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October 31, 2015
This week, we’ll measure enthalpies of dissolution in a constant-pressure calorimeter. We’ll use Hess’s law to determine the enthalpy of formation of magnesium oxide.
Concepts
Review enthalpy , enthalpy of formation , calorimetry , and Hess’s law from your lecture notes and textbook. Define terms and summarize concepts in your notebook.
Hess’s Law The enthalpy of formation of magnesium oxide is the enthalpy of the following reaction:
Mg(s) + 12 O 2 (g) −−→ MgO(s) ∆H = ∆H◦ f (MgO; s) (reaction 1)
Directly measuring the enthalpy of this reaction is challenging, however we can determine this enthalpy from the enthalpy of solution of magnesium metal
Mg(s) + 2 H+(aq) −−→ Mg2+(aq) + H 2 (g) ∆Hrxn2 (reaction 2),
the enthalpy of salvation of magnesium oxide
MgO(s)+2 H+(aq) −−→ Mg2+(aq)+H 2 O(ℓ) ∆Hrxn3 (reaction 3),
and the enthalpy of formation of liquid water
H 2 (g) + 12 O 2 (g) −−→ H 2 O(ℓ) ∆H = − 285 : (^8) molkJ (reaction 4).
Reaction 1 is rewritten using reactions 2 through 4:
Mg(s) + 2 H+(aq) −−→ Mg2+(aq) + H 2 (g) ∆Hrxn2 (reaction 2) Mg2+(aq)+H 2 O(ℓ) −−→ MgO(s)+2 H+(aq) −∆Hrxn3 (reaction 3, reversed) H 2 (g) + 12 O 2 (g) −−→ H 2 O(ℓ) − 285 : (^8) molkJ (reaction 4) Mg(s) + 12 O 2 (g) −−→ MgO(s) ∆H 2 −∆H 3 +
( − 285 : (^8) molkJ
) (reaction 1)
The enthalpy of formation of magnesium oxide is ∆H f◦ (MgO; s) = ∆Hrxn2 − ∆Hrxn3 +
− 285 : (^8) molkJ
We can use a constant-pressure calorimeter to determine the heats of solution of magnesium metal (reaction 2) and of magnesium oxide (reaction 3). With the heat of formation of liquid water, we get a series of reactions to relate the reactants and products of the heat of formation of magnesium oxide.
We’ll measure the enthalpies of solution of magnesium metal and magnesium oxide in a constant- pressure calorimeter (Figure 1).
Thermometer
connected to
Vernier data
interface
2 calorimter
covers
100 mL solution
in 150 mL beaker
Stir bar
2 coffee cups
for insulation
Figure 1. The constant-pressure “coffee cup” calorimeter used to measure the enthalpies of solution of Mg(s) and MgO(s).
Pre-Lab
For the procedure in your pre-lab template, it will be important to recognize that the lab procedure has three steps. In addition to the normal pre-lab requirements bring the following to include in your notebook: 1) Calculate the mass of MgO you will need to have the same number of moles of Mg present in 0.48 grams of Mg metal, 2) write out the 4 reactions on page 1 and explain how they can be combined to determine the enthalpy of formation of MgO, 3) include a copy of figures 2 and 3 which show the extrapolation procedure, and 4) include the data analysis table on page 4 of this document.
Procedure
Data collection will be done using the Vernier data interface and the Logger Pro data collection software. These steps refer to the Logger Pro Instructions, which will be available at your lab station.
Ti
Tf
Time solid is added
20
21
22
23
Temperature (ºC)
Time (minutes)
0 2 4 6
∆T
Figure 2. The extrapolation used to estimate ∆T = Tf − Ti for enthalpy of solution.
Data Analysis
A Calorimetry Report Sheet will be distributed and completed in lab.
Calorimeter Calibration: Heat Lost to the Calorimeter Calculate the heat transfer for the warm water and the heat transfer for the cold water: q = mC∆T with the specific heat of water C = 4: (^184) g·◦JC and the mass of water m = V with density = 0: (^9970) mLg at 25 ◦C and the volume of water, V , from the procedure.
The difference of the heat transfers for the hot and cold water is the heat transfer to the calorimeter. Dividing the difference by the temperature change for the cold water calorimeter gives the heat capacity of the part of the calorimeter—the amount of heat required to raise the temperature of the styrofoam cups, stir bar, the thermometer, etc by 1 : 0 ◦C.
The following is an example of how the heat capacity of the calorimeter is determined:
Extrapolated temperature of 50.0 mL of hot water: Th 37 : 9 ◦C Extrapolated temperature of 50.0 mL of cold water: Tc 20 : 9 ◦C Extrapolated temperature after mixing warm with cold water: Tmix 29 : 1 ◦C Heat transfer for hot water: qh = (49: 9 g) (– 8 : 8 ◦C)
4 : (^184) g·◦JC
– 1837 J
Heat transfer for cold water: qc = (49: 9 g) (8: 2 ◦C)
4 : (^184) g·J◦C
1712 J
Heat lost to the calorimeter: qlost = −qh − qc 125 J
The heat capacity of the parts of the calorimeter is then: Clost = (^) Tmixedqlost−Tc = 15: (^) ◦JC.
Heat Change of the Calorimeter During Solution of the Solids Throughout both reactions, the average concentration of HCl in the calorimeter is approximately 0. M. The heat capacity of the solution in this experiment is approximately the heat capacity of 0.4 M HCl, which is the average molarity during the reaction. Use this solution’s specific heat (Cs = 4: (^07) g·J◦C ) and density (s = 1: (^01) mLg ) to calculate the heat evolved during reactions 2 and 3:
qcal = (Csms + Clost) ∆T
where ms is calculated using the volume and the density of the HCl solution, Clost is determined from the calibration of the calorimeter, and ∆T comes from the extrapolation of the temperature data col- lected for each reaction.
At constant pressure, the enthalpy of each reaction is the heat of the reaction, ∆Hrxn = qrxn, which is absorbed by the calorimeter, qrxn = −qcal.
Post Lab Assignment
Write an abstract for a paper reporting the enthalpy of formation of magnesium oxide. Include a figure with your abstract.
Make sure your lab notebook includes a sketch that shows an example temperature run and explains with words the extrapolation method to determine Tf. If you use any variables, clearly define them. Also, include the completed Calorimetry Report Sheet in your lab notebook.
