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Gases
Solid
Liquid
Gas
States of Matter
Definite Definite
Condensed Fluid
volume Shape
Phase
Definite No
Condensed Fluid
Volume definite
Phase
shape
No
No
Much
definite definite
lower
volume shape
density
Gases at room temperature – can be cooled and compressed into the liquid phase
Vapor – gas formed by evaporation or sublimation
Properties of Gases
1. Gases can be compressed into smaller volumes
↑ P causes ↓ V
therefore: ↑ P causes ↑density
2. Gases exert pressure on surroundings. So to confine a gas, pressure must be
exerted on it. (Newton’s third law: for every action, there is an equal and
opposite reaction)
3. Gases expand without limits, therefore a gas will completely fill the container it
occupies
4. Gases diffuse into one another. Gases mix completely, i.e. gases are miscible.
5. The amount of gas can be known if the pressure, temperature and volume are
known.
Pressure – force / unit area
Units: pounds / square inch (psi)
Newtons / m2 (Pascals)
Atmospheres (atm)
mm Hg
Torr
1 atm = 14.7 psi = 1.01 x 105 Pa = 760 mm Hg = 760 Torr = 33.8 ft water
To measure pressure
a. Barometer – the force of the atmosphere pushes the liquid up the tube.
1 atmosphere pushes a column of mercury 760 mm, water 33.8 feet
b. and c. Manometer – the pressure of a gas is known by comparing it with the
atmosphere
Patm at sea level > Patm at higher elevations
As your elevation increases, the mass of air above you decreases, thus less
pressure
Gas Laws
Boyle’s Law
↑ P causes ↓ V
V is proportional to 1/P
PV = constant so ………………..P1V1 = P2V2
Charles’ Law
↑ T causes ↑ V
This is why a balloon in a hot car pops
By extending this idea, ↓ T causes ↓ V
All volumes approach zero at the same temperature – Absolute Zero
Zero temperature corresponds to zero motion and zero volume
Absolute Zero = -273.15º C = 0 K
Kelvin Temperature = Celsius Temperature + 273.15
Mathematically: T is proportional to V, but you MUST use the Kelvin scale
V1 = V2
T1
T2
Combined Gas Law
P1V1 = P2V2
T1
T2
If temperature is constant: P1V1 = P2V2 (Boyle’s Law)
If pressure is constant: V1 / T1 = V2 / T2 (Charles Law)
If volume is constant: P1 / T1 = P2 / T2
Standard Temperature and Pressure (STP)
T = 0º C = 273.15 K
P = 1 atm = 760 mm Hg = 101.15 kPa
Avogadro’s Law, 1811
At the same temperature and pressure, equal volumes of all gases contain the
same number of molecules (or moles, n)
Mathematically Stated
↑ n causes ↑ V
If you put more gas into a container, it gets bigger
Therefore: n is directly proportional to V
V1 / n1 = V2 / n2
Ideal Gas Law
Combining all gas laws
P1V1 / n1T1 = P2V2 / n2T2 = constant = R = 0.08206 (lit atm) / (mol K)
PV = nRT
You must be careful of units when using the ideal gas law!
Standard Molar Volume
For all ideal gases, at STP
If n = 1 mole, then V = 22.4 L
Kinetic Theory of Gases
Defines the behavior and assumptions of an ideal gas

The particles are small, moving in constant, random, straight-line motion.
Newton’s First Law of Motion – Any object with mass (inertia) will move in a
straight line at a constant speed unless acted on by a force.

Gases are mostly empty space; the particles are separated by large distances. The
particles are considered “point masses” with no appreciable volume. The particles
themselves occupy no volume; the volume of the gas is the result of the motion of
particles and collisions with the container.

Particles collide in perfectly elastic collisions. The collisions are with each other
and the walls of the container.
Elastic collisions – no net loss of energy in the collision, like billiard balls

There are no forces between the particles

Temperature is a measure of the average kinetic energy of the particles.
Substances with higher molecular weight will move with a lower average velocity
at the same temperature.
KE = ½ mv2
Deviations from the ideal gas behavior




Under high pressure, the particles are squished close together and the volume of
the particles themselves affects the total volume. This results in a measured
volume which is higher than expected
At low temperatures, the particles are moving slowly, and the attractions between
the particles become significant. This gives a pressure that is smaller than
expected, because the particles are slowing down as they are drawn to each other.
Polar molecules exert attractive forces on each other – therefore polar gases will
deviate from ideal behavior
Close to the condensation point (boiling point), molecules exert attractive forces
on each other – therefore the gas will deviate from ideal behavior
Diffusion and Effusion



Effusion – the escape of a gas through a small opening due to random motion of
particles
Diffusion – mixing of gases with each other due to random motion of particles
Both effusion and diffusion will happen faster for a lighter molecule which will
have a larger velocity at the same temperature