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A level- Chemistry > Rates of Reaction > Flashcards

Rates of Reaction Flashcards

1
Q

Define rates of reaction

A

Change in concentration of a reactant (or product) in a given time

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2
Q

What units are used to measure rate of reaction

A

Mol dm-3 s-1

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3
Q

What conditions lead to a higher rate of reaction

A
  1. Higher pressure when gas
  2. High concentration
  3. Surface area
  4. Higher temperature
  5. Catalyst
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4
Q

Describe the collision theory

A
  1. Particles must collide in the correct orientation, with significant energy to overcome the activation energy for a reaction to occur.
  2. In order to increase the rate of reaction, the number of successful collisions per unit time must be increased
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5
Q

How does increasing the concentration affect the rate of reaction

A
  1. An increase in concentration increases the number of particles in the same volume.
    2 The particles are closer together and collide more frequently.
  2. Therefore there will be more effective collisions in a given time, so a faster rate of reaction
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6
Q

How does increasing the pressure of a gas affect the rate of reaction

A
  1. When a gas is compressed into a smaller volume the pressure increases and the concentration of the gas molecules increases, as the same number of gas molecules occupy a smaller volume.
  2. The gas molecules are closer together and collide more frequently, leading to more effective collisions in a given time.
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7
Q

On a graph with time on the bottom and reactant concentration on the y axis draw the curve you would expect to see as a reaction progresses

A
  1. A initial steep slope down and then it slowly declines

2. For a product graph it would have an initial fast increase then it plateaus.

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8
Q

What are the two methods that can be used to determine the rate of reaction if a gas is produced

A
  1. Monitoring the volume of gas produced at regular time intervals using gas collection.- graph of time on x-axis and volume of gas produced on y-axis.
  2. Monitoring the loss of mass of reactants using a balance.
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9
Q

List the ways in which the progress of a chemical reaction can be measured

A
  1. Precipitation time
  2. Volume of gas collected
  3. Change in mass
  4. Change in pH
  5. Change in conductivity
  6. colour change
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10
Q

For this reaction suggest to ways to monitor the progress:

Na2S2O3 (aq) + 2HCl (aq) → 2NaCl (aq) + S(s) + SO2(g)+ H2O (l)

A
  1. Change in pH

2. Volume of gas produced

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11
Q

Equation for rate of reaction

A

Rate= change in concentration/change in time

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12
Q

How would the rate of loss of HCl be related to the rate of formation of CO2
CaCO3 + 2HCl → CaCl2 + CO2 + H2O

A

Rate of loss of HCl would be twice the rate of formation of CO2 and hence the rate of reaction

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13
Q

How is the rate of reaction related to the rate of production of a product and the number of moles of product?

A

Rate= rate of production/ moles of product

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14
Q

How is the rate of reaction related to to the rate of consumption of a reactant and the number of moles of a reactant?

A

Rate of reaction will be the inverse of the rate of consumption divided by the moles of that reactant.

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15
Q

Define the term catalyst

A

A substance that increases the rate of a chemical reaction without being used up in the process; a catalyst provides an alternative route for the reaction with lower activation energy.

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16
Q

Describe energy diagrams for endothermic and exothermic reactions

A

X-axis is reaction progress, y-axis is energy

  1. Exothermic- reactant line is above the products. Distance from reactants to peak is activation energy and distance from reactants to products is overall change in energy
  2. Endothermic- product line is above reactant line.
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17
Q

Describe the effect of a catalyst on an energy graph

A
  1. A peak below the original activation energy which shows the lower activation energy using a catalyst
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18
Q

Define the term homogeneous catalyst

A
  1. A catalyst that is in the dame physical state as the reactants-frequently aqueous or gaseous
  2. The catalyst reacts with reactants to form an intermediate.
  3. The intermediate then breaks down to give the product and regenerates the catalyst
  4. a) An example is ozone depletion - 2O3 (g) →(reversible) 3O2(g) Cl. (g) acts as a catalyst.
    b) Making esters with sulfuric acid as the catalyst- all liquids
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19
Q

Define the term heterogeneous catalyst

A
  1. A catalyst that is in a different physical state form the reactants; which is most frequently gaseous reactants and a solid catalyst.
  2. Reactant molecules are adsorbed (weakly bonded) onto the surface of a catalyst, where the reaction takes place.
  3. After the reaction, the product molecules leave the surface of the catalyst by desorption
  4. Examples- Haber process- Fe(s), Hydrogenation of alkenes- Ni(s), Contact process- V2O5 (s)
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20
Q

Describe the importance of catalysts for sustainability and economics

A
  1. Catalysts increase the rate of many industrial chemical reactions by lowering the activation energy.
  2. This then reduces the temperature needed for the process and energy requirements,
  3. If a chemical process requires less energy, then less electricity or fossil fuel is used.
  4. This will cut emissions of CO2 and can cut costs and increase profitability.
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21
Q

Describe a Boltzmann distribution graph

A
  1. X-axis is energy and y-axis is number of molecules with a given energy.
  2. The curve is a straight is line diagonally and then it slopes down but never reaches 0.
  3. Look at a picture and practice drawing
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22
Q

List the features of the Boltzmann distribution

A
  1. No molecules have zero energy- the curve starts at the origin.
  2. The area under the curve is equal to the total number of molecules.
  3. There is no maximum energy for a molecule - the curve does not meet the x-axis at high energy. The curve would need to reach infinite energy to meet the x-axis.
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23
Q

Describe the effect of temperature on the Boltzmann distribution

A
  1. At a higher temperature the curve peak is lower and to the right of the original curve.
  2. At a higher temperature, a greater proportion of molecules can overcome the activation energy.
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24
Q

Describe the effect of a higher temperature on the rate of reaction

A
  1. More molecules have an energy greater than or equal to the activation energy.
  2. Therefore there will be more successful collisions in a given time, increasing the rate of reaction.
  3. Collisions will also be more frequent as the molecules are moving faster, but the increased energy of the molecules is much more important than the increased frequency of collisions.
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25
Q

Describe how to show the effect of a catalyst using a Boltzmann distribution graph

A
  1. Another which is to the left of the original activation energy line.
  2. This shows a greater proportion of particles have an activation energy equal or greater than the activation energy.
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26
Q

Define activation energy

A

The energy required to start a reaction

27
Q

Why do most reactions require an activation energy

A

Bonds need to be broken

28
Q

Why would it not be a good idea to monitor the following reaction by loss in mass (gas produced)?
Mg(s) + 2HCl (aq) → MgCl2 (aq) + H2 (g)

A

Hydrogen has a small Mr so mass would change very little.

Volume of gas produced would be a more suitable method

29
Q

Define rate equation

A

The rate equation for a reaction A + B→C is given by rate=k[A]^m * [B]^n, where m is the order of reaction with respect to A and n is the order of reaction with respect to B.

30
Q

Define Order of Reaction

A

The power to which the concentration of reactant is raised in the rate equation.

31
Q

Define Rate constant

A

The constant that links the rate of reaction with the concentrations of the reactants.

32
Q

Describe Zero order

A
  1. When the concentration of a reactant has no effect on the rate, the reaction is zero order with respect to the reactant.
  2. It’s value is one in the rate equation
  3. Graph with [reactant] on x-axis and rate on y-axis would be a straight horizontal line- wouldn’t touch the y-axis.
  4. The intercept on the y-axis fives the rate constant?
33
Q

Describe first order

A
  1. A reaction is first order with respect to a reactant when the rate depends on its concentration raised to the power of 1.
  2. If the concentration doubles, the reaction rate increases by a factor of 2.
  3. If the concentration triples, the reaction rate increases by a factor of 3.
  4. Graph with [reactant] on x-axis and rate on y-axis would be a straight diagonal line starting from the origin.
  5. Rate constant can be determined from measuring the gradient
34
Q

Describe second order

A
  1. A reaction is second order with respect to a reactant when the rate depends on its concentration raised to the power of 2.
  2. If the concentration doubles, the rate of reaction is increased by a factor of 2^2=4
  3. If the concentration triples, the rate of reaction is increased by a factor of 3^2=9
  4. Graph with [reactant] on x-axis and rate on y-axis would be a curved line which has a gradual incline and then gets steeper as concentration increases.
  5. By plotting a second graph of rate against concentration squared, the result is a straight line through the origin. The gradient of this line= rate constant,k.
35
Q

What is overall order

A

overall order= sum of orders with respect to each reactant

36
Q

For reaction A+B+C→products, the orders are A:0, B:1, C:2.

Write overall order and the rate equation

A
  1. Overall order- 1+2=3
  2. Rate equation- rate=k [A]^0 * [B]^1 * [C]^2
  3. Simplified rate equation- rate= k[B] * [C]^2
37
Q

Describe how to find units for rate constant k

A

k=rate/ [A]^1 - units= s-1
k=rate/ [A] ^2 - units= dm3 mol-1 s-1
k=rate/ [A] ^0- units = mol dm-3 s-1

38
Q

Describe how to carry out an experiment to determine the order of reaction

A
  1. If there are 2 reactants, fist measure the initial rate at set concentrations.
  2. Then double one of the reactants concentration and see the effect on the rate
  3. Then double the other reactant concentration and see the effect on the rate.
39
Q

Methods for continuos monitoring of rate

A
  1. Monitoring by gas collection
  2. Monitoring by mass loss
  3. Colour change using a colorimeter- linked to concentration of the solution
40
Q

Describe the method for measuring rate of reaction using a calorimeter
CH3COCH3 (aq) + I2 (aq) → CH3COCH2I (aq) + H+ (aq) + I- (aq)

A
  1. Prepare standard solutions of known concentration of the coloured chemical iodine (orange-brown)
  2. Select filter with the complementary colour- green-blue
  3. Zero the calorimeter with a blank solution of distilled water
  4. Measure absorbance readings of the standard solution of iodine.
  5. Plot a calibration curve of absorbance against iodine concentration- gives a way of converting absorbance to concentration.
  6. Carry out the reaction, taking absorbance readings of the reacting mixture at measured time intervals.
  7. Use calibration curve to measure the concentration at each absorbance reading.
  8. Plot a second graph of concentration of the coloured reactant against time- from this you can determine the order of reaction with respect to the coloured chemical.
41
Q

Describe concentration-time graphs for 0, 1st and 2nd order reactions

A

0 order- produces a straight line with a negative gradient. The reaction rate does not change at all during the reaction. Gradient=k
1st order- Downward curve with a decreasing gradient over time. the time taken for the concentration to halve is constant- half-life
2nd order- Downward curve, steeper at start, but tailing off more slowly

42
Q

Define half-life

A

The time taken for concentration of a reactant to decrease by half.

43
Q

When measuring half-life why is it necessary to have other reactants present in excess

A

So that their concentration does not fall significantly during the reaction

44
Q

What happens to the half-life of a first order reaction as the reactant concentration falls

A

It stays the same- remains constant

45
Q

What happens to teh half-life of a zero order reaction

A

It gets smaller

46
Q

What happens to the half-life of a second order reaction

A

It gets bigger

47
Q

Write the 2 methods that can be used to determine the rate constant from a concentration-time graph for a first order reaction

A
  1. Using the rate=k[A] equation where rate= gradient

2. Using the half-life k=ln2/(t1/2) t1/2- is the time taken for concentration to halve.

48
Q

Why is it acceptable to assume this reaction is first order overall under these conditions:
C6H5N2Cl + H2O →C6H5OH + N2 + HCl
1. excess water
2. 1st order wrt C6H5N2Cl

A
  1. Under these conditions the concentration of water is effectively constant.
  2. It is not affecting the rate- only reactant affecting rate is C6H5N2Cl which is first order.
49
Q

How would half-life change if the temperature of the reaction were increased

A
  1. If the temperature was increased we would expect the reaction to speed up.
  2. The value of k would be larger and the half-life would get smaller
50
Q

Definition of initial rate of reaction

A

The change in concentration of a reactant or product per unit time at the start of the reaction (t=0).

51
Q

Why is the initial rate of a chemical reaction a good thing to measure in terms of having reproducible and reliable results

A
  1. Concentrations of all reactants are known

2. Temperature is known and controlled.

52
Q

Explain what will happen to the rate constant if the temperature is increased

A
  1. It will increase as rate of reaction will increase

2. A greater proportion of collisions will be successful

53
Q

Explain what will happen to the rate constant if a catalyst is added to the reaction mixture

A

It will increase, because catalyst provides an alternative pathway with lower activation energy so a greater proportion of collisions will be successful.

54
Q

List two ways the initial rate of reaction can be measured

A
  1. Measure the gradient at t=0 on a concentration-time graph

2. A clock reaction- more convenient way by taking a single measurement

55
Q

Describe how to work out the initial rate of reaction using a clock reaction

A
  1. Measure the time taken for a visual change to be observed- e.g. formation of iodine- time taken for formation of orange-brown colour.
  2. Provided there is no significant change in rate during this time, it can be assumed that the average rate of reaction over this time will be the same as the initial rate.
  3. The initial rate is then proportional to 1/t
  4. The clock reaction is repeated several times with different concentrations and values of 1/t are calculated.
  5. Then a graph of 1/t against concentration is plotted and the shape will match one of the order of reactions graph shapes.
  6. Further experiments are carried out in which the concentration of one of the other reactants is changed.- the order with respect to each reaction is determined.
56
Q

Define Rate-determining step

A

The slowest step in a reaction mechanism of a multi-step reaction

57
Q

How do we know if the reaction mechanism is likely to be correct?

A
  1. The rate equation only includes reacting species involved in the rate-determining step .
  2. The orders in the rate equation match the number of species involved in the rate-determining step.
  3. It needs to match the overall reaction equation
  4. Shouldn’t have more than 2 particles colliding- exam boards don’t really care
58
Q

Propose a reaction mechanism for the reaction
2ICl + H2 → 2HCl + I2
Rate= k [ICl] [H2]

A

Slow- ICl + H2 → HCl + HI

Fast- HI + ICl → HCl + I2

59
Q

The rate constant k for a given reaction can be determined at different temperatures, suggest how.

A
  1. A series of experiments are performed where the concentration of each reactant is known. The only variable is temperature
  2. Using an initial rate method, the effect of temperature on rate and therefore k can be determined
60
Q

Describe what the different symbols in the Arrhenius equation mean

A

k= rate constant
A= Frequency factor or pre-exponential factor- represents other factors that affect rate e.g. surface area and orientation of molecules
e= e on calculator
Ea= activation energy
R- the gas constant
T- Temperature in kelvin
e^-Ea/RT= represents the proportion of molecules that exceed Ea and have sufficient energy for a reaction to take place

61
Q

Under what circumstances is it assumed that A is constant

A

A is approximately constant when looking at a fairly narrow range of temperatures

62
Q

Why is the logarithmic form of the Arrhenius equation useful

A
  1. A graph of ln k against 1/T gives a straight line graph of y=mx + c.
  2. Gradient m = -Ea/R
  3. Intercept c on the y-axis = lnA
63
Q

Describe how the following changes effect k:
Increase in T
Increase in Ea
Reduction in A

A

Increase in T- increases k
Increase in Ea- decreases k
Reduction in A- decreases k

A level- Chemistry (31 decks)

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