Introduction Procedure: Calorimeter Constant Heat of Neutralization Heat of Solution Specific Heat of a Metal Make an Adobe Acrobat Copy of this experiment

Heat is a form of energy, sometimes called thermal energy, which can pass spontaneously from an object at a high temperature to an object at a lower temperature. If the two objects are in contact, they will, given sufficient time, both reach the same temperature.

Heat flow is ordinarily measured in a device called a calorimeter. A calorimeter is simply a container with insulating walls, made so that essentially no heat is exchanged between the contents of the calorimeter and the surroundings. Within the calorimeter, chemical reactions may occur or heat may pass from one part of the contents to another, but no heat flows into or out of the calorimeter from or to the surroundings.

SPECIFIC HEAT
When heat flows into a substance, the temperature of that substance will increase. The quantity of heat q required to cause a temperature change Dt of any substance is proportional to the mass m of the substance and the temperature change, as shown in Equation 1. The proportionality constant is called the specific heat, S.H., of that substance.

 q = (specific heat) ( m )( Dt) = S.H. ( m)( Dt)                                    (1)

The specific heat can be considered to be the amount of heat required to raise the temperature of one gram of the substance by 1oC (if you make m and Dt in Equation 1 both equal to 1, then q will equal S.H.). Amounts of heat are measured in joules (historically: calories). To raise the temperature of 1 g of water by 1oC, 4.18 joules of heat must be added to the water. The specific heat of water is therefore 4.18 joules/goC. Since 4.18 joules equals 1 calorie, we can also say that the specific heat of water is 1 calorie /goC. Ordinarily heat flow into or out of a substance is determined by the effect that the flow has on a known amount of water. Because water plays such an important role in these measurements, the calorie, which was the unit of heat most commonly used until recently, was actually defined to be equal to the specific heat of water.
The specific heat of a metal can readily be measured in a calorimeter. A weighed amount of metal is heated to some known temperature and is then quickly poured into a calorimeter that contains a measured amount of water at a known temperature. Heat flows from the metal to the water, and the two equilibrate at some temperature between the initial temperatures of the metal and the water. The amount of heat that flows from the metal as it cools is equal to the amount of heat absorbed by the water and the calorimeter. In thermodynamic terms, the heat flow for the metal is equal in magnitude but opposite in direction, and hence in sign, to that for the water. For the heat flow q,

 q H2O + q calorimeter = - q metal                                                                   (2)

If we now express heat flow in terms of Equation 1 for both the water and the metal M, we get
 (S.H. H2O ) (m H2O ) (Dt H2O ) + (H. C. calorimeter ) (Dt calorimeter ) =                                                                                                                       - (S.H.metal )(mmetal) (Dtmetal)          (3)

In this experiment we measure the masses of water and metal and their initial and final temperatures. (Note that Dtmetal <0 and Dt H2O > 0, since At Dt = tfinal - tinitial.) Given the specific heat of water, we can find the positive specific heat of the metal by Equation 3. We will use this procedure to obtain the specific heat of an unknown metal.  The specific heat of a metal is related in a simple way to its molar mass. Dulong and Petit discovered many years ago that about 25 joules were required to raise the temperature of one mote of many metals by 1oC. This relation, shown in Equation 4, is known as the Law of Dulong and Petit:

where MM is the molar mass of the metal. Once the specific heat of the metal is known, the approximate molar mass can be calculated by Equation 4. The Law of Dulong and Petit was one of the few rules available to early chemists in their studies of molar masses.

HEAT OF REACTION

When a chemical reaction occurs in water solution, the situation is similar to that which is present when a hot metal sample is put into water. With such a reaction there is an exchange of heat between the reaction mixture and the solvent, water. As in the specific heat experiment, the heat flow for the reaction mixture is equal in magnitude but opposite in sign to that for the water. The heat flow associated with the reaction mixture is also equal to the enthalpy change, AH, for the reaction, so we obtain the equation:

 q reaction = DH reaction   =   - [q H2O     +    q calorimeter ]                                       (5)

By measuring the mass of the water used as solvent, and by observing the temperature change that the water undergoes, we can find q H2O by Equation 1 and DH by Equation 5. If the temperature of the water goes up, heat has been given off by the reaction mixture, so the reaction is exothermic; q H2O is Positive and DH is negative. If the temperature of the water goes down, the reaction mixture has absorbed heat from the water and the reaction is endothermic. In this case q H2O is negative and DH is positive. Both exothermic and endothermic reactions are observed.
One of the simplest reactions that can be studied in solution occurs when a solid is dissolved in water. As an example of such a reaction, note the solution of NaOH in water:

 NaOH(s)    Na+(aq) + OH- (aq) ; DH= DHSolution                              (6)

When this reaction occurs, the temperature of the solution becomes much higher than that of the NaOH and water that were used. If we dissolve a known amount of NaOH in a measured amount of water in a calorimeter, and measure the temperature change that occurs, we can use Equation 1 to find qH2O for the reaction and use Equation 5 to obtain DH. Noting that DH is directly proportional to the amount of NaOH used, we can easily calculate DHSolution  for either a gram or a mole of NaOH. In the second part of this experiment you will measure DHSolution  for an unknown ionic solid.

Chemical reactions often occur when solutions are mixed. A precipitate may form, in a reaction opposite in direction to that in Equation 6. A very common reaction is that of neutralization, which occurs when an acidic solution is mixed with one that is basic. In the last part of this experiment you will measure the heat effect when a solution of HCl, hydrochloric acid, is mixed with one containing NaOH, sodium hydroxide, which is basic. The heat effect is quite large, and is the result of the reaction between hydrogen ions in the HCl solution with hydroxide ions in the NaOH solution;

 H+(aq)  + OH- (aq)    H2 O(aq) ; DH = DHSolution                              (7)

A. Purpose
Find answers to the following questions:

1. What is the Heat Capacity of the calorimeter?
2. What is the Heat of Neutralization of a strong acid and a strong base?
2. What is the Specific Heat of a metal?
3. What is the Heat of Solution of a salt in water?

Make copies of the following two links.
DATA AND CALCULATION

 WEAR YOUR SAFETY GLASSES WHILE PERFORMING THIS EXPERIMENT

B. EXPERIMENTAL PROCEDURE

1. Heat Capacity of the Calorimeter
From the workbench obtain a calorimeter and thermistor. The calorimeter consists of two nested Styrofoam cups fitted with a card board cover. There are two holes in the cardboard cover for a thermistor and a glass-stirring rod with a loop bend on one end. Assemble the experimental setup as shown in the figure below. Determine the mass of the empty calorimeter, cover, and stirrer. Place about 50-mLs of deionized water in a clean beaker, and heat the water to boiling. While the water is being heated, prepare LabWorks II according to the instructions given below. Begin recording the temperature of the system about 3 -4 minutes before addition of the boiling water. Use a digital thermometer to record the temperature of the boiling water. Immediately pour the hot water into the calorimeter and continue recording the temperature for at least 200 seconds after the addition of the hot water.
(To go to a discussion of collecting the data, click here. )

2. Heat of Neutralization
From the workbench obtain 50-mL of 1.00 M HCl, and 50-mL of 1.00 M NaOH. Assemble the experimental setup as shown in the figure below. Instructions on using the thermistor to measure the temperature can be found in the accompanying pages.

 Figure 1 Rinse out your calorimeter with distilled water, pouring the rinse into the sink. In a graduated cylinder, measure out 25 cm3 of 1.00 M HCl; pour the acid solution into the calorimeter. Rinse out the cylinder with distilled water, dry, and measure out 25 cm3 of 1.00 M NaOH; pour that solution into a dry 50-mL beaker. Using LabWorks II, measure the temperature of both the acid and of the base to ± 0.1oC, making sure to rinse and dry your thermistor before immersing it in the solutions. Put the thermistor back in the calorimeter containing the HCl solution. Once the themistor is in the acid solution begin and continue recording the temperature for about 3 -4 minutes, pour the NaOH solution into the calorimeter and determine the temperature change. Stir the reaction mixture, and continue recording data until it is apparent from the experimental data that the maximum temperature has been reached by the reacting solution. (To go to a discussion of collecting the data, click here. )

3. Specific Heat of a Metal
Fill a 400-cml beaker two-thirds full of water and begin heating it to boiling. While the water is heating, weigh your sample of unknown metal (~30 grams) to the nearest 0. 1 g on a top loading balance. Weigh the empty test tube;  pour the metal into the test tube.  Now weigh the test tube and metal. Put the test tube and metal into the beaker containing the boiling water. Continue heating the metal in the water bath for at least 10 minutes after the water begins to boil to ensure that the metal attains the temperature of the boiling water. The water level in the beaker should be high enough so that the top of the metal is below the water surface. Add water as necessary to maintain the water level.

While the water is boiling, weigh the calorimeter to 0.1 g. Place about 40 cm3 of water in the calorimeter and weigh again. Insert the stirrer and thermistor into the cover and put it on the calorimeter.

Using LabWorks II, measure the temperature of the water in the calorimeter for about 3 -4 minutes before adding the hot metal. Take the test tube out of the beaker of boiling water, remove the stopper, and pour the metal into the water in the calorimeter. Be careful that no water adhering to the outside of the test tube runs into the calorimeter when you are pouring the metal. Replace the calorimeter cover and agitate the water as best you can with the glass stirrer. Continuously record the temperature of the mixture until the temperature again becomes relatively constant. Repeat the experiment, using about 50 cm3 of water in the calorimeter. Be sure to dry your metal before reusing it; this can be done by heating the metal briefly in the test tube in boiling water and then pouring the metal onto a paper towel to drain. You can dry the hot test tube with a little compressed air.

(To go to a discussion of collecting the data, click here. )

4. Heat of Solution

Place about 50 cm3 of distilled water in the calorimeter and weigh the same as in the previous Procedure. In a small beaker weigh out about 4-grams of the solid compound assigned to you. Make the weighing of the beaker and of the beaker plus solid to 0.1 g. Using LabWorks II measure the temperature of the water for about 3 -4 minutes before adding the unknown salt. The temperature should be within a degree or two of room temperature. Add the unknown salt to the calorimeter. Stirring continuously and occasionally swirling the calorimeter, record the temperature until the final temperature remains relatively constant. Check to make sure that all the solid dissolved. A temperature change of at least 5 degrees should be obtained in this experiment. If necessary, repeat the experiment, increasing the amount of solid used.
Dispose of the solution in Part 2 as directed by your instructor.

(To go to a discussion of collecting the data, click here. )

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DATA AND CALCULATION

Instructions for the use of LabWorks II

- Computer-based measurements (must be completed during the laboratory period).

i. Opening LabWorks II .
To open the LabWorks II program,

Either Click on the Microsoft start button in the lower left corner of the screen. When a pop-up window appears, click PROGRAMS, and then drag the cursor over to LabWorks II 4. Another pop-up window will appear; then click on LabWorks.exe. The LabWorks II program is then activated.

or open LabWorks II 4 by clicking on the ICON found on the Desktop Window.

A view of LabWorks II's Main Menu window should be on your screen.
In the Main Menu screen on the Tool bar , open the Calibrate window by clicking on the button labeled Calibrate.

ii. Calibration Set-up.
The Calibration window appears lists numerous probes which can be utilized in LabWorks II. In this experiment, we will use only the thermistor as a probe:
We will complete three tasks to use the themistor for temperature measurements:

1st    Connect the probe to the LabWorks II interface
2nd  Edit the Sensor so that the data is collected as we choose
3rd  Calibrate the sensor using the Calibration Wizard. The sensor will be calibrated by recording a low temperature (ice water) and warm water) using a digital thermometer.

iii. Design Set-up.
The next step involves the design of the experiment. At this time you give the computer instructions to obtain temperature measurements at some regular time interval.

iv. Acquire Set-up.
The next step involves collecting the experimental data using the probe and parameters previously defined.

v. Analysis Setup View.

The data you have collected will be saved on the 31/2 Floppy (A:) drive or to the desk top for temporary strage. To finish this experiment, you will now analyze this information.

We must accomplish a number of tasks:

Find and open the saved data file,
Open the Analyze window and make the following ajustments,

Set the factors influencing the form of the Graphed data,

Define the General appearence of the Graph,

Define the appearence of the X-Axis,

Define the appearence of the Y1-Axis,

Define the appearence of the Y2-Axis,

Analyze our experimental numbers.
Instructions on graphing in terms of the
assignments for the x- and y-axis.
Use the LabWorks II curve fitting ability to

analyze our data to obtain both the maximum
and minimum temperatures.

iv. Print

Make certain the Spreadsheet and Graph view is on the computer screen. To print your graph, click on the File button found in the menu toolbar.
Then click on Select Print.
Only print the User Graph View rather than the Data Spreadsheet. WHY?