chapter 5 Flashcards
what are gases composed of?
• Gases are composed of molecules or atoms which are spaced very far apart • These particles are in constant motion - colliding with each other - colliding with the walls of the container
what do gases do in a container?
• Gases expand to assume the
volume of the container they are in
• Gases are compressible
what is pressure? formula? what is the amount of pressure felt related to?
Though we aren’t keenly aware of the air around us,
except to breathe it, gas molecules exert a force on the
surfaces that they collide with.
• Pressure is the force per unit area exerted by gas
molecules colliding with the surfaces around them.
Pressure = F/A
The amount of pressure felt is
directly related to the amount
of gas molecules present
units for pressure?
mmHg Torr atm Pa (pascal) psi bar • SI unit is the pascal (Pa) 100 kPa = 750.01 Torr = 1 bar
what is a Torr equal to?
mm Hg is also called a Torr (after Torricelli, inventor of
the barometer)
• While many mercury column instruments are still in use,
most modern instruments use other methods of pressure
sensing due to the cost and toxicity of mercury
what force will a gas with a pressure of 1 Pa exert? what is a Pa?
A gas with a pressure of 1 Pa will exert a force of 1 N on a
wall of area 1 m2
• The Pa is a small unit – 1 atm = 760 mm Hg = 101325 Pa
Many chemists continue to use atm and mmHg
• Meteorologists often use units of bar (1 bar = 100000 Pa)
what did boyles law observe? formulas?
Observed an inverse relationship between volume and
pressure- an increase in one results in a decrease in the
other
p1v1=p2v2
at constant T and n, v=1/p
as pressure increases, volume decreases
reason pressure increases as voulme decreases
We know that pressure from gases is due to the particles hitting the walls of the container • If we decrease the volume the particles are now crowded into a smaller place, and there are more collisions with the container walls • This means that the pressure increases
what is charles’ law?
Charles’ law relates the volume and temperature of a
gas
• Volume is directly proportional to temperature, therefore:
when temperature increases the volume increases
• At lower temperatures the
particles have lower
kinetic energy and occupy
a smaller volume
v1\t1=v2/t2
what is avogadros law?
Describes the relationship between the volume of gas and the amount of gas • Volume is directly proportional to the amount of gas (constant T and P) • When we increase the amount of gas, the volume must increase to maintain the same pressure v1/n1=v2/n2
What is the ideal gas law?
combines them, pv= nrt
what is the ideal gas constant?
The value R is the ideal gas constant The same for every gas R=0.08206 In order to use the ideal gas law, the units on all of our other variables must be the same as found in R Pressure (P) = bar volume (V) = L moles (n) = mol Temperature (T) = K
how do we apply the ideal gas law?
Just as with solution molarity, when we deal with gases we have a molar volume Volume one mol of gas occupies • STP Temperature - 273.15K Pressure - 1 bar Moles - 1 mol `
how do we find total pressure of a mixture?
Most gases exist as a mixture rather than a pure gas
• They each exert their own pressure
Partial pressure, Pn
• The sum of the partial pressures gives us the total
pressure
Pt = Pa + Pb + Pc…
pa= na(rt\v) etc
what is the mole fraction? what represents it?
We can compare the partial pressure, Pa to the total
pressure, Ptotal
• The ratio of pressures is equal to the ratio of the number
of moles:
Pa/Ptotal = na/ntotal
• We can represent the mole fraction with Xa=na/ntotal
• This allows us to represent the partial pressure in terms of
the mole fraction and total pressure
how is density related to molar mass?
• Since we are dealing with a gas we describe density as
g/L
• The density of a gas is directly related to its molar mass
LOOK AT SLIDE FOR EXAMPLE
what is molar density?
Density relates the mass of a substance to its volume
D= m
V
• We know that 1 mole of any gas occupies a volume of
22.7 L under STP
• We also know that 1 mole of any gas has a mass equal to
it’s molar mass
D= molar mass/molar volume
how do we convert L into moles? why?
• When we look at reactions that happen in the gaseous
phase the amounts are given in L, for a specific
temperature and pressure
• We use the ideal gas law to convert L into moles to use
the stoichiometry of a reaction
• We can also relate the change in moles of a reaction to
the change in pressure at a constant volume and
temperature
molar volume and stoich. Under STP conditions what volume does 1 mol of gas
occupy? what can we use this amount as?
Under STP conditions what volume does 1 mol of gas
occupy?
• We can use this amount as a conversion factor when we
have reactions at STP
22.7L
How many grams of water form when 1.24L of H2(g) at
STP completely reacts with O2(g)? According to the
equation:
2H2(g) + O2(g) → 2H2O(g)
what is there between many liquid substances?
• For many liquid substances (and some solids), there is
an equilibrium between the condensed (solid or liquid)
phase and the gas phase
What is the vapor pressure?
• The portion of the substance in the gas phase exerts
pressure like any other gas
• This is called the vapor pressure of the liquid
what does vapor pressure rely on?
The position of this equilibrium depends on the nature of
the substance, and on the temperature
• The more the equilibrium favors the gas phase, the
higher the vapor pressure
• Vapor pressure increases with temperature
what is waters vapor pressure at 100 degrees?
Water has a vapor pressure of 760 mmHg (1 atm) at
100°C
what temp does water boil at? why?
• Water boils at 100°C at standard sea level pressure
• Water (and all other liquids) boils at a lower temperature
at lower pressure, and a higher temperature at higher
pressure
edmonton conditions
• Here in Edmonton, the pressure is about 700 mmHg and
water boils just below 98°C
what happens when vapor pressure exceeds the total pressure of the gas?
If the vapor pressure exceeds the total pressure of the
gas(es) over the liquid, the liquid phase will boil.
• The boiling point of a liquid is thus dependent on the
pressure
• Consider boiling water at various elevations:
At 760 mmHg (1 atm), water will boil at 100°C
What about at higher elevations? Top of a mountain
→Lower pressure
What about higher pressure conditions? Pressure
cooker
what does kinetic molecular theory describe?
The kinetic molecular theory is used to describe the
behaviour of gases
Remember: Theories explain behavior. Kinetic
molecular theory is used to explain the behavior of
gases observed in the three basic gas laws (and the
ideal gas law)
how does KMT model gases?
The kinetic molecular theory models gases as single particles that move in a straight line until it collides with another particle or the walls of the container
kmt assumption on size
1.The size of each gas particle is negligibly small. We
assume that the particles occupy no volume.
Argon gas at STP occupies 0.01% of the volume,
each argon atom is 3.3nm away
kmt assumption on average energy
- The average kinetic energy of a particle is proportional
to the temperature in kelvins. Motion of atoms is due to
thermal energy. At a given temperature there is a
distribution of velocities, but higher temperatures lead to
larger velocities and a higher average kinetic energy.
kmt assumption on collision
The collision of one particle with another (or the walls of
the container) is completely elastic. Energy can be
exchanged during collision, but it is never lost. This
means the molecules are not “sticky” or deformed during
collision.
how does kmt relate to the ideal gas law?
Starting with the kinetic molecular theory we can derive
the ideal gas law. This means that the KMT is a
quantitative model that implies PV=nRT.
• The kinetic molecular theory can also be used to explain
the simple gas laws
kmt and boyles law
Boyle’s Law states that at a constant temperature and
constant quantity of gas pressure and volume are
inversely proportional
• If you decrease the volume of a container, the gas
particles will have less space to move around in before
they collide with the walls or other gas particles
• If the frequency of collisions increases that means the
pressure will increase
• Thus, as volume decreases pressure increases
• P α 1/V (Boyle’s Law)
kmt and charles law
• Charles’ Law states that at a constant pressure and
quantity of gas volume and temperature are directly
proportional
• Temperature of a gas is the average kinetic energy of the
gas molecules
• If you increase the temperature of a gas the average KE
will increase, this means the gas particles will collide
more frequently and will greater force
• In order to maintain a constant pressure the volume must
increase to re-adjust the collision rate back down
• Thus, volume and temperature are directly proportional
• V α T
real gas assumptions and kmt
The theory that explains the ideal gas law (PV = nRT)
assumes that gas particles (molecules) occupy no space
and do not interact with each other
• All real molecules occupy a finite amount of space, so
the zero volume assumption is never strictly true
• All real molecules are also attracted to each other to a
certain extent
so why do we use ideal gas law?
So why do we still use the ideal gas law?
• Under typical laboratory conditions with common gases,
these assumptions are nearly true
• The volume occupied by the gas molecules is small
relative to the size of the container
• Interparticle interactions are weak enough to be ignored
molar volumes under stp
Most gases under STP conditions have molar volumes
close to the ideal gas molar volume
• STP conditions are considered low pressure and high
temperature (relative)
how do real gases differ from ideal? volume deviations?
Gases deviate from ideal behaviour at high pressures
• At high pressure the volume of the actual gas molecules
becomes significant
• The actual volume of the gas increases as pressure
increases (due to the volume of the atoms/molecules)
• The volume occupied by gas molecules results in less
than the full volume of the container being available for
the molecules to move around in
what do deviations result in? how can we correct for them?
This increases the pressure relative to the predictions of the ideal gas law • This volume can be accounted for by introducing a correction factor V-nb • ‘b’ is a constant that is specific for specific gases • ‘n’ is the number of moles of the gas, identical to the “n” in the ideal gas law • The effect is only significant at high particle density (high pressure)
how do real gasses act at low temps?
Gases deviate from ideal behaviour at low temperatures
low temp vs high temp
• All substances experience some amount of
intermolecular attractive forces
• At high temperature the particles have enough kinetic
energy to overcome the forces and thus experience
elastic collisions
• At low temperatures this is not true, as particles collide
the attractive forces temporarily halt the subsequent
motion and the rate of collisions is decreased
what do attractive forces do?
Attractive forces between gas molecules reduce the
pressure relative to the predictions of the ideal gas law
• Because intermolecular interactions are bimolecular,
they are proportional to the square of the molecule
density
how do we correct for attractive forces? when is this effect significant?
The correction for attractive forces involves replacing P
with (P + a(n/V)2
), where n/V is the particle density and
“a” is an empirically determined proportionality constant
that describes the strength of the intermolecular
attractions.
• The effect is only significant at low temperatures, where
the attractive forces are significant relative to the kinetic
energy of the gas molecules
what do these corrections result in together?
Together, these corrections result in the van der Waals
equation
• (P + a(n/V)2
)(V - nb) = nRT
