Concept A: Absolute Zero.
The lowest possible temp is around -273degC.
It is labeled 0K on the Kelvin scale.
K = degC + 273
degC = K - 273
Gas volumes are proportional to K temps.
A temp change from 20K to 40K will double the volume.
A temp change from 20degC to 40degC does not.
(It will increase the volume by 313K / 293K.)
ALL TEMPS USED IN GAS LAWS MUST BE IN KELVINS.
Concept B: Standard Temperature and Pressure. (STP)
Solids and liquids have a defined volume.
Gases have a defined volume if the temperature and pressure are defined.
Standard Temperature = 0degC = 273K
Standard Pressures = 1 atmosphere (atm) = 760mmHg = 760torr = 101.3kPa = 14.7psi (pounds per square inch) = 29.9inHg.
You'll probably have to convert between these units.
Equation 1: The combined gas laws.
P2V2 / T2 = P1V1 / T1
Temps have to be in Kelvins.
Pressures have to have the same units, but it doesn't matter which one.
Volumes have to have the same units, but it doesn't matter which one.
This equation is used to figure out a new value when a fixed amount of gas undergoes a change in conditions of PV or T. Typically a volume of gas will be collected at lab conditions (say 25degC and 740.mmHg) and you will be asked to calculate the volume collected at STP.
V2 = V1P1T2 / P2T1.
Equation 2: Ideal Gas Law
PV = nRT.
This equation is used to calculate one value when the other 3 are known.
T must be in Kelvins.
P and V must have the same units as R. (or R the same units as P and V)
n measures moles.
R is a constant and can be calculated based on the units given in the problem by substituting standard conditions and volume (22.4L or equivalent) for 1 mole of a gas.
R = (PV / nT)
R = (1atm)(22.4L) /[(1mole)(273K)] = 0.0821Latm/moleK
R = (760torr)(22.4dm^3) / [(1mole)(273K)] = 62.4torrdm^3/moleK etc.
By substituting m/M (mass/molarmass) for n, a couple other interesting equations show up.
PV = mRT / M
P = mRT / MV
PM / RT = m/V = D
For molar mass
PV = mRT / M
MPV = mRT
M = mRT / PV
Gas law 3: Graham's Law
v1 / v2 = sqrt(M2 / M1)
v (lower case) stands for velocity not Volume (upper case V).
M stands for molar mass.
This equation applies to two samples of different gases at the same temp and pressure.
If the velocity of one gas is known, the velocity of the other can be calculated.
Alternatively, the "relative rates" of gas1 to gas2 can be calculated.
Or gas1 is how many times faster than gas2?
In the latter two examples it's that ratio that's important not a specific value for v1 or v2.
Note that the subscripts are criss-crossed in the equation.
It's possible you will also have to calculate the volume of a dry gas, but that requires an explanation of open and closed manometers.
You may also have to understand how to do a calculation involving a real gas.
The equation contains corrections that compensate for the assumptions made about ideal gases:
1 that particles are point masses, (have mass but occupy not space.)
2 that collisions are completely elastic (no energy lost.)
See it and a description of the rest of the laws and some examples at