Acid molecules dissociate anion and hydrogen ion

General Chemistry II
Chapter 14 Lecture Notes
Chemical Equilibrium

THE CONCEPT OF EQUILIBRIUM

HA ↔︎ H+1 + A-1

A value called the equilibrium constant can be assigned to the chemical system by measuring, at equilibrium, the concentrations of all the species.

In general, all chemical equilibrium systems obey the law of mass action. For the general equilibrium equation

aA + bB ↔︎ cC + dD

As an example of the use of the exponents, note that if A is the same species as B and if D is not

formed, in other words	1A + 1A ↔︎ cC + 0D or 2A ↔︎ [C]c

HOMOGENEOUS EQUILIBRIA

All reactants and products are in the same liquid or gas phase.

Kp = PB/PA

In general Kc ≠ Kp unless the number of moles of gas does not change during the reaction.

Ex: Write Kp andKc for the reaction of carbon dioxide gas with NaOH solution

eliminated from the equilibrium expression.

NH3(g) + HCl(g) ↔︎ NH4Cl(s)

Keq = 1/[NH3]·[HCl] if the molar concentrations of the gases will be expressed, or

Kp = 1/PNH3 · PHCl if the gas pressures will be expressed.

For A → B Keq = [A]/[B] and

For B → C, Keq = [B]/[C]

H2PO4-1 ↔︎ HPO4-2 + H+1 K2 = etc.

HPO4-2 ↔︎ PO4-3 + H+1 K3 = etc.

For A ↔︎ B + C, Kf = [B][C]/[A]
For B + C ↔︎ A, Kr = [A]/[B][C]
Therefore Kf x Kr = [A][B][C]/[A][B][C] = 1

Remember that the reactant and product concentrations in the Keq expression are raised to the power equal to their coefficients. What happens if the equation is written in more than one way?

Both Keq values will give the same reactant and product concentrations when applied to the corresponding equations.

RELATION BETWEEN CHEMICAL KINETICS AND CHEMICAL EQUILIBRIUM

If ratef = rater, then kf [A] = kr[B]

kf /kr = [A]/[B] which is equal to the equilibrium constant, Kf. A large forward rate constant leads to a high rate of formation of B and a large value of Kf.
Similarly, kr / kf = [B]/[A] = Kr.

Ex: The Ka for the dissociation of acetic acid is 1.8 x 10-5. HA(aq) ↔︎ A-1(aq) + H+1(aq)
If a system contains [A-1] = 2.0 x 10-4M, [H+1] = 1.5 x 10-5 M and [HA] = 0.55 M, determine whether the system is in equilibrium or not and in what direction it must shift to put it into equilibrium. Qc = 5.5 x 10-9 (compared to Ka = 1.8 x 10-5) so the system must shift to the right (more product) to attain equilibrium.

CALCULATING EQUILIBRIUM CONCENTRATIONS

Ex: For the dissociation of any weak acid, HA(aq) ↔︎ H+1(aq) + A-1(aq)

K = [H+1] · [A-1]/ [HA] Suppose K = 4.2 x 10-7 and imagine a system where the initial HA concentration is 0.100M. What are the equilibrium concentrations of HA, H+1 and A-1?

Other calculations may be more difficult and require use of the quadratic formula.

Ex: Ca+2 + EDTA-2 ↔︎ CaEDTA K = 4.80
Suppose initial [Ca+2] = 0.1000 M and initial [EDTA-2] = 0.0500M, what are the equilibrium concentrations of Ca+2, EDTA-2 and CaEDTA?

[CaEDTA] = 0.0032M, [Ca+2] = 0.0068M and [EDTA-2] = 0.0468M

FACTORS THAT AFFECT CHEMICAL EQUILIBRIUM

The effects of concentration changes can be determined qualitatively by inspection of the equilibrium expression or by calculating Qc values for initial reactant and product concentration in the equilibrium expression.

CHANGES IN VOLUME AND PRESSURE

The CH3COOH(aq) + NaOH(aq) → CH3COONa(aq) + H2O(l) equilibrium is unaffected by pressure.

Heat + A ↔︎ B (endothermic)

Increasing temperature causes the equilibrium to shift to the right since heat is needed to force the reaction.