Chemistry
25 flashcards
~13 min
Reaction Kinetics
Understand how fast chemical reactions occur and what controls their speed with this comprehensive reaction kinetics flashcard deck. Kinetics is at the heart of chemistry, biology (enzyme kinetics), pharmacology (drug metabolism), and engineering (reactor desi...
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Understand how fast chemical reactions occur and what controls their speed with this comprehensive reaction kinetics flashcard deck. Kinetics is at the heart of chemistry, biology (enzyme kinetics), pharmacology (drug metabolism), and engineering (reactor design). This deck covers measuring rates of reaction, factors affecting rate (concentration, temperature, surface area, catalysts), the rate law and rate constant, orders of reaction (zero, first, second), half-lives, the Arrhenius equation, collision theory, transition state theory, and reaction mechanisms with rate-determining steps. Essential for AP Chemistry, A-Level, and university physical chemistry.
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The change in concentration of a reactant or product per unit time. Rate = Δ[concentration] / Δtime (mol/L/s or mol dm⁻³ s⁻¹). Rate decreases over time as reactants are consumed.
1. Concentration (more particles → more collisions). 2. Temperature (more energy → more successful collisions). 3. Surface area (more contact → more collisions). 4. Catalyst (lowers activation energy → more successful collisions).
Reactions occur only when reactant particles collide with sufficient energy (≥ activation energy) and with the correct orientation. Increasing any factor that raises collision frequency or energy increases reaction rate.
The minimum energy required for a collision to result in a reaction. Only particles with energy ≥ Ea upon collision will react. A lower Ea means more particles can react → faster rate.
A graph showing the distribution of kinetic energies among particles in a gas/liquid at a given temperature. The area under the curve beyond Ea represents the fraction of particles that can react. Increasing T shifts the curve right — more particles exceed Ea.
Rate approximately doubles for every 10°C rise in temperature (rule of thumb). More particles have kinetic energy ≥ Ea, and collision frequency increases. The Maxwell-Boltzmann curve shifts to higher energies.
Rate = k[A]ᵐ[B]ⁿ where k = rate constant, [A] and [B] = concentrations of reactants, m and n = orders of reaction with respect to each reactant. The rate law must be determined experimentally — it cannot be deduced from the balanced equation.
The power to which a reactant's concentration is raised in the rate law. Zero order: Rate independent of [A] — rate = k. First order: Rate ∝ [A] — rate = k[A]. Second order: Rate ∝ [A]² — rate = k[A]². Overall order = sum of all individual orders.
Using the initial rates method: Run experiments with varying concentrations. Compare how rate changes when one reactant's concentration is doubled (or halved). If rate doubles → 1st order. Rate quadruples → 2nd order. No change → zero order.
A proportionality constant in the rate law. k is specific to a reaction at a given temperature. It increases with increasing temperature. A large k = fast reaction. Units depend on the overall order of reaction.
Zero order: mol L⁻¹ s⁻¹. First order: s⁻¹. Second order: L mol⁻¹ s⁻¹. Third order: L² mol⁻² s⁻¹. Units are derived from: Rate (mol L⁻¹ s⁻¹) = k × [concentration]ⁿ.
The time taken for the concentration of a reactant to fall to half its initial value. First-order reaction: t½ = ln2/k = 0.693/k — constant, independent of initial concentration. Used in radioactive decay calculations.
k = Ae^(−Ea/RT) where k = rate constant, A = frequency factor (collision frequency with correct orientation), Ea = activation energy, R = 8.314 J/mol·K, T = temperature (K). Shows the exponential relationship between k and temperature.
Plot ln k vs 1/T — gives a straight line with slope = −Ea/R. Therefore: Ea = −slope × R. y-intercept = ln A. Used to calculate Ea from experimental rate constant data at different temperatures.
A substance that increases reaction rate without being consumed. It provides an alternative reaction pathway with lower activation energy. Does not change the thermodynamics (ΔG, ΔH) or the equilibrium position.
A catalyst in the same phase as the reactants. Example: H⁺(aq) catalyzing ester hydrolysis in aqueous solution. Often forms an intermediate with the reactant, which then breaks down to give products and regenerate the catalyst.
A catalyst in a different phase from the reactants. Usually a solid catalyst with gaseous or liquid reactants. Reaction occurs at the catalyst surface (adsorption). Examples: Fe in the Haber Process, V₂O₅ in the Contact Process, Pt in catalytic converters.
The slowest step in a multi-step reaction mechanism. It limits the overall rate of the reaction. The rate law reflects the concentrations of species involved in (or before) the rate-determining step.
A step-by-step sequence of elementary reactions showing exactly how reactants are converted to products. Each step is called an elementary step. The overall balanced equation is the sum of all steps.
A species produced in one step of a reaction mechanism and consumed in a subsequent step. It does not appear in the overall balanced equation. Different from a transition state (which cannot be isolated).
An unstable, high-energy arrangement of atoms at the top of the energy barrier between reactants and products. It is transient — it cannot be isolated. The energy difference between reactants and the transition state = activation energy Ea.
A larger surface area exposes more particles to collisions. Dividing a solid into smaller pieces increases surface area dramatically. Example: powdered calcium carbonate reacts faster with HCl than large lumps. Relevant in heterogeneous catalysis.
Rate: Changes during a reaction as concentrations change (usually decreases over time). Rate constant (k): Fixed for a given reaction at a given temperature — independent of concentration. Changing temperature changes k but not the rate law expression.
The catalysed pathway shows a lower activation energy (lower energy hump). If the mechanism involves multiple steps, there may be multiple humps, all lower than the uncatalysed pathway. The start and end energy levels (reactants and products) remain the same.
Enzymes are biological catalysts. At low substrate concentration: rate ∝ [substrate] (first order). At high substrate concentration: rate plateaus at Vmax (zero order — enzyme saturated). Km = substrate concentration at ½Vmax — measures enzyme-substrate affinity.