Module 3: Physical Chemistry V2 Flashcards
1
Q
Define what is meant by “enthalpy change”
A
The heat energy transferred in a reaction at constant pressure. ✓
2
Q
What are the units for enthalpy change
A
3
Q
Define what is meant by “standard states”
A
Physical states they are in under standard conditions ✓
4
Q
State what standard conditions are
A
100kPa and 298K ✓
5
Q
What is meant by “⦵” symbol
A
100kPa and 298K ✓
6
Q
Define what is meant by “standard enthalpy change of reaction” and give an example for this between the reaction of carbon monoxide and nitrogen dioxide to form carbon dioxide and nitrogen monoxide.
A
7
Q
Define what is meant by “standard enthalpy change of neutralisation” and give an example for this i.e. sulfuric acid and sodium hydroxide. Also predict the enthalpy change as being positive or negative.
A
8
Q
Define what is meant by “standard enthalpy change of combustion” and give an example for this i.e. ethanol. Also predict the enthalpy change as being positive or negative.
A
9
Q
Define what is meant by “standard enthalpy change of formation” and give an example for this i.e. ethanol. Also predict the enthalpy change as being positive or negative.
A
10
Q
Define what is meant by the term “activation energy”
A
The minimum amount of energy needed for a reaction to start. ✓
11
Q
Define what is meany by an “exothermic reaction”
A
Reaction that releases energy with a negative ∆H ✓
12
Q
Define what is meany by an “endothermic reaction”
A
Reaction that absorbs energy with a positive ∆H ✓
13
Q
Draw an enthalpy profile diagram for the following.
A
14
Q
Draw an enthalpy profile diagram for the following.
A
15
Q
A student carried out an experiment to measure the enthalpy change of combustion of methanol.The energy from the combustion of methanol was used to heat a beaker containing water.
The student’s calculated enthalpy change of combustion was more exothermic than the value in data books. explain which error could have caused this difference?
A: Some methanol had evaporated from the wick before the final weighing.
B: In the calculation, the student used the molar mass of ethanol instead of methanol.
C: There was incomplete combustion.
D: The water boiled for 5 minutes before the final temperature was taken.
A
B: In the calculation, the student used the molar mass of ethanol instead of methanol.
n becomes smaller, so when delta H = q/n, q is divided by a smaller number, delta H will be larger overall.
16
Q
Explain why standard enthalpy changes of combustion determined experimentally are less exothermic than the calculated theoretical values
A
Heat released to the surroundings ✓
Incomplete combustion ✓
Non-standard conditions ✓
17
Q
Explain why standard enthalpy changes of neutralisation determined experimentally are less exothermic than the calculated theoretical values
A
Heat released to the surroundings ✓
Incomplete reaction ✓
Non-standard conditions ✓
18
Q
Explain why it is not always possible to measure the enthalpy change directly from experiment
A
The activation energy of the reaction is too high ✓
There is a slow reaction rate ✓
More than one reaction may be taking place (and so more than one product is formed) ✓
19
Q
Suggest how enthalpy change can be determined indirectly
A
Average bond enthalpy data ✓
or the standard enthalpy change of combustion or formation data instead. ✓
20
Q
Define the term ‘average bond enthalpy’
A
Average enthalpy change when one mole ✓
of gaseous covalent bonds is broken ✓
21
Q
Explain why bond enthalpies have positive values?
A
Bond breaking is endothermic ✓
22
Q
Explain why bond enthalpies have negative values?
A
Bond forming is exothermic ✓
23
Q
Explain, in terms of bond breaking and bond forming, why a reaction can be exothermic.
A
More energy is released by forming bonds ✓
than energy required when breaking bonds ✓
24
Q
Explain, in terms of bond breaking and bond forming, why a reaction can be endothermic.
A
More energy is required for breaking bonds ✓
than is released by forming bonds ✓
25
Explain why the standard enthalpy changes will be different from that calculated using average bond enthalpies
Actual bond enthalpies may be different from average values ✓
Conditions are not standard ✓
26
27
28
A student is conducting an enthalpy change experiment in the lab. He uses a 0.327 g mass of Zn in the reaction between zinc and copper sulfate. He obtains a percentage error value of 5%.
Suggest a modification, using the same apparatus, that will reduce the percentage uncertainty.
Use a larger mass of Zn ✓
As this gives a larger ΔT reducing the percentage error ✓
29
A student is conducting an enthalpy change experiment in the lab. He uses a 0.327 g mass of Zn in the reaction between zinc and copper sulfate. He obtains a percentage error value of 5%.
He obtains an enthalpy change value of 400 kilo joules per mole
Another student repeats the experiment using the same mass of Zn but double the volume of water. Explain how will this affect ΔT and ΔH?
ΔT would be halved and ΔH would be the same ✓
Same energy released over double the volume ✓
30
A student is conducting an enthalpy change experiment in the lab. He uses a 0.327 g mass of Zn in the reaction between zinc and copper sulfate. He obtains a percentage error value of 5%.
Another student repeats the experiment using double the mass of Zn but the same volume of water. Explain how will this affect ΔT and ΔH?
ΔT would be double ΔH would be the same ✓
as double energy released in the same volume ✓
31
32
Explain why enthalpy changes calculating using average bond enthalpies will be different from standard enthalpy changes
Actual bond enthalpies may be different from average values ✓
Conditions are not standard ✓
33
## Footnote
you're dealing with enthalpy change of "reaction". There are 2 moles of silver nitrate reacting.
34
State which of the following are True/False
a) The activation energy of the forward reaction is 120 kJ mol−1.
b) The activation energy of the reverse reaction is 270 kJ mol−1
c) The reverse reaction is exothermic.
d) The enthalpy change of the forward reaction is −30 kJ mol−1.
e) The reaction is exothermic.
f) The arrow for activation energy and enthalpy point up for the reverse reaction.
a) F
b) T
c) F
d) F
e) T
f) T
35
Hydrogen and chlorine react as shown.
State which of the following are True/False
a) The enthalpy change for the reverse equation is +184.6 kJ mol−1.
b) Less energy is released on bond making than is taken in during bond breaking.
c) The enthalpy change of formation of HCl (g) is −184.6 kJ mol−1.
d) The temperature decreases during the reaction.
a) T
b) F
c) F
d) F
36
37
What is the value of the enthalpy change of formation of ammonia?
## Footnote
Think back to what formation means
38
What is the value of the activation energy for the uncatalysed reverse reaction?
39
40
41
Calculate the enthalpy change of formation using the data provided.
42
Using a Maxwell-Boltzmann Distribution, explain how increasing temperature increases the rate of a reaction.
At a higher temperature, more molecules have energy above activation energy. ✓
And so, there are more frequent and successful collisions. ✓
## Footnote
axis, including activation energy ✓ two curves, labelled (not T1/T2) ✓
43
Describe and explain the effective of increase pressure on the rate of reaction.
Increasing pressure/concentration increases the rate of reaction. ✓
There are more particles per unit volume. ✓
There are more frequent collisions. ✓
44
Explain what is meant by collision theory, in terms of how molecules reaction
Particles in liquids and gases are moving and collide with each other – They do not always react. ✓
They collide in the correct orientation. ✓
They collide with at least a certain minimum amount of kinetic energy, known as activation energy ✓
45
Define what is meant by a "catalyst"
Catalysts increases the rate of reaction ✓
by providing an alternate reaction pathway with a lower activation energy. ✓
The catalyst is chemically unchanged at the end of a reaction. ✓
46
Describe what is meant by a homogenous catalyst. Give an example and the reaction of catalyses.
Same phase as reactants i.e. same physical state ✓
Example: sulfuric acid in some organic reactions i.e. oxidation of ethanol to ethanoic acid ✓
47
Describe what is meant by a heterogenous catalyst. Give an example and the reaction of catalyses.
Different phase from reactants i.e. different physical state ✓
Example: Hydrogenation of alkenes with a Ni catalyst ✓
48
Using a Maxwell-Boltzmann Distribution, explain how using a catalyst increases the rate of a reaction.
Catalyst provides lower activation energy. ✓
More molecules have energy above activation energy. ✓
And so, there are more successful collisions. ✓
## Footnote
xis, including activation energy ✓ axis, labelled with catalysed activation energy ✓
49
Describe four main reasons why catalysts are used in industry
Reactions can be carried out at lower temperatures and pressures. ✓
Leads to less fuels burnt and lower CO2 emissions. ✓
Different reactions can be used with greater atom economy and less waste. ✓
Can reduce the use of toxic substances. ✓
50
Draw an enthalpy profile diagram for the following reaction, with a KI catalyst
## Footnote
axis, with reactants higher than products (labelled too) ✓ Catalysed activation energy arrow ✓ uncatalysed activation energy arrow and key ✓ enthalpy change arrow going down ✓
51
Describe and explain the effective of a decrease in volume on the rate of reaction.
Decreasing volume increases pressure/concentration and so increases the rate of reaction. ✓
There are more particles per unit volume. ✓
There are more frequent collisions. ✓
52
Define what is meant by "rate of reaction"
change in the number of reactants or products per unit time ✓
53
Concentration of reactants is greatest at the start of a reaction ✓
There are greater number of collisions per second and so more successful collisions ✓
The concentration of a reactant decreases as the reaction proceeds. ✓
The rate of the reaction decreases as there are less collisions with sufficient Ea taking place per second between the reactant particles. ✓
54
Using the two coordinates on the tangent, calculate the gradient of the line at t = 20 s.
Give units for your answer.
55
## Footnote
I could not be bothered to draw the Hess' Cycle.
56
State three features of an equilibrium
Reaction is taking place in a closed system. ✓
Rate of the forward reaction is equal to the rate of the reverse reaction. ✓
Concentrations of reactants and products remain constant and do not change. ✓
57
Explain how a reaction reaches equilibrium
As reactants get used up, forwards reaction slows down. ✓
As more product is formed, reverse reaction speeds up. ✓
At equilibrium, both forward and backward reaction are going at the same rate. ✓
58
State Le Chatelier's principle
The position of a dynamic equilibrium shifts to minimise the effect of any change ✓
59
Using Le Chatelier's Principle, What effect would increasing the concentration of ethanol have on the following equilibrium.
60
State and explain the effect of increasing the concentration of water in the following reaction
61
Using Le Chatelier's Principle, What effect would decreasing the concentration of ethanol have on the following equilibrium
62
State and explain the effect of decreasing the volume of the container for the following reaction
63
State and explain the effect of decreasing the pressure of the container for the following reaction
64
Using Le Chatelier's Principle, What effect would increasing the temperature have on the following equilibrium
65
Using Le Chatelier's Principle, What effect would decreasing the temperature have on the following equilibrium
66
State and explain the effect of adding a catalyst, on the position of the equilibrium.
A catalyst has no effect on the position of the equilibrium. ✓
A catalyst increases the rate of the forward and reverse reactions equally. ✓
Therefore, only increases the rate at which equilibrium is established. ✓
67
State and explain the effect of increasing pressure on the equilibrium
No change in equilibrium position ✓
Number of moles of gaseous reactants and products is the same ✓
68
A catalyst is added to a system in equilibrium.
What is the effect on the rates of the forward and reverse reactions?
A There is no effect on the rate in either direction.
B Both rates increase by the same factor.
C The rate in the forward direction increases by a greater factor than the reverse direction.
D The rate in the reverse direction increases by a greater factor than the forward direction.
B Both rates increase by the same factor ✓
69
C ✓
70
B ✓
71
Ethene and steam can be converted into ethanol. The equilibrium is shown below.
Using le Chatelier’s principle, predict and explain the conditions that would give the maximum equilibrium yield of ethanol from ethene and steam.
Low temperature... ✓
... as forward reaction is exothermic. ✓
High pressure... ✓
... as there are fewer moles of gas on the right-hand side. ✓
72
Ethene and steam can be converted into ethanol. The equilibrium is shown below.
The actual conditions used are 60 atmospheres pressure at 3000C in the presence of a catalyst. Why are these conditions used rather than low temperature and high pressures?
Low temperature leads to a slow rate of reaction ✓
High pressure provides a safety risk ✓
High pressure is expensive to generated and uses a lot of energy✓
73
Outline how a catalyst increases the rate of a chemical reaction.
Allows reaction to proceed via a route with lower activation energy. ✓
more molecules have energy above the activation energy. ✓
74
Which statement explains why the rate of a reaction increases when the temperature is increased?
A. The activation energy for the reaction decreases.
B. The activation energy for the reaction increases.
C. The proportion of molecules exceeding the activation energy decreases.
D. The proportion of molecules exceeding the activation energy increases.
D. The proportion of molecules exceeding the activation energy increases. ✓
75
Vanadium oxide is a solid while the reactants are gases
Catalyst is in a different state to reactants
76
Describe and explain the effect of decreasing the pressure on the rate of a reaction.
Decreasing pressure decreases the rate of reaction. ✓
There are less particles per unit volume. ✓
There are less frequent collisions. ✓
77
Explain why the use of a catalyst can reduce the demand for energy.
Allows reactions to take place at lower temperatures. ✓
78
A student investigates the reaction between strontium carbonate and dilute nitric acid.
The rate of reaction is determined from the loss in mass over a period.
Explain why there is a loss in mass during the reaction.
Carbon dioxide gas is produced. ✓
79
A student investigates the reaction between strontium carbonate and dilute nitric acid.
SrCO3 + 2HNO3 → Sr(NO3)2 + CO2 + H2O
The rate of reaction is determined from the loss in mass over a period.
The student plots a graph of total mass (reagents + container) against time.
Describe and explain the change in the rate of the reaction during the first 200 seconds of the experiment.
Rate of reaction decreases ✓
Concentration decreases and reactants are used up. ✓
Less frequent collision. ✓
80
A student investigates the reaction between strontium carbonate and dilute nitric acid.
SrCO3 + 2HNO3 → Sr(NO3)2 + CO2 + H2O
The rate of reaction is determined from the loss in mass over a period.
The student plots a graph of total mass (reagents + container) against time.
Using the graph, calculate the rate of reaction, in g s−1, at 200 seconds.
Show your working on the graph.
Draw a tangent at t = 200s ✓
Calculate the gradient ✓
Answer: 6.9 x 10^-4 g s-1 ✓
## Footnote
Might be hard to draw a tangent, but use your rular to get points on the tangent.
81
A student investigates the reaction between strontium carbonate and dilute nitric acid.
SrCO3 + 2HNO3 → Sr(NO3)2 + CO2 + H2O
The rate of reaction is determined from the loss in mass over a period.
The student plots a graph of total mass (reagents + container) against time.
Outline a method that could be used to obtain the results that are plotted on the graph.
Your answer should include the apparatus required and the procedure for the experiment.
Place a flask on a mass balance. ✓
Record mass at time intervals of 10-15 seconds ✓
using a stopwatch. ✓
82
A student was asked to carry out an experiment to determine the initial rate of reaction of zinc and hydrochloric acid.
The student plans to collect a total of about 72 cm3 of hydrogen at RTP and to use an excess of zinc.
The student selects the following apparatus:
* the apparatus shown in the diagram
* 100 cm3measuring cylinder
* stop clock
* 2 decimal place balance
Outline how the student could carry out the experiment and explain how the results could be processed graphically.
Show all working in your calculations.
Method:
Measure mass of excess zinc using a 2 decimal place balance. ✓
Measure the volume of hydrochloric acid using a measuring cyclinder. ✓
Mix the zinx and acid in the flask. ✓
Measure the gas volumes at time intervals of 10-15 seconds using a stopwatch. ✓
Processing results:
Plot a graph of volume against time. ✓
Draw a tangent at t = 0. ✓
Calculate the gradient of the tangent, which is equal to the initial rate ✓
Gradient = volume/ time. ✓
83
A student is trying to find the initial rate of reaction using calcium carbonate chips and hydrochloric acid.
They set up a flask attached to a gas syringe and measure the change in volume of gas produced over time.
Explain how the results can be processed graphically.
Plot a graph of volume against time.
Draw a tangent at t = 0.
Calculate the gradient of the tangent, which is equal to the initial rate
Gradient = volume/ time.
84
Suggest two reasons why using a lower temperature is beneficial to the environment.
Less fossil fuels used. ✓
Reduction is CO2 emissions. ✓
85
The student investigates the reaction between zinc and dilute sulfuric acid.
Copper(II) sulfate is a catalyst for this reaction.
* The student adds a piece of zinc to each of two test tubes.
* The student adds a few drops of aqueous copper(II) sulfate to one of the test tubes, forming a pale blue solution.
* The student adds an excess of dilute sulfuric acid to each test tube.
Describe two differences the student would observe between the test tubes.
Faster fizzing. ✓
Zn dissolves more quickly. ✓
86
A student is trying to find the initial rate of reaction using calcium carbonate chips and hydrochloric acid.
They set up a flask attached to a gas syringe and measure the change in volume of gas produced over time.
He plots the results on a graph and draws a line of best fit.
He then repeats the experiment but this time crushes the calcium carboante into a fine powder.
Sketch a second line of best fit to show what he would expect in terms of his results.
87
A student is trying to find the initial rate of reaction using calcium carbonate chips and hydrochloric acid.
They set up a flask attached to a gas syringe and measure the change in volume of gas produced over time.
He plots the results on a graph and draws a line of best fit.
He then repeats the experiment but this time uses a more dilute acid solution.
Sketch a second line of best fit to show what he would expect in terms of his results.
88
Which statement(s) explain(s) why reaction rates increase as temperature increases?
1 The activation energy is less.
2 Collisions between molecules are more frequent.
3 A greater proportion of molecules have energy greater than the activation energy.
2 Collisions between molecules are more frequent. ✓
3 A greater proportion of molecules have energy greater than the activation energy. ✓
89
The diagram represents a Boltzmann distribution curve of molecules at a given temperature.
Which statement for this Boltzmann distribution curve is correct at a higher temperature?
A The peak increases in height and moves to the left.
B The peak increases in height and moves to the right.
C The peak decreases in height and moves to the left.
D The peak decreases in height and moves to the right.
The peak decreases in height and moves to the right.
90
At a higher temperature, there is a fast rate of reaction as more molecules have energy above activation energy. ✓
And so, there are more frequent collisions. ✓
Concentration of [Co(H2O)6]2+ is lower. ✓
Forward reaction is exothermic. ✓
91
Carbon monoxide reacts with steam in the following reaction equation:
Which change will shift the position of equilibrium to the right hand side of the equation?
A. decrease in pressure
B. increase in pressure
C. decrease in temperature
D. increase in temperature
C. decrease in temperature ✓
92
Which statement is not correct for a system in dynamic equilibrium?
A. The concentrations of products and reactants are the same.
B. The equilibrium can be achieved from both sides.
C. The rate of the forward reaction is equal to the rate of the reverse reaction.
D. The system is closed.
A. The concentrations of products and reactants are the same. ✓
93
Ammonia is used in the manufacture of nitric acid. The first stage of this process is a dynamic equilibrium.
When the temperature is increased, Kc for this reaction decreases.
State the effect, if any, on the equilibrium yield of NO in this reaction.
Explain your answer.
Equilibrium yield of ammonia increases. ✓
Increasing temperature shifts the equilibrium in the endothermic direction ✓
to oppose the change in temperature. ✓
Equilibrium shifts left. ✓
94
The reversible reaction below is allowed to reach equilibrium.
Which change in conditions would be expected to shift the equilibrium position towards the products?
A decrease the pressure
B decrease the temperature
C increase the pressure
D increase the temperature
B. Decrease in temperature. ✓
95
B: 0.0540. ✓
96
Cl- reacts with Ag+ and AgCl is formed. ✓
97
[CoCl42-] decreases and [Co(H2O)6]2+ increase. ✓
Cl- increase is 4x change in [CoCl42-] / [Co(H2O)6]2+. ✓
Equilibrium shifts to the right. ✓
98
The student’s calculated enthalpy change was less exothermic than the value in data books.
Which of the following errors could have contributed to this result?
Error 1: After the final temperature was recorded, the student removed the burner from under the beaker. The flame burnt for a further 5 minutes before weighing the spirit burner.
Error 2: The student recorded the final temperature 5 minutes after removing the burner.
Error 3: The student spilt some water on the bench when pouring the water from the measuring cylinder into the beaker.
Error 1: After the final temperature was recorded, the student removed the burner from under the beaker. The flame burnt for a further 5 minutes before weighing the spirit burner. ✓
Error 2: The student recorded the final temperature 5 minutes after removing the burner. ✓
99
100
101
102
State the density of water
1g cm-3 ✓
103
104
