Heat of combustion of paraffin wax lab answers

Energy Comparison of Fuels Hands-On Labs, Inc. Version 42-0154-00-02

Review the safety materials and wear goggles when working with chemicals. Read the entire exercise before you begin. Take time to organize the materials you will need and set aside a safe work space in which to complete the exercise.

Experiment Summary:

In this experiment, you will learn how various sources of fuels are formed. You will also assemble a calorimeter and measure the heat released from two different fuel sources. You will learn about the different types of fuels, combustion reactions, and differences between the energy content in fuel sources.

EXPERIMENT

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Learning Objectives Upon completion of this laboratory, you will be able to:

● Describe the process of combustion.

● List different sources of fuels.

● Classify hydrocarbon fuel sources as biofuels or fossil fuels.

● Describe the formation of fossil fuels: coal, petroleum, and natural gas.

● Classify hydrocarbon’s molecules as saturated or unsaturated.

● Describe how energy is measured in calories or joules.

● Diagram a commercial calorimeter and describe the function of its components.

● Calculate the number of calories released per gram of fuel.

● Perform experiments with a homemade calorimeter.

● Calculate the density of water.

● Compare and contrast the energy content of diethylene glycol and paraffin.

Time Allocation: 2 hours

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Experiment Energy Content of Fuels

Materials Student Supplied Materials

Quantity Item Description 1 Bottle of distilled water 1 Clock or timer (with second hand) 1 Matches or lighter

HOL Supplied Materials

Quantity Item Description 2 Aluminium foil, 30” x 24” 1 Beaker, 250 mL, glass 1 Burner-fuel, diethylene glycol 1 Burner stand 1 Candles, tea 1 Goggles 1 Pair of safety gloves 1 Graduated cylinder, 50 mL 1 Ruler, metric 1 Scale, digital 1 Thermometer, analog

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software, such as Microsoft Word® or PowerPoint®, to add labels, leader lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

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Experiment Energy Content of Fuels

Background Combustion, Biofuels, and Fossil Fuels

Fuels are storehouses for energy which exist in many different forms. For example, food is a fuel that supplies energy for metabolism, wood is a fuel that is burned to create warmth, and radioactive materials are fuels that power submarines and electrical generators. However, energy stored in a fuel must first be extracted before it is able to perform work. For fuels such as coal and oil, energy is extracted through a process called combustion. Combustion is an exothermic process, in which the fuel reacts with an oxidizing agent, such as oxygen, to produce heat. In a combustion reaction, heat (energy) is liberated when the fuel reacts with an oxidizing agent, forming different chemical compounds. For example, methane (CH4) reacts with oxygen to form carbon dioxide (CO2), water (H2O), and heat. This combustion equation is shown below:

4 2 2 2CH + 2O CO + 2H O + Heat→ The most common source of fuel is hydrocarbons. A hydrocarbon is an organic compound composed primarily of hydrogen and carbon atoms. Fuel sources composed of hydrocarbons are categorized as either biofuels or fossil fuels. Biofuels are derived from living biological organisms (hence the name “bio”) and can exist as solids, liquids, or gases. Biofuels are generally derived from renewable carbon sources, such as wood, waste materials, or corn. Ethanol, which can be produced from corn or sugar cane, has recently become a popular biofuel. However, there is currently a debate in political, scientific, and economic circles regarding how much net energy is derived from these biofuels after taking into account the fuel required to plant, grow, harvest, and process the crops, then to produce and distribute the fuel.

Fossil fuels are derived from the fossilized remains of organisms that lived millions of years ago; thus, they are not considered to be renewable energy sources. See Figure 1. The use of fossil fuels (coal, petroleum, and natural gas) was not a common practice until after the Industrial revolution. Coal and petroleum, however, were used prior to the Industrial Revolution for a variety of purposes. For example, petroleum was used in medicines and to waterproof boats and coal was used to make jewelry.

Based on fossil fuels used worldwide, it has been approximated

in recent years that it took plant matter somewhere between 400-500

years to yield this energy on our ancient Earth!

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Experiment Energy Content of Fuels

Figure 1. How fossil fuels were formed. © U.S. Energy Information Administration

Energy in Fuel

Fuel sources differ in the amount of energy they contain. While determining the amount of energy in a fuel source can be complex, some generalizations can be made. First, longer-chain hydrocarbon molecules release more energy when they are combusted than shorter-chain molecules; second, even for the same energy source, the amount of energy can vary from batch to batch. Fossil fuels contain inconsistent amounts of carbon, hydrogen, sulfur, nitrogen and oxygen, and two fossil fuel samples may provide more or less energy, and more or less pollution, based on their chemical composition. For example, coal from the eastern United States contains more sulfur than coal from the western United States, so more sulfur dioxide is produced when eastern coal is burned. This has environmental importance, as sulfur dioxide is one of the main precursors of acid rain.

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Experiment Energy Content of Fuels

Hydrocarbon molecules are classified as either saturated or unsaturated, depending on whether the carbon atoms are joined by single bonds or multiple (double or triple) bonds, respectively. A carbon atom can form up to four single bonds. In a saturated hydrocarbon, there are single bonds between all of the carbon atoms, and each carbon atom is completely surrounded (“saturated”) by hydrogen atoms. Butane, C4H10, is an example of a saturated hydrocarbon. See Figure 2A. In an unsaturated hydrocarbon, at least one of the bonds between the carbon atoms is either a double or triple bond. See Figure 2B. A saturated hydrocarbon contains, and therefore releases, more energy after combustion than an unsaturated hydrocarbon with the same number of carbon atoms.

Figure 2. A. Butane, a saturated hydrocarbon, consists of four carbon atoms and ten hydrogen atoms. B. Unsaturated hydrocarbon: Ethene (more commonly known as ethylene) with a double bond. C. Unsatu-

rated hydrocarbon: Ethyne (more commonly known as acetylene) with a triple bond.

Diethylene Glycol and Paraffin

In the LabPaq kit, the fuel canister contains diethylene glycol and the candle contains paraffin. Diethylene glycol is an organic, partially oxidized (contains oxygen) compound. See Figure 3. Diethylene glycol is used in industry as a solvent, as well as in brake fluid, lubricants, and facial creams.

Figure 3. Formula of diethylene glycol.

Paraffin is another term for saturated hydrocarbons. Most paraffins exist as a mixture of different length hydrocarbons and are a commonly used fuel source. Paraffins are also used in crayons, foods, toiletries, and paints. Most candles used today are formed from paraffin wax, although some candles are made of beeswax, soy, or even tallow (rendered beef fat). The hydrocarbon chains in paraffin wax range from C20H42 to C40H82 . See Figure 4.

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Experiment Energy Content of Fuels

Figure 4. Formula of icosane (C20H42), a paraffin wax.

In this activity, you will burn two different fuels (diethylene glycol and paraffin) and compare the amount of energy released. Energy is measured in either calories or joules. One calorie is the amount of heat energy required to raise the temperature of one gram of water by one degree Celsius. The joule (J) is the Système International (SI) unit of energy. To convert from calories to joules, multiply the number of calories by 4.184, the number of joules in a calorie. The term “Calorie” (with a capital C) represents a kilocalorie (1000 calories) which is used when counting food calories.

Calorimeter

A calorimeter is a tool used to determine the number of calories of heat released from a material when it is combusted. See Figure 5. The energy from the fuel is transferred as heat to the water in a container, thereby raising the temperature of the water. The change in water temperature, as a result of the transferred heat, is used to determine the number of calories released by the fuel source.

Figure 5. Schematic of a commercial calorimeter. The motorized stirrer circulates the water located in the chamber, allowing for even heat distribution. The electrodes provide the charge (spark) to start burning the material. The thermometer records the temperature of water. The insulated container reduces the

loss of heat to the external environment, and the water surrounds the fuel container to absorb the heat from the burning material.

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Experiment Energy Content of Fuels

Burning hydrocarbons, whether from fossil fuel or biofuel sources, releases carbon dioxide (CO2), which is an important greenhouse gas. Numerous scientific research studies have determined that the increase in atmospheric CO2 concentrations, from approximately 270 parts per million (ppm) in pre-industrial times to 392 ppm in 2011, is contributing to climate change that is impacting the environment and human society. The types of fuels we use vary greatly in their production of carbon dioxide. Among the fossil fuels, coal releases the largest quantity of CO2 per unit of energy produced, and natural gas produces the least. Nuclear energy does not produce carbon dioxide during electricity generation, but it has other safety concerns; additionally, there are indirect CO2 emissions associated with nuclear energy. Wind energy and solar energy do not produce any greenhouse gases when producing electricity. However, the fabrication, distribution, and installation of wind turbines and solar panels still rely on energy from fossil fuels.

There have been large advances in the creation and

use of renewable energy sources. Vegetable oil is being used in place of gas to run cars; solar energy

panels are becoming more common in residential housing; and car

manufactures are making advances in both hybrid and electric

automobiles.

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Experiment Energy Content of Fuels

Exercise 1: Energy Content of Fuels In this experiment, you will make a simple calorimeter to determine the energy content in two fuels, diethylene glycol and paraffin wax. The energy content will be determined by burning each fuel and measuring the change in temperature of 200 mL of distilled water.

The following equation will be used to calculate the number of calories in each type of fuel:

p Q = T m c∆ × ×

● Q = amount of energy transferred to the water in calories (cal)

● ΔT = change in temperature (final temperature minus initial temperature) in ºC

● m = mass of distilled water (g)

● cp = specific heat capacity of water (1 cal / g × °C)

Note: The transfer of energy from the burning fuel to the water is not 100% efficient; some energy will be lost to the surrounding environment.

Calculating the Density (g/mL) of Water

Note: Under standard conditions (1 atm pressure and 25°C), 1 mL of water has a mass of 1 gram. It is unlikely that you will be performing this experiment at standard conditions, so it is necessary to calculate the density of water at your location.

1. Turn on the scale. Place the graduated cylinder on the scale and tare the scale so that it reads “0.0 g.”

2. Add exactly 10.0 mL of distilled water to the graduated cylinder and record the mass in Data Table 1 of your Lab Report Assistant.

3. Calculate the density of the water by dividing the mass of the water by the volume of the water and record in Data Table 1.

Determining the Energy Content of the Fuels

Fuel Canister

4. Use the graduated cylinder to measure and pour 200.0 mL of distilled water into the glass beaker.

5. Cut out two 20 cm x 20 cm squares of aluminum foil. Set 1 piece of foil aside (to be used in step 17). This will give you 3 pieces of foil: 2 pieces 20 x 20 cm, and 1 full sheet that was not cut.

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Experiment Energy Content of Fuels

6. Place the burner stand on one of the 20 cm x 20 cm squares of aluminum foil. See Figure 6.

Note: The burner stand is placed on a square of foil to protect the surface of the table.

7. Place the 250-mL beaker on the burner stand.

8. Place the thermometer in the water and allow it to equilibrate for 2 - 3 minutes. Record this initial temperature in Data Table 2 of your Lab Report Assistant, under “Initial temperature,” in the “Fuel canister” row.

9. Turn on and tare the digital scale, so that it reads “0.0 g.”

10. Place the fuel canister with the lid on the scale. Measure the mass in grams and record in Data Table 2 under “Initial mass of fuel” in the “Fuel canister” row.

11. Remove the cap from the canister and place the canister under the burner stand. Use the remaining large piece of aluminum foil (not the 20 cm x 20 cm piece put aside) to cover both the beaker (with the thermometer still in it) and the stand. Leave a little room open at the bottom to allow for air flow so that the fuel will have adequate oxygen to burn. See Figure 6.

Figure 6. Construction of calorimeter with foil canopy.

12. Carefully lift the foil off of the burner, and use a match or lighter to light the wick on the fuel canister. Replace the foil, and immediately start the timer.

Note: You may have to make adjustments to ensure that enough air can enter the foil canopy to allow the fuel to continue to burn.

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Experiment Energy Content of Fuels

13. Heat the water for 10 minutes and then carefully remove the foil and place the cap back on the fuel to extinguish the flame. Do not tighten the cap because it will be difficult to remove when it cools.

14. Quickly read the thermometer and record the temperature in Data Table 2 under “Final temperature” in the “Fuel canister” row.

15. Allow the beaker to cool in preparation for the next part of the experiment and then empty the water.

16. When the fuel canister has cooled enough to be safely handled, tare the scale and place the canister back on the digital scale to measure the final mass. Record the mass in Data Table 2 under “Final mass of fuel” in the “Fuel canister” row.

Tea Candle

17. The candle must be suspended from the bottom of the stand so that the flame is close to the wire mesh. Use the piece of foil that was set aside to fasten a hammock-like structure to the legs of the stand. Wrap each corner of the foil around the four legs on the stand, as shown in Figure 7. The candle should be about one inch from the bottom of the stand.

Figure 7. Construction of a “hammock” to suspend the candle.

18. Turn on and tare the digital scale so that it reads “0.0 g.”

19. Place the candle on the scale. Measure the mass in grams and record in Data Table 2 under “Initial mass of fuel” in the “Tea candle” row.

20. Place the candle in the hammock so it sits securely in place.

21. After the beaker has cooled, use the graduated cylinder to measure 200.0 mL of distilled water and pour the water into the beaker. Place the thermometer back into the water and

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Experiment Energy Content of Fuels

allow it to equilibrate for 2 – 3 minutes. Record the temperature in Data Table 2 under “Initial temperature” in the “Tea candle” row.

22. Place the beaker with the water on the burner stand.

23. Cover the experiment with the foil canopy. Adjust the canopy as needed, and make sure that it fits well, leaving a little room open at the bottom to allow for air flow so that the fuel will have adequate oxygen to burn.

24. Remove the foil canopy and use a match or lighter to light the wick of the candle while it is in the hammock. Replace the foil, leaving a small opening at the bottom so air can enter. Immediately start your timer.

25. Allow the candle to burn for 10 minutes.

26. After 10 minutes, carefully remove the foil, blow out the candle, and quickly read the thermometer. Record the temperature in Data Table 2 under “Final temperature” in the “Tea candle” row.

27. After the candle has cooled enough to be safely handled (the wax will have solidified again), disassemble the hammock and remove the candle.

28. Turn on the scale and tare it so the scale reads “0.0 g.”

29. Place the candle on the scale to measure its mass. Record the mass in Data Table 2 under “Final mass of fuel” in the “Tea candle” row.

30. Calculate the ΔT for each fuel type by subtracting the initial temperature from the final temperature and record in Data Table 3 of your Lab Report Assistant.

31. To calculate the number of calories of energy absorbed by the water (Q), use the equation:

● Use the density of water you determined in Data Table 1 to calculate the mass of 200.0 mL of water (multiply the density by 200).

● The specific heat capacity (cp) of water is 1.00 cal / g x °C.

32. Use the temperature change in the equation and calculate Q. Record the calories for both fuel types in Data Table 3 under “Calories absorbed by water.” For example, if the density of the water at your location was 0.95 g/mL and the temperature change was 20 °C, then your equation would be:

× × × ×o o

0.95g 1calQ = (20 C) (200mL ) = 3800cal mL g C

33. Using the mass data in Data Table 2, determine the total mass of fuel consumed by subtracting the final mass from the initial mass, and record the result in Data Table 3 under “Grams of fuel

p Q = T m c∆ × ×

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Experiment Energy Content of Fuels

consumed” for each of the fuel types.

34. For each fuel, divide the calories absorbed by the grams of fuel consumed to determine the amount of energy in calories released per gram of fuel. Record these values in Data Table 3.

35. When the beaker has cooled sufficiently, pour out the water, and clean up your work area.

36. Return all clean items to your LabPaq kit for future use.

Questions A. Which fuel was more efficient (produced more calories per gram)? Explain your answer.

B. Explain the difference in the efficiencies of the fuels you tested based on the chemical struc- tures of the fuels.

C. Conduct research on both diethylene glycol and paraffin wax; for each fuel, summarize in your own words in one or two sentences how these fuels are synthesized. Cite the sources that you use for your research.

D. What are the possible sources of error in this experiment? How could the errors be reduced in future experiments?

E. Research how different fossil fuels compare in terms of the amount of carbon dioxide released per kWh of energy. How much carbon dioxide is produced by burning a gallon of gasoline? A gallon of diesel? Cite the sources that you use for your research.

F. Compare the amount of carbon dioxide released in one year from burning coal to power 10, 65-watt incandescent bulbs with the amount released from powering 10, 13-watt compact fluorescent light (CFL) bulbs. Assume the bulbs are on four hours per day for 365 days. You will need to determine the kilowatt hours (kWh) used. First, multiply the wattage of the bulbs by the number of light bulbs to determine the total watts used in one hour. Then mul- tiply the result by time in hours to obtain the watt hours. Next, divide the result by 1000 to obtain kilowatt hours. On average, 2.1 pounds of carbon dioxide are released for every kWh of electricity produced.

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Experiment Energy Content of Fuels

Energy Comparison of Fuels Hands-On Labs, Inc. Version 42-42-0154-00-02

Lab Report Assistant This document is not meant to be a substitute for a formal laboratory report. The Lab Report Assistant is simply a summary of the experiment’s questions, diagrams if needed, and data tables that should be addressed in a formal lab report. The intent is to facilitate students’ writing of lab reports by providing this information in an editable file which can be sent to an instructor.

Exercise 1: Energy Content of Fuels Data Table 1. Mass and Density of Distilled Water.

Mass of 10.0 mL of distilled water

(g)

Density of water

(g/mL)

Data Table 2. Calorimetric Data.

Fuel Source: Initial mass of

fuel (g)

Final mass of fuel (g)

Initial temperature

(°C)

Final temper- ature (°C)

Fuel canister: (di- ethylene glycol)

Tea candle: (paraf- fin)

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Experiment Energy Content of Fuels

Data Table 3. Calories Released per Gram of Fuel.

Fuel Type ΔT (°C) Calories absorbed by water (cal) Grams of fuel consumed (g) Calories/gram

Fuel canister: (di- ethylene glycol)

Tea candle: (paraf- fin)

Questions A. Which fuel was more efficient (produced more calories per gram)? Explain your answer.

B. Explain the difference in the efficiencies of the fuels you tested based on the chemical struc- tures of the fuels.

C. Conduct research on both diethylene glycol and paraffin wax; for each fuel, summarize in your own words in one or two sentences how these fuels are synthesized. Cite the sources that you use for your research.

D. What are the possible sources of error in this experiment? How could the errors be reduced in future experiments?

E. Research how different fossil fuels compare in terms of the amount of carbon dioxide released per kWh of energy. How much carbon dioxide is produced by burning a gallon of gasoline? A gallon of diesel? Cite the sources that you use for your research.

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Experiment Energy Content of Fuels

F. Compare the amount of carbon dioxide released in one year from burning coal to power 10, 65-watt incandescent bulbs with the amount released from powering 10, 13-watt compact fluorescent light (CFL) bulbs. Assume the bulbs are on four hours per day for 365 days. You will need to determine the kilowatt hours (kWh) used. First, multiply the wattage of the bulbs by the number of light bulbs to determine the total watts used in one hour. Then mul- tiply the result by time in hours to obtain the watt hours. Next, divide the result by 1000 to obtain kilowatt hours. On average, 2.1 pounds of carbon dioxide are released for every kWh of electricity produced.