University of Nairobi Henri Lewis Le Châtelier Chemistry Lab Worksheet

IntroductionWhen the conditions of a system at equilibrium are altered, the system responds in such a way as to maintain
the equilibrium. In 1888, Henri-Lewis Le Châtelier described this phenomenon in a principle that states,
“When a change in temperature, pressure, or concentration disturbs a system in chemical equilibrium, the
change will be counteracted by an alteration in the equilibrium composition.”
This experiment demonstrates Le Châtelier’s principle at work in a reversible reaction between iron(III) ion
and thiocyanate ion, which produces iron(III) thiocyanate ion and the reaction between cobalt(II) hexa-aqua
ion and chloride ions which produces cobalt(II) tetrachloro ion and water. These reactions are shown below.
Fe3+(aq) + SCN-
(aq)
Co(H2O)6 2+
(aq) + 4Cl (aq)
FeSCN2+ (aq)
CoCl 42- (aq) + 6H2O(l)
The concentration of one of the ions is altered either by directly adding a quantity of one ion to the solution
or by selectively removing an ion from the solution through formation of an insoluble salt. Observations of
color changes indicate whether the equilibrium has shifted to favor formation of the products or the reactants.
In addition, the effect of a temperature change on the solution at equilibrium can be observed, which leads
to the ability to conclude whether the reaction is exothermic or endothermic.
To fully understand Le Châtelier’s Principle, a reversible reaction of the sort expressed by the following
chemical equation is considered:
aA + bB
cC + dD
This reaction consists of two competing processes: the forward reaction, in which the products C and D are
formed from the reactants, and the reverse reaction, in which the reactants A and B are formed from the
products. When the rates of these two processes equal each other, there is no net change in the concentration
of either the products or the reactants, and the reaction is said to be at equilibrium. The ratio of the equilibrium
concentrations of the products to the equilibrium concentrations of the reactants is a constant, as shown by
the following equation:
𝐾𝑐 =
[𝐶]𝑐[𝐷]𝑑
[𝐴]𝑎[𝐵]𝑏
where Kc is the equilibrium constant. The brackets signify the concentrations of the various species, and the
lowercase letters represent the number of moles of each substance involved in the balanced equation. In the
case of the reaction between iron(III) and thiocyanate ions shown previously, the equilibrium constant is:
𝐾𝑐 =
[𝐹𝑒𝑆𝐶𝑁2+]
[𝐹𝑒3+][𝑆𝐶𝑁−]
When the concentration of either a reactant or a product in an equilibrium solution is altered, the
concentrations of the other species must change in order to maintain the constant ratio of products to
reactants. These changes are referred to as “shifts” in the equilibrium. The equilibrium can either shift to the
left, meaning it proceeds in the reverse direction and the concentrations of the reactants increase, or shift to
the right, meaning it proceeds in the forward direction and the concentrations of the products increase. In the
reaction between iron(III) and thiocyanate ions, a shift to the left would mean formation of more iron(III)
and thiocyanate ions, while a shift to the right would mean formation of more iron(III) thiocyanate ions.
The equilibrium constant is dependent upon the temperature; thus, a change in the temperature of an
equilibrium solution can also result in a shift to the right or left, depending on whether the reaction is
exothermic or endothermic. For an exothermic reaction, the heat generated by the reaction can be represented
as residing on the product side of the equation, since heat is produced along with the products:
aA + bB
cC + dD + heat
If heat is added to the system by increasing the temperature, the equilibrium shifts to the left, and the
concentrations of the reactants increase. For an endothermic reaction, the addition of heat would result in a
shift to the right.
aA + bB + heat
cC + dD
In this case, the concentrations of the reactants would increase with an increase in temperature.
Part A:
In this part of the experiment we investigate the following equilibrium reaction:
Fe3+(aq) + SCN-
(aq)
FeSCN2+ (aq)
Photos of control samples of solutions of 1.0 M Fe3+ (aq) and 1.0 M SCN-(aq) along with a mixture of equal
volumes of 1.0 M Fe3+(aq) and 1.0 M SCN-(aq) at room temperature are shown below.
Five test tubes were prepared each containing 1 mL of 1.0 M Fe(NO3)3(aq) and 1mL of 1.0 M KSCN were
prepared and treated as described in the table below.
Test tube
Treatment
1
Fe(NO3)3(aq) added
2
KSCN(aq) added
3
AgNO3(aq) added
4
K3PO4(aq) added
5
Heated until 50.0 oC
Questions:
Part A: Fe3+(aq) + SCN-
(aq)
FeSCN2+ (aq)
1. If the equilibrium shifts to the right would you expect the solution to become more red or more yellow?
Explain your reasoning.
2. When Fe(NO)3(aq) is added to test tube 1 would you expect the solution to become more yellow or
more red? Explain your answer in terms of how the equilibrium is shifted.
3. When KSCN(aq) is added to test tube 2 would you expect the solution to become more yellow or more
red? Explain your answer in terms of how the equilibrium is shifted.
4. Write the net ionic equation for the precipitation reaction that occurs when a solution of AgNO3(aq) is
added to a solution containing SCN- ions. Be sure to include correct phase labels.
5. When AgNO3(aq) is added to test tube 3 would you expect the solution to become more red or more
yellow? Explain your answer in terms of the reactions occurring and how the equilibrium is shifted.
6. Write the net ionic equation for the precipitation reaction that occurs when a solution of K3PO4(aq) is
added to a solution containing Fe3+ ions. Be sure to include correct phase labels.
7. When K3PO4(aq) is added to test 4 would you expect the solution to become more red or more yellow?
Explain your answer in terms of the reactions occurring and how the equilibrium is shifted.
8. When test tube 5 is heated to 50.0 oC the solution became significantly more yellow. Is the reaction
exothermic or endothermic? Explain your answer in terms of how the temperature change caused the
equilibrium to shift.
Part B:
In this part of the experiment we investigate the following equilibrium reaction:
Co(H2O)62+(aq) + 4Cl- (aq)
CoCl42- (aq) + 6H2O(l)
2+
2An equilibrium mixture of Co(H2O)
6 (aq) and CoCl 4 ions, in the presence of Cl ions, is purple in color due
to the presence of blue and pink cobalt containing ions. The series of experiments in the following video
were performed in order to determine which of the cobalt containing ions is pink and which is blue. Use the
following link to play the video and be sure to record your observations.

Questions:
Part B: Co(H2O)62+(aq) + 4Cl- (aq)
CoCl42- (aq) + 6H2O(l)
1. When HCl(aq) was added to the equilibrium mixture of ions did the solution become more blue or
more pink?
2. Which of the two cobalt containing ions is blue? Which is pink? How do you know?
3. Why did the solution become more pink upon addition of H2O?
4. Write the net ionic equation for the precipitation reaction that occurs when AgNO3(aq) is added to a
solution containing Cl- ions. Be sure to include correct phase labels.
5. If a solution of AgNO3(aq) were added to an equilibrium mixture of Co(H2O)62+ (aq) and CoCl42(aq), would you expect the solution to become more pink or more blue? Explain your answer in
terms of how you would expect the equilibrium to shift.
6. Is the reaction Co(H2O)62+(aq) + 4Cl- (aq)
CoCl42- (aq) + 6H2O(l) exothermic or endothermic? How
do the results from the experiment allow you to determine this?