5 Steps to a 5 AP Chemistry

acids and bases may be strong or weak. Strong acidscompletely dissociate in water; weak

acids only partially dissociate. For example, consider two acids HCl (strong) and

CH 3 COOH (weak). If each is added to water to form aqueous solutions the following reac-

tions take place:

The first reaction essentially goes to completion—there is no HCl left in solution. The

second reaction is an equilibrium reaction—there are appreciable amounts of both reactants

and products left in solution.

There are generally only two strong bases to consider: the hydroxide and the oxide ion

(OH−and O^2 −, respectively). All other common bases are weak. Weak bases, like weak

acids, also establish an equilibrium system, as in aqueous solutions of ammonia:

In the Brønsted–Lowry acid–base theory, there is competition for an H+. Consider the

acid–base reaction between acetic acid, a weak acid, and ammonia, a weak base:

Acetic acid donates a proton to ammonia in the forward (left-to-right) reaction of the

equilibrium to form the acetate and ammonium ions. But in the reverse (right-to-left) reac-

tion, the ammonium ion donates a proton to the acetate ion to form ammonia and acetic

acid. The ammonium ion is acting as an acid, and the acetate ion as a base. Under the

Brønsted–Lowry system, acetic acid (CH 3 COOH) and the acetate ion (CH 3 COO−) are

called a conjugate acid–base pair. Conjugate acid–base pairsdiffer by only a single H+.

Ammonia (NH 3 ) and the ammonium ion (NH 4 +) are also a conjugate acid–base pair. In

this reaction there is a competition for the H+between acetic acid and the ammonium ion.

To predict on which side the equilibrium will lie, this general rule applies: The equilibrium

will favor the side in which the weaker acid and base are present. Figure 15.1 shows the rela-

tive strengths of the conjugate acid–base pairs.

In Figure 15.1 you can see that acetic acid is a stronger acid than the ammonium ion and

ammonia is a stronger base than the acetate ion. Therefore, the equilibrium will lie to the right.

The reasoning above allows us to find good qualitative answers, but in order to be able

to do quantitative problems (how much is present, etc.), the extent of the dissociation of

the weak acids and bases must be known. That is where a modification of the equilibrium

constant is useful.

Ka—The Acid Dissociation Constant

Strong acids completely dissociate (ionize) in water. Weak acids partially dissociate and

establish an equilibrium system. But as shown in Figure 15.1 there is a large range of weak

acids based upon their ability to donate protons. Consider the general weak acid, HA, and

its reaction when placed in water:

An equilibrium constant expression can be written for this system:

Kc^3

HO A+

HA]

=

[][]−

[

HA(aq) + H O(1) 23 H O aq) + A aq)+((−

CH COOH(aq) + NH aq 333 ()CH COO aq−+()+NH aq 4 ()

NH aq 32 ()+H O(1)OH aq + NH aq)−+() 4 (

HCl(aq) + H O(1) H O aq Cl aq)

CH COOH(aq)

23

3

→++−() (

++ H O(1) 23 H O aq+−()+CH COO aq 3 ()

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