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Chapter 4.6a VSEPR theory

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welcome to the first video for chapter 4

section 6

molecular structure and polarity in this

video we'll be focusing

on uh the molecular structure part we'll

be predicting the

structures of small molecules using

valence shell electron pair repulsion or

vesper theory as we jump in

it's important to just define a couple

of things just to make sure everybody's um

kind of on the same page with this so a

bond angle essentially measures

the angle between two bonds around a

central atom

so around the central carbon for example

here in this formaldehyde molecule we're measuring

measuring

the um this bond angle is measuring the

the angle between these two hydrogen

carbon bonds

usually we'll measure this in degrees we

can also measure

bond lengths and that is essentially

the length of a bond and it's measured

from the center of one atom to the

center of the next atom over

this is usually given in angstroms which

an angstrom is the one angstrom is 1

times 10 to the negative 10th

meters you can also measure bond length

in picometers sometimes

and 100 picometers

is equal to one angstrom

so uh angstroms are useful because they

give us something in the

in the region of uh single digit

bond lengths um picometers you have hundreds

hundreds

of picometers for a bond length all

right so let's dive into vesper theory

so vesper theory so again that's v-s-e-p-r

v-s-e-p-r

theory is basically

a theory that allows us to predict the

geometry of various molecules

uh it works on the premise that

electrons repel each other and molecules

will adopt

a conformation which will minimize that

electron or electron repulsion

by maximizing the distance between

regions of electron density

so essentially what this means is that

lone pairs and

bonding pairs of electrons will get as

far apart from each other around

the central atom in three dimensions remember

remember

as they can in order to minimize the

repulsions of

these electron pairs this ignores

other types of interactions between

nuclei for example so it includes nuclear

nuclear

nuclear repulsion and it also ignores

nuclear electron attraction

so it only focuses on electron electron repulsion

repulsion

but it usually gets the geometry pretty

close anyway so it's a fairly good

model for us to use

um i'm going to show you an example of

uh the vesper theory in practice

looking at um beryllium

dihydride uh and so

what this looks like essentially is

there's a beryllium

in the center and it's attached to two hydrogen

hydrogen

atoms and um beryllium

is uh is in group two and so it can

uh it doesn't have to obey the octet

rule it can

have in fact um only four electrons

around it as is shown in this molecule

um so for example when we're looking at

this we can think about

uh the vesper theory and what it'll

predict and essentially there's two

regions of density

there's two regions of electron density

here and

um and these two regions will repel each

other and what that means is that

they're going to get as far apart from

each other around the central beryllium

as they can and since there's only two things

things

the furthest apart from each other in

three dimensions that they can be

is 180 degrees apart

so in other words vesper theory predicts that

that

this bond angle will be 180.

we can do some similar thought

experiments and think about how far

apart other

numbers of electron regions can get from

each other

um and essentially this chart summarizes

that so if we have two regions of

electron density

um so that's either two bonds or two

lone pairs

those regions of density will get as far

apart from each other and that

will denote a bond angle of 180. if you

have three regions of density

the furthest apart three things can be

in three dimensions is 120 degrees

if you have four regions uh that is uh

the the bond angle

is limited to 109.5 and then if you have

five regions of density

uh you actually adopt a different

conformation where you have

two different bond angle possibilities

so essentially you have

the sort of triangle around the middle

and that is similar to this

this three regions of density around the

middle so those are going to be 120

degrees apart from each other

and then you have these two axial

positions which will be 90 degrees apart

from the triangular on the middle or the

equatorial positions

when you have six regions of density all

the bond angles are

90 degrees uh here uh in this next row

uh there are some uh molecules shown

that that have the

uh the geometry uh that's that's shown

above the spatial arrangement of these

regions of density

um these are shown uh in three

dimensions as best we can

on a two-dimensional piece of paper or

screen um

so some of them work quite well so if

you have just two regions of density

uh 180 degrees apart that works quite

well on a flat sheet of paper

because 180 degrees is still you can

represent that on a flat piece of paper

and same thing with this with three

regions of density 120 degrees

that's a flat plane um things get more

complicated as soon as you have four or

five or six regions of density and so we

use this notation that we call

line dash wedge essentially

what the way that you're going to want

to interpret these is that the lines

are in the plane of whatever you're

looking at so the paper or the

the screen your computer screen that

you're looking at and then anything that's

that's

dashed is going to go back away from you

into the screen

and the wedges are going to be coming

out at you

um so i have a little uh model of of a

tetrahedral thing here with four regions

of density around the central

uh this represents a carbon atom um and

so essentially what you're going to want

to do

is just sort of imagine lining up two

of these things so here the hydrogen

bond and the hydrogen bond

so we'll just go ahead and call that

this white ball and this green ball will

just line up in the

in the plane of the um the screen of the board

board

and then um from your perspective this red

red

ball is going to be coming out at you

from the screen

so uh that would be a wedge

and then this blue uh sphere this blue

one which uh represents usually a sulfur

or some other atoms

is going to be coming back into the screen

screen

so that's going to be represented by a dash

dash

um so that's what that line dash

wedge notation means it may take you a

minute to sort of wrap your head around it

but do spend the time especially this is

a pretty good diagram here

uh it it sort of you can think about the

the dashes and the wedges

um in in relation to these these

triangle ones and the same thing with this

this

this thing where there's a square around

the center around the uh the equator of this

this

of this diagram and that's represented

by these two dash

uh sorry these two wedges and then these

two dashes

so do spend the time to get used to that

uh line dash wedge notation because

you'll be seeing it a lot since we live

in a three-dimensional world but have to

represent it

all right so the next thing that we're

going to think about is what happens if

all of the

things that are attached to the central

atom all of the regions of density aren't

aren't

uh atoms aren't bonding pairs and

essentially we're going to have to define

define

two different types of geometry so we're

going to have the

uh or we'll often abbreviate this as epg

um this can also be called electron

domain geometry

which is going to be edg it's the same

thing electron pair geometry or electron domain

domain

um geometry and essentially this is a

kind of geometry that describes the location

location

of any electron pair all electron pairs

so for example it doesn't matter if these

these

uh these uh electron pairs are

bonding pairs or lone pairs we'll

describe the geometry

using this sort of whole shape in mind

and then if we if we don't want to think

about that if we don't want to think

about the placement of all the electron

pairs but we're really only interested

in the arrangement of the atoms

then we'll be dealing with something

and we will often call that mg for

molecular geometry it's important to

note that this still depends

on the placement of lone pairs right um

just because some of these turn into

lone pairs doesn't mean that the shape

changes it just means that the the way

that we describe the shape changes

because we're going to

ignore uh some some some portions of the shape

shape

but the bond angles are still going to

be um

impacted by the presence of that lone

pair um just because we're not

describing the lone pair doesn't mean

all right so some more important things

about vesper are that uh

it it understands that certain types of

repulsions are going to be

um stronger than others so

this is the order of repulsions

according to vesper theory

and my notation here is lp is lone pair

and bp is bonding pairs so essentially

lone pairs

repel each other more than bonding pairs

and lone pairs repel each other more

than they repel bonding pairs

so if you have two lone pairs repelling

each other that repulsion is stronger

than if you have a lone pair repelling a

bonding pair or a bonding pair repelling

a bonding pair

um since the geometry and the amount of

space that these electron pairs take

up is dependent on the strength of the

repulsion in

uh in vesper theory essentially the

stronger the the repulsion

the repulsive push is um this the more space

space

an electron pair will generate for

itself in in three dimensions

um the the volume taken up by

each of these uh types of electron pairs

is different

so um so this is going to be the order of

uh the size of the different kinds of

electron pairs so a lone pair is going

to take up more

space than a triple bond which takes up

more space than a double bond which

takes up

more space than a single bond um and so

this is interesting because if we're

thinking about our molecular geometry

and our electron pair geometry

uh the presence of a lone pair can

actually change the bond angles

so one way to think about this is is

with ammonia or one example of this is with

with

ammonia so we've got a nitrogen in the

center with three

hydrogens um this is tetrahedral

so this is this four region of density

that we saw up here before where the

bond angles are theoretically

109.5 but what happens is

if one of these things is in fact a lone pair

pair

um so let's talk about this this white

ball on the top being actually a lone pair

pair

rather than a bonding pair um and this

isn't an atom anymore it's just a lone pair

pair

it's actually going to take up more

space than any of these bonding pairs around

around

the bottom part any of these uh that

would be bonded to hydrogens

so essentially when that happens each of

these bonds gets pushed

down a little bit because this lone pair

is taking up so much extra volume

and it actually makes the bond angles

slightly less than

109.5 which is the ideal geometry

the ideal bond angles um for a

tetrahedral structure with four

four regions of density

so essentially what that means is that

we have to think about

uh we have to think about both the

number of regions of electron density

and the number of lone pairs that are

present when we're thinking about the

molecular geometry in the electron

domain geometry

um and the the bond angles that were

that are present in

in all of these structures so there's a

chart here that i think is a really helpful

helpful

thing to use and this is in your

textbook so if you can't see this in

in high quality then go ahead and look

in your textbook and you'll be able to

zoom in there

um but the way that you use this this chart

chart

is uh you first determine the number of

electron pairs or this is

so you're going to count up how many

total bonding regions or lone pair

regions there

are remember that double bonds or triple

bonds count as one

region of density because they're in the

same kind of area

around the molecule so you're going to

go ahead and count up how many regions

of electron density you have

and then from there you're going to have

to determine the electron pair geometry

and then the molecular geometry

thinking about how many lone pairs there

are so as soon as you know how many electron

electron

density regions there are then you know

the electron pair

geometry but you may not know the actual

bond angle until you figure out the

molecular geometry

so this is a really useful chart um you

will need to know

all of the names of these things so

something with only two regions of density

density

180 degrees apart linear something with

three regions of density

and three the electron pair geometry is

going to be trigonal planar

we'll often call this trig planar

something with four regions of density

is going to be called tetrahedral and

then i don't have models for

uh when there's five and six the

hypervalent uh

scenarios but if you have five regions

of density the electron pair geometry is

trigonal bipyramidal

and if you have six regions of density

um the electron pair geometry is octahedral

octahedral

so these uh that's the electron pair

geometry right

this column is the electron pair geometry

geometry

after you determine the electron pair

geometry then you need to think about

are any of the regions of electrons actually

actually

lone pairs instead of bonding pairs

responded to another atom

so essentially you're replacing this x

this terminal atom whatever that is with

a lone pair

so in the case of a trigonal planar

molecule if we were to replace one of these

these

with a lone pair then we would call that

um we would just describe the shape of

those two

um atoms around the central atom

and we would call that bent or angular

is sometimes used for that

and basically it looks something like

this um

right so there will be an imaginary lone

pair on top um but the angle is going to

be in the ballpark of 120 but because

the lone pair takes up that extra little

bit of volume

and pushes these bonds closer together

the angle is going to be slightly less

than 120.

so i will let you go ahead and look at

the rest of these you should again you

should know the names

of of all of these both electron pair

geometries and molecular geometries

you'll be responsible for knowing those and

and

when we get into polarity uh your these

these geometries are going to be

absolutely critical um

they will make differences in the in the

polarity of molecules

so um before we move totally on from

here i just did want to mention

what happens when you replace

one of these terminal atoms with a lone

pair but the axial and the equatorial positions

positions

are not equal like when you have say a

trigonal bipyramidal shape

where some of the bond angles are 120

and some of them are 90.

it's really important to keep this in

mind because

the lone pairs need to occupy the

position with least repulsion

because again that's kind of the

underpinning of vesper theory is that

uh the electrons will the the geometry

depends on the electron pair repulsion

and lone pairs repel each other more

than bonding pairs

so this is what's shown here down here

you're only going to see this when you

have hypervalent molecules

and you'll see this particularly in if

you have five regions of density

when you have one two or three

lone pairs the equatorial positions

disappear first

because those are the more spacious uh

positions for a lone pair to occupy

and you get this saw horse your seesaw

and then you get t shape and then you get

get

um finally a completely linear molecule

if you have

only two actual atoms bonded but three

lone pairs

around that central atom with the um

with the type of hypervalent molecule

with six

regions of density around the center

you'll see this again

uh essentially it doesn't matter anymore

all of these things are 90 degrees apart

but you'll see it when you have

two lone pairs here they will occupy

equatorial positions directly opposite

from each other since that lone parallel

and pair repulsion

is stronger than lone pair bonding pair

the lone pairs will occupy the position

uh that are 180 degrees apart from each other

other

um and then um yeah

so uh you can imagine this there's some

drawings in your book

um and i've i've copied and pasted one

in here just so you can see it

um this is your trigonal bipura middle

structure so this

this green shaded triangle is flat right

across the middle and then you have the

two pyramids on the top and the bottom

and this describes essentially um what

is axial and what's equatorial and it

kind of

hopefully sort of shows you why these

equatorial positions

are more spacious and this is the real

uh geometry this

this um this t-shape geometry is the

real uh

orientation of this of this molecule

with um

with three fluorines and two lone pairs

around the center chlorine atom

whereas c and d are options but are in

fact not

all right so the last thing i'm going to

do here before i move on

is to just basically sum up the steps to

predict the geometry

so the first thing that you're going to

have to do when you're predicting the

geometry of a molecule is draw the lewis structure

structure

once you've got the lewis structure then

you're going to need to count the

regions of electron density

once you've got the regions of electron

density you'll be able to identify the

electron pair geometry

once you have that then you're going to

think about how many lone pairs you have

and use that number of lone pairs to

identify the molecular geometry

which will tell you the final bond

angles um

in addition to the name of the molecular

geometry and electron

geometry for that molecule there will be

plenty of examples

in future videos so we'll be using these

four steps

in future videos so that you can get

some practice

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Chapter 4.6a VSEPR theory