LAB 12:
ENTHALPY CHANGE OF A REACTION AND HESS'S LAW
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This lab does require a lab report.
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You will work with a partner for this lab.
I Introduction
A few minutes of unremarkable heating is followed by a slightly-more-interesting but
none-too-exciting orange glow. Despite a few fervent peeks under the crucible lid,
your reaction seems like just another boring chemistry lab. Suddenly, however, all
that changes. Suddenly a whiff of smoke appears and the shiny turnings begin to glow
a brilliant white, first in one small corner of the pile, and then throughout the entire
crucible. You stand in awe, gazing at the brilliant flash of blinding white light, letting
the crucible lid drop and shatter to the floor. Your mouth is still agape, long after the
glow has ceased. It’s as if you can still see the reaction before your eyes. Oh, wait –
that’s just the blind spot burned onto your retina that is slowly healing. Nevermind
that, you’ve just seen one of the most exciting labs you could possibly see in
introductory chemistry. You’ve just seen the combustion of magnesium.
The combustion of magnesium, as written below, is certainly an impressive sight and
is certainly an exothermic reaction, producing significant amounts of heat and light.
Mg(s) + ½ O2(g)  MgO(s)
H = ?
But how much energy does the reaction produce? That is the question you will answer
in this lab.
The heat of reaction for the combustion of magnesium is somewhat difficult to
measure, especially considering our crude equipment. After all, even if our Styrofoam
cup calorimeters could survive the intense heat of the bunsen burner needed to
initiate the reaction, they would surely succumb to the massive amounts of energy
released in the reaction itself. To work around this problem, we’re not going to do
the reaction at all. Instead, we’re going to do two similar reactions, then use Hess’s
Law to determine the heat of reaction of our combustion reaction. First, however,
let’s look at how we’ll measure the heat of each reaction.
A calorimeter is a device used to measure the heat of a reaction. For reactions
involving aqueous solutions, a simple Styrofoam cup works well. One carries out the
reaction in the cup and measures the temperature change (T). Assuming adiabatic
conditions (no heat loss), the reaction heat all goes into warming the solution and
cup. This heat quantity can be calculated as follows:
H = q = m∙c∙T + ccup∙T
As in our previous lab, the first term of the equation refers to the amount of heat
absorbed by the solution.
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m is the mass of the calorimeter’s contents: the initial solution plus any solids
you added.
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c is the specific heat of the solution. For this lab, assume a specific heat of
3.862 J/g∙oC (as compared to water’s specific heat of 4.18 J/g∙oC).
The second term in the above equation represents the amount of heat absorbed by
the Styrofoam cup calorimeter. For this lab, we will assume a heat capacity of our
cups of ccup=30 J/oC. Thus, for every one-degree increase in the temperature of the
cup, the cup absorbs 30 joules of energy.
In Part 1 of this lab, you and your partner will make measurements to determine the
heats of the following reactions:
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
H1 = ?
MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(l)
H2 = ?
In Part 2 of this lab, you will then use these results to determine the change in
enthalpy for the combustion of magnesium. You will then compare this result to the
value found by using heats of formation and calculate your percent difference.
Equipment and Reagents
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Styrofoam® cup
Beaker
Graduated cylinder (?)
Magnetic stirrer and stir bar
Thermometer OR LabPro with
thermometer probe
Ringstand and thermometer
clamp
Analytical balance
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Distilled water
3 M HCl, Hydrochloric acid
Mg turnings
MgO(s), magnesium oxide
! Warnings!
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The hydrochloric acid is respectably concentrated (3 Molar). Wear goggles and
aprons and immediately clean up any spills with Acid Lock.
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Magnesium is flammable and relatively reactive. Do not expose magnesium
turnings to heat. See Mr. Benter for disposal of excess solid magnesium.
Procedure
Determining the Heat of Reaction of Mg and HCl
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Prepare your lab notebook and don your safety apparel.
Weigh out approximately 0.25 g of magnesium turnings. Record their mass and
set them aside.
Add 50.0 mL of 3 M HCl to a Styrofoam cup calorimeter. Determine and record
the mass of the solution.
Proceed as in Lab 11 to determine the heat released by the reaction. Flush all
waste solutions down the drain with water.
Determining the Heat of Reaction of MgO and HCl
□ Repeat the above procedure, this time using about 1.30 g of magnesium oxide
and 100.0 mL of 3 M HCl.
□ Flush all waste solutions down the drain with water.
Data and Calculations (to be recorded on your lab Data Sheet)
1. Procedure and observations
2. Masses of all reagents used
3. Initial and Final temperatures
4. Heat, q, released both reactions, with supporting calculations
5. Enthalpy change, ΔH (in units of kJ/mol) of both reactions, with supporting
calculations
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Enthalpy change, ΔH (in units of kJ/mol) of combustion of Mg, with supporting
calculations (see next section)
? Questions to Answer
Write a lab report, including the following:
1. Using the enthalpy changes you determined for Reactions 1 and 2 below, as
well as the given enthalpy change for the formation of water (Reaction 3),
determine the heat of reaction for the combustion of magnesium (Reaction 4).
(1)
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
H1 = ?
(2)
MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(l)
H2 = ?
(3)
H2(g) + ½ O2(g)  H2O(l)
H3 = -286 kJ
---------------------------------------------------------------------------------------(4)
Mg(s) + ½ O2(g)  MgO(s)
H = ?
Look up the standard enthalpy of formation (ΔHfo) of magnesium oxide in the
appendix of your textbook. Compare this to your result (ΔH). Calculate a
percent difference for your determination.
2. Identify any sources of error that you recall happening or likely happened. Do
NOT include “bad measurements” (unless you distinctly remember this
happening or strongly believe it happened) or a “mistake in calculations”
(you’re too smart for this, plus you surely went through your work twice (at
least) to verify its correctness).
3. Do you expect that the enthalpy change (ΔH) determined using Hess’s Law,
even if done without error, should be identical to that determined from
standard enthalpies of formation? Why or why not?