Investigating Chemical Equilibrium
(Taken from your lab manual, Experiment 12A)
The chemical reactions you have studied to this point have proceeded to completion. However, many reactions proceed
only to a state of equilibrium, where both reactants and products exist together. This results from some of the product
molecules having sufficient energy to reform the reactants in a reverse reaction. A state of equilibrium is established
when the rates of the forward and reverse reactions are equal. Although both reactions continue to occur, there is no net
change in the observable properties (called macroscopic properties) such as concentration, pressure (for gases), or colour
intensity of a coloured solution. In a reversible reaction, constant macroscopic properties (including temperature)
indicate that equilibrium is present.
Any change to the conditions at equilibrium, such as concentration, pressure, or temperature, is said to produce a stress
on the equilibrium. Le Chatelier’s principle allows us to predict the effects of such as stress. The stress produces a shift
in the system that attempts to counteract the imposed stress and a new equilibrium is established. The macroscopic
properties of the system change, until they become constant once more as the new equilibrium is reached.
You will study two different equilibrium systems involving ions in solution.
The first system, studied in Part I, will be a wet lab and involves the light yellow iron (III) ion, Fe 3+, and a colourless
thiocyanate ion, SCN-. They react together to form the complex ion FeSCN 2+ which exhibits a blood-red colour:
Fe(3aq )  SCN (aq )
FeSCN (2aq )
The effect of changing the concentration of the reactant ions will be studied.
2
The second system, studied in Part II, will be a dry lab, and involves the hydrated cobalt (II) ion, Co( H 2O ) 6 , which can
be converted to the chlorinated complex,
CoCl42 ( aq ) .
Co( H 2O)62 ( aq )  4Cl  ( aq )
pink
CoCl4 2 ( aq )  6 H 2O( l )
blue
The temperature will be changed and the resultant effect of the position on the equilibrium will be observed. This will
enable a prediction to be made as to whether the reaction as written is exothermic or endothermic.
Objectives:
1.
2.
3.
4.
To
To
To
To
recognize the macroscopic properties of chemical systems at equilibrium
observe shifts in equilibrium concentrations as stresses of concentration changes are applied to the system
observe a shift in equilibrium associated with a change in temperature
explain the observations by applying Le Chatelier’s principle
Supplies:
Equipment
5 test tubes (13 mm x 100 mm)
Test tube rack
Graduated cylinder (10 mL)
Beaker (250 mL)
Chemical Reagents
0.2 M FeCl3
0.2 M KSCN
0.2 M KCl
1 M NaOH
Procedure Part I: Equilibrium involving the thiocyanatoiron(III) ion
1. Add 1 mL of 0.2 M FeCl3 to a 1 mL portion of 0.2 M KSCN in a 250 mL beaker. Note the colour change as the
reaction occurs.
2. Add water to dilute the intensity of the colour until the solution is light orange, to make subsequent changes
easier to see.
3. Pour 5 mL of this solution into each of 5 separate test tubes labelled A to E. Nothing will be added to test tube A
– it will serve as a control for comparison with the subsequent colours.
4. For each of the following reactions (Steps 5 – 8) record the results in your copy of Table 1. To record the
“stress” involved, state which ion in the original equilibrium changed concentration and whether this change was
an increase or decrease.
5. To test tube E, add 10 drops of 1M NaOH .... put this test tube aside to give it time to react. Observe it last.
6. To test tube B, add 10 drops of 0.2 M KSCN
7. To test tube C, add 10 drops of 0.2M FeCl3
8. To Test tube D, add 10 drops of 0.2M KCl
Procedure Part II: Equilibrium involving cobalt (III) complexes
1. Place approximately 0.3 g of
CoCl2 6 H 2O into each of t 100 mL beakrs.
2. To the first beaker, add 10 mL of 6M HCl
3. To the second beaker add 10 mL of water. Note the colours of solutions in each beaker. Record your
observations in your copy of Table 2 in your notebook.
4. Gradually add water to the solution in the first beaker until a definite colour change occurs. Record your
observations in your copy of Table 3.
5. Place the first beaker on a ot plate and adjust the heat to gent warm. When a definite colour change is
observed, shut off the hot plate. Record the resulting colour in Table 3.
6. Add approximately 50 mL of cold tap water to a clean 250 mL beaker. Carefully place the warm beaker and
contents from step 4 upright into this cold water bath. Record any additional changes in colour intensity in Table
3.
Post Lab Considerations
In part II, the formation of FeSCN2+ gave such an intense colour that diluting the solution with water was necessary to
enable small changes in the equilibrium to be observed more readily. In explaining your observations for this equilibrium,
remember that spectator ions do not participate in the reaction. In step 8, a precipitate of Fe(OH) 3 was formed. In
explaining your results, analyze how this reaction will affect the concentration of Fe 3+ in the equilibrium solution.
For the reaction in part III, recall that increasing the temperature of an equilibrium reaction favours the process which
counteracts this change, that is, the endothermic direction. From your observations, you will be able to decide whether
the forward or the reverse reaction is the endothermic reaction.
Results
Table 1
Reagent Added
Stress
Colour Observation
Direction of Equilibrium Shift
KSCN (test tube B)
FeCl3 (test tube C)
KCl (test tube D)
NaOH (test tube E)
Table 2
Colour of CoCl2 + HCl
Blue
Colour of CoCl2 + H2O
Pink
Stress
Colour Observation
Direction of Equilibrium Shift
Decreases [Cl-]
Turns pink
left
Step 5 (heated)
Heat
Turns blue
right
Step 6 (cooled)
Cold
Turns pink
left
Step 4 (water added)
Analysis of Results:
1. Apply Le Chatelier`s Principle to explain the esults obtained when 1M NaOH was introduced into the iron (III)
thiocyanate ion equilibrium system.
2. If the hydrated cobalt (III) ion complex were refrigerated, what would you predict as the colour of the
refrigerated solution?
3. Look at the equilibrium equation for the reaction involving the hydrated cobalt (II) ion complex. From your lab
results, which reaction is endothermic? Cite evidence for your answer.
4. Predict how the addition of sodium chloride, NaCl, would affect the hydrated cobalt (II) ion equilibrium. Explain
your prediction in terms of Le Chatelier’s principle.
Follow up Question:
1. When Ag+ ions are added to the red FeSCN2+ ion in Part I, the colour disappears and a white precipitate of
AgSCN forms. Explain why Fe3+ can be used as an indicator in a reaction to determine the concentration of an
unknown solution of Ag+ by reacting it with a solution of KSCN of known concentration.