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CATHERINE DRENNAN: So
that is electron affinity.
But honestly, chemists
don't really talk so much
about electron affinity.
They prefer to talk
about electronegativity.
And these are highly
related terms.
So this was also re-copied,
although completely identical,
I think, between the handouts.
I just thought it was weird
to have re-copied this and not
this.
So your handout for today
is perfect on this point.
So electron negativity,
the net ability
of an atom to attract an
electron from another atom.
So you can see that electron
affinity and electronegativity
are very similar.
In fact, all of these terms are
highly related to each other.
And this idea of
electronegativity,
of this as a term for a
way of thinking about atoms
initiated with Linus Pauling.
But here I have up
a different picture.
I have Robert Millikan.
And the reason why I
picked this picture
is because he helped, a few
years after Linus Pauling came
up with this idea, coming
up with an equation
that help people think better
about what electronegativity
is.
And another reason why
I picked his picture
to put up here instead
of Linus Pauling
is that he was an
MIT undergraduate,
and he was a chemistry major.
I'm not sure in
1917 or whenever--
that was the yearbook picture--
when he took-- I don't think
it was 5.111 at that point.
But at some point, he was
here studying chemistry just
like you.
And then when he got a
faculty position later on,
he did some beautiful work that
had to do with bonding that he
got the Nobel Prize.
So some of you may have a Nobel
Prize in chemistry one day.
And so you want to make
sure your yearbook picture
is at least as good as this
one for other generations
of professors to
show your picture
and describe the work
that you did to contribute
to the field of chemistry.
Oh, and he was born
in Massachusetts too,
so he is a native to this
area in more than one way.
All right.
So the way that he-- and
this is a little bit more
of a squishy definition.
So electron negativity is
proportional to a 1/2--
and IE stands for what?
AUDIENCE: [INAUDIBLE].
CATHERINE DRENNAN:
Ionization energy.
And EA-- our electron affinity
that we just discussed.
All right.
So it's related to
these other terms
that we have already
talked about.
So let's think about
then what this means.
So we can consider an atom
with high electronegativity
and an atom with low
electronegativity.
And we want to think
about whether an atom
with high electronegativity
is going to be an electron
acceptor or an electron donor.
And that you can
tell me, and that is
going to be a clicker question.
So you can try to grab
your handout while clicking
at the same time.
All right.
So we'll take 10 more seconds.
Yes, 88%.
That's great.
Of course, if you looked
and if you didn't believe
it could be a donor, then that
ruled out three of the four,
but that's OK.
Those are good things.
So yes, if it has high
electronegativity,
it's going to be an
electron acceptor.
And part of the
reason for that is
that it has a high
affinity for electrons.
And another part of
the reason for that
is that if you look
at the equation, when
you have a high
ionization energy,
something that has a
high ionization energy
is not going to be a good donor.
So that wouldn't make sense.
So both of those terms having
high in both categories
is consistent then
with this trend.
So let's take a little
bit more of a look at that
and why this is true.
So high
electronegativity, an atom
with high electronegativity
is an electron acceptor,
and then low would be a donor.
And so if we think about this--
and this is our periodic table.
And again, it's not
going to be including
our noble gases, which really
don't want to be accepting
or donating anything.
So in this corner then we had
our high ionization energy,
and we had a high
electron affinity.
And we saw last class we
had high ionization energy.
So it doesn't want to
give up an electron,
but it does want to accept one.
So we have things that are
going to be good acceptors.
And down here, we have
low ionization energy,
so it's easy to
donate an electron.
Oh, let me just put
these up, sorry.
So we have then if you're
high and high up here,
you have something
that's a good acceptor,
and it's going to have a
high electronegativity.
So high high means
high over there.
And then down here,
we have low low,
which means we have
low electronegativity.
Low ionization energy-- it's
easy to give something up.
Low electron affinity--
it doesn't want electrons.
It's happy to give up electrons.
And if it gives up
electrons, then you
can get a complete octet.
It can have a noble
gas configuration.
So on this side,
you need electrons.
This side, it's happy
to give them up.
So if we look then just
at a periodic table again,
this makes sense.
We gain an electron over here.
We get our happy noble
gas configuration.
We lose electrons over here.
We do the same thing.
So that's a way to think
about electronegativity.
All right.
So why should we care
about electronegativity?
And that's because a lot of
atoms that are electronegative
are used in
pharmaceutical molecules,
and that this gives
them special properties.
So to hear in their
own words, we're
going to hear from a former
uropper, Kateryna, talking
about why you should care
about electronegativity.
[VIDEO PLAYBACK]
- My name is
Kateryna Kozyrytska.
I come from Ukraine.
And I'm interested
in how microorganisms
fight each other.
Humans are very smart.
They have found chemical
ways to make new drugs.
And so we spend a
lot of time and money
on making a new antibiotic,
and then we put it into people
and we hope for the best.
Bacteria are also
very, very smart.
And they somehow learn to
resist this new antibiotic
that we just made.
Bugs, on the other hand,
have been fighting each other
with the same
molecules for thousands
and thousands of
years, and we see
no resistance developing there.
So we want to learn what
it is about the antibiotics
that bugs make to
fight each other that
makes them so difficult
to develop resistance to.
Normally, living organisms
use 20 amino acids.
But these bugs get
very tricky in building
their anti-other bugs molecules.
So to add the
functionality, bugs
can do chemistry on
the building blocks,
the amino acids themselves.
And so one of the
things that they can do
is chlorinate carbons, which
activates these carbons
for future chemistry.
The protein I study,
the halogenase SyrB2,
takes chloride ion from the
environment and a molecule
of threonine, the amino
acid, and puts those together
forming a chlorine-carbon bond.
Since chlorine is so
electronegative, when bonded
to carbon it pulls electrons
away from the carbon atom.
And so it makes
the carbon to which
it is attached
much more reactive
toward other molecules.
And this increased reactivity,
at least partially,
accounts for the antibiotic
effect of the molecule.
I am hoping to figure out how
SyrB2 positions all the atoms
in such a way that they react
in this very controlled, very
appropriate manner, so
that later we could maybe
re-engineer this or
other proteins to make
them do chemistry that
we want them to do.
My hope is that we
could understand SyrB2
well enough that we'll
be able to remake it
into a protein
that will actively
participate in synthesis
of new antibiotics.
[END PLAYBACK]
CATHERINE DRENNAN: Yeah.
So that is Kateryna.
And so how adding
halogens, because they
are so electronegative, is
actually a very important area,
and I'll give you a
couple more examples.
I also haven't watched
that video in a little bit,
and she said she was
from the Ukraine.
She's technically
now part of Russia,
I think, from part
of Russia, so.
[LAUGHTER]
Anyway, that's a topic
for a different day.
So this is big business,
actually, putting halogens
on things.
And if you become interested--
if you're taking an antibiotic
or something, start looking
at what the molecule is,
start counting how many
halogens are on that molecule,
you will find a lot.
So a lot of antibiotics have
halogens, either chlorides
or fluorine, as shown here.
Also a very common
antidepressant
has it, another
example of something
that is an anti-diabetic.
Huge numbers of molecules
have halogens added to them.
Some of these are derived
from natural products.
So nature was making these
molecules to kill other bugs,
as you heard about in the video.
Other times, they came
up with this molecule
and they said, well, we need to
make it a little bit different.
It's being consumed too fast.
It's broken down too
fast in the body.
Let's add some halogens.
So sometimes it's sort
of a man-made tailoring,
but often it's a tailoring
that nature came up with.
So why all these halogens?
What's the benefit of having
a carbon-fluorine bond instead
of carbon-hydrogen?
And one reason is that having
a fluorine, this really
electronegative atom, on, say,
an aryl ring like this one,
it actually sucks the electrons
out of the ring and makes it
what we call, or organic
chemists like to call,
electron poor.
So it just kind of hauls
those electrons away.
And when you make
something electron poor,
so by replacing
C-H with C-F, that
can make a potential drug
molecule electron poor.
And what this does
often is make it
harder to oxidize the molecule.
So we're going to talk about
oxidation-reduction much later
in the semester, and we'll
come back to this idea.
But this turns out to
be really important,
because the way that the body
metabolizes or breaks down
these molecules is
that it can oxidize it.
And there are a number
of enzymes in your liver
which will oxidize and
break apart these drugs.
And so if you make
it harder to oxidize,
that makes the drug more
stable in your body.
So if you want to
take a drug, you
want it to last for a while.
And especially if it's
something, an antibiotic,
you want it to last till it
kills all of the bacteria,
not just half of them.
And most medicines, you
need them to be around
for them to have their effect.
So you want to tailor
those molecules
so that they don't get broken
down as easily in the body.
And so this is one reason,
and this is big business.
And a lot of the times,
adding those halogens actually
involves pretty toxic chemicals.
So some people, like Kateryna,
are interested in using enzymes
to do it instead.
Some people who are still using
organic synthesis-- example
is Steve Buchwald's laboratory.
In fact, if you go to almost
any top chemistry department,
I think there's
someone who's trying
to find out new methods of
putting halogens on molecules.
It's a very important area
in designing new molecules
and improving them.
OK.
So that is electronegativity.
So one atom added to
a big number of atoms
can change the property
of the molecule
by sucking electrons away.
And now we're going to talk
about atomic and ionic radius
and also isoelectric atoms.
So these trends are pretty good.
We're back to some pretty
good-- fewer glitches.
So what is the atomic radius?
So here we have 2r, 2
times the atomic radius.
So the atomic radius
is defined by the value
of r that has about 90%
of the electron density.
I mean, technically,
an electron could
be infinitely or close
to infinitely far away,
but pretty much most of them are
going to be about within 90%.
We call that the
radius of the atom.
So they're trends.
Trends are pretty good.
So across the periodic
table, what matters
is the Z effective.
See, everything we've
learned about it comes back.
If you didn't
learn it on exam 1,
still learn it because
we're going to use it again.
Across the periodic table,
Z effective matters.
Down the periodic table, n, or
the principle quantum number,
matters.
So how do these matter?
So across the periodic table, Z
effective is going to do what?
Increase or decrease?
AUDIENCE: Increase.
CATHERINE DRENNAN: So
it's going to increase
across the periodic
table, and this
results in a decrease
in the atomic radius.
So again, it increases going
across the periodic table
because you're adding both
electrons and protons.
But the electrons are not
giving you complete shielding,
so you're not canceling
out every proton
with every electron.
So overall, you get an
increase in the Z effective.
And because of that, when
you have this increased Z,
it's kind of pulling
the electrons in.
And then it isn't until you
go down the periodic table
when n increases that you start
to see the radius increase.
So there, the electrons
are getting farther away,
and so you are getting
this bigger thing.
So I like to think about it as
sort of the mom at the park.
There are some moms, their kids
are sort of running everywhere,
but other ones are sort
of hauling their kids in.
They have this
positive force that
seems to keep them all sort
of in the general area.
And so they're
shrinking the size
of their kids' play
area with this force
that they're exerting.
But if the kids
get too far away,
it's like, yeah, they're
not going to hear you call.
They're not going to hear
you jump up and down.
They're not going to see you.
And so they're just
going to be out there,
and the radius of your kids
is going to be farther away.
So we can look at these trends.
If we're over here
in the beginning,
we're starting on this side.
We're going across
the periodic table.
We go down, then we have a
jump up when we increase n.
And then we go down again,
then we have a jump up
when we increase n.
We go down again, except over
here there's a little glitch.
Those d electrons, they're back.
They're going to give
you a couple glitches.
I love d electrons.
We're going to talk about
them more around Thanksgiving.
That's my favorite
part of the course.
Those d electrons are
always causing trouble.
OK.
Then we go up again, and
then we go down, and go up,
and go down.
Those are pretty good.
Those are pretty good trends.
All right.
OK.
So ions.
Ions are different than their
neutral parent, once again.
So we saw this before, that
when you start filling the 3d,
the energy levels change.
So ions can have
different properties
than their neutral parents.
And so if we have
two kinds of ions,
we can have cations, which
are positively charged.
And so a positively charged
ion will have lost an electron,
and so it's going to be
smaller than its parent.
And so we can see here lithium.
And then in the center,
that's lithium plus.
So when you lose the
electron, the radius
actually shrinks quite a bit.
It's like that electron
was just really kind
of causing a bigger radius.
And when it's
finally gone, you're
at a smaller size over there.
Anions-- negatively
charged ions.
So they're gaining an
electron, and their radius
is larger than their parent.
And so you can see over here,
we have oxygen in the center.
Oxygen minus 2 is much larger.
And again, we can see
some of the other trends.
Some of them are the same.
The ionic radius
also will increase
when we're going down a group,
so when n is increasing,
so from lithium to sodium.
We have an increase
from fluorine
to the top of the periodic
table to chlorine.
So we still, as we increase
n, increase in size.
But you have to think about the
ion-- did it lose an electron,
or did it gain an
electron-- to think
about how its size changed
with respect to its parent.
So why does this matter?
There's one example
of-- if you're
interested in
biology or anything
to do with the brain
or neurons, then
you should care
about ion channels.
So there are channels in
membranes that bring ions in,
and this is really important.
So ion channels are in
muscle cells and in neurons.
So if you want to move or think,
something that MIT students
generally like to do
both of those things,
you need ion
channels to do that.
So ion channels should
be important to you.
And you want ion channels.
They regulate the influx
of ions into the cell
and allow for really
rapid responses,
which is also really important.
And amazingly, they're highly
selective for certain ions.
So this is important.
It needs to be tightly regulated
to say how much sodium you
have in there or
how much potassium
that you have coming in.
And if the ion
channel took potassium
when it was supposed to take
sodium, that would not be good.
So these channels are
designed by nature
to be highly
selective, and so they
care to be highly selective.
And you're talking about a
plus 1 maybe or another plus 1,
they have to think
about the radius.
So why don't you tell
me what the differential
is in the radius from smallest
to largest for these three
different ones.
All right, 10 more seconds.
OK, great.
So most people got that right.
Let's just kind of
take a look at that.
So here we want to
think about the neutral,
and then this one has
one less electron.
So that's going to be smaller.
And then when you compare
potassium with sodium,
you have to think about n.
And so this is down
farther, so that's
going to be bigger than sodium.
So here we're thinking
about the difference
in electron configuration, and
here we're thinking about n.
All right.
So amazingly, these channels
have it figured out,
and so look at this.
This is about
significant figures.
Too many people lost points
on significant figures,
I have to say, on the exam,
so make sure you learn them.
But if you had,
say-- you say, oh,
what's the difference between
1 to 3 significant figures?
The difference is sort of the
potassium versus the sodium
radius.
So 1.38 times 10
to the minus 10,
1.02 times 10 to
the minus 10-- those
seem like pretty small numbers.
Does it really matter
if it's 0.2 versus 0.38?
And the answer is,
yes, you'd be dead
if nature could not distinguish
between these significant
figures for you.
So these channels
are designed to be
selective at that kind of atomic
scale, and only let one ion in.
And so Rod MacKinnon,
who's a crystallographer,
won a Nobel Prize for solving
some of the structures
of these ion channels.
And this just shows
a ribbon drawing,
and this shows an all
atom drawing of a channel,
and there's an
ion going through.
That's its hole, and its
radius is perfect for that ion.
And the other one,
even though it's not
that many significant figures
different, doesn't fit.
And that's pretty amazing.
Nature is truly amazing.
That's the hole.
It makes a perfect hole
just for the one kind of ion
that it's supposed to accept.
All right.
So now, there are
one more definition
that we're going to do.
There are things that can
have the same electron
configuration.
Those are called isoelectronic,
and let's think about those.
They don't necessarily
have the same size,
but they have the same
electron configuration.
And I'm just going
to write these out.
So when we think around other
ones near neon, noble gas, that
would have that exact
configuration-- so how
do we get flourine to
have that configuration?
What does it need to do--
gain or lose an electron,
and how many?
What would its state be?
So what do I write-- what's
the thing for flourine that
is going to be the same electron
configuration just in terms
of its charge?
I'd write F what?
AUDIENCE: Minus.
CATHERINE DRENNAN: Minus.
For O, what am I going to write?
For oxygen, what am
I going to write?
AUDIENCE: 2 minus.
CATHERINE DRENNAN: 2 minus.
And, say, nitrogen?
We'll stop there.
What's that?
AUDIENCE: 3 minus.
CATHERINE DRENNAN: 3 minus.
Great.
Let's go on the other side.
What about for sodium?
What does sodium have to do
to have that configuration?
AUDIENCE: Plus.
CATHERINE DRENNAN: Plus.
What about Mg?
AUDIENCE: 2 plus.
CATHERINE DRENNAN: 2 plus.
Aluminum?
AUDIENCE: 3 plus.
CATHERINE DRENNAN: OK.
And silicon-- 4 plus.
So you get the idea.
And now, we can think about
which will have bigger
and which will have
smaller radii as well.
So are these going to
be bigger or smaller?
AUDIENCE: [INAUDIBLE]
CATHERINE DRENNAN: Right.
So they're going to have
larger radii than their parents
because they've all gained.
And then over
here, these will be
smaller since they've
lost electrons
compared to their parent ion.
OK, so let's just do one,
which should be very fast,
clicker question.
It is a clicker
competition after all,
so we've got to get in some
extra clicker questions.
And this should be
very fast, I think.
All right, let's just
do 10 more seconds.
You have a periodic table
up here in case you need it.
Yeah, OK, that's not going to
distinguish the recitations
very much.
Yeah, so you just have
to look at what is nearby
and think about how many
electrons it needs to gain
or lose to have the
same configuration.
All right, bonds--
now, we're up to bonds.
There are three types of
bonds we will discuss today.
We probably won't get
to them all today.
After all, the handout stuff.
But anyway, then we will
discuss in the class over time.
Now, some of you have probably
figured out-- almost everyone
has probably figured out that
one of the things I love to do
is teach chemistry.
I love to teach chemistry.
Some of you have come to my
office hours or pizza forums
or even paid attention
to some of my slides
about office hours may
realize that I love dogs.
So I love teaching
chemistry, and I love dogs.
What is the most amazing
thing that you can think of?
Dogs teaching chemistry.
[VIDEO PLAYBACK]
CATHERINE DRENNAN: So now,
I'm going to let dogs--
- Welcome to "Dogs
Teaching Chemistry!"
CATHERINE DRENNAN:
--tell you about bonding.
- --is chemical bonding.
Chemical bonds are what
holds atoms together.
A chemical bond is an
attraction between atoms
that allows the formation
of a chemical substance.
The electrons that
participate in a chemical bond
are called valence electrons.
These are electrons that are
found in an atom's outermost
shell.
Let's take a look at the
types of chemical bonds
that can be formed
between atoms.
An ionic bond is formed
when one of the atoms
will lose its electron
to the other atom.
This results in a
positively charged ion
called a cation and a negatively
charged ion called an anion.
Positive and negative
attract, and the result
is an ionic bond.
Covalent chemical bonds
involve the sharing
of a pair of valence
electrons by two atoms.
There is also what is
called polar covalent bonds.
These are covalent
bonds in which
the sharing of the
electron pair is unequal.
The result is a bond
where the electron
pair is displaced toward the
more electronegative atom.
Thanks for watching, and
we'll see you guys next time.
[END PLAYBACK]
[APPLAUSE]
CATHERINE DRENNAN: There's,
I think, one other one,
but that is totally
the best one.
And everything they
say is exactly right,
so it's really exciting.
I even love it when
they had the two--
they put two balls in there
that they were sharing.
It's just really very well done.
OK, so now, you've probably
filled in some of your notes
here, but in case you
missed some of them,
I will tell you exactly
what the dogs just told you.
The dogs had it
completely correct.
So ionic bonds is the transfer
of an electron, as you saw,
and then the generation
of a cation and an anion
that are attracted to each
other due to the charge.
So the bonding comes
from that attraction
between the positively charged
and the negatively charged
atom, and an example that you're
probably all familiar with is
table salt, NaCL.
And so you have Na plus
and CL minus that are
attracted to each other
and form these bonds,
which creates table
salt. So let's see
how far we can get in
thinking about, really,
this interaction
between ionic bonds,
and we'll see if we can
get through ionic bonds.
We might have to wait until
covalent bonds until Monday,
but let's see if
we can finish this.
So the formation of NaCL from
neutral Na and neutral CL
will first involve forming
your cations and your anions.
So you have Na going to
Na plus plus an electron,
and so here, you're
talking about a process
where the energy is going to be
equal to the ionization energy
again.
So we're not moving far
away from these terms
because you're talking about
ionizing a neutral atom to Na
plus, and there's
a value for that.
And then we're talking about
neutral CL, neutral chlorine,
going to CL minus, and so
it's gaining an electron.
So here, the process
you're talking about
is the electron affinity,
and so the energy change here
is equal to the negative
electron affinity, which
is minus 349 in this case.
So this is a
favorable process here
to gain this extra electron,
and so overall, the Delta E here
is negative.
So now, if we're going to
talk about this process here,
we have two of
these, so we're going
to go-- we're going to
put these guys together.
So we need both of
those to ionize,
and so we can add up
what energy difference
we should expect to
form Na plus and CL
minus together
from their parents.
And so we have,
now, a plus-- we've
added these two together-- a
plus 145 kilojoules per mol.
So this seems weird.
It's plus, and so
now, we're seeing
that the formation of these
ions from their neutral atoms
has this positive value, which
means it requires energy.
But we think about NaCL as being
this natural table salt thing,
so why is there so much table
salt if this requires energy
to do it?
And the answer is that this
is only part of the process.
You need to form your
cations and anions,
but then you have energy
of them coming together.
So they're attracted
to each other,
and that's a really important
part of forming the bond.
And they're attracted by a
simple coulombic relationship
here.
So the attraction between
the positively charged
and the negatively
charged ion has
an energy of minus 589
kilojoules per mol,
so overall then, if you consider
both forming Na plus and CL
minus and the
attraction between them,
we have a negative Delta E 444.
So the net energy here is
in favor of forming NaCL.
We have a decrease in energy.
This is a stable compound,
so let's look at where
this number comes from.
So we just put this out.
That's the coulomb
thing, but let's
actually calculate this and see
where that number comes from.
So we're back to
coulombic equations again.
We never get very far away.
They turn out to be very
important in chemistry.
So we have the
coulombic potential.
We have z's, our
charge on our ions.
We have the absolute
value of the charge
of an electron squared
over 4 pi, our permittivity
constant in r, our distance
between those ions.
And for any CL the bond
length, or the distance
between Na plus and CL
minus, is 2.36 angstroms,
so we can use that and just
plug it into the equation.
And we have plus 1 for the
sodium, minus 1 for chloride,
so overall, this will
be a negative term.
And if you work
out the math, it's
minus 9.774 times 10 to
the minus 19th joules,
and we have three
significant figures.
What's limiting our
significant figures?
AUDIENCE: [INAUDIBLE]
CATHERINE DRENNAN: What--
yeah, the bottom length
is limiting it here.
I'm just going to talk
about significant figures
for a while, until Test 2
when you can demonstrate I
can stop talking about
significant figures.
All right, so then we want to
convert to kilojoules per mol
because that was the
number I gave you,
was in kilojoules per mol, so
we have our conversion factor
between joules and kilojoules.
And I guess I should
mention up here--
coulombs cancel, and
our meters cancel,
so we're left in
joules up there.
Units are also important.
Then we can use
Avogadro's number,
because we're given
kilojoules per mol,
and we can get out the
number I gave you before.
So this number really
just comes right out
of this equation minus
589 kilojoules per mol.
So we have our ionization energy
to tell us about forming ions.
We have our electron
affinity, and now, we
have a coulombic relationship.
So I said before-- this
is what I showed you
before-- that we have this
attraction that's favorable.
We have forming
the ions, which had
a positive energy
associated with it,
but overall, this process
has a lower energy.
It forms a bond.
But this is just based
on this calculation.
So we can ask what is the
experimental measurement
for this interaction,
and we note
that it's somewhat different.
So we have, instead
of minus 444,
we have minus 411
kilojoules per mol.
So why the difference?
So again, our ionic model,
which just considers
ionization energy,
electron affinity,
and that positive
coulimbic interaction
in our experimental result--
so problems with this model
that we did up here include
that we only thought
about favorable interactions.
There are also going to be
some that are not favorable;
protons against other protons,
repulsion, electron electron
repulsion.
So there are some negatives.
It's never all positive
in any relationship,
whether it's sodium
chloride or anything else.
Always some negatives,
and we ignore those.
And the result of
this is that you're
going to get a larger
Delta E predicted
than the experimental value.
So it seems like this
is more favorable,
like that's a stronger bond,
a stronger interaction;
but really, if there's
some repulsion,
that's overestimated.
It's really going to be
a lower value, which is
what you see in the experiment.
Also, we just said that sodium
plus was one point charge
and CL minus was another.
It's more complicated.
Their interactions are
more complicated than that.
And we ignored quantum
mechanics, but in doing that,
we did pretty well.
If we had one
significant figure,
we would have been perfect,
so to one significant figure,
these approximations
work really well.
OK, that's it for today,
and I will see you Monday.
Have a great weekend.
All right, let's just
take 10 more seconds
on the clicker question.
Great, so people are
getting the hang of this.
If you hadn't yet,
there's still time.
So as you're going across
the periodic table,
you are increasing z and
increasing the z effective
as well, because you don't
have total shielding,
so aluminum is the
correct answer.
It has a lower z effective,
and so therefore,
a smaller ionization energy.
The electrons aren't
held as tightly.
All right, so we're talking
about chemical bonds,
and it seems like
an appropriate topic
to talk about when we're
also talking about bonding
as a community, so that
seems like a good thing.
So a chemical bond is
an arrangement of atoms
so that they come
together in such a way
that they're lower
in energy than they
were when they were apart.
So they're more stable
together than they were apart,
and that's a chemical bond.
So this is page five of
the handout from last time.
And excitingly, we have
Lecture 10 handouts today,
so there's lots of
things working today.
All right, so this is
lower in energy, i.e.
more negative, when these
atoms come together.
So a chemical bond--
as you saw last time
with those wonderful
dogs sharing a pull toy,
a covalent bond is a bond
where the electrons are shared
between these two
atoms, and each atom
is giving up one bond to share.
So we can think about this
more graphically of what
is happening, and we have this
little plot on your notes,
where you're going to be
filling in a bunch of details.
So we have the internuclear
distance, r, the distance
between the two nuclei.
And we're back to
hydrogen for the moment,
so we're going to talk about
a bond between two H atoms.
And on the axis over
here, we have energy.
So we have energy versus the
distance between these two
hydrogens.
So at 0 energy, we just have
the hydrogens by themselves.
They're not interacting with
each other in any kind of way
that lowers either
one of their energies.
There's no interaction,
no energy change.
They're not interacting.
Down here at this dash
line, we do have a bond,
so we formed H2.
The hydrogen atoms are
interacting with each other,
and this is lower in energy.
So what does this
plot look like then
if you draw energy
versus this distance?
So up here, it's above and
higher energy, above 0--
this is unfavorable-- going
down to this dashed line
and then going back up to 0.
So let's think about
what's happening here,
and there's a bunch
of different kinds
of interactions you can have
between those two hydrogen
atoms.
There are repulsive
interactions,
nuclear nuclear repulsion,
electron electron repulsion;
and there are
positive interactions
like the electron nuclear,
the positive and negative.
So up here, these atoms
are very close together,
and that is-- they're really too
close, and that's unfavorable.
Two objects trying to occupy
the same space at the same time
is unfavorable.
But as you start
separating out these atoms,
then they become
more comfortable,
and you get to a
distance where you
have the sharing of electrons.
They're next to each other.
They're interacting
in a positive way,
and they're sharing.
But then if you bring
them too far apart,
they're no longer
communicating with each other.
We don't even know where one
of them is at this point.
They're just kind of
infinitely far apart.
No interaction.
So we go from too close to
just right to too far away.
So this distance here that
has this minimum energy-- OK,
so we have this
minimum energy here,
and that's the distance at which
you have this really favorable
interaction, this bond length.
That's the bond length between
those two hydrogens right here.
Now, we can think about this
energy difference over here,
and it has a special name.
So this is the dissociation
energy, or Delta E sub d.
Sometimes, it's just called
big capital D in your textbook.
So this is the
energy that's needed
to dissociate those atoms.
So if the atoms come together,
and they're lower in energy,
if you then want to
break them apart,
you need to put energy in so
that they can be broken apart.
And that's called the
dissociation energy.
And if it's a really
big number, that
means it's very hard
to dissociate them,
and if it's a small
number, it's pretty easy.
So this is our
dissociation energy here.
So for hydrogen, this value
is 424 kilojoules per mol.
And if we were putting
it on this axis here,
it would be right
down here, so it's
going to be a negative
value on this axis.
We're below 0, so the negative
of the dissociation energy
is found down here, minus
424 kilojoules per mol.
And if you wanted
to break that bond,
you would need to put in
energy, so dissociation energy
is positive.
It's the energy you need to
put in to break those bonds.
So we can think about this
plot now and consider looking
at the plot and evaluating what
you can and cannot say about
different kinds of compounds,
their bond lengths,
and also their
dissociation energies.
So we good people?
Mostly got this written down?
And today, for some of the
people who wanted to come late,
we can post versions
of this, too.
All right, so now let's
consider this plot
and ask which bond is stronger?
So we have hydrazine,
and we also
have molecular nitrogen, N2.
And in your dashed line,
you have nitrogen, N2,
and in the solid line,
you have hydrazine.
So can you just
look at this plot
and tell me which is stronger?
Is nitrogen or
hydrazine stronger?
And you can just yell
out what you think.
AUDIENCE: [INAUDIBLE]
CATHERINE DRENNAN: Yeah.
So nitrogen is going
to be stronger here,
and people knew that
because it's a deeper well.
So you go way farther down,
there's more stabilization.
It's a lower energy, a lower
negative value of energy,
so that means it's going to have
a greater dissociation energy.
You'd have to put in more energy
to dissociate nitrogen nitrogen
than these two submolecules
here in hydrozine.
All right, so we can
also look at this plot
and ask the question,
which has a shorter bond?
What do you think?
Which bond is shorter?
AUDIENCE: [INAUDIBLE]
CATHERINE DRENNAN:
Nitrogen as well, right.
So nitrogen is also
the shorter bond,
and we know this because
this is increasing distance.
And this is closer to the axis,
so this has a shorter distance.
And later, we're
going to be doing
Lewis structures and
other things of nitrogen
and discover it has a triple
bond, which you may already
know, and we'll talk more about
nitrogen's amazing triple bond
as we go along.
And so that's a very short bond
and a very, very strong bond
as well.
So you should be able
to look at these plots
and evaluate what kind
of dissociation energy
it would have, is it
bigger or smaller,
and also what kind of distance
you expect between them.
And you should be able to
draw these kinds of plots
on the exam, if asked, in
just kind of simple detail.
Nothing too fancy.
All right, so in terms
of bond strengths,
carbon monoxide has one
of the strongest bonds,
so it has a very large
dissociation energy.
And iodide, I2, has
one of the weaker ones.
And later in the
semester, we're going
to be doing a demo that shows
why that's kind of cool,
that weak bond leads to
some cool, cool demos.
OK, so those are covalent bonds.
And, Ashley, could you just
close that door, please?
All right, so let's
finish polar covalent,
and then we'll have
our moment of silence,
and maybe it'll be silent in
the hallway by then as well.
All right, so polar
covalent bonds--
so last time, the dog
showed you that you
can have equal sharing
and unequal sharing.
And of those of
you who've watched
dogs play with pully
toys, most of the time,
the sharing is pretty
unequal, and so whenever
you see that again, you
can think polar covalent
and tell your friends,
and they'll be like,
"I don't know what
happened to you at MIT.
Those are two dogs playing.
What are you talking about?"
So this is unequal sharing of
electrons between two atoms,
and this happens when those
atoms have a very different
electronegativity.
So in general, a bond
between two atoms
is considered a
polar covalent bond
when the difference in electric
negativity between the atoms
is greater than 0.4
and less than 1.7,
and that's Linus Pauling's
scale and works quite well.
So if we look at
this little plot,
we could see that carbon
hydrogen bond only
has a 0.4 difference,
so that would not
be considered a
polar covalent bond.
But nitrogen hydrogen
is a difference
that is greater than 0.4, so
that would be a polar bond.
And so you can use these
values to think about
whether you're going to
have unequal sharing,
and the more
electronegative element
is going to be pulling
on those electrons.
We also can use this to
think about polar molecules,
and this is kind of a little
bit of a flash forward
to Friday's lecture.
We're going to talk
about shapes of molecules
because a polar molecule
has to have polar bonds,
but also has to have those
bonds arranged in such a way
that there's a net difference,
there's a net dipole,
there's a net kind of
pulling of those electrons
in a different way.
So here is carbon
dioxide, a molecule
that causes global warming.
Yes, I said it on the videotape.
Yes, I believe that
human beings are
responsible for some
of the global warming,
and we should do
something about it.
This molecule does
have polar bonds,
so we have carbon in the
middle and oxygen here.
So carbon and oxygen have an
electronegativity difference
greater than 0.4, but it's
not really a polar molecule.
It's a non-polar molecule, and
that's because of its shape.
So shape matters.
So we have pulling
of electrons one way,
but we have equal and opposite
pulling of electrons other way.
This would be really
cool if the dogs
could have done this as well.
So in this case, we have
a non-polar molecule
that has polar bonds.
Now, there's only very few
cases where this would actually
be true, and you really need
to think about the shape.
And so that's why we're
going to talk about shape.
Another molecule that also
has polar bonds is water.
So here, we have polar
bonds between the oxygen
and the hydrogen, so oxygen
and hydrogen, greater than 0.4.
But in this case,
the shape of water
is such that they
don't cancel out,
and you do have a net dipole.
You do have a net
charge on that molecule,
makes it a polar molecule.
So we need to know
about the shape,
and we need to know
about electronegativity.
So in large, organic
molecules, sometimes we
just talk about the
number of polar bonds
and then think
about whether that's
likely to be a polar
molecule or not.
You can't always think about
the shape of something really
complicated and what direction
all of the pulling of electrons
is going, but we can at
least count polar bonds.
So here are two vitamins,
vitamin A and vitamin B-9,
which I think is also B-10
and a number of other B's.
Its name is folic acid.
They kept finding
it again and again,
so there's a whole gap of
b-vitamins where they're like,
oh, B-9, folic acid;
B-10, also folic acid;
B-11, I think also folic acid.
B-12 something
different, though.
So anyway, this is a
very important B vitamin,
and I'm actually going to come
back to this molecule later
in the course.
But we can think about how
many polar bonds it has,
and that's a clicker question.
And there, you have the
molecules up here to look at.
All right, let's just
take 10 more seconds.
It takes a while to
count, but I think
you probably can answer it
without maybe fully counting
all of them.
All right, so over here now,
the answer is the folic acid,
and I'll just highlight.
You might not have found
all of these polar bonds,
but you should have at least
seen that this one really
didn't have many.
Vitamin A only has one.
So we have polar bonds down
here between carbon and oxygen
over here, carbon and oxygen,
nitrogen, hydrogen, carbon
oxygen again, nitrogen with
hydrogen, oxygen with hydrogen,
nitrogen with
hydrogen over here.
So folic acid is quite
polar, and if we're
going to think, now, about
whether it is a water
soluble vitamin or a fat
soluble vitamin-- which
is something that a lot
of times your supplements
will tell you about.
If it's water soluble-- and
we'll talk about this more
later-- like dissolves likes.
So water likes polar
molecules, which
makes folic acid water
soluble and makes vitamin
A fat soluble.
It's not very polar, doesn't
dissolve very well in water.
And this kind of turns
out to be important,
in that if you read your
vitamin supplements,
if you take vitamin supplements,
it will often tell you
interesting things like
how many hundred times
over the daily
recommended allowance
this vitamin tablet is, and
if things are water soluble,
it doesn't matter
so much that you're
taking way more than
your body actually needs.
You just have a
very expensive pee.
But if it's fat
soluble, then it's
going to stay in your body,
and you don't need it,
and it can be a
little bit toxic.
So try to think
about the vitamins.
Not everything-- even though
vitamins are good for you,
they're not good in every kind
of amount that you could take,
so you will now use your
knowledge of polar bonds
to figure out whether you
should be taking certain vitamin
supplements.
