Thermochemistry Laboratory Exercises #1, 2, 3, and 4
Chem 329
Ms. Hobbie
Next year TYPE!!! All of it…
Background – Thermal Equilibrium
In the first experiment you will observe the flow of heat energy from one substance (water) to another (a
metal). You will do this by monitoring the temperature change of the water. The flow of energy will stop
when both the water and the metal have reached the same final temperature. This is called thermal
equilibrium.
According to the law of conservation of energy…
The amount of energy lost by one substance
is equal to the amount of energy gained by another substance.
Another way to state this in terms of heat energy is:
qwater + qmetal = 0
Or, more generally, for any system with two or more components:
q1 + q2 + q3 +q4 + … = 0
Background – Calculating a Value for q: the Amount of Heat Transferred in the Process
From the temperature change of the water, you can determine the amount of heat transferred into or
out of the water using the following relationship:
qwater = mwater ∙ swater ∙ ΔTwater
q ≡ heat energy transferred (J)
m ≡ mass (g)
ΔT ≡ change in temperature, Tf - Ti (°C)
s ≡ specific heat of the substance (J/g·°C)
The value of q can be either positive or negative, depending on whether the water became warmer or
colder. For example, if heat is transferred into the water it will get warmer, ΔT (Tf – Ti) will be a positive
number and the value of qwater will be positive. The opposite is true when heat is transferred out of a
substance.
Background – Coffee-cup Calorimeters
When a warm substance and a cold substance are mixed within
a “coffee-cup” calorimeter, it can be assumed that there is no
loss of heat to the surroundings, that is, to the coffee cup.
Therefore the temperature change of the water is directly
proportional to the energy change of the system.
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Thermochemistry Lab #1 – Specific Heat of Water
Procedure and Analysis
If heat is added to a sample of water (meaning that energy has been transferred to the water), its
temperature will increase (meaning that its average kinetic energy will increase). The relationship
between heat and temperature change is called heat capacity, and is defined as follows:
Heat capacity is an extensive property (it depends on the size of the object). We can also define a
property called specific heat, which is the heat capacity per gram of a substance (this makes it an
intensive property, like density).
In order to determine a value for the specific heat of water, you would need to add a known quantity
of energy to a known mass of water and measure the temperature increase. This work was first done by
James Joule in 1843. He used an apparatus that allowed the change in potential energy of a falling
weight to be transferred to water.
Investigation Procedure
This is a “virtual” lab in which you will watch the workings of Joule’s apparatus and then be provided
data that can be used for the calculations necessary to determine the specific heat of water.
 Watch the following video: https://www.youtube.com/watch?v=PThq8fJpCLw
Data
The masses are 5.0 kg each.
The masses fall through a distance of 1.0 m.
The mass of the water is 2.0 kg
The change in temperature of the water is 11.7 ºC
The following assumptions are made:
 The masses fall at a constant speed, thus have neither gained nor lost kinetic energy
 The paddles stirring the water have negligible mass compared to the water
 The entire system is frictionless
Analysis
1. Use your physics background to calculate the total change in potential energy, PE, of the system
when the two 5.0 kg masses fall 1.0 m.
2. Determine the change in kinetic energy, KE, of the water. Hint: remember that energy is conserved.
3. Calculate a value for the specific heat of water.
4. Comment on how reality (i.e. the 5.0 kg masses did gain some KE; the paddles have mass; the
system isn’t frictionless) would affect the temperature change of the water.
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Thermochemistry Lab #2 – Specific Heat of a Metal
Procedure and Analysis
Investigation Procedure
 Weigh a piece of copper and record this mass on the data table below.
 Place the copper into boiling water for a while. Record the temperature of the boiling water.
Thought question: What will be the temp. of the copper after it’s been in the boiling water for a while?
 Place between 50.0 and 60.0 mL of tap water, precisely measured, in a Styrofoam cup and record
both the volume and the temperature of the water in the cup.
 After the copper has been in the boiling water for at least a minute, use tongs to quickly transfer the
hot copper into the water in the cup. BRING YOUR CUP OVER TO THE BOILING WATER TO DO THIS so
the copper will not cool down during the transfer.
 Gently swirl until thermal equilibrium is reached and record the final temperature.
Thought question: How will you know when thermal equilibrium is reached?
____________________________________________________________________________________________________
Data Table
Mass of copper:
Initial temperature of hot copper:
Volume of water:
Initial temperature of tap water:
Final temperature of the copper-water “system”:
____________________________________________________________________________________________________
Analysis (To be handed in… you may collaborate on all but the Extension questions)
1. Theory:
a. After adding the hot copper to the water, in what direction (water→copper or copper→water)
was heat transferred? Describe, at the molecular/atomic level, how this is accomplished.
b. Assuming no heat is lost to the cup or to the air, how does the heat change in the water (qwater)
compare to that of the copper (qcopper)? What law is at work here?
2. Results:
a. Using the specific heat equation, calculate the heat change in the water: qwater. (Watch signs!)
b. Use this value and your answer to 1(b) above to calculate the specific heat of copper (J/g∙°C).
(Again, watch signs.)
c. The accepted value for the specific heat of copper is 0.385 J/g∙°C. Determine the percent error
of your specific heat value.
3. Error analysis: Consider the following errors separately and explain how each would influence your
final specific heat value. That is, would your calculated value from 2(b) above have been larger
(positive deviation), smaller (negative deviation), or will it have remained unchanged? Explain each
in no more than two sentences.
a. Your Styrofoam cup is not a perfect insulator, and some of the heat from the metal was released
into the surroundings and did not end up heating the water.
b. Your starting metal was wet when you weighed it.
c. Your thermometer was consistently measuring 1.5 ºC lower than it should.
4. Extension: (To be done on your own… no collaboration)
a. Predict what the approximate final temperature of the water would have been if the copper
had been at the cooler temperature and the water in the Styrofoam cup was at the hotter
temperature? Defend your answer in words only, not calculations.
b. What final temperature would you have calculated if gold were the metal used in this exercise
rather than copper? The specific heat of gold is 0.129 J/g∙°C. Comment on your answer – does
it makes sense, compared to the final temperature of the cooled of copper?
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Thermochemistry Lab #3 – Heat of Solution
Procedure and Analysis
Investigation Procedure
 Place between 80.0 and 100.0 mL of tap water, precisely measured, in a Styrofoam cup and
measure the water’s temperature.
 Measure out approximately 10 grams (± 0.1 g) of NaOH(s), record the exact mass, and add it to the
water in the Styrofoam cup.
 Use the temperature probe to stir the NaOH(s) pellets while recording the temperature.
 When the solid has dissolved and the mixture has come to thermal equilibrium, record the final
temperature.
____________________________________________________________________________________________________
Data Table
Volume of water:
Initial temperature of water:
Mass of NaOH(s):
Final temperature of the solution:
____________________________________________________________________________________________________
Analysis (To be handed in… you may collaborate on all but the Extension questions)
1. Results:
a. Calculate the amount of heat transferred to the solution as the NaOH dissolved, qcalorimeter solution.
NOTE – this time the water in the calorimeter is actually a solution. We will assume the density
and specific heat of the solution are the same as that of pure water, but we must take into
account the MASS of any added substance that becomes incorporated into the solution.
b. Calculate the molar heat of solution of NaOH (ΔHsol), that is, the amount of energy exchanged
per mole of NaOH(s) that was dissolved in water. Think -- What should the sign of this value be?
2. Theory:
a. Identify the types of bonds/attractions that are formed and those that are broken whenever any
ionic solid dissolves in water.
b. Based on your observations, what can you say about the amount of energy involved with the
bond breaking process relative to the amount of energy involved in the bond forming process
when sodium hydroxide dissolves in water? Which is greater? Defend your answer.
c. Support your previous explanation with a simple energy diagram.
3. Extension: (To be done on your own… no collaboration)
a. The saturation of sodium hydroxide occurs at 105 g per 100 grams of water. If you dissolved
enough NaOH to reach saturation in 100 grams of water with an initial temperature of 23˚C, how
much heat would be released? Would it be enough for that solution to reach the normal boiling
point of water, 100°C? (Use your calculated molar heat of solution to solve this problem.
Assume specific heat does not change… in reality this is a terrible assumption for such a large
amount of solute. DO take into account that the mass of solution changes as NaOH is added.)
b. Specific heat of water generally decreases with the addition of a soluble compound. How
would this fact change the results of your calculations in #1 above? Explain your answer.
c. In another experiment, a coffee-cup calorimeter containing 100mL of water is used. The initial
temperature of the calorimeter is 23.0°C. If 7.90g of CaCl2 of is added to the calorimeter, what
will be the final temperature of the solution in the calorimeter? The molar heat of solution CaCl 2
is –82.8 kJ/mol.
d. We have assumed so far that the calorimeter is “perfect,” that is, that no heat is lost into or out of
the cup. How would the results of the calculation in (3.c) differ if the calorimeter had a heat
capacity of 275 J/°C? Show your work.
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Thermochemistry Lab #4 – Enthalpy of Reaction (Neutralization)
Procedure and Analysis
In this lab you are asked to determine the heat of reaction for 2 different acid/base neutralization
reactions, described below. The solutions provided are all 1.00 M.
#1: HCl (aq) + NaOH(aq)
#2: HCl (aq) + NH3(aq)
Investigation Procedure
 Add 50.0 mL HCl to a Styrofoam cup and determine the temperature of the solution.
 Add 50.0 mL of the base (NaOH for reaction #1, NH3 for reaction #2)) to a second cup. Rinse and
dry the probe, and determine the temperature of this second solution.
 Combine the two solutions and record the highest temperature reached.
 Repeat the process for reaction #2.
____________________________________________________________________________________________________
Data Table
Initial temp. of acid
Initial temp. of base
Final temperature
Reaction #1:
Reaction #2:
____________________________________________________________________________________________________
Analysis (To be handed in… you may collaborate on all but the Extension questions)
1. Results:
a. Calculate the amount of heat energy transferred during the first reaction (qrxn). Assume specific
heat and density values for the solutions are equal to those of pure water. Note: What should be
the sign of this value?
b. Calculate the amount of heat that would be transferred due to the reaction of exactly 1.00 mol
of NaOH with 1.00 mole of HCl, the amounts that match the coefficients of the balanced
equation. This is referred to as the enthalpy of reaction (ΔHrxn) or in this case, the enthalpy of
neutralization.
c. Repeat these calculations for the second reaction.
2. Theory:
a. Write a net ionic equation for each of these reactions. (Reminder: NH3 is a weak base.)
b. Explain in terms of bond forming what occurs during reaction #1. Is there any bond breaking
going on? Hint: What is the “real” form of H+(aq)? (Note: assume the forming and breaking of
bonds of hydration is negligible here…)
c. What is similar or different about the bond breaking and/or forming processes in reaction #2
compared to reaction #1?
d. Consider the heat released for these two reactions. Do you anticipate the H—O bond in water
to be stronger, weaker, or the same as the H—N bond in the ammonium ion? Defend your
answer.
3. Extension: (To be done on your own… no collaboration) the first question here should be open for
discussion… it is so importnt
a. Consider an experiment in which you use 100.0 mL each of 2.0 M NaOH(aq) and 2.0 M HCl(aq)
i. How would the qrxn of this experiment compare with the qrxn you calculated for reaction #1?
Be specific… if different, by how much? Defend with an explanation only, no calculations.
ii. How would the ΔHrxn you calculate compare with the ΔHrxn you calculated for reaction #1?
Be specific… defend with an explanation only, no calculations.
iii. How would the temperature increase in this experiment compare with the temperature
increase you recorded for reaction #1? Be specific… defend with an explanation only, no
calculations.
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-—OVER—
b. How would the data you collected in this lab (on your data table) differ if you had combined
50.0 mL of 1.0 M NaOH(aq) with 50.0 mL of 2.0 M HCl(aq). (Watch out… is there a limiting
reactant here?)
c. Here is a puzzle. In the fall we learned that when ammonium ions come in contact with
hydroxide ions, ammonia and water are formed. Your results from this lab can be used to
calculate the enthalpy of this reaction.
i. Calculate this value, assuming the ammonia that is formed is aqueous, and explain exactly
how you used your data to determine it. (Section 7.8 would be helpful with this question).
ii. Had the ammonia been formed as a gas, would your answer above have been different?
Howso? More/less negative, more/less positive?
d. Explain how you could use the results of this lab and calorimetry to determine the unknown
concentration of a solution of HCl(aq) if given a 2.0 M NaOH solution. Make sure your method
will work for any acid concentration, both greater than 2.0 M HCl and less than 2.0 M HCl.
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