General Chemistry 1A. Lecture 01. Introduction to General Chemistry.

All right. Welcome to Chem 1A so this is the first class in the general chemistry series here so we’re going to go over it we’re going to start with a general introduction to the course and some fundamentals and these are the things that you’re sort of already expected to know when you walk into the class we have a bunch of homework set up on that so you can go through and learn everything you need to know about it in the meantime I mean the covers of the
important aspects of it today and maybe next class so, over the course of the class we’re
going to cover the first four chapters in your book; in the Atkins book so we’ll start off with just a cursory overview of the fundamentals things that you would have learned in high school or Chem 1P and we’ll go over that just skimming the surface of the and let you
take care of most of that at home then we’ll start into chapter 1 and this talks about quantum mechanics
and atoms at the very very basic level and we’ll do lots of material on that and spend pretty much the first four weeks on that
and your first midterm will cover just chapter one from there we’ll move on and start building molecules out of these and we’ll go in to more complicated structures will learn about geometry and we’ll learn about all of the ways that these look and interact and how they form bonds in more detail rather than saying that they just bond together from that point we’ll also learn how they interact with each other, how one molecule can interact with another molecule and the way in which these these can change based on and the
different properties molecules and that will be what your second midterm is on at that point we move on to chapter four which we start learning about these in both properties
so we start learning about gases and how gases interact with each other and and we will do that for the final So, for the entire fundamentals section these are the things that you already sort of know, and I say that theoretically because it’s probably back there somewhere from three
years ago when you took high school chemistry in tenth grade so go back and review it all that’s A-M in your book. uh… For J, K, and L, we’re not going to really talk about that at all, I gave you a little bit of homework on it that’s going to be really important for 1C and if you have me for one 1C I’ll test you on that when we get there for this class, I’m not going to test you on that at all, and I’m not going to talk about it at all. but everything up and through J is your responsibility and is what we will sort of go over in class today so, with all that in mind, let’s start in on this We’re going to start with significant figures This is really important and sometimes gets forgotten about, and this is a way that we measure how precise we can be in any of our calculations. so for every calculation do in this
class you’re going to have to worry about
significant figures So, there’s a few different rules for this the first one is nice. Any time you have a non-zero number, this is going to be a significant figure so every calculation you do you’re going to count up all your significant figures and you’re going to decide how accurate you can be so, for instance, with this one, it says ’54.’ we have two significant figures and what that means is that we can be accurate to this fifty-four place– or, to this
one’s place right here now, when we get to zeros, that’s when things start to get a little more complicated so first we’re just going to talk about what is and what isn’t a significant figure then we’ll go on to how to calculate with them so if you have a zero that’s between to other digits, so
for instance, this zero, where you have a zero between a five and a four that’s going to also be significant so in this case where we have five
hundred and four the five and the four are significant as well as zero that’s between them now if we have zeros that are left of the first non- zero digit so I have some examples here we have zero point zero zero five zero four– something that you’d see pretty regularly, you know, a decimal point, and then this one which is a little bit
more strange and not something that you would ever see really written out but just in case I included it in this case we don’t count either of
these zeros these zeros, I think we can kind of see
whether or not significant these they’re just placeholders we can’t write point five zero four and have it be
the same number as playing zero zero five zero four they’re simply placeholders so they’re not significant if you have numbers are greater than one and you have zeros letter to the right of the decimal
place those are definitely significant if you think about it, you can think about this without actually memorizing the rule could we just write ‘2’? well sure of course we could just write ‘2’ that wouldn’t change the number so there’s no reason to include this
point zero zero zero zero unless you are doing it for the sake unless you’re doing it for the sake of
showing how precise you can be that you can not only just gage to this
one this one place, but you can go all the
way out to here and so those are all going to be significant now when you compare that to up here
remember I said you can’t just write .504, that changes the number so here, these are place holders. Here, these show precision. now, this next part says trailing zeros aren’t significant unless the decimal point or scientific notation is used so for instance here I just have four thousand two hundred. there’s no decimal point after those two zeros aren’t significant and that’s because you can just write ’42’
there and have it be the same number those zeros are just placeholders anytime you have a zero that’s a placeholder, it’s not going to be significant anytime you have a zero that isn’t a
placeholder like for instance here those are now what if we wanted all of these digits
to be significant we can force them to all be significant by putting a period at the end so if we just put a period at the
end or a decimal point at the end then that forces, that tells whoever is
looking at your work you really can measure out to the ones
place, it just happens to be that in this case both the tens and the ones
place is a zero now, what if we wanted just four-two, and we wanted that first zero to be significant there is a way to do that too scientific notation is the way to say
exactly what you want to be significant and what you don’t so in this case we can write out four
point two zero and that says, ‘hey, everything that we
have written in the scientific notation that’s significant’. so if you have it and you have it written out in
scientific notation whatever is there is going to be significant so if you ever get to an exam and it rounds up to something like four thousand two hundred and you say
‘but I need three significant figures’ how can I write that down how can I get
that point on the test Well, it’s going to be by going though and putting it in to scientific notation so that I know that you know that it’s three
significant figures Now, the next part of this is how did you
calculations with these because just being able to count how many significant figures you have is great and, for something like taking a
measurement that would that would be fine but you’re also going to be doing
actual calculations. You’re going to be plugging things into equations maybe something that has three significant
figures and something that has four significant figures and you need to know how many come out at the end so there’s two different sets of rules There are rules for if you are doing addition and subtraction and there are rules for if you are doing multiplication and division for addition and subtraction you use the lowest number of places or
decimal points so you’re not actually going to be counting significant figures as we’ve just talked about on the previous slide for addition and subtraction so for example, let’s say we are adding up .24, so something with a ones place and a tens place and something like
.345 where you have- up to a thousands place you round this to the lowest number of decimal
places so in this case two decimal points or two decimal places, and so you would
round it to .59 you’re not counting significant figures here, you’re counting decimal places so, what if you don’t have a decimal. What if you’re just sitting in the higher numbers? well in that case you just round to the
number of places so here we are all the way out to the
ones but here, we are only at the hundreds so when you round this, you round to the nearest hundreds and so you end up with 4500 so, that’s for addition and subtraction you’re not coming sig figs just places whether that be decimal places or places before or to the left of the decimal point. with multiplication and division that’s
where you’re going to bring in all that significant figures that we did in the previous slide so, for multiplication and division you go to the lowest number of sig figs. so if we look at twenty three we can see that there are two
significant figures if we look at four hundred thirty six we can see that there are three significant figures so when we want to round this we want to round to two significant figures now if you go ahead and plug this in to your calculator for me real quick you’ll see that when it rounds we end up with a one and a zero and you might say, ‘well how do I round that, can I just put ten thousand’ well you can’t just put ten thousand because
that would then have just one decimal-or, just one place, one significant figure and you need to have two. so you would say, ‘okay, well now I have to convert it into scientific notation so that I can put the point zero there So let’s do another one, with division so here we had 453 divided by
3.2 so, sure, we have a decimal place here but that doesn’t make any difference we don’t have to pay any
attention to the decimals we just go through and count how many sig figs do we have so here we have three and here we have two so when we get all done with that, that means that we are going to have two significant figures and so we’re left with 140. now, some rules for how to actually work with these, so these are the rules for coming up with an answer that you’re gonna turn in and that’s fine if you only have one or
one step you can just say, ‘well, okay I’ve added these two
numbers, I’ve multiplied these two numbers,’ whichever and go ahead and put it down on paper. but if you have five or six different
steps you’re taking an answer from one equation, filling it in another equation, now you have to kind of keep track of this. it is not a good idea to round at every single step. that introduces a little bit of an error
every single time and there’s a good chance you’ll start getting answers wrong, especially on sort of online systematic homework so what you want to do if you want to hold out a few decimal
points. To be honest, when I do these sorts of things, I just keep them in my calculator and stick an answer the whole way through. if you want to write them out that’s fine too, just write out a couple points past the significant figure the last significant figure then you might say, okay, well how am I supposed to keep track of what’s going on then. How can I, if I take the answer from one equation and I add all these extra sig figs, how do I know how many sig figs I have there? just put a little underline under the last significant one. and then you’ll see this when I start doing some work, but every time I get done with an answer that I need to use someplace else I put a little underline under the last significant figure so I can keep track of it. so that some kind of hints for actually going about doing this and real problems. so that’s one of the most basic things that you’re gonna have
to know to do every single calculation problem this class now, the next set of things that you’re
really going to need to have a good background on is something called dimensional analysis or conversion factors and these are very much interrelated to each other the idea with a conversion factor is you have something that says ‘this much
is equal to this much something of else so a hundred centimeters equals one meter. That’s a conversion factor now that the way we use this is through
something called dimensional analysis …it’s one of the main ways anyways where you can go through and you can
convert one item into something else using conversion factors the easiest way to do this is actually
go through and do a few example problems and so that’s what we’re gonna do All right. So we have your worksheet that you’ve printed off from the
internet and we start by saying the average speed
of helium at twenty five degrees is twelve hundred
fifty five meters per second and I say, ‘convert that to miles per hour.’ So the way that you always want to think about conversion factors is writing down what we have and writing down what we’re trying to get to and
seeing what you need to do in between so we’ll start by writing down what we
have so we know we have this and we know, at the end of the day, we need to have miles per hour now how do we get between them using only conversions that we know or can easily look up I asked you to memorize all of the metric conversions, I don’t care about the other conversions and I would get them to you. The
metric ones you have to know so if we have meters on top we need to do something to get rid of
meters having something in the numerator is saying
that you’re multiplying so the opposite of multiplication is division so we put that on the bottom and we know that we can get from meters
to miles through using some sort of conversion factor
that we either look up or have memorize. More likely look up and so we’ll put miles up here now we go and we find that conversion factor and we say well for every one mile there is 1,609.3 meters when you look at these sorts of
conversion factors that aren’t nice round numbers it’s not something like in the
metric system we have one intended 100 in 1,000 you want to keep one place more or one significant figure more
than what your actual number that your starting with is so if we think about how this cancels we now have miles per second if we keep
track of all our all of our units we don’t want miles per second, we want miles per hour so now we have seconds on the bottom and we
want to get rid of that so that means we have to put seconds on
the top to cross that out and that we want hours on the bottom so we’ll do this and then we fill in this conversion factor of course this would be one that you’d
be expected to know and if you didn’t remember that at the
top of your head you know how many seconds are in a minute, so you could convert minutes, and you know how many minutes are in an hour, so you can convert to hours from this point now we can look and we can see that our meters cancel we can see that our seconds cancel we can see that we have miles per hour so our units are all set which means our answer is going to be fine and we can do 2807 so you can see that thats much easier to figure out how to do
that to try to memorize “well am I multiplying by this conversion factor” “am I dividing by this conversion factor,” what am I doing. so this way you can just trace your units
around we have one more of these to do So, how many minutes does it take light from the sun to reach Earth given that the distance from the sun of
ninety three million miles and the speed of light is three times ten meters per
second or three times eight meters per second so too did this one, now this one I’m kind of asking you a real question, but it really turns out to just be a bunch of conversion factors put together if we know that we have a certain amount
of miles and it’s ninety-three million which
means it’s ninety-three times ten to the six now at the end of this we want to get to time So, we need to trace through all the conversion factors that we can find in order to get this well we can start by saying we have we have miles here and we know our speed
is in meters per second so we have a distance here and have a
time that’s gonna be a good way that we can
go through and try to get to time but it’s not going to work with miles so let’s change miles into meters- and we are going to do that the same way we did here now notice on the second problem though
if you look back to your first one that ordering is flip-flopped and that’s because we’re converting from
miles to meters instead of for meters to miles so we can write in those numbers or you could also hold off and write in all the numbers at the very end, too. So now we know that we have meters on top and our miles cancel out and we can look at what else we can do.
We want to get from a distance to a time so, we need something that has both
distance and time in it. from there we need to decide, ‘well, are we going to multiply by this number, or are we going to divide.” so again we trace through our units you have meters on top and we don’t want meters
so we divide by it so that leaves us with seconds on top which is what we want because that will end up with our time unit so now we fill in our numbers the number here goes with the meters
because that’s on top so 3.00 times ten to the
eighth and we put one second here now at this point we want to look at what I
had asked you for So, we can do this in, um… seconds if we wanted but I had actually asked for minutes and so at this point you can see, well, we have seconds so we need to convert that into
minutes and we know that we have sixty seconds in one minute so that gets rid of this unit leaving us with minutes which is what we wanted and we can solve for that. Now it’s time to take a look at our significant figures for both of these. So, now we have them both up Now, notice on this one I left it as four significant figures. Now we need to go back and say ‘okay, for all of these calculations did we do this right.’ we started with four significant figures here we divided by 5 here and we did that on purpose, right? We
looked at that number and we said okay, we have four significant figures here, I
need to keep one extra now what about this one that 3600. Does that mean that we should round to 28 or twenty eight zero zero Well, no. whenever you have a definition you don’t round to that so there is exactly sixty seconds in
one minute there is exactly sixty minutes in one
hour there is exactly thirty six hundred second in one hour; it’s defined that way so you can think of it as being infinite significant figures, however many you needed there to be so that means that this is going to be
left as four because we started with four now when we come down here we started with two here we again looked this up I put it as five basically just because I
had that number handy But you could have rounded it to three if you had wanted we have this number which has three
significant figures and then we have this one, which looks like one, but remember it’s a definition and so because it’s a definition it’s actually insignificant and so the only thing that matters for our
significant figures is the ninety three so we have two significant figures here and so we need to have two significant figures here. So that sort of walks us through some of the calculations that we’re going to be doing in… it’s a very general format where you don’t need to have a lot of chemistry background yet You can always use dimensional analysis, and looking at the units and crossing off the units is a double check whenever you go to do anything so every problem that you do you should write out all the units all the
time and each time go through and look at
where they are look at how they cancel and make sure at the end you have a unit that makes sense if you’re measuring distance and you
come out of the unit of time you did something wrong if you’re measuring velocity, and at the end you need to have a velocity and you come out with something like meters and no per second, you did something wrong so this is a great way to go through and make sure you did everything okay. We’ll get to times later on in the quarter where we have a constant that has tons of
different types of ways or writing it in and all of the differences are with units. How
do you know which one to use? you can memorize it and you can say, ‘well when I use this equation, I’m going to use this version. When I use this equation, I’m going to use this version. or you can just look at the units and say
okay I know I have liters and atmospheres
so I’m going to use this version of R or I know that I have jules so I know I’m going to use this version of E. Things of that sort. So now we’re going to get into the structure of an atom. And we are going to do this at a really basic level right now, and in chapter one we’ll get into it in much, much greater detail. and in some ways I’ll tell you that
we’ve lied a little bit here. so this is the Bohr model of an atom. And I sometimes like to call this the ‘high school model.’ This is sort of the first model that we teach you. it has some very good uses, some of which
were going to take advantage of here, starting this class through next but, it also isn’t 100% accurate. But it’s a good starting point. So, we’ll kind of start from here you can think of it as being a nucleus in the middle that has two different parts in it. It
has protons and it has neutrons the protons have a positive charge the neutrons have no charge it all and then around this nuclei going… right now we’ll think of it as rings and later on we’ll expand that a bit then you have electrons. And those electrons are negatively charged so there’s some things you want to be able to
look at a periodic table and calculate pretty quickly without having to put a huge amount of
thought into it. so if you don’t know how to do this, it’s
fine but you want to get some practice at it so if you have the number of protons you have and you subtract the amount of electrons,
your protons have a plus one charge your electrons have a minus one charge so that’s going to give you the charge of the ion now in any sort of neutral compound, or atom, this is zero because you’re going to have the same number of protons as electrons But, you can also start adding and subtracting protons and electrons as time goes on and making ions out of them. and we’ll talk about that in much more
detail, too so keep this in mind that your charge is
always going to be equal to your protons minus your electrons now if you look at your periodic table
there’s a bunch of different numbers that you can you can deal with one of them is your atomic number; that’s
your number of protons and then there’s also an atomic mass An atomic mass is your number of protons plus your number of neutrons and so you can take those and you can add them up and that gives you your atomic mass because you’ll always have a periodic table
you’ll always know your number of protons you’ll always know your atomic mass for any sort of exam or anything like that and what you’ll see is a lot of times we can calculate and figure out how many neutrons we have We can take your mass and subtract your protons and get your neutrons. and we’ll do that a lot in 1C when we start getting into nuclear chem. now you really have to remember these charges, which ones are positive and which ones are negative and in order to help remember that we
have a little joke lots of great and simultaneously horrible
chemistry jokes out there this is one of them so, you have two neutrons they walk into
a bar and they order a couple of drinks and as one is about to leave, the waiter says, how much does it cost? And the neutron says… let me start over. Okay. A neutron walks into a bar and orders a couple of drinks. As she’s about to leave she asked the
waiter how much and the waiter replies, “For you, no charge.” So that’s the joke to sort of remember what a neutron is we have another one here and this one is
one of the most famous ones that gets repeated over and over and over again and, so you have these two different atoms and they’re talking to each other and the one says, “I’m hit! I’m hit!” I’ve lost an electron! And the other one says, “are you sure?” And the first one says, ‘I’m positive!” They lost an electron, and so they are positive. Right? so, two nice, horrible jokes for you to remember these by. We’ll have lots more of these as the quarter goes on. Okay. So now comes some time for just general definitions. So first of all, we have something called an isotope. what is an isotope? So… any time that you have the same number of protons but you have a different molecular mass. that’s an isotope. And the reason why you get this is because you have different numbers of neutrons and so your neutrons within one particular atom can change and without really changing too much of
the properties well see again in 1C that some of the properties definitely change the mass definitely changes because you are adding in a neutron which has a mass of about one But, most of its properties are very similar. now why are the molecular masses on the periodic table decimal points? so you should probably always have a periodic table handy in this class just kind of sitting out starting next class you probably want to do that. They’re around. So… whenever you look at this you’ll see that your molecular masses
are decimal points. and why is that? well the reason for that is that they’re actually going to be an
average– they’re going to be an average of both or all of isotopes of that compound or that atom. so something like silver has two isotopes that make it up so sometimes you’ll get silver 107, if you were to weigh the mass of that one atom you’d come out with a unit of 107 sometimes it’s 109. so on the periodic table, what they do
is they do something called the weighted average and weighted averages are a good thing
you know how to do if you don’t know how to do that, review it
in your math class– it’s also how your grades are figured out so, you know when you go to figure out your grades and I say figure out a weighted average that’s what I’m talking about and that’s how it’s
determined here So, in this case we would have 107 silver making up fifty one percent 109 silver making up the rest We’ll do this on the dot-cam in just a minute you’ll notice some of these are the even
more complicated. Some of these will end up with two or three different isotopes now, the reason we do a weighted average as opposed to
just averaging it, you may say why can’t I just take this and say, well 107 plus 109 divided by two Well, we want to know what the mass of this is– if we go out into the world and we take some silver out of the ground, clean it up, get rid of all the ore, and purify it how much is that silver what’s the molecular mass of that silver and not all of the silver is split 50/50. 107 and 109. and this is the idea of where weighted averages come in. so we’ll do this one out on the document camera So we can see all worked out Okay. So we have 51.839% of silver is 107 so we need to figure out
how much there is of each so we know this because the problem says so now we also need to know what percent is 109. well, it’s the rest of it so we just take a hundred subtract that so this is a hundred minus the silver 107 percent. which gives us the 48.161 now we do what we call weighted average so I’m going to do it out in two steps you
can do it all in one you’re fine either way so if we know that we have 107 grams of this we can just say we have a hundred and
seven grams per mole and 109 of this What we’ll do is we’ll go through and we’ll multiply that by the percent and the same thing here… so we’ve taken 107 multiplied it by the percent that makes it up in nature this multiplied by the percent that makes it up in nature… and we get those two numbers and then we add them up and we get that. now, with everything in chemistry, you want to be thinking, “does this make any sense?” now in this case we have weighted
average between two things, it’s 107 and 109 and it’s about fifty
fifty, right? one’s fifty one one forty eight about fifty fifty so we would expect the answer to be
close to what the normal average would be or what the real average is which is 108 so since we have this and this being added together in almost equal proportions we want it to be close to 108 with a little bit less because the 107, the lower number, has a little bit higher percentage and that’s exactly what we see. We see
that it’s close to 108 just a little bit less than 108. so that makes sense based on the averages
and what we know about how averages work and so that’s done. now, this is a case where we only had two isotopes that we’re averaging, you could do this for more, something like carbon has one main one and then two smaller ones, you could do that for each and you would just do this three times and then add it all up. So now that we know all these things about atoms– we know their protons, we know their electrons, we know their neutrons their masses we need a nice way of looking at all this data and figuring things out very quickly and this is where the periodic table comes in. we had all the elements arranged an
order increasing atomic number so atomic number, remember, that’s the number of protons that we have. now they are arranged in these repeating patterns and we’ll get into more detail about
exactly how that is but for right now what you can go is
that they have the same number of what we call valence electrons- outside shells. If you think about the Bohr model you can kind of think about those rings, right? And whether they’re filled or how many they have in those outside rings. and so what happens there is that gives certain columns, or groups similar properties so everything that’s going to be in group one
is gonna have a relatively similar property to each other. Know of course
there’s going to be differences because The atomic mass here is much, much bigger than here- the number of protons and number of electrons are much bigger and there’s some trends that we’ll be able to pull out of the periodic table later on but for the most part this group could have the same sort of
trends or properties as each other This group, same sort of properties as each other, all the way across the periodic table so if you go down a column you have very
similar properties now this happens to be my favorite periodic table that I carry around and have in all my books- I actually replace a lot of my book periodic tables with this one you know, find your favorite periodic table from the internet, print it out, and keep it with you all the time. In class and when you’re doing chemistry. I don’t really expected going to the bar with them and such things but keep them around you whenever you’re doing your homework keep them out in class with you because I’ll refer to them a lot. Okay. Now comes something that we’re going to get into a little bit with naming and this is probably all of the freshman chemistry student’s least favorite part about this class. because there’s a lot of memorization in general I say with chemistry you
shouldn’t be memorizing hardly anything if your memorizing things you are, in general, not learning them. And in chemistry that gets dangerous. If you memorize how to do a problem, you are probably going to have problems on an exam because I’m going to give you one that’s a little bit different and if you don’t really know what you’re
doing you’re not going to be able to solve it because it won’t be the exact same as your homework problems. This is sort of the exception to that. you have to do a lot and a lot of
memorization for the ionic naming. it’s a pain just do it You’re going to need it for 1B, you’re going to need it for 1C, and you just need it to be around chemistry in general which includes biology that you’re going to be in. you want to have a good idea of
what’s happening and you want to be able to look at a compound and name it quickly without having to think about it too
much I can test on this in the first midterm
just by saying here’s a name, give me the formula. Here’s the formula, give me the name. I can get on the second midterm by
saying, ‘draw me this compound’ and if you don’t know how to name it you
don’t know how to pull the formula out of the name I give you, and you won’t be able to do it. so make sure you can just go through and do all of this. so before we can get into naming too
much we have to figure out how do we know
what type of naming we’re going to do. So up here I have Ionic, Molecular, Acids, and Organic. now we’re gonna get very much into ionic,
molecular, and acids– those are the ones that you’re testable on in this course for organic you have a homework
assignment on this we’ll talk about it a little bit don’t worry about it too much in general
that’s all going to be covered in really great detail next year but you should have a general idea of
how it works because I’m gonna talk about it. I’m going to say things like methane and ethane, and, you know ethanol and propanol and you should have some idea
of what I’m talking about I’ll also always have the structure there but you don’t want something like that to throw you off Just have a general idea of how it works be able to answer questions on it if you
have the book in front of you we’re just not going to get too into it this quarter or this class that’ll be more for next year these three are the ones were going to focus
on now all of these are going to be named
very differently if you have an ionic compound, it’s named completely differently than a molecular compound You’ll learn to like the molecular ones
for the naming purposes here And acids are going to be based off of the
nomenclature but it’s still very different and so before we can actually get into
the rules for naming we have to get into how we know which one is which type of compound to do this we have to talk a little bit
about bonding and how things bond so whenever something is
trying to bond with another atom is trying to do what we call ‘complete an octet’ now, we’ll see some atoms don’t actually do that exactly, but they’re trying to they’re trying to get a full octet
which means that they would have eight electrons in their outside shell so if you have one atom that has six in it’s outside shell, and another one that has six in it’s outside shell, it can go and it can form a bond. Now in this case, it’s going to want to share those electrons. because you have six here and six here, so this one can’t just give two of it away or it’s going to have some problems. This
one can’t just give two of them away; it’s going to have some problems there are times when you can do that- when you can just trade electrons and you don’t have to share electrons and that’s the difference between your
two different types of compounds your two different types of bonds so, for ionic compounds that have ionic bonds they’re going to trade electrons one atom is going to give away it’s electrons to the other one. so something like sodium chloride So, if you take and you look at your periodic table for a minute, and you look at sodium, right here, and you look at chloride right here, you can see that sodium has one electron in it’s outside shell and the periodic table is really nice for looking at this quickly, because you can look here and say okay, this has one valence electron. This has two valence electrons. Three… Four, five… Six, seven and eight, all the way across. So we can use this to look and see how many valence electrons we have, very quickly. So, sodium has one, and chlorine has seven. So, they can just trade electrons. The sodium can say, well, I don’t really want this one electron sitting there by itself, you take it. And it gives it to the chlorine
and the chlorine says, ‘great! This gives me eight,” So now they both have eight. In covalent bonding, they’re sharing the elections so this would be something like carbonate oxygen or two oxygens or two fluorines, where they don’t, they can’t just give them away. They’d still be too short of electrons, and so instead they’ll share them. Carbon will let oxygen take some of them Oxygen will, you know, lose some of them and then you count the electrons for both. We won’t get too into metallic bonding, But it is important to talk about it and have it sort of in the background now in metallic bonding you have these big networks. in these big networks of electrons that can move back and forth between all of them here and here you sort of have the
electrons that are relatively associated with one or
two groups here they’re completely de-localized which means that you can kind of move them from one side to the other if you do the right sort of thing to it, and we call
this a wire, right? We can take a wire and we can stretch it out, that’s all metallic bonds, we can put electricity through it and the fact that you have these de-localized electrons that are going across this entire group is what allows that to happen. and those are for metallic bonds but these are the two we’ll be focusing
on for the sake of naming. and one last thing we need to talk about
here is something called empirical verses
molecular formula now, this comes up in covalent bonding not in ionic. so keep that in mind; you might want to even write that down. this is just for covalent issues. With ionic compounds, we’re always going to list them as the lowest whole number ratio So we would never say, MG2O2. or, NA2CL2. we always want to reduce them down to the lowest amount. With covalent bonding, that’s not necessarily how things work. Something like hydrazine if you look at this we have N2H4. and you can say, ‘well, can’t I reduce that down to NH2?’ Well you can’t, because it changes the whole compound. N2H4 is not the same thing as NH2. so we have to have some nomenclature for this
that we’re going to refer to from time to time so we have an empirical formula, which is our lowest whole-number ratio. It doesn’t actually tell us a lot about the molecule itself. but it does tell us how many of each atom are in the substance. and then with the molecular formula it’s
going to tell us it’s going to go through and say well this is what the molecule actually has in it. One molecule of hydrazine actually has two nitrogens and four hydrogens or ethane; two carbons and six hydrogens. now, we’ll do some examples using this where you can see that we can find the empirical formula fairly easily experimentally and we can find the molecular mass fairly easily experimentally which is why it’s one of the reasons that
this is so important to have these differences here so that goes into a little bit more of covalent bonding definitions now we’ll spend some time on ionic bonds and then we’ll learn how to actually name them. so this is the day filled with bad jokes so we have another one so the way that ionic bonding works is by the fact that these atoms will trade electrons and when you trade electrons, electrons have a negative charge and so there’s going to be a charge development if you trade electrons like that So, we have this teacher up here and you have these little atoms here, all these positive ions, It says, “perhaps one of you gentlemen
wouldn’t mind telling me just what it is outside his window that you find so
attractive.” so remember, positives and negatives are always going to attract each other. so if you take an electron, you take it
away from one atom, you make it a positive charge. You’re taking a negative charge, away, you’re making it a positive charge. You’re giving it to another atom, which means that that atom is going to be a negative. so you have a positive atom now and a negative
atom those two are going to attract just like a
magnet would. And that’s how you form your ionic bonds. so how do we actually write this out
quickly? Well, we can do this this way. where we have ionic bonding here, we take something like potassium, and if you look at your periodic table you see that potassium has one electron. it’s in that first group, and so it has one valence electron if you look at iodine, so you’re looking at your periodic table, you’re finding it on the periodic table you see iodine is in the seventh group so that means that it has seven electrons so how can we get this so they both
have a full octet well, we’ll take the electron away from potassium we’ll give it to iodine when we do that potassium develops a plus charge. You’ve given away one of it’s electrons Iodine has developed a negative charge because you’ve given it to iodine, and so you get this structure where you form the positive and the minus, they attract, and you form potassium iodine the exact same thing can happen where now instead you have to give one electron away to two different atoms so if you’re trying to combine something
like magnesium along with something like fluorine, now you have an issue where you have two electrons in magnesium’s outside shell, and fluorine only can take one well, you just take double the amount of fluorines. So now magnesium says ‘okay, I’m going to give you one electron and I’m going to give you one electron.’ both of the fluorines develop a
negative one charge the magnesium gets a plus two charge
and they all attract each other and they form this structure which means then you would have MgF2. one Mg and two fluorines. now that’s how you want to know what’s going on how these are trading whether they just trade one and become K+ and I- in this case Or sodium chloride would be the same same form or whether you have MgF2– it’s all about the charges and making sure that the charges balance now there are quicker ways of doing this, though, than trying to write this out each time and think about exactly where the electrons are moving around each time so here’s a helpful trick to remember it so if you can kind of take the chargers and you criss-cross them. now that you might be saying, ‘how do I
know what the charges are?’ basically through memorization, and we’ll get
into that just a moment if you write down the charges here and the charges here and you know that your ions have these charges you can criss-cross them down. You can say, ‘well, I’m going to move this down to the oxygen I’m going to move this down to the
aluminum.’ now what that ends up doing is it gives you a compound where your charges balance you can say okay aluminum has a +3 charge and I’m going to multiply it by two. Now that means I have +6. oxygen has a -2 charge But there’s three of them. So minus two times three that’s minus 6. so now you add up your plus charges and your minus charges and they need to equal zero and in this case you see that they do so this is sort of a helpful trick for
getting you to this neutral compound faster you want your ionic compounds to be neutral. there is one major caveat to this
that you have to watch out for if you’re going to use this little
quick trick and that goes back to this idea of empirical formula and molecular formula and how that’s only
true in covalent; we only deal with that in covalent. Ionics always need to be the lowest whole-number ratio. What if you had something that had the same charge, and it isn’t one or, same magnitude of charges, I should say. and it isn’t one. something like magnesium that has a +2, oxygen that has a -2 you criss-cross them down you form Mg2O2 That’s an ionic compound. You have to have the lowest whole number ratio. so this isn’t okay. So you need to then reduce down to the lowest whole number ratio. So when you reduce that down it becomes MgO so at that so whenever you have this sort of situation you can probably just say ‘well, these are the same charge’ So I’m going to say that that’s just MgO but if you didn’t catch that right away
and you did do the criss-cross trick and you saw that this is Mg2O2 you have to reduce that down something that we are going to just sort
of show you and then move on with is that these form these big, crystal lattices in 1B, when you first start 1B, this is what you’re going to start with You’re going to learn all about these
different shapes and these different forms and you’ll put names to them, and all of that and for this class, I just what you know
that this exists. that when you get these you don’t have
one sodium chloride just stuck together and you have this one little atom with sodium chloride all by itself what you actually have is you have these
crystals If you go to your pantry and you pull out salt you actually have all these little
crystal lattices inside and that’s what’s forming. so just keep that in the back of your head that this is what these ionic compounds look like. okay. now how do you know what the ions are? So I’ve said that it’s memorization, and that wasn’t
really a hundred percent true. so you can look at the periodic table
and you can find out what the charges are for the most part. There’s some
exceptions here but for the most part so if you look at this first row you have one valence electron And you want to find a way to get rid of that so you can go through and you can just take that one electron
off giving it a plus one charge or you can take two electrons off from
this group giving it a plus two charge this group you’d have three valence electrons, so you take three off when you start getting into this group
now it’s not really going to be able to just gain two or lose excuse me, gain four or lose four as easily, those aren’t going to be big on forming ionic compounds at all When we are to the right side of the
periodic table we look at here how many valence electrons does that have? well, it has seven. what’s been easier? Taking away seven electrons or just adding in one well of course adding in one would be easier and so because of that you’re adding an electron, you’re adding in a negative charge and so it’s going to be negative one if you look at this row you have six valence electrons it’s in group six, so you have six, is it going to be easier to pull off six or add in two? it’s going to be easier to add in two, and so you get a -2 charge same thing here-easier to pull off five or add in three We’ll add in three, so that’s going to be a -3 charge. this section in here– your transition
metals in this little group right here for right now, you pretty much just have to
memorize them after we start talking about electron
configurations in more detail and how to go about dealing with those and seeing
where electrons are removed from we’ll actually be able to explain most of
them however for right this moment there’s
really no way to explain it easily so these groups, you just sort of have to memorize them. and when we get into the very end of chapter one we’ll learn why all of those are now, I do want to help you memorize one of this section so this little group right here, we’ll pull out and we’re going to look at it in more detail. There’s something called the inert pair effect, and I don’t want to go in to exactly why this is at the moment– if you’ve had a lot of chemistry and you want me to explain it I can do that later– but for now we’re just going to leave it as “this is how it works” and then at the very end of chapter one we’ll go back and we’ll explain it using electron configurations so for right now you can notice though, that these have these groups of one through three two four three five so they’re always off by two you can form this top lower ion or you can form this bottom ion and they’re always off by a factor of two. so use this to help you remember it for right now. and we’ll explain why it is later on. now we’ll finally be able to get into naming these so, for here, a lot of these listed– there’s copies of all of these sorts of things online– this is from a different book– but we’ve pulled out all of the ions and put them online for you so The idea to get out of this, and the list online, is that there’s also all of these polyatomic ions that we need to talk about. Those for the most part are going to be memorization uh… but there’s some hints for
memorization So everything I’m about to say we’ve also put into an
online study guide that is available so that you can kind of go through and see what i’m saying when I say -ates and -ites because it’s a little bit hard to hear So these are lists of things that you effectively need to have memorized but there’s tricks to memorizing it
that make life a lot easier so there’s these two these different types of endings and these different types of prefixes so you have -ates, -ites, and -ides as your endings if you have something with and -ide ending, that is just the compound– or the atom, on its own. something like phosphide or sulfide
oxide if you have something with an -ate or an -ite ending like this that’s going to be your oxygens so something like phosphate or phosphite are going to have oxygens on the end and the -ite and the -ate refers to how many oxygens, or its oxidation numbers so, what we’ll do here you should memorize one version of these you memorize all of the -ates, or you memorize all of the -ites, you don’t memorize both if you have something like with an -ate ending, that’s what I happened to memorize when I did this back in high school or college and I memorized through and I memorized phosphate and carbonate and sulfate and then I knew that every time that I changed that to an -ite ending, I just took away an oxygen. so if phosphate is PO4^(3-) phosphite is PO3^(3-) so that’s kind of your hints for doing
this now you also see times where you have
something like, if you look at iron, iron has two different oxidation states; two different charges so a 3+ and a 2+ and we have different nomenclature for
certain compounds for instance, iron. Where we call it ferrous or ferrite and those -ous and -ic endings those refer to the charges if you have an -ic ending, you would have a 3+ charge if you have an -ous ending, you have a 2+ charge the -ous always refers to the lower charge the -ic always refers to the higher charge and so keep that in mind what you’re
memorizing those weird named ones like iron and all lead is another one uh… stannous and stannic for tin so there’s these groupings of ones that
have two different names one comes from the original latin route the other one is the name that we know
it by and you’re pretty easy to pick out of the periodic table because they are the ones where the element symbol doesn’t really match up with what we know we know that iron is Fe We know that tin is Sn; and so those are pretty easy to pick out of the periodic table and i have those highlighted on
the site as well so make sure you go, and, start working on memorizing those and start working on making sure that
you can form those ionic compounds using the ions that you have memorized because, again, you’ll be tested on this all through general chem it’s better to just do it now and get all the points starting now rather than waiting ’till the end to it Next class we’ll get in to how to name acids and how to name covalent compounds which are significantly easier than this
just because there’s a lot less memorization to it If this didn’t make a huge amount of sense, make sure you go and look at the study guide online where it’s all written out with the -ous, -ate, and -ite endings. to help.

95 thoughts on “General Chemistry 1A. Lecture 01. Introduction to General Chemistry.

  1. It is because the question was asking for units in minutes not seconds. So you have to remember that one last conversion at the end.


  2. Doctor Brindley,

    It has come to my attention that I believe there is a mistake on the powerpoint slide at 43:39. On the fluoride (F) at the farmost left, there are 6 valence electrons when there should be 7.


  3. There would not be a leap day if time is exact and defined. Time can't be confined to human terms. I believe so atleast. I may be wrong

  4. I don't have access to your website Dr. Brindley, so with that can you post the sheet you referenced that you highlighted Fe, Cr, Mn. All the Atoms with Latin names.Thanks

  5. She does not explain the math computations for the questions after canceling seconds and meters this is where the student gets lost…

  6. Got an issue with ionic bonding termed a "trade" of electrons. If one gives and the other receives, isn't that a donation rather than a trade? 

  7. im learning chemistry online, is there anything that came before general chemistry 1A? or you start from here?

  8. i think some of the math is wrong? .51839 x 107 = 55.46773, which then leads to the the right answer she wrote down for adding the averages together.

  9. ¿Éste es un curso normal de química en la Uni en los estados unidos? Se ve medio aburrido y seco y hasta tal vez atrasado dependiendo del tiempo que van en el curso. Por eso quiero ir de intercambio a Alemania en vez.

  10. It would have been helpfull if she explained WHY the intro math is significant and why one will need to care about such details for chem. Without this proper roadmaping students will spend too much mental energy guessing at the end point of all of it, and that mental churning takes away from understanding the details.

  11. She is a mediocre teacher… and you all (at least the ones that left a comment) are ignorant math-wise: Nobody commented on the wrong answer of 107.97 g/mol (timeline 32:16)… the correct answer is 108 g/mol… as the lowest/final number of significant figures is/should be 3. The irony is that she taught "us" that on the very first slide and didn't apply to her own math/problem!!!

  12. WOW. No, she is not mediocre, she is bad. There is no value added to just reading a chem. book on your own. She made some more mistakes besides the one I mentioned below… she cannot distinguish between "a factor of x" and "a difference of x". How did she get the PhD and the teaching position?!

  13. She gets the first answer by 1255 x (3600/1609.3) = 2807; going to 4 sig figs from 2807.4318 for anyone else who was totally lost like me

  14. Professor Brindley, is the Atkins book you talk about available in full at Google Books or is it just a preview version there?

  15. are there any tests or exams in this course, like a link or video to any? I can find some googling sure, but i want ones specific to this course.

  16. How is 23×436 = 1.0×10^4, 4 sig figs, since she earlier stated that 4.20 x10^3 is 3 sig figs…? Is the sig figs from the actual decimal number or the exponent's value?

  17. Hi guys! Im currently in Grade 12 and taking Chem. My grade at the moment is a 95. What I've recently been doing is when I have some extra time, ill go and watch these lectures and take note on what I'm learning so I can prepare for University next year. Do you guys think this is a good idea or should I not do this?

  18. hi kate……u r tying is good bt if u want interest in chem…i think u have to see the origin or begining of chemicals on earth or in universe…..before study book…
    good luck

  19. hi kate……u r tying is good bt if u want interest in chem…i think u have to see the origin or begining of chemicals on earth or in universe…..before study book…
    good luck

  20. is this really how a chemistry lecture is like. you just finished 2 chaptes of information what are you going to do next time you see the class. are you going to do 1 chapter every lecture. in that case you will be done with chemistry in only 3 months

  21. It's nice you shared the audio with us, but the text is unreadable to the online audience without a scanning electron microscope. Come on AV crew? Take some pride in your work. We need this information out here in poor people land. Better off watching Tyler Dewitt's channel He is "DA MAN", DA CHEMISTRY MAN" >>>

  22. I was excellent in chemistry. I remembered being the top in my class back in high school. I use to love chemistry and read so many textbooks on the subject. Easy A without studying. I wanted to do my bachelors in chemistry back in 2007 but my parents, seeing me good in chemistry, forced me into chemical engineering. I failed the first semester hard. Ended up switching to business and now I work as a banker. I still remember my chemistry quiet well. Also in college, we had to take a course in chemistry and I got an A without attending any class and NOT even using the book. Just sat for the quiz and exams and wrote whatever I remember. Got an A (100/100) not even a 99/100. If I can go back in time, I'd study chemistry but been working for 6 years and almost turning 30. I feel it is late for now.

  23. I’m pretty sure the video would’ve been less than 55 minutes if she didn’t talk so much in between… so much time wasted on small talk

  24. 31:00 58.47+52.50 does not equal 107.97, which is still inaccurate as my periodic table says 107.87. Nothing about that example gives me confidence. And why choose 107g and 109g to represent Ag 107 and 109?! That is confusing.

  25. Thanks really appreciate this kind of lecture. No history, no stories, no bs straight to the point.

  26. A normal conversation

  27. Classrooms either go to fast or too slow. They're not for learning. They're IQ tests in disguise. If I needed this, I couldn't use this. If I didn't need this, I'd merely get a high score and they'd know I was talented and I'd get a job in the field. What's the point of the video. To "help" people with their homework? If you need that, you're in the wrong major. If you're studying for the sake of self improvement, you need a book and/or a tutor where you ask questions as you go. A vid like this or the class is useless for that.

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