[positiveion]

I am having serious trouble understanding it.

Sure I've gotten all the basic definitions of 'oh it's just a number', 'just pretend its not there and replace the amu with the mole' .. but I don't understand WHY

Why is the mole a direct translation of its atomic mass number? What is even the point of having this system of moles if its just the same thing as its amu? WHY is it the same thing as its amu?

[renge ishyo]

The mole is a number that was obtained from late 19th century experiments concerning gases. It is useful because it is a standard number of molecules that can be used to compare the masses of different chemical substances. First came Avogadro's idea that equal volumes of different gases at the same temperature and pressure contain the same number of molecules. Then, someone went out and measured the number of molecules in one mL of gas and came out with the answer that there were about 6.022x10²³ molecules in a mL of gas. The usefulness of measuring the number of molecules in one mL of gas is that the definition of the gram has it that 1mL of water at its point of highest density is defined to weigh one gram. So with this all in place, here is how you can go about obtaining the mass of different chemical substances:

1. Say you take 1 mL of Oxygen gas and 1 mL of Nitrogen gas at STP and weigh them. You know that each contains the same number of molecules (6.022x10²³) because you have 1 mL of each, but their masses are different when you place the gas on a scale...Oxygen weighs in at 32 grams and Nitrogen at 28 grams.

2. You can use this information to find out how much 1 molecule of oxygen gas and 1 molecule of nitrogen gas weighs:

1 single molecule of oxygen (O₂) gas = 32 g/6.022x10²³ molecules = 5.31x10^-23 g/molecule!

1 single molecule of nitrogen (N₂) gas  = 28g/6.022x10²³ molecules = 4.65x10^-23 g/molecule!!

3. So now we can directly compare the weight of one molecule of oxygen gas to one molecule of nitrogen gas, but umm...there is a problem, who wants to do calculations using numbers like 5.31x10^-23? That would be a pain in the rear. So O.K., to make people that don't like math happy, why don't we just define the number 6.022x10²³ molecules to be one of something in the same way that we define 12 donuts to be 1 dozen so that we can count large groups of donuts easier? So lets just say that 6.022x10²³ molecues equals something we call "1 mole". When you do this the above equations become:

6.022 x10²³ molecules of oxygen (O₂) gas = 32 g/1 mole or 32g/mol

6.022x10²³ molecules of nitrogen (N₂) gas  = 28g/1 mole or 28g/mol

And now everyone is happy because the calculations are easy again. Well, almost everyone. Now people are angry because the relative weights need to be absolute, independent of the gram standard. So O.K...you jerks...we will invent the "amu" or "atomic mass unit" which is dimensionless but just "happens" to be equal to 1g/mol, and now everyone is happy. Well...almost everyone (students are of course quite unhappy with all this confusion when they first start studying chemistry, but after some time they do get used to it...the confusion that is, and then it doesn't bother them as much).

Hope this helped.

[Sam (NG)]

renge ishyo's thinking is correct, except that 1 mole of a gas at STP will not occupy 1 ml.

Much closer to 23 litres if you consider it as an ideal gas.

[renge ishyo]

Yes, that's correct. STP slipped in out of habit, but the gases certainly wouldn't be at STP in this case. For that matter, this description isn't historically accurate either. The true history behind how the mole concept developed is decidedly more complex.