is chiral
Square planar complexes are generally NOT optically active — the molecular plane acts as a plane of symmetry

Structural Isomerism

Ionisation Isomerism

These isomers have the same molecular formula but produce different ions in aqueous solution. They can be distinguished by simple precipitation tests:

[Co(NH₃)₅Br]SO₄ → [Co(NH₃)₅Br]²⁺(aq) + SO₄²⁻(aq) → white BaSO₄ precipitate with BaCl₂(aq); NO AgBr precipitate with AgNO₃(aq) [Co(NH₃)₅SO₄]Br → [Co(NH₃)₅SO₄]⁺(aq) + Br⁻(aq) → cream AgBr precipitate with AgNO₃(aq); NO BaSO₄ precipitate with BaCl₂(aq)

Linkage Isomerism

Arises with ambidentate ligands — ligands that possess two different possible donor atoms and can therefore coordinate to the metal through either atom:

[Co(NH₃)₅(NO₂)]²⁺ — N-bonded nitrite (nitro complex, N donor) yellow; thermodynamically stable [Co(NH₃)₅(ONO)]²⁺ — O-bonded nitrite (nitrito complex, O donor) red-brown; kinetic product, less stable

Other common ambidentate ligands: SCN⁻ (can bond via S or via N) and CN⁻ (can bond via C or via N).

In CIE A Level exams, geometric (cis/trans) isomerism is tested most frequently. Quick decision rule: (1) Does the complex have ≥2 different ligand types? (2) Is the geometry square planar or octahedral? If both conditions are met — geometric isomers exist. For optical isomerism — look for bidentate ligands wrapping around the metal and ask: can a plane of symmetry be drawn through the complex?

Complete Summary — Isomerism in Transition Metal Complexes

Category	Type	Key Condition	Classic Example
Stereoisomerism	cis / trans	Square planar or octahedral; ≥2 different ligands present	[Pt(NH₃)₂Cl₂], [Co(NH₃)₄Cl₂]⁺
	fac / mer	Octahedral [MA₃B₃] type	[Co(NH₃)₃Cl₃]
	Optical (enantiomers)	No plane of symmetry; bidentate ligands common	[Co(en)₃]³⁺, cis-[Co(en)₂Cl₂]⁺
Structural isomerism	Ionisation	Different ions released in solution	[Co(NH₃)₅Br]SO₄ vs [Co(NH₃)₅SO₄]Br
Structural isomerism	Linkage	Ambidentate ligand bonds via different donor atom	[Co(NH₃)₅(NO₂)]²⁺ vs [Co(NH₃)₅(ONO)]²⁺

Summary — Key Features of Transition Metals

Definition: d-block element forming ≥1 ion with incomplete d sub-shell
Variable oxidation states → catalytic activity and redox chemistry
Coloured ions → d-d transitions in the visible spectrum
Complex ion formation → ligands donate lone pairs
Fe²⁺: green ppt; Fe³⁺: brown ppt; Cu²⁺: blue ppt; Cr³⁺: grey-green ppt; Co²⁺: pink/blue; Mn²⁺: cream ppt
Isomerism: geometric (cis/trans, fac/mer), optical (enantiomers), structural (ionisation, linkage)

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Chemical Periodicity

1.1 Atomic Radius

1.2 First Ionisation Energy

1.3 Electronegativity

1.4 Melting Points of Period 3 Elements

1.5 Reactions of Period 3 Elements with Oxygen

Formation Equations

1.6 Reactions of Oxides with Water

1.7 Amphoteric Character of Al₂O₃

1.8 Period 3 Chlorides and Their Reactions with Water

Group 2 – The Alkaline Earth Metals

2.1 Physical Trends Down Group 2

2.2 Reactions with Water

2.3 Reactions with Oxygen

2.4 Reactions with Dilute Acids

2.5 Thermal Stability of Group 2 Carbonates and Nitrates

Explaining Thermal Stability via Polarising Power

Mechanism — Why Mg²⁺ Makes MgCO₃ Least Stable

Mechanism — Polarising Power and Nitrate Stability

2.6 Solubility Trends

Hydroxides — Why Solubility Increases Down the Group

Sulfates — Why Solubility Decreases Down the Group

2.7 Important Uses

2.8 Test for Sulfate Ions (SO₄²⁻)

Group 17 – The Halogens

3.1 Physical Properties

3.2 Oxidising Power — Displacement Reactions

Why does oxidising power decrease?

3.3 Reactions with Sodium Hydroxide (NaOH)

Cold, dilute NaOH — Disproportionation

Hot, concentrated NaOH — Further Disproportionation

3.4 Reaction with Water — Chlorine

3.5 Hydrogen Halides (HX)

Preparation from NaCl + Concentrated H₂SO₄

3.6 Reducing Power of Halide Ions

3.7 Tests for Halide Ions

3.8 Uses of Chlorine and Its Compounds

Nitrogen and Sulfur Chemistry

4.1 Nitrogen — Oxides and Their Properties

4.2 The Haber Process — Industrial Synthesis of Ammonia

Source of Raw Materials

4.3 Reactions of Ammonia

As a Base

As a Ligand

Catalytic Oxidation (Ostwald Process)

4.4 Nitric Acid — Properties and Uses

4.5 Sulfur — Allotropes

4.6 Oxides of Sulfur

4.7 The Contact Process — Industrial Synthesis of H₂SO₄

4.8 Properties of Sulfuric Acid

As a Strong Acid (dilute H₂SO₄)

As an Oxidising Agent (concentrated H₂SO₄)

As a Dehydrating Agent (concentrated H₂SO₄)

4.9 Acid Rain and Environmental Impact

4.10 NO₂ as a Homogeneous Catalyst for SO₂ Oxidation

Why This Mechanism Works — Oxidation States

4.11 Nitrogen Cycle — Brief Overview

Chemistry of Transition Elements

5.1 Definition and Properties

5.2 Electronic Configurations

5.3 Variable Oxidation States

5.4 Complex Ions and Ligands

Types of Ligands

Common Complex Ions — Shapes and Colours

5.5 Why Transition Metal Complexes Are Coloured

5.6 Ligand Substitution Reactions

Cu²⁺ + NH₃

Co²⁺ + Cl⁻

Ag⁺ + NH₃ (Tollens' reagent)

5.7 Reactions of Aqueous Transition Metal Ions with NaOH and NH₃

5.8 Transition Metals as Catalysts

Why Are Transition Metals Good Catalysts?

5.9 Chromium Chemistry — Chromate/Dichromate

5.10 Manganese Chemistry

5.11 Copper Chemistry

5.12 Isomerism in Transition Metal Complexes

Geometric (cis/trans) Isomerism

Square Planar Complexes — e.g. [Pt(NH3)2Cl2]

Octahedral Complexes — e.g. [Co(NH3)4Cl2]+

Facial (fac) and Meridional (mer) Isomerism

Square Planar Complexes — e.g. [Pt(NH₃)₂Cl₂]

Octahedral Complexes — e.g. [Co(NH₃)₄Cl₂]⁺