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Q: What does H2 Chemistry Notes: Topic 8 - Reaction Kinetics cover? A: Nail rate laws, experimental determination of orders, Arrhenius analysis, and catalysis narratives for Core Idea 3 (Reaction Kinetics).
Kinetics questions test both experimental planning and mathematical fluency.
This guide summarises rate-law theory, data analysis routines, and catalytic mechanisms tailored to the 2026 examination format.
Status: SEAB H2 Chemistry (9476, first exam 2026) syllabus and Chemistry Data Booklet last checked 2026-01-13. Core Idea 3 Topic 8 is assessed across Papers 1–3.
Quick revision box
What this topic tests: Rate laws, order determination, Arrhenius relationships, and catalyst reasoning.
Top mistakes to avoid: Treating coefficient as order by default; poor graph interpretation; superficial catalyst explanations.
20-minute sprint plan: 5 min order-method recap; 10 min initial-rate/graph questions; 5 min catalyst and Arrhenius explanation.
1 Rate Laws and Orders
General rate law for reaction aA + bB -> products:
m,n : orders with respect to A,B; determined experimentally.
Overall order = m+n.
1.1 Units of Rate Constant
Overall order
Rate expression
Units of k
0
rate = k
mol⋅L−1⋅s−1
1
rate = k[A]
s−1
2
rate = k[A]2
L⋅mol−1⋅s−1
Always deduce units by inspecting the rate equation provided.
2 Determining Orders Experimentally
2.1 Initial Rates Method
Conduct experiments varying one reactant concentration while keeping others constant.
Measure initial rates.
Compare rate ratios to concentration ratios:
If doubling [A] doubles rate → first order; quadruples rate → second order.
2.2 Continuous Monitoring
Plot concentration vs time data:
Zero order: linear decrease; gradient = −k.
First order:[A] decays exponentially; ln[A] vs time linear with gradient −k.
Second order:[A]1 vs time linear with gradient k.
Half-life t1/2 for first-order reactions is constant: t1/2=kln2.
3 Arrhenius Equation
Temperature dependence captured by:
k=Ae−RTEa
Taking natural logs:
lnk=lnA−REa⋅T1
Plot lnk vs 1/T; gradient = −REa, intercept = lnA.
Activation energy Ea in J⋅mol−1; convert to kJ⋅mol−1
4 Mechanisms and Rate-Determining Step (RDS)
Link overall rate law to mechanism:
Identify slow step (RDS); its molecularity often reflects orders.
For fast pre-equilibria, express intermediate concentrations using equilibrium constants, substitute into rate expression.
Ensure proposed mechanism matches stoichiometry.
Example: Reaction between NOX2 and CO. If rate law is rate = k[NOX2]X2, mechanism might involve dimerisation of NOX2 as slow step followed by fast reaction with CO.
Discuss adsorption-reaction-desorption sequence for heterogeneous catalysts, referencing surface area and poisoning effects.
In exam responses, pair catalyst explanations with an energy-profile sketch showing uncatalysed and catalysed pathways: label lower Ea, keep ΔH/ΔG unchanged, and describe adsorption → surface reaction → desorption for heterogeneous systems.
Hydrogen peroxide: catalase substrate commonly used in kinetics-rate demonstrations.
For catalase-linked rate questions, track how quickly HX2OX2 disappears (or OX2 appears) as temperature or catalyst loading changes.
6 Worked Example
Question:
The decomposition of NX2OX5 in gas phase follows first-order kinetics. At 298K, k=3.46×10−5s−1; at 318K, k=1.16×10−4s−1. Calculate the activation energy.
State activation energy as 48kJ⋅mol−1 to two significant figures.
7 Practical Contexts
Clock reactions: Measure time for colour change; use reciprocal time as rate proxy.
Gas collection: Monitor volume evolved vs time (link to mole concept).
Colorimetry: Track absorbance of coloured species to determine concentration (Beer-Lambert law).
Hydrolysis follow-up task: Aspirin hydrolysis can be treated as a practical rate-comparison system when concentration is tracked by titration or spectroscopy.
When planning experiments (Paper 4), mention control of temperature (water bath), use of data loggers, and replicates for reliability.
Aspirin: ester hydrolysis substrate for temperature-dependent kinetics comparisons.
8 Common Missteps
Assuming order equals stoichiometric coefficient without experimental evidence.
Forgetting to convert minutes to seconds when calculating rate constants.
Failing to justify why catalysts do not affect ΔG but lower Ea.
Mixing up rate and rate constant when temperature changes: k changes; order remains constant.
9 Quick Drills
Experimental data: When [A] doubles and rate quadruples while [B] stays constant, deduce orders and write rate law.
Plot ln[A] vs time for sample first-order data; extract k and calculate half-life.
Explain why poisoning of heterogeneous catalysts reduces rate, referencing adsorption theory and industrial examples (e.g. lead poisoning Pt in catalytic converters).
Common exam mistakes
Mistake: Reading off stoichiometric coefficients as reaction orders without experimental evidence - orders must always be determined from data, not from the balanced equation.
Mistake: Carrying rate constant units as mol⋅L−1⋅s−1 regardless of order - units depend on overall order and must be derived each time from the rate equation.
Mistake: Confusing the rate constant k with the rate of reaction - k is temperature-dependent but concentration-independent; rate changes when concentrations change even at fixed temperature.
Mistake: Plotting [A] vs time and concluding first-order from a smooth curve - the correct test is a linear ln[A] vs time graph with constant half-life, not a visual impression.
Mistake: Forgetting to convert temperature to Kelvin in the Arrhenius equation - using Celsius gives a completely wrong activation energy.
Mistake: Claiming that a catalyst changes the enthalpy change ΔH or equilibrium position - catalysts lower Ea only; ΔH and the position of equilibrium are unchanged.
Mistake: Giving only "provides a surface" for heterogeneous catalysis without describing the adsorption-reaction-desorption sequence - full mechanism marks require all three stages.
Frequently asked questions
Is reaction kinetics tested in Paper 1? Yes. Multiple-choice questions can ask you to identify correct rate expressions, read half-lives from graphs, or select the correct units for k. Always apply the experimental evidence rule - never assume order equals coefficient.
What is the difference between reaction rate and rate constant? Rate of reaction is the change in concentration per unit time and varies with concentration. The rate constant k is a proportionality constant that depends only on temperature (via the Arrhenius equation); it does not change when you alter concentration.
How do I know which graph to plot for determining reaction order? Plot [A] vs time (linear → zero order), ln[A] vs time (linear → first order), or 1/[A] vs time (linear → second order). For a first-order reaction, a constant half-life across the concentration–time curve is also diagnostic.
Does a catalyst affect the value of Kc? No. A catalyst speeds up both the forward and reverse reactions equally, so the equilibrium position and Kc are unchanged. Only a temperature change alters Kc.
Struggling with Reaction Kinetics? Our H2 Chemistry tuition programme covers this topic with structured practice, Paper 4 practical drills, and worked exam solutions.
