Today is 11 August 2008 -- Comprehensive Calorimetry Experiment Report

We conducted an experiment about heat capacity.  Specifically we investigated the calorie content of different foods.  Here is a very well written lab report covering the experiment.  The"Reference" section needs more detail, however.  Please study the report and you will learn both how to write a lab report and you will understand a lot about the energy content of substances. 

I.  Purpose:  To determine the heating value (calorie content) of various substances, especially foods.

II.  Materials:  1.  Known volume of water in 250 ml Erlenmeyer flask

                         2.  Thermometer

                         3.  Known mass of vegetable oil (approx. 2 grams)

                         3a  Food-stuff or other combustible material (about 2 grams)

                         4.  Watch glass

                         5.  Cotton wick

                         6.  Apparatus stand and flask clamp

III.  Procedure: A.  Introduction:  The word calorimetry comes from the Latin word "Calor" meaning heat, and "-metry" meaning to measure.  The method depends on burning a known mass of something as fuel and recording the increase in temperature of a known mass of water.  The calorie is a unit of heat energy that is defined at the amount of heat required to raise the temperature of one gram of water (initially at 4 degrees Celsius) by one degree.  Thus, ten calories is the amount of heat to raise ten grams of water by one degree, or one gram of water by ten degrees.  Starting with water at its freezing point, it would take approximately 100 calories per gram to raise it to its boiling point.  To evaporate water at its boiling point takes an additional 540 calories per gram. 

                           B.  Method:  In this experiment, some students will measure the heat released by burning vegetable oil.  Other students will use a small amount of vegetable oil (weighed) as a fuel to start another substance, such as a corn chip, etc., to burn.  All materials are carefully weighed and the masses recorded in the student notebooks.    In all cases, the results will be expressed as "calories per gram."  The heat that is released by the combustion is used to raise the temperature of a known amount of water in a flask.   The final temperature of the water in the flask is read and from it is subtracted the initial temperature.  This temperature difference multiplied by the mass of the water in the flask represents the amount of heat liberated by the fuel in calories.  (In the event the fuel is very energy-rich, or if too much fuel is used, the water in the flask may boil.  In this case, the amount of liquid evaporated in grams must be determined (by subtraction from the amount initially present) and that difference must be multiplied by 540 calories to determine the amount of heat released by the fuel to vaporize the water.  That amount of heat plus the amount of heat to raise all of the water to the boiling point from the starting temperature is the total heating value of the fuel.

                          C.  Example:  50 ml of water initially at 22 degrees Celsius is heated by burning vegetable oil.  The amount of oil initially present was 5 grams.  When the experiment was stopped, 1.5 gram of oil remained and 10 ml of water had evaporated.  What is the heating value of the oil? 

IV.  Results:        Ans:  From the statement of the problem, we know that 10 ml of water, or 10 grams, had been converted to vapor.  Since it takes 540 calories per gram to evaporate at the boiling point, we know the amount of heat to evaporate the water was 540 cal/g X 10 g = 5,400 calories.  We also know that we started with 50 ml (50 g) of water and this went from 22 degrees to 100 degrees, a difference of 78 degrees.  We multiply 50 g X 78 degrees X 1 cal/degree/g = 3,900 calories.  We add 3,900 calories + 5,400 calories to get 9,300 calories.   The amount of oil that was burned was 5 g - 1.5 g = 3.5 g.  Dividing 9,300 calories by 3.5 g yields a heating value of 2,657 calories per gram of oil. 

V.  Error Analysis:  Some heat was lost to the surroundings because the experiment was not conducted inside of a thermally insulated calorimeter.  To give support to the oil as it burned, a cotton wick was used.  Combustion of the wick may have added a little energy to the water as it heated.  Combustion of the oil was not complete.  We know this because a layer of soot accumulated on the bottom of the flask.  Had the oil burned completely, there would not have been the soot layer and more water would probably have evaporated. 

VI.  Conclusions:  To a first approximation, this experiment allowed us to calculate the heating value, or enthalpy, of oil.  Had we used a mix of oil and a food, we could have used the data to determine how much of the heat came from the food and how much came from the oil. 

VII.  References:  Holt Chemistry

Today is 31 July 2009 -- Enthalpy study

Here are more instructions on how to make sodium acetate:

A better way of seeing when the reaction between sodium bicarbonate and acetic acid is complete is when CO2 bubbles stop being produced.  Then, you need to evaporate the solution to dryness.  Except if you do it all the way on a hot plate, you will burn the acetate salt.  When it is almost dry, but you have some liquid, put the beaker in the microwave for 30 seconds at a time.  When dry, the salt is fluffy and has a slightly waxy feel.

Next, add 25 ml of water to a 100 ml beaker.  Add about 100 ml of water to a larger (e.g. -- 250 ml) beaker.  Put the larger beaker on the hot plate.  Put the smaller beaker in it.  This makes a double-boiler and assures that the temperature of the 25 ml of water will not be higher than the boiling point of water.  When the water in the outer beaker is boiling, add the sodium acetate salt little by little until no more will dissolve to the inner beaker.  At this point, you have a supersaturated solution.   Take the smaller beaker out and let it cool.  You can put it in another 250 ml beaker with cool water in it to chill it faster.  If it remains fluid at room temperature, you probably have done everything right.  Then, add a crystal of sodium acetate.  It should start to crystalize and will release heat.  If so, reheat the beaker (in the boiling water bath -- 250 ml beaker) to dissolve and pour into a pouch.  Add the stainless steel initiator strip and seal the pouch.  Now, when it cools, you should be able to flex the stainless steel and see the sodium acetate crystallize with release of heat.

Today is 9 May 2010 -- Enthalpy experiment -- Make sodium acetate handwarmer

This experiment uses the product of the experiment where you reacted vinegar with sodium bicarbonate. It uses the left-over sodium acetate that was formed as a reaction product to make a handwarmer.  When sodium acetate solution is evaporated it becomes supersaturated.  That means there is more material dissolved than would normally occur.  Any disturbance of the solution causes it to crystalize and heat is released in the process.  This allows us to study the heat of crystallization, heat capacity and the concept of enthalpy.  

My intent was for my students to use plastic lay-flat tubing as the case for the solution.  I have a lot of it on hand.  Unfortunately, I've had the tubing for so long that I think it has degraded.  When we tried to demonstrate how to seal the tubing into bags, it leaked.  Until I get some better plastic, we will have to finish the experiment another way.  I still want you to learn about heat capacity so I have revised the experiment.  The procedure follow this note.  Please study it and follow the directions given:

1.  Accurately weigh out some dry sodium acetate crystals.   Use at least a gram or two in a small beaker.

2.   Put the beaker in a hot water bath until the crystals melt. 

3.  Take the beaker out of the bath and immediately put a thermometer into the melt.  Hold the thermometer so it doesn't fall out and shatter.  Record the initial temperature and the time.  Keep the thermometer in the melt until the crystals resolidify.  Record the temperature at time intervals.  You might try two minute intervals.  If that is too long or too short, repeat with another time of your selection.

4.  Prepare another beaker, larger than the one holding the crystals.  Accurately weigh twice the weight of water as crystals and put it in the larger beaker. 

5.  Clean off crystals that may be sticking to the thermometer.  Add those crystals back to the smaller beaker. 

6.  Put the thermometer in the larger beaker. (Mount it so it doesn't fall out.)  Take the initial temperature of the water.

7.  Return the beaker with the crystals to the hot water bath.  As soon as the crystals melt, put it in the beaker with the water.  Swirl the beaker to warm the water as quickly as possible. 

8.  Keep watching the thermometer.  Record the temperature as it rises to a peak (and the time), then starts back down.  Keep measuring time and temperature.  When the sodium acetate in the smaller beaker starts to re crystallize, note it (time and temperature).  You might see an increase in temperature at that point.  That would be due to the heat of crystallization.  If it does increase, great!  Track it.  It will hit a maximum, then start cooling.  The difference between the maximum temperature and temperature when the crystals started to form is due to the heat released from crystallization.

9.  We also need to know the heat capacity of the glass beaker.  To find this, first weigh the beaker and record the weight.  Then submerge the beaker in the hot water bath for a few moments.  Meanwhile, fill another container with cool water.  Take its temperature.  Using tongs, remove the beaker from the hot water bath, pour out the hot water and immediately fill it with cool water.  Insert the thermometer and watch as the temperature of the water rises.  Record the maximum temperature.  Weigh the beaker full of water.  Subtract the weight of the beaker (tare weight) from the filled weight (gross weight).  This tells you the weight of the water.  

10.  You know the heat capacity of water: It is 1 calorie per gram per degree C. (We can convert to joules, but it is more convenient to use calories for this experiment.)   You can calculate the enthalpy of the glass beaker from the temperature rise and the amount of water.  (Dividing this by the mass of the beaker yields the heat capacity of the glass per unit of mass.)   You will subtract the enthalpy of the glass from the overall heat gain of the water in the larger beaker. With this information, the moles of water in the larger beaker and the temperature difference, you can calculate the heat released by the sodium acetate.  That heat is the same needed by the sodium acetate to melt.  If you get two temperature peaks, you can calculate the heat released during crystallization.

11.  Here is the equation again;   q = Cp x moles x (T1 - T2)  where q = enthalpy, Cp = heat capacity per mole and (T1 - T2) is the temperature difference.