The Ionic Equilibria, Slides of Physical Chemistry

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Ionic

Equilibria

Mahfuza Afroz Soma

Lecturer

UAP

Modern theories of acids, bases and salts

• There are three concepts of acids and

bases in current use.

I. Arrhenius concept II. Bronsted-Lowry concept III.Lewis concept

Limitations of Arrhenius Concept

1. Free H

+

and OH

-

ions do not exist in water: The

H

+

and OH

-

ions produced by acids and bases

respectively do not exist in water in the free

state. They are associated with water molecules

to form complex ions through hydrogen bonding.

Thus the H

ion forms a hydronium ion:

• Similarly, OH

ion forms the complex H

3

O

2 −

Limitations of Arrhenius Concept

2. Limited to water only: Arrhenius defined acids and bases as compounds producing H + and OH - ions in water only. But a truly general concept of acids and bases should be appropriate to other solvents as well. 3. Some bases do not contain OH - : Arrhenius base is one that produces OH - ions in water. Yet there are compounds like ammonia (NH 3 ) and calcium oxide (CaO) that are bases but contain no OH - ions in their original formulation. Arrhenius models of acids and bases, no doubt, proved very helpful in interpreting their action. However on account of its limitations the Arrhenius concept needed to be modified.

Examples of Bronsted acids and bases

1. HCl gas and H

2

O: When dry HCl gas dissolves in

water, each HCl molecule donates a proton to a

water molecule to produce hydronium ion.

Thus HCl gas is a Bronsted acid and water that

accepts a proton is a Bronsted base.

2. HCl and Ammonia (NH

3

): HCl gas reacts with

ammonia (NH

3

) to form solid NH

4

Cl.

HCl is a proton donor and hence a Bronsted

acid, while NH

3

is a proton acceptor and a

Bronsted base.

Examples of Bronsted acids and bases

Conjugate Acid-Base pairs

• In an acid-base reaction, the acid (HA) gives up

its proton (H

) and produces a new base (A

The new base that is related to the original acid

is called a conjugate base. Similarly the original

base (B

) after accepting a proton (H

) gives a

new acid (HB) which is called a conjugate acid. A

hypothetical reaction between the acid HA and

the base B

will illustrate the above definitions.

• The acid (HA) and the conjugate base (A -

) that

are related to each other by donating and

accepting a single proton, are said to

constitute a conjugate Acid-Base pair.

• Now let us consider the reaction between acetic acid

and water to form the conjugate base CH

3

COO

and

the conjugate acid H

3

O

• We know that acetic acid is less than 1 % ionised in

water. Since the equilibrium is displaced toward the

left, we can say that : ( i) CH

3

COO

-

is a stronger base

than H

2

O; and (ii) H

3

O

+

is a stronger acid than

CH

3

COOH. Thus we can conclude that :

• ( a) a weak base has strong conjugate acid
• ( b) a weak acid has a strong conjugate base

Amphiprotic substances

• Molecules or ions that can behave both as

Bronsted acid and base are called amphiprotic

substances. For example, with HCl, water acts

as a base in accepting a proton from the acid.

• If the Lewis acid be denoted by A and the Lewis

base by B, then the fundamental equation of the

Lewis theory can be written as:

• It may be noted that :
  1. all cations or molecules short of an electron- pair act as Lewis acids; and
  2. all anions or molecules having a lone electron- pair act as Lewis bases.

Examples of Lewis reactions

1. Between H + and NH 3 : Proton (H + ) is a Lewis acid as it can accept an electron-pair. Ammonia molecule (:NH 3 ) has an electron-pair which it can donate and is a Lewis base. Thus the Lewis reaction between H+ and NH 3 can be written as:

RELATIVE STRENGTH OF ACIDS

• The strength of an acid depends on its ability to transfer its proton (H + ) to a base to form its conjugate base. When a monoprotic acid (HA) dissolves in water, it transfers its proton to water (a Bronsted base) to form hydronium ion (H 3

O

) and a conjugate base. …………….. ( 1 )

• For simplifying our discussion, we take-
• Thus we can write the equilibrium reaction ( 1 ) as- ..…………….. ( 2 )
• This equation represents the dissociation of the acid HA into H + ion and A - ion. Applying the Law of Mass action to the acid dissociation equilibrium, we can write- ..…………... (3)

• where K a is called the acid dissociation constant. In

dilute solution of the acid (HA) we note that the

concentration of liquid water remains essentially

constant.

• The strength of an acid is defined as the

concentration of H

+

ions in its aqueous solution at a

given temperature.

• From the equilibrium ( 3 ), it is evident that the

concentration of H

ions, [H

], depends on the value of

K

a

. Therefore, the value of K

a

for a particular acid is a

measure of its acid strength or acidity.