THERMOCHEMISTRY III: HEAT OF COMBUSTION OF MAGNESIUM | 87
Thermochemistry III: Heat of Combustion of Magnesium
OBJECTIVES:
• Further explore quantitative aspects of constant-pressure calorimetry
• Use Hess’s law for indirect determination of reaction enthalpy
DISCUSSION:
In the Thermochemistry II experiment, you tested the validity of Hess’s law. In this experiment,
you will use this principle to make an indirect determination of the enthalpy of reaction (∆Hrxn)
for the combustion of magnesium ribbon:
Mg (s) + 1/2 O2 (g) → MgO (s)
(1)
This reaction releases so much enthalpy that the magnesium metal incandesces at white heat. It
will not be practical to measure the reaction enthalpy in your foam cup calorimeter, but you can
use Hess’s law and a combination of known and measurable reactions to indirectly determine the
molar heat of magnesium combustion in oxygen. The other reactions appear below:
H2 (g) + 1/2 O2 (g) → H2O (l)
(2)
(3)
Mg (s) + 2 HCl (aq) → MgCl2 (aq) + H2 (g)
(4)
MgO (s) + 2 HCl (aq) → MgCl2 (aq) + H2O (l)
The first of these, reaction (2), is the formation reaction of liquid water, which has a known
standard molar enthalpy value: ∆Hf° = –285.8 kJ·mol –1. Finding the appropriate combination of
reactions (2)–(4) will allow you to use Hess’s law for reaction (1).
PROCEDURE:
Reaction 4: Mg(s) + 2 HCl(aq) → MgCl2 (aq) + H2 (g).
1. Place a calorimeter cup inside a 250-mL beaker as before. Measure out 100.0 mL of 1.00M
HCl solution into the cup. Lower the temperature probe into the solution.
2. Weigh a 0.5-g sample of magnesium ribbon, and record the mass to milligram precision.
3. Record the temperature every 20 s. After you have obtained three readings at the same
temperature, add the weighed magnesium ribbon to the solution. Stir with a glass rod while
the temperature climbs. Once the temperature passes its maximum and starts to drop, suspend
data collection. If you can’t tell, collect data for 450 s.
4. Dispose of the product solution in the proper waste container.
5. Search your data to determine the solution’s initial temperature, T1, and the maximum
temperature, T2. Record your T1 and T2 values on the data sheet.
Reaction 5: MgO(s) + 2 HCl(aq) → MgCl2 (aq) + H2 O(l)
1. Repeat steps 1–7 using 1.0 g of MgO powder. Weigh the white solid to milligram precision
onto creased weighing paper, for easy transfer into the calorimeter. Record the MgO mass on
your data sheet.
Data Analysis
1. Calculate the temperature change, ∆T, for each calorimeter trial.
2. Calculate the heat, q, absorbed by the solution in each trial, using:
q = m CP ∆T
88 | CHM151LL: GENERAL CHEMISTRY I LAB MANUAL
3.
4.
5.
6.
7.
where CP = 4.10 J/g·°C for the product solution, and m is calculated by adding the mass of
the magnesium or magnesium oxide to the mass of the hydrochloric acid solution. A 1.0M
HCl solution has a density of 1.016 g/mL.
Calculate the enthalpy change of the reaction, ∆Hrxn, for each trial, using q + ∆Hrxn = 0. This
assumes that no heat is lost to the calorimeter or the surroundings.
Calculate the amount (in moles) of Mg and MgO reacted, using the known molar masses and
your measured mass values.
Calculate the molar enthalpy of reaction for reactions (3) and (4).
Use Hess’s law to calculate the molar enthalpy of reaction (1), the combustion of Mg, which
is also the formation reaction of MgO.
Look up an accepted value of ∆Hf° for MgO, and compare your value to the accepted value.
Calculate a percent error, using:
⎛ | accepted value – experimental value |⎞
% error = ⎜
⎟ × 100%
accepted value
⎝
⎠
THERMOCHEMISTRY III: HEAT OF COMBUSTION OF MAGNESIUM | 89
Name
Partner
Lab Report III
Data & Calculations Sheet for Magnesium Combustion Calorimetry
Data
Reaction 4
Mg(s) + HCl(aq)
Mass of solid
Reaction 5
MgO(s) + HCl(aq)
g
g
Maximum temperature, T2
°C
°C
Initial temperature, T1
°C
°C
Temperature change, ∆T
°C
°C
Volume of HCl solution
Ml
mL
Mass of HCl soln (calc)
g
g
Total mass (Mg + HCl soln)
g
g
Heat, q
J
J
Enthalpy change, ∆H
J
J
mol Mg
mol MgO
kJ/mol
kJ/mol
Calculations
Amount
Molar enthalpy change
Calculate the molar enthalpy change (kJ per mole of Mg) for reaction (1), using the ∆H values
you found (or were given) for the other reactions.
Calculate the percent error in your experimental value.
90 | CHM151LL: GENERAL CHEMISTRY I LAB MANUAL
Name
Date
Score
Postlaboratory Assignment
1. Use Hess’s law and the standard molar heats of formation from the Appendix in your lecture
textbook to calculate the standard molar enthalpy of combustion for benzene. That is,
calculate ∆H° (in kJ per mole of C6H6 (l)) for the reaction:
2 C6H6 (l) + 15 O2 (g) → 12 CO2 (g) + 6 H2O (l)
Show the individual chemical equation for the formation of each compound, written in the
appropriate direction, multiplied by the correct coefficient (where needed), showing the
enthalpy change of each individual reaction with the correct sign and numerical value.