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.
