Download Thermodynamics

CHM 152LL: THERMODYNAMICS
Pre Lab: In addition to title, purpose, and data/results tables, please include the following two calculations in
your lab notebook and show your work for each.
1. Calculate the mass of ammonium chloride required to prepare 25 mL of a 2.0 M solution.
2. Calculate the mass of calcium chloride required to prepare 25 mL of a 2.0 M solution.
Purpose
In this experiment you will use calorimetry to determine the enthalpy changes, ∆Hrxn, for the dissolution of two
chloride salts in water. You will then use textbook values to calculate ∆Sorxn and your experimental values to
calculate ∆Hrxn. Together, these will allow you to calculate the free-energy changes, ∆Grxn, for these two
processes.
Introduction
The Gibbs-Helmholtz equation expresses the relationship between the free-energy change, the enthalpy change,
and the entropy change at constant temperature and pressure:
∆G = ∆H - T∆S
(equation 1)
From knowing the value of ∆G, you may predict whether a process/reaction will be spontaneous at a certain
temperature. A process is spontaneous if ∆G is negative (∆G < 0), nonspontaneous if ∆G is positive (∆G > 0), and
at equilibrium if ∆G = 0.
The enthalpy change, ∆Hrxn, is the heat gained or lost by a system during a reaction carried out at constant
pressure. Most reactions occur in several steps, with energy required (endothermic, positive ∆H) because energy is
needed to break bonds, and energy released (exothermic, negative ∆H) because energy is released as new bonds
are formed. ∆Hrxn represents the total change in heat energy or enthalpy over the course of the reaction.
In this experiment, you will use a coffee-cup calorimeter to determine the heat absorbed or released during the
dissolution of ammonium chloride and the dissolution of calcium chloride. From observing the contents of the
coffee-cup calorimeter, you will decide whether the dissolution processes are spontaneous or nonspontaneous.
You will also calculate values of ∆Grxn to check your prediction.
Calculations
From the law of conservation of energy (energy is conserved) the total energy for the dissolution process is:
qsys + qsurr = 0
OR
qsys = - qsurr
(equation 2)
where qsys (or qrxn) represents the heat gained or lost by dissolving the solid, and qsurr (or qsoln) is the heat gained or
lost by the solution in the calorimeter. Thus, heat energy is essentially transferred between the dissolved solid and
the solution in the calorimeter. (For this experiment, we will consider the heat absorbed by the cup, probe, and
surroundings to be negligible, so it is not included in the expression above.)
The heat absorbed or released by the contents of the calorimeter is given by:
qsurr = (mass solution)*(specific heat of solution)*(∆T)
(equation 3)
The mass of the solution is the sum of the masses of the water and salt placed in the calorimeter. (Recall that the
density of water is1.00 g/mL.) Because the solution is very dilute, the specific heat of the solution is basically
equal to that of water, which is 4.184 J/g°C. To calculate change (∆) for a variable it is always final minus initial.
The heat of reaction, qsys, can then be calculated from combining equation 2 and equation 3.
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The molar enthalpy of reaction, ∆Hrxn, will then be calculated by dividing the heat of reaction by the experimental
number of moles of salt used in the experiment.
∆Hrxn = qsys / moles salt
(equation 4)
You will need to calculate the ∆Sorxn values for the dissolution of solid ammonium chloride and calcium chloride
using data from Appendix 2 in the back of your textbook. We do not have experimental data for this calculation, so
we will use the textbook values and solve for ∆Sorxn like a homework problem.
Finally you can calculate the experimental change in Gibb’s Free Energy (∆G) using the Gibbs-Helmholtz
equation, equation 1, using the initial temperature for T, the experimental value for the enthalpy of reaction, and
the textbook value for the entropy of reaction.
Procedure
1. Use a 100-mL graduated cylinder to measure about 25 mL of deionized water and add to the blue plastic cup
inside the Styrofoam cups. Record the exact volume of water used, paying attention to significant figures.
2. Tare out the weight of a plastic weighing cup. Remove the plastic cup from the balance and use a spatula to
add the appropriate mass of ammonium chloride (calculated in pre-lab) to the weighing cup. The mass should
be within ±0.2 grams of the calculated value. Record the exact mass used in your lab notebook.
3. Place a thermometer in the deionized water. Record the initial temperature of the water in degrees Celsius
(Time = 0 sec).
4. Add the solid to the water in the calorimeter and replace the lid. Stir the solution vigorously by swirling the
beaker and contents, carefully holding the lid and thermometer in place, for three minutes. Record the
temperature of the mixture every 10 seconds. Do not stir with the thermometer. Do not leave the thermometer
in the apparatus unsupported – it will be top-heavy and could fall over.
5. The highest (or lowest) temperature reached will be the final temperature Tf. Note: Tf is NOT the temperature
after 3 minutes, but the maximum (or minimum) temperature obtained during the three minutes.
6. Observe the appearance of the salt solution in the calorimeter to see if the salt dissolved. Pour your salt
solution in the waste container. Rinse out your calorimeter, rinse the probe, then repeat the experiment using
ammonium chloride again for trial number two.
7. Repeat all steps for calcium chloride for two trials.
Clean-Up: CaCl2 is hygroscopic and very corrosive to our balances. Please use the
brush by the balance to clean up any spills immediately. Any spills left behind might
result in points being deducted (at the discretion of your instructor).
Pour the salt solutions in the waste container. Rinse everything well with tap water
followed by a quick DI water rinse. Return the measuring cups to the reagent stations.
Clean your benchtop. Put all equipment back exactly where you found it.
Data and Results –
Record temperature data (every 10 seconds, for a maximum of 3 minutes) in your lab
notebook. Two trials for each salt will be completed. The table on page 3 is an example of the
data and results that you also need to record in your notebook. Please copy this table into
your notebook twice (one for each salt). All data must be recorded in ink in your lab
notebook as the reaction proceeds. Your calculations will also be completed in the lab
notebook.
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You will conduct two trials for each salt. Pay attention to units, significant figures, and
signs in your tables.
Show all calculations for one trial for each salt. Example calculations should include:
• moles of salt
• mass of water
• mass of solution
• ∆T
• qsys
• ∆Hrxn
• ∆Sorxn
• ∆Grxn
In your conclusion in your notebook, summarize and discuss your
calculated ∆H, ∆S, and ∆G values for both salts. Also discuss at least
2 experimental sources of error and how they could have affected your
results.
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Name: __________________________
Partners’ Names: ___________________________
Thermodynamics Discussion Questions
Please turn in pages 4-5 along with your notebook pages and your conclusion.
Dissolution of Ammonium Chloride Questions (19 pts): Please refer to your results for
ammonium chloride to answer the following questions.
1. Write the balanced chemical equation for the dissolution of ammonium chloride in water.
2. In the experiment, identify the system ______________ and the surroundings _______________.
3. Which one gains heat in this experiment? _______________
4. Is the system endothermic or exothermic? ___________________
5. Explain how the observed temperature change verifies your answer to #4.
6. From the temperature change obtained for the system in the calorimeter, what must be the sign for
∆Hrxn? ______________
7. a) Based on observing the solution in the coffee-cup calorimeter, is the dissolution of ammonium
chloride spontaneous or non-spontaneous at room temperature? ___________________________
b) Based on observing the solution, is ∆G > 0 or < 0 for this process at room temperature? _______
8. Based on your calculated change in entropy, does the dissolution of ammonium chloride create more
order or disorder? __________________
9. Is this salt soluble at all temperatures? Explain based on the signs of the enthalpy change and
entropy change for the dissolution of ammonium chloride.
10. Calculate the temperature above or below which the salt will not dissolve, if applicable. Indicate if
the solid will dissolve above or below this temperature.
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11. (5 pts) Use ∆Hfo and ∆Gfo data in Appendix B of your textbook to calculate ∆Hrxno (pp. 293-294) and
∆Grxno (p. 676) for the dissolution of ammonium chloride. Compare these standard values to your
calculated values in lab.
∆Hrxno: _______________________________________________________________
∆Grxno: _______________________________________________________________
How do your values compare to the textbook values?
Dissolution of Calcium Chloride Questions (16 pts): Please refer to your results for calcium
chloride to answer the following questions.
1. Write the balanced chemical equation for the dissolution of calcium chloride in water.
2. In the experiment, identify the system ______________ and the surroundings _______________.
3. Which one gains heat in this experiment? ______________
4. From the temperature change obtained for the system in the calorimeter, what must be the sign for
∆Hrxn? ______________
5. a) Based on observing the solution in the coffee-cup calorimeter, is the dissolution of calcium chloride
spontaneous or non-spontaneous at room temperature? ___________________________
b) Based on observing the solution, is ∆G > 0 or < 0 for this process at room temperature? _______
6. Based on your calculated change in entropy, does the dissolution of calcium chloride create more order
or disorder? __________________
7. Is this salt soluble at all temperatures? Explain based on the signs of the enthalpy change and
entropy change for the dissolution of calcium chloride.
8. Calculate the temperature above or below which the salt will not dissolve, if applicable. Indicate if
the solid will dissolve above or below this temperature.
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