196 5 Energetics/thermochemistry
■ Table 5.4 Enthalpies
of combustion of fuels
Fuel
Main component
Formula and
standard state
ΔH c of main
component/kJ mol –1
Hydrogen
Hydrogen
H2(g)
−286
Compressed natural
gas (CNG)
90% methane
CH4(g)
−890
Liquid petroleum
gas (LPG)
95% propane
C3H8(g)
−2219
Methanol
Methanol
CH3OH(l)
−726
Alcohol
Ethanol
C2H5OH(l)
−1367
Another relevant property of fuels is the energy density of the fuel (Table 5.5). This is the amount of
thermal energy (heat) released by 1 kilogram of the fuel. The energy density is calculated from the
standard enthalpy of combustion of the fuel and the mass of one mole of the fuel. Petrol has an energy
density of approximately 46 000 kJ kg−1 which makes it a very concentrated energy source. Liquid
hydrogen has a higher density than petrol, but currently storage problems on board a vehicle are one
reason why its use is limited, along with the very low temperatures required to keep it as a liquid.
■ Table 5.5 Energy
densities of fuels
Additional
Perspectives
Fuel
Formula
ΔHc
Mass of one mole/g
Energy density/kJ kg –1
Hydrogen
H2(g)
−286
2
143 000
Methane
CH4(g)
−890
16
27 800
Methanol
CH3OH(l)
−2219
32
22 700
Ethanol
C2H5OH(l)
−1367
46
30 000
Feasibility of reactions
There are many examples of reactions which are spontaneous. The vast majority of these
reactions are exothermic. Hence it appears that the enthalpy change, ΔH, is a reliable guide
to which direction a reaction will go. However, there are examples of endothermic reactions
that occur without the need for heat to initiate the reaction, for example, the reaction between
citric acid and a solution of sodium hydrogencarbonate. Some salts dissolve endothermically
in water. Chapter 15 introduces a factor, known as entropy, that, in conjunction with enthalpy
and temperature, determines whether or not reactions occur at a specified temperature.
■ Examination questions –
a selection
Paper 1 IB questions and IB style questions
Q1 When 0.3205 g of methanol is completely
combusted under a water filled flame combustion
calorimeter, the temperature of 1x103 cm3 of water
is raised by 1.5°C. (Molar mass of methanol =
32.05 g mol−1; specific heat capacity of water =
4.18 J g−1 °C1.)
What is the expression for the molar enthalpy of
combustion of methanol?
1×103 × 4.18 × 1.5 × 32.05
A −
0.3205
1×103 × 4.18 × (273.00 + 1.5) × 32.05
B −
0.3205 × 1 × 103
1×103 × 4.18 × 1.5 × 32.05
C −
0.3205 × 1 × 103
0.3205 × 1 × 103
D −
1 × 103 × 4.18 × 1.5 × 32.05
Q2 Which of the following reactions would you
expect to provide the largest amount of heat?
A C2H6(l) + 7O2(l) → 4CO2(g) + 6H2O(g)
B C2H6(l) + 7O2(g) → 4CO2(g) + 6H2O(g)
C C2H6(g) + 7O2(g) → 4CO2(g) + 6H2O(g)
D C2H6(g) + 7O2(g) → 4CO2(g) + 6H2O(l)
Q3 Why does the temperature of boiling water
remain constant even though heat is supplied at
a constant rate?
A Heat is lost to the surroundings.
B The heat is used to break the covalent bonds
in the water molecules.
C Heat is also taken in by the container.
D The heat is used to overcome the
intermolecular forces of attraction between
water molecules.
Standard Level Paper 1, Nov 2005, Q14
Q4 When 0.050 mol of nitric acid is reacted with
0.050 mol of potassium hydroxide in water, the
temperature of the system increases by 13.7 °C.
Examination questions 197
Calculate the enthalpy of reaction in kJ mol−1.
HNO3(aq) + KOH(aq) → KNO3(aq) + H2O(l)
Assume that the heat capacity of the system was
209.2 J °C−1.
C −2.87 kJ mol−1
A +57.3 kJ mol−1
B +2.87 kJ mol−1
D −57.3 kJ mol−1
Q5 What can be deduced about the relative stability
of the reactants and products and the sign of ΔH,
from the enthalpy level diagram below?
Relative stability
Sign of ΔH
A products more stable
−
B products more stable
+
C reactants more stable
−
D reactants more stable
+
reactants
$H
products
Standard Level Paper 1, May 1999, Q16
B The temperature of the system has fallen to
room temperature.
C The solid which forms insulates the system,
preventing heat loss.
D Heat is gained from the surroundings as
the solid forms, maintaining a constant
temperature.
Q9 Consider the following equation:
6CO2(g) + 6H2O(l) → C6H12O6(s) + 6O2(g)
ΔH = 2824 kJ mol−1
What is the enthalpy change associated with the
production of 100.0 g of C6H12O6?
A 157 kJ
C 508 kJ
B 282 kJ
D 1570 kJ
Q10 N2(g) + O2(g) → 2NO(g)
ΔH = 180.4 kJ mol−1
N2(g) + 2O2(g) → 2NO2(g) ΔH = 66.4 kJ mol−1
Use the enthalpy values to calculate ΔH for the
reaction:
1
NO(g) + O2(g) → NO2(g)
2
A −57 kJ mol−1
B −114 kJ mol−1
C 57 kJ mol−1
D 114 kJ mol−1
Standard Level Paper 1, May 2000, Q18
Q11 The heating curve for 10 g of a substance is given
below. How much energy would be required to
melt completely 40 g of the substance that is
initially at 10 °C?
50
Temperature/°C
Q6 The specific heat capacities of some metals are
given below.
Metal
Specific heat capacity (J g−1 K−1)
copper
0.385
magnesium
1.020
mercury
0.138
platinum
0.130
If 100 kJ of heat is added to 10.0 g samples of
each of the metals above, which are all at 25 °C,
which metal will have the lowest temperature?
A copper
C mercury
B magnesium
D platinum
30
C mercury
D platinum
10
Q7 The bond energy for the H–F bond is equal to the
enthalpy change for which process?
A H+(g) + F− (g) → HF(g)
B HF(g) → H(g) + F(g)
C
1
F (g)
2 2
+
D HF(g) →
1
H (g) → HF(g)
2 2
1
1
F (g) + H2(g)
2 2
2
Q8 When a sample of a pure hydrocarbon (melting
point 85 °C) cools, the temperature is observed to
remain constant as it solidifies. Which statement
accounts for this observation?
A The heat released in the change of state
equals the heat loss to the surroundings.
0
A 4800 J
B 2400 J
400
800
1200
Energy added/J
1600
C 1600 J
D 800 J
Q12 The bond energies for H2, I2 and HI are 432,
149 and 295 kJ mol−1, respectively. From these
data, what is the enthalpy change (in kJ) for the
reaction below?
H2(g) + I2(g) → 2HI(g)
A +9
C −286
B +286
D −9
198 5 Energetics/thermochemistry
Q13 The specific heat capacity of aluminum is
0.900 J g −1 K−1. What is the heat energy
change, in J, when 10.0 g of aluminum is
heated and its temperature increases from
15.0 °C to 35 °C?
A +180
C +1800
B +315
D +2637
Higher Level Paper 1, May 2013, Q14
Q14 Which reaction represents the average bond
enthalpy of the Si–H bond in silane, SiH4?
A
B
C
D
1
SiH4(g)
4
1
SiH4(g)
4
1
SiH4(g)
4
1
SiH4(g)
4
1
1
2
→ 4 Si(g) + H2(g)
→
→
→
1
1
SiH2(g) + 4 H2(g)
4
1
Si(g) + H(g)
4
1
Si(s) + H(g)
4
Q15 Which of the following is the correct equation
for the standard enthalpy change of formation of
carbon monoxide?
1
A C(s) + O2(g) → CO(g)
B C(g) +
2
1
O (g)
2 2
→ CO(g)
C C(g) + O(g) → CO(g)
D 2C(s) + O2(g) → 2CO(g)
Q16 The bond enthalpies for H2(g) and HF(g) are
435 kJ mol−1 and 565 kJ mol−1, respectively. For the
1
2
1
2
reaction H2(g) + F2(g) → HF(g), the enthalpy of
reaction is −268 kJ mol−1 of HF produced.
What is the bond energy of F2 in kJ mol−1?
A 464
C 243
B 138
D 159
Q17 The standard enthalpy change of formation
values of two oxides of phosphorus are:
P4(s) + 3O2(g) → P4O6(s) ΔH f = −1600 kJ mol−1
P4(s) + 5O2(g) → P4O10(s) ΔH f = −3000 kJ mol−1
What is the enthalpy change, in kJ mol−1, for the
reaction below?
P4O6(s) + 2O2(g) → P4O10(s)
A +4600
B +1400
C −1400
D −4600
Paper 2 IB questions and IB style questions
Q1 a i Define the term average bond enthalpy. [3]
ii Explain why the fluorine molecule, F2, is
not suitable as an example to illustrate
the term average bond enthalpy.
[1]
b i
Using values from page 11 of the IB
Chemistry data booklet, calculate the
enthalpy change for the following reaction:
CH4(g) + F2(g) → CH3F(g) + HF(g)
[3]
ii Sketch an enthalpy diagram for the
reaction.
[2]
iii Without carrying out a calculation, suggest,
with a reason, how the enthalpy change
for the following reaction compares with
that of the previous reaction.
CH3F(g) + F2(g) → CH2F2(g) + HF(g) [2]
Q2 In aqueous solution, lithium hydroxide and
hydrochloric acid react as follows.
LiOH(aq) + HCl(aq) → LiCl(aq) + H2O(l)
The data below are from an experiment to determine
the standard enthalpy change of this reaction.
50.0 cm3 of a 0.500 mol dm−3 solution of LiOH
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 = 20.6 °C
Final temperature of the mixture = 24.1 °C
a State, with a reason, whether the reaction is
exothermic or endothermic.
[1]
b Explain why the solutions were mixed rapidly. [1]
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.
[4]
d Identify the major source of error in the
experimental procedure described above.
Explain how it could be minimized.
[2]
e The experiment was repeated but with an
HCl concentration of 0.520 mol dm−3 instead
of 0.500 mol dm−3. State and explain what
the temperature change would be.
[2]
Q3 a Define the term standard enthalpy change
of formation.
[2]
b Define the term standard enthalpy change
of combustion.
[2]
c State Hess’s law.
[1]
d Calculate the standard enthalpy change
of formation of propane, C3H8, given the
following standard enthalpies of combustion:
ΔH c [C3H8(g)] = −2220 kJ mol−1
ΔH c [Cgraphite(s)] = −393 kJ mol−1
ΔH c [H2(g)] = −286 kJ mol−1
Draw an energy cycle and use Hess’s law to
produce an equation for the enthalpy change
of formation.
[4]