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Chapter 3.4 Electron Configurations

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the topic of this video is electronic

structure of atoms

otherwise known as electron

configurations the learning objectives

are also on the screen so go ahead and

pause the video now to write those down

in your notes

um in a previous video we discussed

quantum numbers

and specifically the principal quantum number

number

uh referred to the size and relative

energy of the orbital of the atomic

orbital that we were looking at and

really the principal quantum number defined

defined

um what other orbitals were

available you know with principal

quantum number one there's only

an s-type orbital available but as soon

as you go up to the principal quantum number

number

two now we could think about p orbitals

and so on and so forth

um when we when we want to

describe the electronic structure of an

atom using something like an electron

configuration which we'll

dive into in this video we need to

really think about

relative energy levels okay and so this is

heavily tied into um principal quantum

numbers because the higher the

principal quantum number the higher the

relative energy of the orbital

if we arrange s type p-type

d-type and f-type atomic orbitals

um as a function of

energy we see that we'll start with at

the lowest

this is the lowest energy down here and

this is the

highest energy up here if we were to

fill electrons

into this into these orbitals we would

start from the bottom

up we would start from a low energy

point and work our way

up to a high energy point now

um if we were to put one single electron

down here

in the 1s that is the most favorable

place for it to go

if you were to add another electron we

would have to also add it into the

1s because they're still um

room for one more electron down here

there's room for one more electron because

because

the arrow up represents spin positive

one half

so that means we can also fit a second

one in there with

spin of negative one-half you cannot fit

another electron in here because

there are only two types of spin and we

can't have

um according to the poly exclusion

principle um

electrons with the same four quantum

numbers so these two electrons

don't have the same four quantum numbers

but they're close they're just different

they just differ in their spin there's

no more room for

other electrons in this orbital but you

can see that this is a really nice view

because now we know

okay if we want to add even more

electrons we would jump up to the 2s

and then after that we would jump over

to the 2p before

filling the 3s so

the observation of starting to fill um

atomic orbitals

from the lowest energy uh orbitals first

is called the off-bob principle

the off ball principle states is such it

can be stated such that

each added electron

occupies the subshell

of lowest energy

available subject

to the limitations

imposed by the allowed

quantum numbers according

to the poly exclusion principle

so all this is saying is that you fill

from the lowest energy on up that's it

this extra language about the paul

exclusion principle is what i already

outlined here

you can't fit another even though the 1s

remains the lowest energy atomic orbital

in this figure

we would violate the poly exclusion

principle if we tried to fit a third

electron into that

1s atomic orbital we can't because to do

so means we'd have to

have another electron that has the exact

same four quantum numbers as either

the spin up electron with the up arrow

or the spin down arrow

a spin down electron represented by the down

down

single barbed arrow another really nice

way to lay out

the electron configurations of

different atoms is using a periodic table

table

and so here this is a you hopefully

recognize the overall shape and

structure of this

table but notice that the typical stuff

that you might see in a periodic table are

are

removed and instead this is structured

in such a way that that we

are filling um starting from the 1s

first and then helium is now considered

a 1s 2 electro electron configuration

lithium would be now a 2s1

beryllium would be a 2s2 so the periodic table

table

is really just a road map for electron configurations

configurations

uh that's all it really is so let's

let's pick

um an element that i worked with quite a

bit as an undergraduate

researcher boron and let's write out

an actual electron configuration of

boron and we can use this periodic table

to to help us do so

so when you write out an electron

configuration there's a few different

things that we need

um we need to start for the complete

electron configuration

we need to start always with with a 1s

so for boron we're already past the 1s

orbitals so we know that those must be

filled already

so the 1s orbitals must be 1s2 because

that's the maximum amount of

electrons that oh an s orbital can can accommodate

accommodate

not only have we completely filled the

1s here to get to boron but we've also

completely filled the 2s

so right next to that 1s2 we're going to

write 2s2

so now we're saying that the first and second

second

shells are or orbitals atomic orbitals

are completely filled with 2 electrons

each when we get to p orbitals now there

are actually since there are three p

orbitals each of which can accommodate

two electrons there are six total p

electrons for boron since it's the very first

atom here in this in the p orbital series

series

so this right here is indeed the

electron configuration for boron

now you can also depict this using

a what's called an orbital diagram and

so an orbital diagram just sort of uses boxes

boxes

so this box will represent 1s the

next box will represent 2s and then we

can even have

a box here for the 2p

where we have three p orbitals all of

which are at the same energy

so if we were to draw out the orbital

diagram for boron again so this is

used for um for boron we would put

two electrons

into the 1s two electrons

into the 2s and only one electron

into the 2p so this would look like

this we have two electrons in the 1s

two electrons in the 2s and only one electron

electron

in the 2p let's go ahead and do one more

example here

let's look at the electron configuration

of sodium

so sodium is right over here

circling it in red so now we know let me

go ahead and erase

what i have written before here so we're

looking at sodium

i'll put an x in sodium so we know that

our 1s

is completely filled our 2s is

completely filled our 2p

is completely filled before we get to

sodium so

here i'm going to write out uh let me

change back to white

1s2 is filled 2 s2

is filled 2p6 is filled and now we're at

3s and there's only one electron at

sodium's position 1s in the

one electron in the 3s orbital we could

draw out the entire orbital diagram here

so this is the

1s 2s

this is going to be our two p orbitals

and then also we have our 3s so this is

1 2 3 4 5

6 7 8 9 10

11 electrons total now you can see that

this would get tedious we're only

up to sodium and which is not

one of the heavier elements

you can imagine that if you had to do

this for example

xenon or even

barium you you would the electron configuration

configuration

is going to be really really tedious

it's going to be redundant with many other

other

atoms and it takes a while to write it

all out so

thankfully there are there is a

designation for how to handle this let's

take a look at sodium

so for for sodium in particular i'm

going to

rewrite the electron configuration over here

here

1s2 2s2

2p6 and 3s1 there's a

there's a break here where the principal

quantum number jumps from

one to two and two to three but the

outermost one the one that has

the um the largest size or the highest

relative energy

is the highest uh principal quantum

number value of three

all of the electrons before that break

that principal quantum number three

outermost electrons have a special

name they are referred to as valence

electrons okay

so what we have what we can do is we can

use this shorthand notation

to sort of get we don't have to write

out all the core electrons

every single time what we can do is we

can actually take a look at and see what

element um has the same

electron configuration as these core

electrons so which element has the same

configuration as one 1s2

2s2 2p6 that is the element

i'm going to erase these again the

element there

is going to be so we have 1s

2 is gone 2 s2 is good 2p6

is neon so what we can do

is we can actually say okay forget

writing out all those core electrons

we can just summarize all of those using

a bracket

neon close bracket

notation followed by the valence

electrons 3s1

so this is now the abbreviated

on your screen now i have the electron

configuration table

where the valence electrons are included

with each

element and what makes the periodic

table of elements periodic

and what makes it a very useful

predictor of chemical properties is the

fact that

it is arranged in such a way that

that each element or all the elements

are grouped according to

similar valence electron finally we're

going to end on electron configurations

of ions so if you are asked to

write the electron configuration of ions you

you

will presumably be given the ion so

that's over here

the ions are are here that we'll be um

practicing with

and the first step is to write out the

electron configuration of the atom

itself of not the ion but just the atom

okay so this

is um electron

config of

atom i have chosen to use the

abbreviated notation here because i

don't feel like writing out all of the

core electrons

the next question is to ask yourself to

form the ion in question

were electrons added or removed so if we

have cations which are the first three

samples here that means that the ion is

inherently positively charged in order

to become inherently positively charged

it must have lost a negative charge

and so these all three cases the cations

to get to the cation from the neutral

atom we must

remove electrons how many electrons do

we remove well that depends on the

magnitude of the charge sodium

only has a one plus charge so it lost

one electron so

the ion

electron config

would just be removing one electron and

there's only one electron to remove from

the valence shell the 3s1

so literally the only thing that you

would need to write

for the abbreviated electron

configuration of the sodium ion

is bracket neon the next ions

are both um iron ions iron two plus iron

three plus the rule for transition

metals is that you

remove electrons from s um orbitals

first so iron two plus we need to remove

two electrons and since it's a

transition metal we are going to remove

them from the 4s2 first

so the electron configuration for iron

two for the iron two plus

ion is going to be argon

core electrons 3d6 as our valence electrons

electrons

now the reason that i wanted to look at

iron three plus is because now we can

essentially say well

similar to iron two plus we're going to

remove the four s

two electrons first but we also need to

remove one more because this is a three plus

plus

cation so we actually have the argon core

core

three d for iron 3 plus

now in the case of chloride which is an anion

anion

in order to get a neutro from a neutral

atom to

a negatively charged ion we need to add

one or more electrons

this only has a single negative charge

so we needed to add

one electron to this if we're going to

add the electron we need to add it to

the the subshell that can accommodate it the

the

3s subshell is already occupied completely

completely

so but the 3p can occupy one more

or it can accommodate one more electron so

for chloride we would write the neon core

core

and you can write three s two

three p six but now you notice now we are

are

we have a complete um shell here

uh where the the three p sub shell is

completely filled

so you could write it this way um uh

neon core 3s2 3p6

alternatively that is also the

valence of argon

and so we could also write this as just

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Chapter 3.4 Electron Configurations