We're gonna now discuss Avagadro's Constant and the Mole. I know this doesn't come up until chapter three of our textbook, but I think it's an appropriate place to start discussion about this. And the mole refers to a specific number and we're gonna use that as a unit throughout the rest of this course. So, we see this in our everyday lives. Not to this scale of Avogadro's constant or the mole, but if we look at various units that we typically have in our everyday language, there's a pair or a dozen, and the mole is just like each one of these units. So we can relate this to the number particles in each of these units. A pair is gonna tell us we have two of something. A dozen means we have 12. A mole means we have 6.022 times 10 to the 23rd. Again, exact same concept, but much larger in magnitude, than two or 12. So, keep this in mind, kind of as we go forward. Now, the reason Avogadro came up with this number, or this constant, is because we needed a good way. To measure the chemical reactions that are going on. Everything's happening at an atomic level. And we can't physically weigh every single atom. We need a quantity that is manageable for us to look at. And investigate as we see chemical reactions happen. So, for instance, if we wanted to say I have one dozen baseballs. And, we put these on a balance and take their mass, there's gonna be, it's gonna weigh 1,740 grams. If on the other hand I have one dozen softballs. Put those on a balance, they're gonna weigh roughly 2140 grams. The reason I'm using baseballs and softballs here is because atoms are often visualized as spheres. So here we have two spheres, they're different sizes and obviously they're different masses. But we still have a dozen of them. And these values right here are definitely different because the make up or the composition of each of these particles are also different. Same thing happens with atoms and we need to look at a couple of different things here. So, first, let's look at elements. So, for any elements. One mol is gonna equal the atomic weight. In grams. And again, we've touched on some mass units earlier these atomic mass units. Let's get something that's a little bit more manageable. So, we're gonna look at the atomic weight in grams, and that's what's gonna be our mol. So, remember 1 mol is 6.022 x 10 to the 23rd of something. So, if we're looking at 1 mol of V, that's gonna equal 6.022 times 10 to the 23rd V atoms, and if we put that many vanadium atoms on a balance, we're gonna get a mass of 50.9415 grams. If on the other hand, we took one mole of Mg, that's gonna be 6.022 times 10 to the 23rd Mg atoms. If we put those on a balance, we find that the mass is gonna be 24.305 grams. Again, 1 mol of each of these substances are different just as the mass of the baseballs and softballs were different. One dozen baseballs has a different mass than one dozen softballs. One mole of vanadium is gonna have a different mass than one mole of magnesium. And we see that here and these guys are gonna be different and this is the mass that's also reported on the periodic table. We're gonna come back to this in chapter three and discuss these, but it's kind of a good transition to our next topic, that here we have elements, and elements form together to give compounds. So if we have, for a molecule. Or compound. We can say that one mol is equal to the molecular weight in grams. So, if we have one mol of CO2, that's gonna = 6.022 x 10 to the 23rd CO2 molecules. And if we put those on a balance, that's gonna give us a mass of 44.0 grams. So, here we have CO2 and I list this as a molecule verses describing the elements having atoms. So the question is what is the distinction between. In these atoms and molecules and what happens when compounds come together. So that's gonna be the topic of our next video.