NAME __________________________________ ACTIVITY: CAPTURING THE HEAT EXCHANGE
CALORIMETRY
Calorimetry (kal-ō-rim-eh-tree) is a precise measurement of the heat exchange between the chemical system
and the air or water surrounding the chemicals. Calorimetric measurements can be made for both, chemical
reactions and physical changes.
You know how to interpret the terms, exothermic and endothermic.
You also know that the chemical or physical process is classified from the point of view of the chemicals.
So, we must always establish what are the chemicals involved and then, what is the environment surrounding
the chemicals. In calorimetry the environment is usually, a surrounding mass of water.
When the surrounding water becomes warmer, energy must be flowing from the chemicals to the water. Thus
the process is considered to be EXOthermic (the chemicals are releasing energy, to the environment.)
When the surrounding water becomes cooler, energy is being absorbed by the chemicals, and they are taking
that energy from the water …hence, the process is ENDOthermic.
A natural inference, for instance is that given the correct circumstances, during an exothermic process, the
amount of energy released by the chemicals must be equivalent to the energy gained by the water.
To do this we use the equation q = mcT
where q is the joules of energy
m is the mass of the water surrounding the chemicals
c is the specific heat of water (4.18 J/g•°C)
T is the change is the temperature of the water
**************************
A couple of issues: First, let’s say we want to find out how much energy found in 10.000 grams of Cheetos®
Just for example, in the case of a foodstuff, chemists will burn the sample. But there is a problem… We need
the change in temperature and that is NOT the same thing as putting a thermometer in a flame … A flame has
many different temperature zones, which can be different based upon the material burning.
Max. flame
temperature
(°C, in air,)
Material burned
Animal fat
800–900
Kerosene
990
Wood
1027
Gasoline
1026
Methanol
1200
Charcoal (forced
draft)
1390
temp ~ 1400°C
temp ~ 1200°C
temp ~ 600 to 1000°C
temp ~ 600°C
http://en.wikipedia.org/wiki/Flame
http://www.candles.org/pdf/MatthaeiPetereit2004TheQualityCandle.pdf
and: http://steadyflame.mono.net/9655/BAGGRUNDSVIDEN
The way chemists overcame this problem of “variable flames” is to use a wonderful, consistent sample of water.
Basic Procedure:
1. Mass a sample of water.
2. Take the initial temperature of the water.
3. Initiate the reaction, and allow the heat of that reaction to pass into a sample of water.
4. Take the final temperature of the water
Using q = mcT, the rest is pretty straightforward.
So you want “q” … You have m, c and you can determine T by using:
Substitute in for T and solve the problem for q.
T = Tfinal - Tinitial
In theory, the amount of energy gained by the water sample (qwater) is the same as the energy released by the
reacting chemicals (qhydrate). So we can figure out how much energy a reaction produces (e.g. burning food,
bonding sodium acetate trihydrate), by determining the amount of energy exchanged with the water.
We shall do something similar today … We shall initiate the bonding reaction between the sodium ion and the
acetate polyatomic ion, in a supersaturated solution of sodium acetate trihydrate.
Bonding is an exothermic process, and energy should be evolved. If we can capture that released energy in a
known mass of water, with a known temperature, we should be able to calculate the heat of reaction.
Now, just try this:
1) Put on your goggles
2) Get a large Styrofoam cup. Mass it, recording that mass to 0.001 grams
Using a test tube containing water, as a model for your supersaturated solution, add enough water to
the Styrofoam cup, so that the volume of material (in theory your supersaturated solution) in the test
tube is just submerged beneath the level of tap water in the cup. Remove the test tube and mass the
cup and water, recording the mass to 0.001 grams.
Discard this water, and put the test tube away, when done.
3) Use a thermometer and take the temperature of the cup’s water, by holding the thermometer in the
middle of the water. Read the thermometer to a tenth of a degree. Record this initial temperature
to 0.1°C.
4) Remove the thermometer and put it down on your lab table top
5) Carefully conduct your supersaturated solution to your station and submerge the test tube with its
contents into the water bath. Obtain two or three crystals of sodium acetate trihydrate, from the
chemical hood, using a spatula or spoon. Bring them to your station and remove the rubber stopper
from the test tube. Add the crystals to the test tube contents and re-stopper it.
6) Immediately put the thermometer back into the water bath, and take the temperature of the water bath
every 15 seconds, until the temperature no longer changes, or until it begins to decrease.
7) Record the highest temperature. Be sure to read the thermometer 0.1°C
10) Remove the test tube of solidified NaAc and put it back beneath the hood.
11) Clean up!
NAME _______________________________
DATA TABLE: CALORIMETRY: HEAT OF REACTON
DATA TABLE FOR HEAT OF REACTION: SODIUM ACETATE: SUPERATURATED
1. Mass of the Styrofoam Cup
2. Mass of the Styrofoam Cup and Water
3. Mass of Just the Water (m)
4. Specific Heat Constant for Water (c)
5. Initial Temperature of the Water
6. Final Temperature of the Water
7. Change in Temperature of the Water (T)
8. Energy absorbed by the water (q)
[using q = mcT]
9. Energy released by the supersaturated solution
4.18 J/g °C