Advanced Chemistry Lab Experiment 3
Calorimetry: The Measurement of Heat Changes
REFERENCES
capacity, C, of the calorimeter (or, more commonly, the calorimeter constant, Ccal).
The mathematical expression for heat capacity is
Heat capacity(joules/oC) = specific heat (joules/g oC) x mass(g)
th
“Chemistry and Chemical Reactivity”, 6 Ed., Kotz/Treichel/Weaver
Chapter 6: (Principles of Reactivity: Energy and Chemical Reactions)
J. Chem. Educ., 1991, 68, 229
OBJECTIVES
1. To gain some experience in simple calorimetry (measuring the heat produced by a
chemical reaction).
2. To use graphical representations to obtain data.
3. To determine the enthalpy change, !H, for a simple acid-base reaction.
BACKGROUND
As noted in your text, energy is the ability to move matter, that is, to do work.
Types of energy include chemical, light, heat, sound, electrical, nuclear, and mechanical.
Because much of the discussion of energy in the text is heat-related, this experiment will
introduce you to some very basic heat concepts.
Heat is defined as the form of energy contained by a substance that gives the
substance its temperature. One unit of heat commonly used by scientists is the joule (J).
If we know the mass of a substance and its specific heat, we can determine the
amount of heat leaving or entering a substance by measuring the temperature change.
The mathematical expression for this determination is:
q = (c)(m)( !T)
where q is the amount of heat leaving or entering the solution or surroundings, c is the
specific heat of the substance, m is the mass of the substance, and !T is the change in
temperature, (Tfinal - Tinitial) or (Tf -Ti).
The measurement of heat changes is called calorimetry (the name comes from an
old, but still used unit of heat called the calorie). [Note: The calories (with a small “c)
should not be confused with the Calorie (with a capitol “c”) on a candy bar or other food
item. The Calorie on all food labels are in fact kcals or 1000 calories]. Heat changes are
measured in a calorimeter, a reaction vessel within an insulated container with provision
for measuring a change of temperature. The basic concept that makes calorimetry a
quantitative study is the Law of Conservation of Energy: energy is neither created
nor destroyed. It is only changed in form or transferred between objects. The data
obtained from calorimetric measurements have application to many aspects of chemistry
including equilibria, electrical potentials, thermodynamics, and biochemistry. Although
very accurate calorimeters can be expensive and somewhat tedious to operate, a simple
calorimeter constructed from Styrofoam coffee-cups can be constructed and calibrated for
use in the general chemistry laboratory.
Even though the calorimeter assembly may be well insulated, the calorimeter does
absorb or liberate some heat. Thus, if good results are to be obtained, the calorimeter
must be calibrated by determining the quantity of heat absorbed or liberated by the
calorimeter per degree of temperature change. This quantity of heat is called the heat
or
C = (c)(m)
For a given calorimeter, the calorimeter constant generally can only be considered
to remain constant when (1) reactions take place at about the same speed, (2) the total
volume of the calorimeter contents is approximately the same for each determination, and
(3) temperature changes do not vary too much from one reaction to the next. In short, the
calibration procedure should be carried out under conditions nearly identical to those of
the actual determinations.
In this experiment you will first determine the heat capacity of a coffee-cup
calorimeter and then you will use that calorimeter to determine the enthalpy change for
an acid-base reaction.
Prelaboratory Assignment
Work out this problem on paper. Hand it in as you enter the laboratory.
Exactly 50.0 mL of 1.00 M hydrobromic acid and 50.0 mL of 1.05 M sodium hydroxide
are combined. Heat is liberated and the temperature change is 6.4 °C. Assume the
calorimeter does not absorb any heat, so the heat capacity is essentially that of 100 grams
of water, with specific heat capacity of 4.184 joules/g ºC.
a) Calculate the heat released by the reaction.
b) Calculate the number of moles of each reactant.
c) Identify the limiting reactant.
d) Calculate the enthalpy change per mole of limiting reactant for this reaction.
PROCEDURE
In your laboratory notebook prepare two data tables for Part 1 of this experiment,
(allow space for 14 temperature readings in each trial) and two data tables for Part 2.
Suggested data tables are shown in below.
Check out from the stockroom a calorimeter, stirrer, and a thermometer.
Obtain in a container at least 100 mL of distilled water and set it aside to come to
room temperature (for Part 2).
1. Determination of the Heat Capacity of the Calorimeter
Write all observations in your laboratory notebook.
Configure your tables as shown below:
Time-Temp Data Table
Time (sec)
Temperature (ºC)
Trial 1
Trial 2
-90
-60
-30
ZERO SECONDS
There is no temperature reading at time zero since you should have
been transferring the water from the beaker to the calorimeter at this
time.
30
60
90
etc.
Final Reading at 300 seconds!
Complete a total of SEVEN more readings at 30 second intervals!
Calorimeter Calibration Data Table
Trial 1
Trial 2
Mass of Calorimeter
Mass of Calorimeter and Water
Mass of water (mw)
Room Temperature (Tci)
Data Obtained from a Graph of the Time-Temp Data (graphing done in class)
Twf – Twi
Twf – Tci
Calculate the Calorimeter heat capacity (Ccal) using the equation provided in Part
1 of the calculations section of you lab manual.
Ccal
Average Ccal
*wf = water final; wi = water initial; ci = calorimeter initial; T = temperature
Weigh the calorimeter (The mass of the calorimeter is the mass of the cup, top,
and the stirrer! ) to the nearest mg (0.01 g) using a balance. Record this mass on the
data sheet in your notebook. Mount the calorimeter as shown in provided on the next
page.
Read and record the room temperature, estimating to the nearest 0.1 oC. (This will be
the initial temperature of your calorimeter, Tci).
Heat 100 mL of water in a beaker to a temperature of between 60 oC and 70 oC.
Remove the burner. When the temperature of the water starts to fall, record the
temperature to the nearest 0.1 oC on your first time-temperature data table. Take two
more temperature readings at 30-second intervals. (These three readings will be the -90, -
60, and -30 sec readings.) Then, 30 seconds after you make
your third temperature reading (this will be time 0 on your data
table), pour the water from the beaker into the calorimeter and
immediately replace the lid. Wipe the water off your
thermometer and let the thermometer cool briefly. Then place
the thermometer in the calorimeter. Try to do this so that, as
nearly as possible, there are 60 seconds between your last
temperature reading (at -30 sec) and your next temperature
reading (at 30 sec) of the water in the calorimeter. Stir the
water with the stirrer while you read and record the temperature
of the water in 30-second intervals until you have made 10
more readings (5 more minutes of temperature readings).
Carefully remove the thermometer from the calorimeter.
Weigh and record the mass of the calorimeter and the water to
the nearest mg. The difference between the mass of the
calorimeter with the water and the mass of the calorimeter
Figure 1: Calorimeter Setup
without the water is the mass of the water, mw.
Carefully dry your calorimeter with a paper towel. Repeat this experiment, recording
your data on your second data table.
To determine the temperature of the water in the beaker at time 0 (Twi), and the
temperature of the water cooled by the calorimeter at time 0 (Twf), you must plot your
data. We will do this using a computer spreadsheet. Type your data for the first run into
the spreadsheet template provided; a graph will be plotted as you type. Twi and Twf, will
be determined by the spreadsheet from line-fits of your data. Make sure the data you
typed in is correct then print the graph and repeat the process with your second data set.
The intersections of your straight lines with the time zero line will be Twi (upper one)
and Twf (lower one). Then T = Twf - Twi .
2. Determination of the Enthalpy Change, !H, of an Acid-Base Reaction
You will be assigned a specific acid (HCl) and a specific amount of that acid to
use. The heat produced by the reaction of the acid and the base sodium hydroxide will be
measured in this part of the experiment.
acid(aq) + NaOH(aq) ! salt(aq) + H 2 O(l) (not balanced)
The results and the known concentrations of the reactants will be combined to
calculate the enthalpy change per mole of limiting reagent for this reaction.
An example data table is proved on the next page.
Example DATA TABLE FOR PART 2
Trial 1
Trial 2
HCl(aq)
NaOH(aq)
HCl(aq)
NaOH(aq)
Molarity
Volume
(before dilution)
moles
The amount of heat produced from the reaction is dependent on the
amount of each reactant utilized. In this case one reagent is in
excess. Thus the reagent which runs out first (the limiting reagent)
limits the amount of heat produced. Use the above data and the
balanced reaction to determine the limiting reagent.
Limiting Reagent
Ti
Tf
Calculations
Complete this using the equation provided below or in step two of the
Calculations section of Part II of the lab manual.
Tf -Ti
qrxn (kJ)
! H(kJ/mol)
Average
H
Using your large graduated cylinder obtain your assigned amount (30-50 mL) of acid.
Add enough room-temperature distilled water to the acid to make a total of 50 mL of acid
solution. Stir the solution well with a glass stirrer (the thermometer works well as a
stirrer) (Why not the metal stirrer?). Transfer this solution to a clean beaker. Rinse out
the graduated cylinder thoroughly with distilled water and shake it dry.
Dry your calorimeter. Using a graduate cylinder, add 50 mL of 2.2 M NaOH(aq)
to the calorimeter. Record the initial temperature of this solution as Ti. Then add the 50
mL of acid solution to the calorimeter. Stir the solution well with the thermometer and
promptly put the top on the calorimeter. Insert the thermometer and monitor the
temperature to the nearest 0.1 oC. (The temperature should rise, then fall again.) Record
the highest temperature attained as temperature final (Tf). Discard your solution by
pouring it into the sink and rinse the calorimeter thoroughly with distilled water.
Repeat the procedure at least once.
Do the calculations for Part 1 before leaving the laboratory and have your
instructor check them (Reasonable values for Ccal have a range of about 4 to 25 J/oC).
CALCULATIONS
PART 1:
Background theory:
We will use three basic equations and substitute appropriate values:
(1) q = c x m x !T
(2) qcal + q water = 0
(3)
Ccal = c x m
Combining equation (1) and (3) gives
qcal = Ccal x (Twf - Tci)
and equation (1) gives:
q water = 4.184 x mw x (Twf - Twi)
because the specific heat of water is 4.184 J/(g oC).
Substituting these equations into equation (2) gives:
Ccal x (Twf - Tci) + 4.184 x mw x (Twf - Twi) = 0
which rearranges to
!4.184(mw )(Twf ! Twi )
Ccal =
Equation I
Twf ! Tci
Calculations:
1. For each trial, solve for Ccal using the values you have obtained experimentally in
Equation I.
2. Calculate an average value for Ccal.
(Have your instructor check this value before proceeding. )
PART 2:
Do the following calculations for each trial.
1. Using the average heat capacity of the calorimeter, Ccal, calculated in Part 1, and the
temperature changes that you observed in Part 2, calculate qcal for each trial using the
expression:
qcal = Ccal x (Tf - Ti)
2. Background theory:
The heat of reaction, q, is included in this expression:
qrxn + qcal + qsoln = 0
which rearranges to
qrxn = - (qcal + qsoln)
or
qrxn = ! (Ccal +418.4)(Tf -Ti )
Equation II
(assuming that the mass of solution is 100 g and its specific heat is 4.184 J/(g
o
C).)
Calculation: Calculate the heat of reaction according to equation II.
3. Convert the above qrxn to molar enthalpy change, !H, in units of J/mol.
Finally: Calculate for the two trials an average value of !H in kJ/mol. Compare your
result with the result you get for the following exercise.