chemical equilibrium 2. ionic equilibrium acid & base ionization for weak acids like acetic acid...

CHEMICAL EQUILIBRIUM 2

Ionic Equilibrium

Acid & Base Ionization

• For weak acids like acetic acid there will be an equilibrium according to its ionization in water:

• The equilibrium constant known as acid constant is given by:

OHCOOCHOHCOOHCH 3323

][

]][[

3

33

COOHCH

OHCOOCHK a

• From this equilibrium constant the Henderson-Hasselbalch equation is derived:

which is applied in buffer / pH calculations.• In the same way the base equilibrium constant is:

• The form of Henderson-Hasselbalch equation will then be:

][

]][[

B

BHOHKb

][

][log

base

saltpKpOH b

][

][log

acid

saltpKpH a

Solubility Product

• The solubility of a salt like AgCl

will be having the solubility product

• The solubility product can be written as:

)()()( aqClaqAgsAgCl

][

]][[

AgCl

ClAgK sp

]][[ ClAgK sp

Stability or Formation Constant

• When complexes are formed their formation constant indicates the stability of the complex.

• The Kf or Kstab may be written

2433

2 )()(4)( NHCuaqNHaqCu

43

2

243

]][[

])([

NHCu

NHCuK f

Distribution Coefficient

• In the fields of organic and medicinal chemistry, a partition (P) or distribution coefficient (D) is the ratio of concentration of a compound in the two phases of a mixture of two immiscible solvents at equilibrium. Hence these coefficients are a measure of differential solubility of the compound between these two solvents.

• Normally one of the solvents chosen is water while the second is hydrophobic such as octanol. Hence both the partition and distribution coefficient are measures of how hydrophilic ("water loving") or hydrophobic ("water fearing") a chemical substance is. Partition coefficients are useful for example in estimating distribution of drugs within the body.

• Hydrophobic drugs with high partition coefficients are preferentially distributed to hydrophobic compartments such as lipid bilayers of cells while hydrophilic drugs (low partition coefficients) preferentially are found in hydrophilic compartments such as blood serum

• The distribution coefficient can be written as:

• Hence the value indicates the solubility is s specific phase.

aqu

org

compound

compoundD

][

][

Factors Affecting the Equilibrium

Le Châtelier’s Principle

““If you disturb an If you disturb an equilibrium, it equilibrium, it will shift to undo will shift to undo the disturbance”.the disturbance”.

Change of concentration

• In the equation:

• 2NO(g) + O2(g) <--> 2NO2(g)

• If you add more NO(g) the equilibrium shifts to the right producing more NO2(g)

• If you add more O2(g) the equilibrium shifts to the right producing more NO2(g)

• If you add more NO2(g) the equilibrium shifts to the left producing more NO(g) and O2(g)

• Consider the Haber process

• If H2 is added while the system is at equilibrium, the system must respond to counteract the added H2 (by Le Châtelier).

• That is, the system must consume the H2 and produce products until a new equilibrium is established.

• Therefore, [H2] and [N2] will decrease and [NH3] increases

N2(g) + 3H2(g) 2NH3(g)

Change in Reactant or Product ConcentrationsChange in Reactant or Product Concentrations

Common ion effect

• The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance.

• The presence of a common ion suppresses the ionization of a weak acid or a weak base.

• Consider mixture of CH3COONa (strong electrolyte) and CH3COOH (weak acid).– CH3COONa (s) ↔ Na+ (aq) + CH3COO- (aq)– CH3COOH (aq) ↔ H+ (aq) + CH3COO- (aq)

• CH3COO- common ion

• What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK?

– HCOOH (aq) ↔H+ (aq) + HCOO- (aq)• Initial (M) 0.30 0.00 0.52 • Change (M) - x +x +x• Equilibrium (M) 0.30-x x 0.52 + x

• 0.30 – x ≈ 0.30• 0.52 + x ≈ 0.52

][

][log

HCOOH

HCOOpKpH a

1.4]30.0[

]52.0[log77.3 pH

Pressure/volume change

• In the equation • 2SO2(g) + O2(g) <--> 2SO3(g),

• an increase in pressure will cause the reaction to shift in the direction that reduces pressure, that is the side with the fewer number of gas molecules. Therefore an increase in pressure will cause a shift to the right, producing more product. (A decrease in volume is one way of increasing pressure.)

• Consider the production of ammonia

• As the pressure increases, the amount of ammonia present at equilibrium increases.

N2(g) + 3H2(g) 2NH3(g)

Temperature Change

• The temperature equilibrium constant is governed by van’t Hoff equation:

• Where K is the equilibrium constant, T is temperature and ∆H is the enthalpy.

2

ln

RT

H

dT

Kd

Effect of Temperature ChangesEffect of Temperature Changes

• The equilibrium constant is temperature dependent.

• For an endothermic reaction, H > 0 and heat can be considered as a reactant.

• For an exothermic reaction, H < 0 and heat can be considered as a product.

• Adding heat (i.e. heating the vessel) favors away from the increase:– if H > 0, adding heat favors the forward reaction,– if H < 0, adding heat favors the reverse reaction.

Effect of Temperature ChangesEffect of Temperature Changes

• Removing heat (i.e. cooling the vessel), favors towards the decrease:

– if H > 0, cooling favors the reverse reaction,

– if H < 0, cooling favors the forward reaction.

Using a catalyst

• A catalyst increases the speed in which a reaction takes place, however it never has any effect on the equilibrium.