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Chapter 4.6b Geometry Practice

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welcome to the second video for chapter

four section six on molecular geometry

and polarity

in this video we'll be focusing on

practicing predicting geometry using vesper

vesper

the learning objective is to predict the

structures of small molecules using

valence shell electron

pair repulsion theory which is also

known as vesper theory

in a previous video we talked about some

steps to predict the geometry of some

molecules and the first step of this is

to draw the lewis structure for the molecule

molecule

i'll be working with these molecules

down here

um so what i recommend you do is pause

the video at this point

and spend some time drawing the lewis

structures for these molecules

um and in one ion and just make sure

that you're comfortable

drawing that lewis structure alanine is

a more complex molecule which we will be

talking about at the end

so um if you want to look at the

structure for that molecule you can or

just wait until the end and we'll talk

all right so once you've got your lewis

structures drawn we will go ahead and

walk through how to predict the geometry

of each molecule

our first molecule is water i have drawn

all the structures here with basically

no attempt at representing the shape

and i've done that on purpose so that

you can kind of see

how you might go about drawing these

with some sort of attempted shape

to represent the actual reality of the

bond angles etc

all right so our first step after we've

drawn our lewis structure is to count

the regions of electron density

my favorite way of counting regions of

electron density is trying to draw a

circle around the region

and if i can do that without too much

trouble i'm trying not to you know go

out of my way to include extra electrons

or have to be really focused to avoid

certain electrons

then that's probably one region so

single bonds are pretty clearly

a region of electron density so so far

we have two regions of density

and then each lone pair is also a region

of electron density

so all together we've probably got four

regions of density here

the next step is to use that that count

of regions of density to identify the

electron pair geometry for this molecule

i can do this either by having memorized

um what that means how many how many

regions and what shape

corresponds to that number of regions

which you should do

if you haven't memorized it yet then

we'll just go ahead and use our chart

um so in your textbook there is this handy

handy

chart and if you can't see it well

enough it's in your textbook you can go

ahead and zoom in on it there

the way we're going to use this chart is

by first identifying the number of

electron pairs what this means is the

number of regions of electron density

where there are electron

pairs and we're going to come on down to

this row where we've got

four regions of electron density this

means that our

electron pair geometry is tetrahedral

so we'll go ahead and write down

tetrahedral i'm just going to write down tetra

tetra

as an abbreviation the next thing we're

going to do

is use our lone pair count to see

if the molecular geometry is different

from the electron pair geometry which i

should note here this

is the electron pair geometry so we're

going to see

um if we have any lone pairs that our

molecular geometry will be different

from our

uh our molecular geometry will be

different than our electron pair geometry

geometry

and again that's because uh when you are

describing the shape for the molecular

geometry you don't include

the lone pairs in the description of

that shape

so we've got two lone pairs here so

we'll come back over to our chart

we'll come back down to our our row of

four regions of density and then we'll

go across until we find the spot where

there's two lone pairs

so this tells us that our shape is bent

or angular

and that our bond angle is less than 109 degrees

degrees

so we'll go ahead and write that down

our molecular geometry is bent

and our bond angle is less than 109

degrees all right so now let's go ahead

and try to draw this guy

in some kind of realistic fashion

so we'll start off with our central atom

and then we're going to just go ahead

and try to draw our hydrogens with some

sort of angle

of around 109 degrees um that's an estimation

estimation

and then we'll just put our two lone

pairs um also at something

estimating 109.5 degrees those guys are actually

actually

turned um they are coming into and out

of the page but since they're lone pairs we

don't care too much about representing

them you can draw them at the end of

dashes and wedges if you

if you do want to represent that so i'm

just going to note here that this angle

is less than 109 degrees

all right so we'll move on to our next

molecule which is uh

nitrite the anion one of your polyatomics

polyatomics

so we talked about this one in a

previous video when we were discussing resonance

resonance

and uh so this this molecule or this ion has

has

resonance you can draw it in two

different ways which means that neither

of these is actually the real

representation of the shape of this

molecule but rather

um somewhere in the middle right to an

average of these two so the double bond

isn't here or here it's somewhere in

between it's kind of both at the same time

time

um this is one of the reasons that we

like to draw resonance forms or

resonance contributors because

it helps us understand um the shape

without having to kind of go through the

process of understanding that partial

double bond thing that thing that's not

quite a single bond but it's not really

a double bond either

instead we can just consider the shape

of each of these resonance contributors

and um they will actually be the same

and it turns out that is actually the

shape of this molecule

so we're going to start out the exact

same way we just count up our regions of density

density

so a double bond is one region um i

would have to try pretty darn hard to

circle only some of those electrons so this

this

whole double bond is one region uh our

single bond is a second region

and then our lone pair on that central

nitrogen is a third region

so since we've got three regions of density

density

we will go over to our little chart and

find out that that means our electron

pair geometry

is trig planar or trigonal planar

but i'm going to abbreviate it trig whoa

oh geez pen pen malfunction trig

planar all right

now we're going to look at our lone

pairs and

and uh use that to understand if and how

the molecular geometry is different than

the electron pair geometry

we have a lone pair so it's going to be

different we'll go back over to our chart

chart

um and we come across from our in our

third row or

our row of three regions of density

across to one lone pair

and we find that this is also bent or angular

but the bond angle here is

actually less than 120 instead of 109.5

as it was

above so here we actually have two

different kinds of bent molecules here

we have

um tetrahedral bent with the bond angle

less than 109

and down here we're actually going to

have trig planar bent with the bond

angle less than 120.

so this is one of the reasons why it's

super critical to start with your

electron pair geometry and then think

about your molecular geometry

because the electron pair geometry is

what sort of defines your shape to begin with

with

and then how many lone pairs just kind

of modifies that shape it just modifies

the bond angles and

the name of the shape so if we're going

to draw this guy

um with some sort of reality i'll just

pick one of these resonance structures

to draw although we could draw

both or we could draw the resonance hybrid

hybrid

that would be the same thing but i will

just try to draw this guy

with some sort of representation of the

angle between these two

oxygens as something in the ballpark of

all right so we'll move on to our next

molecule which is carbon tetrachloride

so we'll start off the same way we count

up our regions of density

one two three four

and we can check on our chart but we've

already done one of these so when we

have four regions

we know that that means our electron

pair geometry

is tetrahedral the next thing we're

going to do is see if our molecular

geometry is different than our electron

pair geometry here we have

no lone pairs so our molecular geometry

is not different it is also

tetrahedral and therefore we know that

these angles

are 109.

okay so this guy's a little bit

challenging to draw with some sort of

semblance to reality

um and this is where our dashes and

wedges are going to come in really handy

so i've got this little molecule that

may help you visualize

what's happening here so what you're

going to do is draw your first

atom as your central atom and then

you're going to pick two

two of the surrounding atoms here are

chlorines to

to be in the plane of the paper or the screen

screen

here i'm going to draw this guy and this

guy in the plane of the screen

just because that is uh easy to me

um you can draw any two but whatever you

do they're going to have an angle

between them that's approximately

109 i'll draw the lone pairs on

at the end the next thing you're going

to do is think about

that there are two molecules or two two

atoms at the end of the bond

um these guys are essentially in a plane

that's 90 degrees off

from these guys they're exactly

perpendicular to the screen except

they're also tilted a little bit

so we're going to use our dashes and

wedges to um to sort of understand that

so i'm going to draw a wedge here for

this chlorine and then i will draw a dash

dash

going back for that chlorine and then

i'll just go ahead and add in my lone pairs

pairs

to represent this structure and all of

these bond angles

all right so then we'll move on to our

next molecule which is

iodine pentafluoride

this is a hypervalent molecule it's

clearly got

more than eight electrons around our

central iodine but luckily for us that

doesn't actually change anything about

the way that we assign geometry

we're gonna start off the same way just

counting up the regions so we have one

two three four

five regions um that are bonding and

then we have

a sixth region that's a lone pair so

and we can go to our chart and discover

that that means our electron pair

geometry is octahedral and that all of these

these

species should be something in the

ballpark of 90 degrees off from each other

the next thing we're going to do is use

our lone pair count

to determine if our molecular geometry

is different and since we have a lone

pair it will be different

and that single lone pair means that our

actual geometry or sorry our molecular geometry

geometry

is square pyramid square pyramidal

square pyramidal or square pyramid

um and i'm just going to abbreviate that

square here

okay and our bond angles are actually

less than 90 degrees

because that lone pair squishes

everybody down a little bit

okay so here's another one that's a

little bit challenging to draw

um but we'll give it a shot so we're

going to go ahead and start off same

thing with our central iodine atom

and then the easiest way to do this is

to pick

one guy to be axial

and the other four will be equatorial

and then you can draw the axial one

either up or down

it's the same thing so i will go ahead

and draw this with my axial fluorine

going up

and then the other four fluorines are in

a plane that's perpendicular to the screen

screen

so they're coming like straight in and

straight out so i will go ahead and draw

two of my fluorines going straight back

but at an angle into the plane so i

guess for you guys it's back coming

towards me

and then um the other two are coming straight

straight

with wedges and then i'll go ahead and

and then of course i've got a big lone

pair hanging out in the other axial

position on the iodine

so all of these angles are less than 90

degrees both the equatorial

and the between the axial and the equatorial

all right so we'll move on to our next molecule

molecule

and that is carbon dioxide so hopefully

this guy is familiar to you

at this point we've talked about it a

few times so we're going to start off

exactly the same way

count our regions two regions

so two regions means that our electron

pair geometry

is linear

um we can't have uh

yeah when you have only two regions of

electron density you can't have one of

those be a lone pair because

then you can't it's not a central atom

anymore um so this is

linear and and you can actually see this

on your chart there is no such thing as

a lone pair for a linear

central atom so um so we're done and the

bond uh bond angle is 180

so in fact we don't have to rewrite this

one because it's linear and we tend to

write it as linear

all right our next molecule is

phosphorus pentachloride so again

another hypervalent guy

so we'll just start off the same way as

normal and we're going to count up our

regions of electron density

one two three four

five so we've got five regions

so our electron pair geometry is going

to be we can check our chart

trigonal bipyramidal um

so i'm going to just abbreviate that as

trig by here whoops

pure all right and then we check to see

if we have any lone pairs

uh we don't so that electron pair

geometry is the same thing as our

molecular geometry

and here we actually have a bit of

interesting geometry the

um we have some axial positions and we

have two equatorial positions as well

and um and they're going to have

different bond angles from each other

luckily this is a little bit easy

because we don't have any lone pairs our

electron pair geometry is the sort of

simple model version

so we'll go ahead and try to draw this

before we describe the bond angles

all right so we'll start off with our

phosphorus in the center

and then we're going to have um kind of

the same way that we had

our we had an axial and we had

equatorial positions for

the uh the octahedral guy with with six

regions of density around it

we're going to have axial positions here

as well and then we'll have equatorial positions

positions

so i'm going to go ahead and start off

by putting my axial positions in the

plane of the board

and then my equatorial positions that's

where the other three chlorines are

going to go is directly around

this in kind of a again a flat plain

sort of around

um and so i'm going to choose one of

these guys to be

in the plane of the screen or the paper

and go ahead and pop in that chlorine it

doesn't matter which way it goes

and then i will put the other guys on

the other side

one of them is going to be going back

into the screen at an angle or i guess

sorry back into the screen at an angle

and then the other one would be coming

straight out of the screen

at an angle so i'll go ahead and

and then we'll talk about our bond

angles so between

the um axial position and the equatorial plane

plane

is 90 degrees but between all the positions

positions

in the equatorial plane it's 120.

so here's where you have this uneven

all right and then the last thing we're

going to do is talk about alanine

and so essentially the reason i'm

talking about alanine is because

it's going to help us kind of understand

what to do when you don't have one

single central atom

when you've got more than more than one

single central atom

and essentially the answer is you're

going to consider every single central atom

atom

individually so we're only looking at

local geometry about

each of these atoms um there is

especially with proteins if you go on to

biology biochemistry especially you'll

find that we'll talk about larger structure

structure

um but we're not going to cover that at

all in gen chem we're going to be

focusing on local structure

so i'm going to focus on just a couple

of atoms which i will circle we're going

to focus on

this carbon this nitrogen

um and then we'll do this carbon as well

okay um yep okay

so uh we're going to focus on this

carbon over here

first we'll go ahead and figure out what

its geometry is

we've got one two three four regions of

hey that's tetrahedral it doesn't have

any lone pairs

so its molecular geometry is identical

this is just a tetrahedral carbon

it turns out that we can go ahead and

just see that this guy is also tetrahedral

tetrahedral

since we've just done that one this guy

also has four bonds around it

all right now let's focus on this

nitrogen this nitrogen has one

two three bonds and one lone pair

so this nitrogen has four regions of density

density

so it's electron pair geometry is tetrahedral

tetrahedral

tetrahedral but its

molecular geometry is not tetrahedral

since it has that one lone pair

so if we come over here we can see that

when we have four

uh regions of density but one lone pair

that's trigonal pure middle

and so that lone pair sort of pushes

everybody down a little bit more and the

bond angle is less than 109.

so i will abbreviate this as trig pier

okay and then we'll come and look at

this carbon over here

we've got one two three regions of density

all right so if i'm going to try to

redraw this it gets a little bit complicated

complicated

so what i'm going to go ahead and do is

draw the carbons

in the plane of the board so i'll start

with this guy

and then i'm going to go ahead and say

this hydrogen

and then this carbon-carbon bond is in

the plane of the board

this carbon um has another

thing that's in the plane of or the

screen which is the other

bond to this carbon and then this carbon is

trig planar so since i've drawn it with

one bond in the plane of the

screen or the paper or the board or

whatever um then everything is

in that plane and then this guy is

actually bent as it turns out so we'll

okay so now i'm going to come back to

this tetrahedral carbon

and um sometimes it can help to number

things to just not lose tracks we'll go

ahead and do that

one two three carbons so i'll just go

ahead and

one two three okay so this carbon here

is carbon number one it's tetrahedral

um i've drawn in one of its hydrogens

and so the other two hydrogens need to be

tetrahedral so coming out and going into

the board

okay this guy i have uh drawn

with two bonds in the plane already and

they're going this way so the other two

need to be going

sort of the opposite direction of them

maybe it's better if i use this

so right now i i've got this and this in

the plane

of the screen so i have one guy going

this way and the other guy going this way

way

and i'm just going to go ahead and

decide that the nitrogen

is the guy that's coming out at me

and the hydrogen i will try really hard

to draw going back

in some sort of way okay

so that's carbon number two and then i

will go ahead and draw my nitrogen

um so my nitrogen center is uh

um is tetrahedral but it's actually trig pyramidal

pyramidal

so i've drawn one bond kind of this way

and i'm going to go ahead and uh just

sort of draw another hydrogen kind of

coming out

and then one hydrogen there and then my

last lone pair will be sort of

up and this is really hard to imagine

um and then that's my alanine molecule

so you can see this gets a little bit complicated

complicated

the more things you add to it the harder

it is to draw in three dimensions

which is why sometimes we don't try and

we represent them with these sort of 90

degree angles that

don't represent reality all right so

hopefully that's

given you some practice on assigning geometries

geometries

um if you still need help there are

plenty more problems in the back of the book

book

or just google geometry practice and

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YouTube Transcript:
Chapter 4.6b Geometry Practice