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A Level Chemistry > Topic 16: Kinetics II > Flashcards

Topic 16: Kinetics II Flashcards

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1
Q

What is the rate of reaction?

A
  • the change in concentration of a species divided by the time it takes for the change to occur
  • all reaction rates are positive
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2
Q

What are some methods of measuring rate of reaction?

A
  • measuring the volume of gas evolved:
    1. collection over water into measuring cylinder
    2. collection using gas syringe (more precise)
  • measuring change in mass:
  • most precise when the gas given off has relatively high density, such as CO2
  • colour change:
  • colourimeter results at different known concentrations can be plotted
  • change in pH:
  • used when a reaction produces or uses up H+ ions
  • titration
  • electrical conductivity
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3
Q

What is the rate equation?

A
  • Rate = k [A]m [B]n
  • units are usually mol dm-3 s-1
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4
Q

What are the two ways of determining the orders of reaction?

A
  • continuous method
  • initial-rate method
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5
Q

Briefly describe the continuous method and how orders of reaction can be found

A
  • one reaction mixture is made up
  • samples f the reaction mixture are withdrawn at regular time intervals
  • reaction is stopped by quenching
  • concentration of reactant is determined by titration
  • concentration-time graph drawn
  • find out half-life
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6
Q

How do you work out the reaction rate from a concentration-time graph and for a volume-time graph?

A
  • change in concentration of reactant / time
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7
Q

What is the order of reaction when [X] changes and the rate stays the same?

A
  • order of reaction is 0
  • if [X] double or triples etc, rate will stay the same
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8
Q

What is the order of reaction when the rate is proportional to [X]?

A
  • order of reaction with respect to X is 1
  • if [X] doubles, the rate doubles etc
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9
Q

What is the order of the reaction if the rate is proportional to [X]2 ?

A
  • order of reaction with respect to X is 2
  • if [X] doubles, rate will be 22 = 4 etc
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10
Q

What is the overall order of reaction?

A
  • the sum of all the orders of all the reactants
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11
Q

What is the half-life useful for?

A
  • identifying a first order reaction
  • the half-life is always constant
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12
Q

How do you work out the half-life of a reaction?

A
  • plot a concentration-time graph
  • then draw a line where the concentration is halved
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13
Q

What is the initial rate?

A
  • the time taken for a fixed amount of reactant to be used up or for a fixed amount of product to be formed
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14
Q

Briefly describe the initial-rate method

A
  • several reaction mixtures are made up and the initial rate is measured
  • you can then figure out the order of each reactant
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15
Q

How do you construct a rate-concentration graph?

A
  • using the data from a concentration-time graph, find the gradient at various points
  • now plot each point on a new graph with rate and concentration as the axes
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16
Q

What orders are these concentration-time graphs?

A
17
Q

Describe a clock reaction method

A
  • an example of initial rates method
    1. measure how the time is taken for a set amount of product to form changes as you vary concentration of one of the reactants
    2. there will be a sudden increase in the concentration of a certain product as a limiting reagent is used up
    3. there is usually an observable endpoint e.g. colour change
    4. the quicker the clock reaction finishes, the faster the initial rate of reaction
    5. the following assumptions are made:
  • the concentration of each reactant doesn’t change significantly over the time period
  • temperature stays constant
  • when the endpoint is seen, the reaction has not proceeded too far
18
Q

Describe the iodine clock reaction

A
  • a small amount of sodium thiosulfate solution and starch are added to an excess hydrogen peroxide and iodide ions in acid solution
  • sodium thiosulfate is added to the reaction mixture and reacts instantaneously with any iodine
  • all the iodine that forms in the first reaction is used up straight away in the second reaction, but also sodium thiosulfate is used u[, any more iodine that forms will stay in solution, the starch indicator will turn blue-black - end of reaction
  • varying the iodide or hydrogen peroxide concentration will give different times
  • you can then calculate initial rate with respect to iodide or hydrogen peroxide
19
Q

How do you carry out the iodine clock reaction?

A
20
Q

What is the rate-determining step?

A
  • reaction mechanisms can have one step or a series of steps
  • in a series of steps, each step can have a different rate
  • the overall rate is decided by the step with the slowest rate: the rate determining step
21
Q

How do you know which reactants 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 RDS
  • if a reactant doesn’t appear in the rate equation, then it isn’t involved in the rate-determining step
  • the RDS doesn’t have to be the first step in a mechanism
22
Q

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

A
  • it shows the number of molecules of that reactant which are involved in or before the rate-determining step
23
Q

How can knowing the rate-determining step help predict the mechanism?

A
24
Q

Show the reaction mechanism of the reaction between propanone and iodine

A
25
Q

What are the two different types of mechanism for nucleophilic substitution?

A
  • SN1: reactions only involve 1 molecule or ion in the rate-determining step
  • SN2: involves 2 molecules, 1 molecule and 1 ion or 2 ions in the rate-determining step
26
Q

What reaction mechanism do primary halogenoalkanes undergo?

A
  • only SN2
27
Q

What reaction mechanism do secondary halogenoalkanes undergo?

A
  • both SN1 and SN2
28
Q

What reaction mechanism do tertiary halogenoalkanes undergo?

A
  • only SN1
29
Q

Show an SN2 reaction

A
30
Q

Show an SN1 reaction

A
31
Q

What is a catalyst?

A
  • a catalyst increases the rate of a reaction by providing an alternative reaction pathway with a lower activation energy
  • the catalyst is chemically unchanged at the end of the reaction
32
Q

Describe homogeneous catalysts

A
  • they are in the same phase (solid, liquid, solution or gas) as the reactants
33
Q

Give examples of some homogeneous catalysts

A
  • H+ ion catalyses many reactions in aqueous solution:
  • iodination of propanone
  • Fe2+ or Fe3+ in the reaction between peroxydisulfate ions and iodide ions
34
Q

Describe Fe2+ or Fe3+ catalysts in the reaction between peroxydisulfate ions and iodide ions

A
  • With Fe2+:
  • S2O82-(aq) + 2Fe2+(aq) → 2SO42- + 2Fe3+
  • 2Fe3+ + 2I- → 2Fe2+ + I2
  • the reverse occurs with Fe3+
35
Q

Describe heterogeneous catalysts

A
  • they are in a different physical state from the reactants:
  • solid heterogeneous catalysts provide a surface for the reaction to take place on
  • it is usually a mesh or a fine powder to increase the surface area
  • or it might be spread over an inert support
  • they can be easily separated from the products and leftover reactants
  • they can be poisoned however
  • a poison is a substance that clings to the catalyst’s surface more strongly than the reactant does, preventing the catalyst from working
36
Q

What are the steps in heterogeneous catalysis?

A
  • adsorption: reactants are first adsorbed onto the surface of the catalyst
  • reaction: reactant molecules are held in positions that enable them to react together
  • desorption: product molecules leave the surface
37
Q

Describe the oxidation of ethanedioic acid by manganate(VII) ions

A
  • autocatalysis: the reaction is catalysed by one of its products
  • 5(COOH)2 + 2MnO4- + 6H+ → 10CO2 + 2Mn2+ + 8H2O
  • the reaction is very slow at room temp, but is catalysed by Mn2+
  • as the ions are not present initially, the reaction is very slow at first, but rate increases as ions increase
38
Q

Explain the Arrhenius Equation

A
  • as the activation energy gets bigger, k gets smaller
  • so a larger Ea means a slow rate
  • the equation also shows that as the temperature rises, j increases
  • you can use the equation to create an Arrhenius plot by plotting ln k against 1/T

A Level Chemistry (19 decks)

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