Molecular mass and molecular weight (video) | Khan Academy
Molar mass and molecular weight are often confused, but their values are very different. Molar mass is the mass of one mole of a substance. A tutorial on Molar mass calculations with examples and a molecular weight in stoichiometry (the quantitative relationships between chemical substances in a. Fomula mass and molecular mass are two values that express the size of a molecule. Do you know the difference between formula mass and.
Which is the weighted average of the various isotopes of oxygen as found on earth. Our approximation to 16 is pretty good. Then based on these numbers you would say that this H2O has an atomic mass of approximately Well, two from the hydrogens, where did I get the two from? Each of these two hydrogens have an atomic mass of one. If you have two times one, it's just gonna be two atomic mass units. Then 16 from the oxygen. Two plus 16, which is going to get us Let me do this in another color, since I've been using It's going to give us 18 atomic mass units.
Now, if you wanted to be more precise or if you wanted to say, "Well, I have this big bag "of water, I'm not looking exactly at one water molecule.
Then it is helpful, especially if you're talking about a large number of molecules and you really just wanna take the weighted average of all of those molecules. Then it makes sense to say, "Well, let's "use the atomic weight.
Difference between Molar Mass and Molecular Mass
Which is, once again, 1. We call it atomic weight but it's really just the weighted average, it's not weight in kind of the physics sense of measuring a force. I got my calculator. I'm gonna have two times 1. Let's see, I should go no more than three decimal places to the right. Since I added this, if I don't wanna add precision here If you wanna a review of that, you should look at the video on significant figures.
Since I added something with just three decimals to to the right I shouldn't have more than three decimals in my answer so, So the real one is You could actually consider this the molecular weight. Because once again, we're using atomic weights. Any readily measurable mass of an element or compound contains an extraordinarily large number of atoms, molecules, or ions, so an extraordinarily large numerical unit is needed to count them.
- Molecular mass and molecular weight
- Chapter 1.7: The Mole and Molar Mass
- Molecular mass
The mole is used for this purpose. A mole is defined as the amount of a substance that contains the number of carbon atoms in exactly 12 g of isotopically pure carbon According to the most recent experimental measurements, this mass of carbon contains 6. Just as 1 mol of atoms contains 6. Since the mass of the gas can also be measured on a sensitive balance, knowing both the number of molecules and their total mass allows us to simply determine the mass of a single molecule in grams.
The mole provides a bridge between the atomic world amu and the laboratory grams. It allows determination of the number of molecules or atoms by weighing them. The numerical value of Avogadro's number, usually written as No, is a consequence of the arbitrary value of one kilogram, a block of Pt-Ir metal called the International Prototype Kilogram, and the choice of reference for the atomic mass unit scale, one atom of carbon A mole of C by definition weighs exactly 12 g and Avogadro's number is determined by counting the number of atoms.
It is not so easy. Avogadro's number is the fundamental constant that is least accurately determined.
Chapter The Mole and Molar Mass - Chemistry LibreTexts
The definition of a mole—that is, the decision to base it on 12 g of carbon—is arbitrary but one arrived at after some discussion between chemists and physicists debating about whether to use naturally occurring carbon, a mixture of C and C, or hydrogen. The important point is that 1 mol of carbon—or of anything else, whether atoms, compact discs, or houses—always has the same number of objects: In the following video, Prof.
Most atomic masses are known to a precision of at least one part in ten-thousand, often much better  the atomic mass of lithium is a notable, and serious,  exception. This is adequate for almost all normal uses in chemistry: The precision of atomic masses, and hence of molar masses, is limited by the knowledge of the isotopic distribution of the element.
If a more accurate value of the molar mass is required, it is necessary to determine the isotopic distribution of the sample in question, which may be different from the standard distribution used to calculate the standard atomic mass. The isotopic distributions of the different elements in a sample are not necessarily independent of one another: This complicates the calculation of the standard uncertainty in the molar mass.
A useful convention for normal laboratory work is to quote molar masses to two decimal places for all calculations. This is more accurate than is usually required, but avoids rounding errors during calculations. These conventions are followed in most tabulated values of molar masses.