Inorganic Chemistry

Atomic Radius

• Down a group: Increases (more electron shells)
• Across a period: Decreases (increased nuclear charge)

Ionization Energy

Energy required to remove an electron:

• Increases across periods
• Decreases down groups
• Noble gases have highest IEs

Electronegativity

Tendency to attract electrons:

• Highest: Fluorine (4.0)
• Trend: Increases up and right
• Metals have low electronegativity

Electron Affinity

Energy change when electron is added. Generally more negative (exothermic) for nonmetals.

Metallic Character

Decreases across periods, increases down groups. Metals are good conductors, malleable, ductile.

Ionic Bonding

• Transfer of electrons from metal to nonmetal
• Forms crystal lattice structures
• High melting/boiling points
• Conduct electricity when molten or dissolved
• Example: NaCl, MgO

Covalent Bonding

• Sharing of electron pairs
• Can be polar or nonpolar
• Forms discrete molecules or network solids
• Example: H₂O, CO₂, diamond

Metallic Bonding

• Delocalized electrons in metal lattice
• Results in conductivity, malleability
• Examples: Cu, Fe, Al

VSEPR Theory

Valence Shell Electron Pair Repulsion - predicts molecular geometry:

• Linear: 180°
• Trigonal planar: 120°
• Tetrahedral: 109.5°
• Octahedral: 90°

Components

• Central Metal Ion: Usually transition metal
• Ligands: Electron pair donors (NH₃, H₂O, Cl⁻, CN⁻)
• Coordination Number: Number of ligands attached

Naming

1. Ligands (alphabetical) + metal
2. Use -ate suffix for anions
3. Use Roman numerals for metal oxidation state

Example: [Co(NH₃)₆]Cl₃ = Hexaamminecobalt(III) chloride

Coordination Geometries

• Linear: CN = 2
• Tetrahedral: CN = 4
• Square Planar: CN = 4 (common for d⁸ metals)
• Octahedral: CN = 6 (most common)

Isomerism

• Geometric: cis/trans, fac/mer
• Optical: Enantiomers (mirror images)
• Linkage: Different atoms of same ligand bind

Basic Concept

Ligands create an electric field that splits d-orbital energy levels. This explains color, magnetism, and geometry of transition metal complexes.

Octahedral Field

• d orbitals split into t₂g (lower) and eg (higher)
• Δₒ = crystal field splitting energy
• Strong field: Large Δ, low spin
• Weak field: Small Δ, high spin

Tetrahedral Field

• Inverted splitting: e (lower), t₂ (higher)
• Smaller splitting than octahedral (Δt ≈ 4/9 Δₒ)
• Usually high spin

Spectrochemical Series

Ligand strength (weak to strong):

I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < CN⁻ < CO

Applications

• Explains color of complexes (d-d transitions)
• Predicts magnetic properties
• Determines stability and geometry

Alkali Metals (Group 1)

• Soft, reactive metals
• +1 oxidation state
• Form ionic compounds
• Examples: Na, K

Alkaline Earth Metals (Group 2)

• Harder, less reactive than Group 1
• +2 oxidation state
• Examples: Mg, Ca

Halogens (Group 17)

• Highly reactive nonmetals
• -1 oxidation state common
• Form diatomic molecules
• Examples: F₂, Cl₂, Br₂

Noble Gases (Group 18)

• Inert, full valence shells
• Very low reactivity
• Examples: He, Ne, Ar