Lab 5 (Q) Flashcards
Le Châtelier’s Principle and Position of Equilibrium
Many chemical reactions do not go to completion, instead leaving some reactants and some products. (Other reactions do proceed to completion, consuming all of the limiting reactant.)
When a condition is reached in which the concentrations of the products and reactants do not change with time, a state of chemical equilibrium exists.
This does not mean that all chemical reactions have stopped, but rather that the rate of product formation equals the rate of reactant formation.
Equilibrium is dynamic.
Changing the concentration of one or more of the reactants or products will alter the position of chemical equilibrium.
Changing the temperature or pressure of the system can also alter the position.
Le Châtelier’s principle provides a mechanism to understand how these factors affect the equilibrium.
Le Châtelier’s principle
This principle states that, when an external stress is applied to a
chemical system at equilibrium, the system responds by shifting the equilibrium position in the
direction that minimizes the effects of the external stress.
In other words, the stress created by the addition (or removal) of entities is relieved by a shift in the position of equilibrium in the direction that offsets the imposed stress.
To illustrate Le Châtelier’s principle, consider the generalized reaction
aA(aq) + bB(aq) ⇌ cC(aq) + dD(aq)
If either reactant A or B is added to the solution, Le Châtelier’s Principle predicts that the system responds in a way that counteracts the stress (the increased [A] or [B]); hence, some A and B will react to form more C and D.
As well, If either product C or D is removed, Le Châtelier’s principle predicts that the system responds in a way that counteracts the stress (the decreased [C] or [D]); again, some A and B will react to form more C and D.
reaction quotient (Q):
Q can take on any positive value between zero (all reactants and no products) and infinity (all products and no reactants).
The value of Q under nonequilibrium conditions will change, always moving towards the equilibrium constant, K.
If the reaction favours product formation, (i.e. the forward reaction is spontaneous), then the value of Q will increase.
If the reaction favours reactant formation, (i.e. the reverse reaction is spontaneous), then the value of Q will decrease.
If the reaction is at equilibrium (i.e. neither the forward or the reverse reaction is spontaneous), then the reaction quotient is a constant — the equilibrium constant.
The generalizations, derived from Le Châtelier’s principle, can now be rewritten in terms of Q
- If a stress makes the value of Q less than the value of Keq (more reactants or less products than
required for equilibrium), the position of equilibrium shifts to the right, i.e., the reaction proceeds to form more products until Q = Keq. - If a stress makes the value of Q greater than the value of Keq (less reactants or more products than required for equilibrium), the position of equilibrium shifts to the left; i.e., the reaction proceeds to form more reactants until Q = Keq.
Equilibrium constants are temperature dependent.
Heat may be considered either as a reactant or as a product.
When heat behaves as a reactant, the equilibrium (as written) is endothermic.
Likewise, when heat behaves as a product, the equilibrium (as written) is exothermic.
Consider the reaction:
aA(aq) + bB(aq) ⇌ cC(aq) + dD(aq) + energy
Adding or removing energy (changing the temperature) can be treated by Le Châtelier’s principle,
like any other reagent.
Adding energy (heat) will cause the reaction to shift left to remove energy.
In Part I, you will examine the equilibrium between iron(III), Fe3+, and thiocyanate, SCN–, ions to
form the iron(III) thiocyanate complex, [Fe(SCN)]2+:
Fe3+(aq) + SCN–(aq) ⇌ [Fe(SCN)]2+(aq) (1)
The equilibrium position for reaction (1) will shift right if Fe3+ or SCN– is added.
Conversely, the equilibrium will shift left if either Fe3+ or SCN– is removed.
The concentrations of Fe3+ or SCN– can be decreased by the addition of various reagents that react with either the free Fe3+ or SCN– to form
complex ions in a secondary reaction.
For example, the hydroxide anion (OH–) reacts with free Fe3+ to form insoluble Fe(OH)3(s). reaction (1) responds to a drop in [Fe3+] by shifting to the reactant side replacing the depleted Fe3+ until equilibrium is reestablished.
Similarly, the [SCN–] can be decreased by adding Ag+. The silver ion reacts with the thiocyanate
anion to form AgSCN(s). Equilibrium (1) responds to the decrease in [SCN–] by again shifting to the
reactant side to replace the depleted SCN– until equilibrium is re-established.
You will determine the effect of several reagents on the position of the primary equilibrium and explain your observations in terms of Le Châtelier’s Principle.
In solution [Fe(SCN)]2+ has a rust (reddish) color (in very dilute solutions, it may appear to be a
lighter gold or orange color).
The SCN– is colorless and Fe3+ is pale yellow.
Hence, the color of the solution as well as the appearance of a precipitate can be used to monitor the position of equilibrium in reaction (1).
An increase in color intensity indicates an increase in the concentration of the [Fe(SCN)]2+ complex ion, and vice versa.
The following chemical equations will help you identify the interacting equilibria:
Fe3+(aq) + 3OH–(aq) ⇌ Fe(OH)3(s)
Fe3+(aq) + 3C2O42–(aq) ⇌ [Fe(C2O4)3]3–(aq)
Ag+(aq) + SCN–(aq) ⇌ AgSCN(s)
keep in mind that Fe(OH)3(s)is rust colored, AgSCN(s) is white and [Fe(C2O4)3]3– is colorless.
In Part II, you will determine the value of the equilibrium constant (in 0.05 mol/L HNO3) for the
primary reaction:
Fe3+(aq) + SCN–(aq) ⇌ [Fe(SCN)]2+(aq)
The equilibrium constant expression for the primary reaction, Keq, is defined as the ratio of
equilibrium concentrations:
keq = [[Fe(SCN)]^2+] / [Fe^3+][SCN-]
By using a spectrophotometer to determine the concentration of product [Fe(SCN)]2+ formed, and
by knowing the initial concentration of reactants (Fe(NO3)3 and KSCN), the concentration of all entities at equilibrium can be determined.
When these equilibrium concentrations are substituted
into the expression for Keq, the equilibrium constant can be determined.
This will be done for four runs and the average value will be reported.
The reaction is carried out in 0.05 mol/L nitric acid (excess acid) because:
- the reaction is pH dependent;
- the reaction is sensitive to ionic strength; and
- under basic/neutral conditions Fe3+(aq) tends to form iron(III) hydroxo complexes.
The addition of a small amount of nitric acid prevents these complications from interfering with
your experiment.
[Fe(SCN)]2+ has a maximum absorbance of visible light at a wavelength of 447 nm.
%T = (it / i0) x 100%
and A = - logT = log(100 / %T)
Beer’s Law relates absorbance to concentration, path length and molar absorptivity:
A = ebc
where
- concentration, c: in mol/L of the absorbing entities.
- path length, b : in cm, is the inner width of the cuvette. (b = 1.00 cm)
- molar absorptivity, ε: in (mol/L)–1cm–1, of the entities. At 447 nm, ε = 2200 L/(mol∙cm) for Fe(SCN)2+. Remember, molar absorptivity (ε) is a constant for a given chemical at a given wavelength: ε is not dependent on concentration or path length.
