 So, this lesson is going to be on atomic mass. So what are atomic masses? You see the masses of individual atoms are very small. The mass of a carbon-12 atom is about 1.998 times 10 to the power of negative 26 kilograms. So that is a extremely, extremely small number and you don't want to be dealing with the exponents all the time. So that's why chemists develop a unit called the atomic mass unit. So the atomic mass unit, it is defined as exactly a 12th the mass of a carbon-12 atom. The symbol that you usually see will be either U or DA called Doulton. For example, the atomic mass of a carbon-12 is 12U, exactly 12U or 12 Doulton doesn't matter. They are the same thing. So this is our reference of how to use that in practice, comparing with other elements. For example, you have the atomic mass of a sodium-23 is about 2298U or Doulton, 101 is 1.0783U or Doulton, hydrogen-2 has an extra-neutral, so the weight will be probably doubled. So that will be about 2.014 to 1.0U. So that's how atomic masses are used to describe the weight of the atoms. So the table being shown over here is the Euclideic masses of each of the common isotopes of the common elements. If you look at the carbons on the left, you see the carbon-12, where it's exactly 12U or 12Doulton. Because that's just how atomic mass was defined in the first place, and that is each U is a 12th of the mass of carbon, therefore will be 12 times that, which is 12U. Now, if you look at your periodic table, then below the symbol for each element, you usually find a small number, which of your decimal points, and that is your average mass, or average atomic mass. So let's write a number down. So you have your carbon as an average mass of 12.01 Doulton. Now, how do you get a number? That 12.01 number here is not carbon-12, it's not carbon-13. Now, how does this calculate? Now, the symbol for this is A subscript, little r, with the element in brackets. Now, how does it calculate? It is that you multiply the percentage of the natural abundance of that specific isotope, so in this case, carbon-12. You multiply by the nucleic mass of a carbon-12, which is 12Doulton. That plus the percentage for the natural abundance of the second isotope multiplied by the nucleic mass of the second isotope. Now, if you have more, you're just going to keep going. Now, sometimes you don't need to, specifically for carbon, for example, you don't need to go any further than this. And the reason is because all the isotopes are so rare that you really never encounter them. The percentage of the natural abundance of carbon-12 is 98.90 percent. So that's multiply by 12 plus the natural abundance of carbon-13 is 0.10 percent multiplied by the nucleic mass. Now, that is equal to roughly 12.011Doulton. And that is what this number here is about.