Learn about Thermodynamics - part 4 Video

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Learn about Thermodynamics - part 4

Learn about Thermodynamics - part 4 Before I continue, I want to introduce to what might be an unfamiliar concept although, if you’ve taken chemistry, you might know a little bit about it. It’s called a mole. And this isn’t the thing that grows on your face with a hair on it or the animal that digs in you r backyard. Although, those are also called moles but we’re talking about the SI unit called the mole. And a mole is just a number. It’s like saying a mole of something means a certain number of something just like a dozen or—I don’t know—what other things are there that are numbers. But it’s like saying a dozen eggs is 12 eggs, right? A dozen = 12. Well just like that. A mole of eggs would be 6—I always forget the exact number. It’s 6.023 I think; something of that nature. 6.—you can look it up. I think it’s 6.023 × 1023. Let me look it up to actually the right—the exact number just because I think I’m—that 23. I’m misremembering. Let’s see the mole. I’m on Wikipedia now and I just looked up the mole. Okay there you go. On Wikipedia, see it’s various—oh yeah, I was close. It is 6.022 × 1023 of something. So it’s a very large number of something. So, normally we don’t deal in moles of eggs. I don’t think there have been a mole of eggs ever produced in the history of Universe. 1023 is a very, very, very large number. So where― is it useful? Well a mole is useful for counting atoms. And so what is a mole of atoms? Well it’s that—or molecules. What is that — many molecule? It is 6 followed by roughly 23 zeros of molecules, very, very big number. But what is interesting about a mole is that when I have a mole of something, it’s mass so, let’s say it’s mass in grams. So let’s say a mole of carbon. It’s mass in grams is going to be equal to—so mole of carbon. If I have a mole—so if I have this many carbon molecules, its mass in grams, it will have x grams. It will have a mass of x grams where x is the atomic mass number; atomic mass of a atom of carbon. Although― if I was saying about a mole of a molecule, I would figure out the atomic mass of the entire molecule. So what’s an atomic mass number? Let me see if I can do a web search on a periodic table. See I am showing you what I do here. It’s not fancy. So let me go to Google.com. Let’s look up periodic table. Let’s see if we can find a good one― periodic table. This one looks good. Let’s see what I can do. It’s sometimes― Looks like everything is freezing up too. Oh there you go. So Periodic Table of Elements, good. Let’s see what we can do. So, if we go to Carbon which is right here. We see that its atomic number 6 and that’s the number of protons it has. Let’s see if I can zoom in on carbon, what happens― periodic Table. Oh there you go. That’s pretty neat. So the atomic mass number is the mass of the entire atom. And just so you know, I mean we’re kind of delving into a little bit of chemistry here but most of the mass of an atom is the protons and neutrons. And the neutrons and the protons weigh roughly the same thing. And then the electrons are much, much, much smaller. So, if you pretty much factor in the mass of the protons and the neutrons, you pretty much have the mass of the particle. And then just another, a little more chemistry here is that although on average, most of the atoms have roughly the same number of protons and neutrons. Some don’t, some—you know you could have a carbon atom that has 7 neutrons. You can have another one that has 5, another one with 6. And those are actually all called isotopes and I won’t go into all of that but they’re just the same atom with different numbers of neutrons. But in general, the atomic mass number just, if you had kind of a broad rule of thumb is equal to—oh sorry. The atomic mass is equal to the mass of the protons and the neutrons. And they tend to be equal. So if the atomic number is 6, the atomic mass tends to be 12. So why is this useful? So, we can say if we have—I don’t know. What is this, gallium? If I have a mole of gallium, right? Let’s say Niobium. Let’s say I have a mole of Niobium. If we look here on the Periodic Table, it has an atomic mass number of 41 and then its average atomic mass if we were to average all of the isotopes or kind of how—you know based on their waiting on how they exist to nature. It is 92.9 so roughly 93. So it’s actually a little bit more than double its atomic number. But let’s say it’s 93. So if we had a mole of Niobium. If we had 6.022 × 1023 of Niobium, it would have a mass of 92 grams. So that’s pretty easy. You look at any element; Molybdenum. We see here—let’s say Chromium. We see its atomic mass is roughly 52. We see that there. So, if I have a mole of it, if I have roughly 6 × 1023 of it that much will have a mass of 52 grams. So that’s how we think about a mole. So if tell you I have a mole of something, I am also telling you how many of that a molecule I have and I am also telling you what the mass of that mole the quantity will be if assuming that you have a periodic table in front of you. So with that said, with that out of the way, let’s make some more progress with out thermodynamics. So we said in the last several videos—let me see when I’m running out of time. Well, I have plenty of time—that pressure times volume it’s somehow proportional—you know, let’s call that K and this is an arbitrary number, it’s not some special constant—to the total kinetic energy, kinetic energy total of a system. And we also said that that is roughly proportional. That is some constant, that’s another constant. Times the number of molecules we have times the temperature because temperature we view this kinetic energy per molecule, right? So in general, we could also that this is proportional of this which is proportional of this. That pressure times the volume is proportional and we’ll use R because you’ll see where that’s coming from in a second. It’s equal to some constant times the number of molecules n. And when I write a small n here; see here I’ll just take the absolute number. So that five molecules up to the five here. But now, this n, I am counting in moles, right? So if I say I have—if this n is 1; that means I have 6.022 × 1023 molecules. So let me see. 1 mole = 6.022 × 1023, right? So this is just another way to write the number of molecules and then that’s times temperature. And then if we rearrange it, PV = nRT, we have a relationship that if I know the pressure and the volume and the number of molecules, I can figure out the temperature or if I know the number of molecules, the temperature and the pressure, I can figure out the volume assuming I know what R is. And I am about to tell you what that is. R is called the universal gas constant and it is 8.31 joules/mole Kelvin. And that kinds of tells you what you need in this formula. This should end up being joules, right? So if you have pressure in Pascal and you have volume in m3, you’ll end up with joules there. This should be in moles. This is 8.31 joules per mole Kelvin. And then this as we always said should be in Kelvin. And honestly, if you just memorize two things in all thermodynamics, you’ll probably be able to do 95% of problems but you actually should have the intuition of how they work. But just remember that PV/T = constant or that if you change them, they relate to each other that they all equal to constant. So if P1 × V1 ÷ T1 = P2 ×V2 ÷ T2. And then you also should just need to memorize PV = nRT where R = 8.31 joules per mole Kelvin. And I know you might not have a lot of intuition of this formula. Yeah because I haven’t use it but I am going to do that in the next video but these are literally the two most important things you should know in thermodynamics and hopefully you have a little intuition at this point of what they mean. See you soon.

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Learn about Thermodynamics - part 4

Before I continue, I want to introduce to what might be an unfamiliar concept although, if you’ve taken chemistry, you might know a little bit about it. It’s called a mole. And this isn’t the thing that grows on your face with a hair on it or the animal that digs in you r backyard. Although, those are also called moles but we’re talking about the SI unit called the mole. And a mole is just a number. It’s like saying a mole of something means a certain number of something just like a dozen or—I don’t know—what other things are there that are numbers. But it’s like saying a dozen eggs is 12 eggs, right? A dozen = 12. Well just like that.

A mole of eggs would be 6—I always forget the exact number. It’s 6.023 I think; something of that nature. 6.—you can look it up. I think it’s 6.023 × 1023. Let me look it up to actually the right—the exact number just because I think I’m—that 23. I’m misremembering. Let’s see the mole. I’m on Wikipedia now and I just looked up the mole. Okay there you go. On Wikipedia, see it’s various—oh yeah, I was close. It is 6.022 × 1023 of something. So it’s a very large number of something.

So, normally we don’t deal in moles of eggs. I don’t think there have been a mole of eggs ever produced in the history of Universe. 1023 is a very, very, very large number. So where― is it useful? Well a mole is useful for counting atoms. And so what is a mole of atoms? Well it’s that—or molecules. What is that — many molecule? It is 6 followed by roughly 23 zeros of molecules, very, very big number.

But what is interesting about a mole is that when I have a mole of something, it’s mass so, let’s say it’s mass in grams. So let’s say a mole of carbon. It’s mass in grams is going to be equal to—so mole of carbon. If I have a mole—so if I have this many carbon molecules, its mass in grams, it will have x grams. It will have a mass of x grams where x is the atomic mass number; atomic mass of a atom of carbon. Although― if I was saying about a mole of a molecule, I would figure out the atomic mass of the entire molecule.

So what’s an atomic mass number? Let me see if I can do a web search on a periodic table. See I am showing you what I do here. It’s not fancy. So let me go to Google.com. Let’s look up periodic table. Let’s see if we can find a good one― periodic table. This one looks good. Let’s see what I can do. It’s sometimes― Looks like everything is freezing up too. Oh there you go. So Periodic Table of Elements, good. Let’s see what we can do.

So, if we go to Carbon which is right here. We see that its atomic number 6 and that’s the number of protons it has. Let’s see if I can zoom in on carbon, what happens― periodic Table. Oh there you go. That’s pretty neat. So the atomic mass number is the mass of the entire atom. And just so you know, I mean we’re kind of delving into a little bit of chemistry here but most of the mass of an atom is the protons and neutrons. And the neutrons and the protons weigh roughly the same thing. And then the electrons are much, much, much smaller. So, if you pretty much factor in the mass of the protons and the neutrons, you pretty much have the mass of the particle.

And then just another, a little more chemistry here is that although on average, most of the atoms have roughly the same number of protons and neutrons. Some don’t, some—you know you could have a carbon atom that has 7 neutrons. You can have another one that has 5, another one with 6. And those are actually all called isotopes and I won’t go into all of that but they’re just the same atom with different numbers of neutrons.

But in general, the atomic mass number just, if you had kind of a broad rule of thumb is equal to—oh sorry. The atomic mass is equal to the mass of the protons and the neutrons. And they tend to be equal. So if the atomic number is 6, the atomic mass tends to be 12.

So why is this useful? So, we can say if we have—I don’t know. What is this, gallium? If I have a mole of gallium, right? Let’s say Niobium. Let’s say I have a mole of Niobium. If we look here on the Periodic Table, it has an atomic mass number of 41 and then its average atomic mass if we were to average all of the isotopes or kind of how—you know based on their waiting on how they exist to nature. It is 92.9 so roughly 93. So it’s actually a little bit more than double its atomic number.

But let’s say it’s 93. So if we had a mole of Niobium. If we had 6.022 × 1023 of Niobium, it would have a mass of 92 grams. So that’s pretty easy. You look at any element; Molybdenum. We see here—let’s say Chromium. We see its atomic mass is roughly 52. We see that there. So, if I have a mole of it, if I have roughly 6 × 1023 of it that much will have a mass of 52 grams. So that’s how we think about a mole. So if tell you I have a mole of something, I am also telling you how many of that a molecule I have and I am also telling you what the mass of that mole the quantity will be if assuming that you have a periodic table in front of you.

So with that said, with that out of the way, let’s make some more progress with out thermodynamics. So we said in the last several videos—let me see when I’m running out of time. Well, I have plenty of time—that pressure times volume it’s somehow proportional—you know, let’s call that K and this is an arbitrary number, it’s not some special constant—to the total kinetic energy, kinetic energy total of a system.

And we also said that that is roughly proportional. That is some constant, that’s another constant. Times the number of molecules we have times the temperature because temperature we view this kinetic energy per molecule, right?

So in general, we could also that this is proportional of this which is proportional of this. That pressure times the volume is proportional and we’ll use R because you’ll see where that’s coming from in a second. It’s equal to some constant times the number of molecules n. And when I write a small n here; see here I’ll just take the absolute number. So that five molecules up to the five here.

But now, this n, I am counting in moles, right? So if I say I have—if this n is 1; that means I have 6.022 × 1023 molecules. So let me see. 1 mole = 6.022 × 1023, right? So this is just another way to write the number of molecules and then that’s times temperature.

And then if we rearrange it, PV = nRT, we have a relationship that if I know the pressure and the volume and the number of molecules, I can figure out the temperature or if I know the number of molecules, the temperature and the pressure, I can figure out the volume assuming I know what R is. And I am about to tell you what that is.

R is called the universal gas constant and it is 8.31 joules/mole Kelvin. And that kinds of tells you what you need in this formula. This should end up being joules, right? So if you have pressure in Pascal and you have volume in m3, you’ll end up with joules there. This should be in moles. This is 8.31 joules per mole Kelvin. And then this as we always said should be in Kelvin.

And honestly, if you just memorize two things in all thermodynamics, you’ll probably be able to do 95% of problems but you actually should have the intuition of how they work. But just remember that PV/T = constant or that if you change them, they relate to each other that they all equal to constant. So if P1 × V1 ÷ T1 = P2 ×V2 ÷ T2. And then you also should just need to memorize PV = nRT where R = 8.31 joules per mole Kelvin.

And I know you might not have a lot of intuition of this formula. Yeah because I haven’t use it but I am going to do that in the next video but these are literally the two most important things you should know in thermodynamics and hopefully you have a little intuition at this point of what they mean. See you soon.

Transcription by: Scribe4you Transcription Services