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Part A INTRODUCTION
Reactions are processes where reactants collide together to form
products. It is important to note that while this is occurring
products are also being converted back into reactants. Chemical
equilibrium in a reaction is reached when the rate for the forward
reaction equals the rate of the reverse reaction [Equation 1].
Rate of forward reaction = Rate of backward
reaction (1)
If a reaction is at equilibrium, changes in: concentration;
pressure; volume or temperature may result in the system being
disturbed from its equilibrium position and the system will
then proceed in such a way as to restore equilibrium. The rule
for predicting the direction of change is known as Le
Chatelier’s principle.
Table 1 indicates the result of a number of changes to the Haber
process:
N2(g) + 3H2(g) 2NH3(g) DH=-92.6kJmol-1
Table 1: Effects on the Haber process with changes that
disrupt equilibrium.
Changed imposed on How this system alters to regain
equil
Changed imposed on

How this system alters to regain equilibrium system


system
Add more reactants (N2 or H2)
More products are formed immediately to reduce the amount
reactants. Rate forward reaction > Rate reverse
reaction until equilibrium is
achieved
Add more products (NH3)
More N2 and H2 are formed immediately to reduce the
amount products. Rate forward reaction < Rate reverse
reaction until equilibrium is
achieved
Pressure is increased
More products are formed immediately to reduce the pressure (m
of gas). Rate forward reaction > Rate reverse reaction
until equilibrium is achieved
It is clear from this Table 1, that any externally imposed
changes on a system that disrupt
equilibrium will result in the system opposing those changes in
order to get back to equilibrium. For the following general
system:
aA + bB cC +
dD- (3)
the equilibrium constant, K, of a
system may be defined as follows:
Kc = [C]c[D]d /

- (4)
[A]ab
Kc has a constant numerical value for a given system at a
given temperature. The individual concentrations may change,
but the ratio remains the same (at a given temperature).
In Part A of this experiment you will prepare an equilibrium
system for the following reaction:



Fe3+ + SCN- [FeSCN]2+
- (5)
and study the response of this system to various stresses and
explain these results.
Part A PROCEDURE
1. Into a clean 250 mL beaker add 120 mL of distilled water, 10
drops of 1M KSCN and 10 drops of 1 M Fe(NO3)3 solution. This
will be your stock solution.
2. Label 10 clean
test-tubes A-J.
3. Add 5 mL of this stock solution into each one of the clean test
tubes. 4. Keep test tube A as a
control.
5. Add the following reagents into
test-tubes B through
to I: Test tube #
B : 1 mL 1 M Fe(NO3)3
C: 0.5 mL 1 M KSCN
D: 0.25 mL (5 drops) 0.1 M AgNO3








E : 2 mL conc.
HCl
F : 0.5 mL (10 drops) 0.02 M Hg(NO3)2








G : 1 mL 0.1 M Na3(PO4)









H : 1 mL 0.1 M Na2(C2O4)









I : 1 mL 0.01 M
NaF
Make sure students add these reagents very
carefully.
6. Record all your observations and carefully mix (this may not
even be all that necessary).
7. Heat test tube J in a hot water
bath.
8. Compare the colour intensity
of each of the test tubes with
that of the control (test tube A) and
using the following table explain your observations.
Step 8 table in image
Hg2+ Agt Fe3+ SCN- Ag(SCN) (s) CI- [FeCl4] PO43- FePO4 F [FeF613- C2042- [Fe(C204)3]3- Hg(SCN)2 Test Tube Reagent Added Observation Explanation A Control V No change Darker red Loss of red colour B Fe(NO3)3 N/A Increase in the concentration of reactants pushes forward the reaction leading to greater of products Decreasing the concentration of reactants due to side product formation pushes the reverse reaction . C KSCN D AgNO3 . E ConcHCl . F Hg(NO3)3 G Na3PO4 H Nazc204 . 1 Solid Naf J Heated .