_____University of Puget Sound
Department of Chemistry
Chem 230
EXP. 5–DETERMINATION OF AN EQUILIBRIUM CONSTANT
LABORATORY OBJECTIVES AND ASSESSMENTS
1. Understand the relationship between equilibrium constant and concentration.
a. Recognize the characteristics of chemical equilibrium.
b. Demonstrate that reaction mixtures with differing concentrations of reactant
and product are at equilibrium.
c. Calculate the initial concentration of reactants and products of a chemical
reaction.
d. Write the equilibrium equation for any balanced chemical equation.
e. Describe and demonstrate the use of a calibration curve and Beer’s Law to
measure the concentration of Fe(SCN)2+.
f. Calculate the equilibrium constant, Kc, from the equilibrium concentrations of
products and reactants.
g. Calculate the equilibrium concentrations of reactants and/or products from Kc
and initial concentrations of reactants and/or products using ICE tables.
2. Understand how to use Le Châtelier’s principle to manipulate equilibria.
a. State Le Châtelier’s Principle.
b. Describe and demonstrate on specified chemical reactions how Le Châtelier’s
Principle is applied to equilibrium systems to shift the equilibrium of a
chemical system.
INTRODUCTION
This experiment will give you an opportunity to determine the equilibrium
constant for a reaction and use Le Châtelier’s principle.
Formation of Fe(SCN)2+
When chemical substances react, the reaction typically does not go to completion.
Rather, the system goes to some intermediate state in which both the reactants and
products have concentrations which do not change with time. Such a system is said to
be in dynamic chemical equilibrium.
When the reaction between Fe3+ and SCN- (thiocyanate) ions in an aqueous
solution comes to equilibrium, the solution contains reactants and the product,
Fe(SCN)2+. The chemical equation for this reaction is
Fe3+(aq)
+
SCN-(aq)
Fe(SCN)2+(aq)
(1)
The product is a complex ion that has a coordinate covalent bond between the iron
cation and an atom (probably the S atom) from the thiocyanate anion. The color of this
complex ion is so intense that thiocyanate ions can be used to detect very small
quantities of Fe3+. Interestingly, Fe(SCN)2+ appears to exist solely in solution; solid
compounds containing this cation and some anion have never been isolated.
At equilibrium a reaction mixture obeys the Law of Mass Action, which imposes a
condition on the concentrations of reactants and products. This condition is expressed
in the equilibrium constant Kc. The object of this experiment is to determine the value
for the equilibrium constant for reaction (1). The equilibrium constant is given by the
expression
Kc =
c230 Exp. 5 - Determination of an Equilibrium Constant
(2)
1
where the concentrations of the substances are those at equilibrium. If these
concentrations can be determined, Kc can be calculated easily.
SPECTROPHOTOMETRY
Both reactants, Fe3+ and SCN-, are essentially colorless, whereas the complex ion
product is highly colored. Therefore, you can use a spectrophotometer to monitor the
concentration of Fe(SCN)2+ by measuring the absorbance due to the complex ion
according to Beer’s law.
A = εbc
(3)
LE CHÂTELIER’S PRINCIPLE
To determine the relationship between Fe(SCN)2+ and absorbance, the absorbance
of a series of solutions of known concentrations of complex ion must be measured.
But there is a problem. How can solutions of known concentrations of Fe(SCN)2+ be
obtained? After all, this substance is a participant in the equilibrium with Fe3+ and
SCN- ions. Stoichiometric quantities of reactants yield a specific concentration of
product which can be calculated if the equilibrium constant, Kc, is known. But Kc is
the term that is trying to be determined!
The difficulty of making solutions with known concentrations of Fe(SCN)2+ can be
overcome by using Le Châtelier’s principle. This principle suggests that an
equilibrium position can be shifted to the right, that is, in the forward direction, when
more of a given reactant is added. As more reactant is added, more product will form.
It is possible to add so much of one reactant that essentially all of the other reactant is
converted to product. In this experiment you will use overwhelming amounts of Fe3+
and limited quantities of SCN- to achieve this result. The concentration of Fe(SCN)2+
that is formed will be essentially identical to the starting concentration of the limiting
reactant, SCN-.
THE STANDARD OR CALIBRATION CURVE
In Part 1 of this experiment you will prepare a series of solutions having known,
limited initial concentrations of SCN- and an excess of Fe3+ which allows you to know
the concentrations of Fe(SCN)2+. You will plot absorbances versus concentrations of
Fe(SCN)2+ (at the λmax of Fe(SCN)2+) which should give a straight line (the calibration
curve).
DETERMINATION OF KC
As noted above, when solutions containing Fe3+ ion and thiocyanate ion are
mixed, reaction (1) occurs forming the Fe(SCN)2+ complex ion which has a deep red
color. As a result of the reaction, the equilibrium amounts of Fe3+ and SCN- are less
than they would have been if no reaction occurred; for every mole of Fe(SCN)2+ that
is formed, one mole of Fe3+ and one mole of SCN- have reacted.
The value of Kc in equation (2) is constant at a given temperature. This means that
solutions containing Fe3+ and SCN- will react until the concentrations of remaining
reactants and formed products satisfy the value of Kc. The same value of the Kc will
be obtained no matter what initial amounts of Fe3+ and SCN- are used. Our purpose
in this experiment will be to determine the value of Kc for this reaction using several
solutions made up in different ways, and to show that Kc indeed has the same value
in each circumstance. The reaction is a good one to study because Kc is a convenient
c230 Exp. 5 - Determination of an Equilibrium Constant
2
magnitude and the color of the Fe(SCN)2+ ion makes the analysis of the equilibrium
mixture amenable to spectrophotometric analysis.
In Part 2 of this experiment you will combine solutions of known initial
concentrations of iron(III) nitrate, Fe(NO3)3, and potassium (or sodium) thiocyanate,
KSCN (or NaSCN.) Then you will determine the absorbance at λ max of these new
solutions. The absorbance can be used to determine the equilibrium concentration of
Fe(SCN)2+. Knowing the initial concentrations of the reactants and the equilibrium
concentration of Fe(SCN)2+, you can calculate the equilibrium concentrations of the
remaining reactants. You can then determine the value of Kc.
CALCULATIONS
Since the calculations required in this experiment may not be apparent, we will go
through a specific example. Let us assume that we prepare a solution by mixing 10.0
mL of 2.00 x 10-3 M Fe(NO3)3 with 10.0 mL of 2.00 x 10-3 M KSCN. As a result of
reaction (1) some red Fe(SCN)2+ ion forms. The concentration of the Fe(SCN)2+ in this
solution can be determined by finding the absorbance of the solution at the λ max and
using a standard curve to determine its concentration. Let's say the concentration at
equilibrium is found to be 1.50 x 10-4 M. Our problem is to determine Kc for the
reaction from this information. To do this we can carry out the following steps and
summarize the information on a reaction (or ICE) table.
1.
Write a balanced equation for the reaction.
2.
Calculate the initial concentrations of the reactants and/or products.
3.
Using stoichiometry, find the change during the reaction.
4.
Calculate the equilibrium (final) concentrations.
5.
Write the equilibrium constant and calculate Kc for the reaction.
Here are the steps for solving Kc for reaction (1) from the information given above.
Fe3+(aq)
+ SCN-(aq)
Fe(SCN)2+(aq)
Calculate the initial [Fe3+], [SCN-] and [Fe(SCN)2+]:
and of course,
[Fe3+]init =
= 0.00100 M
[SCN-]init =
= 0.00100 M
[Fe(SCN)2+]init = 0
Set up the reaction table, Table 1, with x = [FeSCN2+], the concentration of
Fe(SCN)2+ at equilibrium.
Table 1. Reaction Table or Ice Table for equilibrium of Fe(SCN)2+.
Concentration (M)
Fe3+
+
SCNInitial
0.00100
0.00100
Change
-x
-x
Equilibrium
0.00100 - x
0.00100 - x
Fe(SCN)2+
0
+x
x
To calculate the equilibrium concentration of all the components we must solve for
x. From the reaction table, x is the equilibrium concentration of FeSCN2+ which you
c230 Exp. 5 - Determination of an Equilibrium Constant
3
will determine from the absorbance versus [Fe(SCN)2+] standard curve. In the above
problem it is given as the following:
x = [Fe(SCN)2+] = 0.000150 M
Therefore, the equilibrium concentrations are:
[Fe3+] = 8.5 x 10-4 M
[SCN-] = 8.5 x 10-4 M
and
[Fe(SCN)2+] = 1.50 x 10-4 M
Substituting into the mass-action expression:
Kc =
=
= 2.1 x 102
In this experiment you will determine the absorbance of different solutions at the
wavelength of maximum absorbance (λmax). That data will be transformed into
concentrations similar to those shown in the example above. The calculations
involved in processing the concentration terms are completely analogous to those
illustrated. (Your value for Kc will differ from the one we obtained, since the data in
our example were obtained at a different temperature and thus relate to a different
value of Kc.
EXPERIMENTAL PROCEDURE
ALL SOLUTIONS IN THIS EXPERIMENT CAN BE DISPOSED IN
THE SINK WITH PLENTY OF WATER. COLLECT IN A WASTE
CONTAINER AT YOU DESK AND DISPOSE AS NEEDED.
WORK IN PAIRS
All stock solutions are prepared in 0.10 M HNO3 and all dilutions are done using
0.10 M HNO3, therefore the H+ concentration is maintained at a constant
concentration. The hydrogen ion does not participate directly in the reaction you are
studying, but its presence is necessary to avoid the formation of brown-colored
species such as Fe(OH)2+, which would interfere with the analysis of [Fe(SCN)2+].
You will be using volumetric and graduated pipets and volumetric flasks to
prepare all solutions, and measuring absorbances for various solutions. We will not be
providing detailed instructions since you have done these types of manipulations and
used the spectrometer for the past two weeks.
Part 1. THE STANDARD CURVE
•
PREPARATION OF SOLUTIONS of varying known Fe(SCN)2+ concentration.
Using the techniques from last week, and the volumes given in Table 1,
prepare the series of solutions. Place each solution into another clean, dry,
labeled container and thoroughly rinse the volumetric flask before making
the next solution. These solutions are used to establish a standard curve of
absorbance at λ max for Fe(SCN)2+ versus Fe(SCN)2+ concentration. (See Table
1 in What-to-Do.)
c230 Exp. 5 - Determination of an Equilibrium Constant
4
Table 1. Composition of Solutions for Standard Curve
•
•
Solution
0.200 M Fe(NO3)3
(in 0.10 M HNO3), mL
0.00200 M NaSCN
(in 0.10 M HNO3), mL
0.10 M HNO3, mL
1
2
3
4
5
10.00
10.00
10.00
10.00
10.00
0.00
1.00
2.00
3.00
4.00
add to total 50.00
add to total 50.00
add to total 50.00
add to total 50.00
add to total 50.00
DETERMINING λ max, for Fe(SCN)2+.
Rinse (x3) cuvette with Solution 1. Calibrate the Spectrophotometer with a
cuvette containing Solution 1. Rinse (x3) cuvette with Solution 3 to prepare a
cuvette with solution 3 and record an Absorbance vs. λ plot. Save your plot
and select absorbance vs. concentration (mol/L). Select the appropriate λ.
STANDARD CURVE.
Collect absorbance-concentration data of solutions 1 through 5 at λmax after
switching to Absorbance vs. Concentration in LoggerPro. Record each
solution's absorbance in both the computer and in your laboratory notebook
(Table 1). Make sure to enter the concentrations determined for Prelab in
Table 1.
Use the application to find the best-fit line linear fit through your data and
use the equation for this line as your standard curve. Record the calculated
equation for the line in your notebook.
LEAVE THE Application OPEN with the data and proceed to Part 2.
Part 2 . Determination of Equilibrium Constant (Kc)
BEFORE PROCEEDING YOU MUST Prepare 0.00200 M Fe(NO3)3.
A solution of 0.00200 M Fe(NO3)3 is not available. To make this solution,
dilute 0.50 mL of 0.200 M Fe(NO3)3, measured with a 1.00-mL pipet, to 50
mL with 0.1 M HNO3 in a volumetric flask.
•
PREPARATION OF TEST SOLUTIONS (Table 2).
Dry and label 5 medium size test tubes as suitable containers for these
solutions. Stir each test solution for approximately one minute until the color
does not change.
With the Spectrophotometer set at λmax for the complex, measure and record
absorbance of solutions 6 through 10.
Table 2. Test Solutions for Kc Determination
Solution
0.00200 M Fe(NO3)3*
(in 0.10 M HNO3), mL
0.00200 M NaSCN
(in 0.10 M HNO3), mL
0.10 M HNO3, mL
6
7
8
9
10
5.00
4.00
3.00
2.00
1.00
1.00
2.00
3.00
4.00
5.00
add to total 10.00
add to total 10.00
add to total 10.00
add to total 10.00
add to total 10.00
c230 Exp. 5 - Determination of an Equilibrium Constant
5
•
DETERMINATION OF FE(SCN)2+ EQUILIBRIUM CONCENTRATION.
From the absorbance value of each solution and the equation for the
standard curve determine the Fe(SCN)2+ equilibrium concentration for each
solution.
Calculate Kc for each solution. Determine the average Kc, its range and 95%
confidence interval of the mean and report with appropriate significant
figures.
ALL SOLUTIONS IN THIS EXPERIMENT CAN BE DISPOSED
IN THE SINK WITH PLENTY OF WATER.
QUESTIONS
Q.1. When Fe3+ and SCN- react to reach an equilibrium with Fe(SCN)2+, what
happens to the initial conc. of Fe3+ in part 2? How are the number of moles
Fe(SCN)2+ produced and the number of moles Fe3+ used up related to each
other?
Q.2. A student mixes 5.0 mL of 2.00 x 10-3 M Fe(NO3)3 with 5.0 mL of 2.00 x 10-3
M NaSCN. She finds that in the equilibrium mixture the concentration of
Fe(SCN)2+ is 1.4 x 10-4 M. Use the data to put together a reaction table and
calculate Kc. Show all work.
Q.3. After calibration of your spectrometer with the blank, you inadvertently put
a fingerprint on the cuvette which affects the absorbance reading for a
Fe(SCN)2+ solution. Will the absorbance read higher or lower than it should
in this case? Briefly explain.
Q.4. If the absorbance of a solution reads greater than 2.0 on the
spectrophotometer, how could the preparation of the solution be modified to
obtain a more accurate reading?
Q.5. Why can you calculate [Fe(SCN)2+] from the amounts of reactants mixed in
Part 1, but you can not calculate [Fe(SCN)2+] from the amounts of reactants
mixed in Part 2?
c230 Exp. 5 - Determination of an Equilibrium Constant
6
WHAT TO DO
To Be Turned in From Exp. 4- Determination of Red#40 and Blue#1 dyes in Koolaid™:
1) Please hand in the following items in a packet STAPLED:
a) A formal Results and Discussion for this lab. (The Appendix can be copy pages
from notebook). See the Guidelines and Exp. 4 Handout for additional
information.
You are collecting data in pairs BUT the written report is individual. You may
work on the figures with your partner but make sure that when it comes to
writing you are working alone.
2) Upload a copy of the report to Turnitin electronically. (Without Appendix if
copy pages were turned in.)
Please Note:
Your Absorbance Spectra for the dyes may be in two separate figures. As well, the
Calibration Curves may also be in two separate figures.
For Experiment 5
Before Lab:
Read this experiment. Fill in the “Title bar” to include the
experiment #, title, your name, partner's name, date, lab section, and TA's name. This
information should be included on each page in your laboratory notebook for every
experiment.
All prelab material should be written in your laboratory notebook; your TA at the
start of the lab will grade the pre-lab assignment.
For this week’s experiment your Prelab write-up should include, in addition to the
completed title bars, the following sections.
1.
Purpose
After carefully reading the experiment, briefly state what you understand to
be the purpose of this lab exercise.
2
Table of Reagents for Fe(NO3)3, KSCN, and HNO3. BE SURE TO NOTE THE
DISPOSAL PROCEDURES.
Your lab book should also have the following information and tables ready for
use:
3.
Prepare the following two full page tables in your notebook before coming
to lab (please note that these tables are extended versions of Tables 1 and 2 in
the experiment). Each of these tables should be one full page and should be
oriented sideways in your notebook (“landscape”). Before lab you should
have calculated all initial concentrations and entered them in the appropriate
tables in your notebook.
c230 Exp. 5 - Determination of an Equilibrium Constant
7
During the Lab: Carry out the procedure as written. Modify your procedure if you
find it necessary, but be sure to record the modifications. Record all observations and
in Part 1 and Part 2. Put the data in the appropriate tables which you have prepared in
your laboratory notebook. Plot the standard curve in LoggerPro and find the best-fit
line for your calibration curve.
To Be Turned In NEXT WEEK
NOTE: Specific pages and order of assignment to be turned in will be state on the
Moodle page for the next experiment.
The report for this experiment includes:
• Part 1. The Standard Curve
A printed copy of Table 1 (prepared in a Word document). Include a title.
Sample calculations with all units for Table 1.
Printed copy of Absorbance Spectrum copied from LoggerPro and pasted into a
Word doc. Include appropriate caption.
Printed copy of Standard Curve copied from LoggerPro and pasted into a Word
doc. Include appropriate caption.
• Part 2. Determination of Equilibrium Constant (Kc)
Printed copy of a Completed Table 2 (prepared in a Word Document). Don’t
forget to include your Mean Kc and 95% confidence interval of the mean.
All sample calculations with all units for Part 2.
• Questions and Sample Calculations
Your responses to questions and sample calculations should be found on
clearly labeled summary pages (from your notebook). Report answers to the
questions in complete sentences.
c230 Exp. 5 - Determination of an Equilibrium Constant
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