R2.3 How far? The extent of chemical change
Practice exam-style IB Chemistry questions for How far? The extent of chemical change, aligned with the syllabus and grouped by topic.
Question 1
A sealed flask contains liquid bromine and bromine vapour at equilibrium.
A.
The rates of evaporation and condensation are equal.
B.
Evaporation has stopped but condensation continues.
C.
The amounts of liquid and vapour must be equal.
D.
No particles move between the liquid and vapour phases.
Question 2
For the homogeneous equilibrium
2NO
A.
K = [NO₂]² / ([NO]²[O₂])
B.
K = [NO₂] / ([NO][O₂])
C.
K = 2[NO₂] / (2[NO][O₂])
D.
K = [NO]²[O₂] / [NO₂]²
Question 3
For a reaction as written, K = 4.0 × 10⁻⁶ at a fixed temperature.
A.
Reactants are strongly favoured.
B.
The reaction is not reversible.
C.
Reactants and products are present in equal concentrations.
D.
Products are strongly favoured.
Question 4
At 500 K, A
A.
156
B.
12.5
C.
−12.5
D.
0.0800
Question 5
What is the effect of adding a catalyst to a reversible reaction mixture at equilibrium at constant temperature?
A.
It decreases the backward rate only.
B.
It increases K by favouring product formation.
C.
It increases the rates of both forward and backward reactions equally.
D.
It shifts the equilibrium toward the exothermic direction.
Question 6
A small amount of liquid ethanol is placed in a sealed container and allowed to reach equilibrium with its vapour.
State one feature of the container required for equilibrium to be established.
[1]
Describe two characteristics of the system once dynamic equilibrium has been reached.
[1]
Question 7
Deduce the equilibrium constant expression for the homogeneous equilibrium:
CH₄
- H₂O
[1]
⇌ CO
[1]
- 3H₂(g)
Question 8
A catalyst is added to a reversible reaction mixture before equilibrium is reached.
State the effect on the time taken to reach equilibrium.
[1]
Explain why the equilibrium composition is unchanged.
[1]
Question 9
For the equilibrium
N₂O₄
A.
The equilibrium shifts right and K decreases.
B.
The equilibrium shifts left and K is unchanged.
C.
The equilibrium shifts right and K increases.
D.
The equilibrium shifts left and K increases.
Question 10
For the equilibrium
CO
A.
There is no shift and K decreases.
B.
The equilibrium shifts to the right and K increases.
C.
The equilibrium shifts to the left and K is unchanged.
D.
The equilibrium shifts to the right and K is unchanged.
Question 11
For the heterogeneous equilibrium
CaCO₃
A.
There is no shift because solids are present.
B.
The equilibrium shifts left.
C.
K increases because the pressure increases.
D.
The equilibrium shifts right.
Question 12
For A
A.
No prediction is possible without pressure.
B.
The forward reaction is favoured.
C.
The mixture is at equilibrium.
D.
The reverse reaction is favoured.
Question 13
For 2X
A.
[X] and [Y] both increase.
B.
[Y] decreases and Q decreases.
C.
[Y] increases and Q increases.
D.
[X] decreases and Q remains constant.
Question 14
At 298 K, a reaction has K = 1.0 × 10³.
A.
Zero, because the reaction is at equilibrium.
B.
Negative, because ln K is negative.
C.
Positive, because ln K is positive.
D.
Negative, because ln K is positive.
Question 15
For A
A.
0.0750 mol dm⁻³
B.
0.450 mol dm⁻³
C.
0.300 mol dm⁻³
D.
0.150 mol dm⁻³
Question 16
At the same temperature, the reaction P
⇌ Q
[1]
has K = 0.025.
[1]
State which side is favoured at equilibrium.
[1]
Determine K for Q
⇌ P(g).
[1]
Explain why the temperature should be stated with a value of K.
[1]
Question 17
Consider the equilibrium:
2SO₂
- O₂
[1]
⇌ 2SO₃
[1]
ΔH < 0
[1]
Predict the effect of increasing pressure on the amount of SO₃.
[1]
Explain the prediction in (a).
[1]
State the effect of increasing pressure on K at constant temperature.
[1]
Question 18
Distinguish between a homogeneous equilibrium and a heterogeneous equilibrium.
Question 19
Carbon dioxide dissolves in water according to the equilibrium:
CO₂
⇌ CO₂(aq)
[1]
Predict the effect of increasing the pressure of CO₂
[1]
on the amount of dissolved CO₂.
[1]
Explain the prediction using Le Châtelier's principle.
[2]
Question 20
For the equilibrium
H₂
- I₂
[1]
⇌ 2HI
[1]
= 64 at a certain temperature. At one instant, [H₂] = 0.20 mol dm⁻³, [I₂] = 0.10 mol dm⁻³ and [HI] = 1.00 mol dm⁻³.
[1]
Calculate Q.
[2]
Determine the direction in which the reaction proceeds to reach equilibrium.
[1]
Question 21
For a reaction mixture at a particular time, Q = K.
State what this indicates about the mixture.
[1]
Explain why no net change in concentrations occurs.
[1]
Question 22
A sealed flask containing a volatile liquid is connected to a pressure sensor. The pressure of vapour in the flask is recorded after sealing.
Describe the change in vapour pressure before equilibrium is reached.
[1]
Identify the evidence from the graph that dynamic equilibrium has been reached.
[1]
Explain, in terms of rates, why the pressure becomes constant.
[1]
Question 23
The table gives equilibrium constants for several reactions at the same temperature.
| Reaction | Equilibrium reaction | K at 298 K |
|---|---|---|
| A | N2(g) + O2(g) ⇌ 2NO(g) | 4.8 × 10^-31 |
| B | N2O4(g) ⇌ 2NO2(g) | 0.20 |
| C | CH3COOH(l) + C2H5OH(l) ⇌ CH3COOC2H5(l) + H2O(l) | 4.0 |
| D | Fe3+(aq) + SCN-(aq) ⇌ FeSCN2+(aq) | 1.4 × 10^2 |
Identify the reaction that is most product-favoured.
[1]
Identify the reaction that is least product-favoured.
[1]
For one reaction, write the value of K for the reverse reaction at the same temperature.
[1]
Explain why comparing K values is only valid here because the temperature is the same.
[1]
Question 24
A student investigates the equilibrium X
⇌ X(aq) in water using gas pressure above the solution. The graph shows dissolved X concentration at different applied pressures.
[1]
Describe the relationship shown by the graph.
[1]
Predict the effect of decreasing pressure on the position of equilibrium.
[1]
Explain your answer using Le Châtelier's principle.
[1]
Question 25
For A
A.
0.100 mol dm⁻³
B.
0.250 mol dm⁻³
C.
0.400 mol dm⁻³
D.
2.00 mol dm⁻³
Question 26
For HA(aq) ⇌ H⁺(aq) + A⁻(aq), K = 1.6 × 10⁻⁵. The initial concentration of HA is 0.100 mol dm⁻³ and ionization is small.
A.
4.0 × 10⁻³ mol dm⁻³
B.
1.6 × 10⁻⁶ mol dm⁻³
C.
0.100 mol dm⁻³
D.
1.3 × 10⁻³ mol dm⁻³
Question 27
At 298 K, ΔG° = +5.70 kJ mol⁻¹ for a reaction.
A.
5700
B.
0.100
C.
10.0
D.
1.00
Question 28
At 350 K, ΔG° = 0 for a reaction.
A.
350
B.
8.31
C.
1
D.
0
Question 29
The equilibrium mixture below is pale brown because NO₂ is brown and N₂O₄ is colourless.
N₂O₄
⇌ 2NO₂
[1]
ΔH > 0
[1]
Predict the colour change when the mixture is heated at constant volume.
[1]
Explain the prediction in (a).
[1]
State the effect of heating on K for the reaction as written.
[1]
Question 30
For A
⇌ 2B(g), K = 0.500 at a fixed temperature. Initially [A] = 1.00 mol dm⁻³ and [B] = 0.
[1]
Complete the equilibrium concentrations in terms of x, where x is the decrease in [A].
[1]
Write the equation that would be solved to determine x.
[1]
State why K is not changed by the initial concentration chosen.
[1]
Question 31
For the weak acid equilibrium
HA(aq) ⇌ H⁺(aq) + A⁻(aq)
K = 2.5 × 10⁻⁵ and the initial concentration of HA is 0.200 mol dm⁻³. Assume the small-K approximation is valid.
Write the approximate expression used for K in terms of x = [H⁺]eqm.
[1]
Calculate [H⁺]eqm.
[1]
Question 32
At 298 K, an equilibrium has K = 2.50 × 10⁴.
Calculate ΔG° in kJ mol⁻¹. Use R = 8.31 J K⁻¹ mol⁻¹.
[1]
Question 33
For the equilibrium
N₂O₄
⇌ 2NO₂
[1]
at a fixed temperature, an equilibrium mixture contains [N₂O₄] = 0.120 mol dm⁻³ and [NO₂] = 0.0600 mol dm⁻³.
[1]
Calculate K.
[1]
If the equilibrium was reached from pure N₂O₄, determine the initial concentration of N₂O₄.
[1]
Question 34
At 400 K, ΔG° = +12.0 kJ mol⁻¹ for a reaction.
Calculate K.
[1]
State whether reactants or products are favoured under standard conditions.
[1]
Question 35
The graph shows the concentrations of reactant R and product P for a gaseous equilibrium. At one time the volume is decreased; at a later time the temperature is increased.
Identify the time at which equilibrium is first reached.
[1]
State which species is favoured immediately after the volume decrease.
[1]
Deduce which side has fewer moles of gas.
[1]
Suggest whether the forward reaction is exothermic or endothermic, using the temperature change.
[1]
Question 36
An aqueous equilibrium contains a yellow species Y and a blue species B. The graph shows absorbance at a wavelength mainly absorbed by B after different solutions are added.
State which addition increases the concentration of B at equilibrium.
[1]
Describe the evidence for your answer in (a).
[1]
Explain how removal of a reactant affects the equilibrium position.
[1]
State the effect of these concentration changes on K.
[1]
Question 37
A reaction A
| Time / s | [A] / mol dm^-3 | [B] / mol dm^-3 | [C] / mol dm^-3 | K / dm3 mol^-1 |
|---|---|---|---|---|
| 10 | 0.750 | 0.550 | 0.050 | 2.00 |
| 30 | 0.680 | 0.480 | 0.120 | 2.00 |
| 75 | 0.580 | 0.380 | 0.220 | 2.00 |
| 150 | 0.520 | 0.320 | 0.280 | 2.00 |
| 300 | 0.500 | 0.300 | 0.300 | 2.00 |
- B
[1]
⇌ C
[1]
is monitored after mixing. The table gives concentrations at several times and K at the experimental temperature.
[1]
Calculate Q for one time before equilibrium using the table.
[1]
Compare Q with K and state the direction of the reaction at that time.
[1]
Identify the row that represents equilibrium.
[1]
Explain why Q changes but K does not during the experiment.
[1]
Question 38
For X(aq) ⇌ Y(aq), K = 1.0 × 10⁻⁴. A student assumes that if the initial [X] is 0.500 mol dm⁻³ and [Y] is 0, then [X]eqm ≈ 0.500 mol dm⁻³.
Explain why this approximation is reasonable.
[1]
Estimate [Y]eqm.
[1]
Question 39
For the equilibrium M
⇌ 2N(g), the graph shows equilibrium concentrations obtained from different initial concentrations of M, with no N initially present.
[1]
Use the graph to determine [M]eqm and [N]eqm for one experiment.
[1]
Calculate K for that experiment.
[1]
Explain why similar K values should be obtained from the other experiments at the same temperature.
[1]
Suggest one reason why an experimental K value might differ slightly between trials.
[1]
Question 40
The graph shows ΔG° as a function of ln K at a fixed temperature for several reactions.
Describe the relationship between ΔG° and ln K.
[1]
Use the graph to identify the reaction for which K = 1.
[1]
Explain what a negative ΔG° indicates about the equilibrium position.
[1]
State why the slope of the graph depends on temperature.
[1]
Question 41
For the reaction 2D
⇌ E(g), the graph shows Q after a sudden compression of the equilibrium mixture. K at the temperature is shown as a horizontal line.
[1]
State the relationship between Q and K immediately after compression.
[1]
Deduce the direction in which the system changes after compression.
[1]
Explain how this direction is consistent with the pressure change.
[1]
Question 42
The Haber process involves the equilibrium:
N₂
- 3H₂
[1]
⇌ 2NH₃
[1]
ΔH < 0
[1]
Explain the effect of increasing pressure on the equilibrium yield of ammonia and on K.
[1]
Explain the effect of increasing temperature and adding an iron catalyst on the equilibrium yield of ammonia, on K, and on the time taken to reach equilibrium.
[1]
Question 43
A saturated sodium chloride solution in contact with undissolved sodium chloride and a reversible gas-phase reaction are both examples of equilibrium systems.
Describe two general characteristics of dynamic equilibrium.
[1]
Compare and contrast physical and chemical equilibria, including phase and microscopic change.
[1]
Question 44
A sealed tube contains the equilibrium:
2NO₂
⇌ N₂O₄
[1]
ΔH < 0
NO₂ is brown and N₂O₄ is colourless. A student suggests that cooling the tube and adding a catalyst are both useful ways to increase the equilibrium amount of N₂O₄.
[1]
Explain the effect of cooling on the colour and K.
[1]
Evaluate the student's suggestion about cooling and adding a catalyst.
[1]
Question 45
A weak acid HA is diluted and its pH is measured. The table gives initial [HA] and measured pH for each solution.
| Trial | Initial [HA] / mol dm⁻³ | pH |
|---|---|---|
| 1 | 0.100 | 2.87 |
| 2 | 0.0500 | 3.03 |
| 3 | 0.0100 | 3.38 |
| 4 | 0.00100 | 3.91 |
| 5 | 0.000100 | 4.52 |
Convert one pH value into [H⁺].
[1]
Use [H⁺] to estimate K for HA ⇌ H⁺ + A⁻ in that trial.
[1]
State whether the small-K approximation is justified for the selected trial.
[1]
Suggest why the calculated K values may vary at very low acid concentrations.
[1]
Question 46
For the equilibrium
A
- 2B
[1]
⇌ C
[1]
the value of K at 600 K is 2.0 × 10⁻³.
[1]
Deduce the equilibrium constant expression and the meaning of the magnitude of K.
[1]
Discuss how K and the equilibrium composition are affected by reversing the equation, changing the initial amounts, and changing temperature.
[1]
Question 47
For the equilibrium
CO
- Cl₂
[1]
⇌ COCl₂
[1]
= 25.0 at a certain temperature. A mixture contains [CO] = 0.200 mol dm⁻³, [Cl₂] = 0.100 mol dm⁻³ and [COCl₂] = 0.300 mol dm⁻³.
[1]
Calculate Q and determine the initial direction of change.
[1]
Evaluate how Q, K and the concentrations change as equilibrium is approached at constant temperature.
[1]
Question 48
At 700 K, the equilibrium
H₂
- CO₂
[1]
⇌ H₂O
[1]
- CO
[1]
has K = 0.640. Initially, [H₂] = [CO₂] = 1.00 mol dm⁻³ and [H₂O] = [CO] = 0.
[1]
Set up the equilibrium concentrations in terms of x.
[1]
Determine the equilibrium concentrations and explain why no pressure shift is expected for this reaction at constant temperature.
[1]
Question 49
For a reaction at 298 K, ΔG° = −17.1 kJ mol⁻¹.
Calculate K at 298 K.
[1]
Discuss how ΔG°, K and the position of equilibrium are related, including the special case K = 1.
[1]
Question 50
A weak acid HA has K = 6.4 × 10⁻⁵ at 298 K. A solution is prepared with initial [HA] = 0.250 mol dm⁻³.
Use the small-K approximation to calculate [H⁺] and pH.
[1]
Evaluate the approximation and explain how the equilibrium law links the value of K to acid strength.
[1]
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