Is the dissolution of nh4cl exothermic or endothermic

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How Can the Thermodynamics of

Dissolution be Defined? Introduction: The Gibbs-Helmholtz equation expresses the relationship between the free-energy change, ∆G, the enthalpy change, ∆H, and the entropy change, ∆S, 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) to break bonds, and energy released (exothermic,

negative ∆H) 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.

From the law of conservation of energy (energy is conserved) the total energy for the

dissolution process is:

qsystem + qsurroundings = 0 or qsystem = - qsurroudnings ……………………Equation 2

where qsystem (or qrxn) represents the heat gained or lost by dissolving the solid (the

system), and qsurroundings (or qsolution) is the heat gained or lost by the solution in the

calorimeter (the surroundings). Thus, heat energy is essentially transferred between the

dissolving solid and the solution in the calorimeter. (For this experiment, the heat

absorbed by the cup, probe, and surroundings can be considered as negligible, so it is

not included in the expression above.)

The heat absorbed or released by the contents of the calorimeter is given by:

q = m· Cs · T ………………………………Equation 3

where m is the mass of the solution, Cs represents the specific heat of solution, and T

is the change in temperature (T=Tfinal-Tinitial)

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For this experiment, the mass of the solution is the sum of the masses of the water and

solid placed in the calorimeter. (Recall that the density of water is 1.00 g/mL). The specific

heat of the solution can be assumed to be equal to that of water because the solution is

very dilute. Cs for water is 4.184 J/gºC. The value of qsurr can be calculated by plugging

in experimental values into Equation 3, and then determining the heat of reaction, qsys,

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 ∆Sºrxn values for the dissolution of solid ammonium chloride

and calcium chloride using data from Appendix C 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.

∆𝑆°𝑟𝑥𝑛 = ∑𝑛𝑆°(𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠) − ∑𝑚𝑆°(𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠) …………………Equation 5

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 for Part 1: Dissolution of NH4Cl (Record questions, answers and data table in notebook)

Objectives upon completion of this section:

 Use calorimetry to determine the enthalpy change, ∆Hrxn, for the dissolution of enough ammonium chloride in water to make 50 mL of a 1.0 M solution.

 Use this value of ∆Hrxn, and theoretical ∆Sºrxn (from textbook) to calculate the free-

energy change, ∆Grxn, for this process.

Before you come to lab, write a rough proposal using the procedural considerations below, calculate a proposed initial mass of salt, and create the following tables in your notebook:

 Develop a procedure that will allow you to determine the enthalpy changes, ∆Hrxn,

for the dissolution of enough ammonium chloride in water to make approximately

50 mL of a 1.0 M solution. You will then use your textbook values to calculate

∆Sºrxn, to finally find the free-energy change, ∆Grxn, for this process.

 During the course of the experiment consider if you will be recording any

observations or sources of potential error in your lab notebook?

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Proposed initial mass of ammonium chloride: To prepare 50 mL of a 1M solution, I will need _________g NH4Cl

Tables to include in your notebook:

 Create data Table 1 in your notebook that includes 14 rows (recordings will be taken every 15 seconds for 3 minutes) with the following column headings:

Time Trial #1 (ºC) Trial #2 (ºC)

1

.

.

.

15

 Create Table 2 in your notebook for experimental data like the following:

Ammonium Chloride Trial #1 Trial #2

Mass of NH4Cl(g)

Volume of Water (mL)

Mass of Water (g)

Mass of Solution (g)

Initial temperature, Ti (ºC)

Final temperature, Tf (ºC)

 Create Table 3 in your notebook for calculations like the following:

CALCULATIONS for Ammonium Chloride

Trial #1 Trial #2

∆T (ºC)

qsurr (J)

qsys (J)

Moles of NH4Cl (mol)

∆Hrxn (kJ/mol)

∆S◦rxn (J/mol∙K) from textbook

∆Grxn (kJ/mol)

Average ∆Grxn

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Experimental Setup:

 Consider the calorimeter set-up. What are all of the components and how do they fit together? If unsure, check with your TA prior to continuing.

 Is taking the mass of an empty calorimeter (with 2 Styrofoam cups, plastic cup, and lid) necessary?

 It is recommended to place the entire calorimeter setup into a 400 mL beaker. Would this help with stability? Would this help with containing heat?

 How will you measure temperature changes? How will this fit into the experimental setup?

Preparation of Reagents: How will you make 50 mL of a 1.0 M solution?

Measurement of the distilled water:

 What volume measuring device should you use? If you choose to use a device like a 100mL graduated cylinder, keep in mind that the actual volume reading should be recorded to the nearest 0.1mL.

 Think about the calculations you will have to perform. Which is more important: the volume of the water or the mass of the water?

 Is taking the mass of the distilled water necessary? How could you get a mass of just the water?

 Is taking the initial temperature of the distilled water in ºC necessary?

Measurement of the solid salt:

 The mass of salt you measure may not be the exact mass you proposed. As long as it is within a range +/- 0.2 g of your estimate is it okay if the mass varies a little from trial to trial? Why should you make sure to record the exact mass of solid you measure?

Experimental Procedure: How will you perform the experiment?

Before Beginning:

 Check your proposal. Is your experimental setup ready? Where is the salt? Where is the water? How will you mix them? What else should be a part of the system?

 Have you recorded all of physical measurements in your notebook? data Table 2)

While performing the experiment:

 Stir vigorously by swirling the beaker and contents, carefully holding the lid and

temperature probe in place

 How often should you record temperature? What units of temperature should you

use? How long are you to keep taking measurements?

 Why should you not stir using the temperature probe?

 How will you determine the highest (or lowest) temperature reached, final

temperature (Tf)?

Note: Tf is NOT the temperature at the end of the trial, but the maximum (or minimum)

temperature obtained during the three minutes.

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Experimental Reset:

 How should you dispose of your salt solution?

 Should you rinse out your calorimeter and probe with distilled water?

 Repeat all steps above for another trial.

Before you Leave Checklist:

 Write the balanced chemical equation for the dissolution of ammonium chloride in

water in your lab notebook.

 For the experiment, identify the system and the surroundings.

 Which one gains heat in this experiment?

 Is the system endothermic or exothermic?

 Explain how the observed temperature change verifies your answer to whether the

system is endothermic or exothermic.

 From the temperature change obtained for the system in the calorimeter, what

must be the sign for Hrxn?

Procedure for Part 2: Dissolution of CaCl2 (Record questions, answers, and data table in notebook)

Objectives upon completion of this section:

 Use calorimetry to determine the enthalpy change, ∆Hrxn, for the dissolution of

enough calcium chloride in water to make 50 mL of a 1.0 M solution.

 Use this value of ∆Hrxn, and theoretical ∆Sºrxn, (from textbook) to calculate the free-

energy change, ∆Grxn, for this process.

Before you come to lab:

 Calculate a proposed initial mass of salt.

 Write a rough proposal using similar procedural considerations as in Part 1.

 Create three new data tables in the same format as the ones used in Part 1.

You will be following the same procedure as used for Part 1, you might have to make

some adjustments or modifications, confirm any changes with your TA before

continuing, and make a note of these changes in your lab notebook!

Record all experimental data, observations, or sources of potential error you notice in your lab notebook.

Clean up and Waste Collection:

 CaCl2 is hygroscopic and very corrosive to our balances. Please use paper towel to clean up any spills immediately. Any spills left behind might result in points being deducted (at the discretion of your instructor).

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 Pour the salt solutions in the waste container. Rinse everything well with tap water followed by a quick distilled water rinse. Return the measuring cups to the TA cart.

 Clean your benchtop. Put all equipment back exactly where you found it.

Before you Leave Checklist:

 Write the balanced chemical equation for the dissolution of calcium chloride in

water.

 In the experiment, identify the system and the surroundings.

 Which one gains heat in this experiment?

 Is the system endothermic or exothermic?

 Explain how the observed temperature change verifies your answer to whether the

system is endothermic or exothermic.

 From the temperature change obtained for the system in the calorimeter, what

must be the sign for Hrxn?

Report Considerations

Special Notes:

 Use correct significant figures for all values

 Follow all guidelines given in the General Lab Report Format on Bb

 Use an equation editor program wherever possible. It gives the report a more

organized, professional appearance.

 Read and confirm that each section throughout the report flows.

 Title page with section, group #, and group members’ names.

 Label each of the sections listed below and write in your own words, without

copying from the handout.

 Reports should be typed in past tense, passive voice, no use of first person, do not

address the reader, and no contractions or slang/text lingo. Double-spaced with

Arialor Times New Roman font.

 Check for proper grammar and spelling.

 Complete contribution sheets and turn-in to your TA.

Introduction

o Introduce the dissolution in water general reaction. o What is thermodynamics?

o What differences/similarities did you see with both solids?

o Explain the importance of taking the mass of each individual substance involved

in a system.

o Explain the importance of units for each variable involved in Gibbs Free Energy

equation.

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o Give a brief description of the major goals of lab.

o Include balanced chemical reactions for NH4Cl and CaCl2 in H2O (include physical

states and use super/subscripts!) as part of the description of approach.

Experimental Clearly describe the steps taken and materials/equipment used in the

context of the description:

o Do NOT copy the lab handout itself! Use your own words in third person, passive voice, past tense!

o Include all volumes and masses of solutions, including the starting mass of both

solids.

o All volumes measured with the 100mL graduated cylinder should be reported to

one decimal place.

o Do NOT round masses, include all digits.

Results

o Description of what it is or what the reader should take note of when studying the provided results or calculations.

o Within context or tables, present experimental observations… what was the

highest/lowest temperatures for each solid during the 3 minute trials?

o Easy to read tables, with units on each heading. Parts 1 (2 tables) and 2 (2 tables)

should be in separate tables.

o Show sample calculations. Show an example calculation for each type of

calculation performed (one example for everything that can be solved with a

calculator.)

o Pay strict attention to significant digits in all the calculations. You only need to show

one sample of each calculation in your formal report.

o

for each salt.

Discussion Discuss each result and/or observation:

o Summarize and analyze the data from the results section

 Do the temperature changes observed for each solid’s dissolution seem

reasonable? Why or why not? Does this mean the reaction is endothermic or

exothermic?

 What did the sign of the entropy for each dissolution reaction mean?

 Based on the ∆Grxn for each solid, what does this mean? Interpret/explain

your results.

- Is the dissolution spontaneous at room temperature?

- Does what you observed and what you calculated for ∆G support this

prediction? Why/Why not?

- Were the values of entropy and enthalpy both good driving forces for

spontaneity or was it just one of them? If so, which one?

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 Use ∆Hfº and ∆Gfº data in your textbook to calculate ∆Hrxn and ∆Grxn for

the dissolution of each salt. Compare and contrast these standard values

to your calculated values in lab.

o Discuss possible sources of error that resulted in your values of ∆Grxn being

different from the standard value. Ex: measurement error for volumes or masses,

adding too much solid to the water in the calorimeter during the reaction (explain

in terms of the temperature changes), loss of heat to hands/beaker, etc....? Then

explain how these errors might have affected the result. Would it have led to an

increase or decrease in final temperature? Would that have lead to a more

negative or more positive ∆Grxn?

o Discuss enthalpy, entropy, temperature, Gibbs Free Energy, and the law of

conservation of energy.

Conclusion Summarize what results were determined:

o Did the experiment succeed? i.e.: Were all of the objectives mentioned in the

"Introduction" achieved?

o Consider your error analysis from the Discussion section, and use them to make

suggestions for improvement of experimental procedure.

 Was it important to repeat multiple trials for each reaction? Why?

o List at least one way that what you learned in this lab could apply or be useful in a

real-life situation. Use credible resources!

References

Cite your references using an ACS format, even if it’s just the lab manual. Do not plagiarize the resources available to you!