A mole is Avagadro’s number of something. You could concieve of a mole of baseballs if you wanted to (you’d have a hell of a time finding a closet big enough to put them in).
The usefulness comes from the fact that when you talk about moles of molecules, the mass of that 1 mole will be the molecular weight in grams. Eg., a mole of hydrogen (H2) will weigh 2 grams (approximately, for the wiseacres out there who are reading the decimal points on there periodic tables).
1 mole = Avogadro’s number of molecules, or 6.023E23 molecules.
So what? A mole of a certain compound will have a weight in grams equal to the compound’s molecular weight. Molecular weight is the total of the atomic weights that make up 1 molecule of the compound. Atomic weight is the total number of protons and neutrons in the nucleus, and is slightly adjusted to account for isotopes (which effect we can ignore for this discussion).
For instance, methane (CH4) has 1 carbon atom, atomic weight 12, and 4 hydrogen atoms, atomic weight 1, making its molecular weight 16. If you put 6.023E23 methane molecules together, their weight will add up to 16 grams.
As I start writing this, I can pretty much guarantee that someone will answer the question before I finish.
Molecular weight is the sum of the atomic weights of all the atoms that make up the molecule.
The gram molecular weight is that same number in grams.
So, if the molecule consists of 2 atoms with an atomic weight of 1, and one atom with an atomic weight of 14, the molecular weight is 16. 16g of this substance would contain Avogadro’s number of such molecules.
The molecular weight of a substance is it’s total number of neutrons and protons (protons and neutrons weigh almost exactly the same). For example, the atomic weight of carbon is 12, the molecular weight of N[sub]2[/sub] is 28, and the molecular weight of H[sub]2[/sub]O is 18. (This ignores complications due to isotopes of the elements.)
A mole of H[sub]2[/sub]O weight 18 grams; a mole of beryllium (atomic weight 9) weight 9 grams. Same number of particles (molecules in the one case, atoms in the other), but different weights.
The molecular weight is basically the total number of protons and neutrons in all the atoms making up a single molecule of the stuff. It’s the mass of a single molecule in Atomic Mass Units (AMUs) where 1 AMU is defined as 1/12 the mass of a carbon-12 atom.
For instance, the atomic weight of Hydrogen is 1 (nucleus contains 1 proton) and the atomic weight of Oxygen is 16 (nucleus has 8 protons and 8 neutrons). Water is H[sub]2[/sub]O, so the molecular weight if water is 18 (=16+21), therefore one mole of water (~6.0210[sup]23[/sup] water molecules) has a mass of 18 grams.
It gets a little more complicated when you factor in different isotopes and binding energy.
The molecular weight of a substance is equal to the atomic weight of all of the atoms in one molecule.
The atomic weight of an atom is equal to the number of nucleons (protons and neutrons) in its nucleus. Not exactly, but close enough for government work.
Thus, ethanol, which has the formula C[sub]2[/sub]H[sub]5[/sub]OH has a molecular weight of 46 (carbon has an atomic weight of 12, hydrogen of 1, and oxygen of 16).
A “mole” of ethanol, therefore, would be 46 grams of ethanol. A mole of water, OTOH, would be 18 grams; of glucose, 180 grams; of table salt, 58 grams (actually, closer to 58.443 grams, I think).
No, no- I could no sooner tell you whose answer was the most helpful than I could choose the most beautiful water droplet of those constituting a rainbow.
This was most helpful for it to be explained a number of different ways.
A lot easier to understand than the encyclopedia definition, with plenty of supporting examples.
molecular weight is (approximately) the weight of a hydrogen atom (i.e. proton). with this weight defined, any other atom or molecule can be expressed as multiples of this unit. therefore a carbon molecule has a molecular weight of 12 because its 6 protons and 6 neutrons have approximately 12 times the weight of a hydrogen atom. there are minor deficiencies in this explaination due to the fact that it is not defined as precisely the weight of a proton (i’m not sure what it is exactly, but it’s pretty close), difference between weight of protons and neutrons, matter converted to energy when seperate protons and neutrons are combined into nuclei, etc. but for our purposes, that’s close enough.
avogadro’s number is essentially a conversion factor (like 2.54 cm to the inch) between the realm of molecular weight and that of conventional metric weight. one hydrogen atom has molecular weight of one, and 6.023E23 (avagodro’s number) hydrogen atoms weigh one gram. one atom of carbon has molecular weight 12, 6.023E23 carbon atoms weigh 12g. glucose (C6H12O6) has molecular weight of something around 180 ((12 x 6) + (1 x 12) + (16 x 6)), so a mole of glucose weighs 180 grams. you get the idea.
Do I least get credit for trying? I said “mass”, then went on to say that an H2 would “weigh” 2 grams. I thought I’d let it pass. When the references all list atomic WEIGHT in atomic MASS units what do they expect?
I disagree. I’ve often heard the term “molecular weight” but never “molecular mass”. In fact, looking in my CRC produces a definition for “molecular weight” but not “molecular mass” (although both “atomic weight” and “atomic mass” are defined).
Admittedly, “molecular mass” makes more intuitive sense, but since “molecular weight” is both defined and understood, I’d use that term exclusively. At least for the sake of tradition.
There’s an underlying grammar question in this thread, to wit: Why do most (all?) chemistry books and other texts use that lame definition “1 mole is the gram molecular weight …” ? I mean, who dreamed up that term and thought it would be the clearest way to explain the idea? Why not just say that “Avogadro’s number of protons or neutrons weigh 1 gram” ???
Any of the posts in this thread would make more sense than the book one.
I think I remember that someone (Dalton) assigned molecular weights, starting with hydrogen. He assigned 1 as the mw of H just to have a reference point. He then called everything else a whole or fractional multiple of H.
I think I’m right in this part because a.m.u.s are units in and of themselves. They are not grams, etc.
As time went on, the measuring got more exact.
So: What’s the story between what I’ve said and the decision to use carbon as the standard?
How was Avogadro’s number determined? Certainly not by Avogadro. Did someone simply determine how many atoms it would take to have 12 grams of carbon 12? When did that happen? I assume calling it “Avogadro’s number” is in respect for his theory that equal volumes of gasses have equal numbers of particles.
Which came first, A’s number, or the idea of m.w.? Which was used to calculate which?