NUTRITION IN SPORT

sis to be maintained. When this is not the case,
and the number of H+ions increases, the pH
of the blood (which is normally around 7.4)
decreases to lower levels (7.0 or lower) (Table
29.1). Muscle pH is normally at 7.0 and decreases
to 6.8 or lower (Robergs & Roberts 1997). The pH
of a given substance is the negative logarithm of
the hydrogen ion concentration (-log [H+]). Since
it is logarithmic, a unit increase of 1.0 means a
tenfold increase in the number of H+ions. Basic
solutions have few H+ions and acids have plenti-
ful amounts of H+ions. Distilled water is consid-
ered a neutral substance at a pH of 7.0 (25 °C).
The pH scale is shown in Fig. 29.1 and was
initially devised by the Danish chemist Soren
Sorensen. Body fluids differ in their pH level,
with gastric fluids being an acidic 1.0 and arterial
and venous blood being slightly basic at c. 7.45.
During normal activity, the blood and extracel-
lular fluid remain at a pH of approximately 7.4, a
slightly alkalotic state. When the number of H+
ions increases, such as during intense exercise,
the blood pH drops to below 7.0 (muscle pH is
even lower), and a state of acidosis exists. As
metabolism is highly H+ion sensitive, the regula-
tion of alkalosis and acidosis is extremely impor-
tant. Figure 29.2 shows the relationship between
pH and [H+] with the extreme physiological
realms.
The body has three basic mechanisms for
adjusting and regulating acid–base balance and
which minimize changes within the body. First,
there are the chemical buffers which adjust [H+]
within seconds. Secondly, pulmonary ventilation
excretes H+through the reaction

394 nutrition and exercise

H++HCO 3 - ́H 2 CO 3 ́H 2 O+CO 2

adjusting [H+] within minutes. Finally, the

kidneys excrete [H+] as fixed acid and work on a

long-term basis to maintain acid–base balance.

We are concerned with the bicarbonate buffer

system (Vick 1984).

The body’s chemical buffer, and more specifi-

cally the bicarbonate buffer, consists of a weak

acid (carbonic acid) and the salt of that acid

(sodium bicarbonate), often termed a conjugate

acid–base pair. Discussion of blood pH regula-

tion has generally focused on the role of bicar-

bonate, since it can accept a proton to form

carbonic acid in the following equation:

HCO 3 - +H+ ́H 2 CO 3

When metabolism produces an acid such as lactic

acid, which is much stronger than carbonic acid,

a proton is liberated, binds with bicarbonate and

Table 29.1Approximate relationship between [H+]
and pH.

[H+] (nmol · l-^1 )pH

20 7.7

30 7.5

40 7.4

50 7.3

60 7.2

70 7.15

H+

1 2 3 4 5 6 7 8 9

10

11

12

13

14

OH–

Gastric fluid

Lemon juice

Tomatoes

Coffee

Urine

Distilled water

Egg white

Milk of magnesia

Cleaning ammonia

Caustic soda

Acidic

Alkaline

Fig. 29.1The pH scale with some typical examples.