0:00 Welcome to general chemistry. This course provides a comprehensive introduction to general chemistry, covering everything from the basic building blocks of matter to the principles governing chemical reactions. You'll explore essential concepts like atomic structure, bonding, thermodynamics, and equilibrium. Gain a solid foundation in chemistry and prepare yourself for future scientific endeavors. Chad Mallister, who teaches this course, spent over 15 years managing the tutoring facility for the chemistry department at a large US university. Matter, that's going to be the topic in this very first lesson of my brand new general chemistry playlist. And it's an appropriate topic because chemistry, simply put, is the study of
0:46 matter. And we're going to learn a lot of vocabulary of how we look at matter and how we classify matter and how we measure matter in this first lesson. Now, before we dive in, I really just want to give you a proper introduction to the entire playlist. This general chemistry playlist is going to span two semesters of college general chemistry. That's general chemistry 1 and general chemistry 2. And we're going to present this in roughly 90 lessons, give or take, spanning about 20 chapters of general chemistry. It is meant to be thorough and comprehensive to just about anything you'd see in your typical college general chemistry course. Now, in addition to having master courses for many of your high school and college science courses, we also have DAT, MCAT, and OAT prep as well. Just an FYI. All right, I have been yammering on enough here. So, let's dive in and let's have a little discussion about matter. So, what
1:33 is matter? So, simply put, matter is anything that has mass and occupies space. So, it has mass and it has volume would be another way of stating it. So, that's what matter is. uh simply put and what you want to know though is that matter is made up of these super teenytiny particles that we call atoms. So and we've kind of you know built on what the some of the ancient Greeks says they were atomists. They believed in these atoms being these tiny building blocks. Well our conception of what an atom now is a little bit different than theirs but we are using the same word they did. We've been using it for a couple hundred years now. And so uh but you should know that matter is made up of atoms. And it turns out there are different types of atoms and there's roughly 118 different types of atoms. And those different types we call elements. But these 118 or so elements,
2:20 that's the number of types of different atoms that are out there. And everything in the universe that we can see, so is made up of these atoms of 118 different types. Okay. Now, it turns out this matter comes in three major phases or three major states. And we're going to call those solid, liquid, and gas. And you might be like, "Oh, Chad, what about plasma?" Well, you're right. There is a fourth phase called plasma. And we're not talking about it. So, you're not going to see it in this class. We'll have a whole chapter dedicated to talking about solids and liquids toward the end of the first semester topics. We'll have a whole chapter talking about gases. We will not have a whole chapter talking about plasma. So, we're not really going to include it. So, but you should be intimately familiar with solids, liquids, and gases, uh, but only marginally familiar now. But once we get there later in the semester, you'll learn a lot more about uh, all three phases. So what you should know though
3:06 is that solids have a definite shape and a definite volume. So both definite shape and definite volume. Whereas a liquid a liquid when you pour it into a container it just takes the same shape as the container. So it does not have a definite shape but it does have a definite volume. And then finally gases. So it turns out gases are going to fill whatever container you put them in and they will expand to fill that container and stuff like that. Or you can compress them to fit into a smaller container and things of a sort. So, not not only do they not have a definite shape like liquids, but they don't even have a definite volume either. And so, that's kind of what you should know in in this description of solid versus liquid versus gas. Again, solids uh definite shape and definite volume. Liquids have only a definite volume but no definite shape. And gases have neither definite volume nor definite shape. And if you
3:53 kind of look at what these look like at the atomic level in a solid, what you'll find is that most solids are what we call crystallin solids. So, and the atoms are kind of arranged in a nice repeating pattern. We call this a crystal structure. And so, if you kind of looked at a a lovely illustration of what a solid might look like, it might look something like this. And it turns out that these don't actually just have to be individual atoms. So, but they might very well be, but they could end up being molecules, which are combinations of atoms as well. So, that's what a solid would look like. It turns out there's also something called an amorphous solid which we'll study much later in the semester where it's actually pretty random where these lovely atoms are. So, but the key is most solids are crystallin and the idea is that typically these atoms aren't really moving to any significant extent.
4:38 Now, a liquid is going to be different. So, you'll find out that uh liquid is a little more random and you'll find out that these atoms are moving around as well. They're flowing, if you will. And so they're not going to be in this crystal uh lovely structure and they're not going to be restricted in their motion. They're be able to move around as compared to a solid. Now gas is going to be a little bit different because now you're going to find out that these atoms are actually spread out and they're moving. And it turns out they're moving typically faster than the corresponding atoms or molecules in a liquid are moving. So, if you're looking at illustration, you might just have a general idea of how to compare a solid versus a liquid versus a gas. Makes a lovely test question. Cool. Now that
5:23 we've gone through the three phases of matter, we've now got to compare and contrast some new vocab terms uh that are going to help us to describe some of the properties of matter uh as well as kind of some of the uh classifications of matter as well. Let's dive in. All right. So, the first way we're going to classify matter is we're going to compare and contrast elements and compounds. And we've already used that word element. We have this lovely periodic table of the elements. And uh in this case all these lovely elements have got their lovely chemical symbol which we use to represent them. So like sodium is Na and carbon is a C and so on and so forth. Uh if we look at what these elements are again this is one of the fundamental types of atoms that exist is one way to look at that element. But there's another way. So it turns out that a substance that is only made up of one type of atom is also
6:09 called an element. But then a compound is going to be a substance that is made up of multiple elements. So a good example here of an element might just simply be that sodium. So we got plain old sodium metal. And if you look at sodium metal, it's a solid at room temperature. And every atom in that metal is by and large going to be sodium. Now there there's a chance there could be some small amounts of impurities. But I mean assuming you've got pure sodium metal, every atom would be sodium. On the other hand, you can have sodium chloride, which is salt, aka table salt. the same stuff you put on your food that your mom probably has lying around in your kitchen. And this is a compound because it actually has two different types of atoms present in a 1:1 ratio. Uh, and there sodium atoms and there are chlorine atoms. Now, it's still just a single substance, but this
6:55 single substance isn't made of just one type of atom. It's made up of two types of atoms. Now, we've got to be careful on how we we look at this because whether you've got an element or a compound, you're typically uh as long as you're not in the gaseous state anyways, you're typically going to have a a bunch of atoms bonded to each other. So, in pure sodium metal, all the sodiums are in a big chunk in the solid phase and in some measure they're bonded to each other. Whereas in sodium chloride, you get a repeating crystal and now you've got all these atoms bonded to each other as well. But now it's not just one type of atom, it's two types. And again in this case present in a 1:1 ratio. So both of these could be used to describe a single substance. But this gets a little bit problematic because now we've used the word element in two different ways. We've here used element as one of
7:42 the components in a compound. So and that's totally uh acceptable. But the other use of the word element is when the element itself is a standalone substance. So like when I have sodium metal. So here I've got the element sodium in a compound. And here I've just got the element sodium itself. And so I just want to make sure you just you realize there's two different very similar ways we use that word element. Uh but not quite exactly the same. Now it turns out that not all the elements on the periodic table in their elemental form come one by one by one like sodium. So it turns out some of them actually come two by two by two. So it turns out like nitrogen, oxygen, florine, chlorine,
8:29 bromine, iodine, and hydrogen all come 2x two x two in their standard what we call elemental form. We refer to these as the diatomics, the two atoms, if you will. So, and there's seven of them, and you should memorize them. And we'll bring this up later on in the semester again, but I figured I'd give you a preview here. So, and what's convenient is that six out of the seven here form the shape of the number seven on the periodic table itself. So, nitrogen, oxygen, florine, chlorine, bromine, iodine, and then hydrogen is the seventh one. Sometimes we'll give you a pneummonic to help you memorize this as well. And we'll say never have fear of ice cold beer. So, either way you memorize it, you're on the hook for memorizing all seven of these lovely diatomics. And so, in their elemental form, they come 2x two by two. But just
9:15 keep in mind again when they're part of a compound they don't have to come by two you know two at a time they could they don't have to. So but this would be elemental chlorine and this would be again non-elemental chlorine part of a compound one just one component of a compound. All right so I just want to make that clear a little bit and elements versus compounds. And now we've got to talk about uh what happens when you've got more than one substance present. Uh what do we call it there? All right. Now, we're going to compare and contrast a pure substance versus a mixture. In a pure substance, you've got one substance, but that one substance could be just one element or it could be just one compound. Either way, we refer to as being a pure substance. So, if I have say
10:01 pure oxygen gas or just pure sodium metal or something like that. So, that's just a single uh element present. Okay, great. But a pure compound, something like water. So water is made up of two elements, but it's still just a single compound, a single substance. And as long as I've got water and no other substance is present, this would be an example of a pure substance as well. So a pure substance is either a pure element or a pure compound. Either one qualifies. Whereas now, a mixture is going to be some combination of multiple substances. Now, we got to be careful here because a lot of students will look at this and be like, well, what's the difference between having two elements like this or having like air? So air, it turns out, is mostly a mixture of nitrogen and oxygen
10:46 gases. Well, the difference is this. In water, the hydrogen oxygen atoms are bonded to each other. So in air, though, air has different components that are together in the mixture, but they're not actually bonded to each other. And so, and that's what makes the difference. It's not a compound. It's actually a mixture instead. Are the atoms bonded together? Yes or no? If they are, then those two elements being bonded together forms a compound. So if they're not, then these two elements together form a mixture instead. And so air is mostly made up of nitroen oxygen, a little bit of carbon dioxide, a little bit of argon, things of a sort. Uh but if you have more than one element that are still separate substances but mixed together, that's a mixture. You could also have multiple compounds like we could have salt
11:32 water. And guess what's in salt water? Well, salt and water. And again salt is NaCCl and water is H2O. In this case again these don't end up bonded to each other with chemical bonds or anything like that but they are present together. And because we have multiple substances together in this case two compounds that once again is going to be a mixture. So you're going to have to be able to uh distinguish between pure substances and mixtures. And now it turns out we're going to make this even more complicated because the next level of classifying matter is you're going to have to distinguish between two different types of mixtures as well. All right. So the two in principal types of mixtures that exist are referred to as homogeneous mixtures and heterogeneous mixtures. And difference are going to be in these prefixes here that help us memorize the
12:18 definitions here between homogeneous and heterogeneous. And homo you might know means the same and heterero means different. And so in a homogeneous mixture, what is the same is the composition throughout the mixture. Like if you had a big mixture and you took a little sample here and a little sample there and a little sample there and a little sample there, no matter where you take the sample, it would have exactly the same amount of every component. The composition would be the same or we use the word uniform. So we often say that the composition is uniform throughout the mixture. Another word for a homogeneous mixture is uh the word solution as well. And when you think of the word solution, you're typically going to think of a liquid solution, but you can also have solid solutions and gaseous solutions as well. And again, as long as the components are spread out
13:04 uniformly uh throughout the entire mixture, so it has uniform composition, then it will be a homogeneous mixture. So a good example of this would be air. If I take a little sample of the air in this room at anywhere in this room, so the likelihood is that there's going to be a set amount or a set percentage of nitrogen like 78% as well as oxygen 21% as well as a little bit of CO2 and things of a sort throughout the room. That's kind of the way it works. It's got a uniform composition throughout. Now heterogeneous is an example of a mixture that does not have a heterero sorry not have a uniform composition throughout. Not doesn't have the same composition throughout. So let's go a little non-chemistry example here first. So in this case, let's go
13:49 with a salad. So if you look at a salad, typical salad, you know, it's got a lot of different components. It's a mixture in this case. Uh and if you take a bite of that salad, so sometimes you might get some tomato in a bite, sometimes you might get cucumber, sometimes you might not. And depending on where you got that bite, where in the bowl you got that bite, you might get a different set of components in every bite. It does not have a uniform composition throughout. And so a salad would be a great example of a heterogeneous mixture. So a nice supreme pizza would also qualify as well. So in some bites of that supreme pizza, you might get some pepperoni and some bites you might not. In some bites you might get sausage, some bites you might not. Some bites you might get olives, some bites you might not. The composition is not uniform throughout.
14:34 Another good example of a heterogeneous mixture. So let's talk about some more chemistry related examples of this. So and here's going to be a little bit strange. We're going to put salt water up here. And then we're going to put salt water over here as well. And this is about to get a little confusing. So, it turns out salt water could be homogeneous or heterogeneous depending on the situation. So, let's say I took a pinch of salt and I dissolved it in water and it all dissolves. Well, it turns out when it all dissolves, it will spread out uniformly throughout the entire solution, in which case it would be a homogeneous mixture. However, it turns out that uh salt is only going to
15:20 be so what we call soluble in water. And it turns out not, you know, if you you put too much salt in, not all of it's going to dissolve. And once you get past that certain point, the rest of it is just going to sink to the bottom and sit on the bottom of your glass or your beaker or your flask or whatever you're uh got for a container there. And if you get to that point, well, then you're going to find out that you have a ton of salt just sitting on the bottom and then you have, you know, lesser amounts dissolved throughout the solution. But depending on if you take a sample right at the bottom or or somewhere else, you might get a different concentration of salt, a different composition for the mixture, in which case it would be a heterogeneous mixture. So that's kind of the deal there. Give me give you one more. And we might use something like milk. So, and milk could be another example of a homogeneous mixture. So, up
16:07 to a point, right? If you leave that milk out on the counter all weekend, then it might curdle. And the parts that are curdled versus the rest, it might actually no longer be homogeneous. So, but milk itself is homogenized. You might hear that word used. And because it's homogenized, uh, the reason they use that word is because it is a homogeneous mixture at that point. So milk in the normal sense though you should consider homogeneous. I realize I use the curdled example but typically if you see this on a test they're giving you normal milk not curdled milk. All right. So now we've kind of looked at some different ways we classify matter. We contrasted elements and compounds and then pure substances versus mixtures and then the two different types of mixtures here. And you might have to classify those. Uh but now we want to talk about some of the properties of matter and how we classify
16:53 some of these properties. All right. All right. So now we want to talk about two different classes of properties. So which are things that describe matter. And it's going to be chemical properties versus physical properties. And if you have some sort of change going on associated with a chemical property, we'll call it a chemical change. And same thing over here with physical change related to a physical property. Now it turns out the big difference between these two different types of properties or changes uh let's deal with the changes typically. If the change converts one substance into a different substance, we're going to refer to that as a chemical change. However, if whatever change is going on does not convert one substance into a different substance, then it's just going to be a physical change. So, for example, if I take the study guide here, so and I tear it in half. So, turns out paper here is
17:40 made of cellulose. And the paper that I've got here is still made of cellulose. So, I haven't turned one substance into a different substance. It's still the same substance. Tearing the paper would be considered a physical change. And if we look here, this isn't exactly cellulose, but it turns out cellulose is made up of glucose. So, I'm going to kind of model it as glucose. And when I tear it, it's still cellulose. So, glucose in this case, uh, it I haven't turned it into a different substance. Just have two pieces of it now instead of one bigger piece. Okay. However, if I took that same study guide and I lit it on fire, things would be a little bit different then instead. And it turns out when you light it on fire, you convert it into
18:25 CO2 and H2O. So, we'll learn how to balance these reactions later, but I just want to show you that where you're turning one substance into different substances. And that's why burning it would be a chemical change rather than a physical change. And so the the chemistry term we might use in we might use the word burning but you might also hear the word combustion used and combustion is a chemical change but every chemical change often has a typically has a property associated with it and the property associated with it here would be called flammability and so here we have a chemical change and a chemical property. All right. So, other examples of physical changes and and definitely more much more likely to see some of
19:11 these uh than tearing a piece of paper uh are things associated with phase changes. We talked about the different phases of matter, solid, liquid, and gas. And it turns out they can change back and forth like when you go from liquid to solid, you're freezing something. And when you're going from solid to liquid, you're melting that something and stuff like that. But it turns out you're not converting one substance into a different substance. So like if we just talk about melting something like let's say we're melting some ice. Well the way we describe that is we would take water that's a solid and convert it into water if that's a liquid. So we often will take the particular phase of matter and write it in parenthesis right after the formula for a substance to ident identify which phase it was in. And so here we just got water that's a solid
19:56 phase going to water in the liquid phase but it's water. It's H2O. It's exact same substance, the exact same compound in this case, both before and after we melt it. And that's why melting is just a physical change. We haven't converted one substance into a different substance. Cool. And it turns out that the property associated with this change we might call the melting point. And in the case of water, it turns out there's a characteristic temperature at which a substance is going to melt. And for water, that's 0° C or 32° F. You'll find out that we're probably going to use Celsius a lot more in this particular class. Uh, and so 0 degrees CC, that's the temperature at which water melts. Now, different substances are all going to melt at different temperatures, but it is a characteristic physical property of a
20:43 substance is its melting point. Another example would be like its boiling point, things of this sort. All right. If we look at some other examples that you might encounter somewhere along the way, you might have like a nail rusting. A nail rusting. Now, a nail is made of iron. And when it rusts, it's actually combining with oxygen to form a new compound, iron oxide. And again, we'll worry about balancing these later on in chapter 3. For now, just realize that you started off with two different elemental substances and made a new compound. And so you turned one set of substances into a different substance. That's definitely a chemical change. So,
21:29 and it turns out when a nail rust, it actually is undergoing oxidation. Not a word you need to know now, but I just want to point out that the property associated with this chemical change might be called oxidizability, things of this sort. So often times again these chemical changes will have a chemical property associated with them. Cool. One last one you probably want to be familiar with and that's dissolving salt in water. And I just want you to think about this for a sec. When you dissolve salt in water, is that a chemical change or is that a physical change? And I asked this one on purpose because it is a little bit tricky. So there's a temptation to think that oh I can see what salt looks like and when you put it in water it looks totally different. So that must be a chemical change. It must be converting a different property or a different substance. Uh, but it's totally not true. If it turns out if you dissolve salt in water, you still have
22:15 salt in there. And if you evaporate off the water, your salt's still there. So, you'll find it sitting on the bottom of your container. And so, it turns out that dissolving salt in water, we might represent it like so. NaCCl solid turning into NaCCl aquous. Now, it turns out aquous here, Aq is the abbreviation for that. AQ just means that something is dissolved in water. But notice it was NaCCl, a compound beforehand, and it's still NaCCl after you dissolve it as well. We haven't converted into it into a different substance, and therefore, it's not a chemical change. It's just a physical change. That's definitely one that's common to show up on your exam in this section. Cool. That's your big difference between chemical changes and physical changes and chemical properties and physical properties. And these are the most common examples you're likely
23:02 to encounter. All right, last piece of vocabulary we got for this lesson uh is just contrasting two other classifications for properties. So an intensive property versus an extensive property. So and your big difference here is really does the size of the sample matter? So does size matter? So for an intensive property, the size of your sample will not affect it whatsoever. For an extensive property, size of your sample will totally affect it. So, if we look at some examples of this, so your melting point that we alluded to just a second ago, we said that for water, the melting point was 0 degrees C. And notice we didn't say, well, if it's a glass of water, it's zero. But if it's a whole swimming pool full of water, it's going to be some other number. No, the melting point is the exact same temperature regardless of
23:48 the size of water sample you have. And so, the melting point and typically your phase change points in general are good examples of intensive properties. Now, on the other hand, for an extensive property, you might have mass. So, you put a a glass of water on a scale versus putting an entire swimming pool of water on a scale, and those scales are going to read something very different. The mass is going to be very different between those two lovely samples. So, the size of the sample is totally going to affect the mass of that substance. Same thing with volume. So the volume of a glass of water is much smaller than the volume of an entire swimming pool full of water. And so the size of your sample is going to totally affect the volume of that sample. So good classic examples of extensive
24:34 properties. So we'll get a little tricky here now. So turns out that density is going to be an example of an intensive property. And what makes that tricky is based on the definition. It turns out that density, we'll talk about this more at the end of this chapter, is equal to mass divided by volume. And so mass is an extensive property. Volume is an extensive property. But if you divide one by the other, you get an intensive property of density. And you should definitely know that density is intensive. It turns out that the density of water at room temperature is about 1 g per milliliter. It's almost exactly 1 g per milliliter. And it doesn't matter if you have a whole swimming pool full of water or just a glass of water, the density is still 1 g per milliliter. It is not affected by the size of the sample. Another good example of an intensive property that you should probably know is temperature.
25:21 So, if I've got a glass of water that's been sitting outside all day and a swimming pool full of water that's been sitting outside all day and they've been uh they've acclimated to the outside temperature, they've reached thermal equilibrium, we might say. So, well, then they're going to have the same temperature. It doesn't matter that one's a smaller sample size and one's a bigger sample size, something like that. So, temperature also does not care about the sample size. It is an intensive property. All right. So, we're going to start with scientific notation here. And scientific notation is just a really convenient way of writing either very large or very small numbers. And that's something we run into quite commonly in general chemistry. So when you're dealing with like the size of an atom being like a fraction of a nanometer, that's a really small fraction of a meter, uh you're dealing with really small numbers or because they're so small, if you dealing with the size of a
26:06 sample that you know fits in the palm of your hand, you've got such a huge number of either atoms or molecules there that you're going to be dealing with very large numbers. And writing really small and really large numbers can be a pain in the butt, but we can make it much simpler with scientific notation. The way this works, so we'll start with a large number here, and we've got a decimal that's not written that's essentially right there. And in this case, we're going to move that till we've got just one number left of the decimal place. Now, we can write a number in scientific notation with more than one number left of the decimal place, but that's not proper. Properly, it should be written with just a single digit left of that decimal. And so we'll rewrite this here as 4.72. So and then we moved it 1 2 3 4
26:52 five places. So in this case that's going to be times 10 to the 5th power. Way this works is when you've got large numbers, numbers that are much larger than one, you're going to end up with positive powers of 10. But we'll see over here when we've got numbers that are smaller than one, we're going to end up with negative powers of 10. So for this one here, we're going to move the decimal now to the right instead. 1 2 3 places and this is going to become 7.349. Notice we've got just one digit left of the decimal as is proper for scientific notation. And then times 10 the power of and in this case it's going to be -3. We moved it three places but again for a number that's much less than one it should be a negative power of 10. So and these are equivalent expressions
27:38 for the same number as are these two right here. And so couple different ways you can look at it. one. Again, if you got a number much bigger than one, positive powers of 10. Number much much smaller than one, negative powers of 10. So, other way to look at this is if you're taking this number and if you're making this number smaller, notice it went from 472,000 to now a coefficient of 4.72. So, then you better have a very large power of 10 to make it uh to make up for it. So, because it's got to be the same number. So, if your number part gets smaller, your power of 10 better get bigger. Same thing over here. In this case, we made a very small number. 0.7349 get bigger to 7.349. And so if the number gets bigger, your power of 10 better get smaller in this case. Cool. This is as much of scientific notation as I want to cover right now. Uh but basically I want to give an introduction first because this is going to be
28:24 helpful when we start talking about significant figures. So now we've got to talk about significant figures or sigfigs for short. And uh significant figures. A lot of students uh memorize the processes of of how we determine how many sigfigs and how we do certain mathematical operations, but they never kind of understand why we're even doing this. And but significant figures deal with how well do we really know a number? How precisely do we really know a number? And so let's say we take a look here and I'm right here in Phoenix, Arizona. And let's say I've got my friend up here in Portland, Maine. So, and somebody asked me, "Hey, how far is it to your friend up in Portland, Maine?" And I say, "Well, you know, it's
29:11 about it's 3,000 mi." So, and in this case, like, you know, that's a pretty rounded number. You know, is it exactly 3,000 miles like to the, you know, to the T or or is it like, you know, give or take? Well, you know, if I had to round it to the nearest,000, it'd be about 3,000 mi. That's a better approximation than 2,000 or 4,000. So, and that's kind of the deal, but it's not, you know, not that exact in this case. And so, in this case, you know, with significant figures, we often look at zeros as not being significant, uh, as the case would be. And so, this would have one significant figure in the thousand's place and be, you know, as precise as it is to plus or minus a thousand, so to speak. All right. So, what if I said, actually, it's like 2700 miles.
29:59 Well, this makes a big difference. You know, in in a lot of vehicles, this would be the difference between a full tank of gas in getting there or something like this. So, all of a sudden, now I've got a more precise number and it's got significant figures now in both the thousands and the hundreds place. And so, now this number would actually be considered to be precise like plus or minus 100. So, maybe that's where we go. Let me take a sec a step further and actually go and say, "Hey, it's actually 20 uh 2,740 miles to my friend's house in Portland, Maine." Okay, so looking at that now, and all of a sudden now we've got significant figures in the thousands, hundreds, and 10 place. So, and again, that zero is not going to be significant, it turns out. But the way we look at this is now it is exact or
30:44 precise all the way to plus or minus 10 miles. So, much more precise now. And finally, if if I take this, you know, to the absurd extent and I go, it's 2739.1 miles. And all of a sudden, now I've got a very precise number relative to the numbers I've given before. And it's precise all the way to the 10th's place on the other side of the decimal. So we'd say this has five significant figures. And now all of a sudden, it's just a much more precise number. And this is important. You know, if you're dealing with the sciences or engineering or something like that, say you're building a bridge. So, and you start using, you know, rounded approximations on setting up uh, you know, some of the supports on your bridge and, you know, because you're rounding to such a great extent, maybe your bridge is not as strong as you think it is or something like this. And so, this is where the
31:30 significant figures become important. It's how well do you really know your numbers. And when you start doing calculations with them, there's way of ways of propagating these significant figures so that you know how much of whatever you should be doing or how much of whatever you should be adding. So when it comes to chemistry and stuff like this. So but this really is important for science and engineering. So for you in your general chemistry course, you're just going to need to know how uh you know how to process you know determine how many sigfigs and how to do uh say addition, subtraction, multiplication, division and stuff like this. So but this really does have real world importance although that might get lost on you in your course. Uh however this is something you do need to know. It is commonly uh questions on your first round of exams and stuff like this. Uh, and it's also really important in the laboratory. Most professors and
32:15 instructors are going to really ding you on sigfigs if you're not doing it properly uh in the laboratory. So, all right. So, let's take a closer look here at sigfigs. All right. So, the first thing we got to deal with with significant figures are the zeros. So, uh any numbers that are non zero are going to be significant and the pesky little things are the zeros. Sometimes they're significant and sometimes they're not. We've got some rules for this. Uh so the first rule is if your number starts off with zeros those zeros are never significant. So here we've got 0.000047 and the number starts off with all zeros. So and in this case those zeros are not significant and as a result this first number only has two significant figures the four and the seven. So it's got a significant figure here in the one. So notice this is the tenth hundredth thousand in the 10,000th
33:02 place and the hundred thousandth's place. That's where the significant figures are. So any zeros at the beginning of a number never significant. Now notice like this next number 40 5400 we'll deal with that in a second. So but let's say I had 5400 and I just put a couple zeros on the front of it. Well we never do that. But if we did they wouldn't be significant either. So just want to point out that zeros at the beginning of a number doesn't matter where you know which side of the decimal place they are on are not going to be significant. In one case we wouldn't even write them. So, and the reason that becomes important is because now zero's at the end of a number, sometimes they're significant and sometimes they're not. And so, it turns out when you end a number in zeros, so but it's to the left hand of the decimal. And often the decimal wouldn't even be written. Notice there's an imaginary decimal right there. So, those zeros are
33:49 not going to be significant either. And so, here we've also only got two significant figures. These zeros are not significant in this case. And that's going to be different than what we see in the next example because if you end a number in zeros but it's to the right hand side of the decimal, those zeros are going to be important. So not only are the the two, the five, and the four significant, but these two zeros are significant as well. And this is a way of saying that, you know, we know this number really, really well. So 2.5400, it's really precise how well we know this number to the right of the decimal here. But this is like, yeah, it's about 5,400 miles, give or take, 100 and stuff like this. So, however, what if it really was like 5400 miles to the nearest mile and it and you actually meant it like you you look it up on
34:35 Google and it was exactly 5,400 to the nearest mile. Well, the way we might show that in such a case to show that it's significant, there's really two ways. We might put a line across the top of that zero which is going to show that it's significant. That might be one way to do it. So, and that gets to be a little bit of a pain in the butt. So, however, another way to pull this off is to use scientific notation. So, if we really did mean 5,400 exactly all the way to plus or - 1 mile here, all the way to the single digits place here, then what we could do is convert this to scientific notation and make it 5.4* 10. In this case, we moved it 1 2 3 10 the 3r and it's a positive power 10 because we're dealing with a number that's bigger than one. And in this
35:21 case, notice now these zeros are to the right of the decimal. And if you end a number in zeros right of the decimal, they're significant. And so now we'd have our four sigfigs. And so what's nice about scientific notation is that, you know, all of your uh zeros and stuff like that are always going to be significant because if you're in proper scientific notation, you have to have a number left of the decimal, not zero. So you can't even start a number with zeros like we did right here. And if you ever do have zeros, they're always going to be right of the decimal, so they're always going to be significant. And so that's one thing that's nice about scientific notation is all of your numbers are always significant. Now, the last rule dealing with zeros is that zeros in the middle of the number are always significant. And so in this case, what I really should properly say is that zeros that
36:07 are surrounded by significant figures. So that two is significant, the one and the six are significant. and zeros that are in between other significant figures are themselves significant. So in this case, we'd have five significant figures in 200.16. Cool. And those are your rules. So again, zeros at the beginning of a number, never significant. Zeros at the end are significant if they're right of the decimal, not significant if it's left of the decimal, and then zeros in the middle of your number surrounded by significant figures are also significant. So that's your rules for zeros. And you got to be able to kind of identify when zeros are and are not significant. But you've also got to be able to do some basic mathematical operations. You got to be able to do multiplication and division, which is by far the most common mathematical process you'll do in general chemistry. But you
36:54 might also have to do a little bit of addition and subtraction. And it turns out so for propagating like number of sigfigs and stuff, there's rules for logs and other mathematical processes that just aren't going to probably come up in your general chemistry course. So, but multiplication, subtra, sorry, multiplication, division for sure and addition, subtraction probably. And so, we're going to cover those here next. So, we're going to start with multiplication and division and and in general chemistry, that's by far the most common operation you're going to do and have to determine significant figures for and it's the easier of the two as well. And so, the way this works in every term you're multiplying and dividing, if all you're doing is multiplying and dividing a big string of numbers, then you can do this all at once at the end. You don't have to do it every step along the way. And so in this case, I've just got two numbers, but what if I was multiplying like four numbers in a row? Well, then you just count how many sigfigs you had in each
37:39 number. I'm going to do that here. So this first one's got three sigfigs. The second one's got four sigfigs. So, and it turns out that your answer, if all you're doing is multiplying and dividing, can't be more precise than your least precise number, so to speak. And so in this case, your whichever of your numbers has the lowest number of sigfigs, your answer can only have that number of sigfigs. So here, one of my numbers has got three sigfigs, one's got four. My answer is only going to be able to have three sigfigs. And if we work this out, we're going to get 17.97. And so I'm going to carry this to four sigfigs. So, but then I'm going to take this last one right here and use it to round the one right before it so that we end up with a final answer and only
38:25 three sigfigs. So in this case, that seven's going to cause us to round up. And in this case, that's going to take this to 18.0. And that zero ends the number right hand side of the decimal. So it is significant. And we've now got three sigfigs there in 18.0. And that would be the answer here in the proper number of sigfigs. And the idea is that, you know, if you're multiplying a string of numbers and if you know some of them very well, very precisely with lots of sigfigs, but one of them is really much more approximate, well, again, you can only be as exact as your least exact or least precise number. That's kind of the way it works. And we'll see the a similar fashion when we're doing addition and subtraction here. Now, you're going to do addition and subtraction much less commonly. Um, but it is going to show up every once in a
39:11 while because it shows up so uncommonly. A lot of students forget there's even a pro a different process and they try to do the same thing they do for multiplication division. But with an addition and subtraction, the way this works, so I'd recommend adding these up kind of the way you did back in grade school. So line them up a little bit differently. So we're going to have 1540 plus 327.4 plus 0.267. And so if you look at these numbers, so this one, your most exact digit here is this seven all the way down here in the thousandth's place. And so we know this number plus or minus 1/ 1,000th. So for this one, it's all the way to the 10th's place. We know this number plus or minus one/10enth. So and then 1540 here is
39:58 only exact. It's most precise significant figure is that four, that zero is not a significant figure. And so in this case, it's in the 10's place. And so it's only exact or precise to plus or minus 10. And so it turns out your final answer when you add all these together can't be more precise than your least precise digit. So and the way actually I've got it right on the sheet there in this case I say the answer is round to the same decimal place as the most precise decimal place in the least precise term. And so in this case this is the least precise term and its most precise decimal place is in the 10's place. And so in this case, we're going to have to round whatever this comes out to to the 10 place with addition subtraction. So it's not about counting how many sigfigs. This has got three and this has got four and this has got three. So we should have three sigfigs.
40:43 It's not how it works. It's actually you have to look at how precise each individual number is in the least precise number. That's as precise as your final answer can be. And so if we work this out now, so 1540 plus 327.4 plus 0.267 267 equals 1867.7. And truth be told, I didn't actually have to carry it any farther than right here because I need to use this digit to round it to the 10 place. And so in this case, this is going to round to 1,870. The seven means we round up. And there's our correct answer. And
41:28 notice it just happens to have three sig figs, but it's not because we needed to make sure it had three sig figs. It's exactly where it's at because we needed to make sure that the most precise significant figure was in the 10's place based on the numbers that we're given. Okay, now this is the last part here is going to show up much less commonly than any of the rest. And here we've got both multiplication and division and addition and subtraction. And you've actually got to propagate your your proper number of sigfigs all along the way. And so you want to follow your order of operations here and stuff like this. And notice we're going to have to add these before we actually divide the sum of those two by.5. And when we add these, we're then going to have to adjust the number of sigfigs. And then when we divide by the bottom number, we'll then go have to go further
42:14 and adjust the number of sigfigs yet again. And so, oh, and I lost a zero. There's a zero on your handout. So I lost a zero right there. That way we got two sigfigs in that number, not just one. All right. So 4.23 + 7.6 is 11.83. So problem is is that here I haven't adjusted my sigfigs to take into account the numbers we have. And when addition and subtraction I can see that this first number is significant all the way to the hundredth's place. But the second number is only significant to the tenth's place. And so I'm going to have to round it to the tenth's place. And
42:59 I'll use the number right after. In this case being a three, I'm just going to keep it as a point8. I'm round down here. So we can essentially just get rid of that three to get it now only significant to that 10th's place. And now we've got to do the process of division. And so we'll take 11.8 divided by.5 and get 23.6. And in this case with doing multiplication, division, in this case specifically division, now it's all about the number of sigfigs. In the numerator, I've got three significant figures. In the denominator, I've only got two. And so my final answer should only actually have two sigfigs. And so we want to round it here at the three. And so that six is going to cause it to round up. And so the actual final answer
43:47 here is 24. Like I said, you know, having to do both addition, subtraction, and multiplication, division in the same problem doesn't show up too often, but it does have a chance of showing up right here in this first chapter. Now, one thing to note about sigfigs is that with sigfigs, we're going to make a real big point of it in this chapter and on the exam uh for this section. However, much of the rest of the semester, it is not going to be the biggest deal. Sort of. And what I mean by that is that you're probably not going to have a bunch of multiple choice answers after this test anyways. Uh and maybe not until your final exam anyways as well. But on, you know, second, third, and fourth exam, you're probably not going to have a bunch of multiple choice answers that, you know, have just the
44:32 same answer, especially in a different number of sigfigs. It's not usually how it works. So, it's not going to have the same level of importance the rest of the semester, at least not until the final exam, as it does on this first test. But once again, don't forget that in the laboratory, sigfigs is typically a pretty big deal for most professors and instructors. So, uh, even though it's not going to get a huge focus in the course, uh, after this first exam, keep in mind it's still got a huge focus in the lab. All right, we're going to finish this lesson off with a really brief discussion of precision versus accuracy. And this is important because precision and accuracy in the everyday vernacular, we kind of use them interchangeably. So, but in the sciences and engineering, they actually have uh similar related meanings, but they are distinct and you definitely need to know the difference. And so, turns out
45:18 precision is not what you'd think. A lot of people think that again in everyday life, we treat precision as if it's just the same thing as accuracy. And let's talk about accuracy first. Accuracy is how close to the true value you are. That's it. So, but precision is different. It's not about being right, so to speak. Uh but it's it goes more to to the repeatability of a measurement. if you will, and it's how close multiple different independent measurements are to each other, not to the true value. That's what precision deals with. And again, it gets down to the repeatability of it. So, and again, in chemistry, we're probably going to, you know, get this in the context of measuring the weight of something or measuring the volume of something or something along those lines, but it's often really convenient to look at a diagram of a bullseye here uh for the way this works.
46:05 And so if I was testing out a a new bow and arrow, so and if I shot four shots and they all were right in the bullseye right there, well, one, we'd say that in this case, this bow was very accurate because I was hitting the the place where I was hoping to hit, right in the bullseye. But we'd also say it's very precise because all four of the arrows are close together. We'd call this a good grouping in archery. So this would be both precise and accurate. Now, on the other hand, if I tried out that same bow, and unfortunately, instead of hitting the bullseye in the middle here, I hit four arrows right here. Well, that's still a rather close grouping. So, but maybe, you know, maybe the sight
46:50 on the uh on the bow is off or something like this because even though it was very repeatable, which makes it very precise, it was not very accurate because I didn't come close to the bullseye. So this would be an example of something that's very precise but not accurate. And then finally here I'm just all over the board. I'm not close to the bullseye, but none of the the arrows are particularly close to each other. And so this is neither precise nor accurate. Cool. So hopefully this demonstrates well that difference between precision and accuracy. And once again, accuracy is how close you are to that true value. But precision is how close the values are to each other, but not necessarily uh to that true value. They can be close to that true value, but they don't have to be close to that true value. They can still be precise in either
47:37 case. Units and conversions going to be the topic of this lesson in my new general chemistry playlist. And we're going to talk about the metric system and SI units, which are smart. And we're going to contrast that with the American system, which is not so smart. Uh we're going to talk about the the different temperature scales and how to interconvert between those. We'll talk about density as well and and we are going to do a number of different conversions in dimensional analysis along the way. We're also going to talk about the most common mistakes students make and and some of the more common errors made when doing conversions. So hopefully you're not one of the students making those errors. So let's dive in. And the first thing I'm going to do is talk about the metric system and SI units. And SI stands for the system international. They're kind of the gatekeepers on units. And uh the metric system has got a number of units. We've
48:22 got some that we call like kind of the base units and some that are kind of the derive units which are combinations of these base units and and we're not going to learn them all in chemistry. So, if you were taking a physics class, you would probably be on the hook for learning all of them. But in chemistry, there's some we use more commonly than others. And there's a couple that you'll find out that we more commonly use some other unit other than the actual SI unit. So, I'll point those out when we get there. So with length, you want to deal with the meter. Just lowercase m for short here, the abbreviation mass. This one's important because the SI unit is the kilogram. But you're going to find out in chemistry, we are much more commonly going to be using just the plain old gram instead. So, but you should know that it's the kilogram that is actually the SI unit, even though we'll be using grams way more commonly in a chemistry class. All right, time.
49:10 We've got seconds, temperature. So something we don't use very commonly outside of the realm of chemistry and physics and things of this sort is the Kelvin scale. So but really handy scale we call it the absolute scale because it actually has a true zero. So 0 degrees Celsius like it's called zero but notice there are negative temperatures. Well on the Kelvin scale there's no such thing as a negative temperature. When you get to zero that's the complete absence of heat and you can't go lower than that. The Celsius scale you know Mr. Celsius or Mr. Centigrade just said look at water over there. It boils. I'm calling that 100, whatever temperature that happens at. And where it freezes, I'm calling that zero. And that's going to be the basis of my scale. Well, that's fantastic. Um, but it's not a real zero, which it turns out is important. Uh, volume, it turns out the SI unit, this is a derived unit. Now, it's derived
49:56 from the unit of length, but it's uh length by length by length. So, it turns out the SI unit is the meter cubed. And once again, we're not going to use this unit, uh, very often at all. we're much more commonly going to use the liter. So, and it turns out there are a thousand lers in 1 meter cubed. And we're just way more commonly going to use liters when we get to like talking about gases uh towards the end of this first semester topics. Uh PV equals NRT. Uh it's the ideal gas law and we'll be talking about volumes almost always in liters. All right. Energy and work. Turns out we commonly measure that in jewels. If you take a look at your you or your parents most recent power bill, uh typically not going to be measuring the energy in jewels there. Usually we use something called the kilowatt hour
50:42 instead. Well, in chemistry here, we're going to be using the jewel quite a bit. We probably won't be using the kilowatt hour in this class. And finally, pressure. The SI unit for pressure is the Pascal. And if you were taking a physics class, you would use this quite a bit. So, but in this class, you're probably not going to use it a whole lot. maybe not even at all depending on who your professor or textbook is. And much more commonly, you're going to find out that we're going to use things like the atmosphere is probably the most common. So, definitely not the SI unit. And it turns out there are uh you know roughly 100,000 pascals is equal to one atmosphere. It's technically 101,325. Not a conversion you need to know at this stage of the game. Uh but we'll use atmosphere is way more common. We'll use other things like millimeters of mercury or to as well probably even
51:27 more often than we'll actually use the Pascal. So, but these are the SI units. You probably just want to know each of these SI units. And these are not all of them. You know, there turns out there's an SI unit for luminosity, how bright a light is, and it's called the candela. And uh for the most part, we'll never use it in this class. So, but these are the ones you definitely should know. And again, couple key places where we're going to use much more prominently some other unit than the SI unit. Now, the next thing you should realize is the beauty of the metric system. Uh the reason I say beauty is everything's nice lovely round powers of 10 and we kind of identify those powers of 10 based on a certain uh prefix. Uh so terra here means 10 the 12th or 1 trillion. Giga means 10 the 9th or 1 billion. So like when you hear like a terabyte or a
52:12 gigabyte. A terabyte is 1 trillion bytes and a gigabyte is 1 billion bytes. Mega 10 the 6th or a million. kilo 10 the 3r or a thousand cent 10 the -2 or 1/100th 001. So just like 1 cent equals 1/100th of a dollar. So centa actually means 1/100th and it's a little bit tricky because people like Chad a centipede has a 100 legs. Yeah, sort of. I mean it does but yeah mill actually means 1,000th or 10 the -3. Micro 1 1 millionth or 10 -6 nano 1 1 billionth or 109 pico 1 1 trillionth or 1012 phento 1015 ato 1018 so depending on who your professor is they may you know not require you to know these last two or
52:59 something like that or they might require you to know another one at the top or things of a sort. So but I tried to kind of be pretty comprehensive to cover just about any class. So, but these are typically the prefixes we make you memorize. And it's super important when we start doing conversions in this in a little bit here that you know what these prefixes mean. So, make some flashcards or do what you need to do, but you need to memorize these. So, because we're going to use them uh and if you memorized what they mean, it'll make the conversions a little bit easier. We'll find out there's two really ways you can do conversions. One of them is to use these prefixes as they are, which is what we're going to do. But it turns out there's a second way and that's really frustrating because then we start, you know, students will often start uh mixing and matching them. Professors sometimes present both and don't realize they're presenting two different ways of doing conversions. So, we're going to present one way and use
53:45 it exclusively except for one example where I show you why there's another way and why we're not going to use it. Uh so that we don't get any confusion here. But we're going to use these lovely prefixes. Okay. I want to contrast this lovely set of powers of 10 with our lovely American system of units. and our lovely American system of units. You know, you might be like, "Well, there's 12 in in 1 foot and there's 3 ft in 1 yard and there's 1,760 yards in a mile or 5,280 ft in a mile and and all of a sudden this gets really confusing and we can talk about, you know, cups and pints and quarts and gallons and, you know, and there's not a nice lovely rhyme or reason that's going to help you memorize it in every given context." So, but here these lovely powers of 10 apply to any
54:31 of these units and stuff like that. So, a kilo g well kilo means a thousand. So, a kilo g is 1,00 g. A kilometer or kilo meter is 1,000 m. A kilo liter, which you're not going to hear that used too often, but a kilo liter would be 1,000 L. And so, it's nice because these prefixes apply to any unit. Whereas in our American system, if we're talking about length, we have a whole bunch of measurements for that. If we're talking about volume, we've got a whole bunch of different measurements for that. And they're not necessarily related, you know, in any specific set way. And so, it's really annoying. So, if you just want to, you know, chalk it up to you Americans, yeah, we're special and we know it. All right. So, uh, for those of
55:17 you that aren't American and are using that metric system, props to you for being smart. All right. So, we want to do some conversions now and use these. So, I'm going to leave these lovely metric prefixes on the board so we can use them. Uh, but we're going to work a number of examples. And like I said, uh, in the intro here, we're going to show where students make the most common errors along the way as well. Let's dive in. All right. The first conversion we're going to do is converting 75 g to milligrams right off the study guide here. And so, if we want to go from grams to milligrams, first thing you got to realize that you're going, you're starting in the metric system. You're already in the metric system and you're staying in the metric system. And that's a beautiful thing. As long as everything's within the metric system, it's all through lovely powers of 10. And so in this case, to go from grams to milligrams, we'll put grams on the bottom. We'll put milligrams on top. And
56:03 uh couple different ways you can actually approach this. So let's first look at what a milligram is. And a mill means 1,000th. And the reason we say that is because we're going to say that one milligram equals 1,000th or 0.001 g. And another way of writing 0.001 001 is 10 -3 and that's where that power of 10 comes from. So we can say
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