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

Section 1 of 39Rate of reaction

Rate of reaction

• Rate of reaction the change in the amount or concentration of a reactant or product per unit time.

Units

• Units mol dm⁻³ s⁻¹.

Rate = change in concentration of reactant or product (mol dm⁻³) ÷ time (s).

The rate at any moment is the gradient of the tangent to a concentration–time graph.

• The initial rate is the tangent at t = 0.

Section 2 of 39Rate and concentration

Rate and concentration

For D(aq) → E(aq) + F(g), rate is directly proportional to [D]: rate ∝ [D], so rate = k[D].

• Doubling [D] doubles the rate.
• Halving [D] halves the rate.

A plot of rate against concentration is a straight line through the origin for a first-order reactant.

Section 3 of 39The rate equation

What the symbols mean:

For A + B → C + D, rate = k[A]^m[B]ⁿ.

What the symbols mean: rate (mol dm⁻³ s⁻¹) · k = rate constant (units depend on the overall order) · [ ] = concentration (mol dm⁻³) · each power is the order with respect to that reactant

[A] and [B] are concentrations; m and n are the orders with respect to each reactant.

Rate equations can only be found by experiment, not from the stoichiometric equation.

Products and catalysts can appear in rate equations; intermediates cannot.

A catalyst can appear in the rate equation but not in the overall equation.

Section 4 of 39Properties used to follow a reaction

• Any property that changes during the reaction and is proportional to concentration can be measured.

Property	Method
Mass lost	Balance
Volume of gas	Gas syringe or collection over water
Colour change	Colorimeter
pH	pH probe
Electrical conductivity	Conductivity meter
Concentration	Titration of quenched samples

• Some reactions are followed as they proceed,
  • giving more data; faster reactions are easier to measure by averaging a measurement over the whole reaction.

Section 5 of 39Colorimetry

1

A colorimeter measures the amount of light passing through a solution.

2

If a solution changes colour, the intensity of light reaching the detector is recorded every few seconds.

3

Light intensity is related to concentration, so the data give concentration against time.

4

Colorimetry cannot be used for coloured precipitates because they scatter or block the light.

Section 6 of 39Mass loss and gas volume

Mass loss and gas volume

Mass loss

• A gas that escapes lowers the mass, e.g. CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂.
• The gas must be dense enough to see on a 2 or 3 d.p. balance: CO₂ (M_r = 44) is suitable, H₂ (M_r = 2) is not.

Gas volume

• Collect the gas in a gas syringe, or over water in an inverted measuring cylinder if it is not water soluble, e.g. Mg + 2HCl → MgCl₂ + H₂.
• Plot volume (amount of product) against time.

Section 7 of 39Sampling, quenching and stop-the-clock methods

Taking a sample and titrating it can change the rate, unless the reaction is stopped by quenching.

Quenching methods include adding sodium hydrogencarbonate to remove an acid catalyst, or cooling and diluting.

A convenient alternative is to stop the clock at a visible point, e.g. a piece of magnesium dissolving completely.

Disappearing cross

• Na₂S₂O₃ + 2HCl → 2NaCl + SO₂ + H₂O + S(s); sulfur obscures a cross viewed through the solution.
• It gives only one piece of data per experiment.

Section 8 of 39More from the specification

1

Polarimetric chiral purity assessment

• Polarimetric chiral purity assessment measuring the optical rotation of monochromatic light provides an analytical means to assess enantiomeric excess.

2

Halogenoalkane

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

3

Activation energy

• The minimum energy colliding particles need to react.

4

Nucleophilic substitution

• Replacement of an atom or group by a nucleophile.

5

Continuous monitoring

• Collecting data throughout a reaction.

6

Elementary step

• A single stage in a reaction mechanism.

7

Arrhenius plot

• Arrhenius plot a graph of ln k against 1/T is a straight line with gradient −E_a/R, so E_a can be calculated.

8

Rate-determining step

• Rate-determining step the slowest step in a mechanism, and the species in it appear in the rate equation.

9

Mechanism

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

10

Graph shapes

• Graph shapes a zero-order concentration–time graph is a straight line, while a first-order one is a falling curve with constant half-life.

Section 9 of 39More on this topic

• A catalyst provides an alternative route with a lower activation energy.

Section 10 of 39Order of reaction

1Order with respect to a reactant

• Order with respect to a reactant the power to which its concentration is raised in the rate equation;
  • it may be positive, negative or fractional.

2

Fractional orders suggest a multi-step reaction.

3Zero order

• Zero order changing the concentration has no effect on rate, and it is not in the rate equation.

4First order

• First order rate is directly proportional to concentration, so doubling concentration doubles the rate.

5Second order

• Second order rate is proportional to the square of the concentration, so doubling concentration quadruples the rate.

6Overall order

• Overall order the sum of the powers of the reactants in the rate equation.

Section 11 of 39Deducing orders from experimental data

1

Take the reactants one at a time.

2

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

3

Compare how the rate changes: unchanged is order 0, doubled is order 1, quadrupled is order 2.

4

Repeat for each reactant.

5

Assemble the rate equation, leaving out order-0 reactants and omitting the power 1.

Deducing orders from experimental data

For (CH₃)₃CBr + OH⁻, doubling [(CH₃)₃CBr] doubles the rate (order 1) and doubling [OH⁻] quadruples it (order 2), so rate = k[(CH₃)₃CBr][OH⁻]².

Read standard-form values carefully, and use any graph given, because it shows the order of one reactant while the table gives the others.

Section 12 of 39Rate constant and its units

Rate constant, k

• Rate constant, k the constant in the rate equation;
  • it stays constant only when concentration is the only factor changed.

k changes if the temperature changes or a catalyst is used.

Calculate k by substituting the rate and concentrations of one experiment into the rearranged rate equation; the other experiments give the same value.

Find the units by replacing each term with its units and cancelling.

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 13 of 39Concentration–time graphs

1Zero order

• Zero order a straight line sloping down; the gradient is the rate and equals k.
• Rate is independent of concentration.

2First order

• First order a downward curve that levels off.

3Second order

• Second order a steeper curve that levels off more slowly.

4

The rate at any point is the gradient of the tangent at that point.

Section 14 of 39Rate–concentration graphs

1Zero order

• Zero order a horizontal line; rate = k.

2First order

• First order a straight line through the origin; rate = k[A].

  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)

3Second order

• Second order an upward curve; rate = k[A]².

4

Do not confuse these with concentration–time graphs: the zero-order concentration–time line is straight but the first-order rate–concentration line is a different straight line.

Section 15 of 39Half-life

1Half-life, t1/2

• Half-life, t1/2 the time for the concentration of a reactant to fall to half its value.

2Zero order

• Zero order successive half-lives decrease as the reaction proceeds.

3First order

• First order the half-life is constant;
  • if several successive half-lives are equal, the reaction is first order.
• k = ln 2 ÷ t_1/2 = 0.693 ÷ t_1/2.

4Second order

• Second order successive half-lives increase.

Section 16 of 39More from the specification

1

Buffer capacity of weak acid solution

• Buffer capacity of weak acid solution equimolar mixtures of ethanoic acid and sodium ethanoate in solution resist pH changes upon small additions of acid.

2

Reaction mechanisms and rate-determining step

• Complex chemical reactions proceed through a multi-step reaction mechanism;
  • the slowest elementary step is the rate-determining step,
• which means its molecularity governs the overall rate equation.

3

Arrhenius equation and activation energy

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

4

Initial rate method

• In experimental kinetics, varying the initial concentration of reactants while maintaining constant total volume allows the determination of individual reaction orders from initial reaction rates.
  • (such as the iodine clock reaction)

5

Energy profile and intermediate steps

• In a multi-step reaction profile, each elementary step has its own activation energy barrier and transition state, with transient reaction intermediates formed between successive steps.

6

Halogenoalkane

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

7

Activation energy

• The minimum energy colliding particles need to react.

8

Arrhenius equation

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

9

Clock reaction

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

10

Nucleophilic substitution

• Replacement of an atom or group by a nucleophile.

11

Continuous monitoring

• Collecting data throughout a reaction.

12

Elementary step

• A single stage in a reaction mechanism.

13

Initial rate

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

14

Mechanism

• The sequence of steps by which reactants form products.

15

Volume

• The space taken up by an object.

Section 17 of 39Initial rate method

• Initial rate
  • the rate at t = 0, when concentrations are known exactly.

1

Run several experiments with different starting concentrations.

2

Plot concentration (or product amount) against time.

3

Draw a tangent at t = 0.

4

Calculate its gradient, which is the initial rate.

Initial rate method

Example reaction: CaCO₃ with HCl, collecting CO₂ and plotting volume against time.

Compare pairs of experiments in which one concentration changes to find the orders.

The limitation is that each experiment gives only one point, so many experiments are needed.

Section 18 of 39Clock reactions

A clock reaction measures the initial rate with one measurement: the time t for a sudden visible change.

The iodine clock reaction gives a blue-black colour; other clocks use a colour change or a precipitate.

Assumption

• Assumption the rate does not change significantly between the start and the time of the measurement.
• Then initial rate ∝ 1/t.

Compare 1/t at different concentrations: when concentration doubles and the rate doubles, the reaction is first order in that reactant.

Limitations

• The measured value is only an estimate of the initial rate, because reactants are used up.
• Shorter times give values closer to the true initial rate.

Keep temperature and all other concentrations constant and repeat to check reliability.

Section 19 of 39Continuous monitoring

Continuous monitoring

Continuous monitoring collects data throughout the reaction to plot a concentration–time graph.

Two common ways are measuring the volume of gas evolved and following colour with a colorimeter.

Draw tangents at several points to find rates; plotting several points is more accurate than using one pair.

Section 20 of 39Iodination of propanone by colorimetry

Reaction

• CH₃COCH₃ + I₂ → CH₃COCH₂I + HI, catalysed by dilute sulfuric acid.

The colorimeter measures absorbance, which is proportional to the concentration of the coloured species, iodine.

1

Choose a filter complementary to the colour of the solution for the strongest absorbance.

2

Measure the absorbance of standard iodine solutions and plot a calibration curve.

3

Fill burettes with 0.02 mol dm⁻³ iodine, 1.0 mol dm⁻³ propanone, 1.0 mol dm⁻³ sulfuric acid and distilled water.

4

Mix propanone, acid and water in a small beaker, varying volumes with the total volume constant.

5

Add the iodine last to start the reaction and the timer, mix and transfer quickly to the cuvette.

6

Record absorbance at intervals, convert to concentration and plot concentration against time.

Section 21 of 39Interpreting the iodination results

1

The iodine concentration falls in a straight line, so the gradient (rate) is constant and independent of iodine concentration.

2

The reaction is zero order with respect to iodine.

3

The rate equation is rate = k[CH₃COCH₃][H⁺].

• It is first order in propanone and first order in acid.
• Iodine does not take part in the rate-determining step.

4

Titration version

• Propanone and acid are in large excess, so their concentrations stay constant.
• Remove portions and quench with sodium hydrogencarbonate.
• Titrate with sodium thiosulfate using starch2S₂O₃²⁻ + I₂ → 2I⁻ + S₄O₆²⁻.

Section 22 of 39Practical points

1

Keep the temperature constant with a water bath, because k changes with temperature.

2

Keep the total volume constant so concentration is proportional to the volume of each stock solution.

3

Use a fresh start for each run and repeat to find anomalous results.

4

The colorimeter uses small volumes, so four burettes can hold the stock solutions.

5

Timing errors arise from starting the clock late and the finite time to mix.

Section 23 of 39More from the specification

1

Spectrophotometric dye quantification

• Spectrophotometric dye quantification applying the Beer-Lambert law relating absorbance to concentration provides an accurate means to determine unknown solution concentrations.

2

Multi-step reaction mechanisms

• Chemical reactions often occur through a sequence of elementary steps termed the reaction mechanism;
  • the rate-determining step dictates the overall rate law for the reaction.

3

Concentration-time graphs and reaction order

• On a concentration-time graph, a linear gradient indicates zero-order kinetics, a curve with constant half-life indicates first-order kinetics,
• and an increasingly shallow curve with doubling half-life indicates second-order kinetics.

4

Arrhenius equation and activation energy

• The Arrhenius equation (k = A e^(-Ea / RT)) links rate constant k to temperature;
  • plotting ln(k) versus 1/T yields a straight-line graph whose slope (-Ea / R) enables determination of the activation energy Ea.

5

Initial rates method

• In experimental kinetics, the initial rate method measures the initial gradient of concentration-time curves under varying initial reactant concentrations to deduce individual and overall orders of reaction.

6

Multi-step reaction mechanisms

• Chemical reactions often occur through a sequence of elementary steps termed the reaction mechanism;
  • the rate-determining step dictates the overall rate law for the reaction.

7

Concentration-time graphs and reaction order

• On a concentration-time graph, a linear gradient indicates zero-order kinetics, a curve with constant half-life indicates first-order kinetics,
• and an increasingly shallow curve with doubling half-life indicates second-order kinetics.

8

Arrhenius equation and activation energy

• The Arrhenius equation (k = A e^(-Ea / RT)) links rate constant k to temperature;
  • plotting ln(k) versus 1/T yields a straight-line graph whose slope (-Ea / R) enables determination of the activation energy Ea.

9

Initial rates method

• In experimental kinetics, the initial rate method measures the initial gradient of concentration-time curves under varying initial reactant concentrations to deduce individual and overall orders of reaction.

10

Directly proportional

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

11

Nucleophilic substitution

• Replacement of an atom or group by a nucleophile.

12

Mechanism

• The sequence of steps by which reactants form products.

13

Halogenoalkane

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

14

Activation energy

• The minimum energy colliding particles need to react.

15

Reaction mechanism

• The sequence of elementary steps in a reaction.

16

Arrhenius equation

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

Section 24 of 39The rate-determining step

1

A reaction can only go as fast as its slowest step, the rate-determining step.

2

If a reactant is in the rate-determining step, its concentration appears in the rate equation.

• CH₃Br + OH⁻ → CH₃OH + Br⁻ with rate = k[CH₃Br][OH⁻] means both take part in the slow step.

3

The order of a reactant equals the number of molecules of it in the rate-determining step (and any fast steps before it).

4

Species in the overall equation but not in the rate equation react after the slow step.

5

Intermediates cannot appear in a rate equation.

Section 25 of 39Mechanisms and elementary steps

1Elementary step

• Elementary step a single molecular event in a mechanism;
  • an intermediate forms in one step and is used up in another.

2

The steps must add up to the overall stoichiometric equation, otherwise the proposed mechanism is wrong.

3

Kinetics can suggest a mechanism but cannot prove it; it can disprove one if the predicted rate equation does not match experiment.

4

NO₂ + CO → NO + CO₂ with rate = k[NO₂]².

• Step 1 (slow): 2NO₂ → NO + NO₃.
• Step 2 (fast): NO₃ + CO → NO₂ + CO₂.
• CO is zero order and does not appear in the slow step.

5

2NO + 2H₂ → N₂ + 2H₂O.

• Step 1 (fast): NO + NO → N₂O₂.
• Step 2 (slow): N₂O₂ + H₂ → H₂O + N₂O.
• Step 3 (fast): N₂O + H₂ → N₂ + H₂O.
• Substituting the intermediate gives rate = k[NO]²[H₂], third order overall.

Section 26 of 39Deducing the rate equation from a mechanism

1

Identify the slow step.

2

Write down the species in the slow step.

3

Give each species an order equal to its coefficient in that step.

4

If an intermediate is in the slow step, replace it using the earlier fast step.

What the symbols mean:

If a slow step A + B → X is followed by fast X + C → products, rate = k[A][B] and C is zero order.

What the symbols mean: rate (mol dm⁻³ s⁻¹) · k = rate constant (units depend on the overall order) · [ ] = concentration (mol dm⁻³) · each power is the order with respect to that reactant

A mechanism with a slow step involving A and B is consistent with an experimental rate = k[A][B].

Catalysts can take part in the rate-determining step and so appear in the rate equation.

Section 27 of 39SN2 mechanism

1

Primary halogenoalkanes react by SN2: S = substitution, N = nucleophilic, 2 = two species in the rate-determining step.

2

One-step reaction

• One-step reaction the nucleophile donates a pair of electrons to the δ+ carbon as the C-X bond breaks and X⁻ leaves.
• Rate = k[halogenoalkane][nucleophile], second order overall; it is bimolecular.
• Bromoethane with hydroxide gives ethanol by attack from the side opposite bromine, causing inversion of configuration.

Section 28 of 39SN1 mechanism

1

Tertiary halogenoalkanes react by SN1: the 1 means only the halogenoalkane is in the rate-determining step.

2

Two steps

• Slow stepthe C-X bond breaks heterolytically to give X⁻ and a tertiary carbocation.
• Fast stepthe nucleophile attacks the carbocation.

3

Rate = k[halogenoalkane]; it is unimolecular, so 2-bromo-2-methylpropane with OH⁻ is first order in the halogenoalkane and zero order in OH⁻.

4

Attack on the planar carbocation can occur from either side, giving a racemic mixture.

Feature	SN2	SN1
Halogenoalkane	Primary (secondary also)	Tertiary (secondary also)
Steps	One step	Two steps via carbocation
Rate equation	k[RX][Nu-]	k[RX]
Intermediate	Transition state only	Carbocation
Stereochemistry	Inversion	Racemic mixture

• Secondary halogenoalkanes react by both mechanisms.

Section 29 of 39Acid-catalysed iodination of propanone

1

CH₃COCH₃ + I₂ → CH₃COCH₂I + HI is catalysed by dilute sulfuric acid.

2

Experiments give rate = k[CH₃COCH₃][H⁺].

3

Iodine is zero order, so it reacts after the rate-determining step.

4

The catalyst H⁺ appears in the rate equation, so catalysts can take part in the rate-determining step.

5

A proposed mechanism must have one propanone molecule and one H⁺ in or before the slow step.

Section 30 of 39More from the specification

More from the specification

Titration equivalence in solution

• Titration equivalence in solution standardized sodium hydroxide neutralizing hydrochloric acid in aqueous solution demonstrates stoichiometric completion at neutral pH.

Reaction Order and Rate Equations

• The order of reaction with respect to a particular reactant indicates how the initial rate changes when that reactant's concentration is modified, appearing as an exponent in the rate equation: Rate = k [A]^m [B]^n.

Graphical Determination of Reaction Rates

• On a concentration-time graph or volume of gas evolved graph, the instantaneous reaction rate is determined by measuring the gradient of a tangent drawn to the curve at any specified time.

Temperature Dependence and the Arrhenius Equation

• The rate constant k increases exponentially with temperature according to the Arrhenius equation: k = A e^(-Ea / RT),
• where Ea is the activation energy and A is the Arrhenius pre-exponential factor.

Arrhenius Plot for Determining Activation Energy

• By plotting ln(k) against 1/T on a graph, a straight line with gradient -Ea/R is obtained,
• allowing the activation energy Ea to be calculated from the change in rate constant over varying temperatures.

Reaction Order and Rate Equations

• The order of reaction with respect to a particular reactant indicates how the initial rate changes when that reactant's concentration is modified, appearing as an exponent in the rate equation: Rate = k [A]^m [B]^n.

Graphical Determination of Reaction Rates

• On a concentration-time graph or volume of gas evolved graph, the instantaneous reaction rate is determined by measuring the gradient of a tangent drawn to the curve at any specified time.

Temperature Dependence and the Arrhenius Equation

• The rate constant k increases exponentially with temperature according to the Arrhenius equation: k = A e^(-Ea / RT),
• where Ea is the activation energy and A is the Arrhenius pre-exponential factor.

Arrhenius Plot for Determining Activation Energy

• By plotting ln(k) against 1/T on a graph, a straight line with gradient -Ea/R is obtained,
• allowing the activation energy Ea to be calculated from the change in rate constant over varying temperatures.

Activation energy

• The minimum energy colliding particles need to react.

Arrhenius equation

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

Directly proportional

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

Clock reaction

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

Continuous monitoring

• Collecting data throughout a reaction.

Initial rate

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

Volume

• The space taken up by an object.

Gradient

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

Section 31 of 39Why the rate constant changes

Why the rate constant changes

k stays constant only if concentration is the only factor changed.

A change of temperature or the use of a catalyst changes k.

At higher temperature a greater proportion of molecules have energy of at least the activation energy.

• The rate of reaction increases.
• The rate constant increases.

Section 32 of 39The Arrhenius equation

What the symbols mean:

k = Ae^−E_a/RT

What the symbols mean: k = rate constant · A = pre-exponential (frequency) factor · E_a = activation energy (J mol⁻¹) · R = 8.31 J mol⁻¹ K⁻¹ · T = temperature (K)

• A — the Arrhenius constant, related to the frequency of collisions with the correct orientation.
• E_a
  • the activation energy, in J mol⁻¹.
• R — the gas constant, 8.31 J K⁻¹ mol⁻¹.
• T — the temperature in kelvin.

e^−E_a/RT is the fraction of molecules with energy equal to or greater than E_a.

A and E_a are constants for a reaction; A varies slightly with temperature but is treated as constant, so k and T are the only variables.

It applies to reactions of gases, in solution and on a catalyst surface.

Section 33 of 39Effect of T and activation energy on k

1

Higher T increases the fraction with E ≥ E_a, so k and the rate increase.

2

A larger E_a reduces the proportion of molecules with enough energy, so k and the rate decrease.

3

A rise of about 10 °C roughly doubles the rate of many reactions with moderate activation energy.

4

A catalyst lowers E_a, which makes k larger at a given temperature.

Section 34 of 39Logarithmic form and the Arrhenius plot

ln k = ln A − E_a/RT, which has the form y = mx + c.

A graph of ln k (y) against 1/T (x) is a straight line.

• Gradient = −E_a/R.
• Intercept = ln A.

Two-point form

• Two-point form ln(k₂/k₁) = (E_a/R)(1/T₁ − 1/T₂).

Section 35 of 39Finding Ea and A from a graph

1

Convert temperatures to kelvin (add 273) 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 the line of best fit.

4

Find the gradient using a large triangle: Δy ÷ Δx.

5

E_a = −gradient × R, in J mol⁻¹; divide by 1000 for kJ mol⁻¹.

6

Read the intercept to get ln A, then A = e^intercept.

• Choose scales so the line fills at least half the graph paper and label axes with quantities and units (1/T in K⁻¹).

Section 36 of 39Interpreting the graph

Interpreting the graph

A steep negative gradient means a high activation energy and a rate very sensitive to temperature.

A catalyst lowers E_a, so the line is less steep.

The gradient is negative because k increases as 1/T decreases, that is as temperature rises.

Section 37 of 39More from the specification

1

Rate-Determining Step and Reaction Mechanisms

• A multi-step chemical reaction mechanism proceeds via intermediate elementary steps;
  • the slowest elementary step is the rate-determining step,
• which limits the overall rate of reaction.

2

Reaction Order and Rate Expression

• The order of reaction with respect to each of the reactants reflects how many molecules or ions of that reactant participate in or precede the rate-determining step, dictating the overall rate equation: Rate = k [A]^m [B]^n.

3

Initial Rates Method and Gas Volume Graphs

• In kinetics investigations, the initial rate of reaction is deduced from the gradient at time t = 0 on graphs of gas volume produced versus time or concentration versus time across different starting reactant concentrations.

4

The Iodine Clock Reaction in Kinetics Studies

• In the iodine clock reaction, the time taken for a dark blue starch-iodine complex to suddenly appear is measured;
  • the inverse of this time (1/t) serves as a direct proxy for initial reaction rate.

5

Halogenoalkane

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

6

Reaction order

• Exponent of a reactant concentration in the rate law.

7

Directly proportional

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

8

Clock reaction

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

9

Nucleophilic substitution

• Replacement of an atom or group by a nucleophile.

10

Continuous monitoring

• Collecting data throughout a reaction.

11

Elementary step

• A single stage in a reaction mechanism.

12

Initial rate

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

13

Mechanism

• The sequence of steps by which reactants form products.

14

Reactants

• The substances that react together at the start of a reaction.

15

Volume

• The space taken up by an object.

Section 38 of 39Key terms

Term	Meaning
Rate of reaction	The change in concentration of a reactant or product per unit time, in mol dm−3 s−1.
Rate equation	An expression rate = k[A]m[B]n found by experiment.
Colorimeter	An instrument that measures how much light a coloured solution absorbs.
Quenching	Stopping a reaction in a sample so it can be analysed.
Initial rate	The rate at time zero, from the tangent at the start of the curve.

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

Term	Meaning
Initial rate	The rate of reaction at time zero.
Clock reaction	A reaction where the time for a sudden visible change gives the initial rate.
Colorimeter	An instrument that measures the absorbance of light by a coloured solution.
Calibration curve	A graph of absorbance against known concentration.
Continuous monitoring	Collecting data throughout a reaction.

Term	Meaning
Rate-determining step	The slowest step in a reaction mechanism.
Elementary step	A single molecular event in a mechanism.
Intermediate	A species formed in one step and used up in a later step.
SN2	Bimolecular nucleophilic substitution in one step.
SN1	Unimolecular nucleophilic substitution via a carbocation.
Molecularity	The number of species taking part in an elementary step.

Term	Meaning
Arrhenius equation	k = Ae^(−Ea/RT), relating the rate constant to temperature and activation energy.
Activation energy	The minimum energy needed for a collision to be successful.
Arrhenius constant	The constant A, related to the frequency of correctly oriented collisions.
Gas constant	R = 8.31 J K−1 mol−1.
Gradient	The slope of a graph; here equal to −Ea/R.

Section 39 of 39Exam technique

1

Writing the rate equation from the balanced equation — orders must be found by experiment.

2

Using colorimetry to follow a precipitate — the solid scatters the light.

3

Using hydrogen for mass loss — its mass is too small to measure.

4

Forgetting to quench a sample before titrating — the reaction continues and the result is wrong.

1

Deducing orders from the coefficients in the balanced equation — orders come from experiment.

2

Giving k without units or with the wrong units — units depend on the overall order.

3

Saying the half-life of a first-order reaction shortens as the reaction goes on — it stays constant.

4

Drawing a tangent at the wrong point on the curve.

5

Including an order-0 reactant in the rate equation — it is left out.

1

Using the final rate instead of the rate at t = 0 — the initial rate is at the start.

2

Varying concentrations without keeping the total volume constant — concentrations would not be proportional to volumes.

3

Writing rate ∝ t in clock reactions — rate ∝ 1/t.

4

Reading concentration straight from absorbance — use the calibration curve.

1

Assuming the rate equation can be read from the overall equation — it comes from experiment.

2

Saying kinetics can prove a mechanism — it can only support or disprove one.

3

Forgetting inversion of configuration in SN2 — the nucleophile attacks from the opposite side.

4

Saying a catalyst cannot appear in the rate equation — it can.

5

Including an intermediate in the rate equation — replace it with the species that form it.

1

Using temperatures in °C — convert to kelvin.

2

Using gradient = Ea/R and forgetting the minus sign — gradient = −Ea/R.

3

Quoting Ea in kJ mol⁻¹ without converting from J mol⁻¹ — divide by 1000.

4

Plotting k against T instead of ln k against 1/T.