Thermo Dyn Anics

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GCC CHM 152LL: Thermodynamics page 1 of 5

CHM 152LL: THERMODYNAMICS

Pre Lab:In addition to title, purpose, procedure, and data and results tables, please include the following twocalculations in your lab notebook and show your work for each.

1. Calculate the mass of ammonium chloride required toprepare 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.

PurposeIn 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 Sorxnand your experimental values to

calculate Hrxn. Together, these will allow you to calculate the free-energy changes, Grxn, for these two

processes.

IntroductionThe 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 - TS (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. Hrxnrepresents 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 Grxnto check your prediction.

Calculations

Sample calculations (for trial 1 for both salts) should be done in your lab notebook.

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 orlost 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 saltplaced in the calorimeter. (Recall thatthe 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/gC. To calculate change () for a variable it is always final minus initial.


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The heat of reaction, qsys, can then be calculated from equation 2.

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

So

rxnvalues for the dissolution of solid ammonium chloride and calcium chlorideusing 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 Sorxnlike a homework problem.

Finally you can calculate the experimental change in Gibbs 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 cupinside the Styrofoam cups. Record the exact volumeof 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 toadd 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 massused 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 vigorouslyby swirling thebeaker and contents, carefully holding the lid and thermometer in place, for three minutes. Record the

temperature of the mixture every 15 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: Tfis NOT the temperatureafter 3 minutes, but the maximum (or minimum) temperature obtained during the three minutes.

6. Pour your salt solution in the waste container. Rinse out your calorimeter, rinse the probe, then repeat theexperiment using ammonium chloride again for trial number two.

7. Repeat all steps for calcium chloride for two trials.

Clean-Up: CaCl2is hygroscopic and very corrosive to our balances. Please use thebrush by the balance to clean up any spills immediately . Any spills left behind mightresult in poin ts being deducted (at the discretion of your ins tructor ).

Pour the salt solutions in the waste container. Rinse everything well with tap waterfollowed 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.


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Data and ResultsRecord temperature data (every 15 seconds, for a maximum of 3 minutes). Two trials for eachsalt will be completed. The table below is an example of the data and results that you need torecord 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. Yourcalculations will also be completed in the lab notebook.

Show all calculations for one trial for each salt.

You will conduct two trials for each salt. Pay attention to units and significant figures inyour tables.


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Name: __________________________ Partners Names: ___________________________

Post Lab QuestionsPlease turn in pages 4-5 along with your laboratory notebook pages.

Dissolution of Ammonium Chloride Questions: 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 ammoniumchloride 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 moreorder or disorder? __________________

9. Is this salt soluble at all temperatures? Explain based on the signs of the enthalpy change andentropy change for the dissolution of ammonium chloride.

10. Calculate the temperature above or below which the salt will not dissolve, if applicable. Indicate ifthe solid will dissolve above or below this temperature.


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11. Use Hfoand Gfodata in Appendix 2 of your textbook to calculate Hrxno(Equation 10.18, p. 385)

and Grxno(Equation 18.12, p. 782) for the dissolution of ammonium chloride. Compare thesestandard values to your calculated values in lab. Discuss possible sources of error that resulted in

your value being different from the standard value.

Hrxno: _______________________________________________________________

Grxno: _______________________________________________________________

Sources of error:

Dissolution of Calcium Chloride Questions:Please refer to your results for calcium chloride toanswer 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 chloridespontaneous 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 moreorder 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 ifthe solid will dissolve above or below this temperature.

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Thermo Dyn Anics