Community College of Philadelphia Module 8 Test Tube Number Chemistry Lab Report

Module 8 – Chemical Equilibrium:Finding a Constant, Kc
MODULE 8 : FINDING AN EQUILIBRIUM CONSTANT
POST-LAB ASSIGNMENT
1. Write the Kc expression for the reaction in the Data and Calculation table below.
2. Calculate the initial concentration of Fe3+, based on the dilution that results from adding KSCN
solution and water to the original 0.0020 M Fe(NO3)3 solution. See Step 2 of the procedure for
the volume of each substance used in Trials 1–4. Calculate [Fe3+]i using the equation:
[Fe3+]i =
Fe(NO3)3 mL
total mL
 (0.0020
M)
This should be the same for all four test tubes.
3. Calculate the initial concentration of SCN–, based on its dilution by Fe(NO3)3 and water:
KSCN mL
[SCN–]i = total mL
 (0.0020
M)
In Test Tube 1, [SCN–]i = (2 mL / 10 mL)(0.0020 M) = 0.00040 M. Calculate this for the other
three test tubes.
4. [FeSCN2+]eq is calculated using the formula:
Aeq
[FeSCN2+]eq = A
std

[FeSCN2+]std
where Aeq and Astd are the absorbance values for the equilibrium and standard test tubes,
respectively, and [FeSCN2+]std = (1/10)(0.0020) = 0.00020 M. Calculate [FeSCN2+]eq for each
of the four trials.
5. [Fe3+]eq: Calculate the concentration of Fe3+ at equilibrium for Trials 1–4 using the equation:
[Fe3+]eq = [Fe3+]i – [FeSCN2+]eq
6. [SCN–]eq: Calculate the concentration of SCN- at equilibrium for Trials 1–4 using the equation:
[SCN–]eq = [SCN–]i – [FeSCN2+]eq
7. Calculate Kc for Trials 1–4. Be sure to show the Kc expression and the values substituted in for
each of these calculations.
Chemistry with Vernier
20 – 1
LabQuest 20
8. Using your four calculated Kc values, determine an average value for Kc. How constant were
your Kc values?
9. Write a lab report for this week’s laboratory.
DATA AND CALCULATIONS
Trial 1
Trial 2
Trial 3
Trial 4
_______
_______
_______
_______
Absorbance
Absorbance of standard (Trial 5)
Temperature
_______
Kc expression
_______ °C
Kc =
[Fe3+]i
[SCN–]i
[FeSCN2+]eq
[Fe3+]eq
[SCN–]eq
Kc value
Average of Kc values
Kc = ________ at ________°C
20 – 2
Chemistry with Vernier
Module 8 – Chemical Equilibrium:
Finding a Constant, Kc
PROCEDURE
1. Obtain and wear goggles.
2. Label four 20  150 mm test tubes 1– 4. Pour about 30 mL of 0.0020 M Fe(NO3)3 into a clean,
dry 100 mL beaker. Pipet 5.0 mL of this solution into each of the four labeled test tubes. Use a
pipet pump or bulb to pipet all solutions. DANGER: Fe(NO3)3 solutions in this experiment are
prepared in 1.0 M nitric acid solution, HNO3. HNO3 may intensify fire. Keep away from heat,
sparks, open flames, and hot surfaces. Causes severe skin burns and eye damage. Do not
breathe mist, vapors, or spray. Avoid contact with acetic acid and readily oxidized substances.
Pour about 25 mL of the 0.0020 M KSCN into another clean, dry 100 mL beaker. WARNING:
Potassium thiocyanate solution, KSCN: Causes eye irritation and mild skin irritation.
Pipet 2, 3, 4 and 5 mL of the KSCN solution into Test Tubes 1–4, respectively. Obtain about
25 mL of distilled water in a 100 mL beaker. Then pipet 3, 2, 1 and 0 mL of distilled water into
Test Tubes 1– 4, respectively, to bring the total volume of each test tube to 10 mL. Mix each
solution thoroughly with a stirring rod. Be sure to clean and dry the stirring rod after each
mixing. Measure and record the temperature of one of the above solutions to use as the
temperature for the equilibrium constant, Kc. Volumes added to each test tube are summarized
below:
Test Tube
Number
Fe(NO3)3
(mL)
KSCN
(mL)
H2O
(mL)
1
2
3
4
5
5
5
5
2
3
4
5
3
2
1
0
Abs
3. The lab instructor should have prepare a standard solution of FeSCN2+ by pipetting 36 mL of
0.200 M Fe(NO3)3 and 4 mL of 0.0020 M KSCN into a flask labeled “standard”. Take 2.0 ml
of this solutions into a new test tube label “5 standard”.
4. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use cuvettes:
•
Wipe the outside of each cuvette with a lint-free tissue.
• Handle cuvettes only by the top edge of the ribbed sides.
• Dislodge any bubbles by gently tapping the cuvette on a hard surface.
• Always position the cuvette so the light passes through the clear sides.
5. Connect the Colorimeter to LabQuest and choose New from the File menu.
6. Calibrate the Colorimeter.
a. Place the blank in the cuvette slot of the Colorimeter and close the lid.
b. Press the < or > button on the Colorimeter to select a wavelength of 470 nm. Press the CAL
button on the Colorimeter. When the LED stops flashing, the calibration is complete.
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20 – 1
LabQuest 20
7. Set up the data collection mode.
a. On the Meter screen, tap Mode. Change the mode to Events with Entry.
b. Enter the Name (Concentration) and Units (mol/L). Select OK.
c. Proceed directly to Step 8.
8. You are now ready to collect absorbance data for the four equilibrium systems and the standard
solution.
a. Leave the cuvette containing the Test Tube 1 mixture in the device. Start data collection.
b. When the value displayed on the screen has stabilized, record it in your note book as the
absorbance of test tub 1.
c. Discard the cuvette contents back into a personal waste beaker. Rinse the cuvette twice with
the water and once with Test Tube 2 solution, fill the cuvette 3/4 full, and place it in the
device. After the reading has stabilizes, record it in your note book as the absorbance of test
tub 2.
d. Repeat the procedure to find the absorbance of the solutions in Test Tubes 3, 4, and 5 (the
standard solution).
e. Stop data collection. To examine the data pairs on the displayed graph, tap any data point.
As you tap a data point, the absorbance and concentration values are displayed to the right of
the graph. Record the absorbance values, for each Trial, in your data table.
f. Dispose of all solutions as directed by your instructor.
THE DATA WAS COLLECTED FOR YOU:
Table 1. This table shows the absorbance of the reactions at different concentrations at a
constant temperature.
Test Tube Number
Absorbance
Temperature (℃)
1(Trial 1)
0.153
25
2 (Trial 2)
0.217
25
3 (Trial 3)
0.286
25
4 (Trial 4)
0.356
25
5 (Standard)
0.559
25
20 – 2
Chemistry with Vernier
Module 8 – Chemical Equilibrium:
Finding a Constant, Kc
MODULE 8 : FINDING AN EQUILIBRIUM CONSTANT
POST-LAB ASSIGNMENT
1. Write the Kc expression for the reaction in the Data and Calculation table below.
2. Calculate the initial concentration of Fe3+, based on the dilution that results from adding KSCN
solution and water to the original 0.0020 M Fe(NO3)3 solution. See Step 2 of the procedure for
the volume of each substance used in Trials 1–4. Calculate [Fe3+]i using the equation:
[Fe3+]i =
Fe(NO3)3 mL
total mL
 (0.0020
M)
This should be the same for all four test tubes.
3. Calculate the initial concentration of SCN–, based on its dilution by Fe(NO3)3 and water:
KSCN mL
[SCN–]i = total mL
 (0.0020
M)
In Test Tube 1, [SCN–]i = (2 mL / 10 mL)(0.0020 M) = 0.00040 M. Calculate this for the other
three test tubes.
4. [FeSCN2+]eq is calculated using the formula:
Aeq
[FeSCN2+]eq = A
std

[FeSCN2+]std
where Aeq and Astd are the absorbance values for the equilibrium and standard test tubes,
respectively, and [FeSCN2+]std = (1/10)(0.0020) = 0.00020 M. Calculate [FeSCN2+]eq for each
of the four trials.
5. [Fe3+]eq: Calculate the concentration of Fe3+ at equilibrium for Trials 1–4 using the equation:
[Fe3+]eq = [Fe3+]i – [FeSCN2+]eq
6. [SCN–]eq: Calculate the concentration of SCN- at equilibrium for Trials 1–4 using the equation:
[SCN–]eq = [SCN–]i – [FeSCN2+]eq
7. Calculate Kc for Trials 1–4. Be sure to show the Kc expression and the values substituted in for
each of these calculations.
Chemistry with Vernier
20 – 1
LabQuest 20
8. Using your four calculated Kc values, determine an average value for Kc. How constant were
your Kc values?
9. Write a lab report for this week’s laboratory.
DATA AND CALCULATIONS
Trial 1
Trial 2
Trial 3
Trial 4
_______
_______
_______
_______
Absorbance
Absorbance of standard (Trial 5)
Temperature
_______
Kc expression
_______ °C
Kc =
[Fe3+]i
[SCN–]i
[FeSCN2+]eq
[Fe3+]eq
[SCN–]eq
Kc value
Average of Kc values
Kc = ________ at ________°C
20 – 2
Chemistry with Vernier
Module 8 – Chemical Equilibrium:
Finding a Constant, Kc
Introduction (Source: https://sites.google.com/a/wrps.net/lhs-ap-chemistry/home/labresources/lab-18
There are many reactions that take place in solution that are equilibrium reactions; that is, they do
not go to completion. In these reactions both the forward and reverse reaction are occurring, and
both reactants and products are always present. Examples of this type of reaction include weak
acids such as acetic acid dissociating in water, weak bases such as ammonia reacting with water,
and the formation of “complex ions” in which a metal ion combines with one or more negative
ions.
In this module, a reaction involving the formation of a complex ion which occurs when solutions of
iron (III) [Fe3+] are combined with solutions of the negative thiocyanate (SCN-) will be studied and
the equilibrium constant determined.
Consider the following general equation for a reversible chemical reaction:
aA + bB ⇄ cC + dD
The equilibrium constant Keq for this general reaction is given by below. The square brackets
refer to the concentration (in molarity, M) of the reactants and products at equilibrium.
The equilibrium constant gets its name from the fact that for any reversible chemical reaction, the
value of Keq is a constant at a particular temperature. The concentrations of reactants and products
at equilibrium vary, depending on the initial amounts of materials present. The special ratio of
reactants and products described by Keq is always the same as long as the system has reached
equilibrium and the temperature does not change. The value of Keq can be calculated if the
concentrations of reactants and products at equilibrium are known.
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LabQuest 20
The reversible chemical reaction of iron (III) ions (Fe 3+) with thiocyanate (SCN-) provides a convenient example for
determining the equilibrium constant of a reaction. As shown in the equation below, Fe3+ and SCN- ions combine to
form a special type of combined or “complex” ion having the formula FeSCN 2+.
The equilibrium constant expression for this reaction is shown below:
Iron thiocyanate – FeSCN information
The value of Keq can be determined experimentally by mixing known concentrations of Fe3+ and SCN- ions and
measuring the concentration of FeSCN2+ ions at equilibrium. As noted in the reversible reaction above,
the Fe3+ ions are pale yellow and the SCN- ions are colorless. The FeSCN2+ ions are blood-red. The concentration
of FeSCN2+ complex ions formed is proportional to the intensity of the red color.
Compounds that are colored absorbed a part of the visible spectrum of light. If a compound absorbs
green light, it will appear red in color, the complimentary color to green. Shown below is an
absorbance spectrum for FeSCN2+ (aq).We will use a spectrophotometer to measures the amount of
light of a given wavelength (in this lab, 470 nm as it corresponds to the wavelength of maximum
absorbance as show below) or color that is absorbed by a solution.
20 – 2
Chemistry with Vernier
Chemical Equilibrium: Finding a Constant, Kc
What You’ll Be Doing:
The purpose of this lab is to experimentally determine the equilibrium constant, Kc, for the
following chemical reaction:
Fe3+(aq) + SCN–(aq) 
⎯→ FeSCN2+(aq)
iron(III) thiocyanate
thiocyanoiron(III)
When Fe3+ and SCN- are combined, equilibrium is established between these two ions and the
FeSCN2+ ion. In order to calculate Kc for the reaction, it is necessary to know the concentrations of
all ions at equilibrium: [FeSCN2+]eq, [SCN–]eq, and [Fe3+]eq. You will prepare four equilibrium
systems containing different concentrations of these three ions. The equilibrium concentrations of
the three ions will then be experimentally determined. These values will be substituted into the
equilibrium constant expression to see if Kc is indeed constant.
You will use a Colorimeter or a Spectrometer to determine [FeSCN2+]eq. The FeSCN2+ ion
produces solutions with a red color. Because the red solutions absorb blue light very well, so
Colorimeter users will be instructed to use the 470 nm (blue) LED. Spectrometer users will
determine an appropriate wavelength based on the absorbance spectrum of the solution. The light
striking the detector is reported as absorbance or percent transmittance. By comparing the
absorbance of each equilibrium system, Aeq, to the absorbance of a standard solution, Astd, you can
determine [FeSCN2+]eq. The standard solution has a known FeSCN2+ concentration.
To prepare the standard solution, a very large concentration of Fe3+ will be added to a small initial
concentration of SCN– (hereafter referred to as [SCN–]i. The [Fe3+] in the standard solution is
100 times larger than [Fe3+] in the equilibrium mixtures. According to LeChatelier’s principle, this
high concentration forces the reaction far to the right, using up nearly 100% of the SCN– ions.
According to the balanced equation, for every one mole of SCN– reacted, one mole of FeSCN2+ is
produced. Thus [FeSCN2+]std is assumed to be equal to [SCN–]i.
Assuming [FeSCN2+] and absorbance are related directly (Beer’s law), the concentration of
FeSCN2+ for any of the equilibrium systems can be found by:
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Aeq
[FeSCN2+]eq = A
std
2
X [FeSCN +]std
Knowing the [FeSCN2+]eq allows you to determine the concentrations of the other two ions at
equilibrium. For each mole of FeSCN2+ ions produced, one less mole of Fe3+ ions will be found in
the solution (see the 1:1 ratio of coefficients in the equation on the previous page). The [Fe3+] can
be determined by:
[Fe3+]eq = [Fe3+]i – [FeSCN2+]eq
Because one mole of SCN- is used up for each mole of FeSCN2+ ions produced, [SCN–]eq can be
determined by:
[SCN–]eq = [SCN–]i – [FeSCN2+]eq
Knowing the values of [Fe3+]eq, [SCN–]eq, and [FeSCN2+]eq, you can now calculate the value of
Kc, the equilibrium constant.
OBJECTIVE
In this experiment, you will determine the equilibrium constant, Kc, for the following chemical
reaction:
Fe3+(aq) + SCN–(aq) 
⎯→ FeSCN2+(aq)
iron(III) thiocyanate
thiocyanoiron(III)
MATERIALS
Vernier LabQuest
LabQuest App
Vernier Colorimeter or Spectrometer
1 plastic cuvette
five 20  150 mm test tubes
three 100 mL beakers
thermometer or Temperature Probe
20 – 4
0.0020 M KSCN
0.0020 M Fe(NO3)3 (in 1.0 M HNO3)
0.200 M Fe(NO3)3 (in 1.0 M HNO3)
four pipets
pipet bulb or pipet pump
tissues (preferably lint-free)
Chemistry with Vernier
6:54
Il LTE C.
< Module 8 Procedure and Collect... Module 8 - Chemical Equilibrium: Finding a Constant, Kc PROCEDURE 1. Obtain and wear goggles. 2. Label four 20 Í 150 mm test tubes 1-4. Pour about 30 mL of 0.0020 M Fe(NO3), into a clean, dry 100 mL beaker. Pipet 5.0 mL of this solution into each of the four labeled test tubes. Use a pipet pump or bulb to pipet all solutions. DANGER: Fe(NO3)3 solutions in this experiment are prepared in 1.0 M nitric acid solution, HNO3. HNO2 may intensify fire. Keep away from heat, sparks, open flames, and hot surfaces. Causes severe skin burns and eye damage. Do not breathe mist, vapors, or spray. Avoid contact with acetic acid and readily oxidized substances. Pour about 25 mL of the 0.0020 M KSCN into another clean, dry 100 mL beaker. WARNING: Potassium thiocyanate solution, KSCN: Causes eye irritation and mild skin irritation. Pipet 2, 3, 4 and 5 mL of the KSCN solution into Test Tubes 1-4, respectively. Obtain about 25 mL of distilled water in a 100 mL beaker. Then pipet 3, 2, 1 and 0 mL of distilled water into Test Tubes 1-4, respectively, to bring the total volume of each test tube to 10 mL. Mix each solution thoroughly with a stirring rod. Be sure to clean and dry the stirring rod after each mixing. Measure and record the temperature of one of the above solutions to use as the temperature for the equilibrium constant, Kc. Volumes added to each test tube are summarized below: Test Tube Fe(NO3)3 KSCN HOAbs Number (mL) (mL) |(mL) 1 5 2 3 2 5 3 2 3 5 4 1 4 5 5 0 3. The lab instructor should have prepare a standard solution of FeSCN²+ by pipetting 36 mL of 0.200 M Fe(NO3)3 and 4 mL of 0.0020 M KSCN into a flask labeled “standard”. Take 2.0 ml of this solutions into a new test tube label “5 standard”. 4. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use cuvettes: • Wipe the outside of each cuvette with a lint-free tissue. · Handle cuvettes only by the top edge of the ribbed sides. · Dislodge any bubbles by gently tapping the cuvette on a hard Dashboard Calendar To Do Notifications Inbox 6:54 4 Il LTE C. < Module 8 Procedure and Collect... F 4. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use cuvettes: · Wipe the outside of each cuvette with a lint-free tissue. • Handle cuvettes only by the top edge of the ribbed sides. · Dislodge any bubbles by gently tapping the cuvette on a hard surface. · Always position the cuvette so the light passes through the clear sides. 5. Connect the Colorimeter to LabQuest and choose New from the File menu. 6. Calibrate the Colorimeter. a. Place the blank in the cuvette slot of the Colorimeter and close the lid. b. Press the < or > button on the Colorimeter to select a
wavelength of 470 nm. Press the CAL button on the
Colorimeter. When the LED stops flashing, the calibration is
complete.
7. Set up the data collection mode.
a. On the Meter screen, tap Mode. Change the mode to Events
with Entry.
b. Enter the Name (Concentration) and Units (mol/L). Select OK.
c. Proceed directly to Step 8.
8. You are now ready to collect absorbance data for the four
equilibrium systems and the standard solution.
a. Leave the cuvette containing the Test Tube 1 mixture in the
device. Start data collection.
b. When the value displayed on the screen has stabilized, record it
in your note book as the absorbance of test tub 1.
c. Discard the cuvette contents back into a personal waste beaker.
Rinse the cuvette twice with the water and once with Test Tube
2 solution, fill the cuvette 3/4 full, and place it in the device.
After the reading has stabilizes, record it in your note book as
the absorbance of test tub 2.
d. Repeat the procedure to find the absorbance of the solutions in
Test Tubes 3, 4, and 5 (the standard solution).
Stop data collection. To examine the data pairs on the displayed
graph, tap any data point. As you tap a data point, the
absorbance and concentration values are displayed to the right
of the graph. Record the absorbance values, for each Trial, in
your data table.
f. Dispose of all solutions as directed by your instructor.
e.
=
Dashboard
Calendar
To Do
Notifications
Inbox
6:54
Il LTE C.
< Module 8 Procedure and Collect... Module 8 - Chemical Equilibrium: Finding a Constant, Kc PROCEDURE 1. Obtain and wear goggles. 2. Label four 20 Í 150 mm test tubes 1-4. Pour about 30 mL of 0.0020 M Fe(NO3), into a clean, dry 100 mL beaker. Pipet 5.0 mL of this solution into each of the four labeled test tubes. Use a pipet pump or bulb to pipet all solutions. DANGER: Fe(NO3)3 solutions in this experiment are prepared in 1.0 M nitric acid solution, HNO3. HNO2 may intensify fire. Keep away from heat, sparks, open flames, and hot surfaces. Causes severe skin burns and eye damage. Do not breathe mist, vapors, or spray. Avoid contact with acetic acid and readily oxidized substances. Pour about 25 mL of the 0.0020 M KSCN into another clean, dry 100 mL beaker. WARNING: Potassium thiocyanate solution, KSCN: Causes eye irritation and mild skin irritation. Pipet 2, 3, 4 and 5 mL of the KSCN solution into Test Tubes 1-4, respectively. Obtain about 25 mL of distilled water in a 100 mL beaker. Then pipet 3, 2, 1 and 0 mL of distilled water into Test Tubes 1-4, respectively, to bring the total volume of each test tube to 10 mL. Mix each solution thoroughly with a stirring rod. Be sure to clean and dry the stirring rod after each mixing. Measure and record the temperature of one of the above solutions to use as the temperature for the equilibrium constant, Kc. Volumes added to each test tube are summarized below: Test Tube Fe(NO3)3 KSCN HOAbs Number (mL) (mL) |(mL) 1 5 2 3 2 5 3 2 3 5 4 1 4 5 5 0 3. The lab instructor should have prepare a standard solution of FeSCN²+ by pipetting 36 mL of 0.200 M Fe(NO3)3 and 4 mL of 0.0020 M KSCN into a flask labeled “standard”. Take 2.0 ml of this solutions into a new test tube label “5 standard”. 4. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use cuvettes: • Wipe the outside of each cuvette with a lint-free tissue. · Handle cuvettes only by the top edge of the ribbed sides. · Dislodge any bubbles by gently tapping the cuvette on a hard Dashboard Calendar To Do Notifications Inbox