Experiment 2

Experiment #2: Determination of an Equilibrium Constant Expression

Experiment #2:The Determination of an Equilibrium Constant ExpressionRevised 6-9-10

Required SkillsTo perform this experiment you will need to have mastered all of the skills from Experiment #1

and also be able to properly …

• obtain solid samples.

• weigh solid samples.

• perform a vacuum filtration.

• perform a gravity filtration.

• perform a titration.

You should thoroughly review this material in the basic skills document available on the Moodle

site. Remember, that the instructor and TAs are constantly assessing your level of preparation.

Students who are disorganized, confused, and uninformed will receive lower laboratory scores.

BackgroundThis laboratory experiment is focused on determining the equilibrium constant expression

for the solubility of an ionic solid. We will consider the process of solid calcium iodate converting

into aqueous calcium and iodate ions...

Ca(IO3)2(s) ↔ Ca2+(aq) + 2IO3

- (aq) (Rxn 1)

When a system reaches chemical equilibrium the concentrations of the chemical species

become constant with respect to time. Furthermore, it has been found that specific ratios of the

chemical concentrations raised to precise exponential powers are always the same (no matter

what the starting concentrations were). The magnitude of the correct ratio is called the

equilibrium constant (often denoted as K) and the mathematical relationship linking the

concentrations to the constant is called the equilibrium constant expression. You will empirically

determine the equilibrium constant expression for Rxn 1 in this laboratory experiment.

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In order to narrow the focus, you will set out to determine which of the following ratios of

concentration is best used as the equilibrium constant expression for Rxn 1…

a. K = ([Ca2+][IO3-])/[Ca(IO3)2]

b. K = [Ca2+] [IO3-]

c. K = [Ca2+]2 [IO3-]

d. K = [Ca2+] [IO3-]2

Your careful experimental work will allow you to answer this multiple-choice question with

certainty.

The laboratory can be broken down into two main parts. In part I, you will first make

known quantities of two of the pertinent species, Ca2+ and Ca(IO3)2. In addition, you will

determine the concentration of thiosulfate in a solution that will serve as your titrating agent. In

part II, you will make eight mixtures with different starting amounts of Ca2+ and Ca(IO3)2 and allow

them to equilibrate. You will then perform a series of titrations to determine the equilibrium

concentrations of Ca2+ and IO3-. And finally, you will use the concentration and titration data to

experimentally determine the correct equilibrium constant expression and the equilibrium

constant.

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Procedure

Part I

Obtain a kit from the stockroom. You will have to return it at the end of the lab period. Make sure

that every item that was originally in the kit gets returned and that all items are clean and rinsed

with deionized water.

Break your team up into two groups. One group will work on Goal #1 while the other group will

work on Goal #2.

Goal #1: Make solid Ca(IO3)2.

You need to make approximately 10 g of solid Ca(IO3)2. You will do this by combining IO3- with a

slight excess of Ca2+ and allowing the following reaction to proceed…

Ca2+(aq) + 2IO3

- (aq) → Ca(IO3)2(s)

You will add calcium in the form of Ca(NO3)2•4H2O(s). When this tetrahydrate of calcium nitrate is

placed in pure water, it fully dissociates to produce calcium ions, nitrate ions, and water

molecules…

Ca(NO3)2•4H2O(s) → Ca2+(aq) + 2NO3

-(aq) + 4H2O(l)

You will add iodate in the form of KIO3(s). When potassium iodate is placed in pure water, it fully

dissociates into potassium and iodate ions…

KIO3(s) → K+(aq) + IO3

- (aq)

PRIOR TO CLASS CALCULATION: Calculate the masses of Ca(NO3)2•4H2O(s) and KIO3(s)

required to make 10 g of Ca(IO3)2(s). Calculate the amount of Ca(NO3)2•4H2O(s) needed so that the Ca2+ is in excess of the iodate ion concentration by 20% (i.e., calculate a weight of calcium nitrate that is 20% higher than the minimum required to produce 10 g of calcium iodate). Thoroughly document all of your calculations for both Ca2+ and IO3

- on the

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laboratory data sheet. Get the approval of the laboratory instructor or TA prior to actually

weighing the solids and beginning the reaction. You will lose the data points for these calculations if it is determined that you are copying from other groups.

Weigh out the correct amounts of Ca(NO3)2•4H2O(s) and KIO3(s). Important note: this is potassium iodate not potassium iodide. There is a difference! Make sure you use the correct solid. Weigh these solids in small, clean, dry beakers. These masses do not have to

agree exactly with your calculated values. Masses within 0.2 g are perfectly acceptable.

Add the KIO3 to a 250 mL beaker containing approximately 150 mL of deionized water (use the

graduations on the side of the beaker to estimate the water volume). Gently warm the mixture

on a hot plate until the KIO3 is fully dissolved. Then add the calcium nitrate. Continue to heat for

15 minutes to coagulate the precipitate. Stir occasionally.

Then let this heterogeneous mixture cool for 40 minutes. Stir for 30 s every 5 minutes.

Isolate the precipitate by vacuum filtration. Make sure you know how to properly perform a

vacuum filtration (see the CHM 157-158 Skills document if you do not).

Assemble the vacuum filtration apparatus. Wet the filter paper with deionized water. First, slowly

pour the liquid that is above the precipitate into the filtering apparatus with the suction on (leave

as much solid in the beaker as possible). Then carefully transfer the solid to the filter paper with

the help of a spatula. Wash the precipitate that is collected on the filter paper by pouring small

amounts of water over the precipitate (approximately 4 washes of 30 mL each). Proper washing

is critical to the success of this experiment. Make sure to allow all of the washing solution to drain

between individual washings.

Transfer the wet Ca(IO3)2(s) to a clean, labeled, 50 mL beaker. Cover 90% of the top of the

beaker with parafilm. Leave a small opening so that the solid can dry further. Store the solid in a

safe place in your drawer until the next lab period.

Goal #2: Make a series of solutions with known Ca2+ concentrations.

This step is critical. Make sure whoever is in charge of this goal is the "detail person" of the

group. You will first make your own Ca2+ stock solution then you will generate two accurately

diluted solutions through a serial dilution procedure.

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PRIOR TO CLASS CALCULATION: Calculate the mass of Ca(NO3)2•4H2O(s) required to generate a 0.1XX M Ca2+ solution in a 100 mL volumetric flask (this should be done prior to coming to lab). It is written as 0.1XX M because you don’t know the actual concentration of

calcium until you weigh the solid and generate the solution, but here is what you do know: (a.) it

should be close to 0.1 M; (b.) the final concentration should be known out to 3 significant figures;

and (c.) you should use the actually generated concentration for all of your subsequent

calculations (not just 0.1 M). Document your calculation of the required mass of

Ca(NO3)2•4H2O(s) on the data sheet and get the approval of the instructor or TA prior to making

the solution.

Weigh out the appropriate amount of Ca(NO3)2•4H2O(s) needed to make the 0.1XX M Ca2+

solution. Carefully transfer all of the solid to a 100 ml volumetric flask. Dilute to mark with

deionized water. Mix thoroughly throughout the dilution process. The solution should be

transferred to a labeled 250 mL erlenmeyer flask after all solid is completely dissolved. Label the

solution as “solution A”.

A serial dilution will be used to generate two additional solutions. Perform a serial dilution using

“solution A” to generate a solution with 25% of the initial concentration and a volume of 100.0 mL.

This new solution is labeled as “Solution B” (approximately 0.02XX M). Remember to mix

throughout the dilution process. Transfer “Solution B” to a labeled 250 mL Erlenmeyer flask.

Then use “Solution B” to generate a 100.0 mL of a solution with a concentration that is 25% of

“Solution B”. This is “Solution C”. Transfer “Solution C” to a labeled 250 mL Erlenmeyer flask.

PRIOR TO CLASS CALCULATION: Document your proposed dilution scheme on the laboratory data sheet (this should be done prior to coming to lab).

Put parafilm over the tops of the three 250 mL erlenmeyer flasks containing Solution A, B and C.

Store them in your lab drawer until the next lab period.

Goal #3: Titrate known quantities of iodate with thiosulfate

In part II of this experiment (i.e., the part you will do next lab period) you will determine the

concentration of iodate in each equilibrated solution. You will do this by quantitatively converting

the iodate into triiodide and then titrating the triiodide with thiosulfate. In order to obtain accurate

measurements of [IO3-] you need to know the thiosulfate concentration of the titrating solution. So

in this lab period we will determine the concentration of the titrating solution. This can be done by

titrating known quantities of iodate. This process is called "standardizing the titrant". You need

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to know how to properly titrate with a buret. This information is available in the Basic Skills

documents associated with this course.

Obtain approximately 150 ml of sodium thiosulfate (Na2S2O3) stock solution in a clean, dry,

labeled beaker.

Obtain about 50 ml of the 0.01XX KIO3 stock solution (record the exact concentration from the

label on the bottle). Also obtain approximately 10 cm3 of solid KI (this is approximately 2

teaspoons), 20 mL of 1 M HCl, and 10 mL of 0.2% starch solution.

Set up a buret (review the basics of proper titration technique described in both skills documents

prior to coming to lab). Rinse the buret with several portions of deionized water. Test the buret

and make sure it does not leak and is not plugged.

Further rinse the buret with small portions (approximately 10 ml) of the titrant (in this case it is the

thiosulfate solution).

Clean a 10.00 mL pipet with deionized water, and then rinse it with a small portion (approximately

2 mL) of the stock 0.01XX KIO3 solution.

Performing a Titration

Fill the buret with the thiosulfate titrant. Drain some titrant out until the level is a milliliter or so

below 0.0 mL. Make sure you don't have a bubble at the bottom stem of the buret (below the

stopcock).

Carefully pipet 10.00 mL of the KIO3 solution into a clean 250 mL Erlenmeyer flask. Add 10 to

20 ml of deionized water. Add about 1 cm3 (i.e., 1 mL) of solid KI. Important note: KI is potassium iodide not potassium iodate. There is a difference! Add 3 ml of 1 M HCL.

Swirl to mix. The solution should be a dark yellow-brown color. The color comes from the

formation of triiodide, I3-. via

IO3- (aq) + 8I-

(aq) + 6 H+(aq) → 3I3

-(aq) + 3H2O(l) (Rxn 2)

The iodate is the limiting reactant (i.e., it determines how much triiodide is produced). That is

why it is carefully pipeted, while the I- and H+ are added in an approximate fashion.

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Record the initial volume on the buret (estimate the volume out to the hundredths of a mL not to

just a tenth of a mL).

Put the Erlenmeyer containing the triiodide underneath the buret and begin slowly adding

titrant. Placing a piece of white paper under the flask is often helpful in observing the color of

the solution. Swirl the flask contents occasionally (or you can checkout magnetic stirring plate

and stir bar and have constant stirring). The solution should turn from a dark yellow-brown to

pale yellow as the titration progresses and the thiosulfate reacts with the triiodide. Stop once

the solution is a very faint yellow (sort of like watered down Mountain Dew).

Add 1 ml of starch solution. This will enhance the color of the triiodide.

Slowly add titrant until the solution just turns colorless. Note: Take your time. Adding titrant

one drop at a time is appropriate near the endpoint. The color change should be dramatic (or as

dramatic as you get in a chemistry laboratory experiment) if you go slowly. Don’t add anymore

titrant once the color is gone.

Record the final volume of titrant in the buret.

Perform three more titrations in the same fashion. The required volumes of titrant should be very

similar if you are doing the solution preparation and titrations properly (i.e., they should be within

0.3 mL of each other). If you are sloppy or inattentive, there will be greater scatter in your results.

The key is to be precise with pipeting the IO3- and precise with adding the titrant.

Goal #4: Determine the concentration of the thiosulfate titrant

Use the data from each titration to calculate the concentration of the thiosulfate solution.

Calculate the average of the 4 values. Show this result to the instructor or TA prior to leaving the

laboratory.

Here are some hints for performing this calculation: Thiosulfate reacts with trioidide according to

2S2O32-+ I3

- → S4O62- + 3I- (Rxn 3)

The stoichiometry between thiosulfate and iodate can be determined by combining Rxn 2 with

Rxn 3. The combined reaction is...

IO3- + 6H+ + 6S2O3

2- → 3S4O62- + I- + 3H2O (Rxn 4)

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Rxn 4 tells you that the number of moles of thiosulfate required to reach the endpoint relative to

the starting number of moles of iodate. The thiosulfate concentration is determined by making

the following conversions...

volume iodate → moles iodate → moles thiosulfate → molarity thiosulfate Scheme 1

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Part II (Second Lab Period)

Goal #1: Prepare a series of mixtures of solid Ca(IO3)2 and Ca2+ solutions.

Weigh the calcium iodate solid prepared in the previous lab period. If the mass is less than 8 g,

then you should inform the instructor.

Get eight 250 mL Erlenmeyer flasks. Clean them thoroughly and rinse them with deionized

water. Label them A1, A2, B1, B2, C1, C2, D1, and D2. Split 80% of the solid Ca(IO3)2 equally

among the flasks ending with a “1” (e.g., A1, B1, etc.). This would be 2 g per beaker if you

obtained a total of 10 g of precipitate. Divide the remaining 20% Ca(IO3)2 equally among flasks

ending with a “2”. This would be 0.5 g per beaker if you obtained a total of 10 g of precipitate.

When done with this step you will have produced 4 flasks with a "large" amount of precipitate

(i.e., each will have approximately 20% of the original Ca(IO3)2 solid) and 4 flasks with a "small"

amount of precipitate (i.e., each will have approximately 5% of the original Ca(IO3)2 solid). Note:

You do not need to carefully weigh out each amount of calcium iodate.

You will now make 8 different mixtures with the calcium solutions and the flasks containing solid

calcium iodate. Below is a description of how to make these solutions.

Flasks A1 and A2: Add approximately 35 ml of Solution A to each flask.

Flasks B1 and B2: Add approximately 35 ml of Solution B to each flask.

Flasks C1 and C2: Add approximately 35 ml of Solution C to each flask.

Flasks D1 and D2: Add approximately 35 ml of deionized water to each flask.

Let these solutions stand for 30 minutes. Swirl each flask by hand for 30 s every 5 minutes.

Some solid should remain in each flask. If no solid remains then notify the TA or instructor.

Break up your team into two groups. One group will focus on Goal #2 and the other group will

focus on Goal #3.

Goal #2: Isolate the liquid above the precipitate for each solution.

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You need to set up a filtration apparatus. Unlike the filtering you did in Part I of this experiment,

this time you are interested in collecting the liquid not the solid. So you will do a gravity filtration.

Set up a gravity filtration apparatus (see the Skills document for information). Use fresh filter

paper and a clean dry funnel for each solution. Carefully pour the solution through the filter.

Collect the filtered liquid (i.e., the filtrate) in a clean, dry, labeled 50 mL beaker or erlenmeyer

flask. You need at least 25 ml of each solution, so be careful. Keep the same labeling system as

before (e.g., the #1 filtered solution should be from the #1 reaction mixture). Slowly transfer all of

the liquid then, all of the solid, to the filter funnel.

When the filtration is completed, the solid Ca(IO3)2 and filter paper should be transferred to the

chemical waste container in the fume hood.

Give the labeled beakers or flasks containing the filtered solutions to the part of your team

performing the titrations as they become available.

Goal #3: Determine the aqueous concentration of IO3- in each sample by titration.

Use the titration procedure from the previous lab part to determine the iodate concentration in

each filtered solution. The only difference is that instead of using the 10.00 mL 0.01XX M IO3-

stock solution you will use 10.00 mL portions of each filtered solution. All other reagents and

quantities are the same.

Determine the iodate concentration using the required titrant volume, the concentration of

thiosulfate determined in part I, and the volume of IO3- sample titrated. This can be easily done

by using a scheme similar to the reverse of Scheme 1 and plugging in the above data obtained in

the lab.

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Goal #4: Calculate the equilibrium concentration of Ca2+ from the starting [Ca2+] and the

equilibrium [IO3-]

You determined the equilibrium concentration of iodate in the previous section, but you did not

directly determine the equilibrium concentration for Ca2+. You need to know both to complete the

experiment. Fortunately, [Ca2+] can be calculated from the data you have gathered. The key is to

set up an equilibrium table like that shown below...

Ca(IO3)2(s) ↔ Ca2+(aq) + 2IO3- (aq)

start solid [Ca2+]o 0

change +x +2x

equilibrium solid [Ca2+]o+x +2x

At the start, the calcium concentration was given by the calcium nitrate concentration in the

solution poured over the precipitate (see results from the first part of this lab). This concentration

is shown as [Ca2+]o in the table. The starting iodate concentration was zero. Some of the solid

Ca(IO3)2 dissolved causing the calcium concentration to increase by an unknown quantity x and

the iodate concentration to increase by 2x. At equilibrium the calcium concentration is given by

[Ca2+] = [Ca2+]o+x and the iodate concentration is given by [IO3-] = 2x. So by solving for x we can

determine that the equilibrium calcium concentration is given by...

[Ca2+] = [Ca2+]o + 0.5 [IO3-]

where [IO3-] is the iodate concentration determined by titration.

Calculate the calcium ion content in each solution using the average value of iodate concentration

determined in the previous goal.

Goal #5: Determine the equilibrium constant expression.

Recall from the introduction that we want to determine which of the following relationships is the

actual equilibrium constant expression.

a. K = ([Ca2+][IO3-])/[Ca(IO3)2]

b. K = [Ca2+] [IO3-]

c. K = [Ca2+]2 [IO3-]

d. K = [Ca2+] [IO3-]2

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Answer a is the only relationship that includes a contribution from Ca(IO3)2. Thus, the remainder

of the answers (i.e., b through d) can only be true if the amount of solid present does not have a

significant effect on the [Ca2+] and [IO3-] concentrations. This can be determined by directly

comparing the results between beakers that had the same calcium solution but a different amount

of calcium iodate (e.g., compare the results of Beaker #1 with Beaker #2, compare the results of

Beaker #3 with Beaker #4, etc.). If solution with the same starting calcium ion concentration but

different amounts of solid produce essentially the same final values of [Ca2+] and [IO3-] then

relationship a cannot be correct.

Assuming that your data shows that the amount of solid does not affect the equilibrium values of

[Ca2+] and [IO3-] you can rule out answer a and focus on b through d. These answers are all

products of the calcium and iodate solution concentrations.

There are a variety of ways that you can determine which product is the best equilibrium constant

expression. Perhaps the most intuitive way is to calculate the values of [Ca2+] [IO3-], [Ca2+]2 [IO3

-],

and [Ca2+] [IO3-]2 for each solution. The correct expression will generate values with the least

amount of relative scatter. This can be determined by calculating the range of the values and

dividing by the average of the values. The result indicates the relative amount that the product

fluctuates between the solutions. Decide which product represents the most consistent result,

and fully discuss the basis of your decision. Use the average of this product as your experimental

value for the equilibrium constant.

Assembling Your Lab Report

You need to include the following items in the order listed

1. Grade Sheet

2. Pre-Lab Questions (If returned to you)

3. Data Forms

4. Graphs and Spread Sheet Tables

5. Answers to Post Lab Questions

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Data Sheets

Part I: Goal #1

Calculation for the required mass of Ca(NO3)2•4H2O(s)

Calculation for the required mass of KIO3(s)

Actual masses

Ca(NO3)2•4H2O :__________________________

KIO3 :__________________________

Summary of Observation during the Precipitation and Filtration Process

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Part I: Goal #2

Calculation for the required mass of Ca(NO3)2•4H2O(s)

Actual Mass Ca(NO3)2•4H2O :__________________________

Describe how you made the three calcium solutions (i.e., solutions A, B, and C).

Calculation for the concentration of “Solution A” (0.1XX M Ca2+ concentration)

Calculation for the concentration of “Solution B”. (0.02XX M Ca2+ concentration)

Calculation for the concentration of “Solution C”. (0.006XX M Ca2+ concentration)

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Part I: Goal #3

Concentration of IO3- stock solution:_______________

Calculation of moles of IO3- added per titration

Theoretical calculation of moles of S2O32- required to reach the endpoint

#1 #2 #3 #4

initial buret reading

final buret reading

titrant volume delivered

[S2O32-]

Average [S2O32-]:_________________________

Example Calculation of S2O32- Concentration. Use the data from titration #1

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Part II: Goal #3

Soln A1 Soln A2 Soln B1 Soln B2

initial buret volume

final buret volume


Moles S2O32- Added

Moles IO3- Reacted

[IO3-] in Mixture

Soln C1 Soln C2 Soln D1 Soln D2

initial buret volume

final buret volume


Moles S2O32- Added

Moles IO3- Reacted

[IO3-] in Mixture

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Example Calculations Using Data From Titration #1

Moles of S2O32- Added:

Moles of IO3- Reacted:

IO3- concentration in mixture:

Part II: Goal #4

Solution # Soln A1 Soln A2 Soln B1 Soln B2

Initial [Ca2+]

Final [IO3-]

Final [Ca2+]

Mixture # Soln C1 Soln C2 Soln D1 Soln D2

Initial [Ca2+]

Final [IO3-]

Final [Ca2+]

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Part II: Goal #5

Mixture # Soln A1 Soln A2 Soln B1 Soln B2

[Ca2+][IO3-]

[Ca2+]2[IO3-]

[Ca2+][IO3-]2

Mixture # Soln C1 Soln C2 Soln D1 Soln D2

[Ca2+][IO3-]

[Ca2+]2[IO3-]

[Ca2+][IO3-]2

Mixture # Average Range Range/Ave

[Ca2+][IO3-]

[Ca2+]2[IO3-]

[Ca2+][IO3-]2

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Final Conclusion: Write down the equilibrium constant expression that is correct based on your

data. Describe how you came to this conclusion. Make sure you comment on how you ruled out

(or didn’t rule out) the appropriateness of answer “a” of the multiple choice question. Also

indicate the value of the equilibrium constant.

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Names: 1)_______________ 2)________________ 3)_______________ 4)________________Group #__________ Date_______________Prelab Questions Hint reading the entire procedure (both Part I and Part II) is helpful.

1. Be able to show the TA or instructor that your group has made all of the “Prior to Class Calculations”.

2. A student pipets 10.00 ml of a 0.100 M solution of reagent A into an Erlenmeyer flask. Reagent A is then titrated with a titrating solution containing reagent B. The following reaction occurs

A + 3B → Products

It takes the addition of 33.24 mL of the titrating solution (i.e., the reagent B solution) to reach the endpoint. What is the molar concentration of reagent B in the titrating solution?

3. Assume that a student places 1.0 mol of solid PbI2 in 1.0 L of water. Part of the lead iodide dissolves according to the following reaction

PbI2(s) ↔ Pb2+(aq) + 2I-

(aq)

After allowing the solution to sit for 30 minutes the student determines that the iodide concentration was [I-] = 0.0032 M. What is the Pb2+ concentration?

4. Assume that a student places 1.0 mol of solid PbBr2 in 1.0 L of 0.0010 M Pb2+. Part of the lead bromide dissolves according to the following reaction

PbBr2(s) ↔ Pb2+(aq) + 2Br-

(aq)

After allowing the solution to sit for 30 minutes the student determines that the bromide concentration was [Br-] = 0.0012 M. What is the Pb2+ concentration?

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Experiment #2 Grading

Names: _______________ _______________ _______________ _______________

Group Number:__________ Date:__________ Day: _________ Time ________

PtsPoss. Earned

Pre-lab Questions

Completed correctly 10 ____

Observations, Calculations, and Conclusions

Calculation for required precipitate reactants 8 ____

Observations of Precipitation 4 ____

Calculation for required calcium for stock soln. 4 ____

Description/Calculations of calcium solutions preparation 6 ____

Proper Titration Standardization data/calcs 10 ____

Proper determination of iodate concentrations 10 ____

Proper determination of calcium concentrations 8 ____

Proper analysis of concentration products 10 ____

Proper final analysis of the results 10 ____

Post-lab Questions

Completed Correctly 10 ____

Total 90 ____

Student Peer Review Impact Factor Lab Report Score

Quiz Score(10 pts)

Total

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Names: 1)_______________ 2)________________ 3)_______________ 4)________________

Group #__________ Date_______________

Postlab Questions - Answer the Following Questions On A Separate Piece(s) of Paper

1. Based on the results from this lab experiment, predict the proper equilibrium constant

expression for the following reaction…

Fe(OH)3(s) → Fe3+(aq) + 3 OH-

(aq)

Explain the reasoning behind your answer.

2. Demonstrate through mole calculations that iodate was indeed the limiting reactant in the

production of triodide in Part I – Goal 3 of this study. You have three reactants to consider KI,

iodate, and H+. Assume that 1 cm3 of KI is approximately 1 g.

3. Did increasing the starting concentration of Ca2+ in part II of this experiment increase or

decrease the amount of Ca(IO3)2 that dissolved. Explain your answer using the concepts of

LeChatelier’s Principle.

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Experiment 2

Documents

equilibrium

laboratory

equilibrium

iodate concentration

serial dilution

added moles

equilibrium

gravity filtration