Good morning we're going to get started.
It takes me a while to get used to being back
in school again, I left my pens in my car
so hopefully I won't do that again.
Alright, hope you had a great weekend.
We left off last time talking about bonding.
We have two kinds of bonding: we have ionic
bonding and covalent bonding.
We were talking about ionic bonding - are
there any questions before we get started?
I want to see a quick show of hands I like
to on Wednesdays try to do some, integrate
some active learning activities into our class
on Wednesday just to see how it goes.
I want to try a bunch of different things
and then I want to give you a survey at the
end to see what you liked.
First of all I want to see a show of hands
of anybody who does not have an iClicker.
You all have them?
Ok so we're going to be trying a couple of
different things.
There's some programs now where you can actually
draw compounds on your smart phone - is there
anybody who does not have a smart phone or
an iPad or a laptop?
A few of you.
So what I would probably have you guys do
is work with the person that's next to you.
And we'll do some group activities and as
long as we stay on time in the class we'll
continue to do that every Wednesday.
Alright -- ok so let's get started.
Alright so we were talking about ionic bonding.
Atoms obtain a full shell by transferring
electrons from one atom to another.
So there are classic examples is lithium fluoride.
So lithium transfers its single valence electron
to fluoride, that makes lithium happy because
now it has a full shell and it makes fluoride
happy because fluoride has a full shell and
these guys are held together by electrostatic
interaction in a crystal lattice.
That's your actual ionic bond.
The electrostatic attraction between the positively
charged species and the negatively charged
species, that's your ionic bond.
Ionic bonding is seen with atoms of widely
different electronegativities and so back
in g chem you learned metal and nonmetal and
there's other things that are not metals and
nonmetals so that's not strictly true but
the reason that's true is because the electronegativity
difference from things on the left hand side
of the periodic table are very different from
the electronegativity of things on the right
hand side.
So one of the periodic trends that you learned
in g chem that you will absolutely need to
know here is that we increase electronegativity
going to the right 
and we also increase electronegativity going
in this direction.
And I'll just abbreviate that EN for increasing
electronegativity going in that direction.
So that leaves fluoride with the largest electronegativity.
Basically, in general, if the electronegativity
difference between the two atoms is greater
than about 1.8 -- so let's write that over
here -- in general, oh it's right here, electronegativity
difference greater than about 1.8, the electron
will be transferred completely and the bond
will be an ionic bond.
There are some exceptions to this rule but
that's a good general rule here.
So that's ionic bonding, let's talk about
covalent bonding.
Atoms attain a filled shell by sharing electrons.
Most of the compounds that we talk about in
organic chemistry have covalent bonding.
So here's an example here.
We have two hydrogens, and these would be
hydrogens that you just plucked right off
the periodic table, so they both have unpaired
electrons.
So they come together and they both share
electrons.
Each hydrogen is happy because it has two
electrons in its valence shell, that would
be a filled shell for hydrogen, and then this
hydrogen is also happy because it has a filled
shell, it has two valence electrons.
Even though they're sharing them, both hydrogens
have a filled valence shell.
And so we usually write that bond as just
a line.
So here's my equivalence symbol here, we have
hydrogen-hydrogen, so there's a hydrogen-hydrogen
covalent bond.
Each hydrogen shares two electrons 
and so each hydrogen has the helium configuration.
So that would be a covalent bond between like
atoms we can also have a covalent bond between
unlike atoms.
So we could have hydrogen combining for example
with a chlorine, straight off the periodic
table again, and they can each share two electrons,
they're going to combine.
Hydrogen is going to share two electrons with
chlorine, chlorine, because hydrogen donated
an electron now it has eight electrons in
its valence shell, hydrogen now has two.
So both atoms are happy.
So hydrogen has two electrons and chlorine
has eight, so they both have filled shells.
This type of bonding is seen with atoms of
the same or closer electronegativities.
So for hydrogens, two hydrogens combining
have the same electronegativity, hydrogen
and chlorine pretty different electronegativities
but we can have all kinds of covalent bondings.
So let's look at how hydrogen, carbon, nitrogen
and oxygen and the halogens satisfy the octet
rule.
So carbon has four valence electrons, it needs
four more to fill its shell and so it's going
to make four bonds with neutral carbon, it's
going to make four bonds.
Nitrogen has five valence electrons so it
needs three electrons and so it's going to
make three bonds, it's going to make three
bonds and it's going to have a pair of nonbonding
electrons.
I want to emphasize the neutral here because
we also work with a lot of charged compounds
in o chem so this is neutral.
Neutral nitrogen has three bonds and a pair
of nonbonding electrons.
Oxygen has six valence electrons, we need
two to complete the shell, so oxygen is going
to make two bonds, and it's going to have
two pairs of nonbonding electrons.
So this is what we will commonly see with
neutral oxygen.
Hydrogen can only make one bond, it needs
one more electron so to make that one more
electron it's going to make one bond.
Halogens, here's an example, chlorine let's
just do chlorine, let's put some atoms here.
Chlorine has seven electrons, it needs one
electron so it's going to make one bond.
So here's an example here for each of these.
Hydrogen, oxygen with six, nitrogen with five,
and carbon with four.
So when you draw it that way you can see,
it's really easy to see that carbon here -- oh
what's going on here?
Oh that, I'm sorry this light needs to go
off, did not notice that.
Let's do [inaudible], is that better?
Ok, sorry about that you'll have to let me
know when that happens, I didn't notice.
Alright, so when we draw it like that you
can see that carbon needs to make four bonds
to make a full valence shell.
Nitrogen has to make three so these unpaired
three electrons are what it's going to use
to make new covalent bonds.
Oxygen needs two, hydrogen one and chlorine
one.
An easy way to remember this is called the
HONC 1, 2, 3, 4 rule: hydrogen and halogens
make one bond, oxygen makes two bonds, nitrogen
makes three bonds and carbon makes four bonds,
and again this is for neutral compounds, I
want to strongly emphasize that, for neutral
atoms.
Questions so far about covalent bonding?
Alright, so that's neutral compounds, what
happens when we have charged atoms?
That's for neutral atoms, what about charged
atoms?
Charged atoms, or charged molecules, the number
of bonds change.
So carbon normally has four valence electrons
when we take it right off the periodic table.
When carbon is positively charged it only
has three and so in order to, when carbon
is cationic it's going to make three bonds.
So notice we have no octet on carbon and we
call this a carbocation.
So no octet on carbon, carbocation, and so
when we see carbon that's positively charged
it's going to have three bonds and an incomplete
octet always.
So the number of bonds to atom when ionic
so carbon with a negative charge is going
to have three bonds and a pair of nonbonding
electrons.
So we call this a carbanion.
So we have carbocations and carbanions.
Nitrogen has number of bonds to neutral atom
three, when it's cationic it's going to have
four bonds, when it's anionic it's going to
have two bonds and two pairs of nonbonding
electrons.
For oxygen number of bonds to neutral atom
we have two, when it's cationic we have three
and a pair of nonbonding electrons, so this
would be similar to hydronium ion which you're
already familiar with, when it's anionic it
has one bond and three pairs of nonbonding
electrons.
Alright, so a couple of different approaches
to understanding why this is: approach number
one you can just memorize this chart, that's
certainly one way to do it; another way to
do it is if you don't like memorizing, I don't
like memorizing, is how can I understand what's
going on here, how can I understand that carbocation
has three bonds, a carbanion has three bonds
and a pair of nonbonding electrons.
So let's look at that.
So we have neutral, let's look at carbon,
we have neutral carbon has four electrons
in the valence shell.
So if we have a carbon with a positive charge
that means we have one of those electrons
is gone now carbon only has three electrons
in its valence shell.
So this is what if we drew electron dot structure
here, this is what a carbocation would look
like.
And so you can see that that carbon can only
really do three shared bonds because it has
three unpaired electrons here.
So it can only bond to three atoms.
Alright, carbon with a negative charge, on
the other hand, neutral carbon has four electrons
carbon with a negative charge is going to
have five electrons.
Alright, so that's negatively charged carbon
and you can see its got three spots to share
an electron pair.
So it's going to make three bonds.
So this bonds to three atoms.
So bonding like in nitrogen, bonding like
in neutral nitrogen.
So let's put this back up for a second here
and take a look at that.
See nitrogen here, we're going to overlap
a little bit hopefully you can see both of
those, neutral nitrogen looks like this.
So this is isoelectronic with neutral nitrogen,
it's got three spots for making bonds just
like nitrogen.
The only difference here is that this is negatively
charged and certainly it's not nitrogen it's
carbon.
Oxygen, let's do O+ here, oxygen normally
has six electrons in its valence shell but
we have one less because it's positively charged
so we have five electrons.
So oxygen with a positive charge is going
to look, it's going to have the same number
of valence electrons as carbon with a negative
charge and nitrogen with a neutral charge.
You can see, when you do that, that we're
going to make three bonds also.
Bonds to three atoms like C- and nitrogen.
Let's do one more, you can do this also for
nitrogen, let's do oxygen with a negative
charge.
So normally oxygen straight off the periodic
table has six, it has a negative charge so
we add an electron, so it's going to have
seven electrons.
So it's going to be like a halogen, isn't
it, it's going to bond like a halogen it's
going to make one bond.
So this one bonds to one atom.
So it's kind of a big deal.
One of the things that's very different about
O chem is that we are very concerned about
where charges are in the atom.
So in G chem you would draw a molecule and
if the molecule had an overall charge you
would put the brackets around it and put the
charge outside the brackets.
We don't do it that way we put the charges
on the atoms.
So if we have a molecule that has a positive
and a negative charge in the molecule, and
so no net charge overall, we put the charges
on the atoms.
So we don't ever put the charge outside of
brackets.
So really big deal about what the charge is
on an atom based on how bonds and how many
nonbonding electrons, that's a big deal here.
However you want to do this, you just want
to memorize the charges, you want to memorize
them here but I can tell you that the points
that you miss with not doing charges correctly
they'll add up a lot, and I still see students
in 51C getting charges wrong.
So if you want good insurance for yourself
you're going to learn the charges right away
and that will really help you so it's very
second nature to you all quarter.
So what I invite you to do is when you get
your first midterm you start keeping a tally
of how many points you've missed because you
didn't have the right charges.
So it does add up, it's a big deal.
So this is really important and completely
different than G chem in this respect.
Alright, exceptions to the octet rule, third
row and higher elements these have d orbitals
available for bonding so they sometimes exceed
an octet, not all the time, in their valence
shell.
We don't deal with too many compounds like
this but the two things that we do use: sulfuric
acid and phosphoric acid.
So we have sulfuric acid here, we have phosphoric
acid here.
Not even so much phosphoric acid, but we do
work with potassium-containing compounds and
so because these are third row or higher elements
for our central atom here they can exceed
the octet very easily.
So oxygen here has an octet, all the oxygens
have octets.
Oxygens, anything in the second row cannot
exceed an octet but sulfur can.
This oxygen here has an octet, has to third
row.
Sulfur on the other hand has how many electrons?
2, 4, 6, 8, 10, 12 electrons.
And that's OK because it's a third row.
And we'll see coming up why that's OK to have
that for sulfur when we talk about orbitals,
we'll see that coming up.
Phosphorus has, remember two for every bond,
2, 4, 6, 8, 10.
10 electrons.
Again OK because it's a third row element.
Alright so that's one exception to the octet
rule.
Yes, question?
"Why is it that on the phosphoric acid on
of them is OH and the other [inaudible]?"
That's the structure of phosphoric acid, there's
several resonance structures, we're going
to get to resonance structures coming up but
that's generally the way we draw that, OK?
We've got three hydrogens to account for so
three of those oxygens are bonded to hydrogen
and so if we had a double bond here, that
would make positively charged oxygen and that
could be a resonance structure but it's not
the best resonance structure for phosphoric
acid.
Any other questions?
Yes?
[Inaudible]
Oh, right here is this what you're talking
about?
[Inaudible]
OK, yeah you know what, you're right.
We're going to talk about that coming up,
we do little short hand things in O chem and
once you've done it it's hard to not do it
and so that's what I just did, thank you.
So you're not even going to bat an eye when
you see that later on because we're going
to talk about short hand ways of drawing.
Thanks for clearing that up.
Alright, so exception number one: molecules,
atoms that are third row and higher elements.
Molecules with open shells sometimes there
are not enough electrons to provide an octet,
so boron is a famous example which you probably
saw in G chem a number of times.
This can make three shared bonds.
So this is just like something we just looked
at, right?
What did we just look at that looked like
that?
Carbocation, right?
So yeah this is like a carbocation.
The difference is the carbocation has a positive
charge but it's the same situation.
So this can make three shared bonds.
So one of the famous compounds that we have
that boron really likes to combine with is
fluorine so we have, and I'm just going to
draw these in a different color so you can
see what's going on easily here it's going
to combine with three fluorines.
When it does that you can see though that
boron only has six electrons in its valence
shell.
So we call this an open shell compound.
We're going to see carbon as an open shell
compound.
Right now it looks really strange but we use
carbon, carbocations are very very common
intermediates in organic chemistry so that's
a common example in organic chemistry that
you'll see.
Open shell compounds can accept a pair of
electrons to complete an octet so one of the
things you might remember from, and we're
just going to draw this with bonds now, one
of the things you might remember from G chem
is compounds with open shells like to accept
a pair of electrons.
So this is a common reaction for boron trifluoride.
And I'm going to use an arrow here, this is
going to be our first arrow, we're going to
see more arrows in chapter 2, but what that's
telling me is that I"m just taking those two
electrons from nitrogen and donating them
to boron.
Anybody remember what this time of reaction
is called from G chem?
Lewis acid-base reaction, if you don't know
that, not a big deal we're going to talk about
that in Chapter 2, but that's a Lewis acid-base
reaction.
So once that happens boron has a negative
charge and nitrogen has a positive charge.
Now since overall that's a neutral in G chem
you wouldn't have put those charges on those
atoms but we do in O chem, we want to know
exactly where all the charges are on each
atom.
So new covalent bond here 
and now boron has an octet.
So we're going to see the same type of activity
with carbocations, carbocations open shell
just like boron so only six electrons in the
valence shell just like this and what we're
going to see with carbocations is carbocations
also accept a new pair of electrons from another
molecule.
Questions so far, anybody?Yes?
"Are you putting the positives and minuses
because of the polar [inaudible]?"
This is a formal charge on boron and a formal
charge on nitrogen, and we're going to talk
about that coming up, how you calculate that
how you figure that out, OK?
Alright, representation of structure.
There's a Lewis structure which is what I've
been drawing so far.
Lewis structures give connectivity and the
location of all bonding and nonbonding electrons,
so when I ask for a Lewis structure on the
exam you need to include all the bonding and
all the nonbonding electrons and any formal
charges on any atoms.
Again we're going to talk about how you calculate
that coming up.
They don't show actual 3D arrangement of atoms
in space, so that's the nice thing about Lewis
structures we don't have to worry about 3-dimensional
orientation in space.
So we know for a fact that in boron trifluoride
we don't have 90 degree bond angles, we just
don't, nor do we have 90 degree bond angles
in ammonia, but if it's a Lewis structure
that doesn't matter we can draw it anyway
we want, we can draw any bond angles we want
because Lewis structure doesn't show orientation
in three-dimensional space.
There are ways to represent that we'll talk
about coming up.
To draw Lewis structures use the traditional
method, see page 13 3rd edition or page 14
4th edition, follow the HONC 1, 2, 3, 4 rule
for neutral compounds to maximize the number
of bonds without exceeding the octet rule
and assign formal charges to any charged atoms.
So I know you drew a lot of Lewis structures
in G chem it's actually a little bit easier
here because in G chem you pretty much had
the whole periodic table you had to think
about we just use a very small bit of the
periodic table so it makes it a little bit
easier for us to draw structures.
What's different is that in G chem you always
had a central atom, we don't always have a
central atom here, so you kind of have to
look at each atom as you go.
So I'm going to work some examples so you
can see how to do this.
So anyway that works for you is what you're
going to use, I don't make you do this a certain
way.
So you can use the traditional method, you
can do the HONC 1, 2, 3, 4 rule or a combination,
so whatever works for you.
And what you're going to find is that after
a while you're going to know how to draw things
without going through the whole rigamarole
to figure out how many electrons you have
just by using the HONC rule and just getting
used to certain structural features in molecules
you'll be able to do it without going through
a lot of work here.
So for example draw the Lewis structure for
CH3Cl there's two ways to do this.
Find the total number of valence electrons
of all atoms and arrange pairs of electrons
around each atom to give each an octet, if
any atom does not have an octet use a lone
pair to form a double bond so this is kind
of the G chem way that you used to do this.
So let's do that.
Carbon has four electrons, three hydrogens
so three times one electron per hydrogen,
chlorine has seven electrons.
So we add those all up, 14 electrons.
So we have carbon, we actually do have a central
atom in this particular example we're just
not always going to have that.
We've got three bonds to hydrogen, that would
be 2, 4, 6, two for every bond, bond to chlorine
that would be eight, and we certainly are
going to have some lone pairs on chlorine,
2, 4, 6, 8, 10, 12, 14.
OK, so everything adds up.
We've got our electrons, nothing is -- everything
has an octet, hydrogen doesn't have more than
one bond, we're good.
Definitely I want to emphasize, don't forget
to draw all nonbonding electrons on Lewis
structures.
Don't forget to draw all nonbonding electrons
on Lewis structures.
We're going to learn some shorthand ways of
drawing things and when we do the shorthand
ways we're not required to put lone pairs,
but on Lewis Structures we are.
So that's point number one, point number two:
don't use lines for lone pairs.
Alright so the way that you would draw this
in G chem and we don't want to do it this
way.
Yes it's faster but it's a problem here and
it's a problem in O chem because, so that
would be lines for lone pairs, so if you haven't
done that before this is going to be easy
for you -- and I'm off the screen here -- if
you haven't done that before this is going
to be really easy for you.
But you see the problem here when we start
dealing with things like chlorine so sometimes
students will leave off a lone pair, and if
they leave off a lone pair then is that a
chlorine, is that a carbon with a pair of
nonbonding electrons?
The other thing is when we're doing charges
we have, say we have atoms with a negative
charge so that a nonbonding electron or is
that a charge?
And the other thing is we're going to be doing
things with arrows when we're doing mechanisms
coming up and you might have an arrow going
to this molecule and that pair of nonbonding
electrons might look like part of the arrow
or not, it just gets too confusing so we don't,
I don't want you to use lines for lone pairs
in this class, so that's maybe a hard habit
for you to break.
Alright, here's a shortcut, so that's sort
of the G chem way, the shortcut is to use
the HONC 1, 2, 3, 4 rule.
That's that we know that carbon makes four
bonds and hydrogen and chlorine one bond each.
And so carbon's going to make four bonds,
so one bond to hydrogen, another bond to hydrogen,
another bond to hydrogen, another bond to
chlorine and then certainly chlorine needs
its lone pairs to make an octet so that would
be a faster way to do that.
And you really are going to get to a point
where you won't even have to think to draw
methyl chloride, you won't have to count electrons
you won't have to do any of that, that's where
we want to get so that makes it a lot easier
for you.
Let's try ethylene, so CH2CH2 for ethylene.
So we've got two carbons, so two times four
equals eight electrons we have four hydrogens,
four times one that equals four electrons
so we've got 12 electrons all together.
We don't have a central atom here, so we're
going to have the two carbons here.
This is a shorthand notation we're going to
talk about coming up it's a condensed structure.
Each carbon is going to be bonded to two hydrogens.
OK so let's count electrons 2, 4, 6, 8, 10,
and one carbon is getting a lone pair and
one is not.
So if we have an unpaired electron we're going
to try to make a bond to get rid of that unpaired
electron because carbon doesn't have an octet,
this carbon does and so we're just kind of
move that over like this.
That gives us this structure and as you can
see carbon is making four bonds, each carbon
makes four bonds, each hydrogen makes one
bond and so that's what ethylene looks like.
Now ethylene does not have right angles in
its structure so I know for a fact that we
don't have right angles here but that's OK
it's a Lewis structure we can draw things
anyway we want in Lewis structures, don't
have to draw correct angles.
Alright, so this can also be done with the
HONC rule we know that carbon has four bonds,
each carbon has four bonds, each hydrogen
has one, so in order for this carbon to have
four bonds it has to have a double bond between
the two carbons, so that's another way to
look at that.
Alright, now we're going to get a little,
we're getting increasingly difficulty here
now we have something that has a charge and
so we're going to draw a Lewis structure for
this compound that has a positive charge and
I would recommend if you're drawing a structure
that has an overall charge that you don't
use the HONC rule that you actually count
electrons because especially at the beginning
it's going to be difficult for you to draw
and to know that you have the right number
of electrons.
Alright, so we have two carbons, and I'll
just do it right underneath, so that's two
times four electrons, we have five hydrogens,
five times one, we have one oxygen, so that's
six electrons and don't forget to account
for the charge, if we have a positive charge
that means we have one less electron, right?
So one positive charge, minus one electron
and if you add that all up it's going to be
18 electrons.
Alright, when we're drawing, I want to make
a strong point here, when we're drawing Lewis
structures we are only doing even electron
pairs, we're not doing any unpaired electrons
with these, so make sure when you add up your
electrons that this number is an even number,
if it's not you've done something wrong, you've
forgotten to account for charge or counted
something wrong, so 18 electrons.
If we forgot that charge this would be 19
and we would know something is wrong, OK?
Right, so let's draw this here.
Carbon here, this carbon is bonded to three
hydrogens because of the way it's drawn in
the condensed structure and I'm going to show
you that coming up.
This carbon here, the second carbon is bonded
to one hydrogen and then we have oxygen which
is also bonded to a hydrogen, so there's sort
of the skeleton now let's count up electrons
because we have some extra electrons we have
to place here.
2, 4, 6, 8, 10, 12, 14, 16, 18.
Now I put those on oxygen because I just naturally
put those on oxygen because it's more electronegative,
oxygen is more electronegative than carbon
so we have to make decisions on where to distribute
electrons because we don't really have a central
atom, which atom is central here?
So I put them on oxygen because oxygen is
more electronegative so when you were doing
this in G chem you'd distribute the electrons
to the outer atoms, right, first before you
did the central.
And I've put them on oxygen because I know
that oxygen is more electronegative.
Alright, could've drawn it that way but notice
carbon only has three bonds here, so we actually
could have drawn, you've might have drawn
it a different way here.
And this is another way to draw it.
This carbon with four, with three hydrogens
and we have this carbon bonded to one hydrogen.
We're concerned that this carbon only has
three bonds so let's make that carbon with
two bonds.
So let's count electrons 2, 4, 6, 8 ,10, 12,
14, 16, 18.
So that would be another way to draw that
and that looks better because all atoms have
octets in that molecule.
But there is more than one way to draw this,
and so I want to make some points about this
example here before you turn the page let's
do -- I'm going to put this up here but you
don't have to turn the page yet.
A formal charge must be assigned to any atoms
that have a formal charge, do we have atoms
with a formal charge in this first way of
drawing that?
Carbon, right?
Carbon with three bonds has a positive charge
and that's if you've memorized that chart,
but we know carbon has four bonds and so if
carbon doesn't have four bonds it has a charge
so we know that.
Oxygen likes to make two bonds by the HONC
rule so it doesn't have two bonds it has three
that means it's going to have a charge, and
so if you've memorized that that chart than
you would know that this has a positive charge.
Alright, so formal charge must be assigned
to any atoms that have a formal charge so
these structures would not be complete without
the charge.
Formal charge indicates deviations from neutrality
and it's a way to locate the charge on an
ion.
So formal charge, if you calculate it, is
number of valence electrons minus half the
number of bonding electrons minus the nonbonding
electrons, that's the formula for calculating
it.
Do you remember that from G chem, did you
have something like that in G chem?
Ok, so that's the same formula.
And you also know from G chem that the sum
of all the formal charges in the molecule
must equal the overall charge.
So let's calculate that.
Now I'm going to save you some time here with
the HONC rule.
We're going to go back and calculate that
positive charge, so that's in case we don't
have that charge memorized we would need to
calculate this charge.
So I'm going to save you some time here, do
I have to calculate a formal charge of any
of these hydrogens?
No because they're bonded to one atom, so
I will see some people on the exam calculate
formal charge for every single atom.
If it has the right number of bonds and it
has the right number of nonbonding electrons
you don't need to calculate it, it won't have
a charge.
Does this carbon have a formal charge?
No it's got four bonds, this oxygen here does
it have a formal charge?
No it's got two bonds, so we're just left
with this carbon here.
So let's calculate that.
Alright carbon has four valence electrons
and we subtract half of the number of nonbonding
electrons, carbon has no nonbonding electrons,
so that's minus -- half the number of bonding
electrons, so half of six, it has zero nonbonding
electrons and if we calculate that it's plus
one.
Now we leave the one out so one is understood,
so don't write plus one just write plus, if
it's two then we would write it but one is
understood.
And then let's go ahead and do oxygen here.
Oxygen has six valence electrons, it has three
bonds so it's going to do half of six and
it's got two nonbonding electrons, minus two
and that's equal to plus one.
So each of these molecules here, each of these
ways of drawing this both have a charge and
that's good because the overall charge of
this molecule is positive and both of these
have a positive charge so that means we've
done really well.
Alright, I'm going to show you another way
to do this without using the formula.
Yeah, question?
"Would it be incorrect to do it, to calculate
[inaudible]"
That's what I'm going to show you right now.
So I think he reads minds a little bit here
in the front here.
So here's another way to do it, because you
know what, when you're nervous taking the
exam it's easy to make a math error, just
doing this calculation here.
Alright, so carbon normally has straight off
the periodic table has four valence electrons,
it owns for the purposes of formal charge
it owns one electron from each of these bonds.
So it owns 1, 2, 3, it's supposed to own four
so that means it's down one electron, it has
a positive charge.
Is that what you were talking about?
Same thing, oxygen six valence electrons it
owns all of its nonbonding and one from each
of these bonds 1, 2, 3, 4, 5.
It's only got five, that means it has a positive
charge, so that's another way.
So I've given you three ways to do this, memorize
the chart, use the formula or just look at
how many electrons that atom actually owns
and when we look at how many electrons that
atom owns that's exactly what we're doing
right here, OK?
Alright so that's point number one, the above
molecule can be represented by more than one
Lewis structure, we've got two here that are
possible.
They are resonance structures for the same
molecule so these two are resonance structures.
Alright, we're ready to turn the page now.
Alright, I just redrew this to save us a little
time.
If we wanted to do arrows we can show how
to go from this resonance structure to this
one or form this one to this one.
To go from this resonance structure to this
one we're going to use an arrow that looks
like this it's going come from the nonbonding
electron pair and we're going to go right
here.
We're going to move that pair of nonbonding
electrons we're going to make a new bond between
carbon and oxygen.
So there's the new bond, so notice what we've
just done here we've moved this pair of electrons,
we're making a new bond here so now over here
we have a new bond, new covalent bond, and
we have we only have one pair of nonbonding
electrons now.
So that's -- and so we can go back and forth,
if we wanted to go the other way, Sapling
has you go the other way, they have you go
each ways, so let's go the other way too.
To go from here to here we would just reverse
that arrow we're going to come from this bond
and we're going to go here over to oxygen.
That would go back and forth from each of
these resonance structures.
Alright so what we've done is we've done something
called a curvy arrow and that's something
new that you didn't use in G chem, curvy arrows
like right here and that shows movement of
a pair of electrons.
A fish-hook arrow shows movement of a single
electron, so this looks exactly like a fish-hook.
So this shows movement of a pair, and this
shows movement of a single.
We have a reaction arrow, we have all different
arrows here.
We have reaction arrow, there's also equilibrium
arrow, which you should remember from G chem,
equilibrium arrows going back and forth, and
then we have resonance arrow.
They all mean something different, we want
to make sure that we use the right one.
So this one right here is the resonance arrow
and the other two here are curvy arrows.
And so the brackets indicate resonance structures
here.
So again rather than having the positive charge
on the outside of the brackets we actually
have it on the various atoms that it belongs
to.
So these two structures here, I want to make
a point about this, we have two resonance
structures, these two resonance structures
are not going back and forth from each other,
that would be if we were going from this resonance
structure to this one and back to this one
and back to this one, we would use equilibrium
arrows, we're not doing that we're using resonance
arrows.
So the two structure of this molecule is a
hybrid of those two resonance structures.
It's not 50 percent this and 50 percent this,
it's sort of a weighted average which ever
one is better which we're going to talk about
coming up and it's sort of a hybrid of the
two, so we only have one true structure.
We can draw that structure, I want to emphasize
first that they are not in equilibrium, and
let's try to draw the hybrid really quick
-- I'm not sure which clock is right I've
got a clock up here and I've got this clock
who knows the real time?
It's 9 what?
It's 9:51 we'll draw the hybrid next time.
