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What is a mole (mol) and g/mol?

  1. Oct 23, 2012 #1
    Can someone explain to me these two things? I'm trying to get the average mass of air and it says it is 29g/mol? But what is a mole? I know it's the amount of carbon-12 atoms in 12 grams. What does 12 grams look like?
     
  2. jcsd
  3. Oct 23, 2012 #2
    One mole is a quantity of N0 atoms or molecules. Therefore, the atomic weight is the weight of one mole of atoms, and the molecular weight is the weight of one mole of molecules.

    And a gram is the mass of one cubic centimeter of water at 4 degrees C. 1 gram = 0.0352739619 ounces. Twelve grams is the weight of a normal letter, a handkerchief, something like that
     
  4. Oct 23, 2012 #3
    Just like a dusin is 12 and a million is 1000000, 1 mol is 6,022[itex] \cdot [/itex] 1023.

    The mol is defined from the massunit "units", which is defined as [itex]\frac{1}{12}[/itex] of the mass of a carbon-12 atom, so that 1 u = 1 g/mol. Thus one water molecule weighs 18.02 u whereas 1 mol of water molecules weigh 18.02 g.

    If you need a unit converter to get a better understanding of the metric system here's a good one: http://www.digitaldutch.com/unitconverter/

    When the average mass is given in units of g/mol (actually it's the molar mass M) you may calculate the density [itex]\rho[/itex] (mass per volume), which I suppose is what you're looking for, if you assume that air is an ideal gas:

    \begin{equation}
    \rho = \frac{p M}{R T}
    \end{equation}

    However, let me save you some trouble and tell you that the density of ordinary air at 0oC and 1 atm pressure is 1,293 g/L.
     
    Last edited: Oct 23, 2012
  5. Oct 24, 2012 #4
    I'm still confused. I'm only 13 and I'm like "huh?"
     
  6. Oct 24, 2012 #5
    The easiest way to look at it is that one mole of ANY substance is defined as 6.02214 × 10^23 molecules or atoms (which is 602,214 followed by 18 zeros). This is called Avogadro's number. A mole is just a convenient way to measure things in chemistry, and it was defined based on carbon because it happened to work out nicely with other things, as well.

    So when someone says air is 29 g/mol, what they're saying is that if you took 6.02214 x 10^23 air molecules, it would weigh 29 grams. If you had 2 moles of air, it would weigh 58 grams, and so on.
     
  7. Oct 24, 2012 #6
    Okay, so a mole of something is just how many molecules or atoms there are. If there is 3 moles of water molecules, it would equal (6.0221415 X 10^23) X 3. Now how do you get the mass of a single molecule? For this example, can you use an air molecule?
     
  8. Oct 24, 2012 #7
    If you look on a periodic table, the atomic weight is actually the same as g/mol. So if you have a molecule of air, which I'm assuming is O2, then you would add the atomic weights of two oxygen atoms, and that's your g/mol. If there's anything else in your air molecule then you'll want to add that atomic weight as well.

    Chemists made it easy for us that way. :)
     
  9. Oct 25, 2012 #8
    Maybe you could be more specific about what you need it for? Someone here can probably provide you with exactly the number you need in units you are familiar with.

    (But please, keep asking questions of this sort anyway!)
     
  10. Oct 25, 2012 #9

    Borek

    User Avatar

    Staff: Mentor

    There is a grain of truth in what you wrote, but in general its is completely wrong.

    Mol is not defined in mass units. As you correctly wrote earlier, mole is a just an overgrown dozen. What you refer to here is the fact that amu (atomic mass unit) is selected in such a way that mass of a single atom (molecule) expressed in amu is numerically identical with a mass of a mole of such atoms (molecules) expressed in grams. Makes calculations easier.
     
  11. Oct 25, 2012 #10
    Hmm, I dindn't write that mol is defined in mass units, but from it... Anyway, we obviously agree on what 1 mol and 1 u are, so maybe we shouldn't confuse the boy further by pedantically contradicting each other.
     
  12. Oct 25, 2012 #11
    Okay, well in the average molar mass (is that what you call it?) of air: 29g/mol, the mass is 29 grams for every mole amount of air molecules. So to get the mass of a single molecule in a molecule, you divide the mass by a mole? So 29/(6.0221415 X 10^23) = 4.81 X 10^-23 [appoximately]. Is that correct?
     
  13. Oct 26, 2012 #12

    Borek

    User Avatar

    Staff: Mentor

    Average molar mass is OK.

    "For every mole of molecules of X" or even "for every mole of X" just like you won't say "for every kilogram amount of water" but "for every kilogram of water".

    Yes.

    You just have to remember this is an average mass, so there is no single molecule of this mass in the real air, but the air on the whole in some aspects behaves as if it was made of molecules of this mass.
     
  14. Oct 28, 2012 #13
    Don't get caught up in all of that right now though. There's these things called isotopes, where atoms have different amounts of neutrons, but that's not important to this discussion.

    A "mole" is just a convenient unit. Most textbooks compare it to "a dozen", because it's actually used in the same way. we call twelve things a dozen, in the same way that we call 6.022x10^23 things "a mole". The only thing is that describing 602,200,000,000,000,000,000,000 (six hundred sextillion? lol) of anything other than atoms isn't very useful, so you won't see "moles" used anywhere other than chemistry, really.
     
  15. Oct 28, 2012 #14

    jedishrfu

    Staff: Mentor

    why use moles?

    Because chemistry is actually based on number ratios of atoms and molecules:

    as in H2O:
    - take a dozen Oxygen atoms and combine them with 2 dozen H atoms, or
    - combine a dozen Oxygen atoms with one dozen H2 molecules.

    moles are a convenient measure for numbers. Chemists know how to convert from weight to moles for each kind of compound or element.
     
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