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 • Cp water • T 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), Cp water is the specific heat capacity of water (4.184
J/g°C), and T 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: Hrxn = 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 Hrxn for each.
A. NaOH (s)  NaOH (aq)
B. NaOH (s) + HCl (aq)  NaCl (aq) + H2O (l).
C. NaOH (aq) + HCl (aq)  NaCl (aq) + H2O (l)
Procedure
You will be working with a partner conducting only one of the experiments A, B or C and then sharing data with
the class to determine class averages for each A, B and C experimental values.
Safety: NaOH and HCl are extremely hazardous and caustic materials. Proper handling and safety of all lab
participants must be observed at all times. Keep your lab station and adjacent areas clean and dry. Goggles
are to be worn at all times and any spill is to be cleaned up immediately. If you spill any solution or solid be
sure to rinse any exposed tissue immediately under running water and notify the instructor.
For your own safety in the lab make sure you identify the location of the acid neutralizer—NaHCO3 and the
base neutralizer—solid citric acid containers prior to beginning this lab. All remaining solutions and solids must
be disposed of in the chemical hood in the appropriately labeled containers. No acid or base may be poured
into any sink prior to neutralization.
Reaction A: Solid NaOH dissolves in water
NaOH (s)  NaOH (aq)
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 cover over the end of a thermometer so that it projects
through the lid and reaches reaches nearly to the bottom of two nested Styrofoam cups (your calorimeter).
2. 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, T f.
5. Dispose of the solution in the waste container in the hood. Rinse and dry the thermometer and calorimeter,
reassemble your calorimeter and remount the thermometer.
Reaction B: Solid NaOH neutralizes aqueous HCl
NaOH (s) + HCl (aq)  NaCl (aq) + H2O (l)
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 cover over the end of a thermometer so that it projects
through the lid and reaches nearly to the bottom of two nested Styrofoam cups (your calorimeter).
2. Measure out 50.0 mL of 0.500 M hydrochloric acid, HCl.
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 solution and solid 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. Caution: Handle the solid NaOH and the HCl
and NaOH solutions with care. Rinse and dry the thermometer and calorimeter, reassemble your
calorimeter and remount the thermometer.
Reaction C: Aqueous NaOH neutralizes aqueous HCl
NaOH (aq) + HCl (aq)  NaCl (aq) + H2O (l)
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 cover over the end of a thermometer so that it projects
through the lid and reaches nearly to the bottom of two nested Styrofoam cups (your calorimeter).
2. Measure 50.0 mL of 1.00 M HCl, into the calorimeter. Then measure out 50.0 mL of 1.00 M NaOH solution
into a graduated cylinder, and determine its temperature. Only after Ti for both the HCl and NaOH solutions
have been determined, add the 1.00 M NaOH 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. Rinse and
dry the thermometer and calorimeter, reassemble your calorimeter and remount the thermometer.
Analysis and Report
1. Be sure to include all experimental data in your lab book.
2. Calculate the Hrxn for each of the reactions A, B, and C, expressing each Hrxn 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 Hrxn 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. (See Figure 5.21 on page 191 in your text)
7. Use the Hf° values in Appendix C in your text to calculate the theoretical values for each Hrxn.
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.