Transcript Chapter 8 Gases

Gases
Kinetic Molecular Theory of Gases
A gas consists of small particles
(atoms/molecules) that move
randomly with rapid velocities
Further Information
They move faster when heated.
The attractive forces
between particles of a
gas can be neglected
Do you think this is accurate?
Why would this be important for
calculations?
The actual volume
occupied by a gas
molecule is extremely
small compared to the
volume that gas
occupies.
Is this true in the real world?
Why would this be helpful with
calculations?
The average kinetic
energy of a gas
molecule is
proportional to Kelvin
temperature
What is kinetic energy?
Why Kelvin temperature and not Celsius
or Fahrenheit?
What does proportional mean?
Gas particles are in
constant motion,
moving rapidly in
straight paths. *
Is this true?
What do we know about their motion?
Why would the real situation make the
calculations more difficult?
Ideal Gases
An imaginary gas that perfectly fits all
the assumptions of the kinetic molecular
theory (KMT).
Expansion
Gases do not have definite shape or
volume.
The expand to any container they are
enclosed in.
A gas in a 1 L container is then put into
a 2 L container. How much volume does
it have now?
Fluidity
In an ideal gas, the gas particles glide
past each other.
This feature allows gases to be referred
to as fluids just like liquids.
Low Density
Density of a gas substance is only
about 1/1000 of the same substance in
liquid or solid state.
Why is this true?
Compressibility
This is a crowding effect of gases when
the volume is decreased
Diffusion
Spontaneous (does not require energy)
mixing of particles of two substances
caused by their random motion
Properties of a Gas
Units of Measure
Pressure
Pressure is not the same as force.
Pressure is a force over an area.
Example: psi = Pounds per in2
Measuring Pressure
A barometer measures atmospheric
pressure.
Units of Pressure
kPa, atm, mm of Hg, torr
Helpful Conversions
1 atm = 760 mm Hg
1 atm = 760 torr
1 mm Hg = 1 torr
1 atm = 101.325 kPa
Volume
L, mL or cm3
Helpful conversions
1000 mL = 1 L
 1 mL = 1 cm3

Temperature
0C
or K
Helpful conversions:
0C = K – 273
K = 0C + 273
moles
Number of moles = n
If you are given grams, how would you
convert to moles?
Standard Temperature
and Pressure (STP)
Standard Temperature is 00C or 273 K
Standard Pressure is 101.3 kPa or 1
atm
Boyle’s Law:
Pressure and volume are inversely
proportional
P1V1 = P2V2
Charles’ Law:
Temperature and Volume are directly
proportional
V1 / T1 = V2 / T2
Gay-Lussac’s Law:
Pressure and Temperature are directly
proportional
P1/T1 = P2/T2
Combined Gas Law
P1V1 = P2V2
T1
T2
If you remember this law, hold constant
the other variables not used and you
have all the gas laws we’ve used so far.
Molar Volume
1 mole = 22.4 L of a gas at STP
Now, we can convert between moles
and grams; moles and
molecules/atoms; and moles and
volume (L)
Dalton’s Law of Partial
Pressure
The total pressure is equal to the sum of
the partial pressures
Avagadro’s Law:
V1 / n1 = V2/n2
Where n = number of moles
How do you convert grams to moles?
Ideal Gas Law
PV=nRT
P = Pressure (kPa)
V = volume (L)
n = number of moles
R = 8.31 kPa x L / moles x K
T = temperature (K)
You must use these units for the R
constant to be correct.
Name the Law!
You will be given a series of laws and
asked to name the law or you will be
given the name and be asked to come
up with the formula!
PV =nRT
Ideal Gas Law
V1 / T1 = V2 / T2
Charles’ Law
Boyle’s Law
P1V1 = P2V2
Combined Gas Law
P1V1 = P2V2
T1
T2
Gay-Lussac’s Law:
P1/T1 = P2/T2
Avagadro’s Law:
V1 / n1 = V2/n2