PLEASE RECORD ALL DATA DIRECTLY INTO YOUR LAB NOTEBOOKS
Introduction
Heating a substance is one of the simplest processes carried out in the chemical laboratory, and is
usually accompanied by a rise in the temperature of the substance being heated. The amount of heat
energy required to raise the temperature by 1°C depends both upon the amount (mass) of substance
and the identity (chemical composition) of the substance. The heat capacity, symbolized by C, is the
ratio of the heat added (q) to the observed temperature rise (T)
q
T
q  CT
C
Chemical reactions are often accompanied by the release or absorption of heat. For example, when
gas and heated the reaction produces an enormous amount of
hydrogen gas is combined with oxygen
heat are said to be exothermic. Reactions that absorb heat are
heat. Reactions that produce (release)
said to be endothermic.
The heat released in a chemical reaction is often determined by measuring the temperature change of
the material surrounding the chemical reactants and products. This works because the amount of heat
associated with a reaction is equal to the amount of heat that is either transferred to or from its
surroundings. The study of heat associated with chemical reactions is called thermochemistry. The
measure of temperature changes associated with chemical reactions is called calorimetry.
To the right is a diagram of a simple calorimeter, a device in which the heat of a reaction is measured.
The calorimeter can be as simple as a styrofoam cup and a thermometer. In a calorimeter the reactants
are placed into the container and allowed to react. The reactants and the products of a chemical
reaction, are called the system. The materials surrounding the system are called the surroundings. As
the reaction proceeds, the temperature of the liquid in which the reaction is occurring changes because
heat is transferred between the system and the surroundings. We can measure the change in enthalpy
ng the process. We use the system as the
reference and use the following sign convention:
Heat is absorbed by the system H > 0 endothermic
Heat is released by the system H < 0 exothermic
If the reaction is exothermic then the temperature
of the solvent (surroundings) will increase. If the
reaction is endothermic then the temperature of
the solvent (surroundings) will decrease. The change
in temperature of the solvent can then be used to
determine the amount of heat transferred:

H = –m * s *T
where H is the change in enthalpy for the reaction
(heat of reaction), m is the mass of the solution in
which the reaction is occurring, s is a constant called
the specific heat capacity, and T is the change in
temperature of the solution as the reaction takes place. The specific heat capacity depends on the
material used. Values of s are given below for several materials.
Material
Bi
Pb
Au
Pt
Hg (l)
Sb
I2 (s)
Sn
Cd
Ag
Se
Ge
Cu
s
c
J/(gK) J/(molK)
0.1221
0.1276
0.1290
0.1326
0.1395
0.2072
0.2145
0.2274
0.2311
0.2350
0.3212
0.3216
0.3846
25.52
26.44
25.42
25.86
27.98
25.23
54.44
26.99
25.98
25.35
25.36
23.35
24.44
Material
Zn
Co
Ni
Fe
Ti
Ca
Si
K
Al
Mg
Na
Li
Water
s
c
J/(gK) J/(molK)
0.3886
0.4210
0.4440
0.4494
0.5223
0.6315
0.7121
0.7565
0.9025
1.0238
1.2284
3.5609
4.184
25.40
24.81
26.07
25.10
25.02
25.31
20.00
29.58
24.35
24.89
28.24
24.77
75.40
Experimental Procedure
The experimental procedure has three parts. You first obtain the heat capacity of the
calorimeter by adding measured portions of hot and cold water. This value is used for the
second part of the experiment in which you are given a sample of an unknown metal. This
sample must be weighed, heated, and placed in the calorimeter. The observed temperature rise
is used to calculate the specific heat capacity of the metal.
Part I: Calibration of the Temperature Probe
Click CALIBRATION OF THE TEMPERATURE PROBE. There is a written set of instructions with
pictures that will take you through the steps of calibrating your temperature probe. THE
WATER THAT YOU USE FOR CALIBRATION SHOULD BE NEAR THE RANGE OF THE DATA THAT
YOU WILL COLLECT, SO USE COLD/COOL WATER AROUND 20 DEGREES AND WARM WATER
AROUND 40 DEGREES. WARM WATER CAN BE OBTAINED FROM THE TAP BY THE BENCH
WHERE THE LAB MATERIALS CAN BE FOUND.
Part II: Specific Heat Capacity of a Metal
1.) Dry out your calorimeter.
2.) Add some cool water (about 20°C) to your calorimeter. Use about 50 g of water, but record
exactly (to 4 decimals) the mass of water you use.
3.) Obtain about 15 g of an unknown metal. Record the mass of the metal that you actually
obtain, as well as which unknown you’re using.
4.) Put your metal in a large test tube, and put the test tube in a beaker of water. Heat the
water on a hot plate to approximately 90 °C. Be careful to place the test tube in the beaker
in such a manner that the water cannot splash into the test tube. Be certain that the metal
sample is in the hot water for at least 5 minutes.
5.) Extend the time of the run by pressing “Control D” and putting in 500 seconds.
6.) Proceed to Part III
Part III: Heat of Solution
7.) Nest two clean, dry Styrofoam cups together inside a clean, dry 400 mL beaker. This is
your calorimeter.
8.) Put about 50 g of deionized water into the calorimeter. Record the actual mass of water
in the calorimeter.
9.) Let the filled calorimeter stand for at least 4 minutes, to allow everything to come to
room temperature.
10.) Weigh out approximately 2 g of potassium nitrate (KNO3) and record its mass to the
nearest 0.0001 g. Grind the salt finely so it will dissolve uniformly.
11.) Observe the temperature of the calorimeter carefully. Run at least 60 seconds of
baseline for the water in the calorimeter. For the dissolution of salts, one of two
temperature trends can occur: the temperature may either rise quickly and then fall
slowly or fall quickly and then rise slowly. Record the maximum or minimum observed
temperature according to the trend observed.
12.) Repeat steps (7.)–11.), using approximately two grams of anhydrous sodium carbonate
salt (Na2CO3) and a fresh portion of water.
Part IV: Heat of Neutralization
13.) Assemble a clean, dry calorimeter.
14.) Add 25 mL of sodium hydroxide solution (NaOH) to the calorimeter using a graduated
cylinder. Cover the calorimeter and let it sit for at least 2 minutes with its temperature
probe in place. An excess of NaOH is being used in this experiment
15.) Using a graduated cylinder, transfer about 25 mL of hydrochloric acid (HCl) to a dry 100
or 150 mL beaker. Record the exact concentration of the acid, and exactly how much
you use.
16.) Run at least 60 seconds of baseline of the temperature of the NaOH.
17.) Use a thermometer to record the temperature of the acid in the beaker to the nearest
0.1°C.
18.) Quickly, but carefully, transfer all the acid from the beaker into the NaOH in the
calorimeter. Gently stir the contents, being careful to hold only the rim of the
calorimeter.
19.) Carefully observe the temperature of the calorimeter. The temperature may either rise
quickly and then fall slowly or it may fall quickly and then rise slowly. Record the
maximum or minimum observed temperature. Again, do this by extrapolation.
20.) Repeat steps 1–6, adding 25 mL of acetic acid (CH3COOH) to 25 mL of 1.10 M NaOH
previously placed in the calorimeter. Be sure to thoroughly rinse and dry the calorimeter
between runs. Make sure you know how much of each solution you use, as well as the
concentration of each solution.
21.) Write the balanced equations for both reactions, and the net ionic equations for both
reactions.
Part II: Specific Heat Capacity of a Metal (continued)
20. You are now ready to add the hot metal to the water in the calorimeter; first you need
to record the temperature of the water bath (This will be the TiM for the hot metal);
pour the hot metal into the calorimeter and allow the system to equilibrate for at least
2-4 more minutes. Do not delay while transferring the metal to your calorimeter. There
are consequences for delay, which you will figure out in the post laboratory questions.
As soon as the metal is added, swirl the calorimeter continuously to allow through
mixing of the water.
21. You may need to fit a straight line to the final linear part of your data. Click and drag the
mouse over a linear segment of the data from the end of your run to allow this section
of the line to be measured for slope and intercept. With the line segment selected, click
on “Analyze” and select “Linear Fit” or click on the “Linear Fit” icon (ask your TA).
22. To Print, click on “File” and then on “Print Graph”. Click on the “footer” box and add the
names and other pertinent information. It is recommended that you add the
experimental run (e.g. Run #4) and initials of yourself and your partner.
DO NOT THROW YOUR COFFEE CUPS AWAY! RINSE OUT
AND DRY.
DATA ANALYSIS
Part II: Specific Heat Capacity of a Metal
Analyze the calorimeter temperature versus time data as you did in Part II. Again
extrapolate to find Tf and obtain ∆TC, the temperature increase.
Use your value for CC to calculate the heat transfer (qC) to the calorimeter:
qC  mC s H2O TC  CC TC 
Your TA will provide you with the Cc value.
Remember that the 
heat that the metal lost is equal in magnitude to the heat the calorimeter
gained, but opposite in sign.
qC  q H
Use Tf to find ∆TH, the temperature decrease of the metal sample.
 metal sample of mass mM is exactly equal in magnitude (but
The heat transfer (qH) from the
opposite in sign) to the heat gained by the calorimeter:
q H  m M s M T H
Insert your experimental values and solve for the specific heat capacity of the metal.

The molar heat capacity (c) of a substance is related to its specific heat capacity through the
molar mass:
c  s M
Use the Dulong and Petit value of 26 J/(mol * K) for c and your experimental value for s to
obtain the molar mass, and hence the atomic mass M of your unknown metal.

Part III: Heat of Solution
The temperature change of the calorimeter is obtained by comparing its initial
temperature (Ti) with the final temperature (Tf) obtained by extrapolation of your data:
T  T f Ti
The amount of heat absorbed during the dissolution can be calculated from this temperature
change using:

qC  msolutions solutionT  CC T
where CC is the previously calculated heat capacity of the calorimeter and the mass of
the solution is given by m(H2O) + m(salt). The specific heat capacities of the KNO3 and
Na2CO3
Remember that the amount of heat absorbed by the calorimeter and water is equal in
magnitude to the amount of heat generated by the dissolution of the salt, but opposite
in sign.
qC  q rxn
Divide qrxn by the mass of the salt to obtain the specific heat of solution; divide qrxn by
the number of moles of salt to get the molar heat of solution.

Part IV: Heat of Neutralization
Find the total temperature change (T) of the calorimeter as in Part III.
Assume a density of 1.00 g/mL to find the mass (m) of the solution in the calorimeter.
Assume the specific heat capacity of the solutions to be equal to that of water. Use the
known heat capacity (CC) of the calorimeter to calculate the amount of heat evolved
during the reaction:
qC  msolutionssolutionT  CC T
Calculate the number of moles of acid contained in 25 mL of the 1.00 M acid solution. Divide
the heat of neutralization by this number. This is the molar heat of neutralization of the acid.
Calculate this quantity for both acids, and include them with your lab report.