Experimenting with Thermochemistry
Introduction
All chemical and physical changes are accompanied by transformations in energy. For example, you must
add energy to boil water, and you must remove energy to freeze water. Many of us heat our homes by
combusting natural gas with oxygen in the air. These types of energy transformations are examples of
thermal energy changes, and the processes by which these changes take place we term as heating or
cooling. Such a process is called sensible if you can measure a change in temperature (warming up
water) or latent if the energy transformation doesn’t involve a temperature change (boiling pure water). If
the system you are studying gains energy then we call it endothermic, and if the system loses energy then
it is called exothermic where the amount of energy transferred is q (in joules). Based on the law of the
conservation of energy, every exothermic process must be paired with an endothermic process, and viceversa. Understanding how processes exchange energy and the energy changes associated with
chemical reactions is an important and useful part of chemistry. Outcomes from this lab exercise will be:
 Preparing an adiabatic apparatus to measure transfers of thermal energy.
 Measuring sensible energy changes by measuring temperature differences to:
o Determine the specific heat of a metal.
o Measure the specific and molar enthalpy of solution for a dissolving process.
o Determine the molar enthalpy of reaction for a strong acid and strong base neutralization.
 Measure the latent enthalpy of fusion for ice melting.
 Further developing skills in lab techniques and chemical calculations.
 Understanding and using new chemical terminology.
Room temperature on earth is essentially established by the amount of energy we receive from the sun. If
you have a material that is hot or cold relative to room temperature, eventually it will come to room
temperature. In chemistry, this is called kinetics--the change in something as a function of time. If we use
a well-insulated system, the energy changes we are interested in will take place much faster than the
eventual change to room temperature. We call this an adiabatic approximation which then allows us to
calculate energy changes in the system we are interested in and ignore the heat transfer to the
surroundings. Using this approximation and the law of the conservation of energy together, we can write
0 = qsystem + qsurroundings ≈ qsystem
(1)
In this lab, we will directly determine the amounts of energy transferred between components of a system.
The data/results pages are written in the sequence you would measure and calculate the various values
requested. The bolded-type items are experimental values to measure and record and the non-bolded
values can be calculated later. This lab may be performed individually or with a partner.
Procedure and Discussion
The change in sensible heat is given by
q = m∙C∙∆T
(2)
Where the thermal energy transferred is q (J), the mass is m (g), the specific heat capacity C (J/g∙K), and
o
the difference in final temperature minus initial temperature is ∆T (K or C).
∆T = Tfinal - Tinitial
The difference function, ∆, in science is always assumed (final - initial) unless
otherwise stated.
Constructing a coffee-cup calorimeter.
Obtain two coffee-cups from the chemical bench and nest them together. Make
sure that the inner coffee-cup has no puncture wounds! You will also use a
o
cardboard lid, stirring rod (either triangular or regular) and a 0.1 C lab
thermocouple. The thermocouple tip is sharp, so don’t press it against the
Styrofoam coffee cup. Assemble the calorimeter similar to the figure on the
right. The energy changes we will be examining will take place inside the
coffee-cup calorimeter, and we will ignore heat transfer to the environment.
Del Mar College
(3)
stirring rod
cardboard
lid
thermocouple
nested cups
Page 1
Experiment 1. Determining the specific heat capacity of a metal
For this experiment the system will have two components involved in the energy exchange: water and
metal. Both components will change temperature (sensible energy changes), so we may combine
equations (1) and (2) to find:
0 = qwater + qmetal
=>
qmetal = -qwater
(4)
mmetal∙Cmetal∙∆Tmetal = -qwater = -mwater∙Cwater∙∆Twater
(5)
The specific heat capacity of water is Cwater = 4.184 J/g∙K. We will measure the mass of the water, mass
of metal, ∆Twater, and ∆Tmetal. This allows us to solve for Cmetal.
 Remove and weigh the inner Styrofoam cup from the calorimeter. Pour about 50 mL of DI water into
the calorimeter cup and weigh again. Record both masses on line 2. This will allow you to calculate
the mass of water in the calorimeter by difference. Measure the initial temperature of the water in the
calorimeter and record on line 6.
 Select a beaker with 5 metal cubes from the chemicals bench, and note their identity on the
data/results form on line 1. The data you need to calculate the mass of the metal by difference should
be written on line 4. Use a Bunsen burner or hot plate to bring water to boiling in a 400 mL beaker.
Place the metal cubes in a large, loosely stoppered test tube and immerse in the boiling water for 7-10
minutes. The level of boiling water should be above the level of metal cubes in the test tube. The
temperature of the boiling water will also be the initial temperature of the metal. Measure and note this
value on line 6.
 You will need thermal protection for your hand when you pour the hot metal in the test tube into the
calorimeter. Immediately replace the lid, carefully mix with the stirring rod and monitor the rise in
temperature in the calorimeter. Note the final temperature of the system on line 7 when the
temperature quits rising.
Calculate the ∆T for both materials. The ∆T water should be positive showing that water gained energy
(endothermic process), and the ∆T metal should be negative showing that the metal lost thermal energy
(exothermic process). Calculate and note q, the energy transferred for both materials. Finally, calculate
Cmetal and the %-difference and note both values on line 10. The accepted Cmetal values are provided.
Experiment 2. Determining the specific and molar enthalpy of a solution process
This system has two components involved in the energy exchange: the calorimeter-solution (cal-soln) and
o
the enthalpy of solution, ∆solutionH . Enthalpy changes are energy differences that occur at constant
pressure. Specific enthalpy refers to enthalpy changes per mass, and molar enthalpies involve these
changes per mole. Many physical changes such as forming solutions may involve significant energy
changes, and these spontaneous processes may be endothermic or exothermic. Equations (1) and (2) are
written as:
0 = qcal-soln + ∆solutionH
o
o
=>
o
∆solutionH = -qcal-soln
∆solutionH = -qcal-soln = -mcal-soln∙Ccal-soln∙∆Tcal-soln
(6)
(7)
The calorimeter-solution is mostly water, so use Ccal-soln = 4.18 J/g∙K. We will measure mass and ∆Tcal-soln
which will then allow us to calculate the enthalpy of solution.
 Remove and weigh the inner Styrofoam cup from the calorimeter. Pour 50 mL of DI water into the
calorimeter cup and weigh again. Record both masses on line 12. This will allow you to calculate the
mass of water in the calorimeter by difference. Measure the initial temperature of the water in the
calorimeter and record on line 16.
 Select a salt compound from the chemical bench and note its identity and molar mass on the
data/results form on line 11. You will want to add about 5 g of salt to the calorimeter. A weighing boat,
and weighing by difference, can be used for obtaining the salt mass. Record these values on line 13.
The mass of the calorimeter solution, mcal-soln, will be the combined mass of the water and salt.
 Add the salt to the calorimeter, and immediately replace the lid. Carefully mix with the stirring rod and
monitor the change in temperature in the calorimeter. Note the final temperature of the system on line
16 when the temperature quits rising.
Calculate ∆Tcal-soln and qcal-soln. From this you can calculate the total enthalpy change, specific enthalpy
change, and the molar enthalpy change. Was the enthalpy of solution endothermic or exothermic?
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Page 2
Experiment 3. Determining the molar enthalpy of a reaction
This experiment measures the molar enthalpy of reaction for the neutralization reaction of a strong acid
with a strong base. The net ionic equation is:
+
-
H (aq) + OH (aq) → H2O(l)
(8)
The system has two energy components involved in the energy exchange: the calorimeter-solution (calo
soln) and the enthalpy of reaction, ∆rH . Equations (1) and (2) are written as:
0 = qcal-soln + ∆rH
o
=>
o
∆rH = -qcal-soln
(9)
o
∆rH = -qcal-soln = -mcal-soln∙Ccal-soln∙∆Tcal-soln
(10)
Since the solutions to be mixed are weak, assume that Ccal-soln = 4.18 J/g∙K and that the density is 1.00
g/mL. We will determine mcal-soln and ∆Tcal-soln which will then allow us to calculate the enthalpy of reaction.
 Pour about 25.0 mL of 1.00 M HCl into the calorimeter and note the volume on line 22. Measure the
initial temperature of the HCl(aq) and record on line 24
 Pour about 25.0 mL of 2.00 M NaOH into a dry, clean 50 mL beaker and note the volume on line 22.
Measure the initial temperature of the NaOH(aq) and record on line 24.
 Add the NaOH(aq) from the beaker to the calorimeter, and Immediately replace the lid. Carefully mix
with the stirring rod and monitor the rise in temperature in the calorimeter. Note the final temperature
of the system on line 26 when the temperature quits rising.
Calculate mcal-soln from the added volumes corrected for density, ∆Tcal-soln and qcal-soln. From this you can
calculate the total enthalpy change for the reaction. Divide by the number of moles of the limiting reactant
to obtain the value of the molar enthalpy of reaction. This value will be the same for all neutralization
reactions of strong acids with strong bases.
Experiment 4. Determining the specific enthalpy of fusion of water (latent heat)
This experiment measures the specific enthalpy of fusion for ice melting (the latent heat of fusion). The
system has three components participating in the energy exchange: the initial calorimeter water, the
enthalpy of fusion of ice, and the ice-water warming up to the final calorimeter temperature. Equations (1)
and (2) are written as:
o
0 = qcal-water + ∆fusionH + qice-water
o
=>
o
∆fusionH = -(qcal-water + qice-water)
∆fusionH = -(qcal-water + qice-water) = -(mcal-water∙Cwater∙∆Tcal-water + mice-water∙Cwater∙∆Tice-water)
(11)
(12)

Remove and weigh the inner Styrofoam cup from the calorimeter. Pour about 50 mL of DI water into
the calorimeter cup and weigh again. Record both masses on line 32. Calculate the mass of water in
the calorimeter by difference. Measure the initial temperature of the water in the calorimeter and
record on line 36.
 Place ice from the top of the cooler into a Styrofoam cup. Weigh and record this mass on line 34.
Measure the temperature of the ice and record on line 36. Use the provided plastic spoon to remove
two spoonfulls of ice and immediately place the ice in the calorimeter. Reweigh the Styrofoam cup
with remaining ice, and record the mass on line 34.
 Immediately replace the lid on the calorimeter, mix with the stirring rod, and monitor the drop in
temperature in the calorimeter. Note the final temperature of the system on line 37 when the
o
temperature quits dropping. If the temperature reaches 0 C, you have added too much ice.
You have directly measured the masses, and can now calculate both ∆T’s. Calculate the enthalpy of
fusion for the system, then divide by the mass of ice to calculate the specific enthalpy of fusion. The
o
accepted value for the ∆fusionH of ice is 335 J/g. Calculate and report the %-difference on line 41.
o
o
o
o
o
Disposal and Cleanup
The calorimeter waste from Experiments 1 and 4 may be poured down the sink drain.
The solutions from Experiments 2 and 3 should be disposed of in the waste crock in the fume hood.
Return the metal cubes to their beaker and place in the drying oven.
Clean your glassware, rinse with DI water, dry and put away for next time!
Clean your desk area and also other areas as directed by the instructor.
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Experimenting with Thermochemistry
Advance Study Assignment
(May be on Canvas)
Problem 1. Determining a heat capacity. (Cwater = 4.184 J/g∙K)
o
A student has a coffee-cup calorimeter with 50.0 g of water at 21.5 C. They introduce 15.525 g of
o
o
magnesium metal initially at 100.0 C. The final temperature of the calorimeter came to 27.2 C.
5.65
What is the value for ∆Twater? _____________
K
-72.8
What is the value for ∆Tmagnesium? _____________
K
1192
What was the sensible heat change of water, qwater? ____________
J
-1192
What was the sensible heat change of magnesium, qmagnesium? ____________
J
1.0551
What is Cmagnesium? ______________
J/g∙K
Problem 2. Using the enthalpy of reaction.
How much did our sodium demonstration raise the temperature of water in the dessicator jar? The
reaction we performed was:
Na(s) + H2O(l) → NaOH(aq) + 1/2H2(g)
o
qreaction = ∆rH = -184 kJ/mol-rxn
Assuming the dessicator jar is well insulated (like our coffee-cup calorimeters). The energy balance is:
0 = qsystem = qreaction + qwater ==> qreaction = -qwater = -mwater∙Cwater∙∆Twater
o
The data was: 1.0 kg water, 22.0 C, Cwater = 4.18 J/g∙K, mNa = 0.21 g
22.99
The molar mass of sodium, Na is ________________
g/mol
0.0091
The moles of sodium added to the water was _____________mol
-1680
The J of energy produced was (watch units!) _______________
J
1000
Mass of water in jar is ____________________
g
0.4020
∆Twater is __________________
K
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Page 4
Experimenting with Thermochemistry - Data / Results Sheet
Stephen Henry
Name: ________________________
CHEM 1411.152FA
Section: ___________________
George Garcia
Partner: __________________________
10/20/2016
Date: ________________________
Experiment 1. Determining the specific heat capacity of a metal
1. Identity of metal cubes
3.900
2. Mass of calorimeter____________
g
Al
________________
52.500
Mass of calorimeter and water ___________
g
3. Mass of water
48.6
________________
g
168.411
4. Mass of container______________
g
183.769
Mass of container and cubes _____________
g
3. Mass of metal cubes
15.358
________________
g
o
20.4
6. Initial temperature water_______
C
o
98.6
Initial temperature of metal ____________
C
7. Final temperature of system
25.4
_______________
g
8. ∆Twater
_______________ K
∆Tmetal
9. qwater
_______________ J
qmetal
_______________ J
10. Cmetal
________________ J/g∙K
%-difference
_______________%
_____________ K
Experiment 2. Determining the specific and molar enthalpy of a solution process
11. Identity of salt
Unknown (2)
_______________
Molar mass salt
110.99
_______________g/mol
3.851
12. Mass of calorimeter ____________
g
53.005
Mass of calorimeter and water___________
g
2.322
13. Mass of weigh boat_____________
g
Mass of boat plus salt
7.345
______________
g
Mass of salt
5.023
______________
g
14. Mass of water
49.154
_______________
g
15. Mass of calorimeter solution (cal-soln)
54.177
______________
g
o
20.9
16. Initial temperature water________
C
o
26.9
Final temperature of cal-soln ____________
C
17. ∆Tcal-soln
________________ K
18. qcal-soln
_______________ J
∆solutionHo
19. Specific enthalpy of solution, ∆solutionHo
_________________ J/g
20. Molar enthalpy of solution, ∆solutionHo
_________________ kJ/mol
_____________ J
21. Is this solution process endothermic or exothermic?
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Page 5
Experimenting with Thermochemistry - Data / Results Sheet
Experiment 3. Determining the molar enthalpy of a reaction
22. mL of HCl(aq)
25.00
_________
mL
mL of NaOH(aq)
25.00
__________
mL
23. mL of solution
_________ mL
Mass of solution
__________
g
o
22.2
24. Initial temperature HCl(aq)________
C
o
22.3
Initial temperature NaOH(aq) _________
C
25. Average initial temperature of solution
______________ oC
26. Final temperature of solution
o
29.3
______________
C
27. Solution ∆T
______________ K
28. qsoln
__________ J
qrxn (= ∆Ho)
__________
29. Initial moles HCl(aq)
_________ mol
Initial moles NaOH(aq)
___________ mol
30. Moles of reaction
31. ∆rxnHo
J
______________ mol-rxn
___________ kJ/mol-rxn
Experiment 4. Determining the specific enthalpy of fusion of water (latent heat)
32. Mass of calorimeter
4.060
___________
g
52.406
Mass of calorimeter and water __________
g
48.346
_______________
g
33. Mass of water
34. Mass of ice and cup
11.569
__________
g
Mass remaining
3.775
____________
g
35. Mass of ice added
7.814
_______________
g
o
21.7
36. Initial temperature water ________
C
Initial temperature ice
37. Final calorimeter temperature
o
13.5
_______________
C
38. ∆Tcalorimeter-water
______________ K
qcalorimeter-water
_____________ J
39. ∆Tice-water
______________ K
qice-water
_____________ J
40. qice (= ∆Ho)
41. ∆fusionHo of ice
Del Mar College
o
1.2
____________
C
_______________ J
______________ J/g
%-difference
_____________ %
Page 6