Materials Included In Kit

Additional Materials Required

Safety Precautions

Although the materials in this experiment are considered safe, wear chemical safety glasses and a chemical-resistant apron as good laboratory practice. Use heat-resistant gloves when handling hot glassware.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The tap water can go down the drain according to Flinn Suggested Disposal Method #26b. The Heat Transfer Kit materials should be dried and stored for future use.

Lab Hints

• Enough materials are included in this kit for one group of students. This laboratory activity can be reasonably completed in two 50-minute class periods.
• Additional Heat Transfer Kits (AP4533) may be purchased for each individual lab group.
• If a hot plate and/or ice are not available, hot and cold water directly from the faucet can be used. The initial temperature of the hot and cold water should be approximately the same for each experiment.
• Do not allow students to handle the hot beakers with their bare hands (Procedure step 7). Hot vessel gripping devices or heat-resistant gloves should be worn. If these are not available, double-folded paper towels wrapped around the beaker or kitchen hot pads are good substitutes. Beaker tongs can also be used. Students may have difficulty gripping the beaker when using heat-resistant materials, or beaker tongs, because their “feel” will be diminished. The proper gripping and pouring technique should be demonstrated to the students before the lab. If necessary, have students practice pouring cold water from a beaker into a graduated cylinder while using the appropriate heat-resistant device before they pour the hot water. Also, a pipet can be used to safely and carefully dispense the last 5–10 mL of hot water into the graduated cylinder, instead of trying to carefully measure and pour from hot glassware.
• The use of other materials, such as isopropyl alcohol or different salt solutions, can also be used to show differences in heat transfer, as well as specific heat. Appropriate safety precautions should be followed if other chemicals are used for this experiment. Consult current Materials Safety Data Sheets for more detailed safety information and additional handling and disposal information for any chemical that may be used. If alcohol or other flammable liquids are used, these should not be heated, but should replace the “cold water” in this experiment. Do not use hydrocarbons, such as acetone, which may dissolve the calorimeter cups.

Teacher Tips

• Have students experiment with different temperature differences between the cups. Alternatively, have them experiment with the same temperature difference, but at higher or lower initial temperatures. Which experiment design produced the largest temperature change between the cups? Does the amount of change depend on the initial temperatures or only the temperature difference?

Sample Data

Answers to Questions

1. Which experiment produces the largest temperature increase in the “cold water” cup? Why?

   Experiment 2 had the largest temperature increase in the “cold” water. This is because it requires less heat to increase the temperature of 175 mL of water compared to 350 mL of water. Since approximately the same amount of heat is transferred from the “hot” cup (the temperature difference is the approximately same in the beginning), the 175 mL of water reaches a higher temperature.

2. Calculate the heat lost (Q_lost) by the “hot water” in both experiments using Equation 1. Are these quantities positive or negative? Assume the specific heat of water is 1.0 cal/g∙°C and the density of water is 1.0 g/mL.

   Results will vary with individual data.
   
   Experiment 1: From sample data

   Q_lost = mcΔT = 350 mL x 1.0 g/mL x 1.0 cal/g∙°C x (69 °C–90 °C) = –7350 cal

   Experiment 2: From sample data

   Q_lost = mcΔT = 350 mL x 1.0 g/mL x 1.0 cal/g∙°C x (71 °C–90 °C) = –6650 cal

   These quantities are negative.

3. Calculate the heat gained (Q_gained) by the “cold water” in both experiments using Equation 1. Are these quantities positive or negative? Assume the specific heat of water is 1.0 cal/g∙°C and the density of water is 1.0 g/mL.

   Results will vary with individual data.
   
   Experiment 1: From sample data

   Q_gained = mΔT = 350 mL x 1.0 g/mL x 1.0 cal/g∙°C x (13 °C–3 °C) = 3500 cal

   Experiment 2: From sample data

   Q_gained = mcΔT = 175 mL x 1.0 g/mL x 1.0 cal/g∙°C x (18 °C–3 °C) = 2625 cal

   These quantities are positive.
4. –Q_lost should equal Q_gained. Was this observed? What are some possible errors that occurred during the transfer of thermalenergy?

   No, the calorimeter cups and aluminum bar are not 100% efficient. Answers will vary, but here is a list of a few possible errors: Unaccounted heat is transferred and lost to the calorimeter cups, metal thermometers and air above the water; heat is lost to the air from the aluminum bar during heat transfer; and some heat is retained by the aluminum bar and not fully transferred to the water.

5. Describe ways to improve this lab by eliminating sources of experimental error or design flaws in the setup.

   Individual answers will vary. A few examples might be to insulate the aluminum bar from the surroundings by wrapping it with different materials like tape, paper, or cloth. Seal any possible holes in the calorimeter lid. Use thicker calorimeter cups for better insulation. Use insulated thermometers to prevent significant heat transfer into thermometer material.

6. Why does a metal spoon feel hotter than a wooden spoon when they are held in a pot of hot water?

   The metal spoon is a good conductor of heat. The water’s heat energy will readily transfer into the spoon, and the spoon will readily transfer the heat to a colder hand that grabs it. The wooden spoon is a poor conductor of heat and will not acquire much heat from the water, or transfer its heat energy readily.

References

Kotz, J. C.; Treichel, P., Jr. Chemistry and Chemical Reactivity, 3rd Ed.; Saunders College: New York, 1996; pp 258–271.

Pedrotti, Leno S. Principles of Technology, Unit 4, 2nd Ed.; Center for Occupational Research and Development: Waco, Texas, 1991; pp 96–100.

Introduction

Why does a metal spoon sitting in a hot bowl of soup get hot? What causes wind? These are both consequences of heat transfer. In this experiment, the transfer of heat will be investigated.

Concepts

• Thermodynamics
• Heat transfer
• Conservation of energy

Background

All matter has internal thermal energy as long as the temperature of the substance is above absolute zero. Thermal energy, also known as heat energy, is the energy retained by a substance due to the continuous motion of the atoms and/or molecules that compose it (from tiny vibrations, rotation and the spinning electrons). Absolute zero is the theoretical temperature at which all motion stops, including atomic motion and electron spin. (This temperature has been calculated to be –273.15 °C, which set the zero, called “absolute zero,” on a new temperature scale called the “Kelvin scale.” The thermal energy of a substance is determined by its temperature and the type and number of atoms or molecules that it is composed of. Thermal energy cannot be directly measured, however, but the transfer of thermal energy between substances can be determined using a special apparatus known as a calorimeter. With a calorimeter, a substance’s change in temperature, as a result of thermal energy transfer, can be measured accurately. The temperature change can then be directly related to the thermal energy change using the following equation:

Q = thermal energy change
m = mass
c = specific heat*
ΔT = change in temperature, T_final – T_initial (“Δ” is the capital Greek letter delta, which means “change in”) *Specific heat is a physical property unique to every substance, and is defined as the heat energy required to raise the temperature of one gram of a substance by one degree Celsius.

The Law of Conservation of Energy states that energy must be conserved. Energy cannot be created or destroyed, it can only changed and/or transferred. Therefore, thermal energy lost by one substance must be equal to the amount of thermal energy gained by another substance (Equation 2). This is the First Law of Thermodynamics. Thermodynamics is the study of thermal energy transfer. The standard convention used when calculating thermal energy is that the temperature of a substance decreases (ΔT is negative) when thermal energy is lost. So Q_lost is negative. This means the expression –Q_lost will actually become a positive quantity in Equation 2. When thermal energy is gained, the temperature increases (ΔT is positive) and Q_gained is positive.

An important property of thermal energy is that it always transfers spontaneously from a high temperature region to a low temperature region. There are three ways in which thermal energy can be transferred: conduction, convection and radiation.

Conduction involves the transfer of thermal energy through the direct contact of hot and cold substances. “Hot” regions have fast-moving particles (atoms and/or molecules), which collide with and transfer some of their energy to slower moving particles in a neighboring “cold” region. The faster moving particles will slow down (and this region will cool down) while the slow-moving particles will speed up (and this region will heat up). This energy transfer will continue to proceed down the temperature gradient, from the “hot” neighbors to the “cold” neighbors, until there is thermal equilibrium, or no temperature difference between any regions. When there is no temperature difference, there is no thermal energy transfer.

Not all materials conduct thermal energy equally. A material’s ability to transfer its heat energy throughout itself, to other substances, or to have heat transferred into it, is known as thermal conductivity. Thermal conductivity is an intrinsic property of a material. Metals conduct thermal energy much more readily than nonmetals, for many of the same reasons metals conduct electricity better. Metals generally have extra electrons which are able to move throughout the material relatively freely and this allows for thermal energy to be transferred readily throughout the material and to other substances. Materials that do not conduct thermal energy well are known as insulators. The calorimeter cups in this experiment are examples of good insulators.

Convection is the movement of thermal bodies of different heat energy from one region to another. Instead of transferring energy between particles, a large number of “hot” particles (a “hot” region) move and displace a large number of “cold” particles (a “cold” region) creating thermal convection currents. Convection generally occurs in fluids like liquids or gases in which the hotter regions are less dense and rise, displacing a colder region above. For example, in the atmosphere, the air heats up at the surface more quickly than in the upper atmosphere. As the air becomes warmer, it expands, becoming less dense and the air mass rises. The colder air-mass above will be displaced, sink and then flow in to fill this space. This convection current is commonly referred to as wind.

Radiation is thermal energy transfer that does not require a medium, and can travel through the vacuum of space. Thermal energy is changed and then transferred in the form of electromagnetic energy called infrared radiation that radiates from all hot bodies. Infrared radiation cannot be seen by the naked eye, but its effects can be seen (and felt) when it strikes and interacts directly with atoms and molecules in substances causing their temperature to rise.

Materials

Safety Precautions

Wear chemical safety glasses, heat-resistant gloves and a chemical-resistant apron for good laboratory practice. Be careful when heating with a hot plate. Do not use bare hands to pour the hot water from the heated beaker. Follow all laboratory safety guidelines

Procedure

Calorimeter Setup

1. Insert each end of the aluminum bar into the off-center slot in the two foam lids. Slide the lids until they just reach the bend in the bar (see Figure 1).
   {12811_PreLab_Figure_1_Calorimeter setup}
2. Carefully insert a metal-backed thermometer into the remaining slots in each foam lid.

Preparation

1. Fill a 600-mL beaker with approximately 400 mL of tap water. Add 4 or 5 ice cubes. Allow the water to cool to between 0 and 10 °C. Measure the water temperature using a classroom thermometer.
2. As the ice water cools, fill another 600-mL beaker with approximately 400 mL of tap water. Place the beaker on a hot plate and heat the water to a temperature between 80 and 90 °C. Measure the water temperature using a classroom thermometer.
3. When the water in each beaker has reached the appropriate temperature ranges, follow the procedure.

Experiment 1

4. If necessary, remove any unmelted ice cubes from the cold water beaker with a spoon or forceps.
5. Using a 500-mL graduated cylinder, measure 350 mL of the cold water. Measure the exact volume to the nearest milliliter (mL) and record this as the water volume for the “cold cup” in the data table.
6. Pour the cold water from the graduated cylinder into one of the foam calorimeter cups.
7. With the same 500-mL graduated cylinder, measure 350 mL of hot water. Warning: Do not handle the hot beaker without appropriate heat-resistant hand protection. Measure the exact volume to the nearest milliliter (mL) and record this as the water volume for the “hot cup” in the data table.
8. Pour the hot water from the graduated cylinder into the other foam calorimeter cup.
9. Immediately place the lids, with the U-shaped aluminum bar and metal-backed thermometers inserted in the slots, into the foam calorimeter cups until the tops of the foam lids are flush with the tops of the cups.
10. Slide the thermometers down as far as they will go to get a good temperature reading.
11. Allow the thermometers to adjust for approximately 30–60 seconds.
12. If necessary, carefully slide the thermometers up until the temperature can be read on each thermometer. Measure, to the nearest degree Celsius (1 °C), the initial temperature in the cold water cup and in the hot water cup. Record these temperatures under the corresponding cups in the data table. Then slide the thermometers back into the cups as far as they will go.
13. Allow the cups to remain undisturbed for 15 minutes.
14. At 15 minutes, slide the thermometers up (if necessary) to measure and record the temperature in each cup. Record the temperature measurements in the data table. Then slide the thermometers back into the cups.
15. At 30 minutes, slide the thermometers up (if necessary) to measure and record the temperature in each cup. Record the temperature measurements in the data table.
16. If time permits, empty and dry the calorimeter cups, and allow the U-shaped aluminum bar to sit for 15 minutes so that its temperature can return to its original starting temperature (room temperature) before performing Experiment 2. Otherwise, Experiment 2 should be performed during the next laboratory class period.

Experiment 2

17. Follow the Calorimeter Setup steps 1 and 2 (if necessary) and Procedure steps 1–15. For Experiment 2 use 175 mL of cold water and 350 mL of hot water. Make sure the cold and hot water are approximately at the same initial temperature (plus or minus 5 °C), respectively, as they were in Experiment 1.
18. Consult your instructor for appropriate disposal procedures.

Student Worksheet PDF

12811_Student1.pdf