Ok, so CO2 (atomic mass 44) is about one and a half times as dense as atmospheric oxygen (atomic mass 16, so O2 = 32).
Then why doesn’t all the CO2 just sink gradually to the bottom, suffocating all the spaniels & rodents underground?
Atomic mass and density are not correlated. CO[sub]2[/sub] is denser than air, but not by much, and the winds tend to keep it pretty well mixed. There isn’t that much CO[sub]2[/sub] in the air anyway…only about .033%.
CO2 and heavier gases can accumulate in deep, still areas, just as released helium and hydrogen will go up, up and away. But one of the qualifiers with the heavier gases is “still”.
Air mixes very easily due to its very low viscosity, and thus any type of circulation and turbulence (such as the wind) will easily stir things up.
Actually they are very well correlated for gases at low pressures and high temperatures (basically Ideal Gases)
PV = nRT where n is the number of moles and P is pressure, V is volume, R the Universal Gas Constant and T the absolute temperature.
n = m/M (m = mass of gass , M = molecular mass)
So, PV = (m/M) RT
Or, P = (m/V) (RT/M)
or (PM/RT) = m/V
Since m/V is density, Density of ideal gases = PM/RT,
**So the higher the molecular mass, the higher the density, for ideal gases **
But what really keeps the CO2 from settling down is diffusion and brownian motion. Just like if you put diesel and gasoline together diesel wont settle to the bottom - but will mix.
** Another caution, CO2 at room temperature is a vapor not a gas - but the ideal gas laws are good enough at these conditions **
Buh? What’s the difference between a vapor and a gas? (This kind of stuff is probably why I got a C- in Statistical Physics.)
Typically, a gas is a vapor if it is below the temperature at which its vapor pressure is less than atmospheric. Thus, water vapor is is vapor at standard temperature and pressure until about 212 F, and after the latent heat of vaporization is overcome. Or, I guess you can say it’s in the “two phase” region at 212. It’s kinda a nitpicky definintion that engineers use, although it’s not 100% agreed upon in its usage and applicability. There are other definitions, many not bourne by the Dictionary.
I’m uncertain why CO2 at normal ambient temperatures would be a vapor. Are we thinking different definitions here? 
I am a chemistry major and I don’t know the meaning of the distinction between vapor and gas in this context.
Or, as Anthracite said.
The difference between gas and vapor is the critical point. You can liquefy/solidify (in case of CO2) vapors at atmospheric temperature by the application of pressure alone because it is below its critical temperature. Whereas oxygen you can never liquefy at room temperature.
That is why you can get dry ice (soli CO2 since CO2 sublimes).
CO2 follows Van Der Waals State Equations pretty well.
OK, that makes sense to me. I can buy that, although all of my handbooks don’t mention that difference in nomenclature.
Interesting. Thanks much to Anthracite and andy_fl, and to Dr. Lao for keeping me company. 
Could you refresh my memory: what is an ideal gas?
A gas that obeys the Ideal Gas Law (PV = nRT).