Chem& 161
General Chemistry w/Lab I

Experiment: Hess’s Law

Purpose
This laboratory will give you some practice in calorimetric calculations, using Hess’s Law, and computing heats of reaction.

Introduction
When a reaction is carried out in an aqueous solution the energy given off (or taken up) by the system is assumed to be exchanged with the water. This allows for a simple calculation of the heat of the reaction by first measuring the temperature change for the water, and then using the equation,

qwater = mwater•S.H. water•DT water,

to calculate the heat, qwater, absorbed (or given off) by the water. In this equation mwater is the mass of the water (for simplicity, assume 1.00 mL of water weighs 1.00 g), S.H.water is the specific heat of water (4.184 J/g°C), and DT water = Tfinal – Tinitial. It is assumed that no heat is lost to the calorimeter.

Since the energy is exchanged between the system and the water we note that

qsystem = – qwater,

and the enthalpy change for the reaction is then

DHrxn = qsystem /moles of limiting reactant

For this experiment the reactions below will be carried out, the temperature changes noted, and the heats of reaction calculated. You will then use Hess’s law to demonstrate the relationship between the reaction equations and the respective DHrxn for each.

  1. NaOH (s) ——> NaOH (aq)
  2. NaOH (s) + HCl (aq) ——> NaCl (aq) + H2O (l)
  3. NaOH (aq) + HCl (aq) ——> NaCl (aq) + H2O (l)

Procedure
You will be working with a partner in this experiment.

    Reaction A

  1. Use a utility clamp and a one–holed cork to suspend a thermometer from a ring stand as demonstrated by your instructor. Slide a one–hole Styrofoam cup, whose rim has been cut off, over the bottom of the thermometer in an inverted position. This will serve as the lid to your calorimeter.
  2. Place two nested Styrofoam cups (your calorimeter) under the lid and thermometer, and measure 100.0 mL of deionized water into the calorimeter. Lower the thermometer into the water, leaving the lid up for the moment. Record the temperature of the water, Ti.
  3. Weigh out about 2 grams of solid sodium hydroxide, NaOH, and record the mass with an accuracy of at least 0.01 g. Since sodium hydroxide readily picks up moisture from the air, it is necessary to weigh it quickly and proceed to the next step without delay. Do not attempt to measure exactly 2.00 grams. Caution: Handle the NaOH and the resulting solution with care as they are caustic.
  4. Add the solid NaOH to the calorimeter, and stir continuously for at least three minutes, or until a maximum temperature has been reached. Be sure to record the maximum temperature attained, Tf.
  5. Dispose of the solution in the waste container in the hood. Rinse and dry the thermometer and calorimeter, and reassemble your calorimeter.

    Reaction B

  6. Repeat steps 2–5 using 100.0 mL of 0.500 Mhydrochloric acid, HCl, instead of the water. Be sure to use less than 2 grams of the NaOH. Caution: Handle the HCl solution and the NaOH solid with care.

    Reaction C

  7. Repeat steps 2–5, measuring 50.0 mL of 1.00 MHCl, instead of the water, into the calorimeter in step 2. Then in step 3, instead of solid NaOH measure out 50.0 mL of 1.00 MNaOH solution in a graduated cylinder, and determine its temperature. Only after Ti for both the HCl and NaOH solutions have been determined, add the 1.00 MNaOH solution to the calorimeter. This reaction will be very fast, so be sure to record the Tf immediately. Caution: Handle the HCl and NaOH solutions with care.

Analysis and Report

  1. Be sure to include all experimental data in your lab book.
  2. Calculate the DHrxn for each of the reactions A, B, and C, expressing each DHrxn in units of “kJ/mole” of the limiting reactant. The calculations must be included in your report.
  3. Examine the three reaction equations, and determine the relationship between them. Demonstrate this relationship by combining the equations in a manner consistent with Hess’s Law.
  4. Write the net ionic equation for each of the three reactions A, B, and C, and then again demonstrate the relationship between the three ionic equations by combining them in a manner consistent with Hess’s Law.
  5. Using this relationship between the equations, apply Hess’s Law to the experimental values for DHrxn for the reactions. Do the values “match”? Include in your report a discussion of how well the experimental results fit Hess’s Law.
  6. Construct a single enthalpy (reaction coordinate) diagram that clearly shows the relationship between the three reactions.
  7. Use the DHf° values in Appendix B in your text to calculate the theoretical values for each DHrxn. Compare these values with the experimental values, and include an analysis of the comparison in your lab report.
  8. In this experiment we neglected any heat that may have been lost to the calorimeter, i.e.,we assumed that qcal = 0. Did this tend to make your experimental results higher or lower that the theoretical values? Explain your reasoning carefully.

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