Enthalpy of Formation of MgO
Revised 3/3/15
ENTHALPY OF FORMATION OF MgO
ELECTRONIC LABORATORY NOTEBOOK (ELN) INSTRUCTIONS
All work for this experiment must be recorded, attached, or answered in the ELN. Create a pre
& inlab page in the Experiment #1’s folder containing the following sections:
1) Equipment and chemical tables. Pictures of typical equipment, instrumentation, special
reaction set-ups should be attached. Chemicals with constants needed in the experiment and
NFPA Fire Diamonds should be listed in the table. (A special Widget is available for making
the Fire Diamonds.) In order to have a complete table, you must look ahead at the
calculations and think about constants that will be needed.
2) Procedures and Observations. PROCEDURAL STEPS MUST BE IN YOUR OWN
WORDS – DO NOT COPY THE LAB MANUAL. All quantitative data should be recorded
with the correct number of significant figures and units. All qualitative observations should
indicate the color, depth of color, and either the crystallinity of a solid or the clarity of a
liquid. Any cues that indicate a reaction has occurred (such as heat evolution, color change,
gas evolution, et cetera) should be noted. It is acceptable to write the procedures and
observations together in one section as you perform the experiment. All Logger/Pro data
files created should be recorded or attached to this ELN page. You should include one or
more pictures of assembled glassware set ups and/or equipment used in lab.
Postlab questions at the end of this document should be answered on a new page in Experiment
#1’s folder.
INTRODUCTION
Enthalpy (ΔH) is the heat change in a reaction carried out at constant pressure. According to the
First Law of Thermodynamics, the enthalpy of a reaction is the difference between the
enthalpy of the products and the enthalpy of the reactants. Since enthalpy is a state function, the
enthalpy change is the same whether the reaction takes place in one step or in a series of steps.
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Enthalpy of Formation of MgO
Revised 3/3/15
Therefore, "enthalpies of reaction are additive in the same way that the reactions to which they
pertain are additive" (Hess' Law).
The enthalpy of formation (ΔH°f) is defined as the heat change associated with the formation of
one mole of a compound from its elements in their standard states. Many enthalpies of
formation are often difficult to measure directly in the laboratory because the reactions may
evolve too much heat or form toxic substances. However, they can be calculated indirectly by
measuring the enthalpy changes in a series of related reactions and applying Hess' Law.
For the formation reaction for MgO(s) written below (1), the heat of the reaction is equal to the
enthalpy of formation.
(1)
Mg(s) + ½ O2(g) à MgO(s)
ΔH°rxn = ΔH°f
Determining ΔH°f for MgO(s) is difficult because the combustion of Mg metal results in a large
evolution of heat and toxic fumes. However, Hess’ Law states that ΔH°f can be calculated by the
summation of the enthalpy changes for the reactions of Mg and MgO with dilute acid solutions
and the enthalpy of formation of liquid water.
(2)
MgO(s) + 2 H+(aq) à Mg2+(aq) + H2O(l)
ΔH°2 = ?
(3)
Mg(s) + 2 H+(aq) à Mg2+(aq) + H2(g)
ΔH°3 = ?
(4)
H2(g) + ½ O2(g) à H2O(l)
ΔH°4 = –286 kJ/mol
The release or absorption of heat by aqueous reactions can be measured in constant-pressure
(“coffee-cup”) calorimeters. The heat absorbed or released by the system (reaction) is
determined by measuring the temperature change (ΔT) of the surroundings (solution and
calorimeter):
(5)
qrxn = −(qsoln + qcalorimeter)
(6)
qrxn = −([m Cp ΔT]soln + [m Cp ΔT]calorimeter)
where m = mass (in grams)
Cp = heat capacity (in J / g°C) at constant pressure
ΔT = Tfinal – Tinitial (°C)
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Enthalpy of Formation of MgO
Revised 3/3/15
In this experiment the heat gained by the calorimeter (qcalorimeter) is assumed to be negligible.
Therefore, the heat evolved by the reaction can be calculated from the temperature change, mass,
and heat capacity of the solution alone:
(7)
qrxn = −(mCpΔT)soln
The relationship between enthalpy and q is given below:
(8)
ΔHrxn = qrxn = −(mCpΔT)soln
In an exothermic reaction heat is released to the surroundings and Tfinal > Tinitial, ΔHrxn will be
negative. In an endothermic reaction, heat is absorbed from the surroundings and Tfinal < Tinitial,
ΔHrxn will be positive.
If the limiting reactant is entirely consumed in the reaction, the heat change will be dependent on
the amount of limiting reactant initially present. Therefore, qrxn can be converted to the standard
molar enthalpy, ΔH°rxn (kJ/mol), using moles of limiting reactant:
(9)
ΔH°rxn = qrxn / moles of limiting reactant
SAFETY PRECAUTIONS
Safety goggles and aprons must be worn at all times. Wear gloves when handling the HCl stock
solution - acids are corrosive and can cause burns; wash all contaminated areas thoroughly.
Hydrogen gas (H2) is released when Mg powder reacts with water; make sure no open flames or
spark sources are present in the lab.
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Enthalpy of Formation of MgO
Revised 3/3/15
PROCEDURES
Before starting the experiment, the TA will randomly ask students to do a quick demonstration or
talk-through of one of the following:
1) How to use a volumetric flask
2) How to use a volumetric pipette
3) How to make an acid solution (what order do you add things?)
4) How to correctly neutralize an acidic solution
Read the technique documents and watch the videos on the course website to prepare for these
demonstrations every week. Everyone will have presented at least one topic by the end of the
quarter. The demonstrations should be short (>1 min) and will be graded.
Part A: Solution Preparation
1. Work in pairs. Wear safety goggles and an apron. In your ELN: Click on Heading, type
Procedures and Observations. Click on Rich Text. In this section procedures with the
appropriate observations will be entered as the steps are performed in lab. (See the ELN
instructions for more detail)
2. Connect the temperature probe to Channel 1 of the LabQuest 2. The vertical axis has
temperature scaled from 0 to 50°C. To set the horizontal axis (time) to 0 to 1000 seconds, go
to the
menu and change the duration to 1000 seconds.
3. Wear gloves for this step only, 2.5 M acid can cause burns. When diluting concentrated
acid, pour acid into water, slowly with stirring; never the reverse. HAVE A DILUTION
PLAN BEFORE COMING TO LAB. (Use the concentration values and volume below in
the dilution equation to find the volume of stock solution needed to create the dilute
solution.) Using only volumetric glassware, prepare 250.00 mL of approximately 1.0 M HCl
solution from the stock HCl solution in the hood (approximately 2.5 M). The exact
concentration for the stock solution must be recorded from the bottle’s label and the exact
concentration for the diluted HCl solution must be calculated.
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Enthalpy of Formation of MgO
Revised 3/3/15
Part B: Heat of Reaction for MgO(s) + 2 H+(aq)
4. Place a Styrofoam cup into a 250-mL beaker. Place the beaker on a stirrer/hotplate. Using a
graduated cylinder, measure out 100.00-mL of the HCl solution created in the step above into
the Styrofoam cup. Suspend a temperature probe from the bars at the back of the lab bench
with a clamp. Place the temperature probe and stir bar in the acid solution in the Styrofoam
cup.
5. Weigh out ~1.0 g MgO. Record the mass to the nearest milligram. Caution: Avoid inhaling
MgO dust.
6. Begin stirring. Click
to begin data collection and obtain the initial temperature, T1.
After 15 seconds at the same temperature have been obtained, add MgO to the solution.
Continue to stir the cup contents until a maximum temperature, T2, is reached and the
temperature starts to drop. Click
to end data collection.
7. Examine the initial readings in the Table window (access by clicking on
) to determine the
initial temperature, T1 and final temperature, T2. E-mail this data to your lab notebook.
8. Pour the solution in a 1 L beaker to be neutralized later.
Part C: Heat of Reaction of Mg(s) + 2 H+(aq)
9. Caution: Do not breathe the vapors produced in the reaction. Repeat Steps 4-7 using
approximately 0.5 g of Mg metal instead of MgO. Record its mass to the nearest milligram.
Use the same digital scale as in Part B to limit calibration errors.
10. Add the Mg-acid solution to the 1 L beaker used in step 8.
Make sure to clear your email address and password of the LabQuest2 so others can’t access
your email account. Shutdown the LabQuest2 and not simply put it to sleep. To shutdown the
LabQuest2: press the home key, select System à Shut Down à OK.
Part D: Neutralizing the Acidic Solutions
11. Place about ¼ in. of NaHCO3(s) in the bottom of a clean 1 L beaker, add water until it barely
covers the NaHCO3, stir to make a slurry. Now slowly pour in the diluted acidic solution
from steps 8 & 10 and stir. (Quick pouring will result in uncontrollable foaming and a mess.)
When foaming stops, check the solution with a small piece (< 1 inch) of pH paper. If pH = 7,
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Enthalpy of Formation of MgO
Revised 3/3/15
the solution is neutral (pH = 7) it is OK to pour the neutralized solution down the sink. Na+,
Cl– and Mg2+ ions are normally present in tap water and are not harmful to the environment
in small amounts. If the pH of the solution is less than 7, prepare a new slurry and neutralize
the solution again in the same manner. Clean workspace and glassware thoroughly before
beginning calculations.
Before leaving lab, go to your TA and click on sign in the Page Tools toolbar near the top right
of the ELN window. This will lock the page so no further edits can be made. If you do not do
this, you will not receive a grade for the experiment.
CALCULATIONS
1. Calculate the change in temperature, ΔT, for the reactions in Part B & C.
2. Calculate the heat change (q) in kJ for each reaction. (Assume the density of the acid
solutions is 1.00 g/mL and the heat capacity is 4.184 J/g-°C.
3. Determine the moles of limiting reactant for each reaction.
4. Calculate molar enthalpy, ΔHrxn (in kJ/mol), for each reaction.
5. Determine the heat of formation, ΔHf , for MgO.
ERROR ANALYSIS
1. Compare your experimentally obtained value for ΔHf (MgO) to the true value in your
textbook by calculating the percent error:
% error = (|true value – expt value| / true value) x 100
2. What modifications could be made to the procedure to better account for random
(indeterminate) errors?
3. List three potential systematic (instrumental, methodological, or personal) errors that could
be made in this experiment. (Note: Be specific, systematic errors are in the details. For
example, losing your solution because you knocked over the cuvette is not a systematic error
– it’s a gross one.)
4. Did any gross errors occur? Did you mess up? Did the equipment or instrumentation fail?
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