F.6/7 Chemistry Practical: Enthalpy of reaction
Objective:
To determine enthalpy of reaction of Zinc displacing Copper from Copper(II) sulphate solution and
Silver from Silver(I) nitrate solution
Group size: Individual
Introduction
In this practical, the following reactions are followed by measuring the change of temperature
using thermometer:
Reaction (1)
Zn(s) + CuSO4(aq) -----> ZnSO4(aq) + Cu(s)
H1
Reaction (2)
Zn(s) + AgNO3(aq) -----> Zn(NO3)2(aq) + Ag(s) H 2
This practical enables a direct determination of the heat of a reaction to be carried out.
-----------------------------------------------------------------------------------------------------------------------Theory:
1. Define the term 'molar enthalpy of reaction'.
2.
3.
Explain the term 'endothermic reaction' and 'exothermic reaction'.
Are the above reactions exothermic or endothermic? How do you know?
4.
In the following procedure, temperature change, , for the reaction of zinc and copper(II)
sulphate is determined from the graph of plotting temperature versus time. Give a formula for
the calculation of 'heat given out, Q' by using the temperature change. (Use the following
symbols in your formula: mw for mass of copper(II) sulphate solution, Cw for specific heat
5.
capacity of water,  for the temperature change)
Zinc powder is the limiting reagent in this practical. If x g of zinc is used to react with 50 cm3
of solutions. Express H1 in terms of Q and x.
---------------------------------------------------------------------------------------------------------------------Chemicals:
0.1 M copper(II) sulphate solution, 0.1 M silver nitrate solution,
Additional Materials
Students should bring with 2 pieces of graph paper, preferable 2 mm per division.
100 cm3 expanded polystyrene cup, 0-50 oC thermometer (0.2 oC readable), clean paper for
weighing zinc dust, access to balance (readable 0.01 g), stopwatch or stopclock,
Procedure
Record your results immediately after you get the readings.
1. Pipette 50.0 cm3 0.1 M copper(II) sulphate solution into a 100 cm3 expanded polystyrene cup.
Record the temperature of the solution using a 0-50 oC thermometer.
2. Weigh 0.5 g zinc dust with accuracy 0.01 g, using a clean paper. Add the zinc dust to the
copper(II) sulphate solution, at a time. Then, start the stopclock immediately.
3.
4.
5.
Stir regularly with the thermometer. Record the temperature at 15-second intervals. [Up to 30
intervals.]
Re-weigh the paper such that you get enough weighing data to calculate the mass of zinc used.
Repeat the experiment using 0.1 M silver(I) nitrate solution.
Further notes
1. The timing interval can be spaced out more after the maximum temperature has been reached,
but the temperature should be taken for at least 5 minutes after the maximum value.
2. In the above time range, the temperature should not be cooled by more than 0.2 oC.
3. If cooling is greater than 0.2 oC, then a graph can be used to make a cooling correction.
Data and results
En01_H_reaction /p.1
A. With copper(II) sulphate
Mass of zinc + paper/g
Mass of paper/g
Mass of zinc/g
Interval
T/oC
0 1
2
3
4
5
6
7
8
9
10 11 12 13 14 15 16 17 18 19 20
Interval 21 22 23 24 25 26 27 28 29 30
T/oC
B. With silver nitrate
Mass of zinc + paper/g
Mass of paper/g
Mass of zinc/g
Interval
T/oC
0 1
2
3
4
5
6
7
8
9
10 11 12 13 14 15 16 17 18 19 20
Interval 21 22 23 24 25 26 27 28 29 30
T/oC
Plot a graph of temperature against time for each titration. Deduce the temperature rise and give
this value on the graph paper.
Calculation:
Calculate enthalpy of reaction for the reactions.
Discussion (With the help of the following questions):
1. Give error sources for this practical and estimate percentage error for each of the enthalpy of
2.
3.
4.
reaction.
If part of your skin touches silver nitrate solution, this part turns black after several hours (It
will fade to original colour after several days). Explain.
Discuss the change of internal energy for the reactions in this practical.
Enthalpy (H) of these reactions may be approximately equal to Gibb's free energy (G) of
them. If the approximation is good, the cell potentials (E) of the above redox reactions (G)
can be estimated (by calculation) by using the following equation: G = - nFE. Estimate
electrode potentials for the above redox reactions. Then look up/ calculate these cell potentials
from standard electrode potentials and comment your results. [Note: n is the number of
electrons transferred according to the chemical equation. F is Faraday constant which is equal
to 96500 coulumbs.]
En01_H_reaction /p.2