Le Châtelier's Principle
(revised 7 March 2023)
Introduction
Most chemical reactions do not go to completion.
The reactants initially produce products but after a period of time the
concentrations of
reactants and products remain constant. This cessation of reaction implies that
the chemical
system has reached a state of dynamic equilibrium. For the reaction,
chemical equilibrium is established when the
forward and reverse reactions occur at the same rate. If the
concentration of one of the species in the equilibrium system is
changed, the temperature is changed or the pressure is modified by
changing the volume of the reaction vessel then the position of the
equilibrium (ratio of reactants to products) shifts to oppose the
change (i.e. if more H2
is added, more
NH3
is formed to use up some of the added H2
).
A general statement governing all chemical systems in a state of dynamic
equilibrium is
that "if an external stress is applied to a chemical system at equilibrium, the
position of
the equilibrium shifts in such a manner as to minimize the effect of the
stress." This is a
statement of Le Châtelier's Principle as proposed by Le Châtelier
in 1888.
Two factors that affect the position of equilibrium in a system will be studied
in this
experiment; namely, a change in concentration and a change in temperature.
Safety
Wear your safety glasses. This experiment involves corrosive and poisonous
compounds - avoid all skin contact with them.
Procedure
For each manipulation, it is essential that you accurately record your
observations. In addition, try to determine if the reaction is
exothermic or endothermic.
A. Saturated NaCl Solution
In this section, the saturated NaCl equilibrium will be examined.
NaCl (s) 
Na+
(aq) + Cl
-
(aq)
To 2 mL of H2O in a test tube, add an excess of NaCl and shake. Observe. Transfer the saturated NaCl
solution to a second test tube. Add two or three drops of concentrated HCl to
the saturated solution. Observe. Note that there are two phases, liquid and
solid. The liquid above the solid is called the supernatant.
Determine its color. The solid is called the precipitate.
Determine its color.
B. Chromate ion – dichromate ion equilibrium
In
this section, the color of both chromate ions in
the equation below will be determined (note other equations may be available in your textbook or from a credible online source).
2 CrO42-
(aq) + H2O (l)
Cr2O72-
(aq) + 2 OH
-
(aq)
Place 2 mL of 1M K2CrO4
in a test tube. Observe. Add 2 or 3 drops of 6M HNO3. Observe and note change. Add 2 or 3 drops of 6M NaOH. Observe and note change. Add 2 or 3 drops of
6M HNO3. Any effect?
C. Hexaaqua cobalt(II) ion – tetrachlorocobaltate(II) complex
ion
equilibrium
In
this section, the color of both cobalt complexes in
the equation below will be determined.
4 Cl-
(aq) + [Co(H2O)6]2+
(aq)
[CoCl4]2-
(aq) + 6 H2O
(l)
Place 1 g of Co(NO3)2 · 6 H2O in a test tube. Add 5 mL of H2O.
Observe. Slowly add concentrated HCl. Observe.
Divide the resulting solution into 4 test tubes. Treat as
follows:
- Solution 1: Blank - leave as is.
- Solution 2: Place test tube in boiling water bath. Observe.
Compare color to solution 1. Now place in an ice water
bath. Observe. Compare color to solution.
- Solution 3 & 4: Saturate with NaCl (about 0.5 g). Observe. Compare with solution 1.
- Solution 4: Using an eye dropper, aspirate the supernatant (liquid
on top) and place into another test tube. Add 5 to 10 drops of 1M
AgNO3
centrifuge. Examine.