THE HEAT OF FORMATION OF MAGNESIUM OXIDE
Chem 1411 Experiment
PRE-LAB CHECKLIST
1. ______ Read the experiment carefully.
2. ______ Study the techniques you will use in lab. Review videos if necessary. Note: There are graphs
in the experiment folder that you will need to use for your data analysis and calculations. You will use
these during the experiment, but you do not need to print these out before the lab. You will input your
data and print the graphs out during lab.
3. _____ _Prepare your Pre-Lab:
a. Pre-Lab Questions
b. Beginning Question
c. Safety summary
d. Table of Constants (this should include a value for the Heat of Formation of MgO
from its elements)
e. Balanced equations for reactions
f. Procedure
g. Data Table
Note: There are graphs in the experiment folder that you will need to use for your data
analysis and calculations. You will use these during the experiment, but you do not need
to print these out before the lab. You will input your data and print the graphs out during
lab.
4. ______ Be able to do the calculations required in the experiment before lab. The Pre-Lab Questions
will help with this so pay special attention to them.
IF YOU DO NOT UNDERSTAND SOMETHING DO NOT WAIT UNTIL THE DAY OF LAB TO SEEK HELP!
FREE TUTORING IS AVAILABLE TO YOU.
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THE HEAT OF FORMATION OF MAGNESIUM OXIDE
Chem 1411 Experiment
INTRODUCTION
Calorimetry involves the study of heat characteristics of substances (heat capacity) and chemical reactions (heat
of reaction). The heat liberated or absorbed is measured by changes in temperature in an apparatus called a
Calorimeter. The most common type of calorimetry experiment done in general chemistry labs is coffee cup
calorimetry which uses nested styrofoam cups as the calorimeter and is run under prevailing atmospheric
conditions making it a constant pressure run.
For reactions involving aqueous solutions, this simple calorimeter works well. The reaction is carried out in the
cup and the temperature change, (T), is measured.
NOTE: The Greek letter delta, , is used in chemistry to indicate “change” in a variable. For example: T =
change in temperature, E = change in energy, and so forth.
Assuming no heat loss, the reaction heat all goes into warming the solution and cup. This heat quantity can be
calculated as follows:
-qrxn = qsoln + qcal
where qrxn is the heat of reaction, qsoln is the heat absorbed by the solution, and qcal is the heat absorbed by the
calorimeter.
It is importatn to recognize that calorimetry is a study of heat exchanges. If an AMOUNT of heat is lost by the
system, the same AMOUNT of heat must be gained by the surroundings. For an exothermic reaction where heat
is released from the system to the surroundings:
-qlost = qgained
Notice the amount of heat has the same numerical value but the sign changes.
The heat absorbed by the solution is calculated as follows:
qsoln = m x C x T
where q is heat (J), m is mass (g), C is specific heat of the solution (J/g∙ oC), and T = Tfinal  Tinitial (oC). For
this experiment, 1M HCl is being used for which C = 3.86 J/g∙oC.
The heat absorbed by the calorimeter is calculated as follows:
qcal = Ccal x T
where Ccal is heat capacity of cup (J/oC), which is determined experimentally, and T = Tfinal  Tinitial (oC). The
heat capacity of the calorimeter has been determined experimentally to be 7.27 J/oC; that value will be used for
these experiments.
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The heat of reaction, (H), is defined as the heat lost or gained as a reaction proceeds from reactants to
products. It is often given as part of a thermochemical equation, such as that shown for the combustion of
hydrogen:
2 H2(g) + O2(g)  2 H2O(l)
H = 572 kJ
The negative sign indicates that this reaction is exothermic. Thus, when two moles of hydrogen gas and one
mole of oxygen gas combine to form two moles of liquid water, 572 kJ of heat energy is released. Conditions
of constant atmospheric pressure (reaction vessel is an open container) and near room temperature (25 oC) are
assumed in the above value, although heats of reaction do not vary appreciably with temperature.
Aside from accurate data collection and interpretation, the key to this experiment lies in knowing HOW each
reaction with a measured value for ∆H is related to determining the ∆H for the reaction of interest; the
formation of magnesium oxide from its elements. This will be determined by using Hess’ Law. The following
example shows how Hess’ Law can be used to find the ∆H for an overall reaction when you know the heats of
reaction for a number of component reactions.
In this experiment, you will measure the H values for the following reactions:
Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g)
MgO(s) + 2 HCl(aq)  MgCl2(aq) + H2O(l)
Using the thermochemical equations for these two reactions, the hydrogen + oxygen reaction,
2 H2(g) + O2(g)  2 H2O(l)
H = 572 kJ
and their associated values of ∆H, you will then calculate the H value for the heat of formation of MgO using
Hess’ Law.
Mg(s) + ½ O2(g)  MgO(s)
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The final heat of reaction (H) calculation requires a stoichiometric conversion from grams of limiting reactant
to moles as expressed in the balanced equation. In this experiment, the Mg and the MgO are the limiting
reactants in each of the reactions.
FOR FURTHER UNDERSTANDING THE DETAILS…
This type of simple calorimeter works exceptionally well for reactions in solution that do not generate gases.
What about reactions that do generate gases?
Symbols E = Internal energy, P = pressure, q = heat exchanged, qp = heat exchanged at constant pressure,
qv = heat exchanged at constant volume, V = Volume, w = work
Definition: Change in internal energy of the system = heat exchanged + work done on or by the system.
This statement can be represented mathematically by: E = q + w, where w = -PV
By substitution: E = q - PV and then rearranged: q = E + PV (eqn. 1)
Enthalpy (H) is defined as E + PV; thus, H = E + PV
At constant pressure and referring to (eqn.1), qp = H. At constant volume (V = 0), E = qv
We have two descriptions of energy changes in the system, E and H.
IF the volume changes (as in the generation of gases), the PV term is no longer “0”; work is done, so
E is no longer equal to H by an amount PV
So how does this affect our coffee cup calorimetry experiment?
In reactions that generate gases, there is a volume change under constant pressure that adds to the energy
change as it does work through the PV term. Even though the PV term affects the value of the heat of
reaction (H), from a practical standpoint the effect is very small making coffee cup calorimetry results
reasonably good. Knowing that there is a small amount of error in H values determined with coffee cup
calorimetry does not invalidate the use of the technique. More precise data can be obtained with a different
calorimeter. For what we are measuring in this experiment, the simple calorimeter is acceptable.
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PROCEDURE - OVERVIEW
You will be assigned a reaction: either Mg or MgO which will be added to the 1M HCl. You will perform TWO
runs of one of the assigned reactions using the simple coffee cup calorimeter and a digital thermometer (set to a
precision of 0.01oC) to make the temperature measurements.
An Excel spreadsheet program will be used to graph the temperature vs. time data for the two reactions. From
the results of these plots, T values will be determined and the calculations can be completed.
You will post your two values for the ∆Hrxn of your assigned reaction on the projected instructor spreadsheet (in
kJ/mol). Record all of the ∆Hrxn data determined by your classmates. Use this data to calculate an average
∆Hrxn for each reaction and then calculate the Heat of Formation of MgO.
EXPERIMENTAL PROCEDURE
CONSTRUCTING THE CALORIMETER:
1.
Obtain the components for the calorimeter (two styrofoam cups [one nested in the other for extra
insulation] and a magnetic stir bar), a magnetic stirrer, a digital thermometer, and a stop watch.
NOTE: You might expect a lid to be used with the calorimeter, but in this experiment it is important for
you to see what is going on in solution and heat losses due to the absence of a lid are minimal.
2.
Place the magnetic stir bar in the calorimeter cup. Weigh the empty nested cups containing the stir bar
(DO NOT weigh the 400mL support beaker). This will be the mass of your empty calorimeter.
TECHNIQUE NOTE: Remember to use the same balance for ALL weighings but especially for those
readings that will be used to determine the mass by difference to reduce the error inherent in the
measurement due to minor differences in calibration between instruments.
3.
Obtain 60 mL of 1M HCl in a graduated cylinder. Record its molarity. Pour the HCl into the
calorimeter. Reweigh the calorimeter containing the HCl. Determine the mass of the HCl by difference.
SAFETY NOTE: 1M HCl is a very dilute solution of hydrochloric acid; however, it is still an acid and
it is still corrosive to skin. Use good laboratory practices in handling this solution.
4.
Place the cups in the supporting beaker and position the calorimeter on the magnetic stirrer provided at
your workstation. Adjust the stirring rate to a fast rate but not vortexing too vigorously and certainly
without splashing. Clamp the thermometer probe securely with a 3-prong clamp. Position the
thermometer probe in the solution such that the probe rests just above the level of the stir bar and does
not touch the bottom or sides of the cup.
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5.
Make sure that the digital thermometer is set to read values at a precision of 0.01 oC. With the stir bar
spinning, allow about five minutes for temperature equilibration.
DATA GATHERING FOR THE REACTION:
6.
Make sure that the balance is tared (reads zero with nothing on the weighing pan). Place the weighing
paper on the pan; note its mass and then add approximately 0.6 grams (NO MORE THAN 0.7 g) of
magnesium turnings OR 1.0 g magnesium oxide (NO MORE THAN 1.2 g) (whichever one you have
been assigned). Remember that Mg or MgO MUST be the limiting reactant.
EXAMPLE: If your weighing paper has a value of 0.247 g, adding approximately 0.6 g of Mg would
make your target mass 0.847 g.
TECHNIQUE NOTE: DO NOT try to weigh out exactly 0.847 g; 50 mg over is probably okay…more
sample will give a greater T so it’s better to be too heavy than too light as long as the Mg OR MgO is
still the limiting reagent. If the Mg or MgO is NOT the limiting reactant, your results are INVALID
and must be repeated with the correct amount of Mg or MgO to be limiting.
7.
Record the exact amount of the mass of the solid plus the weighing paper (to 0.001 g).
TECHNIQUE NOTE: You will reweigh this weighing paper at a later time to determine the mass of
solid actually added to the calorimeter. This mass of solid delivered to the calorimeter will be
determined by difference.
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8.
Take 5 temperature readings 20-seconds apart before beginning the experiment. If the temperature is
varying widely (>0.20oC), wait a few more minutes for thermal equilibrium before continuing. Enter the
last temperature as “0 seconds”.
9.
Quickly add the solid to the vortex of the calorimeter. Be ready to continue QUICKLY to the next
step.
SAFETY NOTE: If your reaction uses Mg metal, you will be generating hydrogen gas. The bubbles
may aerosolize small amounts of HCl. DO NOT stand over the calorimeter as the hydrogen is generated
since it may irritate your nose, lungs, and eyes.
TECHNIQUE NOTE: The importance of adding the sample to the center of the vortex is that the solid
MUST be in the solution to get accurate results for the mass added. Be sure to look into the calorimeter
to make sure that the entire solid is in solution
TECHNIQUE NOTE: MgO is a very light and clumpy powder. If you incline the weighing paper too
steeply so that the sample falls as a clump into the HCl, some of the powder will “puff” out of the
calorimeter and be lost incurring error in the determination. Pour it less steeply at an even rate but as
quickly as possible into the vortex to diminish the losses due to the airborne particles.
10.
Record the time and temperature every 20 seconds. (Set the weighing paper safely to the side and out of
the way so that it can be reweighed to determine the mass of reactant actually added to the calorimeter.)
Getting the first 20-second reading is one of the most important readings taken! DON’T MISS
IT!!!
TECHNIQUE NOTE: It is more important to get the solid in solution and start taking readings than it
is to make sure that the entire solid on the weighing paper is transferred to the calorimeter. The
weighing paper will be reweighed with any sample remaining on it after your run is completed. The
difference in that mass and the initial mass of the weighing paper will determine exactly how much
sample (weight by difference) went into the reaction in the calorimeter.
TECHNIQUE NOTE: The Mg will be delivered very cleanly to the solution probably leaving none
behind on the weighing paper. Some of the MgO WILL adhere to the weighing paper. That’s okay.
You will be determining “how much” made it into the solution when you reweigh the weighing paper.
11.
As the reaction occurs, you should observe the change in temperature carefully continuing to take data
defining the maximum temperature followed by a well-established, defined slope as the completed
reaction begins to reach equilibrium with the surroundings. Continue taking readings for the entire
600 seconds.
12.
TECHNIQUE NOTE: More readings are better than too few! It is the extrapolation of the data to time
zero that defines Tfinal. The better the definition of the slope after the reaction is complete the more
accurate the value of Tfinal will be.
13.
Once your temperature readings are completed, reweigh the weighing paper that you earlier set aside.
Record this mass to the nearest 0.001 g. Determine the amount of solid added to the calorimeter by
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subtracting this mass from the mass from Step 7. Dispose of the used weighing paper after you have
recorded the value.
14.
The solution in the calorimeter can go down the drain. Rinse the cups well; dry the cups individually
with paper towel. Repeat the experiment.
15.
Complete the calculations for Steps 1 through 7 given in the Data Analysis section. Post the Hrxn value
for each of your runs on the spreadsheet on the computer at the instructor’s desk. As these are involved
calculations that you need to do in lab instead of at home, you may print out and bring the Data Analysis
page to lab with you.
CALCULATIONS AND DATA ANALYSIS
1.
Graph your data using the excel spread sheet – temperature vs. time to calculate T for your reaction.
You will need to extrapolate the slope to time zero as shown in the graph below to determine Tfinal.
2.
The heat capacity of the calorimeter (the calorimeter constant) has been determined for you. Using this
heat capacity for the calorimeter, determine the amount of heat absorbed by the calorimeter (qcal) for
your reaction. Ccal = 7.27 J/oC.
3.
Using the mass of the HCl solution (including the amount of solid that you added) calculate the amount
of heat absorbed by the solution (qsoln) for your reaction. Csoln = 3.86 J/g∙oC
4.
Calculate the total amount of heat (qrxn) for your reaction.
5.
Determine the number of moles of the limiting reagent added to the calorimeter (Mg or MgO).
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6.
Calculate the heat of reaction per mole (H) for your reaction. (Remember the relationship between heat
lost and heat gained.)
7.
Repeat the calculation steps for the second run of your reaction and post your results with the rest of the
class data on the board.
8.
Calculate the average heat of reaction for the Mg reaction and the MgO reaction from the class data.
This may be provided on the in-class spreadsheet where you posted your Hrxn value. Use these average
values in further calculations.
9.
Use Hess' Law to calculate H for the heat of formation for MgO.
10.
Using your calculated value of the heat of formation of MgO and the literature value for the heat of
formation of MgO that you obtained from a reliable source, determine the experimental error.
PRE-LAB QUESTIONS
1. In determining the Heat of Formation of MgO, you will be using Hess’Law. Given the following three
reactions that you will be using in performing the lab:
Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g)
(Eqn.1)
MgO(s) + 2 HCl(aq)  MgCl2(aq) + H2O(l)
(Eqn.2)
2 H2(g) + O2(g)  2 H2O(l)
(Eqn.3)
a. Using Hess’ Law, rearrange the equations as necessary so that the summation of the reactions is
Eqn.4.
(Understanding the “Example of Hess’ Law on page 3 will help you with this. HINT:
You will simply rearrange the equations. DO NOT try to calculate the heat of reaction in
this prelab, that part will be done when you generate your data during the lab
experiment.)
Mg(s) + ½ O2(g)  MgO(s)
(Eqn.4)
b. In lab you will be determining the values of ∆H for the Mg and MgO reactions (Eqn.1, Eqn.2).
IF one of those values of ∆H is negative and that reaction is reversed when applying Hess’ Law,
what sign will the value of ∆H have for that reaction?
c. What is the meaning of a negative enthalpy value?
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2. You have gathered the following data for the reaction of Magnesium with Hydrochloric Acid.
EQUILIBRATION
Time (s)
Temp
(oC)
-80
21.09
-60
21.10
-40
21.09
-20
21.08
0
21.08
Time
(s)
0
20
40
60
80
100
120
140
Temp
(oC)
21.08
33.86
42.68
48.08
51.77
54.48
56.50
57.84
Time
(s)
160
180
200
220
240
260
280
300
Mg + HCl SAMPLE RUN
Temp
Time Temp(oC)
(oC)
(s)
58.87
320
58.98
59.47
340
58.18
59.82
360
57.66
59.97
380
57.18
60.03
400
56.76
59.92
420
56.26
59.68
440
55.82
59.36
460
55.32
Using the excel workbook, you have generated the graph below.
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Time(s) Temp(oC)
480
500
520
540
560
580
600
54.88
54.38
53.97
53.45
52.90
52.55
51.80
Given: Mass of 1 M HCl = 61.087 g
Mass of Mg = 0.613 g Specific Heat of 1 M HCl = 3.86 J/g∙oC
Heat Capacity of the Calorimeter: 7.27 J/oC
Calculate:
a. Tf
b. ∆T
c. qsoln
d. qcalorimeter
d. Hrxn
POST-LAB QUESTIONS
1.
2. Identify and explain some fo the potential reasons that your heat of formation value is different from the
published value.
3. Batteries are used to store energy. A 1.5 V battery provides energy equivalent to 144.7 kJ/mole. The
table below provides some information about 1.5 V batteries. Discuss some benefits and drawbacks of
using the reaction you performed in place of a 1.5 V battery based on the information provided.
Battery Type Total Energy Content (J) Battery Volume (mL) Battery Mass (g)
AAA
5071
3.9
12
AA
9360
8.3
24
C
34398
26.5
65
D
74970
55.8
135
Data taken from http://www.allaboutbatteries.com/Energy-tables.html
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