4
chapter 1
Water, Acids, Bases, and Buffers
Thus, water functions both as an acid (donor of H+ or
proton) and as a base (acceptor of H+ or proton). This de-
scription of an acid and a base follows from the Brônsted-
Lowry theory. According to the Lewis theory, acids are
electron pair acceptors and bases are electron pair donors.
The equilibrium constant,
K
, for the dissociation reaction
in Equation (1.1) is
[H+][OH~]
[H20]
(
1
.
2
)
where the square brackets refer to the molar concentrations
of the ions involved.
K
can be determined by measurement
of the electrical conductivity of pure water, which has the
value of 1
. 8
x 10
" 16
M at 25°C, indicative of a very small
ion concentration, where M (molar) is the units of moles
per liter. Therefore, the concentration of undissociated wa-
ter is essentially unchanged by the dissociation reaction.
Since 1 L of water weighs 1000 g and 1
mol of water
weighs 18 g, the molar concentration of pure water is
55.5 M. Substitution for
K
and [H
2
0] in Equation (1.2)
yields
[Н+ПОН-] = (55.5 M) x (1.8 x 10
“ ' 6
M)
[H+][OH~] = 1.0 x 10
~ 14
M
2
=
Kw
KK
is known as the
ion product of water.
In pure water,
[H+] and [OH
.] are equal, so that
[OH- ] = [H+] = 1.0 x 10
~ 7
M.
pH is employed to express these ion concentrations in a
convenient form, where the “p” of pH symbolizes “neg-
ative logarithm (to the base
1 0
)” of the concentration in
question. Thus,
PH = - lo g l()[H+] = log —
Similarly,
pOH = - lo g l
0
[OH-] = l o g ^ i r
Therefore, for water,
log[H+] + log[OH“] = log 10
“ 14
or
pH + pOH = 14.
The pH value of 7 for pure water at 25°C is considered to
be neutral, and values below 7 are considered acidic and
above 7 basic. Table 1-2 illustrates the pH scale extend-
ing from —1 to +15. It is important to recognize that as
the pH decreases, [H+] increases. A decrease in one pH
unit reflects a 10-fold increase in H+ concentration. In
TABLE 1-2
The pH Scale
[H+],M
pH
[OH-],M
10.0
-1
1
1 0 “ 15
1 . 0
0
10-14
0 . 1
1
10-13
0.0Щ0-2)
2
10-12
10-3
3
1 0 ” 11
1 0 “4
4
Acidic
10-10
10-5
5
10-9
10“6
6
10-8
W
1П—
7
iu
^
/
імешхаї
^ 1U
10“8
8
It
10-6
10-9
9
10-5
10-10
10
10-4
10-11
11
Basic
10-3
10-12
12
0.01
10"13
13
0 . 1
10-14
14
1
10-15
15
1
10
discussions of acid-base problems in human biochem-
istry, it is often preferable to express H+ concentration
as nanomoles per liter (nmol/L).
1.2 Buffers
Buffers resist change in pH in solutions when acids or
bases are added. They are either a mixture of a weak acid
(HA) and its conjugate base (A~) or a mixture of a weak
base (B) and its conjugate acid (HB+).
EXAMPLE 1
Acetic acid (CH
3
COOH) and car-
bonic acid (H
2
CC>
3
) are weak acids. Ammonia
(NH
3
) is a weak base. CH
3
COOH/CH
3
COO ,
H
2
C
0 3
/HC
0
j", and NH
3
/NH
4
constitute buffer
systems.
A buffer solution functions in the following manner
to resist changes in acidity or alkalinity. In an acetic
acid/sodium acetate buffer system, the species present
in solution are CH
3
COOH, CH
3
COO- , Na+, and H
2
0.
Amounts of H+ and OH
are initially assumed to be small.
When acid is added to the buffer, almost all of the
H+ ions react with acetate ions to produce weakly ion-
ized acetic acid (H+ + CH3COCT
CH
3
COOH). The
H+ ions are thereby prevented from appreciably changing
the pH.