Forces and Bonding: Calorimetry 2

Goals

• To accurately measure temperature changes in chemical reactions
• To discover quantitative aspects of reaction heats

Background

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 that absorbs heat, endothermic. Most spontaneous reactions (those that 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. For discussion's sake, we consider reaction of two strong acids, HCl, H₂SO₄, with the strong base NaOH. When dissolved in water, HCl, H₂SO₄ and NaOH 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.

The experiment involves measuring reaction heats for several common acids with the base, sodium hydroxide, according to equation 8.

Before beginning, predict the form of reaction 8 (i.e., the value of x) for the acids below.

1. Why is it necessary that the total volume of solution in the calorimeter be the same in all runs?
2. Why is it necessary that the acid and base solutions be at the same temperature before mixing?
3. Plot observed temperature change versus moles of acid for each of the four acids investigated. Determine the best fit slope of each plot, and from the slope calculate temperature change per mole of acid.
4. Is the enthalpy for reaction of HCl and NaOH the same as that for the reaction of HPO₄^2- with NaOH? Why or why not?
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 1 M H₂SO₄ solution and measure DT₂. How are DT₁ and DT₂ related?

• calorimeter (a small wide-mouth plastic thermos with two holes drilled in the lid, or 2 nested styrofoam cups)
• Celsius thermometers (2)
• 100-mL graduate (2)
• buret
• beakers
• 2 M NaOH
• 2 M HCl
• 0.667 M H₃PO₄
• 1 M NaH₂PO₄ (NaH₂PO₄·2H₂O is very soluble in water)
• 2 M Na₂HPO₄ (Na₂HPO₄·7H₂O dissolves to the extent of 104 g per 100 mL of water at 40 ^oC)
• 1 M H₂SO₄
• 1 M NaHSO₄

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 ACID.

Experimental

Your instructor will assign you an acid to study. Carry out the following procedure. Carefully record all data, observations, results, and conclusions in a laboratory notebook.

Reaction of acid and base in the calorimeter must be carried out precisely and in a very specific way:

1. Add distilled water to the calorimeter. The amount must be the difference between 50 mL and the volume of acid being used.
2. Add the appropriate volume of acid from a buret.
3. Measure 50 mL of NaOH solution in a 100-mL graduate.
4. Determine the temperatures of the acid and base solutions using separate thermometers. Adjust the temperature of the base solution by running hot or cold water over the outside of the cylinder until the temperature of base is within 0.2 ^o C of the acid temperature.
5. Record the temperature 60 s, 30 s, and 5s before you intend to add the base solution to the calorimeter.
6. Add the base solution to the calorimeter by removing the lid (if there is one), quickly and smoothly pouring the base solution in, and replacing the lid.
7. Swirl or stir the contents and read the temperature as soon as possible after mixing.
8. Record the temperature at 30 second intervals for 3 minutes after mixing.
9. Carefully pour the calorimeter contents in the sink, rinse the calorimeter with distilled water and dry it, and proceed to the next run.

When all runs have been completed, clean up as follows.

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

• Clean all glassware by recommended procedures and shake dry.
• Rinse Pasteur pipets by inserting the nozzle of your wash bottle in the top end and squirting water through. Aspirate dry.
• Return all borrowed equipment to the instructor before leaving lab.

Solutions of acid and base may be flushed down the drain with plenty of water.

Preparation
Forces and Bonding: Calorimetry 2

Preparation Questions