Determining the Heat of Formation of Magnesium
Oxide
Purpose: In this experiment you are going to use calorimetry to determine the enthalpy
of reaction of several compounds of magnesium. Then, you will combine this data via
Hess’s Law to determine the Heat of Formation of magnesium oxide. In addition, you
will estimate the Kc for the solubility of MgO, and calculate the Free Energy Change for
the formation of MgO from its elements.
Introduction:
As we have seen in class, understanding the energy change associated with a chemical
reaction can provide vital information about the relative stability of the reactants and
products. However, as we have also discovered, actually determining the change in
internal energy, ΔE, is very difficult if you are not using a bomb calorimeter, due to the
work done on the atmosphere any gases consumed or formed. Thus, we measure the
related quantity enthalpy, ΔH, which is just the heat gained or lost at constant pressure as
the reaction occurs.
To measure delta H, we need a “probe” that will quantitatively absorb or release heat to
the reaction as needed. For us, that probe is a fixed volume of solution whose change in
heat can be easily measured via q=mcΔT. Unfortunately, we must also account for any
heat gained or lost to the calorimeter itself and add this missing heat back into ΔH.
In this experiment we wish to measure the heat of formation of magnesium oxide. In the
space below, write down the chemical reaction whose ΔH is the ΔHof of MgO(s):
Heat of formation reaction for MgO(s)
Clearly, this reaction would be difficult to carry out in our lab because of the need for a
bomb. However, judicious use of Hess’s Law will allow us to calculate the delta H for
this reaction by combining the heats of other reactions that are more easily measured in
the lab. For example, consider the reactions of hydrochloric acid with magnesium metal
and magnesium oxide respectively.
(1)
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
ΔHo1 = ?
(2)
MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(l) ΔHo2 = ?
The ΔH0 of each of these reactions should be very easy to measure using standard
calorimetry techniques. However, we need a little more information in order to pull the
ΔHof MgO (s) from this data. From your text, the ΔHof of water is:
(3)
H2(g) + 1/2O2(g)  H2O(l)
ΔHof = -285.840 kJ
In the space below, show how you can use these three reactions to obtain the heat of
formation of magnesium oxide:
How do you get ΔHof of MgO?
We will compare this value to the literature value with a percent error. In addition, we
will consider how to the ΔEof for the reation as well. Thus your goals for the experiment
are as follows:
1) React HCl(aq) with Mg(s) and MgO(s) to determine the ΔHo for these reactions,
correcting for the heat lost to the calorimeter
2) Find the ΔHof of MgO and compare to the literature value
3) Using values from your text, calculate the Free Energy change, ΔGo, for he
formation of MgO
4) Estimate the Kc for the solubility of MgO(s)  Mg2+(aq) + O2-(aq)
Procedure:
Part A: Estimating the Kc of MgO(s) solubility.
Into a small cup, place a scoopful of MgO. Stir and ensure the solution is saturated by
making sure there is some solid in the bottom of the container. At the end of the lab,
decant some of the supernatant liquid into a test tube and record the pH with a pH meter.
In addition, add a few drops of universal indicator to the solution to help confirm the pH
measured by the meter.
Part B: Determining the Heat of Reaction between Mg, MgO and HCl(aq)
Rinse your calorimeter clean with DI water and dry with a Kim-Wipe. Measure out 100
ml of 1.0 M HCl into the dry calorimeter and reweigh the system after addition of the
HCl. Add a magnetic stirrer to the system and place the temperature probe into the
solution. Weigh out 0.50 g of Mg turnings and record the mass of the metal. Start
Logger Pro and when the HCl is equilibrated, add the Mg to the solution. Stop data
collection when the system has equilibrated to the new temperature.
Repeat the procedure with MgO, except use approximately 1.00 g of MgO in place of the
0.50 g of Mg. If time permits rerun both of these experiments.
CALCULATIONS:
1) You have two chemical reactions to find the heats of reaction for
a) HCl(aq) + Mg(s)
b) HCl(aq) + MgO(s)
As discussed in class find the ΔH0 of each reaction per mole of the reaction (taking
into account the limiting reactant, reporting your final value in kJ/mol
2) Using Hess’ Law and the standard heat of formation of water (see the beginning
of this lab if you have forgotten) clearly calculate the standard heat of formation
of MgO(s)
3) Use Appendix L in your text (table of heats of formation) to determine the actual
values for the ΔH of HCl(aq) + Mg(s) and HCl(aq) + MgO(s) and compare to
your values with a percent error calculation. Also compare your calculation of the
heat of formation of MgO to the actual value with a percent error calculation.
4) From your text, calculate the ΔS0 for the formation of MgO from its elements.
Using this information, calculate the ΔG0 for the reaction.
5) You will have to think about why an oxide anion generates hydroxide in solution.
The pH will allow you to calculate the [OH-] in the solution. From this, calculate
the [O2-] that originally dissolved in the solution. Since MgO(s) will dissociate as
MgO(s)  Mg2+(aq) + O2-(aq), the concentration of oxide will give you the
concentration of the magnesium cation, and hence the Kc for this reaction. Look
up the literature value for the Kc (also known as Ksp) and compare with a percent
error calculation.
DISCUSSION
You want to address several points in your discussion section. Answer the following
questions in well constructed paragraphs, remembering the person reading your lab does
not know you had this guide.
1) It is very likely that your heat of formation of MgO is substantially different (810%) from the literature value. Do you attribute this to the reaction of Mg with
HCl or the reaction of MgO with HCl
2) Our MgO that we used is probably more than 25 years old. Interestingly, MgO
will react with CO2 in the air to make magnesium carbonate. If a substantial
amount of what we thought was MgO had converted to MgCO3
a) What physical evidence would you have seen during the chemical reaction
b) Would the fact that you may have had a mixture of MgO and MgCO3
account for the fact that your ΔH of the reaction was higher or lower than
expected?
3) Another thing we neglected was the fact that the calorimeter stole energy intended
for the water. Let us assume that the heat capacity of the calorimeter was 13 J/K.
4)
5)
6)
7)
If this were the case, was not accounting for this “lost heat” enough to explain
your percent error?
Your calorimeter measures delta H, and not delta E (if you are still confused
about the difference, look at Chapt 6 in your text). Based on your reactions,
would you expect the Delta E for each of the reactions measured to be higher or
lower than ΔH? Explain your reasoning.
How would your data and results have changed if you had used twice as much 1M
HCl in each trial (approx. 200 ml to the 100 you used)? Would this have been a
good idea or something that was not needed? Explain.
Using your calculation for ΔG0, does there exist a temperature (or range of
temperatures) for which the formation of MgO(s) is nonspontaneous? Explain.
Explain why the oxide anion creates hydroxide anion when dissolved in solution.
Would you expect the oxide anion to be a good base? Explain.
Finish the lab with a conclusion stated your value for the heat of formation of MgO
with your percent error.