LAB: ENERGY CONTENT OF FUELS
INTRODUCTION
The availability of high-quality energy sources is one of the most important issues facing
the global community in the 21st century. Chapter 4 in Chemistry in Context examines
some of the societal choices we face and the advantages/disadvantages of various fuels.
Hydrocarbons (or mixtures of hydrocarbons) are among the most energy-rich and
cleanest-burning fuels. Many of them are liquids, which makes them attractive for
transportation fuels. Some alternative fuels may be derived from either fossil fuels or
from plant sources. The use of renewable plant sources is attractive because it could
reduce the U.S.’s dependence on foreign oil and could also support U.S. agriculture.
In this experiment you will investigate the energy content of two hydrocarbon fuels by
using them to heat water. The data that you and your classmates obtain will enable you to
compare fuels to see which ones provide more energy for a given mass of fuel burned.
You will also be able to draw conclusions about the energy implications of using
oxygenated fuels, such as ethanol. Finally, you will become familiar with alternative fuel
sources, such as plant matter and animal fats.
BACKGROUND INFORMATION
Fuels are substances that burn to give off a relatively large amount of heat. In an overall
sense, such burning is simply a combustion reaction between the fuel and oxygen. How
much heat is generated depends on what kind of fuel is used and how much of it is
burned. The simplest example of a common fuel is natural gas, which is almost pure
methane (CH4). When it is burned completely, carbon dioxide and water are the only
products:
CH4 + 2 O2  CO2 + 2H2O + energy
Methane is an example of a hydrocarbon fuel. Hydrocarbons are compounds that contain
only hydrogen and carbon. Other hydrocarbons, such as propane (C3H8) and butane
(C4H10), will burn to produce the same products—carbon dioxide and water. Other
hydrocarbon fuels include C10H22 (kerosene), C12H26 (lamp oil), and C40H82 (candle wax).
Some fuels such as ethanol (C2H5OH) contain oxygen in addition to carbon and
hydrogen. In effect, they are hydrocarbons that have already partially reacted with
oxygen:
2 C2H6 + O2  2 C2H5OH
These are alcohols. For example, ethanol is also called ethyl alcohol.
In this experiment, you will measure the amount of heat given off by known amounts of
several alcohol fuels. The fuels you will study are propanol (C3H7OH) and ethanol (C2H5OH).
Most methanol is produced from natural gas. However, it is also possible to produce
methanol from biomass. Biomass includes crop and forest produces and even some
municipal wastes. This is attractive because crops are renewable. Ethanol is another
alcohol fuel that is currently made from biomass. It involves fermenting the sugar from
sugar and starch crops. Alcohol fermentation is the conversion of glucose (C6H12O6) to
two molecules of ethanol and two molecules of carbon dioxide. With advanced
technology being developed, cellulosic biomass, like trees and grasses, are also used as
feedstocks for ethanol production. The idea of using ethanol as fuel is not new. In fact,
Henry Ford’s first blueprint for the Model-T required ethanol fuel! Different biomass
produces varying amounts of ethanol.
WHAT IS THE ENERGY CONTENT OF VARIOUS FUELS?
You be heating some water by burning a measured amount of a fuel sample. It takes
exactly 1 calorie (cal) of heat to raise the temperature of 1 gram of liquid water by 1 C.
This is called the “specific heat” of water. Therefore, if you know the mass of water and
how many degrees the temperature went up, the total amount of heat absorbed by the
water can be calculated as follows:
Heat absorbed (calories) = mass of water (grams) x 1.00 cal/g-C x temp. change (C)
= m x specific heat x T
T is a shorthand way of saying “change of temperature”. The units of specific heat
(cal/g-C) make the units agree on both sides of the equation. (You should be able to
combine and cancel units on the right side of the equation to leave only calories.)
Theoretically, the amount of heat released by the burning fuel should equal the heat
absorbed by the water; however, in practice some of the heat will be lost to the
surroundings. Careful experimental technique will reduce the heat lost.
OVERVIEW OF EXPERIMENT
1. Assemble the apparatus.
2. Obtain burners containing known fuels.
3. Add a measured volume of water to the can, then determine the mass of the water.
4. Record the initial temperature of the water.
5. Weigh the burner.
6. Light the burner and heat the water until the water increases by about 20C.
7. Measure the highest temperature of the water.
8. Weigh the burner again, to determine the mass of fuel burned.
9. Calculate the amount of heat absorbed by the water and the amount of heat released
per gram of fuel burned.
MATERIALS NEEDED
Chemicals
One or more alcohol lamps containing liquid
fuels (to be assigned by the instructor)
Optional: non-alcohol heat source
One or more biomass fuel sources
Equipment
Empty soda can without a top
Thermometer
Ring stand with iron ring attached
Glass rod
100 ml graduated cylinder
Matches
Top-loading laboratory balance
Mineral oil
Erlenmeyer flask
Note about weighing: All weighings will be done on a laboratory balance that you will be
sharing with other students. Your instructor will explain the use of the particular balances
in your lab. For each weighing it is important to know that the balance reads exactly zero
with nothing on the balance pan. Check the “zero” each time you make a weighing.
EXPERIMENTAL PROCEDURE
The following general procedure is to be used for each measurement. Your instructor will
tell you which fuel you are to investigate. You’ll be asked to do several trials with a
single fuel to improve the level of your confidence about an average value for that
particular fuel. Then you will place your findings on the board where the rest of the
class’s findings will be assembled for comparison.
1. Obtain a dry soda can with the top removed and two holes punched on opposite sides
near the top. Slide a glass rod through the holes in the can so that it can be suspended
from the ring attached to a ring stand.
2. Obtain a fuel burner, place it under the can, and adjust the height of the ring so that
the bottom of the can is about 2 centimeters above the top of the wick.
3. Take the empty soda can plus your data sheet to a balance. Check to be sure that the
empty balance reads 0.0 gram; then weigh the can and record the mass to the nearest
0.1 gram.
4. Using a graduated cylinder, add approximately 100 ml of water to the can. Re-weigh
the can plus water to the nearest 0.1 gram. By subtraction, determine the mass of
water in the can.
5. Put a thermometer into the can, stir the water for a few moments, then measure the
temperature of the water, trying to estimate to the nearest 0.1 degree Celsius. (The
thermometer is probably marked in 1-degree intervals, so it is not possible to measure
the temperature accurately to a tenth of a degree. Nevertheless, it is useful to make
the best estimate that you can.)
6. Take the fuel burner plus your data sheet to a balance. If necessary, check to be sure
that the empty balance reads zero. Weigh the burner and record its mass on the data
sheet, reporting is to the nearest 0.01 gram (or nearest 0.001 g, if the balance permits
this).
7. Place the fuel burner under the soda can and light the burner. Observe the flame. If
necessary, cautiously adjust the height of the can so that the top of the flame is just
below the bottom of the can.
8. Stir the water occasionally and continue heating the water until the temperature has
increased about 20 degrees Celsius, then extinguish the flame. Cap the burner
quickly, to limit evaporation of fuel. Do the next two steps quickly.
9. Continue stirring the water gently until the temperature stops rising; record the
highest temperature, again estimating it to the nearest 0.1 degree. Calculate the
temperature change by subtracting the initial temperature from the final (highest)
temperature.
10. Take the burner and your data sheet to a balance. If necessary, check that the balance
reads zero and then weigh the burner and record the mass to the nearest 0.01 (or
0.001) gram. Calculate the mass of fuel burned by subtracting the final weight of the
burner from the original weight of the burner.
11. Before doing another measurement, take a few moments to discuss the procedure
with your partners. Did you encounter any difficulties? Can you think of any
desirable improvements in the procedure?
12. Now repeat the procedure, either making another measurement with the same fuel or
switching to a different fuel, as directed. If the water in the can has cooled to nearly
room temperature, you can use the same water sample. If not, it is a good idea to
empty the can and add a fresh 100 ml portion of water. Note that you recorded the
mass of the dry empty can before you started. This will not change, except for the
possible build-up of soot on the bottom; therefore, you need only to measure the mass
of the can with water in it.
CALCULATIONS
To calculate the amount of heat liberated by 1 gram of burning fuel, you need the
following items of information from each trial, all of which should be on the data sheet:
(a) the mass of water that was heated, (b) the changes in temperature of the water, and (c)
the mass of fuel that was burned.
Since it takes 1.00 calorie of energy to raise the temperature of 1 gram of water 1 degree
Celsius, the total heat absorbed by the water is:
Heat absorbed = degrees heated x grams of water x 1.00 cal/g-C
1.
2.
3.
For each trial, calculate the total calories of heat absorbed by the water. This will be
assumed to be the same as the amount of heat liberated by the burning fuel.
Then calculate the calories of heat per gram of fuel burned.
Finally, convert the answers from calories to joules, so as to permit comparison of
your results with data in the text, which uses joules per gram or per mole. To do this
you need to know that 1 calorie = 4.18 joules, and you will need to know the
molecular weight of your fuel(s).
Data Sheet- Energy Content of Fuels
Fuel Used
Trial Number
Mass of can + water
(g)
Mass of empty can
(g)
Mass of water
(g)
Final temp. of water
(C)
Initial temp. of water
(C)
Temperature change
(C)
Initial mass of burner
(g)
Final mass of burner
(g)
Mass of fuel burned
(g)
Heat absorbed by water
(cal.)
Heat per gram of fuel
(cal/g)
Heat per gram of fuel
(joules/g)
Heat of combustion
(kJ/mol)
Ethanol
Trial 1
Ethanol
Trial 2
Propanol
Trial 1
Propanol
Trial 2
QUESTIONS
1.
2.
3.
4.
5.
6.
What relationships or patterns do you see among the results for the different fuels
(heat per gram of fuel)? Does the amount of oxygen in the fuel appear to make a
difference? If so, why to you think this is the case?
There has been some interest on the part of midwestern grain farmers to promote
ethyl alcohol (ethanol) as a fuel alternative to gasoline, or to blend with gasoline.
Based on this experiment, what sort of problems and/or benefits may accrue from
this substitution?
Suppose a mixture that was 95% alcohol and 5% water was used instead of pure
alcohol. How would the results be affected?
In making a series of measurements on fuels, you had the choice of either keeping
the same water in the can, or adding fresh cold water each time. Suppose you kept
the same water and it did not cool significantly before the second measurement. Do
you think this would affect the results? If so, explain in what way they might
change.
Suppose you put 50 ml of water in the can instead of 100 ml. In what ways, if any,
would this affect the results of the experiment? What would happen if you used 200
ml of water? Can you think of any advantages or disadvantages in using either 50
ml or 200 ml of water in this experiment?
There are many possible sources of error in this experiment. List three that you can
think of. Would each error have a large effect, a medium effect, or a small effect on
the calculated heat content of a fuel?
DELIVERABLES
The report for this experiment will be a full report. For guidance, refer to the basic
laboratory report format provided on the Blackboard site and the ISAT Style Manual for
laboratory reports (see link on Blackboard). One report, jointly prepared by all members
of your team, will be due in two weeks.