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A-Level OCR Chemistry > Module 5: Chapter 18 (Reaction Rates) > Flashcards

Module 5: Chapter 18 (Reaction Rates) Flashcards

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1
Q

What affects the rate of a chemical reaction?

A

concentration / pressure
temperature
surface area of a solid
use of a catalyst

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2
Q

Explain collision theory

A

only effective collisions lead to chemical reactions:

  • particles must collide with the correct orientation
  • the particles have sufficient energy to overcome the activation energy Ea barrier of the reaction (E > Ea)
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3
Q

Define activation energy

A

Activation energy Ea is the minimum energy required to start a reaction by the breaking of bonds.

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4
Q

the rate of a chemical reaction = …?

A

how fast a reactant is used up
OR
how fast a product is formed
(conc/time)

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5
Q

What is the order of a reaction?

A

Changing the concentration often changes the rate of reaction
the rate of reaction is proportional to the concentration of a particular reactant raised to a power

rate α [A]n

Different reactants in a chemical reaction can have different orders and hence affect the rate in different ways
Common orders:
Zero order (0)
First order (1)
Second order (2)
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6
Q

What is zero order?

A

rate α [A]0
The concentration of the reactant has no effect on the rate
any number raised to the power zero is 1: [A]0 = 1

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7
Q

What is first order?

A

rate α [A]1
The concentration of the reactant affects the rate of reaction:
doubling [A] (x2) increases reaction rate by a factor of 2 (21 = 2)
tripling [A] (x3) increases reaction rate by a factor of 3 (31 = 3)

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8
Q

What is second order?

A

rate α [A]2
The concentration of the reactant affects the rate of reaction:
doubling [A] (x2) increases reaction rate by a factor of 4 (22 = 4)
tripling [A] (x3) increases reaction rate by a factor of 9 (32 = 9)

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9
Q

What is the rate equation?

A 
rate = k x [A]m x [B]n
(k = rate constant (is the proportionality constant)
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10
Q

What is the overall order?

A

overall order = sum of orders with respect to each reactant

- gives the overall effect of the concentrations on the rate of reaction

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11
Q

What methods must be taken to produce a concentration-time graph?

A
  • Concentration-time graphs can be taken from continuous measurements of a suitable property during the course of a reaction
  • This property must be proportional to the concentration of the reactant or product
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12
Q

Continuous measuring of rate: gas

A
  • Monitoring by gas collection (volume of gas evolved)
  • using a gas syringe to collect gas produced
  • The gradient of a curve produced tells us the rate of the reaction
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13
Q

Continuous measuring of rate: mass

A
  • Monitoring of mass loss
  • Using a balance and time intervals
  • graph of mass loss against time
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14
Q

Continuous measuring of rate: colour

A
  • Monitoring by colour change
  • If one of the reacting substances or products has a colour, the intensity of this colour will change during the reaction.
  • The absorbance reading of a colorimeter provides a measure of the concentration.
  • The amount of light absorbed by a solution is measured
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15
Q

What is the method for colorimeters?

A
  • Choose a suitable filter (complimentary colour)
  • Zero the colorimeter with water/solvent
  • Make up a range of samples with known concentrations
  • Take absorbance readings and plot a calibration curve
  • Set up your reaction, take a small sample, put in a cuvette and place in colorimeter
  • Take absorbance readings at set time intervals
  • Read concentrations of reaction mixture off curve or just use absorbance values
    (-The order of reaction with respect to a coloured chemical can then be determined from the concentration-time graph)
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16
Q

Describe the concentration-time graphs for orders

A
  • rate of reaction is the gradient of a concentration- time graph
  • order with respect to a reactant can only be determined if all other reactant concentrations remain constant
    Zero order = straight line with negative gradient, rate of reaction does not change, value of gradient is equal to rate constant k
    First order = downward curve with decreasing gradient and the time for the concentration of the reactant to half (half-life) is constant
    Second order = downward curve, steeper at start but tailing off more slowly
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17
Q

Half life and first order

A
  • First order reactions have a constant half-life with the concentration halving every half-life (exponential decay)
  • t1/2 is the time taken for the concentration of a reactant to decrease to half its original value
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18
Q

Determination of k for a first order reaction from a concentration-time graph for a first order reaction

A

rate = k x [A]

k = ln 2 / t1/2

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19
Q

Rate-concentration graphs: Zero order

A
  • produces a horizontal straight line with no gradient
    Rate = k[A]^0
    rate = k
  • the interpet on the y-axis give the rate constant
  • reaction rate doesn’t change with increasing conc
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20
Q

Rate-concentration graphs: First Order

A
  • produces a straight-line graph through the origin
    rate = k[A]^1
    rate = k[A]
  • rate is directly proportional to the conc for a first order relationship
  • rate constant can be determined by measuring the gradient of the straight line of this graph
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21
Q

Rate-concentration graphs: Second Order

A
  • produces an upwards curve with increasing gradient
    rate = k[A]^2
  • rate constant cannot be obtained directly from the graph
  • by plotting a second graph of the rate against conc squared, the result is a straight line (gradient of this line is equal to rate constant k)
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22
Q

What is initial rate?

A

The instantaneous rate at the start of a reaction when t=0

- can be found using a tangent on t=0 on a conc-time graph

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23
Q

What is the clock reaction?

A

Convenient way of obtaining the initial rate of a reaction by taking a single measurement

  • the time t from the start of a reaction to when a visual hange is observed is measured
  • provided that there is no signif change in rate during this time, assumed that average rate over time (1/t) is the same as initial rate (initial rate is proportional to 1/t)
24
Q

What is half-life?

A

The time taken for the concentration of a reactant to decrease to half of its original value

25
Q

What is exponential decay?

A

First order reactions having a constant half-life with the concentration halving every half life

26
Q

How can a first order relationship be confirmed on a concentration-time graph?

A

By measuring successive half-lives, if they are the same the reaction is first order with respect to the reactant

27
Q

How can the rate constant be calculated from the half-life of a first order reaction (k)?

A

k (rate constant) = ln2 / half-life

28
Q

What is an iodine clock?

A
  • type of clock reaction
  • relies on formation of iodine in the presence of starch (colourless to blue-black colour)
  • aqueous sodium thiosulfate is added to delay the colour change (remove iodine as it forms)
  • initial rate proportional to 1/t
  • plot results on a rate-concentration graph
29
Q

Why do thiosulfate ions delay the colour change?

A

Thiosulfate ions react with iodine forming I- ions which are not coloured
As soon as all the thiosulfate has reacted, iodine will start to build up in solution producing colour

30
Q

Why is it important to use the same volume of thiosulfate ions each time?

A

The amount of iodine is removed each time by thiosulfate should be constant
This means that the rate is proportional to 1/t

31
Q

What is the reaction mechanism?

A

The series of steps that make up an overall reaction

32
Q

What is the rate-determining step?

A

The slowest step in the sequence

33
Q

Why is the rate-determining step important?

A
  • rate equation only includes reacting species involved in the RDS
  • order in the rate equation match the number of species involved in the RDS
  • therefore provides important evidence in supporting or rejecting a proposed reaction mechanism
34
Q

N2 + 3H2 -> 2NH3 why would this reaction likely be a multi-step reaction?

A

A one step reaction would require 4 molecules colliding at the same time, which is very unlikely

35
Q 
slow H2O2(aq) + Br-(aq) -> H2O(l) + BrO-(aq)
fast H2O2(aq) + BrO-(aq) -> H2O(l) + Br-(aq) + O2(g)

Determine the rate equation for this reaction?

A

rate = k[H2O2(aq)][Br-(aq)]

  • the rate equation only includes the reacting species in the RDS
36
Q 
slow H2O2(aq) + Br-(aq) -> H2O(l) + BrO-(aq)
fast H2O2(aq) + BrO-(aq) -> H2O(l) + Br-(aq) + O2(g)

Determine the overall reaction?

A

2H2O2(aq) -> 2H2O(l) + O2(g)

37
Q 
slow H2O2(aq) + Br-(aq) -> H2O(l) + BrO-(aq)
fast H2O2(aq) + BrO-(aq) -> H2O(l) + Br-(aq) + O2(g)

Identify the intermediate?

A

BrO-

38
Q 
slow H2O2(aq) + Br-(aq) -> H2O(l) + BrO-(aq)
fast H2O2(aq) + BrO-(aq) -> H2O(l) + Br-(aq) + O2(g)

What is the role of Br-?

A

A catalyst as it is used in the first step but regenerated in the second step, not used up overall

39
Q

2N2O(g) -> 2N2(g) + O2(g)
rate = k[N2O(g)]

Suggest a two-step mechanism

A

slow step: N2O(g) -> N2(g) + O(g)

fast step: N2O(g) + O(g) -> N2(g) + O2(g)

40
Q

What increases when temperature increases?

A

The rate constant, k

41
Q

What factors affect the rate constant k?

A

When temperature increases 2 factors contribute to increased rate and k…

  • increasing energy shifts Boltzmann diagram distribution to the right, increasing proportion of particles exceeding Ea (MAIN DETERMINANT)
  • particles move faster and collide more frequently
42
Q

What is the Arrhenius equation?

A

k = A e^-Ea/RT (e^ button on the calculator)

A is the pre-exponential factor (frequency factor)
e^etc is the exponential factor (linked to activation energy and temp)
R is the gas constant (8.314 Jmol^-1K^-1)
T is for temperature in Kelvin

43
Q

What is the gas constant?

A

8.314 J mol^-1 K^-1

44
Q

What is the exponential factor?

A

e^-Ea/RT

- represents the proportion of molecules that exceed Ea and sufficient energy for a reaction to take place

45
Q

What is the pre-exponential factor?

A

A

  • takes into account the frequency of collisions with the correct orientation
  • essentially constant over small temperature ranges
  • gives the rate if there were no activation energy
46
Q

How can the Arrhenius equation be expressed in a logarithmic form?

A

ln k = -Ea/RT + ln A

47
Q

Why is the logarithmic form of the Arrhenius equation useful?

A
  • enables Ea and A to be determined graphically
  • A plot of ln k on y-axis and 1/T on the x-axis gives a straight line graph, y=mx+c
    ln k = y
    -Ea/R = m
    1/T = x
    ln A = c
48
Q

What does a plot of ln K and 1/T give us?

A

A straight downward line

  • gradient m of -Ea/R
  • intercept c of ln A on the y-axis
49
Q

Describe how the activation energy and frequency factor can be found graphically?

A

A graph of ln k against 1/T is plotted
The gradient of the straight line graph is equal to -Ea/R
The intercept of the y-axis is ln A

50
Q

In the Arrhenius equation, state the effect on k of an increase in A?

A

Increases k

51
Q

In the Arrhenius equation, state the effect on k of an increase in Ea?

A

Descreases k

52
Q

In the Arrhenius equation, state the effect on k of an increase in T?

A

Increases k

53
Q

How do you convert between k and ln k?

A

ln(value of k given)

54
Q

How do you convert between T and 1/T

A

1 divided by the value of T

55
Q

How do you calculate activation energy from the gradient of the straight line?

A

Work out gradient (y-axis divided by x-axis)
Since you now know the gas constant and gradient then:
Ea = R x m = 8.314 x gradient
(units are Jmol^-1, divide by 1000 for kJmol^-1)

A-Level OCR Chemistry (21 decks)

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