Materials Included In Kit

Additional Materials Required

Prelab Preparation

Prepare a boiling water bath by filling a 600-mL borosilicate glass beaker with 450 mL of water and place it on a hot plate (or Bunsen burner equivalent). Add several boiling stones to the bottom of the beaker. Two hot water baths should be sufficient for 10 student groups.

Safety Precautions

Use extreme caution to avoid burns when heating the metal sample in the boiling water. Wear chemical splash goggles, heat-resistant gloves and a chemical-resistant apron. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information.

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. Allow the hot water to cool before rinsing down the drain. The metal samples should be collected, dried and stored for repeat use.

Lab Hints

• Enough materials are provided in this kit for 30 students working in groups of three or for 10 groups of students. This is a Super Value Kit—all the materials included are completely reusable. Data for two metal unknowns can reasonably be collected in one 50-minute class period. The prelaboratory assignment may be completed before coming to lab, and the data compilation and calculations may be completed the day after the lab.
• This lab activity works well using a coffee cup calorimeter, which is made by simply placing two 8-oz. insulated foam hot beverage cups together (for extra insulation) and filling the inside cup with the measured amount of water. An alternative thick-walled Foam Calorimeter with a lid is available from Flinn Scientific, Catalog No. AP6456.
• Digital thermometers work well with this experiment and eliminate the risk of breaking a thermometer. Caution students not to leave thermometers standing up in any cup as the cup may tip over and the glass thermometer may break.
• Enough water must be used in the boiling water baths and in the calorimeters to cover the metal specimens completely. For the long aluminum specimen, approximately 450 mL is required for the water bath and 175 mL of water is required for the calorimeter. For the smaller specimens, less water may be used to obtain a larger temperature change. If students will be comparing data, each calorimeter should have the same amount of water.
• Several of the metals are similar in color and size. Lined up from tallest to shortest the specimens are aluminum, zinc (dull, silver gray, surface slightly furrowed), tin (light silver, shiny and reflective, pock-marked), steel (silver or gray, polished or smooth metal surface) and copper.

Teacher Tips

• Use this activity to enhance understanding of specific heat, heat capacity, and transfer of thermal energy.
• Allow each student group to compare their data with the other groups. They should note that each of the specimens has the same mass (57.9 ±0.1 g).
• Discuss the specific heat of different metals in terms of their use in cookware (pots and pans). Is a low or high specific heat desired in cookware? What other properties of the metals are important in this applications?
• Students may also be interested in determining the amount of energy (caloric content) in solid foods. Use Flinn Scientific’s Economy Choice Calorimeter (Catalog No. AP4533) to explore this concept.

Answers to Prelab Questions

1. Define specific heat.

   Specific heat is the quantity of heat heeded to raise the temperature of one gram of a substance one degree Celsius.

2. Wen are two objects considered to be in thermal equilibrium?

   Two objects are considered in thermal equilibrium when they are at the same temperature.

3. Use Table 1 and Equation 1 from the Background section to calculate the amount of energy (in joules) that is needed to heat an iron nail with a mass of 7.0 grams from 25 °C until it becomes red hot at 750 °C.

   Solution: q = (7.0 g) • (0.448 J/g °C) • (725 °C) = 2274 joules of energy

Sample Data

Sample data for zinc is given.

Answers to Questions

1. Calculate and record the average of the two trials for each value in the data table.

   Student answers will vary.

2. Calculate the
   1. temperature change of the water, ∆T (water), by subtracting the initial temperature of the water from the final temperature. Be sure to include the correct sign with your answer.

      Sample calculations for zinc are given.
      ΔT (water) = 27.5 °C – 25.5 °C = 2 °C

   2. temperature change of the metal, ∆T (metal), by subtracting the initial temperature of the heated metal sample from the final temperature of the water and metal sample in the cup. Be sure to include the correct sign with your answer.

      ΔT (metal) = 27.5 °C – 99.4 °C = –71.9 °C

3. Use Equation 2 to calculate the heat energy gained by the water.

   q(gained by water) = (175.7 g) x (4.184 J/g °C) x (2 °C)
   q(gained by water) = 1470 joules

4. Use Equations 4 and 5 to calculate the specific heat of the unknown metal in J/g °C.

   q(gained by water) = –q(lost by metal)
   q(gained by water) = –[(m (metal) x C_p(metal) x ΔT (metal)
   1470 J = –[(58.0 g) x C_p(metal) x (–71.9 °C)]
   C_p(metal) = 0.353 J

5. Determine the identity of the unknown metal used by comparing the experimental specific heat value to the published literature specific heat values listed in Table 2.

   The experimental specific heat value is closest to the specific heat of both copper and zinc. Since the unknown metal was silver gray, the metal is zinc.

6. Experimental procedures will no doubt lead to some degree of difference from the published literature value. Determine the percent error for specific heat for the metal used. This can be done by comparing the value obtained in the lab (experimental value) with the literature value. Use the equation for percent error.

7. Suggest possible reasons for discrepancies between the experimental and literature values.

   Heat may have been lost to the environment during transfer of the metal from the boiling water to the calorimeter. Heat also may have been lost since the calorimeter did not have a lid and because it is not a perfect insulator. The calorimeter itself also absorbed some energy that was not accounted for. The temperature of the metal was assumed to be in equilibrium with the temperature of the boiling water. This may not have been exactly the case.

8. Calculate the amount of energy needed to heat up 240 mL (8 oz) of ice-cold water (0 °C) to body temperature (37 °C) after drinking it. The mass of 1 mL of water is 1 gram.

   (240 g) x (4.184 J/g °C) x (37 °C) = 37,154 J

Introduction

Five metal specimens of equal mass are heated to the same temperature. When each is added to a precise amount of water, the temperature of the water will be altered to a significantly different extent. What is the reason for this? Each of the metals has a different specific heat!

Concepts

• Specific heat
• Heat capacity
• Calorimetry

Background

Transfer of heat or heat flow always occurs in one direction—from a region of higher temperature to a region of lower temperature—until some final equilibrium temperature is reached. The transfer of heat energy can be detected by measuring the resulting temperature change, ΔT, calculated by subtracting the final temperature from the initial temperature.

In this experiment, heat is transferred from a hot metal sample to a colder water sample. Each metal causes the temperature of water to increase to a different extent. This means that each metal must have a differing ability to absorb energy and then release energy to the water causing the temperature to rise. The ability of any material to contain heat energy is called that material’s heat capacity. The measure of heat capacity, or the quantity of heat needed to raise the temperature of one gram of a substance by one degree Celsius at constant pressure is termed specific heat, and is represented by the symbol, C_p. The SI unit for specific heat is given in J/g•°C and the non-SI unit is cal/g•°C. (Note: 1 calorie = 4.184 joules).

In general, larger metal atoms have lower specific heat. Part of the reason for the variation of specific heat values is that substances composed of larger atoms have fewer atoms for the same amount of material as a substance with smaller atoms. For example, the mass of each lead atom is larger than the mass of each iron atom. Therefore, a given mass of lead has fewer atoms than the same mass of iron. When heat is added to lead and iron, fewer atoms need to be put in motion in lead, and thus less heat is needed to increase the kinetic energy of the atoms in the lead. Therefore, the specific heat of lead is lower than the specific heat of iron (see Table 1). Compare the heat capacities of wood and the metals in general. Wood—with its relatively high specific heat value—is able to absorb more heat than metal before its temperature rises, and therefore does not feel hot to the touch. Metals, on the other hand, will heat up more quickly and feel hot to the touch due to their relatively low heat capacities. In general, the lower the specific heat value, the quicker the temperature will rise. Notice the high heat capacity of water. Water is able to absorb and store large amounts of heat. This has a moderating effect on air temperature near bodies of water.

The amount of heat delivered by a material (q) is equal to the mass of the material delivering the heat (m) multiplied by the specific heat of the material (C_p) multiplied by the temperature change associated with delivering the heat (ΔT). (See Equation 1.) To make accurate measurements of heat transfer and to prevent heat loss to the surroundings, an insulating device called a calorimeter is used. A calorimeter is a device used to measure heat flow, where the heat given off by a material is absorbed by the calorimeter and its contents (often water or other materials whose heat capacities are known).

The heat gained by the water in the calorimeter (or gained by the calorimeter itself if a dry calorimeter is used) must be equal in magnitude (and opposite in sign) to the heat lost by the sample (see Equations 2 and 3). Since then Equation 5 may be used to calculate the specific heat of an unknown metal sample.

Experiment Overview

The purpose of this activity is to measure the specific heat of one or more unknown metals using a calorimeter. Experimental values will be compared to theoretical values in order to identify each metal.

Materials

Prelab Questions

1. Define specific heat.
2. When are two objects considered to be in thermal equilibrium?
3. Use Table 1 and Equation 1 from the Background section to calculate the amount of energy (in joules) that is needed to heat an iron nail with a mass of 7.0 grams from 25 °C until it becomes red hot at 750 °C.

Safety Precautions

Use extreme caution to avoid burns when heating the metal sample in the boiling water. Wear chemical splash goggles, heat-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines.

Procedure

1. Measure the mass of an unknown metal sample on a balance to the nearest tenth of a gram. Record this mass in the data table on the Specific Heat Worksheet.
2. Cut approximately 30 cm of fishing line.
3. Tie one end of the thread to the knob on the metal sample.
4. Once the water bath is boiling, carefully place the metal sample in the boiling water using the tied thread. Make sure the end of the thread hangs over the lip of the beaker but does not touch the heat source. Allow the metal sample to sit in the boiling water for at least 5 minutes so that its temperature reaches equilibrium with the water.
5. While the metal sample is heating, measure the mass of the foam calorimeter cups. Record this mass in the data table.
6. Fill the foam calorimeter with approximately 175 mL of tap water.
7. Measure the mass of the foam calorimeter and water. Record this mass in the data table.
8. Set the foam calorimeter in a 400-mL beaker for extra stability.
9. Calculate the mass of the water by subtracting the mass of the calorimeter from the mass of the calorimeter and water. Record this value in the data table.
10. Measure the initial temperature of the water in the foam calorimeter in degrees Celsius. Record this value in the data table.
11. After 5 minutes, measure the temperature of the boiling water bath. Record this value as the Initial Temperature of Heated Metal Sample in degrees Celsius in the data table.
12. Using the fishing line, quickly and carefully pull the metal sample from the boiling water, allow excess water to drip from the sample for a second or two and then place the metal sample into the foam calorimeter cup and fully submerge it under the water. Begin measuring the temperature changes.
13. Gently stir the water in the calorimeter cup occasionally with a stirring rod to ensure thorough heating of the water in the calorimeter cup.
14. Measure the highest temperature that the water reaches in degrees Celsius. Record this value as the final temperature in the data table.
15. Repeat steps 1–14 for a second trial. 16. If time permits, repeat the experiment using a different unknown metal sample. (See instructor.)

Student Worksheet PDF

12017_Student1.pdf