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Chapter 4.4a Lewis Symbols and Structures

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welcome to the first video for chapter 4

section 4 lewis symbols and structures

in this video we'll be working on

writing and drawing lewis symbols and structures

structures

the learning objectives are given here

let's begin

by uh defining what a lewis symbol is

is a depiction of the valence electron

configuration of atoms and monatomic ions

ions

what that means is it's a symbol that

shows what element you're dealing with

and how many electrons are in its

valence shell and that looks like what

we've got here in this table

we have the elemental symbol so for

sodium that's n a

surrounded by dots that represent the

valence electrons the dots tend to show up

in singles up until we have to pair them

they show up in sort of four

pairs four regions around the uh around

the symbol

and there can be up to eight in general around

around

the elemental symbol we can use them

not only to just show uh elements or ions

ions

but we can show the formation of ions

with these um because we can show the electrons

electrons

that are changing in their configuration

around the elementary ion

so for example we can start off with

sodium which has one electron

and if we're forming the ion it'll lose

that electron and become sodium plus

which is different than sodium because

it no longer has this electron

plus we will generate one electron on

its own

we can also use the same notation to

show the formation of

anions so for example sulfur

which starts out as we have it in the

table it starts out with uh

six valence electrons so two pairs and

then two

electrons hanging out over on the sides

and if we're going to form an

ion with sulfur that's going to form the sulfide

sulfide

anion which has a two minus charge um

because that's going to be most

similar to the noble gas configuration

so we'll add two electrons here

and then we can see that sulfur will have

have

eight electrons so let me just switch

colors really quick and i'll just show

you the

two new electrons

are going to come in here and it's gonna

we need to show it with the two minus charge

so that's how we can use this notation

to show the formation of ions we can also

also

use this same notation to show electron

transfer during the formation of ionic compounds

compounds

so for example if we're going to form

sodium chloride we're going to start off

with sodium

and we're going to add some chlorine and

here we'll just show one chlorine atom

even though we know that

chlorine is generally

a diatomic and we as we form sodium

chloride we know that

the sodium element is going to become sodium

sodium

uh plus the ion so what happens is that

this electron

needs to get transferred over to this

chlorine in some way

and we can show this just by using our

lewis notation so we show that

sodium has become an ion and then we'll

show the formation of the chloride ion

as well

um and we'll just go ahead and put this

in brackets just to denote

that it's different and here we can see

that we've got the

sodium chloride compound formed by the

attraction between the sodium cation and

the chloride anion

so we'll move on to thinking about lewis structures

structures

instead of lewis symbols and lewis

structures are quite similar to lewis

symbols they shared a lot of the same notation

notation

but what's different is that instead of

looking at individual atoms or elements

now we're going to be looking at

molecules and polyatomic ions that are

sharing electrons and lewis structures

essentially are drawings that describe

that bonding in molecules and polyatomic

ions and allow us to kind of

look at what electrons are being shared

so i'm going to show you what this might

look like

with two chlorine atoms that are going

to form

a a a single covalent bond

um when they form their uh their diatomic

diatomic

molecule so i'm drawing the two chlorine

atoms in different colors just so we can

keep the electrons straight

and what this looks like when they are

have formed that bond and are in their

diatomic molecule is that there's just

two electrons that are going to be

shared between

these two chlorine atoms so we can

see that these two electrons here are

being shared uh each of these electrons

or each of these chlorine atoms now has

eight electrons around it which mimics

the noble gas configuration they have a

full p

subshell and um so this is the this is

the sort of more stable

version of chlorine this is why chlorine

exists as a diatomic molecule

um we will often represent this instead of

two dots that are just kind of

sandwiched between these two atoms

just for clarification we'll often

represent this structure with a line

between those two atoms and we'll still

draw all those valence electrons

as dots around them but we will use this line

line

to represent the single bond between

these two atoms

um the single bond of course is the the

shared single pair

we're going to talk about how to draw

lewis structures in the next video

but uh for now there's a few rules that

we should

learn and the first important rule is the

the

octet rule and the octet rule

is the tendency of main group atoms to

form enough bonds to obtain

eight valence electrons and so of course

eight is where the octet

comes from and this is again similar

because that's uh

they tend to do this because that is a

full p subshell uh so that mimics the

noble gas configuration

uh i've bolded tendency here because

it's not a hard and fast

rule although we do call it the octet

rule it is hard and fast for some

elements but we are going to talk about

exceptions to the octet rule

a little bit later down one thing that's

pretty cool about the octet rule is that

the number of electrons that's needed to

reach an octet

is predictive of how many bonds an atom

can form

so what that means is if you're looking

at an element in the periodic table so

for example carbon

you'll note that carbon has four valence electrons

electrons

so in order to form an octet this this

atom is missing

four electrons so uh in order to get to

eight electrons it would need four more

and so we can think that carbon might

tend to form four bonds

and in fact that is the case this is

especially this is especially true for

the top row

uh p elements so carbon nitrogen

oxygen and fluorine this is really predictive

predictive

carbon tends to make four bonds nitrogen

tends to make three bonds

oxygen tends to make two bonds and fluorine

fluorine

tends to make one bond so and again

that's just because

of the electrons that are needed to

reach the octet based on the

the valence electrons that the atom has

in its neutral form

and we will get some practice drawing uh these

these

as i said in the next video but real

quick i just wanted to show you a couple

of really common compounds that

follow this pattern so um for example

for example carbon tetrachloride so a

carbon with four chlorines around it

um and we'll talk about how you

determine to form these structures

so this is a really common uh example of

a carbon

atom with four bonds around it each one

to a different atom

nitrogen a really common molecule that

we see is

ammonia which is nh3 and that's

three hydrogens around essential

nitrogen and again this nitrogen tends

to make three bonds

uh oxygen you have probably seen water

uh so water is h2o so there's two

hydrogens around the central oxygen

the oxygen tends to make two bonds and

then with fluorine

a common example of fluorine is hf

hydrofluoric acid and again this fluorine

fluorine

uh tends to make one bond um because

of uh in the in its neutral state it has seven

seven

uh valence electrons which means it has

one vacancy if it needs to make

if it needs to reach nothing again this

is predictive

this is not always the true this is not

always the case

but if you are looking at a lewis

structure this is a good guess to begin with

with

and then um this is predictive it is not

a rule you can

uh you can break this trend but it's

especially true it tends to be true for

these four

uh atoms so we've talked about single bonds

bonds

where we've got one shared pair of electrons

electrons

but what happens if a molecule needs to uh

you can't you can't uh you can't fill

the octets with just single bonds

um in other words you uh a a pair of

atoms or a molecule needs to

share more than one pair of electrons

to reach the full octet and essentially

what happens then is you have double and

triple bonds so a double bond will

represent with two lines

and that means uh two lone pairs so

essentially two pairs of electrons

and a triple bond will represent with

three lines

and that indicates three shared pairs

of electrons between the two um between

the two atoms

and so i'm just going to show you two

examples of what this looks like so formaldehyde

formaldehyde

which is c h2o and essentially the lewis structure

structure

of this molecule looks like this where

we've got a central carbon

we've got two hydrogens and then in

order to reach the octet of both the

carbon and the oxygen

we have to share two lone pairs of or

two pairs of electrons between

um between those atoms and that's a

double bond

another example of a molecule with

a triple bond is carbon monoxide which

you are probably at least familiar with

um it's co and essentially this one

has a triple bond between the c and the

o and then both

atoms have a lone pair on the outside so

again this is three pairs of electrons

being shared between these two

atoms in order to make uh to fulfill the

octet structure

um i will just point out here that

you'll note that this

where carbon is making three bonds and

so is oxygen breaks this trend of

the prediction that we just talked about

above again that's a prediction not a rule

rule

and that's okay um if you have to do

that that's okay

all right so the last thing we're going

to talk about in this video is

exceptions to the octet rule

so sometimes you can't fulfill an octet

for one reason or another

and there are a few cases when one

that's going to be especially common

and one is if you have an odd number of

valence electrons so nitrogen for

example has

five valence electrons uh and so

it's a problem if you have an odd number

of valence electrons you can't fulfill

an octet because you have a lone

electron hanging out

so a common molecule that has this is no

and we will talk about the structure of

no in the next video

but when you have this uh basically the

structure looks like this where you have

a double bond

oxygen has a full octet nitrogen doesn't

it has a lone

electron hanging out and we call these

free radicals when there's an

unpaired electron there's also electron

deficient molecules

which is a case when the central atom

has fewer electrons than needed for the

noble gas configuration or the full

octet and this is often in group

uh 2 and 13 so that's usually beryllium

and boron are the are the the

general players here so a molecule that

we see this with is bf3

and it actually has boron with three

fluorines around it

um and the boron actually only has six

electrons the fluorines will have the

full octet

i'm sorry this is a little messy but the

boron actually only has six electrons it

does not have a full octet

um this means it's quite reactive but

this is in fact the structure of

uh the boron or of bf3 that we'll find experimentally

experimentally

and then the last exception to the octet

rule that we'll see is hypervalent

molecules and that is when you have a

central atom with

more electrons than needed for the noble

gas configuration

that is to say it has more than eight

electrons it breaks the octet rule in

the other direction

this you cannot do until the third row

um only things in the third row and

below can exceed the octet

the top row carbon nitrogen oxygen

fluorine cannot exceed the octet

and we'll talk a little bit about that

but it has to do with the d orbitals

a common example of this is pcl5

and that's going to be a phosphorus with

five chlorine atoms around it

and the chlorine atoms all have their

valence their uh their octet

and the phosphorus has

five chlorines attached to it um and it actually

actually

has therefore ten electrons around it

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YouTube Transcript:Chapter 4.4a Lewis Symbols and Structures

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YouTube Transcript:
Chapter 4.4a Lewis Symbols and Structures