Inorganic Chemistry
Inorganic chemistry is concerned with the synthesis, structure, properties and reactions of compounds of all the elements except those carbon-hydrogen compounds normally classed as organic. It spans simple ionic salts, covalent acids, organometallic complexes, and the catalysts that drive much of modern industry.
A measure of the tendency of an atom in a molecule to attract a shared pair of electrons. Linus Pauling's scale runs from caesium (0.79) to fluorine (3.98). Differences in electronegativity determine bond polarity and ionic/covalent character.
Periodic trends
The modern periodic table arranges elements by increasing atomic number Z, in seven rows (periods) and 18 columns (groups). Key trends:
- Atomic radius decreases across a period (increasing nuclear charge with same shell) and increases down a group (added shells).
- Ionisation energy (energy to remove an electron) increases across a period, decreases down a group.
- Electron affinity (energy released on adding an electron) shows similar but less monotonic trends; halogens have the most negative values.
- Metallic character increases down a group and decreases across a period.
Block organisation
| Block | Groups | Defining orbital | Examples |
|---|---|---|---|
| s | 1, 2 | s | Na, Mg |
| p | 13–18 | p | C, O, Cl, Ne |
| d | 3–12 | d (transition metals) | Fe, Cu, Pt |
| f | Lanthanides, actinides | f | Nd, U |
Chemical bonding
Three idealised bond types:
- Ionic — electrostatic attraction between cations and anions (e.g., NaCl). Forms between metals and non-metals with large electronegativity differences (>~1.7).
- Covalent — shared electron pairs (e.g., H₂, CH₄). Polar covalent when partners differ in electronegativity.
- Metallic — delocalised "sea" of valence electrons surrounding a lattice of cations; explains conductivity and ductility.
Two complementary models:
- Valence-bond theory (Pauling) — bonds form by overlap of atomic orbitals; hybridisation (sp, sp², sp³) explains molecular geometry.
- Molecular-orbital theory — atomic orbitals combine into bonding and antibonding MOs spanning the whole molecule. Successfully explains paramagnetism of O₂.
VSEPR theory predicts molecular shape by treating bonding pairs and lone pairs as point repellers:
| Steric number | Shape | Example |
|---|---|---|
| 2 | Linear | CO₂, BeCl₂ |
| 3 | Trigonal planar | BF₃ |
| 4 | Tetrahedral | CH₄ |
| 4 (1 lone pair) | Trigonal pyramidal | NH₃ |
| 4 (2 lone pairs) | Bent | H₂O |
| 5 | Trigonal bipyramidal | PCl₅ |
| 6 | Octahedral | SF₆ |
- Noble gases (Group 18) have full outer shells, hence their general inertness.
- Halogens (Group 17) are diatomic non-metals; reactivity decreases F > Cl > Br > I.
- Alkali metals (Group 1) react with water: 2 Na + 2 H₂O → 2 NaOH + H₂.
- Transition metals display variable oxidation states, coloured compounds, and catalytic activity.
s- and p-block highlights
- Hydrogen is unique: it can lose 1 electron to become H⁺ (like alkali metals) or gain 1 to become H⁻ (like halogens).
- Carbon (Group 14) shows catenation — its ability to form long chains underlies organic chemistry.
- Nitrogen is inert as N₂ (triple bond) but its compounds (NH₃, HNO₃, NO_x) are industrially crucial.
- Sulfur exists as S₈ rings; its oxides cause acid rain.
- Oxygen is the second most electronegative element; supports combustion and respiration.
Transition metals and coordination compounds
The d-block elements form a vast family of coordination compounds, in which a central metal ion is surrounded by ligands (electron-pair donors). Werner's coordination theory (1893) introduced the ideas of primary and secondary valences (now oxidation state and coordination number).
Common coordination numbers and geometries:
- 2 → linear (e.g., [Ag(NH₃)₂]⁺)
- 4 → tetrahedral (e.g., [Zn(NH₃)₄]²⁺) or square planar (e.g., [PtCl₄]²⁻)
- 6 → octahedral (e.g., [Fe(CN)₆]³⁻)
Crystal-field theory explains the colours, magnetic properties and stability of transition-metal complexes by considering how ligands split the d-orbitals into sets of different energy.
The "magic" of haemoglobin and chlorophyll lies in coordination chemistry: an iron(II) ion (in haem) or magnesium(II) (in chlorophyll) sits at the centre of a porphyrin ring with four nitrogen donors. Carbon monoxide binds Fe more strongly than O₂ — hence its toxicity.
Industrially important inorganic processes
| Process | Reaction | Use |
|---|---|---|
| Haber–Bosch | N₂ + 3 H₂ → 2 NH₃ (Fe catalyst) | Ammonia for fertilisers |
| Contact process | 2 SO₂ + O₂ → 2 SO₃ (V₂O₅ catalyst); SO₃ + H₂O → H₂SO₄ | Sulfuric acid |
| Ostwald | 4 NH₃ + 5 O₂ → 4 NO + 6 H₂O (Pt-Rh catalyst) | Nitric acid |
| Chlor-alkali | 2 NaCl + 2 H₂O → 2 NaOH + Cl₂ + H₂ | Cl₂, NaOH |
| Hall–Héroult | Al₂O₃ → 2 Al + 3/2 O₂ (electrolysis) | Aluminium metal |
Bioinorganic notes
Metals essential to life include Fe (oxygen transport, electron transfer), Mg (chlorophyll, ATP), Ca (bones, signalling), Zn (enzymes), Cu (cytochromes), Mn (photosynthesis), Mo (nitrogenase), Co (vitamin B12), and Ni (hydrogenases). Toxic heavy metals — Pb, Hg, Cd, As — cause harm largely by displacing essential metals or binding to sulfur-rich enzymes.