Formula Weights and Molecular Weights

I'm back, this time with a set of imaginary weighing scales, for now is the time that I'm going to talk (write?) at you about formula weights and molecular weights!

In Atomic Structure and the Periodic Table, I mentioned the mass number. Here we're going to take that concept and take it to the next level.

Mass number started out as relative atomic mass. The mass of atoms were compared to that of a Carbon-12 atom.

But since an atom is so tiny and has such a small mass, just how did they compare the mass?

Well, some genius took 12g of a sample of Carbon-12 and discovered that the sample had 6.02 x 10^23 atoms in it (a very special number as you will soon find out when we talk about calculations). Taking roughly the same number of atoms of any other element would give the mass of the sample in grams- a number reasonably easy to work with. This number could then be compared to Carbon-12 to get the relative atomic mass of an atom (Carbon-12's relative atomic mass is 12).

For example, if 6.02 x 10^23 atoms of an atom were found to weigh twice as much as the same number of atoms of carbon, its relative atomic mass must be double Carbon-12's relative atomic mass. The atom's relative atomic mass is therefore 24.

And then some smart people worked out ways to take isotopes into account, which is why the mass numbers on your periodic table aren't whole numbers. Carbon-12's mass number, for example, is listed as 12.01 on most periodic tables.

Back to formula weights and molecular weights, which is why we're here. Remember that number 6.02 x 10^23? Well, this number has a special name. It is called the mole. I could make lots of lame jokes about moles here, but chances are that you've heard them all from your chemistry class.

When you work out the formula or molecular mass, you're basically trying to work out the combined mass of all of the atoms listed in a compound's formula in one mole of the compound. You do this simply by adding up the mass numbers of the atoms in the formula. For example, if you want to find the mass of one mole of NaCl, just add the mass numbers of Na and Cl together. Of course, when an atom appears more than once in the formula or has subscripts, you need to take that into consideration when you're working out the molar mass by multiplying that atom's mass number by how much of the atom you need. For example, in N₂O_{4, you have to multiply the molar mass of N by 2 and add it to the molar mass of O multiplied by 4.}

... That was a really long explanation for such a simple concept. I'm going to take a break now.