Topic 5 Energetics Practice Questions
1. Which statements about exothermic reactions are correct?
I.
They have negative H values.
II. The products have a lower enthalpy than the reactants.
III. The products are more energetically stable than the reactants.
a.
b.
c.
d.
I and II only
I and III only
II and III only
I, II and III
2. Which of the quantities in the enthalpy level diagram below is (are) affected by the use of a
catalyst?
a. I only
b. III only
I
II
Enthalpy
c. I and II only
III
d. II and III only
3. Which statements are correct for an endothermic reaction?
I.
The system absorbs heat.
II. The enthalpy change is positive.
III. The bond enthalpy total for the reactants is greater than for the products.
a.
b.
c.
d.
I and II only
I and III only
II and III only
I, II and II
4. When a sample of NH4SCN is mixed with solid Ba(OH)2*8H2O in a glass beaker, the mixture changes
to a liquid and the temperature drops sufficiently to freeze the beaker to the table. Which statement is
true about the reaction?
a. The process is endothermic and ∆H is –
b. The process is endothermic and ∆H is +
c. The process is exothermic and ∆H is –
d. The process is exothermic and ∆H is +
5. Which one of the following statements is true of all exothermic reactions?
a. They produce gases
b. They give out heat
c. They occur quickly
d. They involve combustion
6. If 500J of heat is added to 100.0g samples of each of the substances below, which will have the largest
temperature increase?
a. Gold – 0.126 Jg-1K-1
b. Silver – 0.237 Jg-1K-1
c. Copper – 0.385 Jg-1K-1
d. Water – 4.18 Jg-1K-1
7. The specific heat of metallic mercury is 0.138 Jg-1°C-1. If 100.0 J of heat is added to a 100.0g sample
of mercury at 25.0°C, what is the final temperature of mercury?
8. The temperature of a 2.0g sample of aluminum increases from 25°C to 30°C. How many joules of heat
energy were added? (specific heat of Al = 0.90 Jg-1K-1)
a. 0.36
b. 2.3
c. 9.0
d. 11
9. A sample of a metal is heated. Which of the following are needed to calculate the heat absorbed by
the sample?
I.
The mass of the sample
II. The density of the sample
III. The specific heat capacity of the sample
a. I and II only
b. I and III only
c. II and III only
d. I, II and III
10. What is the energy change (in kJ) when the temperature of 20 g of water increases by 10°C?
a. 20 × 10 × 4.18
b. 20 × 283 × 4.18
c.
d.
20  10  4.18
1000
20  283  4.18
1000
11. In aqueous solution, potassium hydroxide and hydrochloric acid react as follows.
KOH(aq) + HCl(aq)→ KCl(aq)+ H2O(l)
The data below is from an experiment to determine the enthalpy change of this reaction.
50.0 cm3 of a 0.500 mol dm−3 solution of KOH was mixed rapidly in a glass beaker with
50.0 cm3 of a 0.500 mol dm−3 solution of HCl.
Initial temperature of each solution = 19.6°C
Final temperature of the mixture = 23.1°C
(a) State, with a reason, whether the reaction is exothermic or endothermic.
(b) Explain why the solutions were mixed rapidly.
(c) Calculate the enthalpy change of this reaction in kJ mol−1. Assume that the specific heat capacity
of the solution is the same as that of water.
(d) Identify the major source of error in the experimental procedure described above. Explain how
it could be minimized.
(e)
The experiment was repeated but with an HCl concentration of 0.510 mol dm−3 instead
of 0.500 mol dm−3. State and explain what the temperature change would be.
12. The mass of the burner and its content is measured before and after the experiment. The thermometer
is read before and after the experiment. What are the expected results?
a. Mass: decreases; Thermometer: increases
b. Mass: decreases; Thermometer: stays the same
c. Mass: increases; Thermometer: increases
d. Mass: increases; Thermometer: stays the same
13. The heat released from the combustion of 0.0500g of white phosphorous increases the temperature of
150.00g of water from 25.0°C to 31.5°C. Calculate a value for the enthalpy change of combustion of
phosphorous. Discuss possible sources of error in the experiment.
14. When 8.00g of ammonium nitrate completely dissolved in 100 cm3 of water, the temperature fell from
19.0°C to 14.5°C. Calculate the enthalpy change of solution.
15. The data below are from an experiment to measure the enthalpy change for the reaction of
aqueous copper(II) sulfate, CuSO4(aq) and zinc, Zn(s).
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
50.0 cm3 of 1.00 mol dm–3 copper(II) sulfate solution was placed in a polystyrene cup and zinc
powder was added after 100 seconds. The temperature-time data was taken from a data-logging
software program. The table shows the initial 23 readings.
A straight line has been drawn through some of the data points. The equation for this line is
given by the data logging software as
T = –0.050t + 78.0
where T is the Temperature at time t.
(a) The heat produced by the reaction can be calculated from the temperature change, ΔT,
using the expression below.
Heat change = Volume of CuSO4(aq) × Specific heat capacity of H2O × ∆T
Describe two assumptions made in using this expression to calculate heat changes.
(2)
(b)
(i)
Use the data presented by the data logging software to deduce the
temperature change, ∆T, which would have occurred if the reaction had taken place
instantaneously with no heat loss.
(2)
(ii)
State the assumption made in part (b)(i).
(1)
(iii) Calculate the heat, in kJ, produced during the reaction using the expression
given in part (a).
(1)
(c) The colour of the solution changed from blue to colourless. Deduce the amount, in moles, of
zinc which reacted in the polystyrene cup.
(1)
–1
(d)
Calculate the enthalpy change, in kJ mol , for this reaction.
(1)
16. Using the equations below:
C(s) + O2(g) → CO2(g)
∆H = −390 kJ
Mn(s) + O2(g) → MnO2(s) ∆H = −520 kJ
What is ∆H (in kJ) for the following reaction?
MnO2(s) + C(s)  Mn(s) + CO2(g)
a. 910
b. 130
c. −130
d. −910
17. Using the equations below
O2(g) → CuO(s) ∆Hο = −156 kJ
2Cu(s) + 12 O2(g) → Cu2O(s) ∆Hο = −170 kJ
What is the value of ∆Hο (in kJ) for the following reaction?
Cu(s) +
a.
b.
c.
d.
1
2
2CuO(s) → Cu2O(s) +
142
15
–15
–142
18. Calculate the standard enthalpy change ∆Hθ for the reaction:
2NO(g) + O2(g)  2NO2(g)
Using the information below:
N2(g) + O2(g)  2NO(g)
∆Hθ = +180.5 kJ
N2(g) + 2O2(g)  2NO2(g) ∆Hθ = +66.4 kJ
19. Calculate the standard enthalpy change ∆Hθ for the reaction:
C2H4(g) + 3O2(g)  2CO2(g) + 2H2O(l)
Using the information below:
2C2H5OH(l)  2C2H4(g) + 2H2O(l)
∆Hθ =-86 kJ
2CO2(g) + 3H2O(l)  C2H5OH(l) + 3O2(g) ∆Hθ = +1386 kJ
20. Calculate the standard enthalpy change ∆Hθ for the reaction:
2C(graphite) + 2H2(g)  2C2H4(g)
Using the information below:
2C(graphite) + 2O2(g)  2CO2(g) ∆Hθ = -790 kJ
2H2(g) + O2(g)  2H2O(l)
∆Hθ = -574 kJ
2CO2(g) + 2H2O(l)  C2H4(g)
∆Hθ = +1416 kJ
21. Calculate the standard enthalpy change ∆Hθ for the reaction:
C(graphite) + ½ O2(g)  CO (g)
Using the information below:
C(graphite) + ½ O2(g)  CO2(g)
∆Hθ = -394 kJ
CO(g) + ½ O2(g)  CO2(g)
∆Hθ =-283 kJ
1
2
O2(g)
22. What is ∆H for the reaction below in kJ?
CS2(g) + 3O2(g)  CO2(g) + 2SO2(g)
[∆Hf / kJ mol–1: CS2(g) 110, CO2(g) – 390, SO2(g) – 290]
a. −570
b. −790
c. −860
d. −1080
23. The standard enthalpy change of formation of Al2O3(s) is −1669 kJ mol−1 and the standard
enthalpy change of formation of Fe O (s) is −822 kJ mol−1.
2 3
(i)
Use these values to calculate ∆Hο for the following reaction.
Fe2O3(s) + 2Al(s) → 2Fe(s) + Al2O3(s)
State whether the reaction is exothermic or endothermic.
(ii)
Draw an enthalpy level diagram to represent this reaction. State the conditions under
which standard enthalpy changes are measured.
24. Define the term standard enthalpy of formation, and write the equation for the standard enthalpy
of formation of ethanol.
25. Which of the following does not have a standard heat of formation of zero at 25°C and 1.00x105
Pa?
a. Cl2(g)
b. I2(s)
c. Br2(g)
d. Na(s)
26. Which of the following does have a standard heat of formation value of zero at 25°C and
1.00x105 Pa?
a. H(g)
b. Hg(s)
c. C(diamond)
d. Si(s)
27. Write the thermochemical equation for the standard enthalpy of formation of propanone
CH3COCH3.
28. Calculate ∆Hθ (in kJ mol-1) for the reaction:
Fe3O4(s) + 2C(graphite)  3Fe(s) + 2CO2(g)
Substance: Fe3O4(s); Heat of Formation: -1118 kJ mol-1
Substance: CO2(g); Heat of Formation = - 394 kJ mol-1
29. Enthalpy changes of combustion data are tabulated in section 13 of the IB Data Booklet.
a. Give a chemical equation for the formation of benzene (thermochemical equation).
b. Calculate ∆Hθf for benzene in the following reaction. Use the data in section 13 to calculate the
enthalpy of formation for benzene. (section 13 gives the values for the ∆H θrxn.
C6H6(l) + 7 ½ O2(g)  6CO2(g) + 3H2O(g)
30. What energy changes occur when chemical bonds are formed and broken?
a. Energy is absorbed when bonds are formed and when they are broken.
b. Energy is released when bonds are formed and when they are broken.
c. Energy is absorbed when bonds are formed and released when they are broken.
d. Energy is released when bonds are formed and absorbed when they are broken.
31. The average bond enthalpies for O—O and O==O are 146 and 496 kJ mol−1 respectively.
What is the enthalpy change, in kJ, for the reaction below?
H—O—O—H(g)  H—O—H(g) + ½O==O(g)
a.
b.
c.
d.
– 102
+ 102
+ 350
+ 394
32. Which reaction has the most negative ∆Hο value?
a. LiF(s) → Li+(g) + F−(g)
b. Li+(g) + F–(g) → LiF(s)
c. NaCl(s) → Na+(g) + Cl−(g)
d. Na+(g) + Cl−(g) → NaCl(s)
33. For which of the following is the sign of the enthalpy change different from the other three?
a. CaCO3(s) → CaO(s)+ CO2(g)
b. Na(g) → Na+(g) + e−
c. CO2(s) → CO2(g)
d. 2Cl(g) → Cl2(g)
34. The average bond enthalpy for the C–H bond is 412 kJ mol−1. Which process has an enthalpy change
closest to this value?
a. CH4(g) → C(s) + 2H2(g)
b. CH4(g) → C(g) + 2H2(g)
c. CH4(g) → C(s) + 4H(g)
d. CH4(g) → CH3(g) + H(g)
35. Identify the bonds which are broken in the following process.
C2H6(g)  2C(g) + 6H(g)
36. Use the bond enthalpies below to calculate ∆H for the reaction:
2H2(g) + O2(g)  2H2O(g)
O=O; +498 kJ mol-1
H – H ; +436 kJ mol-1
O – H ; +464 kJ mol-1
37.
Using the bond enthalpies from Reference section 11, calculate ∆H for the following reaction:
C3H6(g) + H2(g)  C3H8(g)
38.
Using the bond enthalpies from Reference section 11, calculate ∆H for the following reaction:
C4H8(g) + 6O2(g)  4CO2(g) + 4H2O(g)
39.
(a)
Define the term average bond enthalpy, illustrating your answer with an equation for
methane, CH4.
(b)
The equation for the reaction between methane and chlorine is
CH4(g) + Cl2(g) → CH3Cl(g) + HCl(g)
Use the values from section 11 of the Data Booklet to calculate the enthalpy change for
this reaction.
(c)
Explain why no reaction takes place between methane and chlorine at room
temperature unless the reactants are sparked, exposed to UV light or heated.
(d)
Draw an enthalpy level diagram for this reaction.
40. The concentration of ozone in the upper atmosphere is maintained by the following reactions.
i. O2 2O2·
ii. O2 + O·  O3
iii. O3  O2 + O·
a. Identify the step which is exothermic. Why.
b. Identify with reference to bonding in O2 and O3, the most endothermic step.
41. Which combination of ionic charge and ionic radius give the largest lattice enthalpy for an ionic
compound?
A.
B.
C.
D.
Ionic charge
high
high
low
low
Ionic radius
large
small
small
large
42. The lattice enthalpy values for lithium fluoride and calcium fluoride are shown below.
LiF(s)
∆Hο = +1022 kJ mol−1
∆Hο = +2602 kJ mol−1
CaF2(s)
Which of the following statements help(s) to explain why the value for lithium fluoride is less
than that for calcium fluoride?
I.
The ionic radius of lithium is less than that of calcium.
II. The ionic charge of lithium is less than that of calcium.
a. I only
b. II only
c. I and II
d. Neither I nor II
43. Identify the process, which has the sign of its associated enthalpy change different from the rest?
a. Cl(g) + e-  Cl-(g)
b. K(g)  K+(g) + ec. KCl(g)  K+(g) + Cl-(g)
d. Cl2 (g)  2Cl(g)
44. Answer the following questions using the diagram:
a. Write an equation to represent the lattice
energy of potassium oxide, K2O. The BornHaber cycle shown may be used to calculate
the lattice energy of potassium oxide
b. Identify the enthalpy changes labelled by the
letters W,X,Y and Z
c. Use the energy cycle, and further information
from section 7 and 11 of the IB Data booklet
to calculate an experimental value for the
lattice energy of potassium oxide.
45. The theoretical lattice enthalpies of sodium chloride and magnesium oxide are shown below.
Explain the higher lattice enthalpy of magnesium oxide compared to sodium chloride.
Compound: NaCl; Lattice enthalpy: 769 kJmol-1
Compound: MgO; Lattice enthalpy: 3795 kJmol-1
46. The lattice enthalpies of silver bromide and sodium bromide are given below
∆Hο/kJ mol-1
AgBr: 905
NaBr: 691
Explain the relative values of the lattice enthalpies with reference to the bonding.
47. Discuss the relative enthalpies of hydration of the K+ and F- ions in relation to their ionic radii.
48. Use an energy cycle to calculate the enthalpy of solution of potassium chloride from data in
sections 18 and 20 of the IB data booklet.
Calculate the % inaccuracy of your value by comparing with the value in section 19 and
comment on the disagreement between the two values.
49.
Which reaction has the greatest positive entropy change?
a. CH4(g) + 1½O2(g) → CO(g) + 2H2O(g)
b. CH4(g) + 1½O2(g) → CO(g) + 2H2O(l)
c. CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
d. CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)
50. Some chlorine gas is placed in a flask of fixed volume at room temperature. Which change will cause
a decrease in entropy?
a. Adding a small amount of hydrogen
b. Adding a small amount of chlorine
c. Cooling the flask
d. Exposing the flask to sunlight
51. Identify the process expected to have a value of ∆S closest to zero?
a. C2H4(g) +H2(g)  C2H6(g)
b. H2(g) + Cl2(g)  2HCl(g)
c. CaCO3(s)  CaO(s) + CO2(g)
d. H2O(l)  H2O(g)
52. Predict the entropy change ∆S for the following reactions.
a. N2(g) + 3H2(g)  2NH3(g)
b. 3Fe(s) + 4H2O(g)  Fe3O4(s) +4H2(g)
c. Ba(OH)2*8H2O(s) + 2NH4SCN(s)  Ba(SCN)2(aq) + 2NH3(aq) + 10H2O(l)
53. Which is the best description of the entropy and enthalpy changes accompanying the sublimations of
iodine: I2(s)  I2(g)?
a. ∆S+, ∆H+, reaction is endothermic
b. ∆S+, ∆H-, reaction is exothermic
c. ∆S-, ∆H+, reaction is endothermic
d. ∆S-, ∆H-, reaction is exothermic
54. Calculate the entropy change ∆S for the Haber process shown from tabulated standard molar entropies
at 25°C.
N2(g) + 3H2(g)  2NH3(g)
55. Under what circumstances is a reaction spontaneous at all temperatures?
∆Hο
∆Sο
A.
+
+
B.
+
–
C.
–
–
D.
–
+
56. For the process:
C6H6 (l)  C6H6(s)
The standard entropy and enthalpy changes are:
∆Hο = −9.83kJ mol−1 and ∆Sο = −35.2J K mol−1.
a. Predict and explain the effect of an increase in temperature on the spontaneity of the
process.
b. ∆Hο = −9.83 kJ mol−1 and ∆Sο = −35.2J K–1 mol–1. Calculate the temperature (in °C) at
which ∆G =0 for the above process and explain the significance of this temperature.
57. The ∆Hο and ∆Sο values for a certain reaction are both positive. Which statement is correct
about the spontaneity of this reaction at different temperatures?
a. It will be spontaneous at all temperatures.
b. It will be spontaneous at high temperatures but not at low temperatures.
c. It will be spontaneous at low temperatures but not at high temperatures.
d. It will not be spontaneous at any temperature.
58. Explain in terms of Gο, why a reaction for which both Hο andSο values are positive can
sometimes be spontaneous and sometimes not.
59. Consider the following reaction.
N2(g) +3H2(g) → 2NH3(g)
(i)
Use values from section 11 in the Data Booklet to calculate the enthalpy change,Hο,
for this reaction.
(ii)
The magnitude of the entropy change, S, at 27 °C for the reaction is 62.7 J K−1 mol–
1. State, with a reason, the sign of S.
(iii) Calculate G for the reaction at 27 °C and determine whether this reaction is
spontaneous at this temperature.
60. The equation for the decomposition of calcium carbonate is given below.
CaCO3(s) → CaO(s) + CO2(g)
At 500 K, ∆H for this reaction is +177 kJ mol–1 and ∆S is 161 J K–1 mol−1.
(a)
Explain why ∆H for the reaction above cannot be described as ∆Hfο.
(b)
State the meaning of the term ∆S.
(c)
Calculate the value of ∆G at 500 K and determine, giving a reason, whether or not
the reaction will be spontaneous
61. The standard enthalpy change for the combustion of phenol, C6H5OH(s), is −3050 kJ mol−1
at 298 K.
(a)
Write an equation for the complete combustion of phenol.
(b) The standard enthalpy changes of formation of carbon dioxide, CO2(g), and of water,
H2O(l), are −394 kJ mol−1 and −286 kJ mol−1 respectively.
Calculate the standard enthalpy change of formation of phenol, C6H5OH(s).
(c)
The standard entropy change of formation, ∆Sο, of phenol, C6H5OH(s) at
298 K is −385 J K−1 mol −1. Calculate the standard free energy change of formation,
∆Gο , of phenol at 298 K.
(d)
Determine whether the reaction is spontaneous at 298 K, and give a reason.
(e)
Predict the effect, if any, of an increase in temperature on the spontaneity of this
reaction.