David Gilmour
Date _______
Chem 12U
Identifying two unknown salts using calorimetry
to identify molar enthalpies of solution
Introduction
Enthalpy is a measure of the total energy of a system. As total enthalpy is difficult to
measure, changes in enthalpy are more commonly determined. Enthalpy changes result from the
combination of breaking of reactant bonds (requiring energy) and formation of product bonds
(releasing energy). Enthalpies of solution are a specific type of enthalpy change in which the
change in energy is due to the ionization of a salt when it dissolves in water. A molar enthalpy of
solution is the energy change associated with the solution of one mole of a salt. Calorimetry is
the experimental determination of the energy change associated with a chemical change. In
general, calorimetry involves the measurement of heat gained or lost by the surroundings of a
chemical reaction (system). This heat change is measured by a change in temperature of the
surroundings. If the specific heat capacity of the surroundings is known, then the quantity of
energy lost or gained by the system can be calculated. The purpose of this investigation is to
identify two unknown salts using calorimetry to identify molar enthalpies of solution. A known
salt was also tested to identify the experimental error of the technique.
Materials and Methods
Please refer to p. 247 of the Nelson Chemistry 12 text for materials and methods.
Results
Table 1. Mass used and change in temperature of the calorimeter for two unknown salt solutions
and of KBr.
Salt A #1
#2
#3
Avg
Mass of salt (g)
4.491
4.480
5.883
4.951
∆T (ºC)
-5.5
-5.0
-8.0
-6.2
Salt B #1
#2
#3
Avg
4.412
5.316
5.404
5.044
-5.7
-5.0
-7.0
-5.9
KBr
4.622
4.421
4.218
4.420
-3.7
-3.2
-3.0
-3.3
Trial #
#1
#2
#3
Avg
Calculations of molar enthalpies (not included here)
Results Summary
In this experiment it was found that the enthalpies for all three salts were positive. The
enthalpies per gram of salts A and B were about two times that of the know salt KBr. After
making assumptions about the identities of salt A and B, it was found that Salt B and KBr had
similar molar enthalpies (18 kJ/mol) and were about 1.3 times greater the molar enthalpy value
of Salt A (14 kJ/mol). The percent error for KBr (10%) was twice that of those for Salt A and B
(5%).
Discussion
In this experiment, the enthalpies of solution of all three salts were found to be positive
indicating that their reactions were endothermic. In order for an endothermic reaction to occur
the energy required to break the reactant bonds must be greater than the energy released to form
the product bonds. The dissolving process involves two main steps. First the ionic bonds between
the ions must be broken and then the water forms attractions around each ion, known as the
solvation shell. The endothermic nature of all three salts suggests that the breaking of the ionic
bonds during the dissolving process required more energy than the energy that was released
when the ions formed attractions with the water. Both KBr and salt B had the greatest molar
enthalpies releasing 18 kJ of energy for every mole dissolved. This suggests that more energy is
required for the ionization of KBr and Salt B in comparison to salt A. The enthalpy per gram of
salt A, however, was found to be larger than that of salt B as the molar masses are different.
Since salt B is a heavier molecule, there are fewer of them in a gram than that of salt A which is
a smaller molecule. Comparing enthalpies per mole is considers their energy change based on the
same number of molecules.
Using their molar enthalpies of solution calculated from calorimetry data, the identities of
two unknown salts were determined. The enthalpy per gram of Salt A (0.26 kJ/g) was closest to
that of NH4Cl (0.277 kJ/g). In the case of Salt B, the enthalpy per gram (0.24 kJ/g) was closest to
that of KCl (0.231 kJ/g). The experimental error of this technique, as determined from the results
of the known salt KBr, was 10%. In all three KBr trials, the experimental enthalpy was less than
the accepted enthalpy per gram. Considering this, the experimental enthalpies of the two
unknowns would also be less than the accepted values. As the precision of the experiment
between the salts is unknown, there is a certain degree of uncertainty with the identities of the
two unknown salts. The causes of the errors with these results are due to the systematic errors
associated with the technique. The lower than expected enthalpy for the known is most likely
explained by the contribution of heat to the system reaction from the air and cup as they were not
considered. This contribution of heat which was not accounted for would decrease the change in
temperature. This factor could account for a lower than expected molar enthalpies. A calorimeter
with more insulation would help reduce the error in this lab.