Kinetics — IAL Edexcel Modular (Chemistry) Revision Notes & Topical Questions

Section 1 of 40Rate of reaction and rate equations

Rate of reaction and rate equations

Rate of reaction

Rate of reaction the change in concentration of a reactant or product per unit time, in mol dm^-3 s^-1

• It can be followed by mass loss, gas volume, colorimetry, pH or electrical conductivity.

Rate equation

• Rate equation rate = k[A]^m[B]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)
• [A] and [B] are concentrations and m and n are the orders with respect to each reactant.
• It can only be found by experiment, not from the stoichiometric equation.

Species in the rate equation

• Species in the rate equation products and catalysts may appear, but intermediates do not

Section 2 of 40Orders and the rate constant

1Zero order

• Zero order changing the concentration has no effect on the rate
• The species is not in the rate equation.

2First order

• First order the rate is directly proportional to the concentration
• Doubling the concentration doubles the rate.

3Second order

• Second order the rate is proportional to the concentration squared
• Doubling the concentration quadruples the rate.

4Overall order

Overall order the sum of the orders; orders can be positive, negative or fractional

5Rate constant, k

• Rate constant, k found by rearranging the rate equation and substituting the data from one experiment
• k changes with temperature, and with a catalyst.
• All experiments give the same value of k at the same temperature.

Overall order	Rate equation	Units of k
0	rate = k	mol dm-3 s-1
1	rate = k[A]	s-1
2	rate = k[A][B]	dm3 mol-1 s-1
3	rate = k[A]2[B]	dm6 mol-2 s-1

Section 3 of 40Deducing orders from data

1

Take one reactant at a time.

2

Choose two experiments in which only that reactant's concentration changes.

3

Compare the change in rate: unchanged means zero order, doubled means first order, quadrupled means second order.

4

Combine the orders into the rate equation, omitting zero-order species and the power 1.

• Care
  • standard form values must be read carefully, and a graph may give the order of one reactant while the table gives the others

Section 4 of 40Graphs and half-life

Order	Concentration–time graph	Rate–concentration graph	Half-life
Zero	Straight line sloping down	Horizontal line	Successive half-lives decrease
First	Downward curve that levels off	Straight line through the origin	Constant
Second	Steeper curve	Upward curve	Successive half-lives increase

1Zero order

• Zero order the gradient of the concentration–time graph is the rate and equals k
• Rate = k.

2First order

• First order the half-life is constant and t_½ = 0.693 ÷ k (ln 2 ÷ k)
• The gradient of the rate–concentration line is k.

3Common confusion

• Common confusion the straight-line zero-order concentration–time graph is not the same as the straight-line first-order rate–concentration graph

Section 5 of 40Obtaining rate data

Principle

• Principle follow a property that changes during the reaction and is proportional to concentration

Colorimetry

• Colorimetry measures the light absorbed by a coloured species, with a calibration curve to convert readings to concentration
• It cannot be used for coloured precipitates because they scatter light.

Mass loss

• Mass loss suitable for dense gases such as CO₂ but not for H₂

Gas volume

• Gas volume measured with a gas syringe or an inverted measuring cylinder over water

Titration of samples

• Titration of samples the reaction is stopped by quenching before titration
• Quenching can use sodium hydrogencarbonate, or cooling or dilution.

Section 6 of 40Initial-rate method

Initial rate

• Initial rate the rate at t = 0, found from the gradient of the tangent to the concentration–time graph

Method

• Method several runs with different starting concentrations of one reactant at a time, with the others and the temperature constant

Clock reactions

• Clock reactions time, t, for a fixed visible change, so the initial rate is proportional to 1/t
• The assumption is that the rate does not change significantly before the measurement.
• The iodine clock gives a sudden blue-black colour with starch once thiosulfate is used up.
• A graph of 1/t against concentration is a straight line through the origin for a first-order reactant.
• Short times give a more accurate initial rate.

Section 7 of 40Continuous monitoring method

Continuous monitoring

• Continuous monitoring concentration is measured throughout one run to give a concentration–time graph
• Gas volume, mass lost or colorimetry can be used.

Rate at any point

• Rate at any point draw a tangent at that time and find its gradient

Order

• Order a constant half-life shows first order, and a straight line shows zero order

Advantage

• Advantage one run gives the full curve and several half-lives

Section 8 of 40Rate-determining step and mechanisms

1Rate-determining step

Rate-determining step the slowest step in a mechanism, which controls the overall rate

• A reactant in the rate-determining step appears in the rate equation, and the order is the number of molecules in that step.

2Mechanism rules

• Mechanism rules the steps must add up to the overall equation and be consistent with the rate equation

3Limits of kinetics

• Limits of kinetics it can suggest or disprove a mechanism but cannot prove it

4Intermediates

• Intermediates formed in one step and used up in a later step
• They do not appear in the rate equation or overall equation.

5Catalysts

• Catalysts can appear in the rate equation but not in the overall equation

Section 9 of 40SN1 and SN2 mechanisms

Feature	SN2	SN1
Halogenoalkane	Primary	Tertiary
Rate equation	rate = k[RX][Nu−]	rate = k[RX]
Steps	One step; the nucleophile attacks from the side opposite the halogen	Two steps: slow ionisation to a carbocation, then fast attack by the nucleophile
Molecularity	Bimolecular	Unimolecular
Stereochemistry	Inversion of configuration	Racemic mixture

1Secondary halogenoalkanes

• Secondary halogenoalkanes react by both mechanisms

2Tertiary SN1

• Tertiary SN1 the carbocation is stabilised by three alkyl groups, and the crowded carbon blocks backside attack

Section 10 of 40Iodination of propanone and the Arrhenius equation

Iodination of propanone and the Arrhenius equation

Iodination CH₃COCH₃ + I₂ → CH₃COCH₂I + HI

Iodination CH₃COCH₃ + I₂ → CH₃COCH₂I + HI catalysed by dilute acid

• Rate = k[CH₃COCH₃][H⁺], zero order in iodine.
• Iodine is not in the rate-determining step.

Arrhenius equation, k

Ae^−E_a/RT — links k, temperature and activation energy

• A higher temperature increases k, because a greater proportion of molecules have at least E_a.
• A higher E_a lowers k.

Arrhenius plot

• Arrhenius plot ln k = ln A − E_a/RT, so a graph of ln k against 1/T is a straight line
• The gradient is −E_a/R and the intercept is ln A.

Section 11 of 40More from the specification

1

Initial reaction rate in solution

• Initial reaction rate in solution colorimetric tracking of reactant depletion in aqueous solution yields initial rates to establish kinetic order.

2

Rate of reaction

• Rate of reaction change in concentration per unit time, read from the gradient of a concentration-time graph.

3

Initial rates

• Initial rates measuring the rate at the start with different concentrations shows how rate depends on each reactant.

4

Order

• Order a first-order reaction has a constant half-life, a second-order rate depends on concentration squared,
• such as for propanone with iodine or decomposition of dinitrogen pentoxide.

5

Rate equation

• Rate equation rate equals k times concentrations raised to their orders, found only from experiment.

6

Elementary steps

• Elementary steps the slowest step of the mechanism is the rate-determining step and matches the rate equation.

7

Elementary step

• A single stage in a reaction mechanism.

Section 12 of 40More on this topic

Mechanisms

• Mechanisms the rate-determining step explains the rate equation, and SN1 and SN2 differ in order and stereochemistry

Writing that a catalyst increases yield — it does not change the yield at equilibrium, only the time taken to reach it

Section 13 of 40Rate equations and orders

1Rate equation

• Rate equation rate = k[A]^m[B]ⁿ, linking rate to concentrations

  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)
• The order with respect to a reactant is found only from experiment, not from the balanced equation.
• An order can be zero, positive, negative or fractional.

2Zero order

• Zero order changing the concentration has no effect on the rate, so the species is not in the rate equation

3First order

• First order rate is proportional to concentration
• Doubling the concentration doubles the rate.

4Second order

• Second order rate is proportional to concentration squared
• Doubling the concentration quadruples the rate.

5Overall order

Overall order the sum of the orders; products and catalysts may appear in a rate equation but intermediates do not

Section 14 of 40Deducing orders from experiments

1

Identify two experiments in which only one reactant's concentration changes.

2

Compare the change in rate with the change in concentration.

3

Repeat for each reactant in turn.

4

Write the rate equation, leaving out zero-order species and the power 1.

Section 15 of 40Rate constant and its units

Calculating k

• Calculating k rearrange the rate equation and substitute the concentrations and rate from one experiment
• Every experiment gives the same value of k at one temperature.

Units of k

• Units of k found by substituting the units into the rearranged rate equation
• Zero ordermol dm^-3 s^-1.
• First orders^-1.
• Second orderdm³ mol^-1 s^-1.

Section 16 of 40Concentration–time graphs

Order	Shape	Half-life
Zero	Straight line sloping down; gradient is the rate and is constant	Successive half-lives decrease with time
First	Downward curve that levels off (exponential)	Constant throughout the reaction
Second	Steeper curve that falls quickly then flattens	Successive half-lives increase with time

1First order

• First order the time for the concentration to halve is always the same
• The rate falls in proportion to the concentration.

2Second order

• Second order the rate falls faster as reactant is used up, so each half-life is longer than the last

3Measuring half-lives

• Measuring half-lives find the time for concentration to fall from c to c/2, then from c/2 to c/4, and compare

Section 17 of 40Rate–concentration graphs

1Zero order

• Zero order a horizontal line, because the rate is independent of concentration and rate = k

2First order

• First order a straight line through the origin, because the rate is directly proportional to concentration
• The gradient equals k.

3Second order

• Second order an upward curve, because the rate is proportional to the square of the concentration
• A plot of rate against [A]² would be straight.

4Reaction progress

• Reaction progress the rate falls as reactant is used up when the order is above zero,
  • which is why concentration–time graphs for first and second order are curved

Section 18 of 40More from the specification

1

Rate-Determining Step

• In a multi-step reaction mechanism, the slowest elementary step acts as the rate-determining step,
• which means its molecularity dictates the overall rate equation.

2

Activation Energy and the Arrhenius Equation

• The Arrhenius equation (k = A e^(-Ea/RT)) quantifies how reaction rate constants depend exponentially on activation energy and absolute temperature.

3

Initial Rates Method

• Measuring the initial rate of product formation allows deduction of reactant orders before back-reactions occur.
  • (such as gas volume evolved or colorimetric iodine appearance)

4

Reaction Mechanism Validation

• A proposed reaction mechanism consists of individual elementary steps whose stoichiometric sum equals the overall balanced chemical equation.

5

Activation energy

• The minimum energy colliding particles need to react.

6

Arrhenius equation

• K = Ae^−E_a/RT, relating the rate constant to temperature and activation energy.

7

Clock reaction

• A reaction in which a sudden visible change marks when a fixed amount has reacted.

8

Nucleophilic substitution

• Replacement of an atom or group by a nucleophile.

9

Continuous monitoring

• Collecting data throughout a reaction.

10

Elementary step

• A single stage in a reaction mechanism.

11

Initial rate

• The rate at the start of the reaction, at t = 0.

12

Mechanism

• The sequence of steps by which reactants form products.

13

Volume

• The space taken up by an object.

14

Halogenoalkane

• An alkane in which hydrogen atoms are replaced by halogen atoms.

Section 19 of 40Techniques for following a reaction

Principle

• Principle follow a property that changes during the reaction and is proportional to the concentration of a reactant or product

Colorimetry

• Colorimetry a colorimeter measures the light absorbed by a coloured solution
• It cannot follow coloured precipitates because they scatter light.

Mass loss

• Mass loss the mass falls as gas escapes
• The gas must be dense enough, so CO₂ is suitable but H₂ is not.

Gas volume

• Gas volume measured with a gas syringe or an inverted measuring cylinder over water, for gases that are not soluble

Titration of samples

• Titration of samples withdrawn samples are quenched to stop the reaction before titrating
• Sodium hydrogencarbonate can quench an acid-catalysed reaction.

Section 20 of 40Initial-rate method

• Initial rate
  • the rate at the very start of the reaction (t = 0), before the concentrations change
  • Plot a concentration–time (or volume–time) graph and find the gradient of the tangent at t = 0.

1

Carry out a series of runs, changing the starting concentration of one reactant only.

2

Keep the other concentrations and the temperature constant.

3

Find the initial rate of each run.

4

Compare the initial rates to find the order with respect to each reactant.

Section 21 of 40Clock reactions

Clock reaction

• Clock reaction measures the time, t, for a specific visual change, such as a colour change or the appearance of a precipitate
• A fixed amount of reactant is used up in each run.

Key assumption

• Key assumption the rate does not change significantly between the start and the time of measurement
• The initial rate is then proportional to 1/t.

Iodine clock

• Iodine clock thiosulfate removes the iodine formed until it is used up, then starch gives a sudden blue-black colour

Analysis

• Analysis plot 1/t against concentration
• A straight line through the origin shows first order.
• Doubling the concentration doubling 1/t also shows first order.

Accuracy

• Accuracy a short time gives an initial rate closer to the true value, and the initial rate is only an estimate

Section 22 of 40Continuous monitoring

Continuous monitoring

• Continuous monitoring concentration is measured throughout the reaction to plot a concentration–time graph

Colorimetry

• Colorimetry first make a calibration curve of absorbance against concentration for standard solutions
• Choose a filter whose colour is complementary to the colour of the solution, giving the strongest absorbance.
• The reaction is started by adding the last reagent at a recorded time, then the mixture is transferred to the cuvette.

Preparing mixtures

• Preparing mixtures burettes of the reactants and distilled water give mixtures of known concentration at constant total volume

Measuring a rate

• Measuring a rate draw a tangent at the required time and find the gradient

Advantages

• Advantages one run gives the full curve and several half-lives, and the shape reveals the order directly

Section 23 of 40Sources of error and improvements

Sources of error

• Sources of error starting the stopwatch at the instant of mixing, judging the first appearance of colour, and temperature changes

Reducing uncertainty

• Reducing uncertainty use larger volumes and burettes, a colorimeter instead of the eye, and a thermostatted water bath

Percentage uncertainty

(absolute uncertainty ÷ measured value) × 100

• A burette reading has ±0.05 cm³, and a titre uses two readings, so ±0.10 cm³.

Control

• Control keep the temperature constant and mix reagents at a recorded start time

Section 24 of 40Choosing a method

1Reactions producing gas

• Reactions producing gas gas collection or mass loss

2Coloured species

• Coloured species colorimetry

3Reagents such as iodine

• Reagents such as iodine titration of quenched samples

4Quick initial-rate comparisons

• Quick initial-rate comparisons clock reactions

Section 25 of 40More from the specification

1

Initial rate method as an analytical means

• Initial rate method as an analytical means measuring initial reaction velocities across varying reactant concentrations provides a reliable means to determine rate equations.

2

Reaction Mechanisms and Elementary Steps

• Chemical reactions often occur through a sequence of elementary molecular steps known as the reaction mechanism,
• which together account for the stoichiometry of the overall balanced equation.

3

The Rate-Determining Step (RDS)

• The rate-determining step is the slowest elementary step in a reaction mechanism, governing the overall rate law and kinetic orders with respect to participating reactants.

4

The Arrhenius Equation

• The temperature dependence of the rate constant k is modeled by the Arrhenius equation k = A * e^(-Ea / RT),
• where A is the pre-exponential frequency factor and Ea is the activation energy.

5

Arrhenius Graphs and Activation Energy

• Plotting ln k against 1/T yields a linear graph of form y = mx + c with a negative gradient equal to -Ea / R,
• from which the activation energy of the reaction is calculated.

6

Activation energy

• The minimum energy colliding particles need to react.

7

Arrhenius equation

• K = Ae^−E_a/RT, relating the rate constant to temperature and activation energy.

8

Nucleophilic substitution

• Replacement of an atom or group by a nucleophile.

9

Elementary step

• A single stage in a reaction mechanism.

10

Mechanism

• The sequence of steps by which reactants form products.

11

Halogenoalkane

• An alkane in which hydrogen atoms are replaced by halogen atoms.

Section 26 of 40Mechanisms and the rate-determining step

1Mechanism

Mechanism a series of elementary steps, each describing the breaking and making of bonds in a single event

2Rate-determining step

Rate-determining step the slowest step, which controls the overall rate

• Species in the rate equation are involved in this step or in steps before it.
• The order with respect to each species equals the number of molecules of it in that step.

3Limits of kinetics

• Limits of kinetics kinetics can suggest a mechanism or disprove one, but cannot prove it

4Check

• Check the proposed steps must add up to the overall stoichiometric equation, otherwise the mechanism is wrong

Section 27 of 40Deducing a mechanism from a rate equation

1

Write the experimental rate equation.

2

Decide which species, and how many of each, are in the rate-determining step.

3

Propose a slow step consistent with this and fast steps for the rest.

4

Check that the steps sum to the overall equation and that intermediates cancel.

5

Check that any catalyst is regenerated.

Intermediates

• Intermediates appear in the mechanism but not in the rate equation or overall equation
• For NO₂ + CO → NO + CO₂ with rate = k[NO₂]², the slow step is 2NO₂ → NO + NO₃, followed by fast NO₃ + CO → NO₂ + CO₂.

Using a given mechanism

• Using a given mechanism the rate-determining step is the slow step, and species in it and before it give the rate equation
• Where an intermediate is formed from two molecules in an earlier fast step, count both molecules in the order.

Section 28 of 40SN2 mechanism

1SN2

• SN2 substitution, nucleophilic, with rate depending on two species
• Typical of primary halogenoalkanes.
• Rate = k[halogenoalkane][nucleophile].

2One step

• One step the nucleophile donates a pair of electrons to the δ+ carbon while the C–X bond breaks and X leaves as X⁻

3Transition state

• Transition state has partial bonds to both the nucleophile and the halogen

4Bimolecular

• Bimolecular two molecules are involved in the rate-determining step

5Stereochemistry

• Stereochemistry inversion of configuration, because the nucleophile attacks from the side opposite the leaving group

Section 29 of 40SN1 mechanism

• SN1
  • rate depends on one species, the halogenoalkane
  • Typical of tertiary halogenoalkanes.
  • Rate = k[halogenoalkane].

SN1 mechanism

The C–X bond breaks heterolytically in the slow, rate-determining step, giving a tertiary carbocation and X⁻.

The nucleophile attacks the carbocation in a fast second step.

Unimolecular

• Unimolecular only one molecule is involved in the rate-determining step

Why tertiary

• Why tertiary the carbocation is stabilised by three alkyl groups, and the crowded carbon blocks backside attack

Stereochemistry

• Stereochemistry a racemic mixture forms, because the planar carbocation can be attacked from either side

Section 30 of 40Which mechanism?

1Primary halogenoalkanes

• Primary halogenoalkanes react only by SN2

2Secondary halogenoalkanes

• Secondary halogenoalkanes react by both SN1 and SN2

3Tertiary halogenoalkanes

• Tertiary halogenoalkanes react only by SN1

4From the rate equation

• From the rate equation if the nucleophile is not in the rate equation, think SN1, and if it is, think SN2

Section 31 of 40Acid-catalysed iodination of propanone

1Reaction

• CH₃COCH₃ + I₂ → CH₃COCH₂I + H⁺ + I⁻, catalysed by dilute sulfuric acid

2Rate equation

• Rate equation rate = k[CH₃COCH₃][H⁺]

  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)
• It is zero order in iodine, so iodine is not in the rate-determining step.
• Iodine reacts afterwards in a fast step.

3Finding the order for iodine

• Finding the order for iodine propanone and acid are in large excess, so only iodine changes
• Samples are quenched with sodium hydrogencarbonate and titrated with sodium thiosulfate using starch.
• A straight-line concentration–time graph shows zero order in iodine, with constant rate.

4Catalyst

• H⁺ appears in the rate equation but not in the overall equation

Section 32 of 40More from the specification

1Initial Rates Method

• The initial rate of reaction is determined from the gradient of a tangent drawn at time zero on concentration or gas volume versus time graphs.

2Arrhenius Equation

• Expresses the mathematical dependence of the rate constant on absolute temperature: k = A e^(-Ea/RT),
• where Ea is the activation energy and A is the pre-exponential factor.

3Arrhenius Plot and Activation Energy

• Plotting ln(k) against 1/T yields a straight line with a negative gradient equal to -Ea/R, providing an experimental means to calculate activation energy.

4Continuous Rate Measuring Techniques

• Continuous monitoring of reaction progress involves measuring physical quantities such as gas volume, mass change,
• or optical absorbance over time.

5Initial Rates Method

• The initial rate of reaction is determined from the gradient of a tangent drawn at time zero on concentration or gas volume versus time graphs.

6Arrhenius Equation

• Expresses the mathematical dependence of the rate constant on absolute temperature: k = A e^(-Ea/RT),
• where Ea is the activation energy and A is the pre-exponential factor.

7Arrhenius Plot and Activation Energy

• Plotting ln(k) against 1/T yields a straight line with a negative gradient equal to -Ea/R, providing an experimental means to calculate activation energy.

8Continuous Rate Measuring Techniques

• Continuous monitoring of reaction progress involves measuring physical quantities such as gas volume, mass change,
• or optical absorbance over time.

9Activation energy

• The minimum energy colliding particles need to react.

10Arrhenius equation

• K = Ae^−E_a/RT, relating the rate constant to temperature and activation energy.

11Directly proportional

• A relationship whose graph is a straight line through the origin.

12Clock reaction

• A reaction in which a sudden visible change marks when a fixed amount has reacted.

13Continuous monitoring

• Collecting data throughout a reaction.

14Initial rate

• The rate at the start of the reaction, at t = 0.

15Volume

• The space taken up by an object.

16Gradient

• The change in y divided by the change in x on a graph.

Section 33 of 40What changes the rate constant

Rate equation

• Rate equation shows how each reactant concentration affects the rate and contains the rate constant k.
• k stays constant only when concentration is the only factor that changes.

Temperature

• Temperature changing the temperature changes k.
• At higher temperature a greater proportion of molecules have energy of at least the activation energy.
• Rate of reaction and rate constant are both proportional to this fraction, so both increase.

Catalyst

• Catalyst using or changing a catalyst changes k because it changes the activation energy.

Section 34 of 40The Arrhenius equation

Equation

• Equation k = Ae^-Ea/RT relates the rate constant to temperature and activation energy.
• k = rate constant.
• A = Arrhenius constant, linked to collision frequency and orientation.
• Ea = activation energy in J mol^-1.
• R = gas constant, 8.31 J K^-1 mol^-1.
• T = temperature in kelvin.

Constants

• Ea and A are characteristic of a reaction and R is fundamental.
• A varies slightly with temperature but is treated as constant.
• k and T are the only variables in the equation.

Scope

• Scope it describes reactions of gases, reactions in solution and reactions on the surface of a catalyst.

Section 35 of 40Effect of temperature and Ea on k

Higher T

• Higher T gives a larger value of ln k and so a larger k and a faster reaction.

Higher Ea

• Higher Ea means a smaller proportion of molecules reach the activation energy, so k and the rate decrease.

Catalyst

• Catalyst lowers Ea, so k increases at the same temperature.

Section 36 of 40Logarithmic form

Taking natural logs

• Taking natural logs gives ln k = ln A - Ea/RT, which is easier to use.
• Compare with y = mx + c.
• y = ln k and x = 1/T.
• Gradient m = -Ea/R.
• Intercept c = ln A.

Two temperatures

• Two temperatures subtracting the equations eliminates A: ln(k₂/k₁) = (Ea/R)(1/T₁ - 1/T₂).

Section 37 of 40Arrhenius plot to find Ea

1

Convert every temperature to kelvin and calculate 1/T.

2

Calculate ln k for each rate constant.

3

Plot ln k (y-axis) against 1/T (x-axis) and draw a line of best fit.

4

Choose two points far apart on the line and calculate the gradient.

5

Calculate Ea = -gradient x R, which gives J mol^-1.

6

Divide by 1000 to give Ea in kJ mol^-1.

7

Read A from the intercept: A = e^intercept.

Section 38 of 40Interpreting the plot

Negative gradient

• Negative gradient expected, because a smaller 1/T (higher T) gives a larger ln k.

Steeper line

• Steeper line means a larger Ea, so the rate is more sensitive to temperature.

Catalyst

• Catalyst lowers Ea, so the line is less steep.

Best-fit line

• Best-fit line all points should be plotted rather than using only two data points.

Section 39 of 40Key terms

Term	Meaning
Rate equation	An expression rate = k[A]^m[B]ⁿ found by experiment.
Order of reaction	The power to which a concentration is raised in the rate equation.
Rate constant (k)	The proportionality constant in the rate equation, which depends on temperature.
Half-life	The time for the concentration of a reactant to fall to half its value.
Initial rate	The rate at the start of the reaction, at t = 0.
Rate-determining step	The slowest step in a mechanism.
Intermediate	A species formed in one step and used up in a later step.
Arrhenius equation	k = Ae^−Ea/RT, relating the rate constant to temperature and activation energy.

Term	Meaning
Order of reaction	The power to which a reactant concentration is raised in the rate equation.
Rate constant	The proportionality constant k in the rate equation, which depends on temperature.
Half-life	The time for the concentration of a reactant to fall to half its value.
Overall order	The sum of the orders with respect to each reactant.

Term	Meaning
Initial rate	The gradient of the concentration–time curve at t = 0.
Clock reaction	A reaction in which a sudden visible change marks when a fixed amount has reacted.
Colorimetry	Measurement of light absorbance to find concentration.
Calibration curve	A graph of absorbance against known concentrations.
Quenching	Stopping a reaction in a sample so that it can be analysed.

Term	Meaning
Rate-determining step	The slowest step in a mechanism, which controls the rate.
Elementary step	A single stage in a reaction mechanism.
Intermediate	A species formed in one step and used up in a later step.
SN1	Two-step nucleophilic substitution with a carbocation intermediate.
SN2	One-step nucleophilic substitution involving both reactants in the rate-determining step.
Carbocation	A positive ion with a positively charged carbon atom.

Term	Meaning
Arrhenius equation	k = Ae-Ea/RT, relating the rate constant to temperature and activation energy.
Activation energy (Ea)	The minimum energy particles need for a successful collision.
Gas constant R	8.31 J K⁻¹ mol⁻¹.
Arrhenius plot	A graph of ln k against 1/T with gradient -Ea/R and intercept ln A.
Natural logarithm	Logarithm to base e, written ln.

Section 40 of 40Exam technique

1

Writing orders from the stoichiometric coefficients — orders are found only by experiment

2

Giving units of k that do not match the overall order

3

Confusing a straight-line concentration–time plot (zero order) with a rate–concentration straight line (first order)

4

Including an intermediate in the rate equation

5

Saying SN1 gives inversion — SN2 gives inversion and SN1 gives a racemic mixture

1

Writing orders from the stoichiometric coefficients

2

Giving units of k that do not match the overall order

3

Confusing a straight-line concentration–time plot (zero order) with a rate–concentration straight line (first order)

4

Including an intermediate in the rate equation

1

Using the average rate over the whole reaction as the initial rate

2

Changing two concentrations at once

3

Assuming absorbance is proportional to concentration without calibration

4

Forgetting to keep the temperature constant

1

Writing the rate equation from the stoichiometry

2

Including the intermediate in the rate equation

3

Saying SN1 gives inversion — SN1 gives a racemic mixture and SN2 gives inversion

4

Not checking that the steps add up to the overall equation

1

Using temperatures in °C — T must be in kelvin.

2

Forgetting to convert J mol⁻¹ to kJ mol⁻¹ — divide by 1000 when a kJ answer is needed.

3

Taking the gradient as +Ea/R — the gradient is -Ea/R, so Ea = -gradient x R.

4

Using log₁₀ in place of ln — the equation needs natural logarithms (or a factor of 2.303 must be applied).