unit 5 - Gas Laws Flashcards
Kinetic Molecular Theory (ideal gas laws)
(KMT) is a theory used to explain the usual behavior of gases
- describes the relationship among:
pressure, volume, temperature, velocity, frequency of collisions
(KMT) ideal gas laws are
- gases contain particles that are in constant, random, straight-line motion
- gas particles collide with each other and with the walls of the container
- results in a transfer of energy - gas particles are separated by great distances
- no definite shape, no definite volume - gas particles DO NOT attract each other
ideal gases (vs context)
- particles spread out
- gases are most ideal at low pressure, high temperature
- Hydrogen and Helium are always ideal gases (low density)
real gases (vs context)
- particles are closer together
- gas particles DO attract each other
- gas particles DO occupy volume
- gases are most Real at high pressure and low temperature
avogadro’s hypothesis
if two gases are at the same exact temperature and pressure… AND the volumes that the gases occupy are the same
- the # moles (# molecules) are also the same
STP
standard temperature and pressure
- 273K = 0°C
- 1 atm = 101.3 Kpa
at STP, 1 mole of any gas will always occupy
22.4 L
pressure
the amount of force over a certain area
volume
the amount of space something takes up
robert boyle
discovered the relationship between pressure and volume of a gas
what is the only element that is liquid at room temp
mercury
standard pressure
760 mmHg = 760 Torr
1 atm
101.3 Kpa
boyle’s law
relationship between pressure and volume
- at constant temperature, as pressure is added to a gas (increases), the volume will decrease
- pressure and volume are inversely proportional
- P1V1 = P2V2
boyle’s law equation
P1V1 = P2V2
- 1 = initial state
- 2 = final state
- volume units - measured in mL or L
- pressure units - measured in Kpa, atm, mmHg, or torr
boyle’s law: if the pressure doubles, what happens to the volume
goes down by half
boyle’s law: if the volume doubles, what happens to the pressure
goes down by half
boyle’s law: if the pressure is halved, what happens to the volume
doubled
boyle’s law: if the volume is halved, what happens to the pressure
doubled
charles’s law
the relationship between temperature and volume of a gas when pressure is constant
- at constant pressure of a given gas, its new volume is DIRECTLY proportional to it’s change in temperature
jacques charles
discovered the relationship between temperature and volume of a gas
charles’s law equation
V1/T1 = V2/T2
- at constant pressure!!
- let V1 & T1 represent the initial volume and temp (in K)
- let V2 & T2 represent the final volume and temperature
- temperature MUST be measured in kelvins
why must temperature be measured in kelvin for Charles’s Law
- 0°C would make the fractions undefined
- Negative temps do not work mathematically
lussac’s law
relationship between temperature and pressure of a gas at a constant volume
- as temperature increases, pressure exerted by the gas increases
P1/T1 = P2/T2
combined gas law equation
- when ALL variables (temperature, pressure, & volume) change; nothing is held constant
(P1V1) / T1 = (P2V2) / T2
SIGFIG for this question: “convert 2.75 atm to Kpa”
2.75 atm x 101.3 Kpa/1 atm = 279 Kpa
^
only look at
2.75 atm for
sigfig!!
dalton’s law of partial pressure: mixture of gases
- dalton found that the TOTAL PRESSURE exerted by the mixture of gases is equal to the sum of the Partial Pressures of each gas in the mixture
- Total Pressure = P1 + P2 + P3
Law of Partial Pressure Example
If 5.0 moles of O2 and 7.0 moles of N2 together exert a total pressure of 200.0 Kpa, what is the partial pressure of O2?
5.0/12.0 x 200.0 Kpa = 83.33 Kpa
^
don’t use
this ratio
for sigfigs
Graham’s Law of Diffusion
- the rate at which gases diffuse is inversely proportional to their molar mass
- at STP, gases with a smaller molar mass will diffuse the fastest
graham’s law of diffusion example
Which gas at STP will diffuse the slowest?
1) O2 2) H2 3) Xe 4) Ar
32g/mol 2g/mol 131g/mol 40g/mol
Xenon will diffuse the slowest at STP
Graham’s Diffusion Formula
V1/V2 = sqrt (m2/m1)
KE = 1/2 mv^2
(V1/V2) <- relative velocity
at the same temperature, the KE of two different gases is the…
same
