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Chapter 4.2b Electronegativity

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welcome to the second video of

uh chapter 4.2 covalent bonding in this

video our learning objective is to

define electronegativity and assess the

as we jump in let's go ahead and first

define what's the difference between

a pure covalent versus a polar covalent bond

bond

so a pure covalent bond

is a bond with perfectly equal sharing

of electrons between the atoms that make

up that bond

which is to say the atoms both have

equal electron density or

an equal uh an equal distribution of electron

electron

density around them a polar

covalent bond is

a bond with unequal distribution of electrons

electrons

around the two atoms and this is

characterized by

a positive end and a negative end uh

which is

gives us the sort of terminology of

poles right so polar covalent there's one

one

end that's positively charged and the

other end is negatively charged

so uh we can see these uh this depicted

here with an hcl atom

or molecule sorry an hdl molecule of

hydrogen chloride or hydrochloric acid

if you've dissolved it in water

um it's really important to note however

that we're not dealing ever with camilla

bonds we're not dealing with full charges

charges

we're dealing with partial charges and

so we need to use some notation that

tells us that these are partial charges

as opposed to

full charges like we see with ionic

compounds and that notation is this

lowercase greek delta um

so this basically just means that this

chlorine atom has

a partial negative charge and this

hydrogen atom has a

partial positive charge this is shown

here with this difference in electron

density around

each of these atoms so the chlorine side

has more electron density around it

the electrons are spending more of their

time around this chlorine atom

and again because electrons are

negatively charged that means that the

chlorine has a partial negative

charge and when the electrons are

spending most of their time around the

chlorine they're

not spending very much time around the

hydrogen which leaves that with a

partial positive charge

um and uh and so we have these two poles

this arrow notation is a nice way for us

to express this quickly

the positive uh end of the molecule gets

the tail of the arrow which looks like a

little positive

and then we draw the arrow towards the more

negative atom um the one that has the

so we need a way to sort of describe

this uh in

in and if we can come up with some

numbers that's really helpful

and so uh a very famous chemist called

linus pauling came up with

which is a measure of the tendency of an

atom to attract electrons or

rather electron density towards itself

when it's in a bond so essentially this

is a measure of exactly what we're

hoping to describe here

um this ability of an atom to pull the

density towards itself and away from

um this is a chart of

electronegativities so

there's some important things to note here

here

first off there's no noble gases right

we don't see helium neon

and the reason for that is because they

don't tend to form covalent bonds

you can force noble gases to form

covalent bonds

if you put them in extremely uh well

extreme circumstances but in general

they don't tend to form covalent bonds

and so they don't fit neatly into the uh

into the model that we have

for um for electronegativity and so we

we just don't

include them when we're discussing

electronegativity the trend here

is uh that uh electronegativity

increases as you go

towards the right across the periodic

table and um

it decreases as you go down towards the

bottom of the periodic table

i like to describe this as a diagonal

trend um which means that if you can

remember that fluorine is the most

electronegative atom on the periodic table

table

then you know both directions right you

know that it increases towards the right

and you know that it decreases as you go down

so we need to keep in mind the

difference between electronegativity and

electron affinity

they are related but they're very

different terms and they mean very

different things

they sound very similar but they are

very different so electron affinity

which we talked about earlier

is um the energy released

or absorbed during an electron capture

so it is a physical measured thing it's a

and it represents an actual um

measurable physical quantity

electronegativity is not a measurable

physical quantity it's calculated

um it's unitless

and uh it's a on an arbitrary scale

between zero and four

and it essentially just describes how

tightly an atom

uh pulls electrons towards itself

as it is in a bond

all right um so we want to talk about

types of bonds so what

what makes a bond pure covalent versus

polar covalent versus

all the way to ionic when we have

complete electron transfer

and essentially it's the difference in electronegativity

electronegativity

and uh it's we're going to think about

the absolute value but it's it's the

difference in electronegativity

and so we're going to represent that a

lot by delta e

n and so delta is change or difference

and then en is just a representation of electronegativity

electronegativity

so the larger the difference in

electronegativity that is to say the

larger the

the difference in the ability of these

atoms to

hold electrons tightly towards

themselves in a bond

so the larger that difference the more

likely it is that one

is going to essentially win the tug of

war more often

and when that happens if you have a

really large difference in electronegativity

electronegativity

you're going to have either a polar

covalent bond or a full ionic bond where

the electron is just transferred

so um one example of um

uh these ionic or sorry uh polar

covalent bonds is

uh things between hydrogen and um oxygen

so hydrogen has an electronegativity of

2.1 oxygen has an electronegativity if

so the difference in these guys the

difference between those guys is

1.4 and what this tells us is that

one of these atoms is significantly more

electronegative it holds the electrons

in the bond much more tightly than the

other one

and it will uh basically claim the

electron density

much more often than the hydrogen is

able to when they are

sharing electrons in a bond so what this

looks like is for example if you have

water and we will talk about geometry a

little bit later

but this is a depiction of a water

molecule an oxygen with two hydrogens

bonded to it

we can see that oxygen has a greater

electronegativity and so we're going to draw

draw

this sort of arrow thingy up towards the

oxygen and actually each bond has

this sort of uh we we can represent this

uh this polarity on each bond

so the oxygen is when winning the tug of

war for electrons um versus the hydrogen

the density of electrons is going to be

focused around the oxygen

and it'll the density will be much lower

around the two hydrogens in that water molecule

molecule

if you have uh atoms that have a more

similar electronegativity

like for example two identical atoms

um then you have a pure covalent bond

where there's no difference in

electronegativity we don't draw any

polar representations there's no partial

positives or partial negatives

um and that essentially happens when the

difference in electronegativity is very small

small

um or the atoms are identical so this

table represents

some of the values that we have um

sort of set there it's a it's guidelines

there's exceptions to this for example

hf uh so

hydrogen has an electronegativity of 2.1

fluorine has an electronegativity of 4.0

the difference in these guys is 1.9 if

we just go by the table

um the difference is so big that this

would actually predict

that um the electron would just be

transferred in this bond would be

completely ionic

in fact that's not the case this is a

polar covalent

um it's just very polar um and the same

thing is true there are ionic

uh bonds that you would not predict

would be ionic you would predict they'd

be polar covalent because the difference

in electronegativity is too small

um so this is just guidelines but in

general um

you can think about how big the

difference is between

um the bonding atoms so if there's no

difference they're pure covalent bonds

if there's a medium difference then

they're probably experiencing a polar

covalent bonds and if there's a huge

difference in electronegativities

between those atoms

like for example between the metals and

the non-metals that we see up here

um the metals have very low

electronegativity values

then you are likely to see ionic bonds

in those compounds

all right the last thing we're going to

talk about in this video is polyatomics

so it's important for us to note that

you can actually have

ionic compounds that contain covalent compounds

compounds

and what that looks like is you can have

these polyatomic ions like for example nitrate

nitrate

where the the bonds between these atoms

are covalent bonds but you can also have

an ionic

bond or an ionic compound form if you

have an attraction between

a sodium plus ion and your nitrate

anion you can form sodium nitrate

which is a ionic compound but it

contains this

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Chapter 4.2b Electronegativity