Measuring the Steady State Thermal Efficiency of a Heating System. Last week you had a tour of “Colby Power and Light” and had a chance to see Gus Libby’s really cool, campus-scale, multi-million dollar calorimeter. This week we will use lab-scale calorimeters to measure the heat and carbon content of different fuels (methane, propane, butane, and gasoline). Our simple calorimeters were designed by Whitney King and Chuck Jones, Science Division Instrument Wizard, to simulate the highly efficient HTP Versa Flame fire-tube boiler. The lab has three parts: 1) Calibrate the calorimeter 2) Measure the heat content of different fuels 3) Compare the measured heat content to the calculated heat content 1) Calibrate the Calorimeter. You lab instructor will demonstrate the assembly of the laboratory- scale fire-tube calorimeter. The heat transferred from the burning fuel to the calorimeter can be determined from the mass of the object being heated, the heat capacity of the object and the change in temperature.
� = ���� ∗ �� ∗ ∆� We will begin the experiment by adding a known amount of heat to our calorimeters from a coffee cup heater. The heater is a simple resistor that converts electrical energy into heat. The heater has a resistance of 47.5 ohms and is plugged into a 120 V outlet. Recalling your electricity basics, we know that V=IR where V is the voltage in Figure 1. HTP Versa Flame Fire Tube volts, I is the current in amps, and R is the Boiler. Brown Tubes carry the hot resistance in ohms. The power, P, consumed by combustion gasses inside of the the heater is volts times current, water tank where they condense releasing the maximum available 2 heat. P = V*I = V /R (Watts or J/s) www.htproducts.com/versaflame.html and all of this power will enter the calorimeter as heat. Total heat is simply power times time. Q = P*time = V2*time/R (Joules) We will provide ohm meters for the lab so that you can measure the resistance of your coffee cup heater. If you place the coffee cup heater into the calorimeter for about five minutes it should raise the temperature of the calorimeter about ten degrees. The actual total heat capacity (CT = mass*Cp) of the calorimeter is calculated from
Heat Content of Fuels 2015 1 2 V *time/R = Q = CT*ΔT where we measure the temperature change carefully using a digital thermometer. You should get a value of around 17,000 J/K.
Procedure: 1) Prepare your calorimeter as described by you lab instructor (take pictures for your notebook). 2) Measure the resistance of the heater in ohms and compute the heater power. 3) Turn on the vacuum line connected to the chimney, start logging temperature data, and plug in the heater. 4) Allow the temperature of the calorimeter to increase about 10 degrees K. 5) Unplug the heater. 6) Stop logging temperature data when the final temperature is Figure 2 Lab Calorimeter constant. 7) Record the starting and ending temperatures and the starting and ending times.
8) Compute CT. 2) Measure the Heat Content of a Fuel. Now that you have a calibrated calorimeter we can use it to measure the heat content of fuels. Instead of heating the calorimeter with an electric heater we will heat the calorimeter with a gas flame and try to capture as much of the energy from combustion as possible. Procedure: 1) Prepare your calorimeter as described by you lab instructor (take pictures for your notebook) 2) Weigh the mass of your propane, butane, or gasoline fuel source (stove + fuel) 3) Place the stove under the calorimeter, turn on the vacuum line connected to
the chimney, start logging Figure 3. Top view of Fire Tube temperature data, and light the
Heat Content of Fuels 2015 2 stove. 4) Allow the temperature of the calorimeter to increase about 10 degrees K. 5) Stop logging temperature data when the final temperature is constant. 6) Record the starting and ending temperatures and the starting and ending times. 7) Weigh the stove and compute the mass of fuel burned in grams and pounds.
8) Calculate Q using CT and ΔT for your fuel. 9) Compute the heating efficiency of your fuel and calorimeter by comparing your results to the values listed in Table 1. You will need to perform some unit conversions to make a sensible comparison. Welcome the real world! Should you use the Gross or Net heating values in the table?
Table 1. Energy Content of Fuels. www.engineeringtoolbox.com/heating- values-fuel-gases-d_823.htm
Heat Content of Fuels 2015 3 3) Calculating the heat content of fuels. Under conditions of constant pressure and with work restricted to PV work, the heat produced by a chemical reaction is equal to the enthalpy change of the reaction, ΔH. If a mole of fuel is combusted the heat of the reaction is equal to the molar enthalpy change. Enthalpies of combustion are measured experimentally (Table 1) or may be calculated in several different ways. One way to think of heats of combustion is to consider that chemical reactions are simply the reorganization of the reactants to produce new products. In other words, heats of reaction are the energy of breaking all the bonds in the reactants and then recovering energy by forming new bonds in the products. To get the correct sign for the enthalpy change of the system (reaction) we calculate enthalpy change of a reaction as the energy of all bonds broken minus the energy of bonds formed. Considering the combustion of a
CH4 (g) + 2O2 (g)à CO2 (g) + 2H2O (g) mole of methane we would need to break four moles of C-H bonds in CH4, break two moles of O=O double bonds in two moles of oxygen gas, form two moles of C=O double bonds in CO2, and form four moles of H-O bonds in two moles of H2O. Notice that I use the balanced chemical reaction to calculate the numbers of bonds broken and formed. The following table illustrates the bond accounting. Bond Moles of Energy total broken Bonds kJ/mol kJ/mol reactant C-C bonds 0 346 0 C-H bonds 4 411 1644 O=O 2 498 996
formed C=O 2 805 1610 O-H 4 464 1856
Sum bonds broken - Sum bonds formed - 826
The calculated heat of combustion, -826 kJ/mol, is close to the experimental value obtained using a calorimeter of, -798 kJ/mol.
Another way to calculate heats of reaction is take the heats of formation, fΔH, of the products minus the reactants of the combustion reaction. The heat of formation is the heat given off when a mole of molecules is formed from pure elements at room temperature, Table 2. This is really the same as the calculation above except we break each molecule down to pure stable elements instead of atoms (O2 instead of O, H2 instead of H) and then rebuild the products from the pure elements. For example, the reaction for combustion of methane to carbon dioxide and steam is
CH4 (g) + 2O2 (g)à CO2 (g) + 2H2O (g)
Heat Content of Fuels 2015 4 where the molar enthalpy of the reaction is the heat of formation of products minus reactants.
ΔH combustion = fΔH (CO2) + 2 fΔH (H2O (g)) - fΔH (CH4) - 2 fΔH (O2)
ΔH combustion = −393 + 2∗(−241.8) − (−74.9) − 2∗(0) = −801.7 kJ/mol Notice that this value is in very good agreement with the experimental value (-798) because heats of formation are measured experimentally and are specific to each product and reactant while bond energies are averages. The heats of formation change depending on the phase of the molecule. If water is condensed to a liquid it releases heat so the fΔH (H2O (g) > fΔH (H2O (l)). Making the change to liquid water causes
ΔH combustion = −393 + 2∗(−285.8) − (−74.9) − 2∗(0) = −889.7 kJ/mol the enthalpy of the reaction increases (becomes more negative) because the condensation of water releases more heat. The heat of reaction of fuels to produce liquid water is called the gross heat of combustion and the heat of reaction of fuels to produce steam is called the net heat of combustion. The difference is about 10% depending on the amount of water produced. Boilers that condense the water produced during combustion are about 10% more efficient then conventional boilers but must operate at lower temperatures – typically less than 140 oF. Table 2. Heats of Formation
Formula Name fΔH kJ/mol
H2 Hydrogen 0.0
CO2 Carbon Dioxide -393.5
O2 Oxygen 0.0
H2O Water (g) -241.8
H2O Water (l) -285.8
CH4 Methane -74.9
C2H6 Ethane -83.7
C2H2 Acetylene 226.8
C3H8 n-Propane -104.6
C4H10 n-Butane -125.5
C5H12 n-Pentane -146.9
C6H14 n-Hexane -167.4
C7H16 n-Heptane -187.9
C8H18 n-Octane -208.4
C9H20 n-Nonane -229.3
C10H22 n-Decane -249.4
Heat Content of Fuels 2015 5 Procedure:
1) Write the balanced combustion reactions for hydrogen (H2), CH4 (methane), C3H8 (propane), C4H10 (butane), and C8H18 (octane or “gasoline”). 2) Make ball-and-stick models of the first four fuels and oxygen. Now reorganize the atoms and bonds (sticks) to make the combustion products. 3) Determine the moles of bonds broken and formed for each combustion reaction. 4) Calculate the enthalpy of combustion of each fuel using the energy of bonds broken minus the energy of bonds formed. Use and Excel table like the one shown above to help organize your calculations. 5) Calculate the enthalpy of combustion of each fuel using the enthalpies of formation. 6) Compare both calculations to the values in Table 2. You will need to do some unit conversions to make the comparisons. Since this is a practical course, lets do the comparison using units of BTU/pound of fuel. (see conversions at the bottom of Table 2) 7) Compute the pounds of carbon dioxide produced per BTU (kJ) of heat content for each fuel. Lab Report: 1) In an Excel spreadsheet show your calculations for determining the efficiency of you calorimeter. Compare your results with the other students in the lab. What features of the calorimeter appear to be most important in maximizing efficiency? 2) In a second Excel spreadsheet show your results for part III of the lab. Which fuel is the most economical and which fuel will produce the least carbon dioxide to heat a home? Current fuel prices: http://maine.gov/energy/fuel_prices/index.shtml
Extra Credit: A lot has been written about the hydrogen economy and the advantages of carbon free, high energy content hydrogen. However, the most cost effective industrial method to produce hydrogen gas is a reaction called steam reforming.
Stage 1. CH4 + H2O + heat → CO + 3H2
Stage 2. CO + H2O + heat → CO2 + H2
Net reaction CH4 + 2H2O → CO2 + 4H2
This reaction requires heat (from the combustion of fossil fuels) to proceed. How does the practical source of hydrogen change you opinion of this fuel?
Heat Content of Fuels 2015 6