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CHEMICALEQUILIBRIUM

Teacher's Guide

Steady Unsteadiness

Students should be able to:

Program DescriptionProgram 1 begins with a toboggan sliding down a hill, illustrating thetendency for most changes to involve movement from a state of highenergy to one of low energy. The application of this regularity tochemical systems suggests that only exothermic reactions are possible.It also suggests that chemical reactions must stop. The fallacy in theirconclusions is illustrated and the concepts of steady state and dynamicequilibrium are introduced.

Before ViewingUse Activity 1, The Blue Bottle, and/or Activity 2, The Thionine System,to demonstrate that chemical reactions can be reversed. Most studentsbelieve this cannot be done, so this is time well spent. Note that thedevelopment of the accepted model for either of these activities isbeyond the ability of most students at this stage. However, if the activityis done in groups of four, most groups should come up with a model thati nvolves a set of reversible reactions.

Use Activity 3 to show that reactions proceed and reach a pointwhere macroscopic properties are constant. Ask students to suggestplausible explanations for the constant macroscopic properties, thenshow Program 1.

Explanation for Activity 1: The Blue BottleGlucose, a weak acid reacts with the hydroxide to form the glucoside ionwhich reduces the blue form of methylene blue to a colorless form.

When the flask is shaken, oxygen dissolves in the water and oxidizes thecolorless reduced form of methylene blue, reforming the blue dye. Belowis a simplified mechanism for the system.

Explanation for Activity 2: The Thionine SystemThionine is protonated by the sulfuric acid, and then is electronicallyexcited by the light source. The excited thionine ion oxidizes the iron(II)to iron(III). The reduced form of thionine is colorless. When the lightsource is turned off the excited thionine decays to its ground state. Thisreverses the iron reaction; iron(III) is reduced to iron(II), and the purplecolor of the oxidized thionine returns.

A more detailed explanation is available in L.J. Hardt's article "ThePhotochemical Reduction of Thionine" listed on page 5.

After ViewingUse Activities 4 and 5 to illustrate the difference between steady stateand dynamic equilibrium. Demo 1 shows a typical steady state system,with the rate of water flowing into the can equal to the rate of waterflowing out. The candle and burner flames also are steady state systems,but the input and output are not as obvious. Have students explain the"balance" in the systems.

Demo 2 involves two long-term experiments. The best way to treatthis demonstration is to set it up for next year and use the given obser-vations for your current classes.

Activity 5 is a logical stepping stone to Program 2. It is an excellentanalogy of dynamic equilibrium. Working in pairs, students transfercolored water from their container into their partner's. By starting withdifferent volumes of water and using different capacity vessels totransfer the water, the situation can be used to illustrate many of thefeatures of a chemical system.

Objectives

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1 . State the Second Law of Thermodynamics.2. Identify the implications of the regularity "energy runs downhill" to

chemical reactions.3. Identify systems in a steady state.4. Identify closed systems.5. Describe at least two explanations for macroscopic properties

becoming constant during a chemical reaction.6. Describe what is meant by the term dynamic chemical equilibrium.

Details can be obtained from J.A. Campbell's article "Kinetics - Earlyand Often" listed on page 5.

Activity 5 is written as a student experiment. If you wish to demon-strate this activity, use beakers of different sizes to transfer the waterand aquaria or large battery jars filled with different volumes of coloredwater.

Explanation for Activity 4, Demo 2:Long-Term ExperimentsSet-up 1Since the volumes continued to decrease, water must have continuouslyevaporated from the solutions. Thus the system must have been open togases, in this case water vapor escaping from the jar. Since thephenolphthalein reddened in the cylinder labelled "B," then some of thesodium hydroxide was neutralized, presumably by HCI(g) escaping fromcylinder "A."Note: Phenolphthalein has two color changes as pH increases: colorlessto pink between 8-10 and then back to colorless around pH 14.Set-up 2

because the solution in beaker "C" was absorbing water at a much fasterrate than "D." This was caused by the large difference in concentration.The solution in beaker "C" was 615 times more concentrated than thatin "D" at the start. Thus the beakers tended to exchange water. Theexchange should have continued until the concentrations were equal.The rate of evaporation from "D" should have started out fast and thendecreased until the concentrations were equal. But in this case thesituation was complicated by the water vapor diffusing into the jar. Whenthe rate of diffusion was greater than the rate of absorption by thesolution in beaker "C," the volume in "D" should have slowly begun torise. As the concentration decreased, the rate of absorption also shouldhave decreased; however, the concentration of the water was gettinglarger at the same time so the rate of evaporation should have increaseduntil, at 21 months, the rate of evaporation and condensation becameequal, or perhaps all evaporation and condensation stopped. Moredetails can be obtained from L.W. Bixby's article "Long-Term ChemicalReactions" listed on page 5.

Since the volumes of both beakers overflowed, the system must havebeen open to water vapor. The level in beaker "D" dropped initially

20 g of sodium hydroxide

20 g of glucose

1.5 mL of 1 % alcoholic methylene blue

1 L of water

250 mL Erlenmeyer flask with rubber stopper

Note: Dissolve the ingredients in the litre of water.Add the sodium hydroxide just before using.

Method1 . Fill the Erlenmeyer flask a little less than half

full with the solution. Stopper the flaski mmediately.

2. Shake the flask vigorously, then place it on awhite sheet of paper and observe whathappens. Repeat several times.

Using "OH' - " to represent the sodiumhydroxide, "G" to represent the glucose, and"MB" to represent the methylene blue, writeequations that will account for yourobservations.Make predictions and design experiments totest the models you developed in step 3.Check with your teacher before carrying theseout.

Discussion

Develop a model for the blue bottle experiment.I ndicate the evidence on which your model isbased.

Activity 2: The Thionine System

10 mL graduated cylinder

600 mL beaker

Light source (overhead projector or goose necklamp)

Method

Add 10 mL of the thionine solution, 10 mL ofthe sulfuric acid solution, and sufficient waterto bring the solution to 500 mL.Thoroughly mix in 2.0 g of iron(II) sulfateheptahydrate.Use the overhead projector to light the solutionfrom below or use the goose neck lamp with a250 W photoflood bulb to light it from above.Turn the light source on and observe whathappens. Then turn it off and observe. Repeatseveral times.Using "S" to represent the sulfuric acid, "T" torepresent the thionine, and "Fe 2+" to

represent the iron(II) sulfate in solution, writeequations that will account for your observations.

2.

3.

4.

5.

2

0.001 mol/L thionine

Apparatus

I ron(II) sulfate heptahydrate

Activities 3.

Activity 1: The Blue BottleApparatus

4.

6. Make predictions and design experiments totest the models you developed in step 5.Check with your teacher before carrying theseout.

DiscussionDevelop a model for the thionine system.I ndicate the evidence on which your model isbased.

Activity 3: Acid/BaseDemonstrationsDemo 1Slowly pour equal volumes of 1 mol/L NaOH andHCI solutions containing the indicator bromthymolblue into a beaker placed on the overheadprojector. Repeat for different but equal volumesof NaOH and HCI. Try to explain why the color ofthe mixture remains the same.

Note: There are two plausible explanations: thatthe reaction has stopped, and that there areopposing reactions occurring at equal rates.

Demo 2Use a syringe to inject stock ammonium hydroxidecontaining the indicator bromthymol blue into theapparatus shown in Figure 1. Explain the changes

in the manometer. Again, the same two explana-tions are plausible.

Use a syringe to inject stock hydrochloric acidcontaining the indicator bromthymol blue intoanother example of the set-up. Explain thechanges in the manometer. Again, two explana-tions are plausible.

Now add a little of the hydrochloric acid to theapparatus containing the ammonium hydroxide.Similarly, add a little of the ammonium hydroxide tothe apparatus containing the hydrochloric acid. Inthis case the chemical reaction is obvious whenthe white smoke forms. Account for any changesi n the indicator and manometer.

Note: It will be necessary to fiddle with theconcentrations and volumes of the acid and baseto get reasonable changes in the levels of the

Figure 2

manometer. Be sure to try this experiment beforeyou demonstrate it for the class.

Activity 4: Steady State/DynamicEquilibrium DemonstrationsDemo 1Place a burning candle and a lit Bunsen burner onthe demonstration desk. Thirdly, set up theapparatus as shown in Figure 2. Ask students toaccount for the "balance" in the three systems.

Note: Use a fairly crude knife edge for thefulcrum, otherwise it is very difficult to maintain thebalance. Ask students to compare the threesystems. How are they the same? How are theydifferent?

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Demo 2Make a transparency of the two experimental set-ups described below, or better still, actually setthem up. Include the following:

Set-up 1

Forty millilitres of 0.3 mol/L HCI was added to a100 mL graduated cylinder labelled "A." Ten milli-litres of a 1 mol/L NaOH solution containing theindicator phenolphthalein was added to a second1 00 mL graduated cylinder labelled "B." The twocylinders were sealed in a large jar and observedweekly for two years.

The solutions in both graduated cylindersdecreased continually over the period of 24months. After 13 months the solution in cylinderB reddened.

Set-up 2

One hundred and fifty millilitres of saturated(6.15 mol/L) NaCl was placed in a 250 mL beakerl abelled "C." One hundred and fifty millilitres of0.01 mol/L NaCl was placed in a second 250 mLbeaker labelled "D." The two beakers were sealedi n a large jar and observed weekly for two years.

After 13 months beaker "C" was filled and thenoverflowed by 30 mL or more into the jar. Thel evel in beaker "D" initially dropped, but it too alsoeventually overflowed. After 21 months there wasno apparent change.

Ask students to come up with plausible explan-ations for these observations. Help them focustheir attention on the principles learned in Pro-gram 1 with questions such as "How is the systemclosed?" "How is the system open?" "Were thesystems at equilibrium?" "At a steady state?"

Demo 3Use an equal arm balance and some small bath-room tiles to simulate a system at equilibrium.Label one set of tiles reactants (r) and, if possible,another set - half the mass of the former =

products (p). I n this way you can simulate thereaction where one reactant decomposes into twoproducts.

If possible, move the fulcrum position to oneside. This enables you to simulate a situationwhere there is more product than reactant atequilibrium.

Add the appropriate number of tiles to each panto achieve a balance. Then challenge the studentsto think of another way the balance could bemaintained. Obviously, if one (r) tile were removedfrom one side and two (p) tiles were removed fromthe other side, the system would remain balanced.But that is not how a chemical system works.Reactants are changed into products, and as soonas that happens the system is no longer balanced.However, if as one (r) tile were transferred to theproduct side, two (p) tiles were removed and one(r) tile was added to the reactant side, balancewould be maintained. This would simulate a steadystate with the continuous removal of products andthe continuous addition of reactants.

Another way to maintain balance is, of course,to switch one reactant tile for two product tiles. Inthis way both the reaction and its reverse wouldhave to occur and their rates would have to beequal.

Activity 5: A Model forChemical EquilibriumI n this activity you will simulate a systemapproaching chemical equilibrium. Your task is todetermine if the simulation is a good model for achemical system at equilibrium.

ApparatusTwo 150 mL beakers

Two plastic rulers

Two pieces of glass tubing of different diameters,10-15 cm long

1 0 mL graduated cylinder

Water with blue food coloring

Water with yellow food coloring

Method1 . Put some blue colored water in one beaker

and some yellow colored water in the other.The only restriction is that the total volume ofthe two is 100 mL.

2. Work in pairs. Transfer water from B, the bluebeaker, using one of the pieces of glasstubing, to Y, the yellow beaker. At the sametime your partner should transfer water from Yto B using the other piece of glass tubing. Besure to keep the glass tube vertical at all times.

3. Use the rulers to record the height of thewater in each beaker after every fiveexchanges.

4. Continue transferring water until you havethree successive readings that are the same.

5. When the heights remain constant measurethe volume transferred by each glass tube.

6. Repeat steps 1-5 for different volumes of blueand yellow water, but keep the total volume100 mL.

Observations

TIME HEIGHT OF WATER HEIGHT OF WATERBLUE BEAKER YELLOW BEAKER

( 5 exchanges) ( mm) ( mm)

1

2

3

etc.

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Discussion1. What are the following analogous to in a

chemical reaction? The height of the water inthe two beakers; the transfer of water fromone beaker to another; two pieces of glasstubing of different diameters; the coloring inthe water; the volume of water transferred; thevolume of water transferred when the heightremains constant.

2. What evidence suggests that you reached apoint of equilibrium in this activity? You shouldbe able to think of two key observations.

3. What property of the system determines thefinal height of the water in the beakers? Whatwould this be analogous to in a chemicalreaction?

4. Plot height of water versus time in units of fiveexchanges.(a) Describe the graph in words.(b) What is the significance of the horizontal

part of the graph?(c) What would be the significance of the

slope of a tangent drawn at any point onthis graph?

(d) What happens to the slope of the tangentas the system approaches equilibrium?What does this imply?

(e) The implications of (d) contradict what isactually happening; you can still see waterbeing transferred from one beaker toanother. Explain how it is possible for thegraph to suggest that the reaction hasstopped even though you can see it goingon.

5. Consider the following gaseous reaction:

(a) Plot a graph of the data.(b) Describe the graph in words.(c) Explain why the graph has the shape it

has.(d) What is the significance of the horizontal

part of the graph?(e) What would be the significance of the

slope of a tangent drawn at any point onthis graph?

(f) What happens to the slope of the tangentas the system approaches equilibrium?What does this imply?

( g) Use the analogy to explain the changes inthe slope of the tangents.6. Is this analogy a good model for chemical

equilibrium? Explain your answer.

One mole of C12 was added to 2 mol of CO in a1 L container. The concentration of the threegases was measured using a spectropho-tometer every minute. The following data wererecorded:

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Further ReadingBixby, L.W. "Long-Term Chemical Reactions.

Parts I & If. " Chemistry. September andOctober 1976.

Campbell, J.A. "Kinetics - Early and Often."Journal of Chemical Education. Vol 40. No.11. November 1963.

Hardt, L.J. "The Photochemical Reduction ofThionine: A Reversible Reaction." Journal ofChemical Education. Vol 26. No. 25. 1949.

Parry, Robert W. et al. Chemistry: ExperimentalFoundations. Englewood Cliffs, New Jersey:Prentice-Hall, 1970.

Rowley, Wayne R. E. Matter in Balance: ChemicalEquilibrium. Toronto: Wiley, 1979.

Sienko, Michell J., and Robert A. Plane.Chemistry. New York: McGraw-Hill, 1976.

Toon, E.R., and G.P. Ellis. Foundations ofChemistry. New York: Holt, Rinehart andWinston, 1973.

Dynamic Equilibrium

ActivitiesActivity 1: Testing theDynamic ModelI n this activity you will gather indirect evidence tosupport the dynamic model for chemicalequilibrium. It is not possible to see what ishappening on a microscopic scale, but it ispossibie to check if the reaction and its reverseoccur and if both reactants and products arepresent at equilibrium. The two systems you willstudy are:

ObjectivesStudents should be able to:

1. Describe a system at equilibrium macroscopically and microscopically.2. Use the Kinetic Molecular Theory to explain why a "dynamic" model

for chemical equilibrium is the most reasonable explanation for theobservations.

3. Describe the types of motion of molecules in the gaseous phase.4. Describe the changes in concentration of the reactants and products

as a chemical system approaches and reaches equilibrium.5. Interpret the "double arrow" conventions.6. Identify data that supports a dynamic model for chemical equilibrium.

Program DescriptionProgram 2 begins with a review of the tendency towards minimumenergy and the observation that in a closed system a chemical reaction

The formation of a black precipitate whensodium sulfide is added is used to identify leadi ons in solution. The addition of silver nitrate andthe formation of a pale yellow precipitate is usedto identify the bromide ions in solution. The silverbromide that is formed darkens when exposed tolight. Iodine vapor can be identified by itscharacteristic pink color.

Apparatus

point of equilibrium - can be reached by starting with the product,hydrogen iodide. But the puzzle of how to describe the system at thispoint remains. Has the reaction stopped, as the macroscopic propertiessuggest, or is there a dynamic equilibrium with reactions occurring in twodirections? To resolve this puzzle students are reminded that a kineticmodel is used to describe matter which supports the dynamicdescription of the equilibrium point. In addition, they are presented withempirical data that is difficult to explain without assuming a dynamicequilibrium between the forward and reverse reaction for thehydrogen/iodine/hydrogen iodide system.

Before ViewingUse the discussion of Demonstration 3 in Activity 4, and Activity 5 inProgram 1 as an introduction to Program 2. The questions that remainshould be "Are the analogies a good way to view chemical equilibrium?"and "What evidence is there to support a dynamic model?"

After ViewingHave students complete Activity 1. This enables them to gather indirectevidence for the dynamic model and to practise using it to describe aphysical change and a chemical change.

1 mol/L NaBr solution (102.9 g NaBr/L) * indropping bottles

Short-stemmed funnel and filter paper

Small test tube

150 mL beaker

Wash bottle filled with distilled water

Piece of acetate and a sheet of white paper

A glass rod

* All solutions must be made with distilled water.

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N02 and Nz04 i s used tonappears to have stopped. The reaction betweel ustrate the reversibility of chemical reactions. A quantitative look at theeaction between hydrogen and iodine follows. The analysis of the, ystem shows both reactants and products are present when theeaction appears to have stopped. It also shows that this point - the

1 . Add approximately 1 cm of the sodiumbromide solution to the test tube.

2. Add an equivalent volume of lead nitratesolution to the test tube. Caution: Weardisposable gloves. The white precipitate thatforms is lead bromide.

3. Fold the filter paper into a cone and place it inthe funnel. The funnel is held by the 150 mLbeaker.

4. Stir the contents of the test tube and transferthem onto the filter paper. Use the washbottle to ensure that all of the solid istransferred to the filter paper.

5. Wash the lead bromide on the filter papertwice using approximately 2 cm of distilledwater in the test tube each time.

6. Put the acetate sheet over a piece of whitepaper.

7. Use the glass rod to transfer as much of thel ead bromide as possible from the filter paperto the acetate sheet.

8. Cover the solid lead bromide with distilledwater. Mix the solid and water with the glassrod. Call this mixture 1. Let the mixture situntil you need it in step 13.

9. On another section of the acetate sheet put adrop of the lead nitrate solution. Add to it adrop of sodium sulfide solution. Record anychanges.

On another section of the acetate sheet put adrop of the sodium bromide solution. Add to ita drop of the silver nitrate solution. Recordany changes.On another section of the acetate mix a dropof the sodium bromide and lead nitratesolutions. Call this mixture 2. Allow the solidto settle and then divide the mixture into twoequal parts using the glass rod.

12. To one part of the mixture add a drop of thesodium sulfide solution. To the other part adda drop of the silver nitrate solution. Recordany changes.

13. Divide the mixture produced in step 8( mixture 1) into two equal parts. To one partadd a drop of the sodium sulfide solution. Tothe other part add a drop of the silver nitratesolution. Record any changes.

14. Dispose of the filter paper as directed byyour teacher. Rinse the acetate sheet, testtube, beaker, and glass rod with runningwater. Do not forget to wash your handsafter cleaning up. Lead compounds aretoxic.

Part B

1 . Use the hot plate to heat water until it isbetween 70 0 and 901C. Several pairs ofstudents can use the same hot water bath.

2. Hold the bottom of the test tube in the hotwater until the amount of iodine vapor remainsconstant.

3. Remove the test tube from the hot water andobserve what happens.

4. Repeat steps 2 and 3 several times.5. Return the corked test tube to your teacher.

Do not attempt to clean out the test tube.

ObservationsPart A

Mixture 1 is formed by adding distilled water tol ead bromide.

Mixture 2 is formed by mixing lead nitrate andsodium bromide.

Part B

Describe the original contents of the test tube andany changes that occurred as a result of heatingand cooling the tube.

DiscussionPart A

1 . Describe a test for lead ions in solution. Writea balanced chemical and ionic equation forthe test.

2. Describe a test for bromide ions in solution.Write a balanced chemical and ionic equationfor the test.

3. Write a balanced chemical and ionic equationto describe the formation of a white precip-i tate when solutions of lead nitrate andsodium bromide are mixed.

4. What are two plausible explanations for thefact that the amount of white precipitateformed when lead nitrate and sodium bromideare mixed does not seem to change onstanding?

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Part B 10.

Small corked test tube containing a few crystals ofiodine

Large beaker containing hot (70-90°C) water 1Hot plate

MethodPart A

5. What is implied by the observations madewhen sodium sulfide and silver nitrate areadded to mixture 2?

6. Explain why it is hard to accept the"stopped" model for equilibrium based on theobservations made when sodium sulfide andsilver nitrate are added to mixture 2.

7. What is implied by the observations madewhen sodium sulfide and silver nitrate areadded to mixture 1 ?

8. Explain why the indirect evidence supplied bythe observations made when sodium sulfideand silver nitrate are added to mixture 1 buildour confidence in the dynamic model.

9. Write a balanced chemical and ionic equationfor a reaction that would account for theobservations made when sodium sulfide andsilver nitrate are added to mixture 1 .

10. Based on the observations write a balancedchemical and ionic equation to describe thetwo mixtures. Interpet all symbols used.

11. Describe what you think is happening, on amicroscopic scale, in these mixtures.

Part B

1 . Write an equation to describe the physicalchange that produced the purple vapor whenthe test tube was heated.

2. Write an equation to describe the physicalchange that produced the needle-like crystalsin the upper region of the test tube when itwas removed from the hot water.

3. What evidence suggests the changes in 1 and2 were changes of state and not chemicalchanges?

4. What are two plausible explanations for thefact that the amount of purple vapor remainsconstant even though the test tube remains inthe hot water and there is still solid iodine atthe bottom of the test tube?

5. What evidence makes it hard to accept the"stopped" model?

6. Write an equation to describe the systemwhen the amount of purple vapor remainedconstant. Interpret all symbols used.

7. Describe what you think is happening on amicroscopic scale when the pink color in thetest tube remains constant.

Parts A and B

1. What is meant by the term "dynamic" whenapplied to the concept of chemical equilibrium?

2. Why is a double arrow I used in achemical equation to describe a chemicalsystem at equilibrium?

3. What evidence suggests that the dynamicmodel is the best description of a chemicalsystem at equilibrium?

Further ReadingAlyea, Hubert N., and F.B. Dutton. Tested

Demonstrations in Chemistry. Easton, Penn-sylvania: Chemical Education Publishing Co.,1965.

Parry, Robert W et al. Chemistry: ExperimentalFoundations. Englewood Cliffs, New Jersey:Prentice-Hall, 1970.

Rowley, Wayne R.E. Matter in Balance: ChemicalEquilibrium. Toronto: Wiley, 1979.


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Reaction KineticsObjectivesStudents should be able to:

1 . Describe experimental evidence that supports a molecular distributionof energies within a sample of gas.

2. Define the term activation energy.3. Describe the effect of temperature on the molecular distribution of

energies.4. Use an example to explain what is meant by the term chain

mechanism.5. Draw energy pathways for endothermic and exothermic reactions.6. Label the following on an energy pathway: activation energy; total

energy released; and the net energy required or released by achemical reaction.

7. Use the energy concepts to explain how a reaction and its reverseend up in a dynamic equilibrium.

Program DescriptionProgram 3 begins with a review of the meaning of the term dynamicequilibrium. Then a problem is posed: "Why is it that under one set ofconditions the equilibrium point will involve very few product molecules,but under other conditions it will involve almost complete conversion ofthe reactants to products?" The example used is the chlorine/hydrogen/hydrogen chloride system. A more detailed application of the kinetictheory to chemical systems follows. The molecular distribution of kinetic

energy and its relationship to temperature, activation energy, andreaction mechanisms is introduced and is used to explain howhydrogen and chlorine can exist in the dark, forming very little hydrogenchloride; but add a little daylight and almost all the hydrogen and chlorineis transformed into hydrogen chloride. Energy graphs are used toillustrate energy changes in exothermic and endothermic reactions andhow one influences the other in a system at equilibrium. The programconcludes by asking a question that opens the door to explore theeffects of other factors on the equilibrium point, such as concentrationand pressure. The question provides the logical link to the next program.

Before ViewingReview the meaning of steady state and dynamic equilibrium and whychemists believe most chemical reactions proceed to a state of dynamicequilibrium. Use Activity 1 to show that a collision must occur before areaction can take place. Review the basic assumptions of the collisiontheory and the factors that affect the rate of a chemical reaction. UseActivity 2 to show that many chemical reactions involve a series of steps.These demonstrations set the stage for the introduction of the collisionand mechanism models presented in Program 3.

After ViewingUse Activity 3 to consolidate the ideas presented in the program.Discuss the answers as they will help to form a bridge betweenPrograms 3 and 4.

ActivitiesActivity 1: A Collision Model forChemical ReactionsSet up two stoppered test tubes, each with 2 cmof solid potassium iodide on top of 2 cm of solidlead nitrate. If the test tubes are set up severaldays in advance, a yellow line will appear at thepoint of contact between the two solids.

Pass the two tubes around the class. Allowstudents to shake one but not the other. Theyshould be able to see that yellow lead iodide onlyforms when the two solids come in contact. Thispoint can be made dramatically by adding water tothe test tube, thus dissolving the solids. Whendissolved the contact area is increasedtremendously and the contents of the test tubei nstantaneously turn bright yellow.

This is a good reaction to show studentsbecause most believe reactions do not take placebetween solids. It also illustrates the basic

assumption of the collision theory that reactantsmust come in contact or collide before a reactiontakes place.

Activity 2: A Mechanism Model forChemical ReactionsSince most of the students' experience has beenwith reactions that appear to go in a single step, itis worth the time to show them some of theevidence that suggests that a better way todescribe most chemical reactions is with a series

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1 0

of simple reactions or steps called the mechanismof the reaction.

Demo 1

the dispersion to room temperature. Add sufficientwater to bring the total volume up to 1 L. Add the15 g of sodium hydrogen sulfate(IV) and allow itto dissolve.

Add a drop of aqueous iron(III) chloride to asolution of sodium thiosulfate (concentrations arenot critical) in a petrie dish on the overhead

Demo 2To avoid problems when conducting thisdemonstration, only make as much solution as youwill need for your classes plus a couple of trials;make up the solutions the night before you intendto use them. The concentrations of the threesolutions are as follows:

To produce C make a slurry of the starch andenough water to cover the starch in a mortar. Usethe pestle to mix the two well. Pour the slurry intoapproximately 500 mL of boiling water. Turn offthe heat and stir well. Allow the colloidal starchdispersion that forms to cool while you make upsolutions A and B. If necessary, add ice to cool

Caution: Do not add the sodium hydrogensulfate(IV) to the hot dispersion. If you plan tokeep solution C for more than a day or two add1 g of salicylic acid and 10 mL of ethanol forevery litre of solution.

Label two Erlenmeyer flasks 1 and 2. Use thechart below to obtain the volumes of A, B, C, andwater. Use four graduated cylinders, one for eachof the components, to avoid contamination.

Add the appropriate volumes of A, B, and waterto flasks 1 and 2. Have the students predict therelative rates of the reaction in the two flasks.They should predict flask 1 to be the faster of thetwo because the concentration of B has beenreduced in flask 2.

Review why the total volume was keptconstant. Swirl the flasks to ensure that thereactants are well mixed and add 20 mL of Csimultaneously to each of the flasks. Continue toswirl. An orange precipitate will appear in flask 1first. But just as the students begin to feel thatflush of success that comes from being right, flask2 will turn jet black, followed shortly by flask 1.Who won? How is it possible that the orangeprecipitate could form first in flask 1 but the blackappear first in flask 2?

Divide the students into groups of four. Tellthem to represent the solutions by the letters A,B, and C and ask them to develop a model that willaccount for what they have seen. The following istypical of what they will produce:

When all of B is used up then step 3 can takeplace:

You could tell students the ingredients but thatwould just stifle the creative process as they try torecall appropriate chemical reactions. Rememberthat the purpose of this exercise is not to teach

the mechanism of this reaction but to establishthat it is reasonable to think of a reaction asi nvolving a series of steps.

Activity 3: Reaction KineticsThe following questions will help you review theterms and concepts of the rates of chemicalreactions. Some of them ask you to apply the rateconcepts to a chemical system at equilibrium.

Basic Terms and Concepts1. Define the following terms:

endothermic; exothermic; activation energy;activated complex; rate of reaction; andmechanism of a chemical reaction.

2. Describe a chemical system at equilibrium interms of rates of reactions.

3. (a) List the four factors that determine the rateof a reaction.

( b) Use the collision theory to explain on amicroscopic scale how each of the factorsaffects a chemical reaction.

Step1.2.

forms andslowly disappears, leaving a pale yellow solid,colloidal sulfur. Be prepared to show that theobservations are not due to simple diffusion. Thisis accomplished by adding a drop of iron solutionto a second petrie dish containing just water. Thisis an excellent opportunity to discuss plausibleexplanations and then design experiments to testthem. Do not spend time presenting or trying todevelop the mechanism for this reaction. Just useit to help establish that most chemical reactionscan be pictured as a series of simple steps whichadd up to the overall reaction. The nextdemonstration will help you drive this point home.

projector. A black complex,

Energy Pathways1. (a) Draw and label the energy pathway for the

following reactions:

Label the type of reaction (exo- orendothermic); the position of the reactantsand products; the activated complex; theactivation energy; the total energyreleased; and the net energy produced orused by the reaction.

(b) Use the diagram to explain how a reactionand its reverse end up in dynamicequilibrium.

2. The following energy pathway, Figure 3,describes the energy changes for the reaction:

(a) Complete the following chart for the stepsof the reaction mechanism. Actual valuesare not required. Compare one step to theother using terms such as larger andslower.

(b)

(c)

What is the rate-determining step for thisreaction?Explain why the reaction mechanism is abetter model for what happens on amicroscopic scale during this reaction thanthe net equation.

Molecular Distribution of Kinetic Energy1. (a) Draw and label the curve that shows the

distribution of kinetic energies in a sampleof gaseous molecules at room temperature. (This distribution is called theMaxwell-Boltzmann Distribution.) Showwhat happens to the shape of the curvewhen the temperature is increased. (Use adifferent color.)

( b) Mark the activation energies of the

reactions in question 1(a) EnergyPathways on the distribution curve.

(c) Use the diagram in 1(b) to explain why youwould expect more of the products of theexothermic reaction at equilibrium thanthose of the reverse endothermic reaction.

(d) Use the diagram in 1(b) to predict which oftwo reactions will be most affected by ani ncrease in temperature.

Implications for Equilibrium1. Consider the effects on the rate of formation of

hydrogen iodide if the concentration ofhydrogen were increased. Use these changesto predict what will happen if some hydrogen isadded to a container holding iodine andhydrogen in equilibrium with hydrogen iodide.In your prediction describe the changes in theamounts of iodine and hydrogen iodide atequilibrium.

2. Use an argument based on rates of reactionsto predict the effect of a catalyst on the

1 1

1 2

amounts of reactants and products atequilibrium.

3. The formation and decomposition of hydrogeni odide are second-order reactions.(a) Write a rate expression for each reaction.( b) How will the rates of the two reactions

compare when hydrogen iodide is inequilibrium with hydrogen and iodine?State your answer mathematically and inwords.

(c) Use the mathematical statement to showthat at equilibrium there exists a ratio ofconcentrations of reactants and productsthat equals a constant.

( d) What are the only variables that will alterthe value of the constant developed in (c)?



Huff, George E. Molecules in Motion. Toronto:Wiley, 1976.

Laidler, Keith J. Chemical Kinetics. New York:McGraw-Hill, 1950.

Mahan, Bruce H. University Chemistry. Reading,Massachusetts: Addison-Wesley, 1965.



Reaction TendenciesObjectivesStudents should be able to:

1. State Le ChMelier's Principle.2. Apply Le Chatelier's Principle to predict and explain how the stresses

of heat and pressure affect a given system of equilibrium.

Program DescriptionProgram 4 begins with a review of the concept of dynamic equilibrium.The effect of a stress on a system at equilibrium is introduced using ananalogy. Credit for applying this common-sense approach to chemicalsystems is given to Le Chatelier. The effects of the stresses of heat andpressure are illustrated and accounted for. Le Chatelier's Principle isused to predict the shift in the position of equilibrium. The program endsby establishing a need for knowing the effect quantitatively.

Before ViewingUse the overhead to show the constant macroscopic properties in the

After ViewingReview Le Chatelier's Principle. Have the students review what happensi n the demonstrations you showed them before the program. Redo themif they cannot recall the observations. Complete Activities 2 and 3.Activity 3 provides the link with Activity 1 in Program 5.

ActivitiesActivity 1: Pressure EffectsThe following can be demonstrated with a 50 mLtranslucent plastic syringe on the overhead or itcan be a student activity using 1 mL disposablesyringes held over a white sheet of paper. Thesyringes are sealed by heating the plastic at theneedle end in a flame and squeezing it shut with apair of tongs.

Write the following equation on the chalkboard:

Ask the students for ways to change theconcentration of the participants. Obviouslykeeping the volume constant and adding more ofone will change its concentration. However, it isalso possible to alter the concentration of both bychanging the volume of the container. Ask thestudents which reaction, the forward or thereverse, will be altered the most by a change involume. If their first guess is that both will bechanged equally, do not evaluate it but ask for therate expression for the formation of N02 and forN204, assuming that the reactions occur micro-scopically as written. Write the following on thechalkboard:

Now ask the question again. It should be moreobvious that a compression that halves the volumewould double the concentration of each species,but that the rate of formation of N204 would befour times what it was while the rate of formationfor N02 would only be doubled. Thus there wouldbe a build up of N204 until the rates of formationwere once again equal. At this new equilibriumthere would be more N204 and less N02 thanbefore the compression. Reasoning this way, thestudents would predict that a compression,

1 3

tubes. Have the studentsdescribe the systems on both a macroscopic and microscopic scale.Have them write an equation to describe the processes occurring in thetubes. Review the effects of temperature, a change in volume, andconcentration on the rate of a chemical reaction, and use this data topredict what will happen if these same variables are applied to a systemat equilibrium. Make the predictions and then demonstrate whathappens. As part of the predictions and demonstrations completeActivity 1 . Ask the students to deduce a general principle based on whatthey have observed, then show Program 4.

1 4

quickly push the piston in halfway and hold it. Atfirst glance the only apparent change is anincrease in the intensity of the reddish-browncolor, the opposite to what was predicted. Repeatthe procedure. Push the piston in quickly, hold itat the halfway point, pause until the color remainsconstant, then quickly pull it out. Be careful not tocompletely remove the piston. Careful observationwill reveal that the intensity increases when youpush the piston in and then fades. Similarly,quickly pulling the piston out causes the intensityto fade initially and then darken. Ask the studentsfor plausible explanations but don't evaluate them.Program 4 uses an excellent analogy that will helpthem account for their observations.

Caution: Nitric acid is corrosive; wear protectivegloves. NO2 is toxic; fill the syringes in afumehood. Add small pieces of copper so youonly produce as much gas as you want. Thereaction can be stopped by adding water to thetest tube. Be sure to wash your hands when youare finished.

Because the syringes are very quickly stainedby the gas it is not possible to fill them in advance.This also means that a disposable syringe can beused only once or twice. Thus you will need alarge supply of them. If your doctor gives allergyshots she or he may be willing to save used 1 mL

syringes for you. The large 50 mL syringes canbe purchased through a pharmacy.

Activity 2: Pressure TemperatureEffectsDemo 1Review Le Chatelier's Principle. Have the studentsgive you the equation that describes the equilib-rium in a water/ice mixture. Ask them to predictthe effect of increasing the pressure on thissystem. Have a model of ice available so you canillustrate why water expands when it freezes. Thisi s a good opportunity to review some of the ideasrelated to the Kinetic Molecular Theory.

After the prediction is established support ablock of ice between two iron rings on stands.Hold a thin wire with 1 kg masses on each endover the block. Make sure the masses will not reston the desk even after the wire passes throughthe ice. To avoid a mess it is also advisable tohave some sort of catch basin for the water that isproduced as the ice melts.

Have the students apply their prediction to thissituation and then let the wire hang on the block.The wire passes through the block and themasses clatter to the desk surface, but the blockremains in one piece. Ask the students forplausible explanations. The melting as a result ofi ncreased pressure is easy; students expect that.But why does the water refreeze? If you havei ncluded the energy requirements of this system inthe equation, students should soon realize that thetemperature of the ice drops when it melts. Theenergy required for melting has to come fromsomewhere. In this case it comes from the iceitself. Be sure to point out that at this stage thesystem is not at equilibrium but is attempting toshift to a new equilibrium. Thus when the pressurei s returned to normal, as the wire passes thewater will freeze at the lower temperature,producing heat so the temperature returns to

0 ° C. It will take several minutes for the wire topass through, depending on the thickness of theblock. While you are waiting proceed withDemonstration 2.

Demo 2

ethanol. Shake vigorously until most of the solidhas dissolved. If the solution is not pink, add dropsof water until it just turns pink. Show the studentsthe color by holding the test tube over the stageof the overhead projector or by pouring some ofthe solution into a petrie dish on the overhead.

Write the following on the chalkboard:

Ask the students for evidence that suggeststhe system is at equilibrium.

Ask the students to use Le Chatelier to predictwhat will happen if the system is heated, then heatthe solution gently in a cool Bunsen flame until itchanges color.

Caution: Alcohol is flammable. Place a book overthe mouth of the test tube to smother the flames ifthe ethanol catches fire.

Ask what will happen if the system is cooled,then cool the test tube by immersing it in a beakerof ice water.

Activity 3: Concentration EffectsThis activity leads into Activity 1 in Program 5.Thus it provides a natural bridge between the twoprograms.

The demonstration is best done on an overheadprojector using petrie dishes. It can, of course,

Place about 0.3 g (a few crystals) of crushedin a test tube and add 5 mL of

increasing the pressure, would cause the intensiof the reddish-brown color to fade.

Once the nrediction is estahlished fill theTo demonstrate the effect,syringes with

l arge test tube fitted with a one-hole rubberstopper. Insert a glass tube that has been bentapproximately 100 ° and drawn out to a point intothe rubber stopper. This allows you to place theglass tube into the barrel of the small syringe,l etting it fill from the bottom up. Add a smallvolume of concentrated nitric acid and smallpieces of copper wire to the test tube. Stopperthe test tube and fill the syringes with qas.

To generate the use a gas generator or a

be done on the demonstration desk withtest tubes or beakers but larger volumes ofreagents will be required. The solutions requiredare the same as those required for Activity 1 inProgram 5.

Five petrie dishes

150 mL beaker

Stirring rod

Add 25 mL of KSCN solution to an equalvolume of water in the beaker. Add three or fourdrops of the iron(III) solution into the beaker andstir. The number of drops can vary. The importantpoint is that the color is intense enough to beseen but not so intense that the equilibrium ispushed essentially to completion. Ask thestudents for the ions present in the mixture andthe possible combinations that might produce the

effectively tie it up, so it is removed fromparticipating in the thiocyanoiron(III) equilibrium.Label the petrie dish, "remove Fe 3 +," and add afew crystals of Na2HP04. Now ask if anyone hasnoticed any other changes during the demonstra-tion that you haven't discussed. Someone usuallynotices that the reference has faded. Ask for plau-sible explanations. Some, of course, will think youare imagining things. Show them the fifth petriedish to establish the effect of the stress of addingheat from the projector on the thiocyanoiron(III)equilibrium. If this is not dramatic enough warmsome of the solution in a Bunsen flame.

Further ReadingAlyea, Hubert N. and F.B. Dutton. Tested


Choppin, Gregory R. et al. Chemistry. Morristown,New Jersey: Silver Burdett, 1978.

Othen, Clifford. Rates of Reaction and Equilibria.London: Heinemann Educational Books,1968.


Parry, Robert W. et al. Chemistry: ExperimentalFoundations. Teacher's Guide. EnglewoodCliffs, New Jersey: Prentice-Hall, 1975.


Toon, E.R., and G.P. Ellis. Foundations inChemistry. New York: Holt, Rinehart andWinston, 1973.

1 5

stress involved, and identify the reaction thatcould reduce that stress. After the prediction isestablished demonstrate what happens. Also pointout that the increase in color indicates that someiron(III) ions were present in the solution.

Label the petrie dish used, "add SCN' - . " Usethe same line of questioning for the stress "addFe3+." After establishing that both reactantsare present at equilibrium ask students whatshould be done to get the color to fade. Theyshould reply that removing one of the reactantswill shift the equilibrium to the left and the color willfade. Again, review the reasoning behind thestudents' answer. Point out that this would requirethat the reaction he reversible. Tell them that

forms a complex with that will

1 6

The Equilibrium ConstantObjectivesStudents should be able to:

a smooth curve of "best fit." Read off the graph the "best" values foreach vial number and have all students use these ratios in the remainderof the calculations.

1. Apply Le Chatelier's Principle to predict and explain how the stress ofconcentration affects a given system of equilibrium.

2. State experimental evidence for Le Chatelier's Principle.3. Use the "Dance Hall" analogy or the Kinetic Molecular Theory to

account for Le Chatelier's predictions.4. Write the mass action expression for the equilibrium constant for a

given chemical system.5. Identify evidence that supports the Law of Chemical Equilibrium.6. State the implications of the size of the equilibrium constant.

Program DescriptionProgram 5 reviews the application of Le Chatelier's Principle to chemicalsystems by illustrating and accounting for the changes produced byaltering the concentration of the reactants or products. Le Chatelier'spredictions are then shown to be consistent with the dynamic model ofequilibrium. The program reestablishes the need for having a morequantitative expression for the changes and then uses empirical data toi ntroduce the equilibrium constant. The result is generalized as the Lawof Chemical Equilibrium. The law is then applied to several systems andthe meaning of the size of the equilibrium constant is introduced.

Figure 4

Before ViewingDo Activities 1 and 2. Use Activity 3 in Program 4 as part of the prelab toActivity 1. See the Teacher's Guide to Chemistry: Experimental

Foundations (see page 24) for suggestions to help students cope withthe serial dilution required in this lab. Use Program 5 as part of the post-l ab for Activities 1 and 2. Before the post-lab have students plot theirdepth ratio values versus the vial number on a set of axes drawn on aditto. (See the following graph of typical results.) Run the ditto off soeveryone has a set of the data. The points will vary over a fairly largevertical range. Don't be concerned; there is bound to be considerablevariation due to the way the dilutions were carried out and the methodused to compare the colors. Use the general shape of the curve to draw

Activity 1 enables students to discover that the mass actionexpression for a system at equilibrium is independent of the concentra-tions of the reactants and products. It also introduces the idea of anequilibrium constant using data gathered by the students. Activity 2 doesthe same thing, using data gathered by others.

After ViewingDo Activities 3 and 4. Activity 3 provides students with additional appli-cations of Le Chatelier's Principle and another method for determiningthe equilibrium constant. Activity 4 provides practice interpretingequilibrium constants and using them to make predictions.

0.9

0.8

0.7

0.6

0.5

0.4

0.3

Vial Number

ActivitiesActivity 1: Chemical Equilibrium -A Quantitative Description, Part 1I n this experiment you will take a quantitative lookat the reaction

which was used qualitatively in an earlierdemonstration. This time you will determine theconcentration of each of the ions at equilibrium,and then seek an expression that relates theconcentrations mathematically in a simple,convenient manner. Such an expression wouldenable the quantification of Le Chatelier's predic-tions, a necessity if you wish to be able to predictprofits and losses for an industrial process.

The concentrations are determined colon-metrically. Color intensity of a solution dependsupon the concentration and depth of the solution.If you add water to a cup of tea the color intensityremains constant. This is because the depth of thetea increases as the concentration decreases.The relationship is expressed as follows:

color intensity = kCd

C = concentration expressed in mol/L

d = depth in mm

k = proportionality constant

Thus if two solutions appear to have the samecolor intensity

the depths of solution required to make the colorintensities equal.

To prepare the standard solution in this activityyou will use a small amount of thiocyanate ion,SCN' - , and add a large excess of ferric ion,Fe 3+(aq). You can assume that all of the thio-cyanate ion will be used in forming the complexthiocyanoiron(III) ion, FeSCN 2 +(aq). Thus theconcentration of FeSCN 2 +(aq) in the standard willbe equal to the starting concentration of theSCN' - ( aq). The validity of this assumption will bediscussed in the post-lab discussion.

Apparatus

Distilled water

Five flat bottom vials (must hold 10 mL)

Two 10 mL Mohr pipets

Pipet bulb

25 mL graduated cylinder

150 mL beaker

Diffused light source or white paper

Ruler

Medicine dropper

Method

Add 10.0 mL of 0.2 mol/L Fe(N03)3 to the25 mL graduated cylinder, then fill the

cylinder to the 25.0 mL mark with distilledwater. Pour the solution into a dry, clean 150mL beaker to mix it. Calculate the concentra-tion of this solution.

5. Rinse the second Mohr pipet with water andfinally with a little of the solution in thebeaker.

Caution: You only have 10 mL excess; onlyuse 2-3 mL to rinse the pipet. Do not returnthe rinse solution to the beaker; discard it.

6. Use the second Mohr pipet and bulb totransfer 5.0 mL of the solution in the beakerto vial 2. Calculate the initial concentrationsof the Fe 3 +(aq) and SCN' - i n the 10 mL ofsolution formed in the vial.

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Thus if you know the concentration of a solu-tion in one situation (let's call it the standardsolution), you can calculate the concentration ofthe solution in a second situation by comparing

Line up five clean, dry vials. Label them 1, 2,etc.Use a clean, dry Mohr pipet and pipet bulb tctransfer 5.0 mL of 0.002 mol/L KSCN i ntoeach of the vials.

1.

2.

3.

4.

of your preparation for the activity.13. Wrap a strip of white paper around vials 1 and 2.1 4. Look vertically down through the solutions at

a diffused light source. Use the medicinedropper to remove a dropper full of the standard.

for each vial as part

Note: One mole of yields one moleof

7. Rinse the graduated cylinder with water andfinally with a little (2-3 mL) solution from thebeaker.

8. Pour 10.0 mL of solution from the beakeri nto the graduated cylinder then fill thecylinder to the 25.0 mL mark with distilledwater.

9. Discard any solution remaining in the beakerand rinse it with water. Dry the beaker with apaper towel.

1 0. Pour the solution in the graduated cylinderi nto the beaker to mix it.

1 1. Repeat steps 5 and 6, only this time transferthe solution to vial 3, then repeat steps 7-10.

1 2. This process is called serial dilution. Continuei t until you have 5.0 mL of successively morediluted solution in each vial.Calculate the initial concentrations of

1 8

15. Return some of the standard in drops until thecolors match. Hold the vials close togetherand blink your eyes between "looks" to helpavoid eye fatigue. Put the unused standard inthe dropper into a clean dry beaker, sinceyou may have to use some of this solutionlater.

16. When the intensity of the color in the vialsmatches, record the height of solution ineach tube to the nearest millimetre.

17. Repeat the matching procedure with vials 1and 3, 1 and 4, and finally 1 and 5.

( d) Calculate the equilibrium concentration ofFe3 +(aq) and SCN1 -(aq).

[SCN'"] E = [FeSCN2+]E(e) Now try to find some simple mathematical

relationship between the equilibriumconcentrations that could be used to make

quantitative predictions. Try calculating thefollowing:

Observations

Discussion

Activity 2: Chemical Equilibrium -A Quantitative Description, Part 2This activity provides you with data to discover ageneral relationship for the concentrations of thereactants and products at equilibrium. Perform the

necessary calculations to complete the data tablesfor each of the following systems.

System I

System I involves the following equilibriumbetween ethyl acetate, water, acetic acid, andethanol:

The data in the table was gathered by startingwith known amounts of ethyl acetate and water.The two reactants were placed in an Erlenmeyerflask and swirled. The concentration of acetic acidat equilibrium was determined with the aid of a pHmeter. Since the production of one mole of aceticacid also involves the production of one mole ofethanol (ethyl alcohol), then the concentration ofacetic acid and ethanol at equilibrium would haveto be the same. All trials were done at roomtemperature.

(b) Calculate the radio of depths from the colorcomparison. Subscript "x" is used todenote the vial number.

2. Calculate the following for vials 2 through 5:(a) Calculate the initial concentrations of

Initial meansbefore any chemical reaction has takenplace. Subscript "I" is used to denote initialconcentrations.

1. Use Le Chatelier's Principle to explain why the

reasonable.2. Which of the combinations, (i), (ii), or (iii), gives

the most constant value? To help you decide,calculate the ratio of the largest value to thesmallest for each expression.

3. Restate the most constant expression in wordsusing the terms reactants and products. Ageneral statement of the regularity you havenoted is called the Law of ChemicalEquilibrium.

4. Assume that the reactions for this systemoccur on a microscopic scale as written. Writerate expressions for the forward and reversereactions. Show that the idea of a dynamicequilibrium is consistent with the expressionfound above.

assumption made to calculate

Note: The subscripts "I" and "E" used in thetable refer to initial and equilibrium concentrationsrespectively.

For column "i" calculate:

trations is not a function of concentration (i.e.,it doesn't change when the concentrationschange)?

3. Which of the three expressions is the mostuseful? Explain your choice.

4. Use Le Chatelier's Principle to predict theeffect when water is added to the ethylacetate system at equilibrium. Which trialssupport your prediction?

System If

System II involves the following equilibrium:

The equilibrium concentration of N02 wasmeasured colorimetrically, similar to the methodused in Activity 1. The only difference is that aspectrophotometer was used to measure thecolor intensity.


For column 'T' calculate:

Discussion

gases was put in the vessel (at least that isthe way it appears on the data table).However, this was not the way the trial wasdone. After equilibrium was reached in trial 4the volume of the container was adjusted to 3 L.

( a) What would a change in the volume do tothe pressure exerted on the gases?

(b) Use Le Chatelier's Principle to predict the

1 9

For column "ii" calculate:

1. Show a sample calculation for

2. The two expressions of the concentrations ofthe products and reactants that you calculatedare called mass action expressions. Whichmass action expression for equilibrium concen-

For column "ii" calculate:

For column "iii" calculate:

1.2.

3.

What stress caused the change infrom trial 1 to trial 8? Use Le Chatelier'sPrinciple to predict the effect of the stress.Does the data in the table support yourprediction?I n trials 4 and 5 1.00 mol of each of the

Show a sample calculation for

effect of this stress on the(c) Does the data in the table support your

prediction?4. Which of the three mass action expressions

of concentrations at equilibrium is not afunction of concentration?

5. Which of the three expressions is the mostuseful? Explain your choice.

6. What would be the mass action expression

20

for the following system that would not be afunction of concentration?

Note: Follow the convention of putting theconcentration of the products over thereactants:

endothermic reaction. In general, what effectdoes an increase in temperature have on theequilibrium constant for an endothermicreaction? Write a generalized statement todescribe the effect of temperature on theequilibrium constant. Is your statementconsistent with Le Chatelier's Principle?Explain your reasoning.

1 0. Predict the mass action expression that willequal the equilibrium constant for thefollowing equilibrium:

System III

System III involves the following equilibrium:

The equilibrium concentration of 12 wasmeasured colorimetrically, similar to the methodused in Activity 1. The only difference wasthat a spectrophotometer was used to measurethe color intensity. All trials were carried out at698 K.


For column "i" calculate:

Discussion

2. (a) Which mass action expression was equalto the equilibrium constant?

(b) Did you predict correctly in question 9 ofSystem II? If not, where did you go wrong?

3. The symbol for the equilibrium constant iscapital "K." Write the mass action expressionfor the equilibrium constant for the followingequilibrium:

Note: The convention is to put the concentra-tion of the products over the reactants.

4. The mathematical expression you wrote in theabove question is a statement of the Law ofChemical Equilibrium. State the law in words.

5. The equilibrium constant is independent ofconcentration. What does this mean?

6. What two variables cause K, the equilibriumconstant, to vary?

Activity 3: Chemical EquilibriumAntimony(III) chloride reacts with water to formwhite insoluble antimony(III) oxychloride. Thishydrolysis reaction can be represented by thefollowing equilibrium:

6 mol/L with respect to HCI. In this solution youcan assume the concentration of SbOCI is almostzero.

Do discussion questions 1-4 on this page aspart of your preparation for this activity.

Apparatus10 mL graduated cylinder

50 mL buret

Two 125 mL Erlenmeyer flasks

20 mL 0.5 mol/L SbC13 in 6 mol/L HCI

25 mL 6 mol/L HCI

1. Show a sample calculation for

7. The value of the mass action expression thati s independent of concentration is called theequilibrium constant. What is the equilibriumconstant for the ethyl acetate equilibrium?What is it for the nitrogen dioxide equilibriumat 142 42 0 C? At 11 5 0 C? At 70OC?

8. What two variables will alter the value of theequilibrium constant? Use the data presentedi n Systems I and II to answer this question.

9. The reaction was written as an

The effects of adding water and hydrochloricacid to the system will be investigated. Also K, theequilibrium constant for this reaction, will bedetermined.

One litre of the solution used contains 0.5 mol

trations are 0.5 mol/L with respect tcand 6 mol HCI. That is, the solution concen-

and

Caution: The soluble compounds of antimony arealmost as poisonous as the compounds ofarsenic. Wear protective gloves and wash yourhands thoroughly after this activity.

Method1 . Transfer 5.0 mL of antimony(Ill) chloride

solution into an Erlenmeyer flask from a buretat one of the dispensing stations yourteacher has set up.

2. Add 5 mL water from a graduated cylinderand swirl the mixture. Record any changes.

3. Add an additional 10 mL of water in portionsof 2 mL at a time until a total of 15 mL havebeen added. The 15 mL includes the 5 mLadded in step 2. Swirl the mixture and recordany changes with each addition.

4. Use the graduated cylinder to add 2 mL of6 mol/L HCI. Swirl and record any changes.

5. Add an additional 6 mL in 2 mL portions,swirling the contents and recording anychanges each time.

6. Dispose of the mixture as your teacherdirects.

7. Add 5.0 mL antimony(III) chloride solutionfrom the buret at the dispensing station to aclean, dry 125 mL Erlenmeyer flask.

8. From another buret add water to the flaskuntil the white precipitate begins to form.Swirl the flask continuously as you add thewater slowly. Try not to add too much. Youcan use your results from steps 2 and 3 as aguide. Record the volume of water requiredto the nearest 0.1 mL.

9. Dispose of the mixture as your teacherdirects.

10. Calculate the concentration of HCI in the totalvolume of solution when the precipitate wasjust starting to form.

11. Calculate the concentration of antimony(III)chloride present in the total volume,

Discussion

2. Predict the effect of adding water on theantimony(III) chloride hydrolysis equilibrium.Describe any changes you would expect tosee.

3. Predict the effect of adding hydrochloric acidon the antimony(III) chloride hydrolysisequilibrium. Describe any changes you wouldexpect to see.

4. (a) Write the mass action expression for theequilibrium constant for the antimony(III)chloride.

( b) Since the concentration of a pure solid orliquid is directly proportional to its density itwill only change with temperature, Thus theconcentration of a pure solid or liquid inequilibrium with other substances remainsconstant. Since the concentration of thewater and the antimony(III) oxychloride donot change they can be grouped with theequilibrium constant to form a new constantcalled K'. Write the mass actionexpression for K'.

(c) What would K' be equal to in terms of K?5. (a) At what point in the addition of water (steps

2 and 3) was the system at equilibrium?(b) Why did the system not appear to change

after several additions of water?6. Account for any change or lack of change that

occurred when the acid was added.

7. Based on your observations, is the assumptionmade in step 12 reasonable? If it isn't, how willit affect your value of K'?

8. (a) What is the value for K'?( b) What does the value for K' imply about the

hydrolysis of antimony(Ill) chloride?(c) Pedict the relationship between your value

for K' and the accepted value. Explainyour reasoning.

Activity 4: Equilibrium ApplicationsThe following questions will help you learn to applythe concept of equilibrium to chemical systems.

1. Use Le Chatelier's Principle to predict theeffects of heat, pressure, and concentrationon the following chemical systems:(a) The production of methanol (methyl

alcohol). Stresses: add heat, increasenressure by compressing the container.

(b) The production of chlorine. Stresses:remove heat, decrease pressure byexpanding the container, add oxygen, adda catalyst, add a noble gas that doesn'treact with any participant in theequilibrium.

Stresses: add heat. increase the pressure

2 1

removal of methanol.

(c) The formation of

of add solid

1. Explain why it is reasonable to assume that[SbOCI] = 0 in the solution that contains

assuming that a negligible amount of SbCI3was used in the formation of the whiteprecipitate.

1 2. Determine the value of K', the modifiedequilibrium constant (see discussion question4), from the concentrations of HCI andfor the solution in which the SbOCI was juststarting to form.

what will be the concentration ofhydrogen iodide at equilibrium?

(b) At a certain temperature hydrogen iodidei s 20 percent dissociated. Determine theconcentration of each component presenti n the mixture at equilibrium, if 0.50 mol ofHI are placed in a 1 .00 L vessel at thistemperature.

(c) Determine the equilibrium constant for theformation of hydrogen iodide at thiscertain temperature.

(d) Is the certain temperature lower or higherthan 699 K? Explain your reasoning.

( e) In a mixture of hydrogen iodide,hydrogen, and iodine at 699 K, the partialpressures are 70.9 kPa, 2.03 kPa, and2.03 kPa respectively. Will there be anychange in the partial pressure if themixture is maintained at 699 K? If so, willHI be consumed or formed?

(f) 1.00 mol of H2 and 2.00 mol of 12 gas areput into a 10 L container at 699 K.Determine equilibrium concentrations ofthe three gases at equilibrium.

(g) What would happen to the equilibriumconcentrations in (f) if the volume werecompressed to 5.0 L?

6.7 g of S02CI2 and 101.3 kPa of Cl2 areplaced in a 1.0 L bulb at 375 K.

( d) Compare parts (b) and (c). In order to getthe situation in (c), what stress must beapplied to the system in (b)? Use LeChatelier's Principle to predict the effect.Compare your predictions with thecalculated concentrations. Are theyconsistent?

9. Ammonium hydrosulfide decomposes in thefollowing manner:

( a) Determine the partial pressures ofammonia and hydrogen sulfide at equilib-rium when excess solid ammonium hydro-sulfide decomposes in an excavatedchamber at 25 ° C.

(b) Use Le Chatelier's Principle to predict theeffect of injecting ammonia into thesystem at equilibrium.

(c) Check your prediction by completing thefollowing problem:Excess solid ammonium hydrosulfide isplaced in a flask with 50.7 kPa ofammonia. What will be the partialpressures of ammonia and hydrogensulfide at equilibrium?

1 0. The equilibrium constant for the reaction of"A" and "B" to produce "C" is 4.0.

( a) Determine the equilibrium concentrationof each species if the starting conditionsare 0.50 mol of "A" and 1.0 mol of "B" ina 10.0 L container.

( b) If it took 10 minutes to reach equilibrium,accurately sketch a concentration timegraph for the attainment of equilibrium.

(c) On separate sketches of the graph, showthe effects of the following on the originalsystem at equilibrium:

(i) adding a catalyst(ii) adding more "A"(iii) raising the temperature of the system(iv) reducing the volume of the container

to 5.0 L(v) adding the noble gas neon.



CHEM Study. Chemistry: An ExperimentalScience. Teacher's Guide. San Francisco:W.H. Freeman, 1963.

Choppin, Gregory R. et a/. Chemistry. Morristown.-Silver Burdett, 1978.

NSCM Project Practical Activities Cl 0. Principlesof Chemical Equilibrium. Milton, Australia:Jacaranda Press, 1973.

Othen, Clifford. Rates of Reaction and Equilibria.London: Heinemann Educational Books,1968.

Parry, Robert W. et a/. Chemistry: ExperimentalFoundations. Englewood Cliffs, New Jersey:Prentice-Hall, 1970.




2 3

is 243.1 kPa when the partial pressures areexpressed in kilopascals.(a) 6.7 g of S02CI2 are placed in a 1.0 L bulb

at 375 K. Determine the pressure

it dissociated.exerted by the assuming none of

(b) Determine the partial pressures ofat equilibrium.and

, c) Determine the partial pressures ofat equilibrium ifand

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The Haber ProcessObjectivesStudents should be able to:

1. Describe the Haber Process for the production of ammonia.2. Discuss the societal implications of the Haber Process.3. Apply the concepts of rate and equilibrium to the development of the

conditions for an industrial process.

Program DescriptionProgram 6 begins by establishing the historical events that created theneed for the Haber Process. It then goes through a development of theprocess using the concepts from the previous programs. The develop-ment requires not only an application of the concepts of chemicalequilibrium but also of those of reaction kinetics. The program ends witha brief description of the "rewards" reaped by Haber for his mastery ofchemical equilibrium.

AmmoniaProduced

(%)

0 20 40 60 80

Figure 5 Pressure (mPa)

Before ViewingHave students complete Activity 1. They are asked to apply LeChatelier's Principle in selecting conditions for an industrial process.However, because an industrial process continually removes the wantedproducts, the reaction never reaches equilibrium. Thus in most casestheir predictions will be quite different from the actual conditions used.Because an industrial process is better described as a steady state thanan equilibrium system, an application of factors that affect the rate of areaction are also important. As the program points out, both conceptsmust be applied.

Take up the students' answers to the questions but don't evaluatetheir thoughts on question 6. Program 6 will help them with this one.

After ViewingDo Activity 2 and consider using Activity 3 as one of the projects thatstudents may choose to work on this term.

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ActivitiesActivity 1: Le Chatelier andthe Haber ProcessExperimental studies have shown that thepercentage of ammonia formed from hydrogenand nitrogen at equilibrium varies as a function ofthe pressure exerted on the system as a whole.The results for a series of different temperaturesare shown in Figure 5.

Discussion

1 . Write an equation to describe the formation ofammonia. The name given to the industrialprocess using this chemical reaction is theHaber Process.

2. Is the formation of ammonia an endo- orexothermic reaction? Use data from the graphto support your decision.

3. Use Le Chatelier's Principle to predict theeffects of the following stresses on theammonia equilibrium:

(a) removing ammonia( b) increasing the applied pressure(c) raising the temperature(d) adding a catalyst( e) adding an inert substance.

4. State evidence from the graph to support yourpredictions in 3(b) and (c).

5. List in general terms the conditions that wouldyield the most ammonia.

6. (a) At 20 mPa, what is the yield of ammoniaat:(i) 473 K(200 ° C)(ii) 873 K(600 ° C)

(b) At which temperature would you choose torun the Haber Process? Why?

7. The conditions universally employed byindustry are 873 K(600 ° C), 20 mPa, and theuse of a catalyst. Suggest plausible reasonsfor this.

Activity 2: Le Chatelier andOther Industrial ProcessesThe Production of PolystyreneAs part of the reaction sequence in whichpolystyrene is produced from benzene, ethylbenzene is dehydrogenated to produce styrene,the monomer of polystyrene. The followingequation describes the reaction:

1 . Use Le Chatelier's Principle to predict theeffects of the following on the styreneequilibrium:

(a) increasing the applied pressure(b) removing styrene(c) reducing the temperature( d) adding a catalyst(e) adding an inert substance.

2. Based on your predictions in question 1,select the conditions that would maximize theyield of styrene.

3. (a) Why are the conditions selected inquestion 2 unlikely to be those used byi ndustry?

(b) What additional considerations must bemade when selecting conditions tomaximize the yield of an industrial process?

4. In practice, the pressure of the ethyl benzeneis kept low, as Le Chatelier would suggest, butan inert substance - super-heated steam - isadded to keep the total pressure of the mixtureat atmospheric pressure. What advantagedoes this method have over running thereaction without the steam at a pressure lessthan atmospheric?

5. The super-heated steam serves two otherfunctions: First it supplies energy to thesystem. Why is this desirable? Second, itreacts with any carbon formed as a byproductat high temperatures, preventing the carbonfrom contaminating the catalyst. What productswill be formed as a result of the reactionbetween carbon and steam?

6. What is the advantage of using super-heatedsteam at 873 K(600°C) rather than ordinarysteam?

7. In practice, an iron oxide catalyst is used.Explain why this is desirable.

8. The process is run at 873 K(600 ° C) with ani ron oxide catalyst. The conversion is onlyabout 35 percent complete. The yield couldbe increased by raising the temperature.Suggest two plausible reasons why this is notdone.

9. The ethyl benzene (boiling point 136 ° C) thathas not reacted and the styrene (boiling point146 ° C) are separated by fractional distillation.The styrene, of course, is used to make poly-styrene. If you were the industrialist, whatwould you do with the ethyl benzene?

The Production of Sulfuric Acid -The Contact Process

1 . Write a set of equations to describe the majorreactions taking place in the contact process.

2. The system sulfur dioxide, oxygen, and sulfurtrioxide is in equilibrium at 101.3 kPa and400 ° C in the presence of a catalyst. Statewhether the amount of sulfur trioxide would bei ncreased, decreased, or unchanged by eachof the following. Give reasons for youranswers.

(a) decreasing the applied pressure at 400 °C(b) adding 5 mol of oxygen at 400 ° C(c) decreasing the concentration of SO2 at

400°C(d) adding 2 mol of helium, V and T constant(e) removing the catalyst.

3. Use Le Chatelier's Principle to suggest threeways to increase the yield of sulfur trioxide.

4. State the conditions you would use as ani ndustrialist. Explain the reason for choosingthese conditions if they differ from those yousuggested in question 3.

5. List the actual industrial conditions used (e.g.,temperature; pressure; catalyst; for removingsulfur trioxide). Explain the reason for choosingthese conditions if they differ from those yousuggested in question 4.

Activity 3: IndustrialProcess ProjectThe task is to present, in an attractive fashion, thechemistry involved in an industrial process.

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Pictures, drawings, color, letter size, etc., can beused to present the information in as attractive apackage as possible. In effect you are going toproduce an ad for the industrial process youselect.

Description of the Task

1. The industrial process chosen must involvesome chemical reaction-

2. The maximum area covered by your ad shouldbe 1500 cm2.

3. It must be possible to display your ad on thewall.

4. The ad must contain the following information:(a) name of industry(b) chemical reactions involved(c) conditions under which reactions are run(d) location of plants(e) source of raw materials.

5. The following may be included if desired:(a) pictures of plants(b) drawings or flow charts of the process(c) labels used on final products(d) anything else you feel makes your ad more

attractive.

Further ReadingAshman, A., and G. Cremonesi. Sulfuric Acid.

London: LongmanslPenguin Books, 1968.

Bradford, Derek. Chemistry and the World FoodProblem. London: Heinemann EducationalBooks, 1971.

Haber, L.F. The Nitrogen Problem. London:LongmanslPenguin Books, 1966.





OrderinginformationTo order the videotapes or this publication, or foradditional information, please contact one of thefollowing:

TVOntario Sales and LicensingBox 200, Station 0Toronto, Ontario M4T 2T1(416) 484-2613

Untied StatesTVOntarioU.S. Sales Office901 Kildaire Farm RoadBuilding ACary, North Carolina27511Phone: 800-331-9566Fax: 919-380-0961E-mail: [email protected] g

Example Industrial Processes

Several industrial processes are listed. They arei ncluded to give you an idea of the range ofprocesses from which you can select. The list isnot meant to be restrictive; the possibilities areendless.1. Production of iron2. Production of aspirin3. Film developing and printing4. Electroplating5. Brewing

Videotapes BPNProgram 1: Steady Unsteadiness 240701Program 2: Dynamic Equilibrium 240702Program 3: Reaction Kinetics 240703Program 4: Reaction Tendencies 240704Program 5: The Equilibrium Constant 240705Program 6: The Haber Process 240706

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