Human Physiology, 14th edition (2016)

Respiratory Physiology 547

Calculation of P O

2

With increasing altitude, the total atmospheric pressure and the

partial pressure of the constituent gases decrease ( table  16.5 ).

At Denver (5,000 feet above sea level), for example, the atmo-

spheric pressure is decreased to 619 mmHg and the P^ O 2 i s

therefore reduced to 619  3  0.21  5  130 mmHg. At the peak of

Mount Everest (at 29,000 feet) the P^ O 2 is only 42 mmHg. As one

descends below sea level, as in scuba diving, the total pressure

increases by one atmosphere for every 33 feet. At 33 feet there-

fore, the pressure equals 2  3 760  5 1,520  mmHg. At 66 feet

the pressure equals three atmospheres.

Inspired air contains variable amounts of moisture. By the

time the air has passed into the respiratory zone of the lungs,

however, it is normally saturated with water vapor (has a rela-

tive humidity of 100%). The capacity of air to contain water

vapor depends on its temperature; since the temperature of the

respiratory zone is constant at 37 8 C, its water vapor pressure is

also constant (at 47 mmHg).

Water vapor, like the other constituent gases, contributes

a partial pressure to the total atmospheric pressure. The total

atmospheric pressure is constant (depending only on the height

of the air mass), so the water vapor “dilutes” the contribution

of other gases to the total pressure:

P^ wet atmosphere 5 P^ N 2 1 PO^2   1 P^ CO 2 1 P^ H 2 O

When the effect of water vapor pressure is considered, the

partial pressure of oxygen in the inspired air is decreased at sea

level to

P^ O 2 (sea level) 5 0.21 (760 2 47) 5 150 mmHg.

As a result of gas exchange in the alveoli, there is an

increase in the P^ CO 2 , while the P^ O 2 of alveolar air is further

diminished to about 105 mmHg. The partial pressures of the

inspired air and the partial pressures of alveolar air are com-

pared in figure 16.19.

16.4 GAS EXCHANGE

IN THE LUNGS

Gas exchange between the alveolar air and the pulmonary

capillaries results in an increased oxygen concentration

and a decreased carbon dioxide concentration in the blood

leaving the lungs. This blood enters the systemic arteries,

where blood gas measurements are taken.

LEARNING OUTCOMES

After studying this section, you should be able to:

7. Explain how partial gas pressures are calculated,
   and their significance in measurements of arterial
   blood gases.


8. Describe the factors that influence the partial
   pressure of blood gases and the total content of
   gases in the blood.

The atmosphere is an ocean of gas that exerts pressure on all
objects within it. This pressure can be measured with a glass
U-tube filled with fluid. One end of the U-tube is exposed
to the atmosphere, while the other side is continuous with a
sealed vacuum tube. Because the atmosphere presses on the
open-ended side, but not on the side connected to the vacuum
tube, atmospheric pressure pushes fluid in the U-tube up on the
vacuum side to a height determined by the atmospheric pres-
sure and the density of the fluid. Water, for example, will be
pushed up to a height of 33.9 feet (10,332 mm) at sea level,
whereas mercury (Hg)—which is more dense—will be raised
to a height of 760 mm. As a matter of convenience, therefore,
devices used to measure atmospheric pressure (barometers)
use mercury rather than water. The atmospheric pressure at sea
level is thus said to be equal to 760 mmHg (or 760 torr, after
Evangelista Torricelli, who invented the mercury baro meter in
1643), which is also described as a pressure of one atmosphere
( fig. 16.18 ).
According to Dalton’s law, the total pressure of a gas mix-
ture (such as air) is equal to the sum of the pressures that each
gas in the mixture would exert independently. The pressure
that a particular gas in a mixture exerts independently is the
partial pressure of that gas, which is equal to the product of
the total pressure and the fraction of that gas in the mixture.
Dalton’s law can be restated as follows: The total pressure
of the gas mixture is equal to the sum of the partial pressures of
the constituent gases. Because oxygen constitutes about 21%
of the atmosphere, for example, its partial pressure (abbrevi-
ated P^ O 2 ) is 21% of 760, or about 159 mmHg. Nitrogen con-
stitutes about 78% of the atmosphere, so its partial pressure is
equal to 0.78  3  760  5  593 mmHg.
These two gases thus contribute about 99% of the total
pressure of 760 mmHg:

P^ dry atmosphere 5 PN^2   1 P^ O 2 1 P^ CO 2 5 760 mmHg

Figure 16.18 The measurement of atmospheric

pressure. Atmospheric pressure at sea level can push a column

of mercury to a height of 760 millimeters. This is also described

as 760 torr, or one atmospheric pressure.

Vacuum

Sea level

Atmospheric

pressure

760 mm

Hg