Name _________________________________________ Period ___________ WS # _________
Chemical Equilibrium &
LeChâtelier’s Principle
Pre-Lab Discussion
In most of the chemical reactions you have studied thus far, at least one of the reactants has been
“used up”. The point at which a reactant is used up marks the end of the reaction, and the reaction is
said to have “gone to completion”. Under ordinary circumstances, the product(s) of such reactions are
not able to react to re-form the original reactants. Thus, these are “one-way” reactions. They proceed in
one direction only.
Many other chemical reactions do not go to completion. Rather, the products of these reactions
remain in contact with each other and react to re-form the original reactants. Such reactions are said to
be reversible. In a reversible reaction, the forward and reverse reactions proceed at the same time.
When the rates of the two reactions are equal, a state of chemical equilibrium is said to exist. Under
such conditions, both the forward and reverse reactions continue with no net change in the quantities of
either products or reactants.
A state of equilibrium is affected by concentration and temperature and, if gases are involved, by
pressure. If a system at equilibrium is subjected to a change in one or more of these factors, a stress is
placed on the system. According to LeChâtelier’s principle, when a stress is placed on a system at
equilibrium, the equilibrium will shift in the direction that tends to relieve the stress. Equilibrium will
be re-established at a different point, that is, with different concentrations of reactants and products.
In this experiment you will study the equilibrium system involving the yellow chromate ion, CrO4-2, and
the orange dichromate ion, Cr2O7-2. The equilibrium equation is:
H+ (aq) + 2CrO4-2 (aq)
Cr2O7-2 (aq) + OH- (aq)
The addition of an acid (H+) to this system increases the hydrogen ion concentration and causes the
equilibrium to shift towards the formation of more dichromate, orange. The addition of a base (OH-)
causes the removal of H+ ions (formation of water) and causes equilibrium to shift towards the
formation of more chromate, yellow.
By studying this chromate-yellow
dichromate-orange equilibrium system, you
should achieve a better understanding of equilibrium systems and their responses to stress.
Purpose
Study equilibrium systems and their responses to stress as described by LeChâtelier’s principle.
Prodedure
1.
In each of 4 100 mL beakers, mix 10 mL each of 0.1 M potassium (or sodium) chromate and 0.1 M
potassium (or sodium) dichromate.
2.
Add 3 M hydrochloric acid to beaker #1 until the solution turns bright orange.
3.
Add 3 M sodium hydroxide to beaker #2 until the solution turns bright yellow.
4.
Show that these changes are reversible by adding hydrochloric acid to the yellow solution and
sodium hydroxide to the orange solution.
5.
Add 3 M sodium hydroxide to beaker #3 until the solution turns yellow.
6.
Add a volume of 0.1 M barium nitrate equal to about half the volume of the solution in the beaker.
What happens?
The solid formed establishes a new equilibrium:
Ba +2 (aq) + CrO4-2 (aq)
BaCrO4 (s)
7.
Slowly add 3 M hydrochloric acid to the beaker which contains the precipitate until the precipitate
disappears.
8.
Add 1 M sodium hydroxide to this beaker until the precipitate reappears.
9.
Add 3 M hydrochloric acid to beaker #4 until the solution turns orange.
10.
Add an equal volume of 0.1 M barium nitrate. What happens?
11.
Predict what you need to do to beaker #4 in order to make the precipitate appear.
12.
Try it! What happened?
HINT: The appearance and disappearance of the precipitate is another example of LeChâtelier’s
principle. The addition of hydrogen ion converts the chromate ion to dichromate ion. The
solid barium chromate breaks down in an attempt to replace the missing chromate ion. If
enough hydrogen ion is added, the barium chromate is eventually consumed.
DATA CHART
BEAKER
HCl Added
NaOH Added
Solution turned a
darker orange
Solution turned
from orange to
yellow
NaOH Added
HCl Added
Solution turned
from light orange
to yellow
Solution turned
from yellow to
orange
Number
1
BEAKER
Number
2
BEAKER
NaOH Added
Ba(NO3)2 Added
HCl Added
Solution turned
yellow
A white/yellow
precipitate formed
(BaCrO4 )
Solution turned
orange and the
precipitate
dissolved
HCl Added
Ba(NO3)2 Added
Prediction
Try It!
NaOH must be
added to create
chromate ions for
the barium ions to
react with
When NaOH is
added, the solution
turned yellow, and
the white/yellow
precipitate forms.
Number
3
BEAKER
Solution turned
dark orange
Number
4
No Change
NaOH Added
Solution turned
yellow and the
precipitate
reappeared
Post Lab Write-Up
The equilibrium equation for the chromate-dichromate
system is ___1____.
2. ______________________
3. ______________________
1. ______________________________________________
4. ______________________
When hydrogen ions (H+; an acid) are added to this system,
equilibrium shifts towards the formation of __2__ which is __3__
in color.
5. ______________________
6. ______________________
-
When hydroxide ions (OH ; a base) are added to the system,
equilibrium shifts towards the formation of __4__ which is __5__
in color.
7. ______________________
8. ______________________
The yellow solution contains an excess of __6__ ions. When
barium nitrate is added to this solution, the barium ions react with
the __7__ ions to form a precipitate of __8__. When HCl is added
to the beaker containing this precipitate, the precipitate __9__.
The addition of the acid causes equilibrium to shift to convert
chromate to __10__ in order to use up the excess __11__ ions.
This in turn causes the barium chromate to dissolve in order to
replace the lost __12__ ions.
Adding sodium hydroxide now to this same beaker causes
equilibrium to shift to convert __13__ ions to chromate ions. This
excess of CrO4-2 ions react with the barium ions to cause the
precipitate to __14__.
9. ______________________
10. _____________________
11. _____________________
12. _____________________
13. _____________________
14. _____________________
15. _____________________
The addition of __15__ to the fourth beaker caused the
solution to turn __16__ in color. This is because equilibrium
shifted to produce __17__ ions in order to use up excess __18__
ions. When barium nitrate was added to this beaker, __19__
precipitate formed. This is because there was an abundance of
__20__ ions, and very few __21__ ions. Thus no __22__
precipitated in this beaker. You could add __23__ ions to this
solution to cause the precipitate to reform.
16. _____________________
17. _____________________
18. _____________________
19. _____________________
20. _____________________
Word List:
21. _____________________
Orange (2)
Yellow
Cr2O7-2 (5)
CrO4-2 (5)
BaCrO4 (2)
no
H+ (2)
OHHCl
forms
dissolves
22. _____________________
23. _____________________