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Spectrophotometric Determination of an Equilibrium Constant | CH 223, Exams of Chemistry

Material Type: Exam; Class: General Chemistry; Subject: Chemistry; University: Portland Community College; Term: Spring 2004;

Typology: Exams

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Chemistry 223 Spring 2004
Spectrophotometric Determination of an
Equilibrium Constant
Introduction:
In this experiment, you will study the reaction between aqueous iron (III) nitrate,
Fe(NO3)3, and potassium thiocyanate, KSCN. They react to produce the blood-red
complex [Fe(SCN)2+] according to the following net ionic equation:
Fe3+ + SCN- [FeSCN2+]
The equilibrium constant expression may be expressed as:
K =
]][SCN[Fe
][FeSCN
3
2
You will prepare a series of standard solutions that contain known concentrations of
[FeSCN2+]. Because the red solutions absorb blue light very well, the blue LED setting
on the colorimeter is used and you will determine their absorbances at 470
nanometers (blue). The concentrations and absorbance values will be used to
construct a Beer’s Law calibration curve for [FeSCN2+].
Part I: The Beer’s Law Calibration Curve
One question you might have is “How can I know the equilibrium concentration of
FeSCN2+ before I make a Beer’s Law calibration curve?” The “trick” lies in utilizing
Le Chatelier’s principle as follows: 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 the [SCN–1]i in the equilibrium
mixtures. According to Le Chatelier’s principle, this high concentration forces the
reaction far to the right, using up essentially all of the SCN
ions . According to the
balanced equation, for every one mole of SCN reacted, one mole of FeSCN2+ is
produced. Thus equilibrium moles FeSCN2+eq are assumed to be equal to
initial moles SCNi for the Beer’s Law standard solutions only.
To reiterate: equilibrium moles FeSCN2+eq = initial moles SCN–1i
For Beer’s Law standard solutions only
The equilibrium concentration of FeSCN2+ is simply this number of moles divided by
the total volume after mixing.
Part II: Determining Keq
In the second part of the experiment, various combinations of Fe(NO3)3 and KSCN will
be combined. The amount of product formed at equilibrium, [FeSCN 2+
], will
be determined from the calibration graph prepared earlier. 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 balanced chemical
equation). The [Fe3+]eq can be determined by:
[Fe3+]eq = [Fe3+]i – [FeSCN2+]eq
In words: The equilibrium concentration of Fe3+ is the initial concentration of Fe3+
(after dilution) minus the amount that reacted to form FeSCN2+ (determined by
spectroscopy).
The [FeSCN2+]eq is determined by using spectroscopy and the Beer’s Law calibration
curve.
Spectrophotometric Determination of an Equilibrium Constant 1
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Chemistry 223 Spring 2004

Spectrophotometric Determination of an

Equilibrium Constant

Introduction:

In this experiment, you will study the reaction between aqueous iron (III) nitrate,

Fe(NO 3

3 , and potassium thiocyanate, KSCN. They react to produce the blood-red

complex [Fe(SCN)

2+

] according to the following net ionic equation:

Fe

3+

  • SCN

[FeSCN

2+

]

The equilibrium constant expression may be expressed as:

K =

[Fe ][SCN ]

[FeSCN ]

3

2

 

You will prepare a series of standard solutions that contain known concentrations of

[FeSCN

2+

]. Because the red solutions absorb blue light very well, the blue LED setting

on the colorimeter is used and you will determine their absorbances at 470

nanometers (blue). The concentrations and absorbance values will be used to

construct a Beer’s Law calibration curve for [FeSCN

2+

].

Part I: The Beer’s Law Calibration Curve

One question you might have is “How can I know the equilibrium concentration of

FeSCN

2+

before I make a Beer’s Law calibration curve?” The “trick” lies in utilizing

Le Chatelier’s principle as follows: A very large concentration of Fe

3+

will be added to

a small initial concentration of SCN

(hereafter referred to as [SCN

] i

). The [Fe

3+

] in

the standard solution is 100 times larger than the [SCN

] i

in the equilibrium

mixtures. According to Le Chatelier’s principle , this high concentration forces the

reaction far to the right, using up essentially all of the SCN

ions. According to the

balanced equation, for every one mole of SCN

reacted, one mole of FeSCN

2+

is

produced. Thus equilibrium moles FeSCN

2+

eq

are assumed to be equal to

initial moles SCN

-

i

for the Beer’s Law standard solutions only.

To reiterate: equilibrium moles FeSCN

2+

eq =

initial moles SCN

-

i

For Beer’s Law standard solutions only

The equilibrium concentration of FeSCN

2+

is simply this number of moles divided by

the total volume after mixing.

Part II: Determining K eq

In the second part of the experiment, various combinations of Fe(NO 3

3 and KSCN will

be combined. The amount of product formed at equilibrium, [FeSCN

2+

], will

be determined from the calibration graph prepared earlier. Knowing the

[FeSCN

2+

] eq

allows you to determine the concentrations of the other two ions at

equilibrium. For each mole of FeSCN

2+

ions produced, one less mole of Fe

ions will

be found in the solution (see the 1:1 ratio of coefficients in the balanced chemical

equation). The [Fe

3+

] eq

can be determined by:

[Fe

3+

] eq

= [Fe

3+

] i

  • [FeSCN

2+

] eq

In words: The equilibrium concentration of Fe

3+

is the initial concentration of Fe

3+

(after dilution) minus the amount that reacted to form FeSCN

2+

(determined by

spectroscopy).

The [FeSCN

2+ ] eq

is determined by using spectroscopy and the Beer’s Law calibration

curve.

Because one mole of SCN

is used up for each mole of FeSCN

2+

ions produced,

[SCN

]

eq

can be determined by:

[SCN

] eq

= [SCN

] i

  • [FeSCN

2+

] eq

Knowing the values of [Fe

3+ ] eq

, [SCN

  • ] eq

, and [FeSCN

2+ ] eq

, you can then calculate

the value of K c

, the equilibrium constant. When these concentrations are substituted

into the equation for the equilibrium constant, numerical values for the equilibrium

constant are determined. An average value for the constant will then be determined.

Purpose :

The purpose of this experiment is to determine a value for the equilibrium constant

for the reaction between iron (III) nitrate and potassium thiocyanate.

MATERIALS

computer 0.0020 M KSCN

Vernier computer interface 0.0020 M Fe(NO 3

3 (in 1.0 M HNO 3

Logger Pro 0.200 M Fe(NO 3

3 (in 1.0 M HNO 3

Vernier Colorimeter pipets and bulbs

1 plastic cuvette buret

five 20  150 mm test tubes 25-mL volumetric flasks

thermometer tissues (preferably lint-free)

Safety and Disposal Guidelines :

 Goggles (and apron) should always be worn in the lab.

 The Fe(NO 3

3 solutions are dissolved in 1.0 M HNO 3 , a strong acid. 0.050 M

HNO

3 is used as well. In case of eye contact, flush with water for 15 minutes

in an eye-wash. In case of skin contact, flush with water for 15 minutes.

 Dispose of all solutions in the waste bottle in the hood. Do not flush solutions

down the drain.

Procedure :

Special Note on Data Collection : At this point in the year, you will be expected to

prepare well-organized tables to contain your data. Pay attention to detail and what

you have learned in the past about column headings (quantity and unit) as well as

proper attention to significant figures. In addition, a good chemist will always write

down qualitative observations with every step. If you have been lacking this in the

past, now is the time to start. Think ahead: If you do not wish to re-write your data

section in the final report, be neat and organized now, it will save you time later.

Also, since you are expected to keep good lab notebooks by this point in the year, a

poor notebook may result in a reduced score for the lab report.

PART I: BEER’S LAW CALIBRATION CURVE

1. Clean, dry, and label 4 small beakers to contain the necessary reagent solutions. 2. Obtain the necessary reagents and dispense about 30 mL of each into a small

beaker. DO NOT pipet directly from a reagent stock bottle.

3. NOTE : You may prepare the following standard solutions to be used by 2 pairs of

students at the same table. Divide up the work of preparing the solutions among

the 4 people at a table. However, each pair of students wil be performing the

experiment below (you are just sharing solutions!!)

a. Open the Colorimeter lid.

b. Holding the cuvette by the upper edges, place it in the cuvette slot of the

Colorimeter. Close the lid.

c. The newer colorimeters only have a CAL button. Press the < or > button on the

Colorimeter to select a wavelength of 470 nm (Blue) for this experiment. Press

and hold the CAL button until the red LED begins to flash. Then release the

CAL button. When the LED stops flashing, the calibration is complete. Proceed

directly to Step 7.

If your Colorimeter does not have a CAL button, continue with this step to

calibrate your Colorimeter.

First Calibration Point

d. Choose Calibrate  CH1: Colorimeter (%T) from the “Experiment” menu and

then click.

e. Turn the wavelength knob on the Colorimeter to the “0% T” position.

f. Type “0” in the edit box.

g. When the displayed voltage reading for Reading 1 stabilizes, click.

Second Calibration Point

h. Turn the knob of the Colorimeter to the Blue LED position (470 nm).

i. Type “100” in the edit box.

j. When the displayed voltage reading for Reading 2 stabilizes, click , then

click.

8. You are now ready to collect absorbance data for the Beer’s Law standard

solutions.

a. Click to begin data collection.

b. Empty the water from the cuvette. Rinse it twice with ~1 mL portions of the

Test Tube 1 solution.

c. Wipe the outside of the cuvette with a tissue and then place the cuvette in the

Colorimeter. After closing the lid, wait for the absorbance value displayed in

the meter to stabilize. Then click , type the FeSCN

2+

concentration in

edit box, and press the ENTER key.

d. Discard the cuvette contents as directed by your teacher. Rinse the cuvette

twice with the second solution and fill the cuvette 3/4 full. Follow the Step-c

procedure to find the absorbance of this solution. Type the FeSCN

2+

concentration in the edit box and press ENTER.

e. Repeat the Step-d procedure to find the absorbance of the remaining standard

solutions.

f. Record the concentration and absorbance values in your notebook.

g. Determine the best fit line by clicking the Linear Fit button,.

h. Record the equation of the line and print a copy of the graph.

i. Dispose of all solutions as directed by your instructor.

PART II: DETERMINATION OF EQUILIBRIUM CONSTANT.

9. Clean and dry 5 large test tubes and label them 1-5. Use a buret or pipet to

measure the target volumes of the reactants listed below. Note that this set of

Note: If you

have an

older “box”

type

colorimeter,

you will

need to do

steps d. – j.

combinations uses the more dilute Fe(NO 3

3 solution. As usual, record

actual volumes in your data table.

Trial

Volume

0.0020 M KSCN (mL)

Volume

0.0020 M Fe(NO 3

3

(mL)

Volume

0.050 M HNO

3 (mL)

10. Re-calibrate the Colorimeter, but this time use 0.0020 M Fe(NO 3

3

11. Measure the absorbance of each equilibrium solution following the same steps as

in Part I. NOTE: You are not making a graph this time. Simply record the

absorbance values in your notebook for further analysis.

Analysis

  1. Using you’re equation of the Beer’s

Law calibration curve, determine the concentration of [Fe(SCN)

2+

] eq for each of the

trials.

  1. From the concentration of

[Fe(SCN)

2+

] eq produced and the original concentrations of the reactants, construct

tables to determine the equilibrium concentrations of all species. NOTE: The

initial concentrations should be calculated based on the fact that they are diluted

(what they would be right after mixing). Since each solution is mixed with the

other solution and HNO 3 (aq), they are no longer at 0.0020 M.

  1. Use these values to calculate the

equilibrium constant (K c ) for each trial.

  1. Calculate your standard deviation

(STD) for the K c values, and your relative (%) STD. Comment on your level of

precision. What sources of error might there have been?

Lab Report

Prepare a lab report in the standard scientific manner. Your report must be typed

and grammatically correct. You may hand-write equations instead of typing them if

you prefer. Consider the following points when creating your report, they

are the basis for which your grade will be assigned:

 Abstract: a 1 paragraph overall summary of the experiment. NOT an

explanation of the procedure! This paragraph would be similar to how you

would explain the overall purpose and theory of the lab to another person.

 Introduction: describe and explain Beer’s Law, show equations, discuss

equations used. Describe a calibration curve, discuss what a calibration curve

is used for and why they are useful, describe and discuss chemical