Topic 7: Equilibrium and LeChatelier’s Principle Experiments A, B and C Evaluation and conclusion questions
LeChâtelier’s Principle
What happens when the balance is disturbed—due to changes in reactant and product concentrations,
temperature, or pres- sure, for example? LeChâtelier’s Principle provides an intuitive explanation for how balance
can be restored:
“If an equilibrium system is subjected to a stress,
the system will react in such a way as to remove the stress.”
To remove a stress, one of two things can happen. The reaction can shift in the forward direction to make more
products, thus using up reactants. Alternatively, the reaction can shift in the reverse direction to re-form the
reactants, thus using up products.
Adding a reactant to an equilibrium mixture of reactants and products always shifts a reversible chemical
reaction in the for- ward direction, to use up the excess reactant and make more products. Removing a product
from a mixture already at equilibrium has the same effect. Conversely, adding a product or removing a reactant has
the opposite effect, shifting a reaction in the reverse direction to use up the excess product and make more reactants.
The effect of temperature on a system at equilibrium depends on whether a reaction is endothermic or
exothermic. If a reac- tion is endothermic, heat may be considered as a reactant in the forward reaction.
(Endothermic reactions require heat in order to proceed.) Increasing the temperature of an endothermic, reversible
chemical reaction shifts the position of equilibrium in the forward direction, to use up the excess heat and make
more products. In the case of an exothermic chemical reaction, heat may be thought of as a product of the forward
reaction. Increasing the temperature of an exothermic, reversible chemical reaction shifts the position of
equilibrium in the reverse direction, in order to remove the excess heat represented by the higher temperature. The
ammonia synthesis reaction is exothermic (Equation 2). Increasing the temperature of this reaction shifts the
equilibrium in the reverse direction, to use up some of the excess heat. Ammonia decomposition becomes more
significant at higher temperatures.
N2(g) + 3H2(g) 
2NH3(g) + heat
Equation 2
Equilibria governing reactions in aqueous solutions play a vital role in environmental and biological chemistry.
The pH and alkalinity of drinking water depend on acid–base equilibria. The factors that control the movement of
solutes in groundwater or surface water depend on solubility equilibria. The binding of oxygen to hemoglobin
depends on complex-ion equilibria involving iron(II) ions bound to porphyrin complexes in the protein.
Overview of the Experiments
The properties of an acid–base indicator illustrate the nature of acid–base equilibria. An indicator is an
intensely colored compound that can gain or lose a hydrogen ion to form substances that have different colors.
Equation 3 illustrates the reaction of the indicator bromcresol green (HIn). HIn represents the uncharged
indicator molecule and In– is the indicator ion after the molecule has lost a hydrogen ion. The position of
equilibrium for this reaction depends on the H+ concentration and can easily be shifted in one direction or
another by adding or removing H+. In examining the behavior of equilibria involving H+ ions, it is important to
remember that OH– reacts irreversibly with H+ to give water (Equation 4). Adding excess OH– to a reversible
acid– base reaction has the effect, therefore, of removing H+ from the reaction mixture.
HIn(aq) 
H+(aq) + In–(aq)
Yellow
Blue
Equation 3
H+(aq) + OH–(aq) H2O(l)
Equation 4
The solubility of magnesium hydroxide in aqueous solution illustrates the nature of solubility equilibria. A
saturated aqueous solution of magnesium hydroxide is prepared by the reaction of magnesium chloride with
sodium hydroxide in water (Equation
5). The solubility equilibrium observed in a saturated aqueous solution of magnesium hydroxide involves solid
salt on one side and dissolved aqueous ions on the other side (Equation 6). What are the predicted effects of adding
excess OH– ion, adding H+ (in the form of HCl), and adding EDTA (a reagent that reacts irreversibly with Mg2+
ion, according to Equation 7)?
MgCl2(aq) + 2NaOH(aq) Mg(OH)2(s) + 2NaCl(aq)
Mg(OH)2(s) 
Mg2+(aq) + 2OH–(aq)
Equation 5
Equation 6
Topic 7: Equilibrium and LeChatelier’s Principle Experiments A, B and C Evaluation and conclusion questions
Mg2+(aq) + EDTA(aq) MgEDTA(aq)
Equation 7
Reaction of ferric nitrate with potassium thiocyanate illustrates the nature of complex-ion equilibria. Aqueous
iron(III) ion (Fe3+) is pale yellow. Adding colorless thiocyanate ion (SCN–) results in the formation of a complex
ion in which the anion is tightly bound to the metal ion (Equation 8). The resulting iron–thiocyanate complex ion
is a deep red color. The effects of adding excess reactants on the position of equilibrium will be observed. The
effect of adding sodium phosphate (NaH2PO4), which under- goes a competing reaction with Fe3+, will also be
examined. Finally, the effect of temperature will be studied in order to determine whether the reaction is
endothermic or exothermic in the forward direction.
Fe3+(aq) + SCN–(aq) 
FeSCN2+(aq)
Equation 8
Answers these questions on a separate paper to HAND IN TO ME.
1.
Acetic acid (CH3COOH) is the principal ingredient in vinegar. Consider the following reversible
reaction of acetic acid at equilibrium:
CH3COOH(aq) 
CH3COO–(aq) + H+(aq)
a.
b.
Use LeChâtelier’s Principle to predict the effect of adding excess sodium acetate (Na+CH3COO–)
on the concentration of H+ present at equilibrium.
Why will the addition of OH– ions (in the form of sodium hydroxide) cause the concentration of
CH3COO– to increase?
2. Write the reaction equation for the reversible chemical reaction that occurs when the indicator
bromcresol green is dis- solved in water (Part A).
3. Use LeChâtelier’s Principle to explain the color changes observed upon addition of HCl and NaOH,
respectively, to a solu- tion of bromcresol green in water (steps 3 and 4).
4. a. What is the likely composition of the indicator solution when the intermediate or transition color
is observed in step 6?
b. How does this observation provide visual proof of the idea that not all reactions “go to completion”?
Explain.
5. Write the chemical equation for the reversible chemical reaction that occurs when magnesium
hydroxide is dissolved in water (Part B).
6. Use LeChâtelier’s Principle to explain the changes observed upon addition of HCl and NaOH,
respectively, to a saturated solution of magnesium hydroxide in water (steps 11 and 12).
7. Predict the direction in which the solubility equilibrium should shift when EDTA is added to a
saturated solution of mag- nesium hydroxide in water (step 13). Do your observations support this
prediction?
8. Write the reaction equation for the reversible chemical reaction that occurs when iron(III) and
thiocyanate ion are mixed together in aqueous solution (Part C).
9. How was the composition of the solution (relative amounts of reactants and products) affected
when the solution was cooled (step 19)? When the solution was heated (step 20)?
10. Based on the observed effect of temperature on the position of equilibrium, is the forward reaction
for this reaction endo- thermic or exothermic? Explain, using LeChâtelier’s Principle.
11. Use LeChâtelier’s Principle to explain the color changes observed upon addition of excess Fe3+ and
SCN–, respectively, to an equilibrium mixture of reactants and products in this complex-ion reaction
(steps 21 and 22).
12. After observing the effect of NaH2PO4 on the equilibrium mixture (step 23), a student was
skeptical that both Fe3+ and SCN– were still present in solution. Can you think of additional
experimental steps that could be done to prove that both reactants are still present at this point?