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Pressure
Volume
Moles
Temperature
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Common units
Atmospheres (atm)
 mm Hg
 Torr
 kiloPascals (kPa)
 Pounds per square inch (psi)
 Bar or millibar (b or mb)
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Influences
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Weight of air or water above you
Force caused by gas particle collisions with the
container wall
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1 atm
760 torr
760 mmHg
101.325 kPa
14.69 psi
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Typically in Liters
mL = cc or cm3
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Gas laws are independent of mass, but are
dependent on moles, which represents number
of particles
Exception is Graham’s Law
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Must use Kelvins
K = oC + 273.16, but we can get away with
saying 273
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PV/nT = PV/nT
Cross out any factors that are either constant or
unmentioned
Use this only if a factor is changing
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Boyles Law PV
Charles’ Law VT
Avogadro’s Law Vn
Gay-Lussac: PT
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PV=nRT
Use when you have all factors but one and the
system is not changing
R = ideal gas constant
Values for R: depends upon your unit of pressure
0.08206 L atm/mol K
 62.36 L mmHg/mol K
 62.36 L torr / mol K
 8.314 L kPa/mol K
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Pressure Tot = S partial pressure
Mole fraction = pressure fraction
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Pressure Tot = P coll gas + P water vapor
Total Pressure = atmospheric pressure
Water vapor pressure can be looked up in a
table by temperature
Pressure of dry gas = Atmospheric pressure –
water vapor pressure
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Have a mixture of 2 mol He, 4 mol Kr, and 4
mol Ar. If the total pressure is 760 mmHg,
what is the partial pressure of Ar?
Mole fraction of Ar = 4/(2+4+4) or .4
Partial pressure of Ar = 760 x .4
Partial pressure of Ar = 304 mmHg
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Two gases in a mixture are at the same
temperature
Same temp means same avg kinetic energy
½ mv2 = ½ mv2
The lower mass will have the higher velocity
A lower mass gas will diffuse or effuse faster
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May have to use PV=nRT to find moles of
known or unknown
Be careful with your phase symbols. ALL
gases must be considered