Topic 6 studies how fast reactions occur and what factors control the rate. SL covers qualitative explanations using collision theory. HL introduces quantitative rate expressions, reaction orders, and the Arrhenius equation.
Rate = change in concentration / time. Collision theory: reactions occur when particles collide with sufficient energy (≥ Ea) and correct orientation. Factors increasing rate: higher temperature (faster particles, more exceed Ea), higher concentration/pressure (more frequent collisions), larger surface area, catalysts (lower Ea).
Graph shows distribution of molecular kinetic energies at a given temperature. Area under curve = total molecules. Area to the right of Ea = molecules with enough energy to react. At higher temperature: curve shifts right, flattens, and peak is lower — more molecules exceed Ea, so rate increases. A catalyst lowers Ea, so more molecules exceed it without changing temperature.
Catalysts increase rate by providing an alternative pathway with lower activation energy. Homogeneous: same phase as reactants (e.g., H⁺ in esterification). Heterogeneous: different phase (e.g., Fe in Haber process, V₂O₅ in Contact process). Catalysts are not consumed and do not change ΔH or equilibrium position — they only speed up the approach to equilibrium.
Rate = k[A]ᵐ[B]ⁿ where m, n are orders (determined experimentally, NOT from balanced equation). Overall order = m + n. Half-life: for first order, t₁/₂ = ln2/k (constant). Arrhenius equation: k = Ae^(−Ea/RT). Mechanisms: series of elementary steps; rate-determining step is slowest. The rate expression is consistent with the RDS.
The balanced equation shows the overall stoichiometry, not the mechanism. Most reactions occur in multiple steps, and the rate depends on the slowest (rate-determining) step. The order with respect to each reactant reflects its involvement in the RDS, which may differ from its stoichiometric coefficient. Order must be determined experimentally (e.g., initial rates method or graphical methods).
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