Summary Sheet: SL Topic 6. Energetics
Resources: Chapter 12: sections 1-11.
Online: www.dorjegurung.com/chemistry/IB_year1/Energetics.htm
Online practice tests: www.dorjegurung.com/chemistry/IB_year1/IB1_practice_tests/index.htm
Concepts to be mastered:
To master a concept, you must be able to do three things:
1. define the concept
2. explain the concept, and
3. give an example of the concept.
NB: Ignore all references to internal energy, internal energy change, and pressure-volume work.
• exothermic, endothermic, bond breaking, bond making
• heat, source, sink, state, universe, system, surroundings, boundary,
• isolated system, closed system, open system, Laws of Thermodynamics, First Law of
Thermodynamics, conservation of energy, equilibrium, thermal equilibrium,
• calorimetry, calorimeter, calorie, Calorie, Joule
• specific heat capacity (specific heat), heat capacity, molar heat capacity,
• enthalpy change, Hess’s Law (of Heat Summations), Hess cycle (energy cycle)
• enthalpy (heat) of reaction, formation reaction, enthalpy (heat) of formation, combustion reaction,
enthalpy (heat) of combustion, neutralization reaction, enthalpy (heat) of neutralization,
• standard enthalpy of formation, standard enthalpy of combustion, standard enthalpy of
neutralization
• thermodynamic stability, kinetic stability
Skills to be mastered:
To master a skill, you must be able to
1. recognize when the skill is needed,
2. recognize what information is needed to execute the skill,
3. execute the skill, and
4. assess whether the skill has been executed correctly.
• Express the first law of thermodynamics in terms of thermal sources and sinks.
• Given three of mass, specific heat, heat gained or lost, and temperature change for a homogeneous
system, compute the fourth.
• Given two of heat capacity, heat gained or lost, and temperature change for a system, compute the
third.
• Given three of mass, molar heat capacity, heat gained or lost, and temperature change for a pure
substance, compute the fourth.
• Perform calculations involving heat exchanges among multiple substances and systems.
• Perform calculations to obtain heats of reaction from calorimetric information.
• Evaluate results of experiment on enthalpy changes
• Write the formation reaction for a compound or element.
• Compute heats of reaction from heats of formation.
• Compute heats of reactions from bond enthalpies.
• Construct enthalpy diagrams featuring a particular chemical process.
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Use enthalpy diagrams or Hess cycles (energy cycles) to compute an enthalpy change for a
particular process.
Relate bond breaking to endothermicity and bond making to exothermic
relate exothermic reaction to overall the bonds made being stronger than the bond broken
relate endothermic reaction to overall the bonds broken being stronger than the bonds made.
Relate a chemical reaction’s energetics to interconversions between kinetic energy and potential
energy and the response of a thermometer
Distinguish between kinetic and thermodynamic stability
Describe compounds as stable or unstable with respect to specified substances, using the enthalpy
of reaction or of formation
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Further problems:
1. In an exothermic reaction, what changes cause energy to be released?
2. Magnesium is used in fireworks, and incendiary bombs because it burns fiercely, liberating a great
amount of heat. The reaction of magnesium with oxygen is
2 Mg (s) + O2 (g) → 2 MgO (s)
(a) Draw an enthalpy level diagram for this reaction.
(b) Is this reaction exothermic or endothermic?
(c) Which has the larger potential energy, reactants or products?
3. (a) A 100-ml of water in a beaker made of glass (specific heat capacity = 0) is heated from 25°C to
3
−1
40°C. Assuming that the density of water is 1 g/cm and specific heat capacity of water is 4.2 J K
−1
g , determine the energy gained by the water.
(b) An endothermic reaction is conducted inside a metal calorimeter submerged in a beaker of
cold water. As a consequence, the 250 ml of water and the 100.0 g calorimeter goes down in
temperature to 13°C from 25°C. Determine the total heat energy lost by the water and the
metal calorimeter to the reaction mixture if the specific heat capacity of the metal were 13.5
−1
−1
kJ °C kg .
(c) 50.0 g of water at a temperature of 24.1°C are contained in a calorimeter. To this is added
50.0 g of water at 41.7°C. The mixture is stirred immediately and a temperature of 32.6°C is
recorded as the maximum temperature reached. What is the heat capacity of the calorimeter?
4. Given the thermochemical reaction
θ
2 NO(g) + Cl2(g) → 2 NOCl(g)
ΔH = − 77.4 kJ
θ
determine the ΔH for the following reactions
(a) NO(g) + ½ Cl2(g) → NOCl(g)
(b) 6 NOCl(g) → 6 NO(g) + 3 Cl2(g)
5. Write the chemical reaction corresponding to the enthalpy of formation of Na2S2O3(s).
6. Calculate the enthalpy of combustion of methane using standard enthalpies of formation.
CH4(g) + 2 O2(g) → CO2 (g) + 2 H2O (l)
7. The sugar arabinose, C5H10O5, is burned completely in oxygen in a calorimeter.
C5H10O5 (g) + 5 O2(g) → 5 CO2 (g) + 5 H2O (l)
The calorimeter is made of metal and contains water. Burning a 0.548g sample caused the
temperature to rise from 20.00°C to 210.54°C. The heat capacity of the calorimeter and its
contents is 15.8 kJ/°C. Calculate ΔH for the combustion reaction of one mole of arabinose.
8. Write the chemical reaction corresponding to the bond enthalpy of Cl2(g).
9. Use bond enthalpy values in Table 10 of the Data Booklet to calculate the enthalpy change for the
following reactions.
(a) CH4(g) + Cl2(g) → CH3Cl (g) + HCl (g)
(b) CH2=CH2 (g) + H2O (l) → CH3CH2OH (g)
10. Cis-but-2-ene and trans-but-2-ene can both be converted to butane by the addition of one mole of
hydrogen. The enthalpies of these processes are −6.8 and −6.6 kJ/mol respectively. Which of the
two compounds is thermodynamically more stable? Explain your answer.
11. M01/1 (a) Hess.s law states that, whether a reaction occurs in one or several steps, the total
enthalpy change is the same. Illustrate your understanding of this law by using the data below to
calculatethe enthalpy change(ΔH) when one mole of solid carbon is converted into carbon
monoxide.
C (s) + O2 (g) → CO2 (g)
ΔH=−393.5 kJ
CO (g) + ½ O2 (g) → CO2 (g) ΔH=−283.0 kJ
(b) State what is meant by the term endothermic reaction. [1]
(c) Enthalpy changes may also be calculated by using bond enthalpies, some values of which
-1
(kJmol ) are provided below:
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C== C 612; C−H 412; O−H 463; C==O 743; O==O 496.
The balanced equation for the complete combustion of one mole of ethene, C2H4, in oxygen is
shown below:
C2H4 + 3O2 → 2CO2 + 2H2O
(i) Use the equation and the bond enthalpy data above to calculate the enthalpychange for the
complete combustion of one mole of ethene. [4]
(ii) State, with a brief explanation, whether the reaction is endothermic or exothermic. [1]
12. M00/5. (a) (i) Explain what is meant by the term standard enthalpy change of reaction.
[3]
(ii) Describe an experiment to determine the enthalpy change of the reaction between dilute
hydrochloric acid and aqueous sodium hydroxide. Show how the value of ΔH would be
calculated from the data obtained.
[2]
(iii) Draw an enthalpy level diagram for the neutralization reaction above. Indicate on your
diagram the enthalpy change of the reaction and hence compare the relative stabilities of
reactants and products.[4]
(b) Explain, giving one example, the usefulness of Hess’s Law in determining ΔH values.
[4]
13. M99/2. (a) Define average bond enthalpy.
[2]
−1
(b) Using the following bond enthalpies (kJ mol ):
C⎯H 412, O=O 496, C=O 743, O⎯H 463
calculate the enthalpy change for the following reaction:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
14. M98-4. Enthalpy changes of reaction can be calculated from average bond enthalpies.
(a) For the combustion of propane, C3H8:
(i) Describe the changes in bonding that take place during this reaction.
[2]
[4]
(ii) Write a balanced equation for this reaction and calculate the enthalpy change using the bond
enthalpies in the data book.
[6]
(b) It is possible using Hess’s Law, to determine the enthalpy change, ΔH, for a reaction such as
the combustion of poprane by combining reactions such as:
C(s) + O2(g) → CO2(g); ΔH = x kJ mol-1;
H2(g) + ½O2(g) → H2O(l); ΔH = y kJ mol-1;
3 C(s) + 4 H2(g) → C3H8(g); ΔH = z kJ mol-1
Show how these enthalpy changes may be arranged to give the enthalpy change for the combustion
of propane.
[3]
(c) (i) Draw a labelled relative energy diagram for the combustion of propane, showing clearly the
enthalpy of reaction, ΔH, and including its algebraic sign.
[3]
(ii) State the conditions needed when calculating a standard enthalpy change for a reaction. [1]
(iii) If the combustion of propane were incomplete, what other products, apart from carbon dioxide
and water, might be formed? State also how this would affect the magnitude of the enthalpy
change of reaction.
[3]
15. N99/1. In order to determine the enthalpy change of reaction between zinc and copper(II) sulfate, a
3
−3
student placed 50.0 cm of 0.200 mol dm copper(II) sulfate solution in a polystyrene beaker. The
temperature was recorded every 30 seconds. After two minutes 1.20 g of powdered zinc was
added. The solution was stirred and the temperature recorded every half minute for the next 14
minutes. The results obtained were then plotted to give the following graph:
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29
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point A
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Temperature / C
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25
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23
22
21
20
19
18
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16
0
1
2
3
4
5
6
7
8
9
10 11 12 13 14 15 16
Time / min
(a) Write the equation for the reaction taking place.
[1]
(b) Determine which of the two reagents was present in excess. [2]
(c) The highest temperature is reached at point A. Explain what is happening in the system at this
point.
[1]
(d) By drawing a suitable line on the graph estimate what the rise in temperature would have been
if the reaction had taken place instantaneously. [2]
(e) Calculate how much heat was evolved during the reaction. Give your answer to three
significant figures.[2]
−1
(f) What is the enthalpy change of reaction in kJmol ? [1]
−1
(g) The accepted value for the enthalpy change of reaction is −218 kJ mol . What is the
percentage error for the value obtained in this experiment?
[1]
(h) Suggest one reason why there is disagreement between the experimental value and the
accepted value. [1]
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