164
Gases
and conclude that, at a given volume and temperature, the fraction (or percentage) of the total pressure exerted by each component is just equal to its fraction
(or percentage) of the total moles present. Thus, to draw on a previous problem,
if we know that 21.0% of the moles of air is O 2 , we know that 21.0% of the air
pressure is due to O 2 . If the barometric pressure is 740 torr, then the partial
pressure of O 2 is
P 02 = (0.210)(740 torr) = 155 torr
A corollary of this observation is Dalian's law of partial pressures, that the
total pressure (P) of a gas mixture is equal to the sum of the partial pressures of
the components, i.e.,
P= P A + P B + P c + • • •
(11-4)
In the laboratory work of general chemistry, we have important applications
of partial pressures. Gases (such as oxygen) that are not very soluble in water
are collected in bottles by displacement of water. As the gas bubbles rise
through the water, tney become saturated with vapor, and the collected gas is a
mixture of water vapor and the original gas. When the bottle is filled, it is at
atmospheric pressure, or
^"gas + ^H20 = barometric pressure = P B
To obtain the partial pressure of the gas (P sa& ), we must subtract the watervapor pressure from the barometric pressure
•fgas
= PB ~ PtizO
Fortunately, water-vapor pressures are known accurately over the entire liquid
range of water (see Table 11-1) and do not have to be determined experimentally each time.
When gases are collected over mercury, no correction is needed for the vapor
pressure of Hg because it is so small (about 2 x 10~
3 torr at room temperature).
Collection over Hg has the further advantage that gases are insoluble in it.
Disadvantages are its high cost and the toxicity of its vapor.
PROBLEM:
A 250 ml flask is filled with oxygen, collected over water at a barometric pressure
of 730 torr and a temperature of 25°C. What will be the volume of the oxygen
sample, dry, at standard conditions?
SOLUTION:
Changing the temperature and pressure of a gas sample changes only the volume,
not the number of moles (that is, n remains constant). As a consequence, the ideal
Gases
and conclude that, at a given volume and temperature, the fraction (or percentage) of the total pressure exerted by each component is just equal to its fraction
(or percentage) of the total moles present. Thus, to draw on a previous problem,
if we know that 21.0% of the moles of air is O 2 , we know that 21.0% of the air
pressure is due to O 2 . If the barometric pressure is 740 torr, then the partial
pressure of O 2 is
P 02 = (0.210)(740 torr) = 155 torr
A corollary of this observation is Dalian's law of partial pressures, that the
total pressure (P) of a gas mixture is equal to the sum of the partial pressures of
the components, i.e.,
P= P A + P B + P c + • • •
(11-4)
In the laboratory work of general chemistry, we have important applications
of partial pressures. Gases (such as oxygen) that are not very soluble in water
are collected in bottles by displacement of water. As the gas bubbles rise
through the water, tney become saturated with vapor, and the collected gas is a
mixture of water vapor and the original gas. When the bottle is filled, it is at
atmospheric pressure, or
^"gas + ^H20 = barometric pressure = P B
To obtain the partial pressure of the gas (P sa& ), we must subtract the watervapor pressure from the barometric pressure
•fgas
= PB ~ PtizO
Fortunately, water-vapor pressures are known accurately over the entire liquid
range of water (see Table 11-1) and do not have to be determined experimentally each time.
When gases are collected over mercury, no correction is needed for the vapor
pressure of Hg because it is so small (about 2 x 10~
3 torr at room temperature).
Collection over Hg has the further advantage that gases are insoluble in it.
Disadvantages are its high cost and the toxicity of its vapor.
PROBLEM:
A 250 ml flask is filled with oxygen, collected over water at a barometric pressure
of 730 torr and a temperature of 25°C. What will be the volume of the oxygen
sample, dry, at standard conditions?
SOLUTION:
Changing the temperature and pressure of a gas sample changes only the volume,
not the number of moles (that is, n remains constant). As a consequence, the ideal