F.6/7 Chemistry Practical
To Determine Equilibrium Constant
Objective:
To determine the equilibrium constant for the following reaction
Ag+(aq) + Fe2+(aq) <====> Fe3+(aq) + Ag(s)
Group size: Individual
Introduction
If silver nitrate solution and iron(II) nitrate solution are mixed, silver ion and iron(II) ion may react
to form iron(III) ion and metal silver. Silver ion and iron(II) ion form a redox pair. Iron(III) ion and
silver form another redox pair. In this reaction silver ion oxidizes iron(II) ion to form iron(III) ion. To
predict the position of equilibrium one must look up their standard electrode potentials:
Reactions
Standard electrode potential
+
Ag (aq) + e <====> Ag(s)
E1o = 0.799 V
Fe3+(aq) + e <====> Fe2+(aq)
E 2o = 0.771 V
and calculate the overall standard cell potential of the reaction. It is given by
This value is almost equal to zero. This value suggests that, at the state of equilibrium under s.t.p.,
concentrations of all the species, Fe2+(aq), Fe3+(aq), Ag+(aq) and silver are not negligible amounts.
Hence it is possible to determine these concentration / activity. Once they are evaluated, the value of
equilibrium constant, Kc can be calculated.
The Method
Silver nitrate solution and iron(II) nitrate solution are mixed in a flask and silver and iron(III)
nitrate are formed. The reaction mixture is allowed to set aside for the establishment of equilibrium.
After then metal silver residues stay in the bottom of the flask and the supernatant liquid is pipetted out.
Standard solution of potassium thiocyanate is used to titrate with Ag+(aq) in the solution. The end-point
is indicated by the blood red colour of iron(III) thiocyanate which is a complex formed by the following
reaction.
Fe3+(aq) + SCN-(aq)
FeSCN2+(aq)
Iron(II) nitrate is very unstable and it is susceptible to be oxidized by air. It is prepared by the following
reaction.
Fe(s) + Cu(NO3)2(aq)
Fe(NO3)2(aq) + Cu(s)
If the above reaction is filtered, an aqueous solution of iron(II) will be obtained.
Additional Materials required (per student)
0.20 M copper(II) nitrate solution (80 cm3), 0.20 M silver nitrate solution (80 cm3), 0.1 M standard
potassium thiocyanate solution (50 cm3), iron powder, 50 cm3 pipette, 50 cm3 burette, stopper for
conical flask
Pre-laboratory questions
(1)
(2)
(3)
(4)
(5)
(6)
Read the experimental steps carefully and try to list all safety precautions.
Give an expression for the equilibrium constant, Kc, for the following reaction.
Why should a dry conical flask should be used in step 1?
Why is it necessary to stopper the bottle in step 2?
Why is it necessary to let the mixture stand for about half an hour.
Explain the principle of the detection of the point where all silver ion has just used up by
thiocyanate ion [i.e. the equivalent point].
F.6/7 Chem.Prac./Eq2_eqCon/p.1(2)
Procedure
Record your results immediately after you get the readings.Wear goggle in the course of the practical.
Distilled water may be replaced by de-ionized water.
1. Pipette 50.0 cm3 of 0.20 M silver nitrate solution into a dry 250 cm3 conical flask.
2.
Add a large spoon of finely divided iron powder into a beaker. Add 80 cm3 of copper(II) nitrate
solution to the beaker. Stir for 2 minutes.
3.
Add a little iron powder to the filter paper. Use this filter paper to filter the reaction mixture
obtained in step (2). Once you get 30 cm3 of filtrate, pipette 25 cm3 of it and add the aliquot of
filtrate into the conical flask containing silver nitrate solution.
4.
Pipette another 25 cm3 of the filtrate into the flask containing silvere nitrate once again. Mix the
content well and then stopper the flask. Allow to stand for about half an hour. Keep several cm3 of
the filtrate for further investigation (step 5).
5.
Add excess bench ammonia to about 4 cm3 of your filtrate such that all iron ions are precipitated
out. Filter the mixture and get filtrate B. Observe the colour of filtrate B. What ion(s) is/are
present? Give reason.
6.
Pipette 25.0 cm3 of the equilibrium mixture into a conical flask and titrate it with standard
potassium thiocyanate solution. The end point is the first sign of a blood red colour following the
precipitation of all the silver ions as white AgSCN.
Data
Concentration of standard KSCN solution
= ______________
Titration Data
Run
Final burette reading/cm3
Initial burette reading/cm3
Volume of KSCN used/cm3
1
2
3
Calculation
1. What are the initial concentrations of silver and iron(II) ions before the solutions are mixed? What
assumption(s) have you made?
2. How many moles of silver ions were present in the 25.0 cm3 sample of the equilibrium mixture?
3. Calculate the concentration of silver ions in the equilibrium mixture.
4. Hence, or otherwise, determine concentration of iron(III) and iron(II) at equilibrium.
5. Calculate value of Kc.
6. Estimate the error of the value of Kc.
Discussion and Conclusion
Discuss the method used in this practical (That is, how good is it? How poor is it? Give reasons to your
opinions. You should consider error sources.)
The Report: Hand in your report to Mr. Lai (in the staff room) before 5:00 p.m.
F.6/7 Chem.Prac./Eq2_eqCon/p.2(2)