Q: What does H2 Chemistry Notes: CORE IDEA 3, 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.
General rate law for reaction aA + bB -> products:
rate=k[A]m[B]n
k
: rate constant (temperature-dependent).
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 = \(k\)
\(\pu{mol.L-1.s-1}\)
1
rate = \(k [A]\)
\(\pu{s-1}\)
2
rate = \(k [A]^2\)
\(\pu{L.mol-1.s-1}\)
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.
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.
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).
When planning experiments (Paper 4), mention control of temperature (water bath), use of data loggers, and replicates for reliability.
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 Delta G but lower E_a .
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).
Kinetics requires a balance of algebra and chemical storytelling. Keep solving dataset questions and align your answers with the explanation templates found throughout https://eclatinstitute.sg/blog/h2-chemistry-notes.
