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Inorganic Chemistry > Topic 16 - Kinetics II > Flashcards

Topic 16 - Kinetics II Flashcards

1
Q

What is the reaction rate?

A

The change in the amount of reactants or products per unit time.

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2
Q

What are the different continuous monitoring methods of following the rate of a reaction?

A
  • Gas volume
  • Loss of mass
  • Colour change
  • Change in pH
  • Titration
  • Electrical conductivity
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3
Q

Remember to revise the different ways of following the rate of a reaction.

A

Pg 180 of revision guide

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4
Q

When continuously monitoring the volume of gas produced by a reaction, how can you work out the moles of a reactant at any point?

A
  • Look at the volume of gas produced
  • Use the ideal gas equation to work out how many moles of gas this is
  • Use the molar ratios in the chemical equation to work out the moles/concentration of the reactant
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5
Q

When continuously monitoring the loss of mass in a reaction, how can you work out the moles of a reactant at any point?

A
  • Use mole calculations to work out how much gas you’ve lost

* Use the molar ratios in the chemical equation to work out the moles/concentration of the reactant

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6
Q

What device is used to measure the colour change of a reaction?

A

Colorimeter

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7
Q

What does a colorimeter measure?

A

The absorbance of a solution.

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8
Q

When continuously monitoring the colour change in a reaction, how can you work out the moles of a reactant at any point?

A
  • Plot a calibration curve of known values (absorbance values against concentration of the coloured solution)
  • During the experiment, take a small sample from the reaction and read the absorbance
  • Use the calibration curve to convert the absorbance to a concentration
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9
Q

When continuously monitoring the change in pH in a reaction, how can you work out the moles of a reactant at any point?

A
  • Measure the pH
  • Calculate the concentration of H⁺
  • The rate can be calculated from this
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10
Q

When continuously monitoring a reaction, how can titration be used to work out the reaction rate?

A
  • Small samples can be taken at regular intervals
  • These can be titrated using a standard solution
  • The rate can be found from measuring the change in concentration of the products or reactants over time
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11
Q

When can electrical conductivity be used to continuously monitor a reaction?

A

If the number of ions changes, so will the electrical conductivity.

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12
Q

How can a reaction rate be found from a concentration-time graph?

A

It is the gradient at any point.

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13
Q

When using a concentration-time graph to measure the reaction rate, what are the units?

A

moldm⁻³s⁻¹

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14
Q

What does the order of a reaction with respect to a given reactant tell you?

A

How the reactant’s concentration affects the rate.

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15
Q

What does it mean if the order of reaction with respect to reactant X is 0?

A

The rate is not affected by [X].

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16
Q

What does it mean if the order of reaction with respect to reactant X is 1?

A

The rate is proportional to [X].

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17
Q

What does it mean if the order of reaction with respect to reactant X is 2?

A

The rate is proportional to [X]².

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18
Q

If the rate of a reaction with respect to reactant X is 0, what happens to the rate if [X] doubles and triples?

A
  • Doubles -> Rate stays the same

* Triples -> Rate stays the same

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19
Q

If the rate of a reaction with respect to reactant X is 1, what happens to the rate if [X] doubles and triples?

A
  • Doubles -> Rate doubles

* Triples -> Rate triples

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20
Q

If the rate of a reaction with respect to reactant X is 2, what happens to the rate if [X] doubles and triples?

A
  • Doubles -> Reaction is 4 times faster

* Triples -> Reaction is 9 tunes faster

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21
Q

Can orders of reaction be be worked out from a chemical equation?

A

No, but they can be worked out experimentally.

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22
Q

What two types of graph can be drawn in terms of orders of reactions?

A
  • Concentration against time

* Rate against concentration graph

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23
Q

How can you construct a rate-concentration curve from a concentration-time graph of a reaction?

A
  • Find the gradient of the concentration-time graph at multiple points
  • This will give you a series of rates at various concentrations
  • Plot each of these points on a new graph of rate against concentration
  • Draw a line or curve of best fit
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24
Q

Describe the concentration-time graph for a 0 order reaction.

A
  • Starts at positive y-intercept
  • Straight line of negative gradient

(See diagram pg 182 of revision guide)

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25
Q

Describe the concentration-time graph for a 1st order reaction.

A
  • Starts at positive y-intercept
  • Exponential decay (not very steep decline)

(See diagram pg 182 of revision guide)

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26
Q

Describe the concentration-time graph for a 2nd order reaction.

A
  • Starts at positive y-intercept
  • Curve down of decreasingly negative gradient -> Steeper than 1st order

(See diagram pg 182 of revision guide)

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27
Q

Describe the rate-concentration graph for a 0 order reaction.

A

Horizontal line at positive y-intercept

See diagram pg 182 of revision guide

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28
Q

Describe the rate-concentration graph for a 1st order reaction.

A
  • Starts at origin
  • Straight line of positive gradient

(See diagram pg 182 of revision guide)

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29
Q

Describe the rate-concentration graph for a 2nd order reaction.

A
  • Starts at origin
  • x² graph

(See diagram pg 182 of revision guide)

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30
Q

Remember to practise drawing out the concentration-time and rate-concentration graphs for the different orders of reaction.

A

Pg 182 of revision guide

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31
Q

What is the half-life of a reaction?

A

The time for half of the reactant to be used up.

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32
Q

What can you work out the half-life of a reaction?

A
  • Plot the concentration-time graph
  • Draw lines across from the y-axis at points where the concentration has halved
  • Read off the time taken
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33
Q

For what order reaction is half-life always constant?

A

First order

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34
Q

What is the experimental method you could use to work out the order of a reaction?

A

Initial rates method

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35
Q

What is the initial rates method?

A

A technique that lets you use the initial rate of an experiment to work out the order of a reaction.

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36
Q

Describe in general how the initial rates method works for calculating the order of a reaction.

A

1) Carry our separate experiments using different concentrations of a given reactant. Change only this one variable.
2) Work out the how the change in initial concentration affects the initial rates and therefore the order of the reaction. (e.g. If the rate doubles when the concentration doubles, it is first order with respect to that reactant)

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37
Q

What are the two ways the initial rates method can be done experimentally?

A
  • Continuous monitoring

* Clock reaction

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38
Q

Describe how the initial rates method of determining the order of a reaction can use continuous monitoring.

A
  • A continuous monitoring method is carried out at various concentrations of a reactant (keeping all other variables the same).
  • Concentration-time or volume-time graph is plotted for each one
  • Initial rate is found by finding the gradient at the start of each line
  • The change in initial rate between concentrations is used to find the rate of the reaction (e.g. If the rate doubles when the concentration doubles, it is first order with respect to that reactant)
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39
Q

Describe how the initial rates method of determining the order of a reaction can use a a clock reaction.

A
  • A reaction is chosen so that there is a clear end point (e.g. a colour change) when a sufficient amount of product has been formed
  • The reaction is carried out with various concentrations of a reactant and the time for the endpoint to be reached is measured
  • This gives the initial rate at each concentration
  • The change in initial rate between concentrations is used to find the rate of the reaction (e.g. If the rate doubles when the concentration doubles, it is first order with respect to that reactant)
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40
Q

What does a clock reaction usually have?

A

A clear, sharp endpoint (such as a colour change)

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41
Q

What are some assumptions of a clock reaction?

A
  • Concentration of rah reactant doesn’t change
  • Temperature stays constant
  • When the end point is seen, the reaction has not proceeded too far
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42
Q

What is an example of a clock reaction?

A

Iodine clock

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43
Q

What is the equation for the iodine clock reaction?

A

H₂O₂(aq) + 2I⁻(aq) + 2H⁺(aq) -> 2H₂O(l) + I₂(aq)

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44
Q

What is the reaction for the instantaneous reaction reaction in the iodine clock (that stops there being an instant colour change)?

A

2S₂O₃²⁻(aq) + I₂(aq) -> 2I⁻(aq) + S₄O₆⁻(aq)

45
Q

What are the reactants in the iodine clock reaction?

A 
• Hydrogen peroxide
• Iodine ions (e.g. potassium iodide)
• Acid
• Sodium thiosulfate
(• Starch indicator)
46
Q

Describe the principle on which the iodine clock works.

A

1) A small amount of sodium thiosulfate solution and starch are fed to an excess of hydrogen peroxide and iodide ions in acid solution. Starch acts as an indicator.
• H₂O₂(aq) + 2I⁻(aq) + 2H⁺(aq) -> 2H₂O(l) + I₂(aq)
2) The sodium thiosulfate instantaneously with any iodine formed.
• 2S₂O₃²⁻(aq) + I₂(aq) -> 2I⁻(aq) + S₄O₆⁻(aq)
3) This means that all of the sodium thiosulphate has to be used up before the iodine can be detected.
4) Once this happens, the starch turns the solution blue-black.
5) Varying the concentration of the reactants will give different times for the colour change.
6) This allows the rates to be calculated for each concentration, so the order can be worked out.

47
Q

What colour change is observed in the iodine clock reaction?

A

From colourless to blue-black (due to the starch indicator).

48
Q

Describe how to carry out the iodine clock reaction in a lab to find the order with respect to potassium iodide.

A

1) Rinse a clean pipette with sulfuric acid. Then, use this pipette to transfer a small amount of sulfuric acid, of known concentration (e.g. 0.25 mol/dm³) to a clean beaker.
2) Using a clean pipette or measuring cylinder, add distilled water to the beaker.
3) Using a dropping pipette, add a few drops of starch solution.
4) Measure a known amount of potassium iodide solution of a known concentration, using either a pipette or a burette, rinsed with potassium iodide solution. Transfer this volume to the reaction vessel.
5) Next, using a pipette rinsed with sodium thiosulfate solution, or clean measuring cyclinder, add sodium thiosulphate to the reaction vessel. Ensure everything else is in excess. Swirl the contents of the beaker so all the solutions are evenly mixed.
6) Finally, rinse a pipette with hydrogen peroxide solution. Then, use the pipette to transfer (an excess of) hydrogen peroxide solution to the reaction vessel while stirring the contents and simultaneously start a stop watch.
7) Continue to stir, and stop the stop watch when the contents of the beaker turn from colourless to blue-black. Record this time in a results table, along with the quantities of sulfuric acid, water, potassium iodide and sodium thiosulphate solutions you used in the experiment.
8) Repeat the experiment, varying the volume of potassium iodide solution. Keep the volume of sulfuric acid, sodium thiosulphate and hydrogen peroxide constant and use varying amounts of distilled water in each experiment so the overall volume of the reaction mixture remains constant.
9) Process the data to work out the order with respect to potassium iodide.

49
Q

In the iodine clock experiment, what is the purpose of the water?

A

It allows the concentration of the reactants to be varied in situ, while maintaining the volume of the mixture.

50
Q

What is the chemical equation for the reaction between propanone and iodine?

A

CH₃COCH₃(aq) + I₂(aq) —(H⁺)—> CH₃COCH₂I(aq) + H⁺(aq) + I⁻(aq)

(NOTE: This is done in acidic conditions!)

51
Q

Describe how you can investigate how the rate of reaction between propanone and iodine, as the concentration of iodine varies.

A

Continuous monitoring titrimetric method:
• Set up the reaction (in acidic conditions)
• Take samples at regular intervals. Stop the reaction in each by adding sodium hydrogencarbonate to neutralise the acid.
• Then titrate each sample against sodium thiosulphate and starch to work out the concentration of the iodine.
• The whole thing is done several times, changing the concentration of only 1 reactant each time.

52
Q

In the continuous monitoring titrimetric method for the reaction of propane and iodine, what is the chemical used to stop the reaction in each sample?

A

Sodium thiosulphate

53
Q

Remember to revise continuous monitoring titrimetric methods.

A

Pg 185 of revision guide

54
Q

In the reaction between propanone and iodine, what is it important to remember?

A

It must be done in acidic conditions.

55
Q

What do rate equations tell you?

A

How the rate is affected by the concentrations of the reactants.

56
Q

For the reaction A + B -> C + D, what is the general rate equation?

A

Rate = k x [A]^m x [B]^n

Where:
• k = Rate constant
• m = Order of reaction with respect to A
• n = Order of reaction with respect to B

57
Q

What is the symbol for the rate constant?

A

k

58
Q

CH₃COCH₃(aq) + I₂(aq) —(H⁺)—> CH₃COCH₂I(aq) + H⁺(aq) + I⁻(aq)

The reaction is first order with respect to propanone and H⁺, and zero order with respect to iodine. Write down the rate equation.

A

Rate = k[CH₃COCH₃][H⁺]

59
Q

Are spectator ions usually included in rate equations?

A

No

60
Q

How can you work out the rate equation for a reaction from an initial rates method?

A
  • Look at two experiments where the concentration of everything is the same, except for 1 reactant
  • Look at how the rate changes when the concentration of this reactant changes
  • This can help you work out the order with respect to that reactant
  • Do this for all of the reactants
61
Q

Remember to practise working out the orders of reaction using an initial rates method.

A

Pg 186 of revision guide

62
Q

Is the rate equation a fixed quantity?

A

No, it changes if the temperature changes.

63
Q

How can you work out the rate constant and it’s units?

A
  • Write our the rate equation.
  • Insert the concentrations and the rate.
  • Rearrange the equation and calculate the value of k.
  • Find the units for k by putting the other units in the rate equation.
64
Q

The reaction below was found to be second order with respect to NO and zero order with respect to CO and O₂. The rate is 1.76 x 10⁻³ mol dm⁻³ s⁻¹ when [NO] = [CO] = [O₂] = 2.00 x 10⁻³ mol dm⁻³.

NO(g) + CO(g) + O₂(g) -> NO₂(g) + CO₂(g)

Find the rate constant.

A
  • Rate = k[NO]²[CO]⁰[O₂]⁰ = k[NO]²
  • 1.76 x 10⁻³ = k x (2.00 x 10⁻³)²
  • k = (1.76 x 10⁻³) / (2.00 x 10⁻³)² = 440
  • mol dm⁻³ s⁻¹ = k x (mol dm⁻³)²
  • k = dm³ mol⁻¹ s⁻¹

So k = 440 dm³ mol⁻¹ s⁻¹

65
Q

What is the rate-determining step?

A

The slowest step in a multi-step reaction, which determines the rate.

66
Q

What are the rules for picking out which reactants from a chemical equation are involved in the rate-determining step?

A
  • If a reactant appears in the rate equation, it must affect the rate. So this reactant, or something derived from it, must be in the rate-determining step.
  • If the reactant doesn’t appear in the rate equation, then it isn’t involved in the rate-determining step (and neither is anything derived from it).
67
Q

If a reactant is in the rate equation, what does this tell you about the rate-determining step?

A

This reactant, or something derived from it, must be in the rate-determining step.

68
Q

If a reactant is not in the rate equation, what does this tell you about the rate-determining step?

A

This reactant isn’t involved in the rate-determining step (and neither is anything derived from it).

69
Q

What are some important points to remember about the rate-determining step and mechanism?

A

1) The rate-determining step doesn’t have to be the first step in a mechanism
2) The reaction mechanism can’t usually be predicted from just the chemical equation

70
Q

What does the order of a reaction with respect to a reactant tell you about the rate-determining step?

A

The order of a reaction with respect to a reactant shows the number of molecules of that reactant which are involved in or before the rate-determining step.

71
Q

The mechanism for the reaction between chlorine free radicals and ozone, O₃, consists of two steps:

1) Cl• + O₃ -> ClO• + O₂ (slow)
2) ClO• + O -> Cl• + O₂ (fast)

Determine the rate equation.

A

Rate = k[Cl•][O₃]

72
Q

Can you predict the mechanism for a reaction from its rate equation?

A

Yes

73
Q

Remember to revise how to predict the mechanism for a reaction from the rate equation.

A

Pg 188 of revision guide

74
Q

CH₃COCH₃(aq) + I₂(aq) —(H⁺)—> CH₃COCH₂I(aq) + H⁺(aq) + I⁻(aq)

The rate equation is:
Rate = k[CH₃COCH₃][H⁺]

What can you say about the reaction?

A

1) Propanone and H⁺ (or something derived from them) must be in the rate-determining step
2) Iodine is not involved until after the rate-determining step
3) The rate-determining step must use 1 molecule of propane and H⁺ (or something derived from them)
4) H⁺ is a catalyst, so it must be regenerated in another step

75
Q

Remember to practice doing the mechanism for the example of rate-determining steps on pg 189 of revision guide.

A

Do it.

76
Q

What is produced when a halogenoalkane reacts with sodium hydroxide?

A

Alcohol and halide ions

77
Q

What type of reaction in the reaction of a halogenoalkane with sodium hydroxide?

A

Nucleophilic substitution

78
Q

What type of reactions do halogenoalkanes tend to undergo?

A

Nucleophilic substitution

79
Q

What are the two types of mechanism for nucleophilic substitution?

A
  • SN1

* SN2

80
Q

What is an SN1 reaction?

A

A nucleophilic substitution that involves only 1 molecule or ion in the rate-determining step.

81
Q

What is an SN2 reaction?

A

A nucleophilic substitution that involves 2 species in the rate-determining step.

82
Q

What type of nucleophilic substitution do primary halogenoalkanes react by?

A

SN2

83
Q

What type of nucleophilic substitution do secondary halogenoalkanes react by?

A

SN1 and SN2

84
Q

What type of nucleophilic substitution do tertiary halogenoalkanes react by?

A

SN1

85
Q

What type of nucleophilic substitution do primary, secondary and tertiary halogenoalkanes react by?

A
  • Primary -> SN2
  • Secondary -> SN1 and SN2
  • Tertiary -> SN1
86
Q

What is the equation for the reaction of bromoethane with hydroxide ions?

A

CH₃CH₂Br + OH⁻ -> CH₃CH₂OH + Br⁻

87
Q

Describe and explain the rate equation for bromoethane with hydroxide ions.

A
  • Rate = k[CH₃CH₂Br][OH]

* This is because the reaction is SN2 and therefore both reactants are involved in the rate-determine step

88
Q

Describe the mechanism for a primary halogenoalkane reacting with OH⁻ ions.

A

ALL STEP 1:
• Arrow from lone pair on OH⁻ to positive carbon in halogenoalkane
• Arrow from C-X bond to the X (halogen)
• Alcohol and X⁻ are formed

(NOTE: This is SN2. See diagram pg 190 of revision guide.)

89
Q

When drawing out an SN2 mechanism, what is it important to remember?

A

There is technically a transition state.

See diagram pg 190 of revision guide

90
Q

What is the equation for the reaction of 2-bromo-2-methylpropane with hydroxide ions?

A

(CH₃)₃CBr + OH⁻ -> (CH₃)₃COH + Br⁻

91
Q

Describe and explain the rate equation for 2-bromo-2-methylpropane with hydroxide ions.

A
  • Rate = k[(CH₃)₃CBr]

* This is because the reaction is SN1 and therefore only 1 reactant is involved in the rate-determine step

92
Q

Describe the mechanism for a secondary halogenoalkane reacting with OH⁻ ions.

A

STEP 1:
• Arrow from the C-X bond in the halogenoalkane to the X (halogen)
STEP 2:
• Arrow from lone pair on the OH⁻ to the positive carbon in the halogenoalkane
• Alcohol and X⁻ are formed

(NOTE: This is SN1. See diagram pg 191 of revision guide.)

93
Q

In an SN1 reaction between a halogenoalkane and OH^- ions, what is the thing that the rate is dependent on?

A

Only on the concentration of the halogenoalkane.

94
Q

Remember to practise drawing out SN1 and SN2 mechanisms.

A

Pg 190 and 191 of revision guide.

95
Q

State the Arrhenius equation.

A

k = Ae^(-Ea / RT)

Where:
• k = Rate constant
• Ea = Activation energy (J)
• T = Temperature (K)
• R = Gas constant (= 8.31 J/K/mol)
• A = Another constant
96
Q

What happens to the rate constant (k) as the activation energy (Ea) gets bigger?

A

It gets smaller.

97
Q

Make the Arrhenius equation linear.

A
  • k = Ae^(-Ea / RT)
  • ln(k) = ln(A) - (Ea / RT)
  • ln(k) = Constant - (Ea / RT)
98
Q

How can the activation energy of a reaction be calculated using a graph derived from the Arrhenius equation? Why?

A
  • Linearising the Arrhenius equation gives: ln(k) = Constant - (Ea / RT)
  • Plot ln(k) against 1/T.
  • This means the gradient is -Ea/R.
  • From this, work out Ea.
99
Q

What is an Arrhenius plot?

A

A graph of ln(k) against 1/T.

100
Q

Remember to revise the Arrhenius plot example on pg 192 of revision guide.

A

Do it.

101
Q

In Arrhenius calculations, if no rate constant (k) value is given, what can be used instead?

A

1/t

Where t is the time for initial reaction to be complete. 1/t is essentially just the rate.

102
Q

For the iodine clock reaction, how can the activation energy be calculated using an Arrhenius plot?

A
  • The reaction is carried out at multiple temperatures (T) and the time for the colour change (t) is recorded for each one.
  • A graph of ln(1/t) against 1/T is plotted (since k is replaced by 1/t).
  • The gradient is equal to -Ea/R.
  • From this, calculate Ea.
103
Q

Remember to the iodine clock activation energy example on pg 192 and 193 of revision guide.

A

Do it.

104
Q

What are the two things that can be plotted on the axis of an Arrhenius plot?

A
  • ln(k) against 1/T

* ln(1/t) against 1/T

105
Q

How do catalyst increase the rate of a reaction?

A

By providing an alternative reaction pathway with a lower activation energy.

106
Q

What is an advantage and disadvantage of a heterogeneous catalyst?

A
  • Advantage: Easy to separate from the products and leftover reactants
  • Disadvantage: Can be poisoned
107
Q

What is poisoning of a catalyst?

A

When a substance clings to a catalyst’s surface more strongly than the intended reactant, preventing the catalyst from doing its job.

108
Q

Give an example of a catalyst being poisoned.

A

Sulfur can poison the iron catalyst used in the Haber process.

Inorganic Chemistry (21 decks)

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