Orbital Diagrams and Electron Configuration: Aufbau, Pauli, and Hund's Rules Explained

Introduction to Orbital Diagrams and Electron Filling Rules

This video provides a comprehensive guide to constructing orbital diagrams and understanding the fundamental principles that govern electron placement in atoms. The instructor uses a clear, step-by-step approach to demonstrate how electrons fill energy levels and sublevels.

Key Concepts Covered

What Are Orbitals?

Energy Levels: Represented by the principal quantum number (n = 1, 2, 3...)
Sublevels: s, p, d, and f sublevels within each energy level
Shapes: Each sublevel has a distinct shape (s is spherical, p is dumbbell-shaped)
Orientation in Space: p orbitals are oriented along x, y, and z axes; d and f orbitals have more complex orientations
Electron Spin: Each orbital can hold up to two electrons with opposite spins

The Orbital Diagram Structure

1s: One orbital (s sublevel)
2s and 2p: Two sublevels, one s orbital + three p orbitals
3s, 3p, and 3d: Three sublevels, one s, three p, and five d orbitals
4s, 4p, 4d, and 4f: Four sublevels (f sublevel not shown in diagram)

The Three Core Rules for Filling Orbitals

1. Aufbau Principle

"Electrons occupy the lowest energy orbital available first."

Think of it like filling a building from the ground floor up
Electrons seek the lowest energy state possible
Fill from bottom to top on the orbital diagram

2. Pauli Exclusion Principle

"Each orbital can hold a maximum of two electrons, and they must have opposite spins."

Represent electrons as arrows: ↑ for spin up, ↓ for spin down
Two electrons in the same orbital must be paired with opposite spins
No orbital can contain more than two electrons

3. Hund's Rule

"When filling orbitals of equal energy (degenerate orbitals), electrons occupy each orbital singly before pairing begins."

The "Roommate" Analogy: Everyone gets their own room first before sharing
Spin Consistency: All unpaired electrons in degenerate orbitals must have the same spin
Application: Important for p, d, and f sublevels where multiple orbitals share the same energy

Step-by-Step Example: Bromine (35 Electrons)

Electron Count: 35 (atomic number 35)

| Step | Orbital | Electrons Added | Running Total | Notes | |------|---------|-----------------|---------------|-------| | 1 | 1s | 2 (↑↓) | 2 | Lowest energy level | | 2 | 2s | 2 (↑↓) | 4 | Next lowest s orbital | | 3 | 2p | 6 (↑↓ ↑↓ ↑↓) | 10 | Hund's rule: 3 singly, then pair | | 4 | 3s | 2 (↑↓) | 12 | s orbital fills completely | | 5 | 3p | 6 (↑↓ ↑↓ ↑↓) | 18 | Follow Hund's rule again | | 6 | 4s | 2 (↑↓) | 20 | Note: 4s fills before 3d | | 7 | 3d | 10 (5 ↑↓ pairs) | 30 | Five d orbitals, each with 2 electrons | | 8 | 4p | 5 (↑ ↑ ↑ ↓ ↓) | 35 | Hund's rule: 3 singly, then 2 paired |

Visual Representation of Bromine's Electron Configuration

1s: ↑↓
2s: ↑↓    2p: ↑↓  ↑↓  ↑↓
3s: ↑↓    3p: ↑↓  ↑↓  ↑↓    3d: ↑↓  ↑↓  ↑↓  ↑↓  ↑↓
4s: ↑↓    4p: ↑  ↑  ↑  ↓  ↓

Important Observations

Energy Level Ordering

4s fills before 3d: The 4s orbital has lower energy than 3d
Energy increases: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s...

Maximum Electron Capacity by Sublevel

s sublevel: 1 orbital × 2 electrons = 2 electrons maximum
p sublevel: 3 orbitals × 2 electrons = 6 electrons maximum
d sublevel: 5 orbitals × 2 electrons = 10 electrons maximum
f sublevel: 7 orbitals × 2 electrons = 14 electrons maximum

Practical Applications

Understanding orbital diagrams is essential for predicting chemical bonding, reactivity, and magnetic properties of elements
This knowledge forms the foundation for studying the Understanding the Classification of Elements and Periodic Properties in Chemistry and electron configurations
The rules apply universally to all atoms, regardless of the number of electrons

Common Mistakes to Avoid

❌ Filling 3d before 4s (remember: 4s is lower energy)
❌ Pairing electrons in degenerate orbitals before filling singly (violates Hund's rule)
❌ Giving two electrons in the same orbital the same spin (violates Pauli exclusion principle)
❌ Placing electrons in higher energy levels when lower ones have available space

Before constructing orbital diagrams, it's helpful to have a solid foundation in Understanding Atomic Structure: From Atoms to Subatomic Particles. Additionally, for an overview of concepts like sublevels and electron capacity, refer to the Comprehensive Overview of Periodic Table and Key Concepts in Chemistry. To reinforce your understanding of the building blocks of matter, you can also explore Understanding Atoms: Structure, Particles, and Elements. Furthermore, this material builds directly on concepts introduced in Understanding Atomic Structure: Protons, Electrons, and Electron Configuration.

hello everyone and welcome back my name is mr covald and in this video we're going to go over orbital diagrams and

the principles and rules on how to fill them up so in a previous video we talked about

orbitals and how there's different levels of uh energies for orbitals so the

the few things that we need to know about orbitals is that uh the they are described by their energy level

their energy sub level uh they are described by their shape uh and they are described by

their um orientation in space as well as the number of electrons in there or the spin of the electrons which is

describing more of the electron than the orbital itself but these are things that we need to keep in mind

um so in this video what i'm going to show you how to do is how to fill in an orbital diagram uh using the rules so

i'll give you one example so let's get into this so uh here is an orbital diagram

and so what you'll notice is uh we have this uh broken up into different levels and

so the first number that you see here is the energy level um of the of the orbital

and then the letters here we talked about last last time in a video

that the letters here represent the shapes of the orbitals and you can see that

they are broken up into levels so we have levels and sub levels so we have level number one

level number one only has one sub level and that sub level is the s sub level and each line here represents an orbital

in that sub level okay so you can see here in this uh

energy level here number one that is only one sublevel

and so it's the s sub level in energy level number two notice we have two

sub levels here so we have the s sub level and we have the p sub level and you'll notice that we have three

lines here those lines represent the three p orbitals that are at that sub level and

so we talked about in a different video about the orientation of the orbitals so

again they're oriented again along the different axes the x-axis y and z-axis and so this would be maybe this is the

px orbital this is the py orbital and maybe this is the p

z orbital so we can label those here if you'd like so that would maybe be x this is the y

and that's the z so three different orbitals p orbitals at the p sublevel at

energy level number two and so remember the energy levels are those n numbers n equals one n equals two and so on so

here we have n equals one here's we have the energy level number two with two

sublevels and then we have energy level number three

and this is where the d sub levels come in this is where your d orbitals are and you'll notice that there are

three again three sub levels notice the energy level number three and

we have three sub levels we have the s the p and the d and again notice here we have

five lines that's the five d orbitals and the different orientations that they have

in space and then of course we have four the energy level number four we have

four s p

d and then there's an f sub level that i didn't i didn't draw in

this um in this uh diagram because uh i just want to uh go over just for the purposes

of this video um i cut it off there and then you go to 5s 5p and 6.

and so this is your uh your orbital diagram

and again like i said i cut it off short but this it goes it goes further the other thing i want to point out is again

what we i said before is that the energy of the sub levels increases

going up so the higher the number the higher the energy and so remember that the electrons

uh closer to the nucleus if the nucleus is down here if i was to draw out a nucleus this is the nucleus of the atom

so that nucleus the electrons that are closer closest to the nucleus are going

to have the lowest energy as you get farther up in energy you're going to get farther

away from the nucleus so the farther away you are from the nucleus the more energy so here we have energy increasing

going up okay so

here's our orbital diagram so how do the electrons go into the orbital diagram so what rules do we

follow and so over here i have a few rules i have two principles in a rule that we're going to follow and so let's

go over that right now so first is the aufbau principle the

outbound principles basically says an electron occupies the lowest energy orbital that it can first

so it's going to reach the lowest energy so it wants to go as low as possible if

there's space so basically electrons are going to fill from the

bottom up so think of this as a building and uh the elevator is broken there's no elevator and you have to go up the

stairs right so are you going to go all the way up 7 6 10 flights of stairs no so you're going to want to be in the

lowest floor of your building so think of the energy levels as floors

of your building maybe that analogy works for you so

so enter electrons don't want to go all the way up they want to be lowest to the ground they want to be down here

if there's room so if all of the space down here is filled then they have no choice but to go to the next floor and

if that's filled they have no choice but to go to the next floor but they want to have the lowest

energy possible so when you put when you start putting electrons in your atom into the different energy levels you're

going to start low and work your way up okay so that's the alpha principle

the second principle is what is called the poly exclusion principle and in a nutshell that basically says that each

orbital can only hold a maximum of two electrons and

they must have different spins so we talked about how electron is either spin up or to spin

down so those two electrons must be obsolete spin and they

in the same orbital so if you have two electrons in the same orbital they have to have opposite spins and you can't fit

any more electrons into that orbital and so that's only two

electrons in an orbital and then third we have is hung's rule and so hon's rule basically says is that

if you have orbitals that are of the same energy level like for example here we have

three p orbitals and these are the three p orbitals so three p orbitals they all have the same energy

these three p orbitals have the same energy you could see the d orbitals here the d

orbitals have the same energy so all five of these d orbitals have the same energy so when

you have orbitals that are of the same energy then

the electrons are going to occupy the orbitals one at a time before you pair them up so you don't pair up electrons

into an orbital until you've actually filled one each first so think about having your own room

everyone wants to have their own room first you don't really want to share room with someone maybe you do but for

this analogy you don't really want to share room you have your own private room

so people are going to fill in and take the single rooms first before you start bunk uh bunking up with

someone in your same room or having a roommate right so um you want you prefer to have a single

room until we uh don't have enough uh we have too many people and so we need to bunk people up into rooms so that's

analogy to help you remember hund's rules so hun's rule says you got energy you got orbitals of the same

energy you put electrons singly in each one one at a time here and then if you have other electrons

that need to go into an orbital then you start pairing them up but again remember that

they have to be opposite spin the other thing you want to be aware of with hunt's rule is that all

have the same spin so when you put them into the orbitals here you have they all have

to have the same spin okay so those are the two principles and the rule for orbital diagrams and how do we

fill those in so you want to think of an orbital diagram of basically an address of where the electrons are

where are they placed in in the orbital diagram right so where we

place those electrons we'll uh we'll place them in the orbital diagram according to these rules all right so

let's go over an example and this might be very uh unclear but going over an example

will be uh helpful in making this clear so let's go over um what will be in a good example let's

go uh through boron i'm sorry not bromine bromine is number 35 on the periodic table

so that means that the bromine atom has 35 electrons because

we're assuming a neutral atom so the uh bromine has atomic number 35

it therefore it has 35 protons and we're going to assume that it has enough electrons to balance out the charge so

35 electrons where do those electrons go where are they housed in the electron

this is the structure this is the electron structure of the atom so

when you have electrons in your atom the electrons are going into orbitals remember orbitals are those spaces where

you can have a probability of finding the electrons somewhere within that space and depending on what orbital

you're in you have a different space or different probability space that you are occupying

so how do we fill these electrons in we have 35 electrons so again this is bromine

right so let's let's see how we fill these in so again we're going to i'll we're going to follow the outbound

principle that means that electrons want to be in the lowest energy possible lowest energy orbitals first

and then they uh and then they're going to work our way up so when we put our electrons in the elect the first

electrons are going to go in a low energy until those are occupied and then when this is occupied we put

them in the next lowest energy and then the next low synergy and then next so these are building up you can see where

the energy is increasing up so this is lowest energy this is the next lowest this is next and then foul here

and then this one and then this one is next and then this one over here is next and then followed by that one and so on

so we're following the order of the energy level so let's start with the first electrons so the first

electrons are going to be put into here when we put them into the orbital diagram we

draw arrows to represent spin so the first arrows spin up

and the second arrow spin down so those are two electrons one has a spin up one has spin down notice

that they have opposite spins that is following the poly exclusion principle so the poly exclusion principle says

that we can only have a maximum of two electrons and those electrons have to be opposite spins so here's our one orbital

in the 1s sublevel so we have an s orbital here so we're putting those two electrons

there the next one is another s orbital and now it's at the higher energy level

which is energy level number two that's going to have two electrons so now we're going to put a spin up

electron there and a spin down electron there so here now we have two electrons there so we've used up four of

electrons but we got 35 total we got 31 more electrons to put in next we're going to put electrons in

this one this is the next energy level the next lowest energy uh sublevel right we have a sublevel s

and we have a sublevel p so now this is the next one so we have three p orbitals p x p y p z that are

the same energy so we need to follow hund's rule here so owen's rule says if there are more than one orbital with the

same energy we need to put the electrons in each one singly first before we pair them up and they have to have

the same spin so we are going to put three electrons in

here so we first do this we have one electron spin up we next fill this one up

spin up and then this one spin up and we keep going so

this is how many electrons is seven electrons again we have 35 to fill in so we got to keep going so now the next

electrons are going to be paired up so remember we fill them in one at a time singly and then we pair them up if we

have more but we have more so then this one's going to be spin down this one's going to be spin down and that one's

going to be spin down so again following ons rule we filled one in one at a time same spin and then

we paired them up with opposite spins so now we have uh what is this we have 10 electrons we

have 25 more to go so we keep filling them in so we're following the alpha principle

we're adding electrons to the lowest energy levels building ourselves up until we run out of electrons so we've

got 10 we've got 25 more to go so the next energy level is the 3s so we're gonna put two electrons here

spin up for that one spin down again we only have one orbital at the s

sublevel here that's one orbital here we have the sub level p sublevel at the third energy level and

again we have three equal energy p orbitals the p x p y p z so again we're going to follow hund's

rule and we're going to put them in one at a time same spin so that's going to be one

two three so those three electrons go there we got more electrons so we're going to

have to fill them up so we're going to pair that one up spin down pair that one up spin down

pair that one up spin down and so now we have 18 electrons used up we've got

more so we gotta keep filling up the next energy level is here

so we're gonna fill this up and so uh this is the 4s so again s

this sublevel only has one s orbital so we're going to fill that up with two electrons spin up spin down

and so now we have 20 electrons we have 15 more to go so let's keep going so now where do we go so we have the 4s

what's the next highest energy level right here so this is the next energy level so

we're going to fill this one in but notice that all five of them are the same

energy so we're gonna have to follow hund's rule here so we fill them in one at a time

so spin up that's one spin up two spin up three

spin up four spin up five and now we're gonna pair them up spin

down spin down spin down

spin down and spin down

so that's 10 electrons in there so we had 20 now we put 10 as 30 we have five more

electrons because bromine is number 35

number 35 of the periodic table so we have 30 now we need to fill in five more electrons what's the next one

it's right here so that's the next energy level so those electrons have to go here

again hund's rule we're going to spin up spin up spin up

okay that's three electrons we need two more those two have to go here and we're

gonna pair up two of them so that's spin down and spin down and so now we have

all the electrons put into the orbital diagram for bromine and so this is the how the electrons are

structured in bromine we have electrons in the 1s the 2s the 2p orbitals

the 3s the 3p 4s

3d and 4p now there's a couple of things i want to point out here

we said that there could be two electrons in each orbital so each line remember is an orbital

and so the 1s is the lowest so each

s sublevel can hold a maximum of two electrons because there's only one orbital

in that sublevel right so 1s only has the one sublevel the 1s that 2 enter the energy level number 2

has two sublevels this sublevel only has one orbital like one line

so that only has two electrons but this sub sublevel the p sublevel here has

three p orbitals each orbital can hold a maximum of

two electrons so that means at this sublevel the p sub holds a maximum of six because there's three orbitals in in

the sub level and each orbital holds two electrons so that's a maximum of six and it's the same for all of these p sub

levels this one can hold a maximum of six this one can hold a maximum of six this one can hold a maximum of six

and then when we get to the d sub level so this 3d sub level has one two three four five orbitals five d

orbitals and so since all five d orbitals can hold a maximum of two electrons that

means the total number of electrons is ten same thing for

this one as well so this has five d orbitals in this sub level the d the four d sublevel and so

this can also hold a maximum of ten and i didn't draw the the

f sub level it's somewhere over there but if i was to we we saw before that the

number of orbitals in the f sub level is seven there's seven f orbitals each orbital holds a maximum of two

electrons so that means that an f sub level with with seven elec uh with seven orbitals can hold a maximum of fourteen

and that's going to be very important later on when we talk about the periodic table in the next in in a future video

i uh hope this was helpful uh so these are the rules this is how you use the rules to

to fill in your orbital diagram depending on how many electrons you have so different atoms have

different number of electrons and so they're going to fill in the table a little bit differently but they all

follow the same two principles and the same

rule so these are always followed i hope you enjoy this video i hope this is insightful if you liked this video

please like the button smash that like button please

hit the notification bell for uh notifica be notified of other videos i put out

uh hit uh please subscribe to my channel put a comment in the comment section let me know what you think if you have any

questions please ask thanks for joining me and have a great day

Keep this summary

Save it to LunaNotes and it becomes a real note in your library — editable, searchable, and ready to turn into flashcards or a diagram. Free to start.

Save to LunaNotes

Or summarise for another video.

This summary and transcript were automatically generated using AI with the Free YouTube Transcript Summary Tool by LunaNotes.

Related summaries

Understanding Atomic Structure: Protons, Electrons, and Electron Configuration

Explore the fundamental building blocks of atoms, including protons, electrons, and neutrons, and learn how electron configurations determine the properties of elements. This video breaks down the first ten elements of the periodic table and explains why atomic number defines an atom's identity.

Understanding the Classification of Elements and Periodic Properties in Chemistry

This video provides an in-depth exploration of the classification of elements and their periodic properties in chemistry. It covers the periodic table's structure, trends in atomic radius, ionization energy, electron affinity, and electronegativity, along with historical context and exceptions to general trends.

Comprehensive Overview of Periodic Table and Key Concepts in Chemistry

This video provides an in-depth exploration of the periodic table, focusing on the classification of elements, atomic structure, and key concepts such as ionization energy and electron affinity. It covers essential topics for students preparing for chemistry exams, including the trends in the periodic table and the significance of effective nuclear charge.

Understanding Atoms and Elements: A Comprehensive Overview of Chapter 4 in Chemistry

This lecture provides an in-depth exploration of Chapter 4, focusing on the fundamental concepts of atoms and elements in chemistry. Key topics include chemical symbols, the periodic table, atomic structure, and the significance of isotopes and electron configurations.

Oxidation Numbers and How to Find Them: Complete Guide

Learn how to find oxidation numbers with clear rules and practical examples. This step-by-step tutorial covers pure elements, ions, compounds, polyatomic ions, and electronegativity's role, helping you master oxidation states for chemistry success.