Forces and Bonding: Calorimetry

• To accurately measure temperature changes during chemical reaction
• To discover quantitative aspects of heat produced by reactions
• To understand that enthalpy is extensive

Chemical reactions are accompanied by energy changes. Most often, the energy manifests as heat. Heat may be either produced or absorbed, depending on the reaction. A reaction that produces heat is exothermic; one which absorbs heat, endothermic. Most spontaneous reactions (those which occur unaided) are exothermic. If a reaction is carried out in a closed, insulated container, the heat produced or absorbed causes a temperature change (DT) of the container contents. Measurement of DT allows quantitative determination of the heat produced per mole of chemical reaction. This experimental technique is called calorimetry.

A chemical reaction between reagents A and B in aqueous solution is shown in equation 1. For discussion, we assume the reaction is exothermic--produces heat.

To measure the heat produced, we follow a very careful and precise procedure:

1. Prepare solutions of reactants A and B of precisely known concentration.
2. Allow the temperatures of solutions A and B to stabilize.
3. Place a precisely known volume of reactant A in the calorimeter.
4. Measure and record the temperature of reactant A in the calorimeter at precise intervals for a specific total time (for example, record temperature every 30 seconds for a total of two minutes). During this time adjust the temperature of reactant B until it is the same as that of reactant A.
5. At time t_o, quickly add a precisely known volume of reactant B to the calorimeter with good stirring.
6. Measure and record the temperature of the solution in the calorimeter at precise intervals for a specific time (for example, once every 30 seconds for 3-5 minutes total).
7. Plot the temperature-time data. A typical plot for an exothermic process is in Figure 1.
8. Measure DT for the reaction by extrapolating the pre- and post-reaction linear portions of the T-t curve to time t_o.

Hess' Law. Hess' Law states that enthalpy, H, is a state function. The value of DH for conversion of a set of reactants to a set of products is independent of how the conversion is done. For example, consider reaction of AB with C. Two different ways of carrying out this reaction are shown below.

Path 1, direct reaction:

Path 2, indirect reaction, involving two steps:

The sum of reactions b and c is reaction a. Since paths 1 and 2 start and end with the same substances, the total change in enthalpy along path 1 is the same as the total change in enthalpy along path 2.

If any two of these enthalpies are known or measurable, the third may be readily calculated, even if it is not directly measurable! This is the value of Hess' Law.

Aqueous Solutions of Acids and Bases. In this experiment, we work with two strong acids, HCl and H₂SO₄, and the strong base NaOH. When dissolved in water, these three substances ionize completely, as indicated below.

Thus a 0.1 M solution of HCl actually contains 0.1 mole H⁺ and 0.1 mole Cl^- per liter of solution. It contains no unionized HCl molecules. Similarly, a 0.1 M solution of H₂SO₄ contains 0.1 mole H⁺ and 0.1 mole HSO₄^- per liter; and a 0.1 M solution of NaOH contains 0.1 mole Na⁺ and 0.1 mole OH^- per liter.

When we mix a solution of HCl with a solution of NaOH, a neutralization reaction occurs between H⁺ and OH^-:

The Cl^- ion from HCl and the Na⁺ ion from NaOH undergo no reaction, so are not explicitly written in the neutralization equation. When we mix a solution of H₂SO₄ with a solution of NaOH, the situation is more complex. Two reactions must take place during neutralization of H₂SO₄:

The total enthalpy change for neutralization of a solution of H₂SO₄ by NaOH, DH_total, is therefore DH₆ + DH₇, by Hess' Law.

In this experiment, you will use calorimetry to discover an answer to the focus questions below.

1. In Figure 1, why does T stay constant before addition of reagent B?
2. In Figure 1, why does T steadily decrease after reaction of A and B?
3. In Figure 1, why does the full temperature change, DT, not take place instantaneously when B is added?
4. Suppose that a plot of DT versus volume HCl in the calorimeter was constructed from class data. What would such a plot look like?
5. Suppose you add 50 mL 2 M NaOH solution to 20 mL 2 M HCl solution and measure DT₁. Then you add 50 mL 2 M NaOH solution to 20 mL 2 M H₂SO₄ solution and measure DT ₂. How are DT₁ and DT₂ related?

• graph paper
• thermos calorimeter
• bent glass stirring rod
• timer
• 2 thermometers
• 50-mL burets
• 50-mL graduated cylinder
• 100-mL graduated cylinder
• standard NaOH, »2M
• standard HCl, »2M
• standard H₂SO₄, »2M

Safety glasses must be worn at all times in the laboratory. You will work with solutions of acids and bases. Avoid ingestion and contact of these solutions with the skin. If you spill an acid or base solution on your skin, wash the area immediately with plenty of cold water. BE PARTICULARLY CAREFUL WHEN FILLING THE BURET WITH 2.0 M ACID.

Record all data in your lab notebook. Each team will be assigned 2 or 3 volumes of acid solution to react with 50 mL of 2.0 M NaOH. Obtain about 160 mL of NaOH solution and a volume of acid equal to the sum of your assigned volume + 15 mL. Rinse a 50-mL buret with 2 5-mL portions of your acid, then fill the buret with acid. Measure the DT for addition of 50 mL of 2.0 M NaOH to each assigned acid volume in the calorimeter, using the general procedure outlined in the Background. For results to be comparable, all teams must have the same total volume of solution--100 mL--in the calorimeter at the end of the reaction. Add a volume of distilled water equal to (50 - assigned acid volume) mL to the calorimeter using a 50-mL graduated cylinder. Then add the assigned volume of acid solution to the calorimeter using a buret. Pour 50 mL NaOH solution into the 100-mL graduated cylinder. Measure the temperature of the solution in the calorimeter and the NaOH solution in the graduate. They should be the same to within 0.2 ^oC. If necessary, adjust the temperature of the NaOH solution by running hot or cold water over the outside of the graduate. BE CAREFUL NOT TO GET ANY OF THIS WATER INTO THE GRADUATE. After temperature equilibration, measure and record the temperature (to the nearest 0.1^oC) of the solution in the calorimeter at t = -60 sec, t = -30 sec, and t = -5 sec. At t = 0, quickly remove the calorimeter lid and pour the 50 mL NaOH solution into the calorimeter from the 100-mL graduated cylinder. Total volume in the calorimeter will be 100 mL. Record temperature at t = 10, 20, and 30 sec and subsequently at 30 second intervals for a total time of 3 minutes. Repeat for the second and third runs. Make all three T-t plots on the same sheet of graph paper, and determine DT for each run. Process data as directed by the instructor. Work on Questions until the postlab discussion.

Clean-up. When you have finished all of your work:

• Clean Labkit glassware by recommended procedures, shake off excess water, and return to the Labkit.
• Clean Pasteur pipets, graduated pipets, and vials by recommended procedures, shake off excess water, and place in the drying oven.
• Rinse calorimeters, stirrers, and graduates thoroughly with tap water, shake off excess water.
• Return all borrowed equipment to the instructor before leaving lab.
• Clean up your work area before leaving lab.

Pour spent reagents carefully into the sink, and flush down the drain with plenty of water.

1. "A Laboratory-Centered Approach to Teaching General Chemistry" by Robert W. Ricci and Mauri A. Ditzler. Journal of Chemical Education, 1991, 68, 228-231.

Preparation Questions