Reaction Kinetics

A reaction rate can be measured as the change in a reactant or product concentration per unit time. Reactant disappearance rates are reported positively: -Δ[A]/Δt; product appearance rates use +Δ[P]/Δt. A tangent to a concentration-time curve gives the instantaneous rate, while a chord gives an average rate over the interval. State which species is being measured because stoichiometric coefficients connect their different disappearance and appearance rates.

For rate = k[A]^m[B]ⁿ, m and n are the orders with respect to A and B; m + n is the overall order. The rate constant k is constant for a given reaction under specified temperature and catalytic conditions. Changing concentrations changes the rate through the concentration terms, not k, within the conditions where that rate law applies.

If both concentrations change, first use another pair to find one order, then divide out its known contribution. For example, if doubling both A and B raises the rate eightfold and B is first order, the remaining factor of four comes from A, which is second order. Check the final equation against every row, not just the pair used to derive it.

The half-life is the time for a reactant concentration to fall to half its value at the start of the interval. For first order, each equal time interval removes the same fraction, so half-life is independent of the starting concentration. Look for several successive halvings on a graph; a single halving interval alone cannot establish first-order behaviour.

When another reactant is maintained effectively constant, an observed simple decay can describe the dependence on the changing reactant under those conditions. State that constraint before attributing the entire reaction an overall order. Integrated rate equations are not required for this syllabus.

A mechanism is a sequence of elementary steps. The rate-determining step is the slow step that controls the overall rate under the specified conditions. For an elementary slow step involving A and B, the simplest collision model gives rate = k[A][B]. A later fast step can consume another reactant without that reactant appearing in the observed rate law.

A proposed elementary step cannot have an arbitrary concentration exponent unrelated to the species participating in that step. Conversely, an observed zero order does not mean the substance is absent from the reaction: its concentration may not control the slow process over the measured range. Always use the mechanism supplied rather than inventing a slow step from the overall coefficients.

To compare concentration effects, change one initial concentration while holding total volume, temperature and other reactant concentrations constant, often by replacing part of a solution volume with water. Use an appropriate initial gradient, because reactant concentrations change during a run. Keep solid particle size or surface area constant if a solid participates.

This is a valid theoretical planning method. The 9476 practical syllabus excludes gas collection by gas syringe or water displacement as a hands-on examination operation; interpreting supplied gas-volume data or discussing an appropriate method remains useful chemical reasoning.

The activation energy E_a is the minimum energy barrier that must be overcome along the reaction pathway. Colliding particles need sufficient energy and an appropriate orientation to react. At higher reactant concentration, more particles occupy a given volume, usually producing more frequent relevant collisions; the measured rate law tells how strongly rate responds.

The distribution begins at zero and has a long tail: particles do not all have the same energy, and no sharp maximum energy exists. Raising temperature shifts the population towards higher energies and greatly increases the fraction able to cross the barrier. Together with a moderate increase in collision frequency, this raises k and hence the rate at fixed concentrations. Do not say that heating lowers activation energy.

Adding more particles at the same temperature does not change their normalised energy distribution or the fraction above E_a. It increases the number of potentially successful collisions. A temperature change and a concentration change therefore operate differently even when both make a reaction faster.

A catalyst participates in intermediate steps and is regenerated by the end of its cycle. It accelerates the approach to equilibrium in both directions and does not change ΔH, the equilibrium constant or the equilibrium composition at the same temperature. It can be deactivated or poisoned in practice even though it is not consumed in the ideal overall equation.

NO and NO₂ cycle between forms and cancel from the overall equation. This gas-phase scheme illustrates homogeneous catalysis of atmospheric SO₂ oxidation; real atmospheric sulfur chemistry also includes other oxidants and aqueous pathways. Formation and hydration of sulfur oxides contributes to acidic deposition.

The uncatalysed reaction requires encounters between two negative reactants, I^- and S₂O₈^2-. The iron redox cycle offers a different pathway involving positive iron ions, with a lower effective activation barrier. Fe²⁺ is oxidised and then regenerated; its concentration may be small compared with the reactants.

The solid catalyst is in a different phase from the reacting gases. A larger exposed surface gives more accessible active sites, while a poison that binds strongly can block them. Binding must be strong enough to activate reactants but not so strong that products cannot leave. Catalysis permits a useful rate at a lower temperature than the uncatalysed route; it does not independently increase the equilibrium yield.

In a vehicle catalytic converter, gases contact solid precious-metal surfaces such as platinum and rhodium. Adsorption and surface reaction help remove nitrogen oxides, with carbon monoxide acting as a reducing agent. A representative reaction is 2NO(g) + 2CO(g) → N₂(g) + 2CO₂(g). Other catalytic reactions oxidise CO and unburnt hydrocarbons. The converter requires a suitable operating temperature; it is less effective immediately after a cold start.

Specificity has two aspects: which substrate binds effectively and which reaction is catalysed. A substrate that does not fit or interact appropriately is not efficiently converted. The lock-and-key model captures this selectivity; it is a simplified model rather than a claim that proteins are perfectly rigid.

Orders come from rates or an explicit elementary mechanism, not overall coefficients. A zero-order concentration-time plot is straight; first order has constant successive half-lives. Derive k units from the measured rate units and overall order.

Concentration changes collision frequency. Temperature changes collision frequency and the energetic fraction, increasing k. A catalyst changes the pathway and lowers the barrier while leaving the same-temperature energy distribution and equilibrium constant unchanged. Enzymes add active-site specificity and temperature/pH sensitivity.