Transcript
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
Part II: Goal #3
Soln A1 Soln A2 Soln B1 Soln B2
initial buret volume
final buret volume
titrant volume delivered
Moles S2O32- Added
Moles IO3- Reacted
[IO3-] in Mixture
Soln C1 Soln C2 Soln D1 Soln D2
initial buret volume
final buret volume
titrant volume delivered
Moles S2O32- Added
Moles IO3- Reacted
[IO3-] in Mixture
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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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: Determination of an Equilibrium Constant Expression
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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Experiment #2: Determination of an Equilibrium Constant Expression
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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