C1502 Notes
Chemical Equilibrium 2
Chemical Equilibrium
The magnitude of equilibrium constant (Kc)
(Implication of the magnitude of equilibrium
constant)
By looking at the magnitude of equilibrium
constant Kc, one can tell whether a particular
equilibrium constant favours products or
reactants the magnitude.
• Equilibrium constant Kc, does not neglect how
fast the reaction goes.
• A large value of equilibrium constant Kc,
means that mostly products will be present at
equilibrium.
• A small value of equilibrium constant, means
that mostly reactants will be present at
equilibrium.
• In general; if K >> 1, the equilibrium lies to the
right, the products are favoured
• If K << 1, the equilibrium lies to the left, the
reactants are favoured.
*When the equilibrium is neither large or small
(around 1), neither the products nor the reactants
are strongly favoured.
i.e. the equilibrium mixture contains appreciable
amounts of both reactants and products.
Predicting the Direction of the Reaction
(Or redirecting the direction of the chemical
equation and K)
• Determining the reaction Quotient Q, allows one
to predict the direction an equilibrium reaction
will proceed.
Definition
The reaction Quotient ,Q; is an expression that
has the same form as equilibrium constant
expression but whose concentration values are
not necessarily those at equilibrium.
• Q (reaction Quotient) is formed by
substituting given concentrations or pressures
into an equilibrium expression.
Note: Q is not K, but if Q = K, then the reaction
mixture is at equilibrium.
If Q > K, the reaction will go to the left,
approaching the equilibrium.
Note:
• The equilibrium expression for a reaction
written in one direction is the reciprocal of the
one for reaction written in the reverse
direction.
• It is necessary then to known in which
direction the original equation was written;
since the values of the equilibrium constants
will be different in each case.
i.e. substrates on the right hand side of the
chemical equation will react to form substrates
on the left.
If Q < K, the reaction will go to the right in
approaching equilibrium
i.e. the reaction will achieve the equilibrium by
forming more products.
For the general equation:
aA + bB
pP + qQ
[𝑃]𝑝 [𝑄]𝑞
Q=
[𝐴]𝑎 [𝐵]𝑏
Note: Concentrations are not at equilibrium
Example:
A 50,0 ml reaction vessel contains 1,00 mol N2,
3,00 mol H2, and 0,500 mol NH3. Will more
ammonia formed or dissociates when the
mixture goes to equilibrium at 400 ℃?
The reaction equation is given as;
N2(g) + 3H2(g) 2NH3(g)
Kc is 0.500 at 400 ℃
Ans: ammonia will dissociate.
Calculating Equilibrium Concentrations
Example; Reaction Quotient
Consider system(reaction)
PCl5(g) PCl3(g) + Cl2(g) at 250 ℃
The equilibrium constant Kc = 4.0x10-2
If the concentrations of both Cl2 and PCl3 are
0.3 M while the concentration of PCl5(g) is 3.0 M.
1. Is the system at equilibrium?
2. Is Q larger than, equal to or smaller than K?
3. If the system is not at equilibrium, in which
direction does the reaction proceed?
ans: Q < K
(i.e. Q = 3.0x10-2)
• The reaction is not at equilibrium
• The reaction proceed to equilibrium by
converting more reactant PCl5 to products
(PCl3 and Cl2).
External Factors Affecting Equilibrium
Chemical equilibrium is affected (disturbed) by
changes in the following three factors:
1. Concentration of one of the components
2. Temperature
3. Pressure (for gaseous systems only)
Le Chatelier’s Principle
States that; when system in chemical
equilibrium is disturbed by change in
temperature, pressure or concentration, system
shift in equilibrium position in away that tends
counteract this change of a variable.
Concentration Changes
Let us look at the effect of removing the
products or adding reactants.
If the chemical reaction is at equilibrium and we
add a substrate (either a reactant or product)
the system will shift so as to re-establish
equilibrium by consuming part of the added
substance.
Conversely, removal of a substance will result in
a reaction moving in a direction that forms more
of a substance.
In summary;
• Add more reactants
• Remove reactants
shift to products
shift to reactants
• Add more products
shift to reactants
• Remove products
shift to products
Recall: Reaction Quotient for an equilibrium
system is calculated from the same expression
as the equilibrium constant, but the
concentrations are NOT at equilibrium.
e.g. N2O4(g)
2NO2(g)
[𝑁𝑂2 ]2
𝑄=
[𝑁2 𝑂4 ]
Change in concentrations are best understood in
terms of what would happen to ′′𝑄′′ if the
concentrations are changed.
Note:
𝑄 = K at equilibrium.
If 𝑄 < K then there too many reactants, the
reaction will shift in the forward direction
(products).
If 𝑄 > K then there too many products, the
reaction will shift to reactants.
Effect of Changing Temperature
Exothermic Reaction
In exothermic reaction consider heat as product
aA + bB
pP + qQ + heat
• Add heat
shift to reactants
• Remove heat
shift to products
Note: Heat is absorbed as the products are
converted to reactants, so the equilibrium shift
to the left, equilibrium constant decreases.
For an exothermic reaction (-∆H), the amount of
products are increasing at equilibrium by a
decrease in temperature.
Kc, equilibrium constant increase.
Endothermic Reactions
In endothermic reactions, consider heat as
reactant.
aA + bB + heat pP + qQ
• Add heat
shift to products
• Remove heat
shift to reactants
In an endothermic reaction heat is absorbed as the
reactants are converted to products, so an increase
in temperature (heat) causes the equilibrium shift
to the right (toward products), and K, increases.
For an endothermic reaction (+∆H), the amount of
products are increased at equilibrium by an
increase in temperature. Kc, is larger at higher T.
Effect of Changing Pressure (and Volume)
• Pressure changes only after equilibrium
systems with unequal moles of gaseous
reactants and products.
If pressure is increased by decreasing the
volume of the reaction mixture, the reaction
shift in the direction of fewer moles of gas.
Note: at constant temperature, reducing the
volume of a gaseous equilibrium mixture causes
causes the system to shift in the direction that
reduce the number of moles of gas.
Increase the volume causes the shift in the
direction that produces more moles of gas.
Example:(Increase in pressure)
N2(g) + 3H2(g) 2NH3(g)
If the pressure is increased (by reducing volume)
then the reaction shifts the product (ammonia).
Example (decrease in pressure)
PCl5(g)
PCl3(g) + Cl2(g)
If the pressure is decrease (by increasing the
volume) the reaction shifts to the products.
The Effect of the Catalyst
A catalyst increase the rate at which the
equilibrium is achieved, but it does not change
the composition of the equilibrium mixture.
Therefore, a catalyst has NO effect on a system
at equilibrium. It just gets the reaction faster to
the equilibrium.
Presence of Inert Substance
• An inert substance is the substance that is not
reactive with any species in the equilibrium
mixture.
• Therefore an inert substance will not affect
the equilibrium mixture.