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Chapter 4.5a Formal Charges

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welcome to the first video for chapter 4

section 5

formal charges and resonance in this

video we'll be focusing on formal charges

charges

and we'll be focusing on these first two

learning objectives which are given here

so you should pause and take a moment to

write these in your notebook

as we jump in let's first begin by defining

defining

formal charge so formal charge

is the hypothetical charge an atom in a

lewis structure would have if we could

redistribute the electrons in the bonds evenly

evenly

this is not a real charge it's not a representation

representation

of any sort of real distribution

of electron density but it's really

useful for us to think about

um both as just a bookkeeping measure

to make sure that we have the right

number of electrons and um

and it's also helpful for us to think

about how the electrons are distributed

in a molecule

and compare to the sort of electronic

structure of that atom

the formal charge i'm going to be

abbreviating through this video with

fc to represent formal charge the

equation for formal charge essentially

compares the number of valence electrons

uh to the number of atoms electrons that

the atom has

uh in in the structure that we're

looking at so we're going to get the

number of valence electrons that the

atom has

from the periodic table then we're going

to compare it

uh by subtracting the number of

electrons that the atom has

and what that means is we're going to

subtract the number of

and then we're also going to subtract

the number of electrons if we could

divide up the bonds evenly

so we'll take however many electrons are

in all of the bonds and divide those by two

two

because again those are going to be

shared between the two atoms that are

making the bond

so we'll subtract half of the number of electrons

electrons

a really important point here is that

the sum of any formal charges must equal

the charge

on the molecule or ion so if you're

dealing with a molecule the charge must

be zero so the sum

of all of the formal charges must be

zero or if you're dealing with an ion

the sum of the char

the sum of the formal charges must equal

the charge on the ion

we'll do an example uh we'll look at formaldehyde

formaldehyde

and hopefully you are familiar with the

lewis structure for this molecule already

already

if you're not go ahead and pause now and

try to draw it on your own

i'm going to go ahead and just draw it

out for us since drawing lewis structures

structures

is not the focus of this video

so we're going to go ahead and calculate

the formal charges for each atom in this structure

structure

we will note that the two hydrogens are identical

identical

in terms of their their situation in

this molecule they're both making one

bond to a carbon and have no lone pairs

so we can treat them equivalently so

let's start out by calculating the

the formal charge on these hydrogens

we'll start by figuring out how many

valence electrons the hydrogen has from

the periodic table

that's going to be one and then we will

subtract the number of electrons

uh in the lone pairs uh here

i'm sorry i'll scroll up so you can

actually see it so we'll subtract the

number of electrons in the

lone pairs in this molecule

the hydrogens have no lone pairs so we

will subtract zero electrons

and then we will subtract one half of

the number of electrons that are

in a bond or or in bonds

with this hydrogen there's one single

bond which is two electrons

so we will subtract half of two uh this

means that basically the formal charge

on this hydrogen is one minus one

which is zero all right

let's move on to looking at the formal charge

charge

so we'll begin again by figuring out how

many valence electrons the carbon has by

looking at the periodic table and we

will find that

carbon has four valence electrons then

we'll subtract

any electrons and lone pairs here that's

zero so we'll subtract zero electrons

and then we will also subtract half of

the number of electrons involved in bonds

bonds

and here we have four bonds that's eight

electrons so we'll subtract half

of eight so this simplifies down to four

minus four which is zero

so this carbon also has a formal charge

of zero

lastly we're going to do the oxygen so

we'll look at the periodic table and

discover that oxygen has six

valence electrons then we'll subtract

the number of electrons in lone pairs

here we've got two lone pairs which is

four electrons

and then we will also subtract half of

the number of electrons in bonds

uh here we've got two single bonds which

is four electrons

this uh simplifies down to

six minus four minus two which is

zero so this oxygen also has a formal

charge of

zero we should always take a moment and

check once we have finished drawing

uh finished our calculating our formal

charges to make sure that they all sum

up to the net charge on the structure

here we have a molecule with a net

charge of zero and it turns out that all

of our formal charges are zero so they

will all sum to zero

all right so the next thing we're going

to talk about is why this is important

to us

um and and there are several reasons why

we might care about formal charge

but one is that sometimes there are

different molecular structures that you

can draw

and the molecular structure is just the

arrangement of atoms

in a molecule or ion and the question is

sometimes you can draw several so how do

you know which structure is the best

for example if we look at carbon dioxide

here we can see that we can draw three

pretty much equivalent structures of

carbon dioxide

if we just sat down and started drawing

a lewis structure we could easily get this

this

structure with a carbon in the center

and two oxygens double bonded on either side

side

we could also draw this structure same

number of electrons with one carbon in

the center

one of the oxygens triple bonded and one

of them single bonded

or we could perhaps start with putting

an oxygen in the center and then double

bond a carbon and double bond the oxygen

to the other oxygen

um this is this perhaps is somewhat less

equivalent this one with the oxygens

double bonded to each other but these

two structures how do we know if we

should have a triple bond or two double bonds

bonds

and one of the ways that we can figure

that out is through formal charges

so we'll go ahead and calculate all the

formal charges on the oxygens and the carbons

carbons

just so we can kind of get a sense of of

what the formal charges might look like

for these various atoms

we've actually already calculated the

formal charge of an oxygen with two lone

pairs and two bonds

we did that above for the formaldehyde

so we know that for an oxygen with two

lone pairs and two bonds the math

looks like this so i'm going to go ahead

and just use this

to to determine that the formal charge

on the oxygen in the exact same situation

situation

with two lone pairs and two double bonds

is also zero

same thing here also zero and this oxygen

oxygen

same thing two lone pairs two double

bonds is going to be zero

we're going to need to calculate the

formal charge for the oxygen

that does not that any oxygen that's not

making two bonds with two lone pairs

so we'll go ahead and do that this

oxygen the

formal charge will be given by that equation

equation

number of valence electrons which is six

minus the number of electrons in lone

pairs which is two

minus half the number of electrons

involved in bonds here we have three

bonds which is six electrons

uh this simplifies down to six minus two

minus three which is equal to

positive one so this oxygen has a formal

charge of positive one

we also need to calculate the formal

charge of this oxygen

that is uh only making one bond and has

three lone pairs so we'll go ahead and

do that

um that oxygen has the same starting

valence electrons six um here it's got

actually six

lone pairs or six electrons in the lone

pairs and then we'll subtract

half the number of electrons in bonds

which is uh

two so we have six minus six minus one

um and this is going to be

all right and then this middle oxygen

here uh we need to calculate

its formal charge as well so here we'll

go ahead and start off with our valence

electrons which is six

subtract any lone pairs which is zero

and subtract half the number

of the electrons in the bonds here we

have four bonds so eight electrons

so six minus zero minus four is equal to

and then we will move on to the carbons

um so carbon making four bonds

we'll just go ahead and calculate it um

for practice here uh so carbon with four bonds

bonds

uh we start out with valence electrons

there are four valence electrons in a carbon

carbon

and then we will subtract lone pairs and

then the

half the number of electrons in the

bonds and we'll figure out that this

is zero uh here we see this carbon

although the arrangement of the bonds is

different with one single bond and two

or in and one triple bond this carbon is

still actually making

four bonds and has no lone pairs so it's

going to have the same formal charge as

the other carbon with four bonds

and then lastly we'll calculate this the

formal charge on this carbon

with two bonds and two lone pairs so

again we will start with the number of

valence electrons subtract the number of

electrons and lone pairs which here

is four and then we will subtract half

the number of electrons involved in bonds

bonds

which is also four so this is going to

simplify to four minus four

minus two which is equal to negative two

so this carbon has a charge of negative two

two

we will always take a moment to uh check

and make sure

that all of our formal charges sum to

zero uh since we have a molecule whose

net charge is zero

in this structure easily we see zero

plus zero plus zero

is zero um here we see plus one

plus zero and then oh i forgot to write

uh this formal charge from down here is

negative one so plus one

minus one the formal charge the sum of

the formal charges here is zero

and then our last structure is zero plus

two minus two

the net uh the sum of all the formal

charges is zero so our formal charges

are correct

and now we can start to kind of

understand why we might care about

formal charges

essentially we can use a set of rules to understand

understand

which structure is best based on the

formal charges

we can already make a guess that this

structure over here

is probably not good because of the

oxygen in the center when we're drawing

lewis structures we tend

to start by putting the less

electronegative atom in the center

and here we see that we didn't do that

and we wind up with kind of crazy formal

charges a plus two and a minus two

that's quite large

um and this is not going to be a good

structure so it turns out that the

best structure is going to be one where

if we can draw

a structure where all of the formal

charges are equal to zero

um then that's going to be preferable

those electrons are going to be in their

most stable configuration

uh when we can calculate the formal

charge to be zero for all of the atoms

here we see right away we have one

structure where

all three formal charges are zero so

this is going to be the best structure

um that's compared to this structure

which uh there's a plus one and a minus

one so that's not going to be as stable

there are other rules that we need to

keep in mind if you can't draw all of

the formal charges to be zero

if you can't draw a structure where all

the formal charges are zero then there

are more

guidelines that we need to take into

account and when we're going to think

about those i'm going to

look at another example molecule which

is thiocyanate

i've already drawn out the three

configurations um three molecular

structures that you can

draw for thiocyanate which is an anion

and then let's go ahead and really

quickly assign all the formal charges to

these three structures so that we can

start to think about

um which of these structures is best

all right so we will start out by calculating

calculating

the uh formal charges for nitrogens with

two lone pairs and two bonds

if we look on the periodic table we'll

see that nitrogen has five valence electrons

electrons

then we'll subtract the number of

electrons in lone pairs

so here that's four and then we will

subtract the number

half the number of electrons in the

bonds which here is four

uh here so this winds up with five minus

four minus two

and we're going to uh wind up with a net

charge of negative one here

on this nitrogen um so this uh

two bonds two lone pairs is identical to

this nitrogen over here so i will go

ahead and just assign that a negative

one as well

we need to calculate the formal charge

on this nitrogen

which has four bonds and no lone pairs

so this

nitrogen has five valence electrons from

the periodic table

it has no lone pairs and then it has

half of the number of electrons in the

bonds which is eight and we're going to find

find

that this is plus one so this nitrogen

has a formal charge of plus one

all right we'll move on here we already

know that a carbon making four bonds has

a formal charge of zero

so we'll go ahead and call that zero

we'll move over here to calculate

the formal charge of a carbon that has

two bonds and two lone pairs

so we start out with the number of

valence electrons from the periodic

table which for carbon is four

and then we will subtract the number of

lone pairs

that this electron or that this atom has

which here is four

and then also subtract half the number

of bonding electrons which

is four and we're going to find that

this carbon has a

uh formal charge of negative two

all right so then we need to calculate

the formal charge of a sulfur with

sulfur lives right below oxygen on the

periodic table so we'll start out with

six valence electrons

and then we will subtract the number of

electrons and lone pairs which is four

and then subtract half the number of

electrons in the

bonds which is four and we will find out

that this has a formal charge of zero

this is the same as this configuration

of sulfur

also has two bonds and two lone pairs so

this is also zero

and then lastly we need to calculate um

the formal charge of a sulfur with four bonds

bonds

so sulfur has six valence electrons

we'll subtract

no electrons and lone pairs and then

half the electrons

in bonds and we'll find that this has a

formal charge of

plus two and then i forgot to copy over

here a carbon with two lone pairs and

two bonds is

has a formal charge of negative two all

right so now that we've got all of our

formal charges

assigned we can go ahead and think about

how we can use these rules to understand

which of these structures might be the best

best

so a molecular structure with

all zero formal charges is the most

stable but we don't have any of those

options so we have to move on to our next

feature which is if we can't draw a

structure with entirely zero formal charges

charges

then the one with the smallest formal

charges is going to be preferable

so we can take a look and we see that

this structure is probably in the lead

here because

it has a negative one um and then zero

and zero

these other two structures both have a

higher formal charge

we also should check to make sure that

we don't have any positive or negative

formal charges all packed in at one end

or another

here we've got a one and a zero sorry

negative 1 and next to a 0 it's not next

to another negative

so that is a good sign we actually don't have

have

any uh structures here with any

positive formal charges next to other

positives or negative

other next to other negatives but

it's still a good thing to check and

then the last thing that we should check

is that if we have to choose among

several structures

one of the things we can look for is

that any negative charges any negative

formal charges should

be assigned to more electronegative

atoms here

nitrogen is the most electronegative

atom in this structure uh carbon is less

electronegative and sulfur is also

less electronegative and so the negative one

one

being assigned to the nitrogen is uh

more stable than if we have say a

positive charge assigned to that nitrogen

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YouTube Transcript:
Chapter 4.5a Formal Charges