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a level chemistry > 3.5 chemical kinetics > Flashcards

3.5 chemical kinetics Flashcards

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1
Q

what are the different ways you can follow and measure the rate of a chemical reaction?

A
  • measure volume of gas given off (using gas syringe)
  • measure change in mass (using scale)
  • measure concentration of coloured compounds in solution (using colorimeter)
  • measure pressure change in a reaction
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2
Q

how do you find the initial rate of a reaction?

A
  • plot what you are measuring (pressure, volume, mass) against time
  • take a tangent at t=0
  • gradient = change in y / change in x
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3
Q

rate definition

A
  • a change in concentration of a substance over time
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4
Q

sampling definition

A
  • when a small amount of reaction mixture is removed at timed intervals
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5
Q

quenching definition

A
  • sudden stopping or slowing down of a chemical reaction by adding the mixture to a beaker of ice or diluting
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6
Q

why is sampling and quenching useful?

A
  • in some reactions, further analysis is needed mid-way through a reaction
  • some reactions take too long, so this allows us to study reactions as time passes without having to wait for the reaction to complete fully
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7
Q

what are some cons with sampling and quenching?

A
  • sampling is only appropriate when the reaction mixture is homogeneous, otherwise you may be removing a sample that is not fully representative of the whole reaction mixture
  • time consuming - as each sample is analysed individually and the process cannot be automated
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8
Q

can stop a reaction which uses a heterogenous catalyst by:

A
  • removing the catalyst
  • destroy the catalyst (e.g add OH- to H+ catalyst)
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9
Q

what is a reaction order used to describe?

A
  • the relationship between the concentration of a substance and its effect on the rate of a reaction
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10
Q

zero order:

A
  • the concentration of A has no effect on the rate of reaction
  • rate ∝ [A]^0
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11
Q

first order:

A
  • the rate of reaction increases at the same rate as [A] increases
  • r.o.r is directly proportional to the concentration
  • rate ∝ [A]^1
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12
Q

second order:

A
  • the rate of reaction is proportional to the concentration of A^2
  • rate ∝ [A]^2
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13
Q

what does a zero order rate-concentration graph look like?

A
  • rate stays constant even as concentration increases
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14
Q

what does a first order rate-concentration graph look like?

A
  • the rate increases as the concentration increases
  • directly proportional
  • linear
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15
Q

what does a second order rate-concentration graph look like?

A
  • as the concentration increases, so does the rate at an increasing rate
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16
Q

what is the rate equation?

A

rate = k[A]^m[B]^n

k= rate constant
[A] = conc of A
m = order of reaction with respect to A

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17
Q

what are the units of rate?

A

moldm^-3s^-1

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18
Q

e.g rate = k [OH-]
units:
moldm^-3s^-1 = s^-1 x moldm^-3

A
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19
Q

the orders have nothing to do with the stoichiometric coefficients in the balanced equation. they are worked out experimentally

A
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20
Q

how do you work out the total order for a reaction?

A
  • by adding all the individual orders together (e.g m+n)
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21
Q

rate constant definition

A
  • this is the constant of proportionality in the rate equation (linking the rate of reaction and the concentrations of the reactants raised to the power of their orders in the rate equation)
  • these are constant for any concentration of reactants
  • but they change if temperature changes
  • k
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22
Q

order definition

A

the power to which the concentration is raised in the rate equation

23
Q

the value of k is independent of:

A
  • concentration
  • time

(and is constant at a fixed temperature)

24
Q

if you increase the temperature, what happens to k?

A

it increases

25
Q

what happens to k if the concentration or pressure is changed at constant temperature?

A

it will not change

26
Q

if a catalyst is added, what happens to k?

A
  • the value of k will increase as the rate increases (as there is no concentration change in reactants)
    (rate = k[A]^m[B]^n)
27
Q

calcualte the units of the rate constant, k, for: rate = k[PCl][Cl2]^2

A

moldm^-3s^-1 = k [moldm^-3][moldm^-3]^2
k = moldm^-3s^-1 / mol3dm^-9
= mol^-2dm^6s^-1

28
Q

what’s the significance of the rate equation?

A
  • it defines the rate determining step (RDS) (slowest step in a mechanism/reaction)
29
Q

rate determining step definition

A
  • the slowest step in the reaction or set of reactions
30
Q

e.g A + 2B + C —> D + E
rate equation: r=k[A]^1[B]^1[C]^0

step 1: A+B —> X+D
step 2: X+C —> Y
step 3: Y+B —> E

which is the rate determining step?

A

1 because both reactants are in the rate equation

31
Q

e.g A—>E
1st : A—>B fast
2nd: B—>C fast
3rd: C—>D slowest = RDS
4th: D—>E fast

rate equation describes C—>D

A
32
Q

e.g
1st : A+B —> X+D fast
2nd: X +C —> Y slow
3rd: Y + B —> E fast

  • generate the rate equation
A

rate = k[X]^1[C]^1

(order number + balancing number in rate determining step)

33
Q

suggesting mechanistic steps using the rate equation:

A
  • a mechanism is a series of steps through which the reaction progresses, often forming intermediate compounds
  • If all the steps are added together they will add up to the overall equation for the reaction
  • the idea behind suggesting a mechanism is to draw up potential reaction routes a reaction might progress by
  • the rate equation and the overall reaction equation helps you massively when doing this
  • you might have to make up hypothetical fragments of molecules you may not have seen before in your course
  • there are often multiple ways of getting the answer correct as there is more than one way a reaction can proceed by
34
Q

e.g overall reaction NO2(g) + CO(g) —> NO(g) + CO2(g)
rate equation: rate= k[NO2]^2

suggest a two-step mechanism for this reaction

A
  1. NO2 + NO2 (bc in rate equation) —> N2O4
  2. CO + N2O4 —> NO + CO2 + NO2
    (one NO2 and N2O4 cancel out leaving overall equation)

OR

  1. NO2 + NO2 —> N2 + 2O2
  2. CO + N2 + 2O2 —> NO + CO2 + NO2
35
Q

what is arrhenius equation?

A

k = A. e^(-Ea / RT)

k = rate constant
A = frequency factor/pre-exponential factor/a-factor (rough measure of how many collisions per second)
e = exponential
Ea = activation energy (JOULES PER MOLE)
R = 8.31
T = temp in K

(it relates how the value of the rate constant changes with regard to temperature and how this also relates to the activation energy for a chemical reaction)

36
Q

-!the units of A in the arrhenius equation are the same as the units for the rate constant for a reaction

A
37
Q
  • the value of A is constant over a small temperature range
  • however at larger differences in temperature, this number greatly increases
A
38
Q

what are the units for the activation energy in the arrhenius equation?

A
  • joules per mole
  • Jmol^-1
39
Q

calculate the activation energy for a reaction which was known to have a rate constant of 0.349s^-1 at a temperature of 450°C. the frequency factor for this reaction is 7.11×10^12s^-1

A

k = A x e^(-Ea / RT)
k/A = e^(-Ea/RT)
ln(k/A) = -Ea / RT
ln(k/A) x RT = -Ea
= 184 kJmol^-1
= 184,000 Jmol^-1

40
Q

how can you take the arrhenius equation and turn it into the straight line graph formula of y= mx + c

A

k = A. e^(-Ea/RT)
(take logs of both sides)
lnK = lnA - Ea/RT
lnK = lnA - Ea/R 1/T
y = c + m x

(plot with y axis= lnK, x axis = 1/T)

41
Q
  • questions in the past have often asked about reactants in excess
  • in reactions where there are several reactants, if the concentration of one of the reactants is kept in a large excess then that reactant will appear not to affect rate and will be pseudo-zero order
  • this is because its concentration stays virtually constant and therefore does not affect rate
A
42
Q

what is the method of calculating the unit of a rate constant?

A
  • rearrange the rate equation to make k the subject
  • substitute units into the equation
  • cancel the common units to find the units for k
43
Q

activation energy definition

A

the minimum amount of energy required for a reaction to occur

44
Q

what are some different experimental techniques that allow you to obtain rate data?

A
  • titration
  • measuring the volume of gas releases from a reaction over time
  • measuring the change in mass of a reactant over time
  • colorimetry
  • measuring the time taken for a colour change
45
Q

what equation can be used to calculate activation energy?

A
  • the Arrhenius equation
46
Q

2N2O5 (g) —> 4NO2 (g) + O2 (g)
(graph: as time increases, concentration of N2O5 decreases)
explain why the rate of reaction is lower at t=60mins than it was at t=40mins [1]

A

as reaction proceeds less collisions (per unit time) occur

47
Q

state giving a reason, how the value of k will alter, if at all, if the temperature is increased [2]

A
  • increases
  • if temperature is increased rate increases
48
Q

hydrogen peroxide reacts with acidified potassium iodide according to the equation:
2H+ + 2I- + H2O2 —> I2 + 2H2O
this reaction was studied using an iodine clock reaction. describe the principles of how the rate of a clock reaction is determined. experimental details are not required [2]

A
  • measure time taken for a sudden colour change
  • rate = 1/time
49
Q

some experiments are undertaken at e.g pH1 and some at pH2. give the difference in the concentrations of H+ ions in these two solutions [1]

A

pH1 has a concentration of H+ ten times higher than pH2

50
Q

rate equation: rate = k[H2O2][I-]
the reaction is repeated at a higher temperature. state how the increase in temperature affect the rate equation and rate constant [1]

A
  • rate equation unchanged (bc rate equation shows rate determining step so still slowest)
  • increasing temperature increases the value of the rate constant
51
Q

rate = k[NaOH][A]
NaOH conc = 0.04 A conc = 0.03

in a further experiment at a different temperature, the initial rate of reaction was found to be 9.0x10^-3 moldm^-3s^-1 when the initial concentration of A was 0.020moldm^-3 and the initial concentration of NaOH was 2moldm^-3

under these new conditions with the much higher concentration of sodium hydroxide, the reaction is first order with respect to A and appears to be zero order with respect to sodium hydroxide

suggest why the order of reaction with respect to sodium hydroxide appears to be zero under these new conditions [2]

A
  • (large) excess of OH- / [OH-] is large/high (1)
  • [OH-] is (effectively) constant
    OR
  • [A] is the limiting factor
52
Q

explain what is meant by sampling and quenching [1]

A
  • removing samples from the reaction mixture at set intervals and then stopping the reaction/adding large volume of water
53
Q

explain what is meant by the term ‘quenching’ and why it needs to be carried out [2]

A
  • quenching is the sudden stopping/significant slowing of a chemical reaction in a sample
  • ensure sample composition doesn’t change between taking sample and analysis OR during analysis

a level chemistry (17 decks)

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