Reactivity (energetics, kinetics, equilibrium, redox)

Reactivity encompasses four interconnected themes that determine how and why chemical reactions occur.

Energetics studies energy changes in reactions. Enthalpy change (ΔH) measures heat energy transferred at constant pressure. Exothermic reactions (ΔH < 0) release energy, while endothermic reactions (ΔH > 0) absorb energy. Key equation: ΔH = H(products) - H(reactants). Hess's Law states that total enthalpy change is independent of pathway taken. Bond enthalpy is energy required to break one mole of bonds in gaseous state.

Kinetics examines reaction rates. Rate of reaction = change in concentration/time. The collision theory requires particles to collide with sufficient activation energy (Ea) and correct orientation. Catalysts lower Ea without being consumed, providing alternative reaction pathways.

Equilibrium applies to reversible reactions where forward and reverse rates become equal. At dynamic equilibrium, concentrations remain constant but reactions continue. Le Chatelier's Principle: systems respond to counteract changes in concentration, temperature, or pressure. Kc = [products]/[reactants] at equilibrium (concentrations raised to stoichiometric coefficients).

Redox involves electron transfer. Oxidation = electron loss (increase in oxidation state); Reduction = electron gain (decrease in oxidation state). Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain). Oxidation state rules: elements = 0, monatomic ions = charge, oxygen usually -2, hydrogen usually +1. Half-equations show electron transfer explicitly.

These concepts govern everyday chemistry around us.

Energetics in Action: Your body's metabolism is controlled enthalpy changes. Glucose combustion (ΔH = -2800 kJ/mol) releases energy for cellular work. Hand warmers use exothermic crystallization of sodium acetate, while instant cold packs exploit endothermic ammonium nitrate dissolution. Industrial ammonia synthesis via the Haber process must balance exothermicity with reaction kinetics.

Kinetics in Practice: Think of activation energy as a hill between reactants and products. Food refrigeration slows molecular collisions, reducing spoilage rates. Catalytic converters in vehicles use platinum/rhodium catalysts to accelerate harmful gas breakdown at lower temperatures. Enzyme catalysts (biological catalysts) enable reactions at body temperature that would otherwise require extreme conditions.

Equilibrium Applications: The Haber process (N₂ + 3H₂ ⇌ 2NH₃) demonstrates Le Chatelier's principle brilliantly. Increasing pressure shifts equilibrium toward fewer gas molecules (right), favoring ammonia production. Your blood maintains pH equilibrium through carbonic acid/bicarbonate buffer systems.

Redox Everywhere: Rusting is slow iron oxidation (Fe → Fe²⁺ + 2e⁻). Batteries operate via controlled redox reactions—zinc oxidizes at anode, copper reduces at cathode. Photosynthesis reduces CO₂ to glucose while oxidizing water. Bleach oxidizes colored molecules, breaking chromophores.

Analogy: Think of equilibrium as a busy two-way street where traffic (molecules) moves both directions at equal rates—the street looks stationary but cars continuously flow both ways.

Example 1: Enthalpy Calculation using Hess's Law

Calculate ΔH for: C(s) + ½O₂(g) → CO(g)

Given:

• C(s) + O₂(g) → CO₂(g), ΔH₁ = -394 kJ/mol
• CO(g) + ½O₂(g) → CO₂(g), ΔH₂ = -283 kJ/mol

Solution: Target equation = Equation 1 - Equation 2

ΔH = ΔH₁ - ΔH₂ = -394 - (-283) = -111 kJ/mol

Examiner note: Always check stoichiometry matches target equation. Show algebraic manipulation of equations clearly.

Example 2: Equilibrium Calculation

For N₂O₄(g) ⇌ 2NO₂(g), at equilibrium: [N₂O₄] = 0.042 mol/dm³, [NO₂] = 0.016 mol/dm³. Calculate Kc.

Solution: Kc = [NO₂]²/[N₂O₄] = (0.016)²/(0.042) = 2.56 × 10⁻⁴/0.042 = 6.10 × 10⁻³ mol/dm³

Examiner note: Include units for Kc—derived from concentration terms. Square the NO₂ concentration due to coefficient 2.

Example 3: Oxidation States

Determine oxidation state changes in: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O

Solution:

• Cr in Cr₂O₇²⁻: 2x + 7(-2) = -2, so x = +6
• Cr in Cr³⁺: +3
• Change: +6 → +3 = reduction (gain of 3e⁻ per Cr)

Examiner note: Always verify total charge balances. Reduction occurs when oxidation state decreases.