CHEM161 Experiment 7: Calorimetry
Introduction
Calorimetry is the science of measuring heat flow, and heat is defined as thermal energy
flowing from an object at a higher temperature to one at a lower temperature. For example, if a
chunk of metal at room temperature is placed in a beaker of boiling water, the metal will absorb
heat from the water until it is at the same temperature as the boiling water.
Scientists also often study the heat associated with different physical and chemical changes.
Ideally, the heat changes resulting from physical and chemical phenomena can be harnessed to
do work. For example, the burning of gasoline and fossil fuels can be used to run our
automobiles and heat our houses. However, in some cases, the heat associated with some
processes is transferred to the environment; for example, traditional incandescent light bulbs use
almost as much (or maybe more) electricity to produce heat rather than light.
Most physical and chemical changes are either exothermic or endothermic. Exothermic
reactions release energy or heat to increase the temperature of the surroundings; thus, the
surroundings are hotter after an exothermic change. For example, nitroglycerine exploding is an
extremely exothermic reaction. Endothermic reactions absorb energy or heat to decrease the
temperature of the surroundings; thus, the surroundings are colder after an endothermic change.
For example, cold packs used to relieve swelling joints or muscles often use chemicals that
absorb heat when mixed, so the packs feel cold.
When a reaction is carried out under constant pressure, the heat of a reaction is defined as the
enthalpy change for the reaction (H). Since most reactions occur under constant atmospheric
pressure, the heat of a reaction is equal to H, which is generally reported in units of kilojoules
(kJ) per mole of the reactants and products as shown in the balanced thermochemical equation.
For example, the reaction between hydrogen gas and oxygen gas to produce water is represented
by the following thermochemical equation:
2 H2(g) + O2(g)  2 H2O(l)
H = -571.6 kJ
[1]
Note that the sign of the enthalpy change indicates whether the reaction is exothermic or
endothermic. If H is negative, the reaction is exothermic. Since H for the formation of
water is negative, it is exothermic, so this reaction releases heat. The surroundings become
hotter, which is observed as an increase in temperature.
For endothermic changes, H is positive. Heat is absorbed from the surroundings, and because
the surroundings lose heat, they become colder, which is observed as a decrease in temperature.
For example, solid ammonium chloride dissolving in water is an endothermic process which is
represented by the following thermochemical equation:
NH4Cl(s)  NH4+(aq) + Cl(aq)
H = +14.9 kJ
[2]
To measure the enthalpy change for a chemical reaction in the lab, a calorimeter is used. A
calorimeter is any well-insulated apparatus that can hold the reactants and products, so the
amount of heat that is absorbed or released by a reaction or process can be measured.
Thus, the heat of the reaction (q) can be determined by measuring the temperature change
(T = Tfinal - Tinitial) for the calorimeter contents and using their specific heat—i.e., the heat
required to raise the temperature of 1 gram of the substance by 1C.
q = specific heat  mass  T
[3]
Note that the sign for the q is determined by the change in temperature (T). If the temperature
increases, T is positive, and q is positive, indicating the calorimeter contents absorbed heat.
Thus, the reaction must have released heat, so it is exothermic. If the temperature decreases,
T is negative, and q is negative, indicating the calorimeter contents lost heat. Thus, the
reaction must have gained heat, so it is endothermic.
In this experiment, the calorimeter used will consist of two nested Styrofoam coffee cups with a
cover and a thermometer (see Figure below).
However, instead of a glass thermometer, the Vernier LabQuest system will be used to measure
and record the temperature change for the calorimeter contents.
Ideally, the calorimeter is so well insulated, that all of the heat gained or lost during the reaction
is completely contained within the calorimeter. In reality, the coffee-cup calorimeter is not a
perfect insulator, so it will actually absorb and lose heat; however, in this experiment, the heat
absorbed or lost by the calorimeter is often negligible given the temperatures measured.
Heat transfer does not occur instantaneously, so using a calorimeter to determine the enthalpy
change requires measuring the temperature of the calorimeter contents as the physical or
chemical change occurs and for several minutes afterwards. The maximum temperature reached
by the calorimeter contents can rarely be measured directly because of drastic temperature
fluctuations near the probe or heat lost at the time of mixing. If the contents are mixed uniformly
while data is recorded, the contents should equilibrate at a consistent rate. Thus, the timetemperature data can be plotted, and the resulting regression line can be extrapolated to the time
of mixing to get the maximum or final temperature. The initial and final temperatures can then
be used to calculate the change in temperature (T) and the resulting heat gained or lost, as
shown in the example below:
Example for Determining the Enthalpy Change (H) for a Neutralization Reaction:
When 1.500 g of solid potassium hydroxide is added to 25.00 mL of 2.010M hydrobromic
acid at 22.5°C in a coffee-cup calorimeter, the temperature of the resulting solutions is
measured and recorded for 180 s. (The density of the hydrobromic acid is 1.159 g/mL, and
the specific heat of the solution is 3.70 J/g·°C.) The temperature data is plotted with respect
to time, and a regression line for the peak temperature and subsequent readings is obtained.
Figure 1: Temperature versus Time Plot for the HBr(aq) + KOH(s) Reaction
Temperature vs Time for the HBr(aq) + KOH(s) Reaction
Temperature (°C)
49.1
48.9
y = -0.0033x + 49.125
R2 = 0.8752
48.7
48.6
48.4
0
50
100
150
200
Time (s)
The equation for the regression line in the plot shows the y-intercept is 49.125°C, which is
rounded to 49.1°C to reflect the accuracy of the thermometers used—i.e., the thermometers
used to collect the data for the plot are only accurate to 0.1°C, so the final temperature is
also only accurate to 0.1°C.
The initial and final temperatures of the hot water and the cold water were measured, so the
change in temperature for each can be calculated as shown below:
T = 49.1°C – 22.5°C = 26.6°C
[4]
The heat absorbed by the solution can be calculated using the specific heat of resulting
3.70 J
solution,
, the volume and density of the hydrobromic acid along with the mass of the
g  C
solid sodium hydroxide to get the total mass of solution, and the change in temperature.
Calculating
the Heat
Absorbed
heat absorbed by the solution = qsol’n = cs · msoln · T
 3.70 J 

 g  C 
qsoln = 
[5]


 1.159 g 
(25.00 mL HBr) mL   1.500 g KOH (26.6°C) = 2999.35 J




Note that the heat calculated is not yet rounded to the correct significant figures (sig figs) but
must eventually be rounded to 3 sig figs, as indicated by underlining the first three digits.
Assuming the calorimeter is a perfectly insulator and absorbs no heat itself, the heat absorbed
by the solution should equal the heat lost by the reaction:
qsoln = -qrxn
Rearranging Equation [6] results in
[6]
qrxn = -qsoln = -2999.35 J
Enthalpy changes are generally reported in kilojoules per mole (kJ/mol) of a reactant or
product; for example, in acid-base neutralization reactions, like the example above, the heat
of neutralization (Hneut) is generally reported in kilojoules per mole of water produced
(kJ/mol H2O).
To determine the number of moles of water produced, the limiting reactant must be
determined using the balanced chemical equation:
HBr(aq) + KOH(aq)  H2O(l) + KBr(aq)
1.500 g KOH 
0.02500 L 
mol KOH
56.108 g KOH

mol H 2 O
mol KOH
 0.0267341 mol H 2 O
2.010 mol HBr mol H 2 O

 0.05025 mol H 2 O
L
mol HBr
Thus, KOH is the limiting reactant, and 0.02673 moles of H2O are produced. The heat of
neutralization (Hneut) for the reaction can then be calculated as follows:
- 2999.35 J
1 kJ

 - 112 kJ/mol H 2 O
0.0267341 mol H 2 O 1000 J
Thus, rounded to the correct number of sig figs, the heat of neutralization (Hneut) for
hydrobromic acid reacting with solid potassium hydroxide is -112 kJ/mol of water produced.
The heat for any reaction can be calculated in the same way, by measuring and plotting the
temperature change for the system. Again, the regression line is extrapolated to the time of
mixing to obtain the final temperature and calculate the overall temperature change.
In this experiment, the specific heat is 3.70 J/g∙°C for the resulting solutions in both reactions,
HCl(aq)-NaOH(aq) and HNO3(aq)-NaOH(aq). Thus, the heat absorbed or released by the
calorimeter contents is calculated as follows:
 3.70 J 

qsoln = 
 g  C 
 total mass 

 (Tsolution)
 of solution 
[7]
The initial temperature is taken as the average temperature of the reactants before mixing
assuming equal amounts of reactants are mixed, and the final temperature is obtained by plotting
the temperature change for the calorimeter contents as shown in Figure 1 on the previous page.
The mass of the solution is measured directly by weighing the calorimeter with and without the
solution and taking the difference.
The enthalpy change for both acid-base neutralization reactions (Hneut) will be reported in units
of kilojoules per mole (kJ/mol) of water produced. Thus, the amount of water that can be
produced given the amounts of each reactant must be calculated, and Hneut is calculated as
follows:
Hneut =
q re action
# of mole s of wate r produce d
[8]
In this experiment, the heat of neutralization for the reaction between hydrochloric acid and
sodium hydroxide will be determined experimentally using a coffee-cup calorimeter. In addition,
the heat of neutralization for the reaction between nitric acid and sodium hydroxide will be
determined experimentally in the same way.
The resulting salt solutions for both neutralization reactions (HCl-NaOH and HNO3-NaOH) will
contain a sufficient number of ions to change the specific heat to be slightly lower than water’s.
3.70 J
The specific heat used for each solution is
.
g  C
Procedure
SAFETY PRECAUTIONS: NaOH, HNO3(aq), and HCl(aq) are toxic and corrosive and can
all cause chemical burns upon contact with skin. NaOH can damage eyes on contact. Any acid or
base spilled on your skin or clothing must be rinsed immediately with water for 15 minutes.
Inform your instructor of an acid or base spills.
Students will work in pairs to complete the experiment.
Heats of
Neutralization
Experimental Determination of the Hneut for HCl(aq)-NaOH(aq) and for HNO3(aq)NaOH(aq)
1.
Obtain four Styrofoam coffee cups (to make two sets of nested calorimeters) and one
lid/cover. Use a pencil to mark one calorimeter #1 and the second calorimeter #2. Use a
centigram or milligram balance to weigh empty calorimeter #1 with the cover, and record
the mass.
2.
Check out a Vernier LabQuest system from the stockroom.
Note: The Vernier LabQuest systems are expensive and not covered by
standard lab breakage fees. The displays can be easily damaged by chemical
spills and must be kept as far away as possible from any chemicals. Students
will be charged the full cost to replace damaged LabQuest systems in addition
to any lab fees already paid.
3.
Record the molarity of the NaOH solution. Use the self-zeroing buret to measure out 50.00
mL of NaOH solution directly into calorimeter #1. Record the molarity of the hydrochloric
acid. Use the self-zeroing buret to measure out 50.00 mL of hydrochloric acid directly into
calorimeter #2.
4.
Plug the two temperature probes into “CH1 and CH2” of the LabQuest. Connect and plug in
the LabQuest’s AC adaptor. Turn on the LabQuest by pressing the Power button on the top
left corner (surrounded by a blue light). The system will take a few minutes to boot up.
5.
Swirl the NaOH solution in calorimeter #1, cover it with the lid, put the “CH1” probe
through the lid and into the NaOH solution in calorimeter #1. Measure the temperature of the
solution until the temperature remains constant for at least 30 seconds. Record the
temperature.
6.
The probes should be automatically detected, so when the screen comes on, it should display
two temperatures. On the right hand side are three windows: Mode, Rate, and Length. The
Mode should be “Time Based”; the Rate should be “2 samples/s”; and the Length should be
“180.0 s”. To make changes use the stylus (attached to the string) to tap on the appropriate
window to correct the setting.
Note: The tip of the LabQuest probes are sharp and can easily pierce the Styrofoam
cups. Be sure to keep the tip submerged in water or solution but do not touch the
bottom of the Styrofoam cup with the probe.
7.
8.
9.
Save the
LabQuest
Data
Place the “CH2” probe into the hydrochloric acid solution in calorimeter #2. (Don’t use the
lid for the acid.) Measure the temperature of the hydrochloric acid until the temperature
remains constant for at least 30 seconds. Record the temperature. Unplug the “CH2” probe.
Lift the lid and probe then press on the collect button,
, to collect the temperature of
the NaOH solution in calorimeter #1. Quickly but carefully transfer the acid to calorimeter
#1, replace the lid and probe, and swirl the solution to insure complete mixing. The LabQuest
will record the temperature of the solutions in calorimeter #1 for 180 s and plot the data on
the screen.
After 180 s of data, save the data to your thumb/Flash drive by tapping on the “File”
dropdown menu and choose “Save”. Click on the Name box on the top of the screen, and
enter the file name. The .qmbl format is the file type for the Vernier LoggerPro software. To
save your data in your thumb/Flash drive, tap on the “File” dropdown menu and choose
“Save”, then tap on the USB icon, and type in the file name. Choose a descriptive and
distinct filename for each trial.
10. Weigh and record the mass of calorimeter #1 with the solution and cover.
11. The LoggerPro software is installed on any student computer at SCCC. Retrieve the data
from the thumb/Flash drive by double-clicking on the .qmbl file, and the LoggerPro software
should open automatically. (If it does not, launch the LoggerPro software in the Physics
folder in All Programs from the Start menu., then open the file.) The LoggerPro software
should display both the data table and the plot.
Obtaining
the Tfinal
using
LoggerPro
12. Place the cursor on any point in the plot, and it will display the time and temperature. Use the
data table to determine when the temperature peaks and begins to fall, and place the cursor at
that point. Press down on the left mouse button and drag the cursor all the way to the right to
Adding Titles
to the Plot
and Data
Table
13. Double-click on the graph to open the Graph Options window, and enter a Title for the plot
(e.g. “Temperature Change for the Calorimeter”). To change the title of the Data Table, click
on the title, and a Data Set Options window will open.
Printing the
Data Table
and/or Plot
15. Dispose of the contents of calorimeter #1 in the sink, then wash and dry both calorimeters
completely. Wash and dry the cover. Use your water bottle to completely rinse the LabQuest
probes with DI water, then dry each with a Kimwipe.
highlight the points recorded after the temperature peaks. Tap on the “linear fit” icon,
,
and a text box will appear. Move the box if it shows up on top of your points. The value of
the y-intercept is the final temperature, Tfinal, of the water in the calorimenter.
14. For a record of the final temperature for the first trial of the HCl-NaOH heat of neutralization
determination, print a copy of the graph showing an appropriate title and the regression line
data with the y-intercept. Have your instructor approve your graph.
16. Repeat steps 2-8 for Trial #2 for hydrochloric acid.
17. Repeat steps 2-8 for Trials #1 and #2 for the HNO3(aq)-NaOH(aq) reaction.
18. LoggerPro allows the data table and graph to be printed together or separately. Follow the
instructions below for each:

To print both the data table and graph, click on the File dropdown menu, choose Page
Setup, and change Orientation to Landscape, then press OK. Next, type Ctrl+p or click on
the File dropdown menu and choose Print…. The Printing Options window should open.
Click on the Print Footer option, and type your name and your partner’s name, then OK.

To print just the graph, click on the File dropdown menu, choose Page Setup, and
change Orientation to Landscape, then press OK. Next, on the File dropdown menu,
choose Print Graph…, and the Printing Options window should open. Click on the Print
Footer option, and type your name and your partner’s name, then OK.

To print just the data table, click on the File dropdown menu, choose Print Data
Table…, and the Printing Options window should open. Click on the Print Footer option,
and type your name and your partner’s name, then OK.
19. Click on the numbers on the y-axis until the Format Axis window appears. Click on the Scale
tab. Change the Minimum from zero to be just under the lowest temperature, and if
necessary, change the Maximum to be just above the trendline/regression line.
Dispose of the inside cup in each calorimeter in the trash, and return the outside cup to the
stockroom. Wash and dry all glassware used. Take a wet paper towel, and wipe down your
entire work area.
DATA ANALYSIS AND CALCULATIONS
Average the initial temperature of the acid and base for each trial.
Use LoggerPro to prepare plots for the temperature change for the remaining neutralization
reactions between HCl(aq)-NaOH(aq) and HNO3(aq)-NaOH(aq).
For each plot,
 Provide a descriptive title for the plot.
 Label each axis (including units).
 Include a trend/regression line, showing the equation for the line with the y-intercept.
Use
3.70 J
for the specific heat of the salt solutions resulting from both neutralization
g  C
reactions (HCl-NaOH and HNO3-NaOH).
For each trial, calculate the heat absorbed by the solution then calculate the heat released by the
reaction using Equations 7-8. Next, calculate the number of moles of water produced using the
amount of each acid and base used for each trial. Calculate the enthalpy of neutralization (Hneut)
for each reaction in kJ per mole of water produced.
Experiment 7: Calorimetry
Name:
__________________________
Pre-Laboratory Assignment
Section Number:
__________________
1. Predict the products (including physical states), and write the total ionic and net ionic equations for the
reaction between hydrochloric acid and aqueous sodium hydroxide.
balanced
chemical
equation:
total
ionic
equation:
net ionic
equation:
2. Predict the products (including physical states), and write the total ionic and net ionic equations for the
reaction between nitric acid and aqueous sodium hydroxide.
balanced
chemical
equation:
total
ionic
equation:
net ionic
equation:
3. Given the net ionic equations for the reactions above, explain if the enthalpy of neutralization (Hneut)
should be the same for both reactions.
4. Briefly explain why the “final temperature” of the calorimeter contents is not simply the last or highest
temperature recorded during a calorimetry experiment.
Experiment 7: Calorimetry
Name:
________________________________
Partner:
_______________________________
DATA
Mass of empty Calorimeter #1 + cover: ________________________________
Experimental Determination of the Hneut for HCl(aq)-NaOH(aq) and for HNO3(aq)-NaOH(aq)
sodium hydroxide
hydrochloric acid
nitric acid
50.00 mL
50.00 mL
50.00 mL
Molarity of the solution
Volume of solution used
in each trial
HCl(aq)-NaOH(aq)
Trial #1
Trial #2
Initial Temperature of HCl(aq)
Initial Temperature of NaOH(aq)
Average Initial Temperature
Mass of Calorimeter #1 + solution
+ cover (after reaction)
Mass of empty Calorimeter #1 + cover
(from part A)
Mass of solution*
* Use Mass of empty Calorimeter #1+ cover from Part A to calculate
HNO3(aq)-NaOH(aq)
Initial Temperature of HNO3(aq)
Initial Temperature of NaOH(aq)
Average Initial Temperature
Mass of Calorimeter #1 + solution
(after reaction)
Trial #1
Trial #2
Mass of solution*
*Use Mass of empty Calorimeter #1+ cover from Part A to calculate
For each trial, attach the Data Table obtained using the Vernier LabQuest system, along with the titled
graph with labeled axes that shows the regression line and the y-intercept with the maximum
temperature (Tfinal).
Show calculations for each step of the HCl(aq)-NaOH(aq) and the HNO3(aq)-NaOH(aq) neutralization
reactions in your lab notebook.
Results Table to Determine Heat of Neutralization for the HCl(aq)-NaOH(aq) Reaction
Trial #1
Trial #2
Maximum Temperature
(y-intercept) from the plot
Change in Temperature (T)
qsolution (in kJ)
qreaction (in kJ)
# of moles of H2O produced
Hneut (in kJ/mol H2O)
Average Hneut (in kJ/mol H2O)
Results Table to Determine Heat of Neutralization for the HNO3(aq)-NaOH(aq) Reaction
Trial #1
Maximum Temperature
(y-intercept) from the plot
Change in Temperature (T)
qsolution (in kJ)
qreaction (in kJ)
# of moles of H2O produced
Hneut (in kJ/mol H2O)
Average Hneut (in kJ/mol H2O)
Trial #2
CHEM 161 Experiment 7: Calorimetry Post-Laboratory
Questions
Post-Laboratory Calculations to include in the Formal Lab Report
1. Write the net ionic equation for the reaction between hydrochloric acid and aqueous sodium
hydroxide, then calculate the heat of neutralization (Hneut) for the reaction using the standard
enthalpies of formation ( H f ) in the Appendix of the textbook for each reactant and product.
Do the same for the reaction between nitric acid and aqueous sodium hydroxide.
2. Calculate the percent error for your average experimental average Hneut for each reaction
using the following formula:
| experiment al value  true value |
 100%
true value
Post-Laboratory Questions to address in the Discussion Section of the Formal Lab Report
% error =
1. Use the observed temperature change to explain if acid-base neutralization reactions are
endothermic or exothermic. Explain if the sign for the calculated heat of neutralization (Hneut)
should be positive or negative based on the observed temperature change.
2. Are the heats of neutralization for the reaction between hydrochloric acid and aqueous sodium
hydroxide and for the reaction between nitric acid and aqueous sodium hydroxide the same?
Explain why. Would you expect the heats of neutralization for the reactions to be the same if
barium hydroxide were used instead of sodium hydroxide? Explain why.
3. How would the temperature change and the calculated heats of neutralization vary if the
concentrations of the acids and the base used in this experiment were doubled? Explain.
4. Would the temperature change and the heats of neutralization vary if the concentrations
remained the same but the volume of each acid and base used in this experiment was cut in
half? Explain.
5. If your percent error was greater than 5%, discuss possible sources of error in the experiment.
If you carried out more than two trials and/or omitted data from a trial to calculate your
average, discuss why this was necessary. (Provide specific sources of experimental error that
could have caused your incorrectly high or incorrectly low experimental results. Do not discuss
calculation errors.)
Note: Your formal lab report should include the four graphs (two each for the acid-base
neutralization reactions), each with a descriptive title and labeled axes. In addition,
your report should also include the purpose, procedure, data, calculations, results in
summary tables, discussion, and conclusion.