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Chapter 4.6c Polarity

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welcome to the third video for chapter

four section six molecular structure and polarity

polarity

in this video we'll be focusing on

polarity the learning objectives are to

explain the concepts of polar covalent

bonds and molecular polarity

and to assess the polarity of a molecule

based on its bonding and structure

um we've discussed previously when you

can have a polar covalent bond what what

that sort of

is but as a quick refresher essentially

a polar covalent bond happens when two

atoms are bonded

together and they have a different

there's a difference in electronegativity

electronegativity

essentially what happens when there's a

difference in electronegativity is that

one end of the bond there's more

electrons spending time around that end

of the bond because of the electronegativity

electronegativity

and again the electronegativity sort of

tells you how stable the electrons can be

uh when they're spending time around

that atom so for example fluorine is

quite electronegative

and that just means that electrons can

be fairly stable

when they are spending time around the

fluorine as opposed to different molecules

molecules

that the electrons cannot be can't enter

as stable

an energy state when they're around the other

other

molecule or the other atom and what that

means is that

in this bond the electrons will be

spending more time around the fluorine

which is going to give rise to a partial negative

negative

and again we're going to use this greek

delta um the other end of the bond the

atom that's going to have

less electron density is going to have a

partial positive

so the difference between this partial

positive and this partial negative

is going to give rise to what we call a

a bond

dipole moment and that is essentially

just the separation of charge in a bond

uh so we'll notice that boron and

fluorine are quite different in their

electronegativities but even

carbon and hydrogen are slightly

different and they they also there is a

bond dipole in this carbon-hydrogen bond

um it's just very small and we can tell

that it's very small because we have a

relationship that tells us the strength

of the dipole moment

so the strength of the bond dipole is

represented by mu

and it is given by q which is the magnitude

magnitude

of the uh the partial charges

or the the magnitude of the separation

of charge and then

r which is the distance between

the two atoms so in other words the

greater the difference

in uh in in electronegativity and

therefore the greater the difference in

the partial charges the larger the net

separation of charge between these two

uh atoms

the stronger the bond dipole moment will

be and also the longer the bond

the stronger the dipole will be

fluorine and boron have a huge

difference in electronegativity so their

bond dipole moment is represented as a

kind of large arrow

and carbon and hydrogen have a much

smaller difference in electronegativity

so it's represented with a much smaller arrow

arrow

as a refresher the arrow goes from the

side of the bond with the partial positive

positive

to the side with the partial negative

and the way that i remember that is that

the back of the arrow looks like a

little mini plus sign

so it goes over the atom that is

we can represent this bond dipole as a

vector and that is essentially what's

happening with these little arrows

a vector is a quantity with both

magnitude and

direction so in other words it has a size

size

like the length of the arrow and also a

direction which is it's pointing towards

one atom versus the other one and this

is really important for us

um because we need to keep track of both

the size of the separation of charge

and the distribution of the electrons

and how that might

so a molecule with a net separation of charge

charge

uh what what that is to say is that um

there is some sort of

general net uh separation of charge

within that entire molecule as opposed

to just in

bonds so the net separation of charge in

a molecule is going to be called the

which is different than the bond dipole

so the bond dipole is just for a single bond

bond

the dipole moment is for the entire molecule

for diatomic molecules essentially the

dipole moment

is equal uh to the um

to any bond dipole because there's only

so if you have two atoms that have

different electronegativities like for

example hydrogen

and fluorine this molecule is absolutely

going to have

a there's definitely a bond dipole here

which means that the

the molecule will also have a dipole

moment because they are the same thing

and we can draw that um based on

the uh the electronegativities of

hydrogen and fluorine

so we know that hydrogen is less

electronegative than fluorine

so we'll go ahead and give fluorine the

partial negative and we'll give hydrogen

the partial positive

and then we can draw our bond dipole

and this in fact is also the dipole

moment of this entire molecule

if we're looking at a different kind of

diatomic for example

we can think about if this molecule is

going to display

any bond dipole and we should say no

this molecule will not have any bond

dipoles this can be a little bit

confusing because

we know from our periodic trends that

nitrogen is a fairly electronegative

atom and the tendency is to say ah if

there's an electronegative atom it must

be polar

but in fact nitrogen and nitrogen

have the exact same electronegativity

there is no difference in

electronegativity here

so there is no dipole moment there is no

bond dipole and there is no dipole

moment there is no separation of charge

this is a perfectly

nonpolar molecule because there is no

difference in electronegativity between

so this gets a little more complicated

when we're dealing with molecules that have

have

more than one bond and this is where our

vector notation is going to come into play

play

so again a vector is just a quantity

with both a magnitude or a size like the

length of the arrow and a direction

like the way the arrow is pointing

um when you have molecules with more

than one bond you have to pay attention

to the three-dimensional structure

um and and i do mean three-dimensional

um we have to think about

uh these molecules as three-dimensional objects

objects

that may have things happening in the

directions that we can't draw too well

because they are into and out of the paper

paper

here are two molecules this one is co2

and this is water

and i'm going to go ahead and walk you

through how we can determine the overall

dipole moment for these two molecules

um co2 is fairly straightforward

we've got our two bond dipole moments c

and o have different electronegativities

so there will be

um bond dipoles as shown here and

they'll be pointing towards the oxygens

which will have partial negative charges

however when we think about the vectors

we need to add

the vectors together and if you've had

calculus you have probably heard that before

before

if you have not had calculus don't worry

about it i'm going to show you how to do it

essentially what you're going to do is

you're going to take these arrows which

represent our bond dipoles

and you're going to redraw them and

we're going to redraw them in a very

special way that i'm going to call tip

to tail

so what's going to happen is you're

going to take the first one and you're

going to draw it

and you're going to try to draw it in

about the same size

and the same direction that the original

one had

and then you're going to take the second

one and you're going to draw it

tip to tail so wherever the last vector ended

ended

is right where you're going to start

your next vector

and then you're going to draw your next vector

vector

and you're going to see what was the net

change right what was the

net difference and here it turns out

that we started

over here at this end we went over there

with this vector this vector showed us

that the electrons

were spending more time over towards

this end of the molecule and then we had

our next vector and it took us straight back

back

so the net difference between the

original place where we started our tail

and then the final place where we wound

up with the tip of our last vector

is uh the difference between those two

spots is your

net dipole moment for the molecule here

there is no difference

and so the net dipole is zero

okay so next we're going to look at water

water

so water has a different geometry than

co2 right carbon dioxide is linear

and water is bent and this is really

crucial right

there's two identical bond dipole

moments here and there's two identical

bond dipole moments here

and we're going to see that this does

not yield

the same result here so we're going to

do the exact same process we're going to

take our

first arrow and we're going to try to

redraw it in

the same size and orientation

i'm going to try real well i already

i'm going to try to draw in the same

size and orientation

all right well that's backwards um

one more try and then we will call it

good okay

size and orientation all right that's

pretty good so that's my first

vector my first bond dipole moment

and then i'm going to take the second

guy and i'm going to try to draw it tip

to tail so i'm going to wherever my

first vector wound up

i'm going to take the tail of my next

vector and start right there

and i'm going to try to draw in the same

size and orientation as my first one

um and also itself um so uh

that was well you know close enough all

right i'm gonna just get rid of this little

little

um arrow just to make things a little

bit clearer

okay so the last thing that i'm going to

do is try to figure out what was my

net what's the net distribution of this

charge right is there a

net difference in how these electrons

are distributed through the molecule

as represented by my vector and so i'm

going to draw

my net arrow from wherever my first tail was

was

and i'm going to draw it with the tip

ending where the the tip of my last

vector ended

and so here i see yes in fact i did have

a dipole moment right and it goes

straight up

it's exactly right what this is showing

right it's straight up through the

center of the molecule

um because there is a net displacement

of electrons towards this oxygen

so we can say that there is a partial negative

negative

on the end of the molecule of the oxygen

and there's going to be a net partial positive

positive

on the end of the of the molecule with

the hydrogens

so what we'll say is that this molecule

water is polar

it is a polar molecule it has a next

dipole moment

and co2 even though the bonds were polar

this molecule is linear the bond moments

are identical they're symmetrical

they're identical and opposite

all right so we're going to do a couple

more examples

just to let you practice with this

um so we're going to look at two different

different

tetrahedral molecules we're going to

look at carbon tetrachloride

which is whoops a carbon with four

chlorines around it

i'm going to try to draw this in three

dimensions i'm going to ignore the lone

pairs just for clarity here but there

are lone pairs around

okay and then i'm going to draw a

a different molecule which

is this guy where there's three

hydrogens and just

okay so we want to figure out what are

the net dipole moments on these two molecules

molecules

so um i will just go ahead and tell you

that chlorine is more electronegative

than carbon

you're going to need to know that in

order to draw all of these dipole moments

moments

i'll go ahead and draw them in black

actually i'm going to draw them in

a different color for now so you can see

them so

uh all of these bonds are going to have

a bond dipole

because chlorine is more electronegative

than carbon

so we can go ahead and draw that and

it's going to be

some arbitrary length um it's not super crucial

crucial

as long as we try to draw them all about

the same and on the same relative scale so

all of these bond dipoles are between

carbon and chloride so they're going to

be about the same

strength here we have carbon in the

chlorine so we're going to have that

same kind of dipole moment i'm going to

try to draw it around the same

and then between hydrogen and carbon we

actually have a pretty tiny dipole moment

moment

bond dipole between the hydrogen and the

carbon and it's actually going towards

the carbon

because the carbon is more

electronegative than the hydrogen

so we have these little baby arrows that

point towards the carbon

and again the the actual size of the

arrows isn't too important as long as

the relative magnitude is around the

same so

the carbon chloride bond dipole is

stronger than the carbon hydrogen bond

dipole so i need to draw that arrow

bigger than i draw the carbon hydrogen

bond dipole moment

okay now i'm going to try to add them together

together

and this is where things get really

wonky because if you're thinking

in three dimensions um it's it's sort of hard

hard

to think about this right so we can try

to draw this tip to tail

and we can absolutely do this and i will

try to do this and it's going to be ridiculous

ridiculous

but we can also sort of look at these

and see

hey this carbon has four bond dipoles

and they are distributed evenly around

that carbon

they're all pulling out towards those chlorines

chlorines

but they are all symmetric so

if we uh use a an analogy which kind of

works if we think about this as a tug of war

war

the chlorines are all trying to pull on

the carbon

and nobody's winning there's no net tug

of the electrons anywhere and so we're

going to say that this

and this is going to be a nonpolar molecule

and this is true anytime you have a

tetrahedral thing with four identical

atoms attached to it um this is also

true if you have a trigonal bipyramidal

thing with five or six

if you have octahedral with with

identical atoms symmetrically arranged

around it

it will be totally nonpolar because

there is no net dipole moment

so i'm going to go ahead and try to draw

these bond dipoles just to

attempt to convince you that this is in

fact the case um but

bear with me it's going to be ridiculous

all right so i'm going to start with the

two easy ones that are um

in the plane of the board so we've got

one going straight up

and then we have our next guy going sort

and then we've got this guy which is

coming kind of out at us

so i'm going to try to draw it with a

bit of a wedge

but it's sort of weird and this is all

kind of wonky right this is in space

and then we have our last one which is

kind of going

back in

and if you can kind of imagine those four

four

vectors all do add up to zero you have no

net electron tug there's no net

partial positive or partial negative

within this molecule

all right so now if we look at this guy

um we can see that these are definitely

not symmetric right if you just sort of

look at it you can imagine that

well this chlorine is pulling up right

this chlorine is pulling

this vector is kind of pulling straight

up and then each of these

hydrogen carbon bonds um maybe it helps

if we look at this

uh model we've got these three

guys on the bottom are kind of hydrogens

they're all sort of equivalent

they're all kind of pushing

symmetrically in at this

carbon so there's actually kind of a net

symmetric tug up from or push

up from these guys they're all pushing

electron density at the carbon or

or allowing the carbon to take the

electron density is another way to think

about it

if i try to draw this it's going to be

ridiculous again but i will give it a shot

shot

so we'll have this guy which is in the

plane of the board this little vector

and then we have

this one which is kind of going back away

away

it's also very small and then we have

our last guy which is coming kind of

um back straight in

ah so i'll try to draw it

kind of i don't know this way sort of

i don't know anyway it's basically a net upwards

upwards

from these three things kind of

contribute a net upwards push

and then this chlorine just pulls

straight up so this actually does have a

net dipole moment and it's going to be

straight up through the molecule

in the direction of the chlorine which

is the more electronegative atom

tetrahedral things are hard to draw they

can be difficult to visualize

if it helps you to make a little model

out of marshmallows and toothpicks is

one way you can do this at home

or balloons you can tie some balloons

together and get kind of a model

if it helps you to have a little model

all right so i'm going to talk about one

more molecule which looks like it should

be symmetric but it isn't

um and that is nh3 ammonia

when we draw this guy what we're going

to find is that it

is essentially there's a lone pair and

all we'll just pretend that that's kind

of in the plane of the paper

then there is a hydrogen there's a

hydrogen coming kind of straight out at us

and then there's another hydrogen kind

of coming down so again think of it like

this right this is tetrahedral except for

for

the molecular geometry is trigonal

pyramidal and these angles are slightly

less than 109.

okay if we draw our bond angles

we might be tempted to say listen this

is pretty darn

um this is this is really quite

symmetric right there's three hydrogens

nicely arranged around the central

nitrogen right it looks pretty symmetric

to me

the problem is that um they're all going

in one direction

it's not complete symmetry so all of

these bond dipoles are pointing towards

the nitrogen

because it is more electronegative than

the hydrogen

and essentially what happens is we have

that same kind of thing that was

happening above

where the three hydrogens are sort of

all pushing electron density or rather

allowing your electron density to spend

more time around that central atom

and when that happens essentially um there's

there's

no there's nothing coming to sort of

balance out that net

upwards uh deviation of the electron density

density

and so this actually does have a net

dipole moment and it's straight up

through the molecule just like we saw

with the

um with the molecule previously and

and so we can actually say that this is

a polar molecule it has a partial negative

negative

at the nitrogen end and it has a partial

positive down

um around the hydrogen end of the molecule

all right so the last thing that i want

to talk about with polar molecules

is the properties of polar molecules so

there's a couple properties that we are

interested in as chemists

um and one of these is that

if you just have polar molecules

floating around in space

they tend to just orient themselves randomly

randomly

but if you apply an electric field you

can actually orient all of the molecules

in one direction

which is is pretty cool they all line up

so that the uh the positive ends are

aligned with the negative side of the

electric field and

the negative ends are aligned with the

positive side um

and the fact that you can do this maybe

gives us a little bit of insight into

some of the properties that our polar

molecules might have

when we're just thinking about how those

polar molecules interact with other molecules

molecules

um and it turns out that one thing that

we'll think about

is is exactly that we'll think about it

much later this semester but

we'll think about intermolecular forces

and the forces of attraction that we

might see between

molecules and polarity is going to be

one of the key things that we think

about when we get there

um for now i think it's enough to say that

that

like molecules interact with like molecules

molecules

so polar molecules tend to dissolve

polar molecules in other words polar

molecules tend to interact more strongly

with polar molecules and you can get a

bit of a sense for why right you've got

this electrostatic interactions hey

um maybe that's a maybe that's part of

what this is so if you have a polar substance

substance

um and you want to dissolve it a polar

solvent a

polar liquid to dissolve that polar

substance is a pretty good a pretty good

um this is light dissolves like this is

why oil and water don't mix

oil is nonpolar fairly and water is polar

polar

um so polar things

do not dissolve non-polar things

in general so um for now you can just

remem remember that like dissolves like

uh it's a good

shortcut for this and we'll talk about

the reasons behind it

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YouTube Transcript:Chapter 4.6c Polarity

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YouTube Transcript:
Chapter 4.6c Polarity