chapter 5

Question 1

what are gases composed of?

Answer 1

• 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

Question 2

what do gases do in a container?

Answer 2

• Gases expand to assume the
volume of the container they are in
• Gases are compressible

Question 3

what is pressure? formula? what is the amount of pressure felt related to?

Answer 3

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

Question 4

units for pressure?

Answer 4

mmHg
Torr
atm
Pa (pascal)
psi
bar
• SI unit is the pascal (Pa)
100 kPa = 750.01 Torr = 1 bar

Question 5

what is a Torr equal to?

Answer 5

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

Question 6

what force will a gas with a pressure of 1 Pa exert? what is a Pa?

Answer 6

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)

Question 7

what did boyles law observe? formulas?

Answer 7

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

Question 8

reason pressure increases as voulme decreases

Answer 8

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

Question 9

what is charles’ law?

Answer 9

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

Question 10

what is avogadros law?

Answer 10

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

Question 11

What is the ideal gas law?

Answer 11

combines them, pv= nrt

Question 12

what is the ideal gas constant?

Answer 12

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

Question 13

how do we apply the ideal gas law?

Answer 13

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
`

Question 14

how do we find total pressure of a mixture?

Answer 14

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

Question 15

what is the mole fraction? what represents it?

Answer 15

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

Question 16

how is density related to molar mass?

Answer 16

• 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

Question 17

what is molar density?

Answer 17

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

Question 18

how do we convert L into moles? why?

Answer 18

• 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

Question 19

molar volume and stoich. Under STP conditions what volume does 1 mol of gas
occupy? what can we use this amount as?

Answer 19

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)

Question 20

what is there between many liquid substances?

Answer 20

• For many liquid substances (and some solids), there is
an equilibrium between the condensed (solid or liquid)
phase and the gas phase

Question 21

What is the vapor pressure?

Answer 21

• The portion of the substance in the gas phase exerts
pressure like any other gas
• This is called the vapor pressure of the liquid

Question 22

what does vapor pressure rely on?

Answer 22

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

Question 23

what is waters vapor pressure at 100 degrees?

Answer 23

Water has a vapor pressure of 760 mmHg (1 atm) at

100°C

Question 24

what temp does water boil at? why?

Answer 24

• 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

Question 25

edmonton conditions

Answer 25

• Here in Edmonton, the pressure is about 700 mmHg and | water boils just below 98°C

Question 26

what happens when vapor pressure exceeds the total pressure of the gas?

Answer 26

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

Question 27

what does kinetic molecular theory describe?

Answer 27

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)

Question 28

how does KMT model gases?

Answer 28

``` 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 ```

Question 29

kmt assumption on size

Answer 29

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

Question 30

kmt assumption on average energy

Answer 30

2. 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.

Question 31

kmt assumption on collision

Answer 31

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.

Question 32

how does kmt relate to the ideal gas law?

Answer 32

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

Question 33

kmt and boyles law

Answer 33

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)

Question 34

kmt and charles law

Answer 34

• 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

Question 35

real gas assumptions and kmt

Answer 35

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

Question 36

so why do we use ideal gas law?

Answer 36

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

Question 37

molar volumes under stp

Answer 37

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)

Question 38

how do real gases differ from ideal? volume deviations?

Answer 38

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

Question 39

what do deviations result in? how can we correct for them?

Answer 39

``` 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) ```

Question 40

how do real gasses act at low temps?

Answer 40

Gases deviate from ideal behaviour at low temperatures

Question 41

low temp vs high temp

Answer 41

• 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

Question 42

what do attractive forces do?

Answer 42

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

Question 43

how do we correct for attractive forces? when is this effect significant?

Answer 43

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

Question 44

what do these corrections result in together?

Answer 44

Together, these corrections result in the van der Waals equation • (P + a(n/V)2 )(V - nb) = nRT