INVESTIGATING LE CHATELIER’S PRINCIPLE
Introduction
When a reaction has reached dynamic equilibrium, it tends to want to stay there and will resist changes to the
equilibrium. This is the basis of Le Chatelier’s principle which states that:
A closed system at dynamic equilibrium will respond to disturbances in such a way as to minimise the
disturbance.
For example:
 If you increase the concentration of a reatant, the equilibrium will shift so as to lower the
concentration to somewhere near where it was
 If you lower the temperature, the equilibrium will shift in such a way as to bring the temperature back
up a bit
 If you increase the pressure, the equilibrium will shift in such a way as to reduce the pressure again.
This comes about not through sorcery but as a result of the rules that govern the rates of the forwards and
backwards reactions. For example, if you increase the concentration of a reactant, this would increase the
rate of the forward reaction and thus use up a lot of the extra you just added.
In this practical you will be looking at a number of different equilibrium systems in order to better understand
Le Chatelier’s Principle. The experiments are qualitative rather than quantitative so don’t worry so much
about exact quanities and focus more on what you can see happening.
Reaction 1: Nitrogen dioxide/dinitrogen tetroxide
Intro:
You are going to look at the following equilibrium involving the gasses nitrogen dioxide (NO 2) and
dinitrogen tetroxide (N2O4):
2 NO2(g)
red/brown
Method:

N2O4(g)
colourless
H –ve
STEPS 1-3 MUST BE DONE WEARING SAFETY GLASSES AND PROTECTIVE GLOVES IN A FUME HOOD
1. Place about 10 cm3 of concentrated nitric acid in a flask.
2. Carefully add a piece of copper.
3. Carefully fill two gas syringes with 50 ml of the gas and then seal the end. One of these will be
used in steps 4-8, the other will kept as a control to compare against.
4. Immerse the gas syringe in a hot water bath at 80-90OC and note any changes.
5. Immerse the syringe in an ice bath and note any changes.
6. Return the syringe to room temperature and note any changes.
7. Push the syringe in and hold, note any changes. (Do not pop the seal!)
8. Pull the syringe out and note any changes
https://www.youtube.com/watch?v=ZVzrijJdl4c
Analysis:

Write an expression for Kc and explain all observed changes. If you observed no changes (this
can be a bit temperamental) state the changes you would have expected to see.
SOLUTION:
Kc = [N2O4]/[NO2]2
For an exothermic reaction cold  hot.
So decreasing temperature in an ice bath shifts equilibrium to the hot side, hence favours the product side. In this
case getting more colourless as N2O4 is produced. Increasing temperature in a warm water bath shifts equilibrium to
the cold side, hence favours the reactant side. In this case getting more red/brown as more NO2 is produced.
Reaction 2: The Reaction of Copper(II) ions with ammonia
Intro:
Copper 2+ ions form two separate equilibria with aqueous ammonia depending on how much we add:
[Cu(H2O)6]2+(aq)
+
transparent light blue
[Cu(H2O)4(OH)2](s) +
cloudy pale blue
Method:
2 NH3(aq)
4 NH3(aq)


[Cu(H2O)4(OH)2](s)
cloudy pale blue
[Cu(NH3)4(H2O)2]2+(aq)
+
dark transparent blue/purple
+
2 NH4+(aq)
2 H2O(l)
+
2 OH-(aq)
1.
Place approximately 2 cm3 of 0.1 M CuSO4 in a test tube. CuSO4 is a light blue solution.
2.
3.
4.
Add several drops of NH3(aq) and note any changes.*
Add several drops of NH3(aq) and note any changes…keep going until there is no further change.
Add several drops of HCl(aq) and note any changes.*
*If several drops does not produce a change, keep going until there is a change
Analysis:

Write an expression for Kc and use Le Chatelier’s principle to explain all observed changes
SOLUTION:
For each equilibrium, Kc is equal to the product concentrations raised to the power of the coefficient divided by the
reactant concentrations raised to the power of the coefficient. (See solution 1).
1st equilibrium, a cloudy pale blue precipitate is formed as more reactant (NH3) is added. The system opposes this
change by using up the reactant and shifting the equilibrium to the right. The ammonia is behaving as a weak base in
this reaction.
2nd equilibrium, a dark transparent blue/purple solution is formed as adding excess NH 3 results in ligand exchange
with the water ligands (see the equation). Again this reaction occurs, because the system is trying to counteract the
addition of NH3 by using it up.
Reaction 3: The Reaction of Cobalt (II) ions with chloride ions
[Co(H2O)6]2+(aq)
Intro:
+
4 Cl-(aq)
pink

[CoCl4]2-(aq)
+
6 H2O(l)
blue
H +ve
Method:
1. Begin with 10 cm3 or so of 0.2 M cobalt chloride hexahydrate (pink) in a 100 cm3 conical flask.
2. In a fume hood, wearing gloves: carefully add concentrated HCl to the solution until a
noticeable change takes place, keep adding until the change is (just) complete.
3. Note any changes observed. Pink  Blue
4. Place a small sample of the mixture in a test-tube and leave to cool in an ice-bath. Record
changes after the other steps are complete. Blue  Pink
5. Half fill another test-tube and place it in a hot water bath. Leave for a few minutes and record
any changes observed. Pink  Blue
6. Add distilled water to the solution in the flask until a change is observed (and just completed).
Record any changes observed. Blue  Pink
Analysis:

Write an expression for Kc, and use Le Chatelier’s principle to explain all observed changes
SOLUTION:
Kc is equal to the product concentrations raised to the power of the coefficient divided by the reactant
concentrations raised to the power of the coefficient. (See solution 1).
For an endothermic reaction hot  cold.
So decreasing temperature in an ice bath shifts equilibrium to the hot side, hence favours the reactant side. In this
case going pink as the reactant is formed. Increasing temperature in a warm water bath shifts equilibrium to the cold
side, hence favours the product side. In this case going blue as the product is formed.
This reaction is an example of ligand exchange.
Reaction 4: Indicators
Intro:
Indicators function by equilibrium processes. Use your new found understanding (and love) of
equilibrium and Le Chatelier’s principle to develop a theory to explain how they work.
Method:
1.
2.
3.
4.
5.
Half-fill a boiling tube with distilled water and add a large squirt or phenolphthalein indicator
Add a drop of 0.1M NaOH(aq) and record any changes Pink
Add additional drops until the change is completed. Dark Pink
Add 0.1 M HCl(aq) gradually and record any changes. Colourless
Repeat steps 1-4 with bromothymol blue indicator ( Step 2 = Blue, Step 3 = Dark Blue, Step 4
= Yellow)
Analysis:

Use your new-found understanding of equilibrium and Le Chatelier’s principle to propose a
mechanism for how indicators work. This is about you trying to think things through for
yourself, and developing ideas, so don’t worry about being ‘right’ or ‘wrong’.
HINT: Look up proton acceptors and proton donors to try and explain this equilibrium reaction.