Kinetics II (Edexcel A Level Chemistry): Flashcards

True.

For a second-order reactant, rate ∝ [A]². Tripling [A] gives 3² = 9, so the rate increases by a factor of 9.

1. Identify two experiments where the concentration of the reactant of interest changes but all other concentrations remain constant.

2. Compare how the rate changes as the concentration changes.

3. If rate doubles when [A] doubles → first order. If rate quadruples → second order. If rate is unchanged → zero order.

Overall order of reaction is the sum of the individual orders of all reactants in the rate equation.

For rate = k[A]^m[B]ⁿ, the overall order is m + n.

For rate = k[BrO₃^-][Br^-][H⁺], the concentration of bromate ions is doubled while the concentration of bromide ions is halved. The overall effect on the rate is no change (the two effects cancel, as both species are first order).

k is the rate constant — the proportionality constant in the rate equation.

Its value is fixed for a given reaction at a given temperature and changes when temperature changes. The units of k depend on the overall order of the reaction.

Rearrange: k = rate / ([A][B]²)

Substituting units:

k = mol dm^-3 s^-1 / (mol dm^-3 × (mol dm^-3)²)

= mol dm^-3 s^-1 / mol³ dm^-9

= mol^-2 dm⁶ s^-1

(Overall order = 3, so units are mol^-2 dm⁶ s^-1.)

1. Colorimetry — monitors light intensity through a coloured solution

2. Mass loss — tracks decrease in mass as gas escapes

3. Gas collection — measures volume of gas produced over time

4. Titration — samples are analysed for concentration (requires quenching)

Colorimetry is a technique that measures the intensity of light passing through a coloured solution.

As the reaction proceeds and the concentration of the coloured species changes, the light intensity at the detector changes, allowing rate to be monitored continuously.

The disappearing cross method typically generates only one data point per experiment (the time for the cross to disappear), making it difficult to track how rate changes throughout the reaction.

False.

Colorimetry cannot be used to monitor coloured precipitates because the precipitate scatters or blocks the light, preventing an accurate reading.

Hydrogen has a very low molar mass (M_r = 2), so the mass change as hydrogen escapes is too small to be detected accurately on a 2 or 3 decimal-place balance.

A denser gas such as CO₂ (M_r = 44) produces a measurable change in mass.

In the disappearing cross experiment, sodium thiosulfate reacts with hydrochloric acid to form a yellow precipitate of sulfur, which obscures a cross viewed through the solution.

Quenching is deliberately stopping the reaction before taking a sample for analysis by titration.

It is necessary because removing a sample and analysing it can change the concentration of reactants and affect the rate if the reaction is still proceeding during analysis.

Draw a tangent to the curve at the time of interest and calculate the gradient of the tangent:

Rate = Δy / Δx = Δ[concentration] / Δtime

The units of rate are mol dm^-3 s^-1.

The initial rates method is an experimental technique that determines the rate of reaction at the very start of the reaction (t = 0).

It is found by drawing a concentration-time graph, adding a tangent at t = 0, and calculating the gradient of that tangent.

1. Plot a concentration-time graph using experimental data.

2. Draw a tangent to the curve at t = 0.

3. Calculate the gradient of the tangent:

rate = Δy / Δx

The gradient gives the initial rate in mol dm^-3 s^-1.

A clock reaction is a method of measuring the initial rate of a reaction using a single measurement — the time t for a specific visual change (colour change or precipitate) to occur.

The initial rate is proportional to 1/t.

In a clock reaction, the initial rate is proportional to 1/*t*.

For the iodine clock, if the time for the colour change to occur doubles, the initial rate is half as large.

False.

A clock reaction usually generates one data point per experiment — the time taken for the visual change to occur. This is a major limitation of the method.

Clock reactions assume that there is no significant change in the rate of reaction between the start of the reaction and the point at which the visual change is observed (i.e., the rate is approximately constant during the measurement period).

As the reaction proceeds, reactant concentrations decrease, so the rate is not constant — it changes on a curve.

The longer the measured time, the greater the deviation from the initial rate, making the calculated value a less accurate estimate of the true initial rate.

False.

The initial rate from a clock reaction is an estimate. It relies on the assumption that the rate does not change significantly during the measurement period, which is only approximately true.

The continuous monitoring method is a technique in which experimental data is collected throughout a reaction to plot a concentration-time graph, allowing the rate to be measured at any point.

Two common methods are measuring the volume of gas evolved over time and measuring the mass lost by the reaction mixture over time.

In the iodination of propanone, a colorimeter measures colour absorbance, which is proportional to the concentration of the coloured species.

True.

A calibration curve is obtained from standard solutions of known concentration. It allows the colorimeter's absorbance readings to be converted into concentration values during the experiment.

The rate of reaction at a specific point is found by drawing a tangent to the curve at that point and calculating the gradient of the tangent. The gradient equals the rate at that moment.

In the iodination of propanone, the iodine decolourises during the reaction as it is converted into iodopropanone and hydrogen iodide.

The filter is chosen to match the complementary colour of the solution because it gives the strongest absorbance reading, making the measurements more sensitive and accurate.

True.

By keeping the total volume constant with distilled water, different concentrations of each reactant can be studied without changing the overall volume of the reaction mixture.

A zero-order reaction is one in which the rate of reaction is independent of the concentration of the reactant. The rate equation is rate = k, and the rate-concentration graph is a horizontal line.

A first-order concentration-time graph is a downward exponential curve. The concentration decreases by the same fraction over each successive equal time interval, giving a constant half-life.

True.

In a first-order reaction, each successive half-life takes the same amount of time. The time required for the concentration to halve does not change as the reaction proceeds.

For a second-order reaction, the rate-concentration graph is a curve and the rate equation is rate = k[A]².

The half-life increases as a second-order reaction proceeds. It takes progressively longer for the concentration of the reactant to halve as the reaction continues.

The half-life of a reaction is the time taken for the concentration of a reactant to fall to half of its initial value.

In a zero-order reaction, the concentration-time graph is a straight line going downwards, and the gradient of this line equals the rate of reaction.

True.

In a second-order reaction, concentration decreases more sharply with time than in a first-order reaction. The curve falls more steeply because rate depends on the square of concentration.

A zero-order reaction produces a horizontal line on a rate-concentration graph because the rate is independent of the reactant's concentration. As concentration changes, the rate stays constant.

The rate-determining step is the slowest step in a reaction mechanism. It controls the overall rate of the reaction, and only reactants involved in this step appear in the rate equation.

A reactant's concentration appears in the rate equation because that reactant is involved in the rate-determining step. If a reactant is not in the rate-determining step, it does not appear in the rate equation.

The rate equation for CH₃Br + OH^- → CH₃OH + Br^- is rate = k[CH₃Br][OH^-], which means that CH₃Br and OH^- both take part in the rate-determining step.

False.

Intermediates cannot appear in the rate equation. If an intermediate is involved in the rate-determining step, it must be substituted back in terms of the original reactants using an earlier step.

The rate equation shows the reaction is second order in NO₂ and zero order in CO. This means the rate-determining step involves two molecules of NO₂ and no CO molecules.

The molecularity of a reaction step is the number of molecules involved in that step. A step involving one molecule is unimolecular; a step involving two molecules is bimolecular.

In the reaction of NO with H₂, the rate-determining step involves N₂O₂ (formed from two NO molecules) and one H₂ molecule, giving the rate equation rate = k[NO]²[H₂].

False.

A zero-order reactant does not take part in the rate-determining step. Its concentration has no effect on the rate of reaction, so it does not appear in the rate equation.

A proposed mechanism is valid if two conditions are met: the elementary steps must add up to the overall stoichiometric equation, and the rate-determining step must be consistent with the experimental rate equation.

An S_N1 reaction is a two-step nucleophilic substitution in which the C–X bond breaks first to form a carbocation intermediate (slow, rate-determining step), followed by attack from the nucleophile (fast step).

In S_N1, the rate-determining step is the unimolecular breaking of the C–X bond. The nucleophile is not involved in this slow step, so only the halogenoalkane concentration appears in the rate equation: rate = k[halogenoalkane].

In an S_N2 reaction, the nucleophile attacks the δ+ carbon of the halogenoalkane at the same time as the C–X bond breaks, making it a one-step reaction.

True.

In the first step, the C–X bond breaks heterolytically to form a carbocation (slow). In the second step, the nucleophile attacks the carbocation (fast).

An S_N2 reaction is a one-step nucleophilic substitution in which the nucleophile attacks the δ+ carbon and the C–X bond breaks simultaneously. The rate depends on both the halogenoalkane and the nucleophile concentrations.

Primary halogenoalkanes react exclusively by the S_N2 mechanism. The carbon bonded to the halogen is attached to only one alkyl group, making it accessible to nucleophilic attack.

The rate equation for an S_N1 reaction is rate = k[halogenoalkane] and the reaction is described as unimolecular because only one molecule is involved in the rate-determining step.

False.

Secondary halogenoalkanes can react via both the S_N1 and S_N2 mechanisms. Only tertiary halogenoalkanes react exclusively by S_N1, and only primary halogenoalkanes react exclusively by S_N2.

Kinetics can only show which species are involved in the rate-determining step and the overall order of reaction. It is possible for more than one mechanism to be consistent with the same rate equation, so kinetics cannot confirm which mechanism is correct.

The Arrhenius equation is k = Ae^(−E_a/RT), where k is the rate constant, A is the pre-exponential factor, E_a is the activation energy, R is the gas constant, and T is the temperature in kelvin.

A is the pre-exponential factor (also called the frequency factor). It is a constant characteristic of the specific reaction and relates to the frequency of collisions with the correct orientation.

Taking natural logarithms of the Arrhenius equation gives ln k = ln *A − E_a/RT, and the gradient of a graph of ln k against 1/T equals −E_a/R*.

False.

The rate constant k increases with temperature. At higher temperatures, a greater proportion of molecules have energy equal to or greater than the activation energy, increasing the rate of reaction.

At higher temperatures, a greater fraction of molecules has energy equal to or greater than the activation energy E_a. Since k is directly proportional to this fraction, a higher temperature gives a higher k and a faster rate of reaction.

Activation energy (E_a) is the minimum energy that colliding molecules must possess for a reaction to occur. It is characteristic of a specific reaction and has units of kJ mol^-1.

In the Arrhenius equation, the only two variables are k and T, while all other quantities are constants for a given reaction.

True.

The linearised Arrhenius equation has the form y = mx + c, where the gradient m = −E_a/R. Multiplying the gradient by −R gives the activation energy E_a.

Values of k at different temperatures T are measured experimentally. A graph of ln k against 1/T is plotted, giving a straight line. The gradient of the line equals −E_a/R, so E_a = −gradient × R.