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periodicity

Periodic trends, oxidation states, transition metals & complex ions, color.

trends · groups 1 & 17 · oxidation states · transition metals & complexes

Table structure. Period number = highest occupied \(n\); group = valence-electron pattern.

Across a period. Radius decreases; IE and EN increase; metallic character decreases; oxides go basic → amphoteric → acidic.

Down a group. Radius and shielding increase; IE and EN decrease; metallic character increases.

Groups 1 & 17

Group 1. Reactivity increases down the group; forms \(\mathrm{M^+}\); basic oxides/hydroxides. With water: \(\ce{2M + 2H2O -> 2MOH + H2}\).

Group 17. Oxidizing power and reactivity decrease down the group; forms \(\mathrm{X^-}\). Displacement: the more reactive halogen oxidizes the less reactive halide.

Oxidation states

Assignment rules. Element 0; monoatomic ion = its charge; sum of OS = species charge. Group 1 +1, group 2 +2, Al +3, F −1; O usually −2 except peroxides −1 and \(\ce{OF2}\) +2; H +1 except metal hydrides −1; halogens usually −1 except with O/F.

Oxidation vs reduction. Oxidation = OS increases / loss of e− / gain of O / loss of H; reduction = OS decreases / gain of e− / loss of O / gain of H.

Transition metals & complexes

Definition & features. A transition element forms at least one ion with an incomplete d sublevel. Features: variable oxidation states, colored complexes, catalytic activity, complex-ion formation. Variable OS arises because \(4s\)/\(3d\) ionization energies are similar.

Ligands & geometry. Ligand = Lewis base donating a lone pair; coordination number = number of coordinate bonds. Common geometries: CN 2 linear; CN 4 tetrahedral or square planar; CN 6 octahedral. Examples: \(\ce{[Cu(H2O)6]^2+}\), \(\ce{[Fe(CN)6]^3-}\), \(\ce{[Ag(NH3)2]+}\).

Complex formation. General: \(\ce{M^n+ + xL <=> [ML_x]^n+}\). \(\ce{Cu^2+(aq) + 4NH3 <=> [Cu(NH3)4]^2+(aq)}\) deep blue; \(\ce{Fe^3+ + SCN- <=> [FeSCN]^2+}\) blood red.

Color. The ligand field splits the d orbitals; the complex absorbs \(\Delta E=hf\); observed color is complementary to the absorbed color. Strong-field ligands cause larger splitting; oxidation state, ligand identity, and geometry all change the color.