MYP 10 Chemistry 2012-13
Energetics Worksheet 3
Name: _________________________________ ( )
Class: _________ Date: _____________
_________________________________________________________________________________
1. A student was required to determine the enthalpy change for the reaction between iron and copper sulphate
solution.
The student produced the following account of their experiment.
A piece of iron, mass about 3 g, was placed in a glass beaker. 50 cm3 of
0.5 mol dm–3 aqueous copper sulphate solution was measured using a
measuring cylinder and added to the beaker. The temperature of the
mixture was measured immediately before the addition and every minute
afterwards until no further change took place.
Fe + CuSO4
Timing
before
addition
Temperature/°C
22
FeSO4 + Cu
1 min
2 mins
3 mins
4 mins
5 mins
27
29
26
24
22
(a) Suggest two improvements you would make to this experiment. Give a reason for each of the
improvements suggested.
(b) In an improved version of the same experiment a maximum temperature rise of 15.2 °C occurred
when reacting excess iron with 50.0 cm3 of 0.500 mol dm–3 aqueous copper sulphate solution.
(i) Using this data and taking the specific heat capacity of all aqueous solutions as 4.18 Jg–1 deg–1
calculate the heat change.
(ii) Calculate the number of moles of copper sulphate used.
(iii) Calculate the enthalpy change of this reaction in kJ mol–1.
An excess of zinc powder was added to 20.0 cm3 of a solution of copper(II) sulphate of
concentration 0.500 mol dm–3. The temperature increased by 26.3 °C.
(a) How many moles of copper(II) sulphate were used in this experiment?
(b) Calculate the enthalpy change, ΔH, in kJ mol–1 for this reaction given that:
Energy change =
/J
Specific heat
capacity
 mass of
solution
/J g1 K1
/g
 temperature
change
/K
Assume that the mass of solution is 20.0 g and the specific heat capacity of the solution is 4.18 J g–
1K–1.
1
2. Urea, which is used as a fertillser in much of mainland Europe, Asia and Africa, is manufactured by
the reaction of ammonia and carbon dioxide.
2NH3(g) + CO2(g)  NH2CONH2(s) + H2O(l)
(a) Define the term standard enthalpy of formation, Hf
(b) Calculate the enthalpy change, H
of formation.
Substance
, of urea.
for the reaction above, given the following standard enthalpies
Hf
/ kJ mol–1
NH3(g)
–46.2
CO2(g)
–393.5
NH2CONH2(s)
–632.2
H2O(l)
–285.8
3. This question is about a self-heating can of coffee.
The bottom of the can has a compartment containing copper(II) nitrate solution. When a button on the
bottom of the can is pressed, magnesium powder is released into the compartment where it reacts with
the copper(II) nitrate solution.
(a)(i) Write an ionic equation for the reaction between magnesium powder and copper(II) ions. Include
state symbols, but omit any spectator ions.
(ii) Show how the standard enthalpy change for this reaction could be calculated from the standard
enthalpies of formation of copper(II) ions and magnesium ions. You should include a Hess cycle
in your answer.
(b) The can contains 150 g of a solution of coffee in water.
The temperature of the solution needs to increase by 60 °C to produce a hot drink.
(i) Calculate the energy change needed to produce a temperature increase of 60 °C in the coffee, using
the relationship
Energy change = 4.2 × mass of solution × temperature change.
Remember to include a unit in your answer.
(ii) The standard enthalpy change for this reaction is –530 kJ mol–1.
Calculate the number of moles of reactants needed to produce the energy change in (i).
(iii) A solution of copper(II) nitrate of concentration 8.0 mol dm–3 is used.
Use your answer to (ii) to calculate the volume, in cm3, of copper(II) nitrate solution needed.
Your answer should be given to two significant figures.
(c) Suggest TWO reasons why the temperature of the coffee may not increase by as much as 60 °C.
2
4. The enthalpy change for the thermal decomposition of calcium carbonate cannot be measured directly,
but can be found by carrying out two reactions as shown in the Hess cycle below.
CaCO 3 (s)
H reaction
H 3
CaO(s) + CO2 (g)
H 4
Elements in their standard states
(a) Suggest ONE reason why it is difficult to measure  Hreaction directly by experiment.
(b) In an experiment to find  H1 a student added 2.00 g of finely powdered calcium carbonate to
20.0 cm3 of 2.50 mol dm–3 hydrochloric acid solution (an excess) in a polystyrene container.
The temperature rose from 20.5 °C to 23.0 °C.
(i) Why is the calcium carbonate used in this experiment finely powdered, rather than in lumps?
Explain why this is important for an accurate result.
(ii) Calculate the energy change using the relationship below.
Energy change = 4.2
×
mass of solution
× temperature change
–1 –1
/J
/J g K
/g
/K
Assume that the mass of the solution is 20 g.
(iii) Calculate the enthalpy change, H1. Include a sign and units in your answer.
[The molar mass of CaCO3 is 100 g mol–1]
(iv) In another experiment, the value of  H2 was found to be –181 kJ mol–1.
Use this result and your answer to (iii) to calculate the value of  Hreaction.
(c) The student checked the experimental results using information from the Book of data in another
Hess cycle.
CaCO 3 (s)
H reaction
H 3
CaCO(s) + CO 2 (g)
H 4
Elements in their standard states
Name the enthalpy change represented by H3.
3