1
Q
miscible
A
- capable of being mixed in any proportion without reacting chemically
- gases are miscible
- many liquids are immiscible, such as oil and water
2
Q
How do gases differ from liquids and solids?
A
- Unlike with liquids and solids, the volume occupied by a gas changes significantly with pressure
- altitude impacts volume
- volume of a gas changes with temp
- gases are miscible
- gases are typically much less dense than liquids or solids.
- we can tell b/c gas densities are g/L while liquids are g/mL
*these 4 ideas are consistent with the fact that particles of a gas (molecules OR atoms) are farther apart than those of liquids and solids.
3
Q
atmospheric pressure(Patm)
A
- the force exerted by the gases surrounding Earth on Earth’s surface and on all surfaces of all objects
- Pressure = Force/Surface Area
4
Q
devices used for measuring gas
A
- barometer: measures atmospheric pressure
- 1 meter tube is inverted and open end is immersed in mercury
- gravity pulls mercury down, atmospheric pressure pushes mercury up into tube due to force of collision of N2 and O2 molecules with mercury
- the net effect of these opposing forces causes the heigh of the mercury in the tube which is the measurement
- manometer: measures the pressure exerted by a gas
5
Q
Boyle’s Law
A
- constant T and quantity of gas
- P = 1/V
- As the volume of a constant amount of gas(at constant temp) increases, the pressure of the gas decreases b/c of fewer collisions
- As the volume of the gas decreases, its pressure increases because of more collisions
- PV = constant
- P1V2=P2V2
6
Q
Charles’s Law
A
- the principle that the volume of a fixed quantity of gas at constant pressure is directly proportional to its absolute temperature
- temperature in Kelvin
- Kelvin are always positive!
- V/T = constant
- V1 / T1=V2 / T2
- temperature in Kelvin
7
Q
Avogadro’s Law
A
- amount of gas is not constant
- obviously volume depends of amt of gas
- constant P and T
- V/n = constant
- when P and T are constant
- V/n = constant
8
Q
ideal gases + ideal gas law
A
- atoms or molecules are assumed not to interact with one another
- instead they move independently with speeds that are related to their masses and the temp of gas
- most gases exhibit ideal behavior at the P and T typically encountered in the atmosphere
- STP: P=1atm T=0˚C
- n=m/mw
- PV=nRT
- R=0.08206 Latm/molK
- combined gas law
- P1V1/T1 = P2V2/T2
9
Q
partial pressure
A
- the contribution to the total pressure made by one gas in a mixture of gases
10
Q
mole fraction (Xx)
A
- the ratio of the number of moles of a componrnt in a mixture to the total number of moles in the mixture
- 1) unlike molarity, mole fractions have no units
- 2) unlike molarity, mole fractions are based on numbers of moles and can be used for any mixture or solution
- 3) the mole fractions of all the componenets in a mixture must sum to 1
- relationship btwn partial pressure and total pressure
- Px=XxPtotal
11
Q
kinetic molecular theory
A
- a model that describes the bahvior of ideal gases
Assumptions
- gas molecules have tiny volumes compared to the total volume they occupy so their individual volumes are negligible
- they are point masses(have essentially no volume)
- separated by large distances(mostly empty spaces)
- gas molcules move constantly and randomly
- the motion of these molcules is associated with an average KE that is proportional to the absolute temp of the gas
- all populations of gas molecules have same avg KE
- KE is determined by temp
- gas molecules continuously collide elastically (no net energy transfer to the walls)
- avg KE is NOT affected by these collisions and is constant as long as temp is unchanged
- each gas molecule acts independently of all other molecules. we assume there are no forces of attraction or repulsion btwn the molecules
12
Q
molecular speeds and KE
A
- KE=1/2(mu2)
- m is mass
- u is speed
- collisions btwn gas molecules cause molecules in a sample to have a range of speeds
- root-mean-square speed(urms): speed of a molecule possessing the avg KE.
- At a given temp, the pop. of molecules in a gas has the same avg. KE as every other pop. of gas molecules at the same temp this KEavg is
- KEavg=1/2(m)(urms)2
- urms=√(3RT/M)
- M=molar mass
- R=8.314
13
Q
effusion vs. diffusion
A
- effusion: the process by which a gas escapes from its contained through a tiny hole into a region of lower pressure
- Graham’s law of effusion: rate of effusion of a gas is inversely proportional to the √molar mass
- rx/ry = √My/Mx
- gas with greater rms speed leaks at a higher rate
- the escape of gas molecules requires encountering the microscopic holes in the balloon. The faster a molecule moves, the more likely it is to find one of these holes.
- Graham’s law of effusion: rate of effusion of a gas is inversely proportional to the √molar mass
- diffusion: the spread of one substance through another
14
Q
Why the ideal gas law isn’t good enough
A
- The ideal gas law considers gas molecules to have no volume at all, however, at high pressures, more molecules are squeezed into a given volume and it becomes significant
- KMT assumes molecules don’t interact, but real molecules do attract one another, so ideal gas law only works when molecules are really far apart
- as pressure increases gas molecules are pushed together and IMF become significant
- likewise, lowering temp slows molecule movement resulting in more attractions and more deviations from ideal behavior
15
Q
van der waals equation
A
Accounts for the fact that:
- the free volume of a real gas is less than the total volume b/c its molecules occupy significant space
- the observed pressure is less than the pressure of an ideal gas b/c of intermolecular attractions
a and b are constants!
- a term: related to attractive/repulsive forces
- b term: takes into account volume of gas particle
16
Q
Dalton’s Law
A
- Since gas molecules are far apart we assume they behave independently
- Mixture of gases: each individual gas exerts its own pressure(partial pressure)
- Dalton’s Law: Ptotal = P1+P2+P3+……
