Matter & Energy: Temperature & Heat in Physical Processes

Miramar College Chemistry page 1 Energy Lab

Your Name: ______

Lab Partner’s Name: ______

Objectives: 1) To observe changes in temperature and heat energy which occur during physical processes such as freezing and dissolving. 2) To become familiar with the use of a simple calorimeter. 3) To calculate the specific heat of a metal and to compare the result with a known value.

Background: Read the section on matter and energy in your lecture textbook.

Equations:

Equation (1)

Heat Energy = specific heat (J/g•°C) x mass (g) x ∆T (°C)

Equation (2)

∆T = (Tfinal - Tinitial)

Equation (3)

% error = x 100

Conversions:

1 calorie = 4.184 Joules

specific heat of water = 4.184 J/g•°C

Equipment Needed:

Alcohol thermometer Styrofoam cup with lid Beaker (400-mL) Test tubes (2, 25 x 150 mm) Graduated cylinder (100-mL) Utility clamp Hot plate Weigh boat (plastic) Iron ring Wire gauze Laboratory balance Wire stirrer Ringstand Graph paper

Chemicals Needed:

Acetic acid, glacial (appx. 15 mL) Calcium chloride, CaCl2 (appx. 10 g)

Miramar College, Chem 152L, Energy Lab Revised 6/18/2007; page 1 Miramar College Chemistry page 2 Energy Lab

Potassium nitrate, KNO3 (appx. 10 g) Metal with unknown specific heat (appx. 20 g)

Pre-lab activity:

1. Define the following terms in your own words:

a. melting point

b. cooling curve

c. change of state

d. specific heat

e. endothermic

f. exothermic

2. When a solid substance was dissolved in water originally at room temperature (approximately 20 °C), the resulting mixture had a final temperature of 30°C. Was the dissolving process endothermic or exothermic? Explain.

3. A certain physical process requires 32.0 kilojoules (kJ) of heat. Convert this quantity to calories.

Post-lab questions:

1. A student performed Part C and did not allow the metal to heat long enough, so that it did not attain the temperature of the boiling water bath. Would this result in a calculated specific heat that was higher or lower than the true value? Explain.

2. A student performed Part C and did not wait for the water and metal mixture to reach a temperature minimum. Would this result in a calculated specific heat that was higher or lower than the true value? Explain.

3. If the heat transfer for dissolving one gram of calcium chloride is reported to be 670 joules, how many joules of heat energy was transferred when you dissolved the calcium chloride sample in Part A2? (Show supporting calculations below. Note: This does not involve Equation 1.)

4. During a change of state, as was seen in Part B., temperature remains constant. Is this also true of heat energy? That is, does the heat transfer also stop while change of state is occurring? Explain.

5. A 25.000 g metal sample is cooled from 100.0 °C to 25.0 °C. If the heat energy released by the metal is calculated to be 2.50 kJ, what is the specific heat of the metal in J/g•°C? Show calculations below.

Procedure

Part A: Dissolving: An Endothermic or Exothermic Process?

1. Dissolving Potassium Nitrate

Weigh approximately 10 grams of potassium nitrate (KNO3) and record the actual mass (to the nearest milligram) in your data table.

Using a 100-mL graduated cylinder, obtain 100.0 mL of de-ionized (DI) water and transfer it into a Styrofoam cup.

Measure the initial temperature of the water using your alcohol thermometer and record this temperature in your data table as the initial temperature.

Quickly add all of the potassium nitrate to the water and place the lid on the Styrofoam cup with the thermometer through the hole in the lid so that it is immersed in the water.

Swirl gently to dissolve the solid, stopping periodically to record the temperature of the mixture. When the temperature has reached either a minimum or maximum, record that temperature extreme in your data table as the final temperature.

Calculate the change in temperature (∆T) for the dissolving process using equation (2).

2. Dissolving Calcium Chloride

Repeat the process above substituting 10 grams of calcium chloride (CaCl2) for the potassium nitrate.

Data and Calculations

Part A: Dissolving: An Endothermic or Exothermic Process?

1. Dissolving Potassium Nitrate

Mass of KNO3 ______

Tinitial ______

Tfinal ______

∆T ______

Show your calculation for ∆T.

Is the dissolving of potassium nitrate in water an endothermic or exothermic process? Explain.

2. Dissolving Calcium Chloride

Mass of CaCl2 ______

Tinitial ______

Tfinal ______

∆T ______

Show your calculation for ∆T.

Is the dissolving of calcium chloride in water an endothermic or exothermic process? Explain.

Part B: Freezing of Acetic Acid (A Cooling Curve)

Assemble a change of state apparatus as in Figure 2, p. 38, in the lab manual and fill the 400-mL beaker half full with room temperature tap water.

Remove the cork and thermometer and pour about 25-30 mL of glacial acetic acid into the test tube. (CAUTION: This acid may cause burns. Do not let the acid come in contact with your skin or clothing!)

Insert a wire loop stirrer into the acetic acid and replace the cork and thermometer so that the thermometer bulb is fully immersed in the acetic acid and the wire loop goes around the thermometer.

Measure the initial temperature of the acetic acid and record this temperature in your data table at t = 0:00 s.

Add crushed ice to the tap water in the 400-mL beaker until you form an ice-water slurry. Make sure that the level of the acetic acid in the test tube is below the water level in the beaker.

Begin stirring the acetic acid with the wire stirrer (and start the stopwatch) as soon as you place the test tube into the ice water. Maintain a constant rate of stirring and record the temperature every 10-15 seconds until the first few crystals appear in the test tube and the temperature remains constant for two or three time points.

Note: If the acetic acid freezes too quickly, you may not produce an accurate cooling curve. Make sure to stir continuously and do not allow the entire solution to become solid. The freezing point is measured when both solid and liquid coexist.

Using the graph paper provided on pages 45 and 46 of your lab manual, make a graph of the temperature vs. time relationship (the cooling curve) for acetic acid using the data you have collected.

Refer to Figure 1, p. 38, in the lab manual and determine the freezing point for acetic acid. The freezing point is the temperature where the graph levels off after the initial cooling. It is possible (perhaps likely) that some supercooling will occur resulting in a slight dip in the curve prior to the actual leveling off. The freezing point is not the lowest point in the dip; rather, it is the horizontal portion of the graph, during the time when crystallization is occurring.

Thaw the acetic acid by heating the beaker full of water (with the tube inside) using a hot plate. Do not allow the water to boil. Heat it just long enough for the acetic acid to melt.

Repeat the process.

Part B: Freezing of Acetic Acid (A Cooling Curve)

Trial 1 Trial 2

Time (min:sec) T (°C) Time (min:sec) T (°C)

t = 0:00 ______t = 0:00 ______

______

Graph your data and attach the graph to the end of the report. Use your graph to answer the following questions:

What is the freezing point of acetic acid that you determined from your graph?

Is freezing an endothermic or exothermic process? Explain.

At the freezing point, temperature remains constant even though heat continues to be transferred as part of the freezing process. Does your graph help to validate this statement?

Part C: Specific Heat of a Metal

Fill a 400-mL beaker approximately half full with tap water and bring the water to a boil on a hot plate.

While waiting for the water to boil, do the following:

Weigh approximately twenty grams of your assigned metal into a plastic weigh boat that has been tared (i.e. the mass of the weigh boat has been “zeroed out” by pressing the tare button). Accurately record the mass of metal in the data table.

Add the metal to a 25 x 150 mm test tube and place in the boiling water bath for ten minutes.

While your metal sample is being heated for ten minutes, do the following:

After the metal has been heated in the boiling water bath for ten minutes, record the temperature of the boiling water bath as the initial temperature of the metal.

Using a graduated cylinder, obtain 25.0 mL of de-ionized (DI) water and transfer it into a Styrofoam cup.

Measure the initial temperature of the water using your alcohol thermometer and record this temperature in your data table as the initial temperature of the water.

Using a test tube holder, very quickly add all of the metal sample to the water in the Styrofoam cup and place the lid on the Styrofoam cup with the thermometer through the hole in the lid so that it is immersed in the water.

Swirl gently, stopping periodically to record the temperature. When the temperature has reached either a minimum or maximum, record that temperature extreme in your data table as the final temperature for both the water and the metal.

Calculate the change in temperature (∆T) for the metal and for the water, using equation (2).

Part C: Specific Heat of a Metal Sample # ______

Volume of water ______

Mass of metal ______

Tinitial for metal ______

Tinitial for water ______

Tfinal for metal & water ______

∆T for metal ______

∆T for water ______

Show your calculations for ∆T for the water and for the metal.

Use equation (1) to calculate the heat energy absorbed by the water. (Assume that the density of water is 1.00 g/mL.) Show your calculations below.

Assume that no heat is lost to the calorimeter. Therefore, the heat energy that is absorbed by the water should be equivalent to the heat energy that is released by the metal. Only the sign should change (from positive to negative).

Heat absorbed by water = ─ (heat released by metal)

Assuming the above equivalency is valid, calculate the specific heat of the metal by rearranging equation (1) to solve for specific heat. Show your calculations below.

Obtain the true value for the specific heat of your metal and calculate your percent error using equation (3).

Miramar College, Chem 152L, Energy Lab Revised 6/18/2007; page 12