University of Nairobi Chemical Equilibrium and Kc Determination Lab Report Sheet

Chemical Equilibrium and Kc DeterminationChemical equilibrium plays an important role in chemical and biological systems. Consider the generic reaction
at equilibrium below:
aA + bB  cC + dD
Here A and B are the reactants, C and D are the products, and a, b, c, and d are the coefficients in the chemical
equation. The equilibrium constant (Kc) for this reaction is given by:
𝑲𝒄 =
[𝑪]𝒄 [𝑫]𝒅
[𝑨]𝒂 [𝑩]𝒃
Kc is a constant for a given reaction at a specified temperature and pressure, regardless of the starting
concentrations of each reactant and product. If a reaction is not at equilibrium, then the ratio instead gives the
more general reaction quotient Q:
[𝑪]𝒄 [𝑫]𝒅
𝑸=
[𝑨]𝒂 [𝑩]𝒃
Notice that Q takes on the same general form as Kc (ratio of products over reactants). However, this ratio is only
equal to the equilibrium constant Kc when the reaction is at equilibrium, while it is equal to Q at any other point
during the reaction. This is an important difference!
In this experiment, the spectophotometric skills developed at the beginning of this semester will be utilized to
determine the equilibrium constant for a reaction. The equilibrium system that will be studied is:
Fe3+(aq)
+
HSCN(aq)
⇄
Fe(SCN)2+(aq)
(blood-red colored)
+
H+(aq)
The equilibrium constant (Kc) for this reaction will be determined after mixing together varying amounts of iron
(III) nitrate (a source of Fe3+ ions) and HSCN. The thiocyanatoiron(III) species, FeSCN2+, is a complex ion that
has a vivid dark red color, while all other species at equilibrium are mostly colorless. Since the FeSCN 2+ is the
only highly colored species, its final (equilibrium) concentration can be determined by using the UV-visible
spectrophotometer to read its absorbance at 450 nm and comparing the result to a calibration curve. The
equilibrium concentrations of Fe3+ and HSCN can then be determined using an ICE table. (Note: In this procedure
the concentration of H+ will be held roughly constant since the entire reaction takes place in 0.50 M HNO3.) The
equilibrium concentrations can be used to determine the equilibrium constant (Kc) for the reaction.
Objectives
1. Determine equilibrium concentrations of FeSCN2+ from spectrophotometric data
2. Complete an ICE table for each experimental run
3. Calculate Kc for the above reaction
STEM Skills Practiced in this Experiment
1. Technical skills – utilize quantitative laboratory techniques to accurately prepare solutions; utilize the
UV-visible spectrophotometer to accurately measure the absorbance of solutions
2. Data Analysis – determine the concentration of a species in solution using a Beer’s Law calibration curve
and UV-visible absorbance values
3. Problem Solving – complete ICE tables for each experimental run and use these to calculate Kc
4. Critical Thinking – compare Kc for each run and confirm that Kc is not affected when different initial
concentrations of reactants are used
Procedure
Part I: Prepare equilibrium solutions
NOTE: Do not waste solutions. Before going to get solutions from the stock bottles, determine the total volume
needed and do not take extra. Pour what is needed into an appropriately sized beaker or flask. Do not put pipets
into stock bottles. Read labels carefully.
Table 1: Ratio of Reactants to Be Used in Test Solutions
Tube #
2.00 x 10-3 M Fe(NO3)3
2.00 x 10-3 M HSCN
0.50 M HNO3
Total Volume
Run 1
5.00 mL
1.00 mL
4.00 mL
10.00 mL
Run 2
5.00 mL
2.00 mL
3.00 mL
10.00 mL
Run 3
5.00 mL
3.00 mL
2.00 mL
10.00 mL
Run 4
5.00 mL
4.00 mL
1.00 mL
10.00 mL
Run 5
5.00 mL
5.00 mL
0.00 mL
10.00 mL
Blank
5.00 mL
0 mL
5.00 mL
10.00 mL
Prepare duplicate medium test tubes according to Table 1. Click here to view a video showcasing how these
solutions are prepared. Clearly label each test tube. Mix the solutions thoroughly. DO NOT USE DISPOSABLE
PIPPETTES AS THEY ARE NOT ACCURATE ENOUGH. Click here to view an image of the final reaction
mixtures and describe them on your lab report sheet. It can be assumed for each run that the equilibrium is
established as soon as the solutions are mixed together.
IMPORTANT NOTE: The stock solutions were prepared in acidic conditions. The actual contents of each stock
solution are:
• 2.00 x 10-3 M Fe(NO3)3 in 0.50 M HNO3
• 2.00 x 10-3 M HSCN in 0.50 M HNO3
• 0.50 M HNO3
Note that when two or more solutions are mixed together, this dilutes each solution and therefore changes
(lowers) their concentrations. Use Table 1 and the dilution equation (M1V1 = M2V2) to calculate the initial
concentrations (i.e. the concentrations after mixing) of Fe(NO3)3 and HSCN for each run. Insert these
concentrations into Table 2. Show your work for at least one run.
Table 2: Initial Concentrations of Reactants and Absorbance at 450 nm for each Run
A450 –
A450 –
Tube #
[Fe(NO3)3]0 (M)
[HSCN]0 (M)
Sample 1
Sample 2
Blank
0.000
0.000
Run 1
0.046
0.047
Run 2
0.089
0.090
Run 3
0.131
0.131
Run 4
0.170
0.169
Run 5
0.209
0.210
A450 Average
Part II: Determine the concentration of FeSCN2+ spectrophotometrically
Use your blank to zero the spectrophotometer. (Instructions for the spectrophotometers are available near the
instrument.) Click here to view a video detailing how absorbance values are measured in this lab. Measured
absorbance values at 450 nm (A450) of each solution are recorded in Table 2 below. Based on these, calculate the
average absorbance (A450 – Average) for each run and record all these values in Table 2 on the Lab Report Sheet.
Waste Disposal:
Dispose of all solutions in the Hazardous Waste container
Calculations
1. A Beer’s Law calibration curve (also called a standard curve) has been prepared for you to relate the
concentration of Fe(SCN)2+ to its absorbance at 450 nm. Use the standard curve below and the average
absorbances in Table 2 to determine the equilibrium concentration of Fe(SCN)2+ in each run. Insert these
concentrations in Table 3. Show your work for at least one run.
Figure 1: Standard Curve for [Fe(SCN)2+].
Table 3: Equilibrium Concentrations of [FeSCN2+]
Tube #
Run 1
Run 2
Run 3
Run 4
Run 5
[FeSCN2+]eq (M)
2. Write an expression for the equilibrium constant (Kc) for the reaction studied here:
Fe3+(aq)
+
HSCN(aq)
⇄
Fe(SCN)2+(aq) +
H+(aq)
Kc =
3. Complete the following ICE tables and use them to calculate the value of Kc for all five trials. The initial
concentrations can be found in Table 2, and the equilibrium concentration of Fe(SCN)2+ can be found in Table
3. Assume the initial concentration of Fe(SCN)2+ is zero.
Table 4: ICE Table for Run 1
Average
Concentration
(M)
ICE
[Fe3+] (M)
[HSCN] (M)
[FeSCN2+] (M)
[H+] (M)
I
0.50
C
–
E
0.50
Kc =
Table 5: ICE Table for Run 2
Average
Concentration
(M)
ICE
[HSCN] (M)
[FeSCN2+] (M)
[H+] (M)
I
0.50
C
–
E
Kc =
[Fe3+] (M)
0.50
Table 6: ICE Table for Run 3
Average
Concentration
(M)
ICE
[Fe3+] (M)
[HSCN] (M)
[FeSCN2+] (M)
[H+] (M)
I
0.50
C
–
E
0.50
Kc =
Table 7: ICE Table for Run 4
Average
Concentration
(M)
ICE
[Fe3+] (M)
[HSCN] (M)
[FeSCN2+] (M)
[H+] (M)
I
0.50
C
–
E
0.50
Kc =
Table 8: ICE Table for Run 5
Average
Concentration
(M)
ICE
[HSCN] (M)
[FeSCN2+] (M)
[H+] (M)
I
0.50
C
–
E
Kc =
[Fe3+] (M)
0.50
Calculate the average of the equilibrium constant using all 5 runs.
Analysis and Discussion:
1. The equilibrium constant should be a constant for a given reaction at a set temperature and pressure,
regardless of the starting concentrations of each reactant and product.
a. Compare the values of Kc for runs 1 – 5. Do the results support the above statement? Keep in mind
that random experimental error will always prevent them from being exactly the same, but it is
possible to see if the values are consistent or not.
b. If the answer to part (a) was “no”, what errors might have caused this? “Human error” is not an
acceptable error; be specific about the type(s) of error that may have impacted the results.
2. Consult with another group and calculate the percent difference between your Kc values. Explain any
discrepancies between your results.
|𝑽𝒂𝒍𝒖𝒆 𝟏 − 𝑽𝒂𝒍𝒖𝒆 𝟐|
𝑷𝒆𝒓𝒄𝒆𝒏𝒕 𝒅𝒊𝒇𝒇𝒆𝒓𝒆𝒏𝒄𝒆 =
× 𝟏𝟎𝟎 %
𝑽𝒂𝒍𝒖𝒆 𝟏 + 𝑽𝒂𝒍𝒖𝒆 𝟐
(
)
𝟐
3. Assume in a different experiment, you prepare a mixture containing 10.0 M FeSCN2+, 1.0 M H+, 0.1 M
Fe3+ and 0.1 M HSCN. Is the initial mixture at equilibrium? If not, in what direction must the reaction
proceed to reach equilibrium? (Hint: You will need to use the value of Kc you determined in the lab!)
Conclusion
Summarize today’s lab by answering the following questions in your laboratory notebook:
1. What was the main topic/concept investigated today?
2. How was this concept investigated?
3. What were the main results found in this experiment? Include both quantitative results (i.e. important
calculated values, including percent error) and qualitative results (i.e. important properties that were
investigated).
4. Summarize the main errors encountered in this experiment and how they may have affected the results.
Chemical Equilibrium and Kc Determination Lab Report Sheet
Title:
Date:
Name:
Lab Partner(s):
Purpose:
Prepare for the lab by answering the following questions. Use complete sentences, avoiding the use of “I” or
“we”. Check with your instructor to see whether you should just respond to the questions or if you should place
your responses in a single paragraph.
1. What is the main topic/concept being investigated today?
2. How will you investigate this topic/concept?
3. Write the main definitions and equations needed for today’s experiment.
Data:
Describe the appearance of all reactants.
Describe experimental observations as the solutions mix.
Use Table 1 and the dilution equation (M1V1 = M2V2) to calculate the initial concentrations of Fe(NO3)3 and
HSCN for each run. Insert these concentrations into Table 2. You must show your work for at least one trial.
Table 2: Initial Concentrations of Reactants and Absorbance at 450 nm for each run
Tube #
[Fe(NO3)3]0 (M)
[HSCN]0 (M)
A450 – Sample 1
A450 – Sample 2
A450 – Average
Run 1
Run 2
Run 3
Run 4
Run 5
Calculations
1. A Beer’s Law calibration curve (also called a standard curve) has been prepared for you to relate the
concentration of Fe(SCN)2+ to its absorbance at 450 nm. Use the standard curve below and the average
absorbances in Table 2 to determine the equilibrium concentration of Fe(SCN)2+ in each run. Insert these
concentrations in Table 3. Show your work for at least one run.
Table 3: Equilibrium Concentrations of [FeSCN2+]
Tube #
Run 1
Run 2
Run 3
Run 4
Run 5
[FeSCN2+]eq (M)
2. Write an expression for the equilibrium constant (Kc) for the reaction studied here:
Fe3+(aq)
+
HSCN(aq)
⇄
Fe(SCN)2+(aq) +
H+(aq)
Kc =
3. Complete the following ICE tables and use them to calculate the value of Kc for all five trials. The initial
concentrations can be found in Table 2, and the equilibrium concentration of Fe(SCN)2+ can be found in Table
3. Assume the initial concentration of Fe(SCN)2+ is zero.
Table 4: ICE Table for Run 1
Average
Concentration
(M)
ICE
[Fe3+] (M)
[HSCN] (M)
[FeSCN2+] (M)
[H+] (M)
I
0.50
C
–
E
0.50
Kc =
Table 5: ICE Table for Run 2
Average
Concentration
(M)
ICE
[HSCN] (M)
[FeSCN2+] (M)
[H+] (M)
I
0.50
C
–
E
Kc =
[Fe3+] (M)
0.50
Table 6: ICE Table for Run 3
Average
Concentration
(M)
ICE
[Fe3+] (M)
[HSCN] (M)
[FeSCN2+] (M)
[H+] (M)
I
0.50
C
–
E
0.50
Kc =
Table 7: ICE Table for Run 4
Average
Concentration
(M)
ICE
[Fe3+] (M)
[HSCN] (M)
[FeSCN2+] (M)
[H+] (M)
I
0.50
C
–
E
0.50
Kc =
Table 8: ICE Table for Run 5
Average
Concentration
(M)
ICE
[Fe3+] (M)
[HSCN] (M)
[FeSCN2+] (M)
[H+] (M)
I
0.50
C
–
E
Kc =
4. Calculate the average of the equilibrium constant using all 5 runs.
0.50
Analysis and Discussion:
1. The equilibrium constant should be a constant for a given reaction at a set temperature and pressure,
regardless of the starting concentrations of each reactant and product.
a. Compare the values of Kc for runs 1 – 5. Do the results support the above statement? Keep in mind
that random experimental error will always prevent them from being exactly the same, but it is
possible to see if the values are consistent or not.
b. If the answer to part (a) was “no”, what errors might have caused this? “Human error” is not an
acceptable answer; be specific about the type(s) of error that may have impacted the results.
2. Consult with another group and calculate the percent difference between your Kc values. Explain any
discrepancies between your results.
3. Assume in a different experiment, you prepare a mixture containing 10.0 M FeSCN2+, 1.0 M H+, 0.1 M
Fe3+ and 0.1 M HSCN. Is the initial mixture at equilibrium? If not, in what direction must the reaction
proceed to reach equilibrium? (Hint: you will need to use the value of Kc you determined in the lab!)
Conclusion
Summarize the lab by answering the following questions in your lab notebook. Check with your instructor to see
whether you should just respond to the questions or if you should place your responses in a single paragraph.
1. What was the main topic/concept investigated today?
2. How was this concept investigated?
3. What were the main results found in this experiment? If applicable, include both quantitative results (i.e.
important calculated values, including percent error) and qualitative results (i.e. important properties that
were investigated).
4. Summarize the main errors encountered in this experiment and how they may have affected the results.

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