unit 5 Flashcards
gas properties
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low density
changes dramatically with a change in pressure/temp
volume, pressure, temperature are interrelated
pressure
force exerted per unit area as gas molecules strike the surfaces around them
unit equalities
- mm Hg = 760. torr = 1 atm = 101.3 kPa = 14.7 psi
mm Hg (millimeters of mercury)/torr
atmospheric pressure that can support a column of mercury that is 760. mm high in a barometer
atmospheres (atm)
avg. pressure at sea level
pascal (Pa)
SI unit of pressure
also expressed as kPa
psi (lbs per square inch)
measure of force per unit area
boyle’s law
p1v1=p2v2
volume increases = pressure decreases
Charles law
v1/t1 = v2/t2
volume increases = temperature increases
always positive
avogadro’s law
v1/n1 = v2/n2
volume increases = moles increase
Gay-Lussac’s law
p1/t1 = p2/t2
pressure increases = temperature increases
combined gas law
p1v1/t1 = p2v2/t2
ideal gas law
PV = nRT
P: pressure in atm
V: volume in L
n: moles in mol
R: ideal gas constant (0.08206 when Latm/molK)
T: temperature in K
molar volume
volume occupied by 1 mol of a substance
standard temperature and pressure (STP)
for gases, molar volume is usually STP (1.00 atm)
density =
molar mass / molar volume
molar mass increases = density increases
1 mol of an ideal gas occupies:
22.4 L at STP
density of a gas equation
d = PM / RT
d: density in g/L
P: pressure in atm
M: molar mass in g/mol
R: ideal gas constant (0.08206 when Latm/molK)
T: temperature in K
to find molar mass of a gas:
measure mass and volume under known conditions of T and P
use ideal gas law to determine amount of gas in mol
divide mass by moles to get molar mass
partial pressure
pressure due to an individual component of a gas
Pn = Nn RT/V
Dalton’s Law of Partial Pressures
sum of partial pressures in a gas mixture = total pressure
P(total) = P(a) + P(b) + P(c)
mole fraction (X(a))
X(a) = N(a)/N(total)
or
P(a) = X(a)P(total)
gas mole fraction
percent by volume / 100
Vapor pressure
gas mixture with partial pressures
temp increases = Vapor pressure increases
kinetic molecular theory
simplest model for the behavior of gases
KMT postulates
-negligibly small
-avg. kinetic energy is proportional to Kelvin temp
-collision of one particle with another is elastic
pressure =
force / area
Boyle’s law: proportion
volume is inversely proportional to pressure
Charles law: proportion
volume is proportional to temperature
Avogadro’s law: proportion
volume is proportional to moles
particles of different masses have the (same/different) average kinetic energy
same
kinetic energy =
(1/2)mv^2
do lighter particles travel faster than heavier particles?
yes
temperature and velocity distribution proportion
temp increases = higher velocity
diffusion
gas molecules spread out due to a concentration gradient
-heavier molecules diffuse more slowly
effusion
gas escapes from a container into a vacuum through a small hole
-heavier molecules effuse more slowly
rate is inversely proportional to
1 / square root of the molar mass
effusion rate a / effusion rate b =
square root of molar mass b / square root of molar mass a
gases behave ideally when:
-volume of gas particles is small compared to the volume between them
-attractions between particles are insignificant
finite volume
actual size
when is finite volume important?
high pressure because then the particles occupy a significant portion of the total gas volume
when are intermolecular forces important?
low temps: collisions occur with less kinetic energy and weak attractions decrease collisions, also causes a decrease in pressure
Which postulate of the kinetic molecular theory breaks down under conditions of high pressure?
the volume of a gas particle is small when compared to the space between the particles
