x-le chatelier

Introduction

Many chemical reactions reach a state of equilibrium if conditions are right. In an equilibrium system, forward and reverse reactions occur at equal rates so that no net change is produced. When equilibrium is reached by a reaction in a test tube, it appears that changes have stopped in the tube. Once equilbrium has been reached, is it possible to produce further observable changes in the tube? If so, can you control the kinds of changes? If not, why are further observable changes impossible? You will observe several chemical systems in this laboratory activity. A careful study of your observations will enable you to answer these questions.

Purpose

To study factors which can disturb an equilibrium system.

Safety Considerations

Wear protective glasses and an apron at all times. Avoid skin contact with solids and solutions. Exercise caution and use proper technique to handle hot materials safely. Dispose of all solutions in the containers provided by your teacher. Wash your hands before leaving the laboratory.

Procedure

1. Obtain a test tube rack, six small (13 x 100 mm) test tubes that are clean but don't have to be dry, and a test tube clamp. The test tubes should be placed open end up in the test tube rack.

2. Prepare a hot water bath: Half-fill a 250 mL beaker with tap water. Start to heat the water (as your teacher directs) so that the water will be near boiling when you are ready to use it.

3. Prepare an ice water bath: Fill a 250 mL beaker with crushed ice. Add enough tap water to make "slush".

4. Set up a data table with column headings as indicated below (The last column will be completed after data have been collected.)

System Disturbance Observed Change Direction of Shift1 2 etc.

5. As you set up equilibrium systems and add disturbances to them in the procedure, enter appropriate information in each of the first three columns of your data table.

6. Mix chemicals in test tubes by holding the top of the tube with one hand while you flick the bottom of the tube with your other hand until the tube contents.

System 1: Iron(III) and thiocyanate

Setting Up the Equilibrium

1. Half-fill the first tube in your rack with distilled water.2. Add two drops of 0.1 M Fe(NO3)3 and two drops of 0.1 M

KSCN to this tube. Mix the contents thoroughly.3. If the contents of the tube are not red-orange, repeat

Step 2 until the solution is red-orange.4. Divide the red-orange solution in the first tube among six

tubes so each tube contains the same volume.

Chemical Equation for the Equilibrium System

Fe3+(aq) + SCN-(aq) FeSCN2+(aq) + heatColorless Colorless Red-orange from Fe(NO3)3

from KSCN

Disturbing the Equilibrium

1. Leave Tube 1 undisturbed; use it as a control.2. Use a clean, dry spatula to add a small crystal or two of

solid iron(III) nitrate, Fe(NO3)3, to Tube 2. Mix.

Under Disturbance on your data table, record what you did or added to the system to cause the change you

observed. In this and all other observations, pay particular attention to color and color change. Always compare with the control tube or you may miss slight color changes. Phrase your Observed Change so the kind of change you observe is indicated, e.g., "lighter red" or "from grey to pink."

3. Use a clean, dry spatula to add one or two small crystals of solid potassium thiocyanate, KSCN, to Tube 3. Mix. Record observations.

4. Add 5 drops of 0.1 M sodium hydroxide, NaOH, to Tube 4. Mix, observe, and record.

5. Use a test tube clamp to place Tube 5 in a hot water bath. When the contents of the tube are hot, observe and record.

6. Use a test tube clamp to place Tube 6 in an ice water bath. When the contents of the tube are cold, observe and record. (Data check: Obtain your teacher's initials.)

7. Discard all test tube contents in the waste container provided by your teacher. Do not pour anything in the sink. Rinse the tubes with tap water; remove as much water as possible by shaking before standing the tubes upright in the test tube rack. Follow these same disposal and rinsing procedures after you complete each system below.

System 2: Bromothymol blue


1. Half-fill three test tubes with distilled water.2. Add three drops of bromothymol blue indicator to each

tube. Mix thoroughly.

Chemical Equation for the Equilibrium

Bromothymol blue is a weak organic acid with a complex formula. For our purpose, its formula can be abbreviated to HBb.

HBb(aq) H+(aq) + Bb-(aq)

Yellow Colorless Blue

(Green can be observed if approximately equal amounts of yellow and blue forms are present.)


1. To Tube 2 add two drops of 0.1 M hydrochloric acid, HCl, and mix. Observe and record.

2. To Tube 3 add two drops of 0.1 M sodium hydroxide, NaOH, and mix. Observe and record.

3. Explore what happens when you now add NaOH to Tube 2 or HCl to Tube 3. See whether your observations are in agreement with observations you have already recorded.

System 3: Complex Ions of Copper(II) (Cu2+)


1. Half fill a test tube with 1.5 M copper(II) chloride, CuCl2, solution.

2. Divide so five tubes contain approximately equal volumes. Equilibrium has already been established in the solution.


CuCl42-

(aq)+ 4 H2O(l) Cu(H2O)4

2+(aq) + 4 Cl-(aq) + heat

Green soln

Colorless Light blue soln Colorless


1. To Tube 2 add a small quantity (the size of a rice grain) of solid calcium chloride, CaCl2. Mix to dissolve the solid. Repeat the addition and dissolving of solid CaCl2 until no more solid will dissolve. Observe and record.

2. To Tube 3 add enough ethyl alcohol, C2H5OH, to triple the volume of the solution. Mix, observe, and record.

3. Place Tube 4 in a hot-water bath. When the solution is hot, observe and record.

4. Place Tube 5 in an ice-water bath. When the solution is cold, observe and record.

System 4: Dinitrogen tetroxide (N2O4)


Dinitrogen tetroxide, N2O4, can decompose into nitrogen dioxide, NO2, a reddish brown poisonous gas. So that you may work with these substances safely, your teacher will provide two sealed tubes each containing a mixture of these subtances. Equilibrium between N2O4 and NO2 has already been established in the tubes.


N2O4(g) + heat 2 NO2(g)

Colorless Reddish brown


1. (Caution: N2O4 and NO2 in the sealed glass tubes are poisonous. Handle the tubes carefully to avoid breaking the tubes and releasing the gases.) Place one sealed tube containing the equilibrium system in a hot water bath. When hot, compare to the unheated tube and record.

2. After removing the tube from the hot water bath, cool it under running cold tap water. Then place the tube in an ice-water bath. When cold, compare to the unchilled tube and record.

System 5: Complex Ions of Cobalt(II) (Co2+)


1. Half-fill a test tube with 1.5 M cobalt(II) chloride, CoCl2.2. Divide the solution so five tubes contain approximately

equal volumes. Equilibrium has already been established in the solution.


heat + Co(H2O)62+(aq) + 4 Cl-(aq)

CoCl42-

(aq)+ 6 H2O(l)

Red Colorless Blue Colorless


1. To Tube 2 add a small quantity (the size of a rice grain) of solid calcium chloride, CaCl2. Mix to dissolve the solid. Repeat the addition and dissolving of solid CaCl2 until no more solid will dissolve. Observe and record.

2. To Tube 3 add enough acetone, CH3COCH3, to double the volume of the solution. Mix, observe, and record.

3. Place Tube 4 in a hot water bath. When the solution is hot, observe and record.

4. Place Tube 5 in an ice water bath. When the solution is cold, observe and record.

5. Wash hands thoroughly before leaving the laboratory.

Data Analysis, Concept

1. To complete the fourth column on the right side of your data table (headed Direction of Shift), decide whether each disturbance caused the equilibrium system to shift left or right. Record the direction of shift in this column. How do you decide direction of shift? Consider the equilibrium system

A BYellow Green

2. If a disturbance causes the system to become more yellow, chemists would say that the equilibrium position has shifted to the left because the system must have moved to produce more of the yellow molecules shown on the left side of the chemical equation. If the system shifted to the right you would observe more green in the system. The direction of shift is "right". Use these ideas to decide and record the direction of shift caused by each disturbance.

3. Use your data table to find all cases where a disturbance was caused by heating. After you have found all of these cases, answer the following:

a. How does the direction of shift relate to the side of the chemical equation on which the heat term is written?

b. Write a rule which would allow you to predict how other equilibrium systems would shift when disturbed in this way.

4. Use your data table to find all cases where equilibrium systems were disturbed by cooling.

a. How does the direction of shift relate to the side of the chemical equation on which the heat term is written?


5. Use your data table to examine all cases where a disturbance was caused by increasing the concentration of a substance already present in the equilibrium system. Hint: Adding solid Fe(NO3)3 to System 2 increases the concentration of Fe3+(aq) and NO3

-(aq) when the solid dissolves. Adding HCl solution to System 3 increases the concentration of both H+(aq) and Cl-(aq) in the system. Write a rule which would explian how the direction of shift relates to the side of the chemical reaction on which the substance with increased concentration is written.

6. In some cases the equilibrium system was disturbed by decreasing the concentration of a substance in the system. Usually this is done by adding another substance not involved in the equilibrium which reacts with a substance in the system, changing it to a different substance. For example, in System 1 you added 0.1 M NaOH (containing aqueous Na+ and OH- ions). OH-reacts with Fe3+ to form the precipitate Fe(OH)3(s). This decreases the concentration of Fe3+(aq) remaining in the solution. Concentration can also be decreased by adding another solvent (acetone or alcohol) to dilute the water in the system. Identify substances whose concentration is

decreased in as many cases as you can. For each, explain what causes the concentration of a particular substance to decrease. Write chemical equations where possible.

The equation for the example above is:

Fe3+(aq) +3 OH-

(aq)Fe(OH)3(s)

For each case involving a decrease in concentration, identify the substance that is decreased in concentration, on which side of the equation this substance is found, and which way the equilibrium is observed to shift.

7. Consider cases where equilibrium was disturbed by decreasing the concentration of a substance in the equilibrium system.

a. How does the direction of shift relate to the side of the chemical equation on which the substance with altered concentration is written?


8. Write a general rule that would cover all of the types of disturbances you have observed. Write your rule so it can be used to predict the effect of any temperature or concentration disturbance on an equilibrium system.

Imply, Apply

1. Explain why no visible changes can be observed when a system is at equilibrium. Express your answer in terms of the rates of the forward and reverse reactions.

2. What effect would each of the following have on the rate of a reaction?

a. increasing the concentration of a reactantb. decreasing the concentration of a reactantc. increasing the temperature of the systemd. decreasing the temperature of the system

3. Are the rates of both the forward and reverse reactions still equal immediately after an equilibrium system is

disturbed? Support your answer with observations you have made.

4. Which direction of shift would you observe if only

a. the rate of the forward reaction is increased?b. the rate of the reverse reaction is increased?c. the rate of the forward reaction is decreased?d. the rate of the reverse reaction is decreased?

5. Can shifts in equilibrium systems be explained by considering the effect of a disturbance on the separate rates of forward and reverse reactions? Support your answer with evidence.

6. Consider the equilibrium system

PbSO4(s) + 2 I-(aq) + heat PbI2(s) + SO42-(aq)

White Colorless Yellow Colorless

7. Indicate the Direction of Shift and the Predicted Observable Change, given the following:

a. add NaI solutionb. add AgNO3 solutionc. Ag+(aq) + I-(aq) AgI(s)d. add Ba(NO3)2 solutione. Ba2+(aq) + SO4

2-(aq) BaSO4(s)f. add Na2SO4 solutiong. heat the systemh. cool the system

8. Ammonia, NH3, is produced industrially by the reaction:

N2(g) + 3 H2(g) 2 NH3(g) + heat

9. Discuss at least three ways a chemist could shift this equilibrium system so that a greater amount of ammonia is produced at equilibrium.

10. Assume that you add water to a certain ionic solid and stir, but observe that the solid will not completely dissolve. The solubility equilibrium of the ionic solid is represented by the equation:

AB(s) + heat A+(aq) + B-(aq)

11. Discuss two things you could possibly do to get the solid to dissolve more completely.

Teachers Guide

Preparing for the Laboratory Activity

Major Chemical Concept Level Expected Student Background Time Safety Materials Advance Preparation

Conducting the Laboratory Activity

Pre-Lab Discussion Teacher/Student Interaction Answers to Data Analysis Answers to Imply, Apply Post-Lab Discussion Possible Extensions

Assessing the Laboratory Learning

Laboratory Practical Paper/Pencil Items Self-Assessment (Practice Test) Answers to Self-Assessment (Practice Test) Final Assessment Answers to Final Assessment Performance Checks

Preparing for the Laboratory Activity

Major Chemical Concept

LeChatelier's Principle is a useful generalization regarding the response of equilibrium systems to changes in temperature and concentration.

Level

While this laboratory activity is qualitative, it requires excellent reasoning ability. It is most appropriate for general and advanced chemistry.

Expected Student Background

Student should be able to:

Handle and mix solids and solutions Write and interpret chemical equations Explain the particulate model of matter Predict how temperature and concentration changes

affect the rates of chemical reactions Describe the characteristics of the equilibrium state (both

macroscopic and molecular level) Describe the dissolving and precipitation of ionic solids at

the molecular level

Time

One to two class periods

Students can stop work after any of the five systems and can complete their work during the next laboratory period. It is essential that students observe all five systems or they will not have all the data they need to complete the Data Analysis.

Safety

Chemicals with specific hazards:

0.1 M NaOH - corrosive to skin and clothing 0.1 M HCl - corrosive to skin and clothing

Solid CaCl2 - skin irritant, avoid breathing dust Ethyl alcohol - flammable NO2 - poisonous fumes if demonstration tubes are broken Acetone - flammable

Materials

Non-Consumables (per lab team)

2 Beakers, 250 mL 1 Dropping pipet Equipment for heating a hot water bath: hot plate or

burner, ring stand, ring and wire gauze 1 Spatula 1 Test tube clamp 1 Test tube rack 6 Test tubes, 13 x 100 mm

Non-Consumables (for class)

2 NO2/ N2O4 Demonstration tubes Beaker (600 mL) of hot water kept on a hot plate Beaker (600 mL) of ice-water mixture

Consumables (per lab team)

Distilled water, 25 mL Ice, crushed (to fill a 250 mL beaker) Litmus paper, 2 pieces of red, blue and full range

Consumables (for class)

The following liquids and solids should be divided among four sets of labeled dropping bottles. Solutions are aqueous and should be prepared with distilled water unless otherwise indicated. 20 mL are sufficient for five classes.

0.1 M Fe(NO3)3, 8 g Fe(NO3)3•9H2O per 200 mL solution (add concentrated HNO3 drop by drop with mixing to make colorless solution )

0.1 M HCl, 2 mL concentrated HCl per 200 mL solution

0.1 M KSCN, 2 g KSCN per 200 ml solution 0.1 M NaOH, 0.8 g NaOH per 200 ml solution 1.5 M CoCl2, 71 g CoCl2•6H2O per 200 mL solution 1.5 M CuCl2, 51 g CuCl2.2H2O per 200 mL solution Acetone, 100 mL Bromothymol blue indicator, 10 mL (buy as a prepared

indicator or prepare by dissolving 0.1 g in 16 mL 0.01 M NaOH + 234 mL water). [Color change pH range is 6.0 (yellow) - 7.6 (blue)]

Ethyl alcohol, 100 mL, 95% (denatured) Iron(III) nitrate, Fe(NO3)3•9H2O, solid, 10-20 g Potassium thiocyanate, KSCN, solid, 10-20 g Calcium chloride, CaCl2 (abstruse), solid, 10-20 g

Advance Preparation

Directions for preparing solutions are indicated above. All solutions are quite stable except bromothymol blue indicator solution, which must be fresh. Conduct a lab check on System 2 (bromothymol blue and distilled water) to be sure the mix is green. Varying pH values in distilled water can cause problems. If the mix is not green, prepare a quantity of bromothymol blue and water that is green by adjusting the pH appropriately. Let students use the prepared mixture instead of mixing it themselves. Hot water and ice baths for NO2 should be ready for student use at the start of the period.

Conducting the Laboratory Activity

Pre-Lab Discussion

Before students begin laboratory work, raise the question, "Can anything be done to disturb an equilibrium system so a net change can be observed?" Let students suggest possibilities if they can. Students should explore the reversible nature of litmus as follows:

1. Give the chemical change for litmus as:

HLit(aq) H+(aq) + Lit-(aq)

2. Tell students that they can establish this equilibrium merely by dipping either red, blue, or full range litmus paper strips in distilled water.

3. Suggest that students try to disturb the equilibrium system of the moist litmus by dipping their moistened litmus strips in 0.1 M HCl and 0.1 M NaOH.

4. Suggest that students explore whether the color change of litmus is reversible.

5. Ask students to draw a conclusion about the reversibility of the litmus. See if they can propose tentative explanations of what they have seen and/or relate their observations to the chemical equation they have been given. Don't confirm or deny any explanations they propose. Rather, tell students they will explore several chemical equilibrium systems in this activity and should be able to judge the correctness of their proposed explanations for themselves. Note: Besides providing a good student experience in proposing hypotheses, the litmus activity helps eliminate the common student misconception that red and blue litmus are two totally separate, unrelated substances. As a result, they also think that once litmus has changed, you throw must the litmus paper away because it can't "change back" again.

Teacher/Student Interaction

1. Students should work in pairs to share work, learn organization and efficient division of labor, and to discuss ideas, observations, and predictions with each other. If students work together to interpret and understand the data, the additional in-class time is well worth the investment.

2. Tell students they can work on the five equilibrium systems in any order, making access to materials easier for them.

3. NO2 tubes, hot water bath, and ice bath should be easily accessible to students. Suggest that students use the NO2 tubes whenever they are unused by other students.

4. At the end of System 1 in the procedure, students encounter a Data Check that requires your initials. This gives you an opportunity to confirm that students are comparing each "disturbed" tube with the control tube and are noting relative color changes "from red-orange to dark red".

Answers to Data Analysis

1. See table below

Anticipated Student ResultsSystem Disturbance Observed Change Direction of Shift

1

Fe(NO3)3(s)

KSCN(s)

0.1 M NaOH(aq)

heat

cool

darker red

darker red

lighter red, orange

lighter red, orange

slightly darker red

right

right

left

left

right

20.1 M HCl

0.1 M NaOH

to yellow

to blue

left

right

3

CuCl2(s)

C2H5OH

heat

cool

more green

more green

more green

more blue

left

right

left

right

4heat

cool

more reddish-brown

less reddish-brown

right

left

5

CaCl2(s)

acetone

heat

cool

more blue

more blue

more blue

more red

right

right

right

left

2. Students should examine the effect of heating Systems 1, 3, 4, and 5.

a. Shift is to the other side (side opposite the heat term).

b. When heated, equilibrium systems shift to the side opposite where the heat term appears in the chemical equation.

3. Students should examine the effects of cooling Systems 1, 3, 4, and 5.

a. Shift is to the side of the equation on which the heat term appears.

b. When cooled, equilibrium systems shift to the side of the chemical equation where the heat term appears.

4. Students sometimes have difficulty recognizing cases involving an increase in concentration, because they forget that ionic compounds exist as dissociated ions in solution. The disturbance involving an increased concentration of one of the species in the equilibrium system are:

System Disturbance Actual Effect1 Fe(NO3)3(s) Increases conc. of Fe3+(aq)1 KSCN(s) Increases conc. of SCN-(aq)2 0.1 M HCl Increases conc. of H+(aq)3 and 5 CaCl2(s) Increases conc. of Cl-(aq)

5. Direction of shift is to the side opposite where the formula of the substance (whose concentration was increased) is shown in the equation.

6. Students will really need your help here. They have difficulty because they can't determine what chemical reaction is caused by the disturbance and therefore, which substance decreases in concentration. You may need to provide the following chemical equations.

System Disturbance Reaction which occurs Actual effect Shift1 0.1 M NaOH Fe3+(aq) + 3OH-(aq) Fe(OH)3(s) Fe3+(aq) conc. decreases left2 0.1 M NaOH H+(aq) + OH-(aq) H2O(l) H+(aq) conc. decreases right3 C2H5OH dilution of water H2O conc. decreases left5 acetone dilution of water H2O conc. decreases right

7. Direction of shift is to the side in which the formula of the substance (whose concentration is decreased) is shown.

8. Equilibrium systems shift in such a way as to minimize a disturbance. For increased heat or concentration, the system shifts to use up a portion of the added heat or increased concentration. For decreased heat or concentration, the system shifts to create more of the heat or substance that was decreased. Note: Significant post-lab discussion may be needed to help some students generalize to this extent.

Answers to Imply, Apply

1. Forward and reverse reactions occur at equal rates at equilibrium; the effect of the forward reaction is exactly cancelled by the reverse reaction, occurring to the same extent.

2. a) increases b) decreases c) increases d) decreases3. No. Visible changes occur when a disturbance is added,

thus the rate of the forward and reverse reactions cannot be equal.

4. a) right b) left c) left d) right5. Yes. Disturbances of temperature and concentration can

alter the rates of the reaction. If the forward reaction is faster than the reverse reaction, a shift to the right will be observed because products will be produced by the forward reaction faster than they can be consumed by the reverse reaction.

6. Direction of shift; Direction of Shift, Predicted Observable Change

a. right (I- increases); more yellowb. left (I- decreases); more whitec. right (SO4

2- decreases); more yellowd. left (SO4

2- increases); more whitee. right; more yellowf. left; more white

7. Increase concentration of N2(g); increase concentration of H2(g); decrease concentration of NH3(g) or cool

8. Add more water to decrease the concentration of dissolved ions or add heat to increase the rate of dissolving. Both disturbances shift the equilibrium to the right.

Post-Lab Discussion

Students should understand the connection between their work in this activity and the classical statement of LeChatelier's Principle: An equilibrium system shifts in a way that partially relieves a stress applied to it. Students attempted to write a "general rule" in Data Analysis, Question 7. Various student responses should be shared and discussed as a class so that the consensus "general rule" is equivalent to LeChatelier's Principle. Students should be made aware that "disturbance" is the same as "stress" and that the common name for this "general rule" is "LeChatelier's Principle".

Review again the types of disturbances (stress) that students have studied, how each type affects an equilibrium system, how rate changes can be used to explain shifts, and how each type of disturbance (stress) fits the more general statement of LeChatelier's Principle.

Possible Extensions

1. Discuss other factors which affect reaction rates:

a. pressure (gaseous systems),b. surface area (heterogeneous systems),c. nature of reactants andd. catalysts. Discuss in terms of whether or not these

factors can cause disturbances that make equilibrium systems shift.

2. Relate equilibrium to tendencies to minimum energy and maximum entropy.

3. (Advanced level) Relate observations to calculations involving multiple equilibria.

Assessing the Laboratory Learning

Laboratory Practical

See written exercises. You may allow students to check their responses to practice test Questions 5 - 8 by laboratory observations.

Paper/Pencil Items

A self-assessment (practice test) and a final assessment are provided for your use. If time is short, you may elect to omit the practice test. If so, you can use the two sets of questions as alternate test forms.

Chemical systems involved in practice test Questions 5-8 can be observed by students in the laboratory and may be used as a lab practical. If students observe these systems, Nature provides the answers rather than you. Such observations are highly recommended. If students do the practice test on a non-graded basis, they can learn from their mistakes and assess their strengths and weaknesses before they take the final assessment.

If Questions 5-8 of the practice assessment are used as a lab practical, student work can be made more efficient by:

1. preparing equilibrium mixtures for each of the four systems in sufficient quantity and providing them to the students. Equilibrium mixtures should be shifted to a "middle" color so students can observe shifts in either direction, or

2. preparing a sample of each of the four equilibrium systems and asking students to match the color of the samples when they set up each equilibrium.

Needed if students observe systems involved in Practice Test Questions 5-8:

0.1 M Cu(NO3)2, 6 g Cu(NO3)2•6H2O per 200 mL solution 1 M NH3, 14 mL concentrated NH3 per 200 mL solution

3 M NH3, 41 mL concentrated NH3 per 200 mL solution 1 M HNO3, 13 mL concentrated HNO3 per 200 mL solution 0.1 M HNO3, add 20 mL 1 M HNO3 to 180 mL distilled

water 1 M HCl, 17 mL concentrated HCl per 200 mL solution Solid Na2S2O3•5H2O, 250 g

Self-Assessment (Practice Test)

1. Describe how you could set up a working model of an equilibrium system using ten students (each with a ping-pong ball and paddle) and a ping-pong table with a net.

2. Describe how you could disturb the ping-pong equilibrium in Question 1; predict the result of the disturbance.

3. Consider the equilibrium:

A(g) + B(g) AB(g) + heatRed Colorless Blue

4. You set up this equilibrium in a sealed flask and observe that the system has a purple color. You place it on the counter overnight. When you return in the morning the contents of the flask are completely blue. What might have happened to the flask overnight? Explain.

5. Do you agree or disagree with the following statement? "Any disturbance in an equilibrium system that alters the rate of either the forward or reverse reaction will cause the equilibrium system to shift." Support your answer with written reasons.

6. The equilibrium system:

Cu2+(aq) + 2 OH-(aq) Cu(OH)2(s)Light blue soln Colorless Light blue solid

7. can be set up by mixing 0.1 M Cu(NO3)2 and 0.1 M NaOH. After the equilibrium is established, predict the direction the equilibrium would shift if you

a. added 0.1 M NaOH, sodium hydroxide.b. added 0.1 M HNO3 to the original equilibrium

system.

(HNO3 releases H+; then H+ + OH- ---> H2O)

8. The equilibrium system

CuCl42-(aq) + 4 NH3(aq) Cu(NH3)4

2+(aq) + 4 Cl-(aq)Green Colorless Dark blue Colorless

9. can be set up by mixing 1.5 M CuCl2 with 1 M NH3. After the equilibrium is established, predict what color change you would observe if you:

a. add more 1 M NH3

b. add 1 M HNO3: (H+ + NH3 NH4+)

c. add solid CaCl2d. add 1 M HCl

10. If 0.1 M Cu(NO3)2 is mixed with 0.1 M NaOH, a light blue precipitate of Cu(OH)2 forms. If 3 M NH3 is added to this mixture, the following equilibrium results:

Cu(OH)2(s) + 4 NH3(aq) Cu(NH3)42+(aq) + 2 OH-(aq)

Light blue Colorless Dark blue Colorless

11. Explain why 3 M NH3 can be used to dissolve solid Cu(OH)2.

12. Given the equilibrium system:

heat + Na2S2O3•5H2O(s) 5 H2O(l) + 2 Na+(aq) + S2O32-(aq)

White Colorless Colorless Colorless

13. Describe what you would see (include color and phase change) if

a. solid Na2S2O3•5H2O were heated.b. the result of Part (a) were cooled.

After completing the self-assessment questions, check your answers with your teacher to see how well you understand disturbances of temperature and concentration and how they affect an equilibrium system. Your teacher may supply you with the materials needed to observe the systems in Questions 5-8. If so, you can check your answers to these questions with your own observations.

Answers to Self-Assessment (Practice Test)

1. Divide students so groups on opposite sides of the table continue to hit ping-pong balls across the net. Equilibrium is reached when balls move across the net in equal numbers in both directions.

2. Disturb by

a. adding more ping-pong balls or more people with paddles to one side.

b. removing ping-pong balls or people or paddles from one side.

c. any other reasonable answer (blindfold some people, etc.).

3. Equilibrium shifted to the right as room temperature decreased.

4. True because if the rate of either forward or reverse reaction is altered, the reaction in one direction will be faster than in the other direction.

5. a) right (increased OH- concentration) b) left (decreased OH- concentration)

6.

a. more dark blue (increased NH3 concentration causes shift right)

b. more green (decreased NH3 concentration causes shift left)

c. more green (increased Cl- concentration causes shift left)

d. more green (decreased NH3 concentration and increased Cl- concentration causes shift left)

7. Increased concentration of NH3 causes a shift to the right. Solid Cu(OH)2 reacts faster than it forms, thus decreasing the amount of solid Cu(OH)2.

8.

a. Some or all of the white solid changes to a colorless liquid.

b. Some or all of the colorless liquid changes to a white solid.

Note: Thoughtful students may wonder about the effect of reducing all reactant and product concentrations in an equilibrium system when a solution is added to disturb the system. Congratulate the students on their insight and suggest that they can assume dilution due to volume increase is negligible when they make predictions requested in this self-assessment.

Final Assessment

1. Describe two ways a chemist can decrease the rate of a reaction.

2. Carbon monoxide gas, CO, and oxygen gas, O2, react to form carbon dioxide, CO2, and heat. Under appropriate conditions, the reaction is reversible and reaches equilibrium. Write a chemical equation that describes this equilibrium.

3. Chemicals are mixed and an observable color change occurs. When changes are no longer observed, what can you infer about the rates of the forward and reverse reactions? Explain.

4. When the equilibrium system

A BRed Colorless

5. is heated, the intensity of red color decreases. The heat term is missing from the equation. On which side of the equation should the heat term be written? Explain.

6. Given:

X(aq) + Y(aq) Z(aq)Colorless Colorless Green

7. The first four times that a student adds X to this equilibrium system, it becomes darker green each time. The fifth time the student adds X, the system does not become noticeably darker green. What can you conclude about what changes in the concentration of Y during the five additions of X?

8. Given:

heat + CaCl2(s) Ca2+(aq) + 2 Cl-(aq) White Colorless Colorless

9. What observations would you expect to make if this equilibrium system were a) heated? b) cooled?

10. Consider the equilibrium system:

CaCO3(s) + 2 H+(aq) Ca2+(aq) + CO2(g) + H2O(l)

11. Predict the direction this equilibrium system would shift if you

a. add 0.1 M HClb. add 0.1 M NaOHc. add 0.1 M Ca(NO3)2

d. allow CO2(g ) to escape12. Due to the presence of carbonic acid, H2CO3, the

equilibrium present in an unopened bottle of soft drink is relatively complex. A simplified representation of the equilibrium system is:

H2CO3(aq) CO2(g) + H2O(l)

a. Which direction does this equilibrium shift when the closed bottle is opened briefly and then re-sealed? Why?

b. Dissolved carbonic acid is responsible for the "tangy" taste of a soft drink. Explain why a bottle of soft drink loses its tangy taste if left opened.

Answers to Final Assessment

1. Lower the temperature of the reaction system or decrease the concentration of one or more of the reactants.

2. 2 CO(g) + O2(g) 2 CO2(g) + heat3. Rates of forward and reverse reactions are equal. Change

in one direction is offset by an equal change in the opposite direction.

4. The heat term should be written on the left. When heated, equilibrium systems shift to the side opposite the heat term in the chemical equation.

5. Equilibrium has shifted so far to the right that very little Y remains in the system.

6.

a. Some solid dissolves (disappears from view), producing a colorless solution.

b. More white solid appears.7.

a. Right (increased concentration of H+)b. Left (H+ + OH- --> H2O, causing decreased

concentration of H+)c. Left (increased concentration of Ca2+)d. Right (decreased concentration of CO2)

8.

a. Equilibrium shifts right (decreased concentration of CO2).

b. Concentration of carbonic acid decreases.

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