Experiment 30C
FV 2/5/14
CHEMICAL CHANGE AND ENERGY:
What Fuel Makes the Best Energy Source?
MATERIALS: Week 1: 12 oz. Styrofoam cup with lid, thermometer, 100 mL graduated cylinder, 1.0 M HCl,
1.0 M NaOH, stir motor, stir bar.
Week 2: 12-oz. aluminum beverage can with top cut out and holes on side, thermometer, 100 mL
graduated cylinder, 800 mL beaker, long-stem lighter, three fuel burners (filled with ethanol, n-octane,
or 2-pentanol), steel wool, glass rod, ring stand, rubber cork, paper clip, room-temperature water.
PURPOSE: The purpose of this experiment is to learn how chemical reactions transfer heat and to determine which
fuels transfer the most heat during reaction with oxygen.
LEARNING OBJECTIVES: By the end of this experiment, the student should be able to demonstrate the
following proficiencies:
1. Construct and use a calorimeter.
2. Calculate the heat of a reaction from calorimetry data.
3. Calculate the fuel value for several fuels.
4. Compare the fuel value of an oxygenated and non-oxygenated fuel and explain differences in the values.
DISCUSSION:
Energy is essential for life. We fill our cars with energy-rich fuels so that they can transport us across town.
We burn fuel to create electricity to power the electronic gadgets on which we depend. We use fuel to heat the
buildings where we live and work. It is also the essence of life; we eat foods made up of carbohydrates, proteins and
fats to power our bodies.
Most of the fuels used today for power and heat are derived from petroleum oil, often referred to as crude
oil, but these have limitations: they are non-renewable resources, and their combustion produces large amounts of
pollution. Because of these problems, alternatives to petrol-based fuels are being developed, and their overall
effectiveness is being investigated. Each alternative chemical fuel (1) has a different chemical composition and
structure, (2) provides different amounts of energy, (3) produces different types and amounts of pollution, and (4)
costs a different amount.
How do we decide what is the best possible energy source when we account for all of these aspects? This
is one of the questions that you will investigate during the next two weeks as you explore the topic of chemical
change and energy by investigating alternative fuel sources that produce energy via chemical changes. You will
look at the energy content of different fuel sources as well as other issues that determine whether a fuel is a viable
energy source.
During the first week of this laboratory module, you will conduct calorimetry experiments to measure the
energy changes involved in the reaction between NaOH (aq) and HCl (aq). In the second week, you will compare
the combustion behaviors of different chemical compounds to determine which compound has the highest energy
content per gram. Using the information from this module, you should be able to construct an argument to address
the question “What fuel makes the best energy source?”
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WEEK 1:
How Do Chemical Reactions Transfer Heat?
The neutralization of an acid with a base in aqueous solution is an exothermic process. The heat transferred in
this process (qrxn) can be measured using a calorimeter. In this experiment, we will assume that no heat escapes or is
absorbed by the coffee cup. Making this assumption, the energy transferred to the liquid inside the cup approximates the
total energy released by the reaction. Therefore the heat of the calorimeter (qcal) approximates the heat transferred to the
liquid (qsoln) and it is the negative of qrxn.
qcal ≈ qsoln ≈ – qrxn
The heat evolved by the neutralization reaction is transferred to the solution and can be measured as a
temperature change where:
– qrxn = qsoln = msoln · ssoln · Tsoln
We will assume that the specific heat of the solution is the same as that of the solvent (water), therefore, ssoln = 4.184 J/g °C
and msoln is the mass of the solution and Tsoln = Tfinal soln – Tinitial soln. Tfinal soln is the maximum temperature reading on the
thermometer during the reaction.
The enthalpy change of the reaction (ΔHrxn) in kJ/mol can be found using:
qrxn = n ΔHrxn
where n is the number of moles of reaction that occurs.
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PROCEDURE:
Energy Changes During an Acid Base Reaction
1. Rinse the graduated cylinder with water and then rinse it with about 10 mL of 1.0 M NaOH (coat the sides by turning
the graduated cylinder). Discard the rinses. With the graduated cylinder, place 50.0 mL of 1.0 M NaOH into the
calorimeter. Measure the temperature (in °C) of the NaOH solution using the thermometer provided. Record this value
in the Data Section. Rinse the thermometer end after obtaining the temperature.
2. Rinse the graduated cylinder with water and then rinse with 1.0 M HCl. Discard the rinses. Measure 50.0 mL of
1.0 M HCl into the graduated cylinder. Measure the temperature of the HCl solution in the graduated cylinder. Record
this value in the Data Section. The initial temperature, Ti, will be defined as the average temperature of the two
solutions. Record this average temperature in the Data Section.
3. Note the time (time = zero). Then add the HCl solution quickly to the NaOH solution in the calorimeter, swirl the
solution then add the stir bar, and replace the lid on the calorimeter. Add the thermometer making sure it does not touch
the stir bar. You can use a clamp to hold the thermometer at the proper height (in the solution but not touching the stir
bar or sides of the calorimeter). Start the stir motor (on low) to gently mix the solution. Immediately begin collecting and
recording temperature readings at 30-second intervals for 7 minutes. Keep the lid on the calorimeter.
4. When data collection is completed, rinse the calorimeter and the thermometer with distilled water and dry as
completely as possible.
Clean up:
1.
All solutions are dilute aqueous solutions and may be poured down the drain.
2.
Wash all glassware. Rinse your calorimeter well and invert it on a paper towel to dry.
3.
Return all equipment to their original location.
4.
Turn OFF the thermometer.
WEEK 1 DATA SECTION
Energy Changes During an Acid Base Reaction
Initial temperature of the NaOH solution:
_____________
Initial temperature of the HCl solution:
_____________
Average of the two temperatures:
_____________ = Ti
Time (minutes)
Temperature (°C)
Time (minutes)
0.5
4.0
1.0
4.5
1.5
5.0
2.0
5.5
2.5
6.0
3.0
6.5
3.5
7.0
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Temperature (°C)
Week 1 Data Analysis:
1. What was the maximum temperature you recorded after starting the acid-base reaction?
Tf = ________
2. Calculate qrxn (in J) and ΔHrxn (in kJ/mol) for the reaction between HCl (aq) and NaOH (aq). Show your work.
3. Think about the qualitative observations you have made and the quantitative data that you have collected for the
reaction between NaOH (aq) with HCl (aq).
a. Write the molecular and net ionic equations for the chemical reaction that occurred.
b. Draw a molecular-level picture of the reactant and product species, and describe what is happening in the reaction
from a molecular-level perspective.
c. Was the reaction endothermic or exothermic? Explain using your observations.
d. How do you think the changes you observed/measured on the macroscopic-level relate to changes that you
proposed in your molecular-level picture in b. above?
E30C-4
WEEK 2:
Which Fuels Transfer the Most Heat During Reaction with Oxygen?
DISCUSSION:
Fuels
Combustion reactions are utilized in converting stored chemical energy into other forms of energy. Although mechanical
and biological systems are quite different, they both utilize combustion.
Combustion of a hydrocarbon produces carbon dioxide and water. This reaction releases energy (exothermic reaction).
Mechanical systems utilize this energy to do work. Two specific combustion reactions are shown below.
Natural gas, methane:
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (l)
Hcomb = -890.3 kJ/mol CH4
C3H8 (l) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
Hcomb = -2200 kJ/mol C3H8
Propane:
Carbohydrates and fats are examples of biological fuels (food). Although these are not hydrocarbons since they contain
oxygen, they both undergo the same type of combustion reactions.
Glucose (simple carbohydrate):
C6H12O6 (s) + 6 O2 (g)  6 CO2 (g) + 6 H2O (l) Hcomb = -2816 kJ/mol C6H12O6
Glycerol tristearate (fat):
2 C57H110O6 (l) + 163 O2 (g)  114 CO2 (g) + 110 H2O (l)
Hcomb = -75,520 kJ/mol C57H110O6
Comparing Fuels
Comparing fuels can be difficult. Enthalpy of combustion ( Hcomb) are given in the previous examples, but these depend
on the moles of CO2 and H2O formed. Therefore, when comparing fuels it may be more useful to compare the energy
content or fuel value of each fuel. Fuel value (kJ/g) is defined as the amount of energy released per gram of fuel. The
fuel value for methane is 55.5 kJ/g while that of glucose is 15.6 kJ/g.
In this lab, the heat of combustion can be measured by a constant pressure calorimeter. The amount of fuel burned can be
determined by difference in mass. Both of these measurements will be used to get the fuel value.
Calorimeter Efficiency
The efficiency of the calorimeter must be determined prior to calculating fuel values. By burning n moles of ethanol with
a known fuel value, the qcombustion can be calculated.
qcombustion = n · Hcombustion
(1)
The qwater is determined from the mass of the water, it’s specific heat, and the difference in temperature.
qwater = m · s · 
(2)
In this experiment it will become evident that all of the heat released by the burning of the fuels is not transferred to the
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water. The ratio of the heat absorbed by the water to the heat from the combustion of the fuel is the efficiency of the
calorimeter.
(3)
Efficiency = - qwater
qcombustion
Once the efficiency of the calorimeter has been determined, the qwater can be measured for other fuels, and then the
qcombustion of each fuel is calculated.
(4)
qcombustion = - qwater
Efficiency
The ratio of the qcombustion to the mass of fuel used is the fuel value (kJ/g).
Fuel value =
qcombustion
mass of fuel used
(5)
PROCEDURE:
SAFETY: The fuels used in this experiment are very flammable and care must be taken to avoid spillage.
Part A. Determining the Efficiency of the Calorimeter
1.
Light the ethanol burner. If the flame is not an inch or less in height, extinguish the flame by replacing the
burner cap. Adjust the height of the wick and recheck the flame. Once the wick is adjusted properly, extinguish
the flame by recapping.
2.
Weigh the ethanol burner with cap on a top-loading balance and record the mass. Place the capped burner in an
800 mL beaker.
3.
Use steel wool to clean your aluminum can if it is sooty. Gently push a glass rod through the pre-drilled
holes in the can. Set up an iron ring on a ring stand to suspend the can assembly as in the figure. Adjust
the height such that the bottom of the can is approximately an inch above the burner. AFTER THIS
ADJUSTMENT DO NOT CHANGE THE HEIGHT OF THE RING.
4.
Using a graduated cylinder, place 100 mL of water into the
aluminum can. Record the temperature of the water.
5.
Lift out the can, remove the burner cap, and then light the burner
and replace the can as quickly as possible. Stir the water with the
thermometer. (Don’t just leave it sitting on the bottom!)
6.
When the water temperature is about 40°C above its initial
temperature, remove the can and quickly cap the burner using
tongs. Re-suspend the can and keep stirring the water; record the
highest temperature the water reaches.
7.
Weigh and record the mass of the ethanol burner and cap.
8.
Pour the water out of the can. If the can is sooty, clean it with steel
wool. Readjust the wick if necessary.
9.
Repeat the measurement.
Part B. Determining the Fuel Value
1.
Repeat Part A using the n-octane burner and the 2-pentanol burner.
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Name __________________________________
Section ____________
Partner __________________________________
Date ______________
WEEK 2 DATA SECTION
Experiment 30C
Part A. Determining the Efficiency of the Calorimeter using Ethanol
Ethanol Burner
Trial 1
Trial 2
Final mass of burner and cap, g
Initial mass of burner and cap, g
Mass of ethanol combusted, g
Final temperature of water, oC
Initial temperature of water, oC
T, oC
Observations about the flame (i.e. sooty,
clean, etc.)
Part B. Determining the Fuel Value of n-octane and 2-pentanol
n-octane Burner
Trial 1
Trial 2
Trial 1
Trial 2
Final mass of burner and cap, g
Initial mass of burner and cap, g
Mass of n-octane combusted, g
Final temperature of water, oC
Initial temperature of water, oC
T, oC
Observations about the flame (i.e. sooty,
clean, etc.)
2-pentanol Burner
Final mass of burner and cap, g
Initial mass of burner and cap, g
Mass of 2-pentanol combusted, g
Final temperature of water, oC
Initial temperature of water, oC
T, oC
Observations about the flame (i.e. sooty,
clean, etc.)
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WEEK 2 DATA TREATMENT
Experiment 30C
Part A. Determining the Efficiency of the Calorimeter, using Ethanol
(A.1)
Using the relationship qwater = specific heat  mass  T, show the calculation for qwater for Trial 1.
(A.2) Using the fuel value for ethanol determined in your prelab and the mass of ethanol combusted, show the
calculation for qcombustion for Trial 1.
(A.3)
Show the calculation for the efficiency of energy transferred for Trial 1.
(A.4)
Calculate qwater, qcombustion, and efficiency for your other trial(s) and summarize the results.
Ethanol
Trial 1
Trial 2
qwater, kJ
qcombustion, kJ
efficiency
Average efficiency
NOTE: This average efficiency will be used to calculate the fuel value of other fuels.
Part B. Determining the Fuel Value of n-octane
(B.1)
Calculate the qwater for each n-octane trial. Record in the table on next page.
(B.2)
Use the average efficiency for the ethanol burned and qwater for each n-octane trial to calculate the qcombustion for
each n-octane trial. Record in the table on next page.
(B.3)
Use the qcombustion and mass of n-octane combusted to calculate the fuel value for each n-octane trial and
summarize the results.
n-octane
Trial 1
Trial 2
qwater, kJ
qcombustion, kJ
Fuel value, kJ/g
Average fuel value
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Part B. Determining the Fuel Value of pentanol
(B.4)
Calculate the qwater for each 2-pentanol trial and record in the table below.
(B.5)
Use the average efficiency for the ethanol burned and qwater for each 2-pentanol trial to calculate the
qcombustion for each 2-pentanol trial. Record in the table.
(B.6)
Use the qcombustion and mass of 2-pentanol combusted to calculate the fuel value for each 2-pentanol trial.
Record in the table.
pentanol
qwater, kJ
qcombustion, kJ
Fuel value, kJ/g
Average fuel value
Trial 1
Trial 2
Mass Percent of Oxygen in each Fuel
(B.7)
Calculate the mass percent of oxygen in ethanol (C2H5OH), n-octane (C8H18), and 2-pentanol
(C5H11OH).
Week 2 Data Analysis:
1.
Write a balanced chemical reaction for the reaction of ethanol with oxygen.
2.
Draw a molecular-level picture of the reactant and product species, and describe what is happening in the
reaction from a molecular-level perspective.
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3.
Was the reaction endothermic or exothermic? Explain using your observations.
4.
How do you think the changes you observed/measured on the macroscopic-level relate to changes that you
proposed in your molecular-level picture in 2. above?
5.
Compare your visual observations about the combustion of the three different fuels. Consider the
characteristics of the flame and the amount of soot produced on the can.
6.
Oxygenated fuels (compounds containing C, H, and O) are used as motor vehicle fuels or fuel additives because
they burn cleaner, thereby reducing air pollution. They also affect the miles per gallon. Compare the fuel value
and mass percent oxygen for each fuel, and decide if oxygenation results in an increase or decrease in miles per
gallon. Explain why.
7.
How do the sets of observations and data from this lab affect your understanding of what makes a good
fuel?
E30C-10
Name _____________________________
Section _______________________
Date _________________________
WEEK 1
PRE-LAB QUESTIONS
Experiment 30C
1. When 100 mL of 0.500 M HCl is mixed with 100 mL of 0.500 M NaOH in a coffee cup calorimeter, the final
temperature of the mixture is measured to be 25.86 °C. The initial temperature of both solutions was 22.50 °C.
a. Assuming that the densities and specific heats of the solutions are the same as for water (1.00 g/mL and
4.184 J/g °C), calculate the heat released by this reaction in kJ.
b. Report your answer from part a. as an energy per mole value (kJ/mol HCl neutralized).
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Name __________________________________
Section ______________________
Date ________________________
WEEK 2
PRE-LAB EXERCISES
Experiment 30C
1. In this experiment a liquid fuel burner is used to heat water in an aluminum beverage can. Not all of the energy from
the burner is transferred to the water, i.e., qcombustion + qwater  0. Because of this, the efficiency of the energy transfer must
be determined using the relationship:
efficiency 
 q water
q combustion
a. In this experiment what will be the sign of qwater?
b. What is the sign of qcombustion?
c. Why is the negative sign included with the qwater?
2. Write a balanced chemical equation for the complete combustion of one mole of liquid ethanol, C2H5OH. (Assume the
water produced is in the liquid state.)
3. Using standard enthalpy of formation values found in Appendix 3 of your textbook, calculate H for this reaction.
4. The fuel value of a substance is kilojoules of energy released per gram of fuel burned. Calculate the fuel value for
ethanol.
E30C-12