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1.11Physical Chemistry

Rate Equations — Oxford AQA International A Level Chemistry Revision Notes & Topical Questions

15 min read 28 parts Lesson 11 of 13

📚 Oxford AQA International A Level Chemistry Hub / Topical Revision
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Complete PDF with figures and original page layoutPage 1 Oxford AQA International Your notes A Level (IAL) Chemistry Rate Equations Contents Rate Equations The Arrhenius Equation Determination of Rate Equation Measuring the Rate of a Reaction …

Part 1 of 7Rate of reaction

Definition
  • Definition the change in amount or concentration of a reactant or product per unit time
Unit
  • Unit mol dm-3 s-1
Measurement
  • Measurement measure the decrease in reactant concentration or the increase in product concentration over time

Part 2 of 7The rate equation

1Origin
  • Origin rate equations can only be found experimentally, not from the stoichiometric equation
2General form
  • General form for P + Q → products, rate = k[P]m[Q]n
    What the symbols mean: rate (mol dm−3 s−1) · k = rate constant · [ ] = concentration (mol dm−3) · each power is the order with respect to that reactant
    • (units depend on the overall order)
3
Terms
  • k is the rate constant (proportionality constant).
  • [P] and [Q] are reactant concentrations.
  • m and n are the orders, which can only be 0, 1 or 2.
4Products
  • Products their concentrations are not included, as they do not affect the rate
5First-order example
  • First-order example rate = k[D], because rate ÷ [D] is constant at 6.67 × 10-4 s-1 for [D] of 3.00, 2.00 and 1.00 mol dm-3

Part 3 of 7Nitric oxide and hydrogen

Reaction
  • 2NO(g) + 2H2(g) → N2(g) + 2H2O(g), with rate = k[NO]2[H2]
Method
  • Method keep one concentration constant to see the effect of the other
With [H2] constant
  • With [H2] constant the rate is proportional to [NO]2
With [NO] constant
  • With [NO] constant the rate is proportional to [H2]
Key point
  • H2 is not second order even though its coefficient is 2, because orders come from experiment, not from the equation

Part 4 of 7Order of reaction

Order = the power to which the concentration of a reactant is raised in the rate equation
Zero order
  • Changing the concentration has no effect on the rate.
  • The reactant does not appear in the rate equation.
First order
  • The rate is directly proportional to the concentration.
  • Doubling the concentration doubles the rate.
Second order
  • The rate is proportional to the square of the concentration.
  • Doubling the concentration increases the rate fourfold.
Overall order
  • Overall order the sum of the individual orders, so rate = k[NO]2[H2] is second order in NO, first order in H2 and third order overall

Part 5 of 7Using the rate equation

N2O5 decomposition
  • N2O5 decomposition rate = k[N2O5]
  • The order is 1.
  • Tripling the concentration triples the rate.
Bromate and bromide
  • Bromate and bromide rate = k[BrO3-][Br-][H+]
  • Each reactant is first order, so the overall order is 3.
  • Doubling [BrO3-] doubles the rate and halving [Br-] halves it, so there is no overall change.

Part 6 of 7Deducing orders from a table

  • Approach
    • use a graph or a table of results with calculation
      • (its shape shows the order)
1
Find two experiments in which only one reactant concentration changes.
2
Calculate how that concentration changed.
3
Calculate how the rate changed.
4
Decide the order for that reactant.
5
Repeat for each reactant.
2-bromo-2-methylpropane and hydroxide
  • Experiments 1 and 2: [(CH3)3CBr] doubles and the rate doubles, so first order.
  • Experiments 1 and 3: [OH-] doubles and the rate increases fourfold, so second order.
  • Rate = k[(CH3)3CBr][OH-]2.
Writing the equation
  • Writing the equation omit zero-order reactants, and omit the power 1 for first-order reactants
Caution
  • Caution read standard form values carefully

Part 7 of 7Calculating the rate constant

Method
  • Method rearrange the rate equation to k = rate ÷ , and use one set of initial values
    • (concentration terms)
Example
  • Example rate = k[Na2CO3][Cl-] with an initial rate of 4.38 × 10-6 mol dm-3 s-1, [Na2CO3] = 0.0250 and [Cl-] = 0.0125 mol dm-3
Result
  • Result k = 1.40 × 10-2 dm3 mol-1 s-1
Consistency
  • Consistency the other experiments give the same value of k
Units
  • Units they depend on the overall order, as for first order (s-1) and second order (dm3 mol-1 s-1)

Key terms

TermMeaning
Rate equationAn expression linking rate to reactant concentrations, rate = k[A]m[B]n.
Rate constantThe constant k in the rate equation.
Order of reactionThe power of a reactant concentration in the rate equation.
Overall orderThe sum of the orders of all reactants in the rate equation.
Initial rateThe rate at the start of a reaction, at known concentrations.

Avoid these mistakes

1
Using the stoichiometric coefficients as orders — orders are found by experiment.
2
Including product concentrations in the rate equation — they do not appear unless they affect the rate.
3
Saying doubling a second-order reactant doubles the rate — it quadruples it.
4
Quoting units of k without checking the overall order — units depend on the order.

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