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A level Chemistry > Unit 3.5 - Chemical Kinetics > Flashcards

Unit 3.5 - Chemical Kinetics Flashcards

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1
Q

Equation for the rate of a reaction

A

Rate of reaction = change in concentration/unit time

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2
Q

Unit of rate if measuring change in concentration over time

A

moldm-3s-1

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3
Q

Unit of rate if measuring change in volume over time

A

cms-1

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4
Q

Unit of rate if measuring change in mass over time

A

gs-1

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5
Q

Relationship between concentration and rate

A

Directly proportional

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6
Q

What does it mean that concentration is directly proportional to rate?

A

Double concentration = double rate

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7
Q

Which word do we always use when explaining rates

A

Time

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8
Q

How do we measure the rate of reaction for a change in concentration of reagent or product

A

By kinetic experiments

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9
Q

The change in concentration of what do we measure for the rate of a reaction?

A

Concentration of reagents or reactants

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10
Q

What’s the problem with measuring the rate of reaction through the change in concentration and what is done therefore?

A

Not always easy to follow the change in concentration
Usually another property that changes during the reaction is measured

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11
Q

How can we measure the rate of reaction at any instant in a reaction?

A

If continuous results are recorded, or many results over time, a graph is drawn
Th rate at any instant is measured using a tangent and a gradient at that point

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12
Q

What do we do to calculate the initial rate of a reaction?

A

Tangent at time = 0

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13
Q

How do we calculate the instantaneous rate of a reaction?

A

Tangent at that specific time and gradient

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14
Q

When is the rate of a reaction highest?

A

At the start of the reaction

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15
Q

Why is the rate of a reaction highest at the start of the reaction?

A

This is when there are the most reactants available for collision

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16
Q

List the things that he rate of a reaction is dependent on

A

Concentration
Pressure (for gas)
Temperature
Particle size
Catalysts
Light (sometimes)

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17
Q

What is the state that an element has to be in for it to be affected by pressure changes?

A

Gaseous

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18
Q

What does increasing the concentration or pressure (for a gas) do to the rate of reaction and why?

A

Increases the rate
Less distance between molecules in a given volume = increased number of collisions per unit time

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19
Q

What does a different particle size do to the rate of reaction? Explain

A

Smaller = increased rate
Increased surface area = closer together = more collisions per unit time

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20
Q

How do we measure the rate of a reaction?

A

We need to work out how much reactant has been used up or how much product has been produced

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21
Q

What properties can change with time that we can use to measure the rate of reaction?

A

Concentration of reactants or products
Mass of reactant
Volume or pressure of gas produced
pH
Colour change (or other electromagnetic absorption)

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22
Q

How do we work out the means of following a reaction?

A

Study the equation

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23
Q

Means of following this reaction:
CaCO3 + 2HCl —> CaCl2 + H2O + CO2

A
  1. Gas syringe to measure the change in volume over time
  2. Measure change in mass over time with scales (CO2 is a heavy gas)
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24
Q

What property about CO2 makes it useful for following the rate of a reaction? Explain

A

CO2 is a heavy gas
Can measure the change in mass over time with scales to measure the rate of the reaction

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25
Q

Means of following this reaction:
CH3CO2C2H5 + H2O —> CH3COOH + C2H5OH

A
  1. Measure change in pH over time (ethanoic acid is involved)
  2. Sampling and quenching
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26
Q

Means of following this reaction:
CH3COCH3 + I2 —> CH3COCH2 + HI

A

Measure change in colour over time using a colorimeter
(Brown solution of I2 disappears during the reaction)

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27
Q

Describe I2

A

Brown solution

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28
Q

Means of following this reaction:
2KI + H2O2 + 2H+ —> I2 + 2H2O

A
  1. Monitor change in pH over time
  2. Iodine-clock reaction - measure change in colour over time using a colorimeter
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29
Q

Example of a spectroscopic method to follow the rate of a reaction

A

NMR has been used to study the rates at which drugs act within the body

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30
Q

What type of method is using NMR to study the rate of a reaction?

A

Spectroscopic

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31
Q

When is the iodine-clock technique used to measure the rate of reaction?

A

In a reaction where iodine is produced

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32
Q

Give a description of the theory behind the iodine-clock reaction

A

Iodine ions can be oxidised to iodine at a measurable rate
Iodine gives a strongly coloured blue/black complex with starch solution
If a given amount of thiosulfate ion (with which iodine reacts very rapidly with to reform iodide ions) is added, the solution will be colourless until enough iodine has been formed to react with all of the thiosulfate
The time taken for this to occur acts like a clock to measure the rate of iodide ions being oxidised

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33
Q

What are iodide ions oxidised to at a measurable rate?

A

Iodine

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34
Q

In what solution does iodine give a strongly coloured blue/black complex?

A

Starch solution

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35
Q

Colour of iodine in starch

A

Blue/black

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36
Q

Iodine indicator

A

Starch

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37
Q

What does iodine do with thiosulfate ions?

A

Reforms iodide ions

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38
Q

When is the solution colourless during the iodine clock reaction?

A

When the iodine is reacting rapidly with the thiosulfate ions to reform iodide ions - it’s colourless until enough iodine has been formed to react with all of the thiosulfate

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39
Q

The rate of what is essentially being measured during the iodine-clock reaction?

A

The rate of iodide ions being oxidised

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40
Q

When does the indicator remain colourless during the iodine clock reaction?

A

When the thiosulfate reacts with the iodine, the indicator remains colourless

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41
Q

At which point does the iodine in the reaction mixture and the starch turn blue black during the iodine clock reaction?

A

The instant the last thiosulfate ion has reacted, there is iodine in the reaction mixture and the starch turns blue/black

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42
Q

Thiosulfate ion

A

S2O32-

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43
Q

How could reaction rates be compared using the iodine clock reaction?

A

If the amount of thiosulfate is kept constant, can compare the times taken to change colour

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44
Q

Suitable example of the iodine clock reaction

A

Oxidising iodide ions by hydrogen peroxide in acid solution

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45
Q

Equation for oxidising iodine ions by hydrogen peroxide in acid solution

A

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

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46
Q

Equation for iodine reforming iodide ions with thiosulfate ions

A

I2 (aq) + 2s2O32- (aq) —> 2I- (aq) + S4O62- (aq)

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47
Q

Why is measuring the time taken to produce excess iodine in the iodine clock reaction only an approximation for the rate as the reaction proceeds?

A

The reaction is faster at first (higher concentration of reactant)
The average rate is measured

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48
Q

When is a reaction fastest and why?

A

At the start
Higher concentration of reactant

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49
Q

What gives a measure of the initial rate of reaction in the iodide clock experiment?

A

The time taken from mixing the reactant to the formation of the dark blue starch solution

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50
Q

What does the time taken from mixing the reactant to the formation of the dark blue starch solution in the iodine clock reaction give a measure of?

A

The initial rate of reaction

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51
Q

Initial rate equation

A

Initial rate ∝ 1/t

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52
Q

What is proportional to the reciprocal of time (1/t)?

A

initial rate

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53
Q

When does the beginning of the concentration-time curve approximate to a straight line in the iodine clock reaction?

A

If the amount of sodium thiosulfate used is small and is equivalent to m moldm^3 of iodine

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54
Q

Describe the beginning of the concentration time curve of a concentration of iodine against time graph

A

Straight line

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55
Q

How can we measure the rate of reaction of the isomerisation of cyclopropane to propane?

A

Can measure concentration over time using gas chromatography at a constant temperature

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56
Q

Under which condition must gas chromatography be done?

A

Constant temperature

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57
Q

Sampling

A

A process involving removing small samples of the reaction mixture at regular time intervals

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58
Q

Quenching

A

Where samples are placed in ice water to lower the concentration of the reactants and to stop the reaction continuing

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59
Q

How can the samples be tested after sampling and quenching?

A

Using other techniques like titration

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60
Q

What is sampling and quenching used to measure?

A

Changes in concentration

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61
Q

Why do we use ice water when sampling and quenching?

A

Cools the sample
Slows down the motion of particles (significantly - don’t actually stop completely)
Can monitor the exact concentrations at that moment in time

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62
Q

What can we titrate CH3CO2C2H5 with and what is the indicator?

A

NaOH
Methyl orange

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63
Q

What happens to the concentration of the product during the sampling and quenching method?

A

Concentration increases throughout the reaction
Starts to flatline when the reaction is stopping

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64
Q

What do we plot on a graph after sampling and quenching?

A

Concentration against time

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65
Q

What is the ratio between the moles of NaOH (titrant) and ethanoic acid produced during a sampling and quenching reaction example?

A

1:1

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66
Q

Advantages of quenching

A

Accurate
Easy to perform
No need for sophisticated equipment

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67
Q

Disadvantages of quenching

A

Laborious process (involves 5 titrations)

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68
Q

Equation for the reaction between 1-bromobutane and sodium hydroxide

A

CH3CH2CH2CH2Br + OH- —> CHCH2CH2CH2OH + Br-

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69
Q

What type of reaction is the reaction between 1-bromobutane and sodium hydroxide?

A

Nucleophillic substitution reaction

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70
Q

What does the fact that rate ∝ [OH-] in a reaction mean?

A

The rate of reaction depends on the concentration of OH-
If the concentration of OH- were doubled, the rate of reaction would double

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71
Q

Decomposition of hydrogen peroxide equation

A

2H2O2 —> 2H2O + O2

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72
Q

What is rate proportional to in the decomposition of hydrogen perioxide?

A

The concentration of the hydrogen peroxide

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73
Q

How can we form a rate equation?

A

By changing the proportionality sign for an equal sign and adding k, the rate constant

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74
Q

Rate equation for the decomposition of hydrogen peroxide

A

Rate = k[H2O2]

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75
Q

What does a rate equation give?

A

The dependence of the rate of the reaction on the concentration of the species in the rate equation

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76
Q

What is k in rate equations?

A

The rate constant

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77
Q

What is k (in rate equations) unique for?

A

Every set of chemical reactions

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78
Q

What is k (rate constant) independent of?

A

Concentration and time

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79
Q

What is the only thing that affects the value of k (rate constant)?

A

Temperature

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80
Q

What does the value of k (rate constant) depend on?

A

The reaction being studied
The temperature

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81
Q

What odes a higher k value mean?

A

Faster reaction

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82
Q

Example of a reaction with a high k value

A

Fireworks

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83
Q

Example of a reaction with a low k value

A

Rusting

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84
Q

How can the rate equations deduced?

A

From kinetic experiments

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85
Q

How can rate equations not be deduced?

A

From chemical equations

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86
Q

What does it mean if a reaction is first order with respect to something?

A

The rate of reaction depends on the concentration of that thing to the power one

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87
Q

What is the reaction 2H2O2 —> 2H2O + O2’s order with respect to hydrogen peroxide?

A

First order

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88
Q

How is the order of reaction found?

A

By experiment

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89
Q

How is the order of a reaction not found?

A

By chemical equations

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90
Q

General rate equation for a reaction whose rate depends on the concentration of reagents A and B

A

Rate = k[A]^x[B]^y

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91
Q

In Rate = k[A]^x[B]^y
What is x?

A

The order of the reaction with respect to A

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92
Q

In Rate = k[A]^x[B]^y
What is y?

A

The order of the reaction with respect to B

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93
Q

Order of reaction

A

The power to which the concentration term must be raised to fit the rate equation - is the exponent

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94
Q

What values do the order of a reaction usually have?

A

0, 1 or 2 for each reagent

95
Q

What does it mean if something has an order of zero?

A

It doesn’t affect the rate

96
Q

How do we work out the overall order of a reaction?

A

Is equal to the sum of the powers of the concentration terms in the rate equation

97
Q

Can the overall rate of a reaction be more than 2?

A

Yes

98
Q

The concentrations of what do we show in a rate equation?

A

The reactants

99
Q

How do we know if something is to the power of one in orders of reaction?

A

If something is to the power of 1, it isn’t usually shown

100
Q

Can the rate equation be deduced from the chemical equation?

A

No

101
Q

How can we find orders?

A

Perform experiments

102
Q

Why can’t we deduce rate equations from chemical equations?

A

Chemical reactions may be a one step process or may involve a number of steps

103
Q

Rate determining step

A

The slowest step in a reaction

104
Q

What is the slowest step of a reaction known as?

A

The rate determining step

105
Q

Which step determines the rate equation? Explain

A

Chemical reactions may be a one step process or may involve a number of steps
In any reaction, one step will e slower than all the other steps and it is this step which determines the rate equation

106
Q

How do you determine the rate of any clock reaction?

A

Measure time taken for the colour change to occur
rate = 1/time

107
Q

Effect of an increased temperature on a rate equation + explain

A

No effect
Rate equation is independent on temperature

108
Q

Catalyst of the decomposition of ammonia

A

Platinum

109
Q

Decomposition of ammonia equation

A

2NH3 —> N2 + 3H2

110
Q

If the decomposition of ammonia is zero order with respect to ammonia, what does this mean?

A

The rate is independent on the concentration of ammonia

111
Q

Rate equation for zeroth order reactions

A

Rate = k

112
Q

Units of k in zeroth order reactions

A

moldm^-3s^-1

113
Q

What happens to the rate of reaction when the concentration of ammonia is doubled in the decomposition of ammonia? Why?

A

The rate of reaction is unchanged
The reaction is zero order with respect to the ammonia

114
Q

What is the relationship between the concentration of the reagent and the rate of reaction in zeroth order reactions and what does this mean?

A

The rate of reaction is independent of the concentration of the reagent - rate stays the same as the concentration changes

115
Q

Describe the line on a rate v.s concentration graph for a zeroth order reaction

A

Horizontal straight line

116
Q

Give 2 examples of first order reactions

A

Decomposition reactions
Radioactive decay

117
Q

Rate equation for first order reactions (e.g - decomposition of dinitrogen perioxide)

A

Rate = k[N2O5]

118
Q

What order is the dinitrogen perioxide in its decomposition reaction?

A

First order

119
Q

Units of k in first order reactions

A

s^-1

120
Q

How do we know that the units are s^-1 with first order reactions?

A

k = rate/[N2O5] = moldm-3s-1/moldm-3 = s-1

121
Q

Relationship between rate and concentration in first order reactions

A

Directly proportional

122
Q

What happens when the concentrate of dinitrogen perioxide is doubled in its decomposition reaction and why?

A

Reaction is first order with respect to dinitrogen peroxide
As concentration is doubled, rate is doubled

123
Q

Decomposition of ninitrogen peroxide

A

2N2O5 —> 4NO2 + O2

124
Q

Explain the graph of rate v.s concentration for first order reactions

A

Straight line through the origin (as rate is zero when concentration is zero)

125
Q

Why is the line through zero on rate v.s concentration graphs?

A

Rate is zero when concentration is zero

126
Q

Half life of first order reactions

A

Constant

127
Q

2 ways to show that a reaction is first order

A
  1. Tangent at 0.1 and tangent at 0.05, and compare rates
  2. Work out the half life and check that it’s constant
128
Q

Decomposition of nitrogen dioxide

A

2NO2 —> 2NO + O2

129
Q

What order is the decomposition of nitrogen dioxide?

A

3rd order with respect to the nitrogen dioxide

130
Q

Rate equation for 2nd order reactions (e.g - the decomposition of nitrogen dioxide)

A

Rate = k[NO2]^2

131
Q

Units of k in second order reactions

A

mol-1dm3s-1

132
Q

Why is the unit of k mol-1dm3s-1 in 2nd order reactions?

A

k = rate/concentration
k = moldm-3s-1/(moldm-3)2
= 1/moldm-3 x s-1
=mol-1dm3s-1

133
Q

What is rate proportional to in second order reactions?

A

Concentration^2

134
Q

In the decomposition of nitrogen dioxide (a 2nd order reaction), if the concentration of nitrogen dioxide is doubled, what happens to the rate of reaction?

A

Is quadrupled, increased fourfold

135
Q

Describe the curve on a concentration v.s rate graph of a 2nd order reaction

A

A curved line starting at the origin as rate is zero when concentration is zero

136
Q

What will happen to the rate of a first order reaction if the concentration is increased by a factor of 3?

A

The rate will increase by a factor of 3 too

137
Q

What happens to the rate of a 2nd order reaction if the concentration is increased by a factor of 3? Explain

A

3^2
= increases by a factor of 9

138
Q

Describe the half life of second order reactions

A

Isn’t constant
Half life increases as the concentration decreases (exponential decay curve)

139
Q

Describe the curve of rate v.s concentration for a second order reaction

A

Exponential decay curve

140
Q

If we’re asked to work out he order of reaction with respect to a specific element, what do we need to do?

A

Choose the experiments where the concentration of the other reactant is constant

141
Q

Equation to work out k (and where does this come from?)

A

k = rate/concentration
Rearranged rate = k[concentration]

142
Q

What do we need to do with a rate-based question where we’re given pH values?

A

Should work out [H+]

143
Q

Equation to work out [H+]

A

[H+] = 10^pH

144
Q

How many steps do chemical reactions involve?

A

May involve one stop or a number of steps

145
Q

Which step determines the rate equation?

A

The slowest step - the rate determining step

146
Q

What does the rate determining step determine?

A

The rate equation

147
Q

How can we identify the rate determining step of a chemical reaction?

A

It’s the slowest step

148
Q

What does the rate determining step tell us?

A

Which particles are reacting in the reaction

149
Q

Which species specifically is a rate dependent on?

A

The species in the rate determining step

150
Q

What is dependent on the species in the rate determining step?

A

The rate

151
Q

What information do we obtain about the rate determining step?

A

According to collision theory:
1.) the reacting particles must collide
2.) the particles must have sufficient energy for reaction (the activation energy)

152
Q

What must colliding particles have in order to reacti?

A

Sufficient energy - the activation energy

153
Q

How much energy is sufficient for a reaction to occur?

A

An amount of energy equivalent to the activation energy of a reaction

154
Q

If something is first order with respect to something, what does it mean?

A

That only this element is involved in the rate determining step

155
Q

How do we know if only one element is involved in a rate determining step?

A

If the rate equation shows only that element with first order respect to it

156
Q

What does a rate equation tell us (relating to the rate determining step)?

A

How many particles must collide in the rate determining step

157
Q

How many particles collide in 2nd order reactions?

A

2

158
Q

How many particles collide in 3rd order reactions?

A

3

159
Q

How many particles in the rate determining step of first order reactions?

A

1

160
Q

When is there only 1 particle in the rate determining step?

A

1st order reactions

161
Q

How many particles in the rate determining step for 1st order reactions?

A

1

162
Q

Write a possible mechanism by which explains the rate equation of
Rate = k[A]
For the reaction
A + B —> C + D

A

A —> A* slow
A* + B —> C + D fast

163
Q

What does the rate determining step of a reaction determine?

A

The rate of a reaction

164
Q

When is something not involved in the rate equation?

A

If it’s not involved in the rate determining step

165
Q

What is an element also not included in if it’s not included in the rate determining step?

A

The rate equation

166
Q

What does a species have to be present for a reaction to be dependent on it?

A

in the rate equation

167
Q

What will the rate determining step and the rate equation include for a one-step process?

A

Both reagents

168
Q

When will the rate determining step and the rate equation include both reagents?

A

One step processes

169
Q

Equation for the hydrolysis of 1-bromobutane by aqueous sodium hydroxide

A

CH3CH2CH2CH2Br + OH- —> CH3CH2CH2CH2OH + Br-

170
Q

What type of process is the hydrolysis of 1-bromobutane by aqueous sodium hydroxide?

A

One step reaction

171
Q

Give the rate equation for the following one step reaction
CH3CH2CH2CH2Br + OH- —> CH3CH2CH2CH2OH + Br-

A

Rate = k[CH3(CH2)3Br][OH-]

172
Q

When can the rate equation be deduced from a single step?

A

If that step is a rate determining step

173
Q

If the single step of a one step reaction is the rate determining step, what an be deduced?

A

The rate equation

174
Q

What can’t be determined from a chemical equation?

A

Rate equations

175
Q

What can’t rate equations be deduced from?

A

Chemical equations

176
Q

If the rate equation is
Rate = k[CH3COOCH3][H+]^2
What are the reactants of the rate determining step?

A

CH3COOCH3 + 2H+ —> products

177
Q

Is it possible to identify conclusively the product of the rate determining step in every case?

A

No

178
Q

What can we do as we can’t always conclusively identify the products of a rate-determining step?

A

Can suggest products

179
Q

What must we do when suggesting products for a rate determining step?

A

Make sure that the equation balances just like any other

180
Q

Suggest products for the following rate determining step

A

C3H7I + Br- ——> C3H7Br + I-

181
Q

What must combine to form the overall equation of a reaction?

A

All of the steps in the mechanism, including the rate determining step

182
Q

If a rate determining step has the following reactants:
CH3COOCH3 + 2H+ —> products
What will the rate equation be?

A

Rate = k[CH3COOCH3][H+]^2

183
Q

How can we prove or disprove a reaction mechanism?

A

Can use the information derived from the kinetics of a reaction

184
Q

What can we prove or disprove from the kinetics of a reaction?

A

A reaction mechanism

185
Q

How can we actually work out what the reactants are for a rate determining step?

A

By looking at the kinetics of a reaction

186
Q

What can we work out using kinetics?

A

The reactants for the rate determining step of a reaction

187
Q

How do we know if a proposed mechanism for a reaction is correct?

A

If our proposed mechanism does not have a step with the reactants for the rate determining step, it cannot be the correct mechanism

188
Q

Which reactants is it vital that a mechanism has in it for it to be correct?

A

The reactants for the rate determining step

189
Q

What do the coefficients in a chemical reaction show?

A

All of the reactants in all steps, including those not in the rate equation

190
Q

Which species do orders show?

A

Only those present in the rate determining step

191
Q

Explain why the orders of reaction in a rate equation do not always correspond to the coefficients in a chemical equation

A

Coefficients —> all reactants in all steps, including those not in the rate equation
Orders —> species present in the rate determining step

192
Q

When considering the reaction
BrO3- + 6H+ + 6Br- —> 3Br2 + 3H2O
With rate equation
Rate = k[BrO3-][Br-][H+}^2
How can you tell that this reaction involves more than one step?

A

Coefficients of equation don’t match up with the orders of the rate equation
13 species in chemical equation (LHS), 4 in rate equation

193
Q

Which side of a reaction do we look at the coefficients to compare to them with the orders of the rate equation?

A

Left

194
Q

What’s the only thing we need to remember to consider when writing out possible mechanisms for reactions using rate equations?

A

Make sure that the equation balances

195
Q

Consider the following reaction:
NO2 + CO —> NO + CO2
Rate equation:
Rate = k[NO2]^2
Suggest a likely rate determining step for the reaction and hence suggest a two-step mechanism for this reaction. Explain.

A

NO2 + NO2 —> NO + NO3 slow step (rate determining)
NO3 + CO—> NO + CO2
or
NO2 + NO2 —> intermediates
Intermediates —> NO + CO2

We needed to make sure that the reactants had 2 NO2’s since it’s squared in the rate equation

196
Q

If something is squared in a rate equation, what do we need to ensure we do for the rate determining step reaction?

A

Include it twice

197
Q

What does this rate equation show?
Rate = k[A][B]^2

A

It shows how the rate is affected by the concentrations of the reactants [A] and [B]

198
Q

What does the rate constant actually show in a rate equation?

A

How other variables other than the concentration of the reactants affect the rate

199
Q

What effects the value of k (rate constant)?

A

Change in temperature
Addition of catalyst

200
Q

How is k expressed mathematically?

A

Using the Arrhenius equation

201
Q

What is the Arrhenius equation for?

A

Expressing k mathematically

202
Q

Arrhenius equation

A

k = Ae^-(Ea/RT)

203
Q

What does the whole of e^-(Ea/RT) represent in the Arrhenius equation?

A

The fraction of the molecules in a gas which have energies equal to or greater than the activation energy at a stated temperature

204
Q

What does the frequency factor A in the Arrhenius equation include?

A

Factors like the frequency of collisions and their orientation

205
Q

What includes factors like the frequency of collisions and their orientations in the Arrhenius equation?

A

A, the frequency factor

206
Q

What does A, the frequency factor, vary with and to what extent?

A

It varies slightly with temperature, although not much and is taken as constant across small temperature ranges

207
Q

When is A, the frequency factor, taken as constant?

A

Across small temperature ranges

208
Q

What is the only thing that A, the frequency factor, is effected by?

A

Large temperature changes

209
Q

Name for this equation
k = Ae^-(Ea/RT)

A

The Arrhenius equation

210
Q

Units of k (rate constant)

A

Depends on the overall order

211
Q

Units of T in Arrhenius equation

A

K

212
Q

R in the Arrhenius equation

A

Gas constant

213
Q

Ea in Arrhenius equation and unit

A

Activation energy
Jmol-1

214
Q

e in Arrhenius equation

A

Mathematical constant

215
Q

How do we obtain “e” for the Arrhenius equation?

A

Shift ln

216
Q

A in Arrhenius equation and units

A

Frequency factor
No units

217
Q

What is the frequency factor related to?

A

The frequency of collisions between particles

218
Q

When do we treat the frequency factor as a constant?

A

Over a limited range of temperatures

219
Q

When can a frequency factor vary?

A

If temperatures change significantly

220
Q

What type of term is ln?

A

Logarithmic
(Loge not log10)

221
Q

Write the Arrhenius equation in logarithmic terms

A

ln A = ln k + Ea/RT

222
Q

What is the only bit we put in brackets when putting in Arrhenius equation calculations into a calculator?

A

Multiplying R and T

223
Q

What do we do with the R and T when typing in Arrhenius equation calculations into a calculator?

A

Use brackets

224
Q

How would we rearrange the Arrhenius equation to find temperature?

A

T = Ea/(lnA-lnK) x R

225
Q

What happen when we take the natural log when rearranging the Arrhenius equation?

A

It cancels out e

226
Q

How do we rearrange the Arrhenius equation to work out the activation energy?

A

Ea = RT(lnA - lnK)

227
Q

How do we rearrange the Arrhenius equation to find A?

A

A = k/e^-Ea/RT

228
Q

How do we use a graphical method with the Arrhenius equation?

A

Convert and re-arrange the equation int the straight line formula, y = mx + c

229
Q

Put the Arrhenius equation into the straight line formula and explain what each feature now means

A

ln (k) = (Ea/R) 1/T + ln A
Y = mx + c

Y-axis = plot ln K
X-axis = plot 1/T
Gradient (m) = gives - (Ea/R)
Y-intercept (c) = gives ln A

230
Q

How do we work out the activation energy of a reaction using the Arrhenius equation in graphical form? Explain

A

Arrhenius equation in the equation of a straight line
ln (k) = (Ea/R) 1/T + ln A
y = mx + c

Gradient = Ea/R
So Ea must be gradient x R

231
Q

Do we use the big numbers from the equation when working out rate constants?

A

No - ignore them

232
Q

Another word for frequency factor in the Arrhenius equation

A

Arrhenius constant

233
Q

Why do we need to work out the concentration of hydrogen ions if we’re given pH in a determination of the rate equation question?

A

To work out whether [H+] needs to be included in the rate equation

234
Q

Order of a reaction

A

The power to which the concentration is raised

A level Chemistry (39 decks)

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