Since we’ve already covered the other parts of an atom (the nucleus) in O-Levels, I’ll focus on some new info on electrons.
Today we’ll be learning about:
- Shells & Sub-Shells
- Rules for Filling Orbitals
- Electronic Configurations
Let’s jump right in!
So far, we’ve learnt about the structure of atoms using a simplified BOHR MODEL:
- a nucleus (consisting of protons & neutrons), surrounded by various shells of electrons
- each shell is a circular orbit
- each shell can be occupied by a certain number of electrons (188.8.131.52 etc)
This model is handy when learning about ionization, bonding, valency, & other theories in high school-level chemistry, BUT it does NOT paint a complete picture of the true model of an atom.
Let’s expand our view a bit, & get ourselves acquainted with shells & orbitals.
Before anything, let’s break down Bohr’s Model:
Bohr’s Model states:
- Electrons orbit the nucleus
- The distances of the electrons from the nucleus are QUANTISED
- The further the distance from the nucleus, the HIGHER the ENERGY LEVEL of said electron
What is Quantisation?
The fact that certain values MUST be at DISCRETE values, and CANNOT exist as CONTINUOUS values.
In our everyday macroscopic (big) world, we measure things in continuous values, Things can have masses of 1g, 1.093g, 0.838g etc.
In the microscopic (extremely small) world however, things behave in a different way.
Certain quantities can only exist at certain values.
For example, electrons can only have an energy level of 1, 2, etc, but cannot have an energy level of 1.5.
This world of quantisation is known as the quantum world.
What are Shells?
Different ENERGY LEVELS of electrons inside an atom.
In Bohr’s Model, a Shell is an orbit.
For example, the 1st Shell is the nearest Orbit to the nucleus. The 2nd Shell is the 2nd Orbit nearest to the nucleus, etc.
The distance of an orbit from the nucleus corresponds to the ENERGY LEVEL of the electrons in said orbit.
But now, let’s update our knowledge.
No longer is each Shell a simple ORBIT, but a group of electrons existing in ORBITALS (regions of different shapes) around the nucleus.
Each Shell has a MAIN energy level, with smaller divisions in energy levels in each shell.
This is why Shells are also known as PRINCIPAL ENERGY LEVELS or PRINCIPAL QUANTUM SHELLS.
The symbol for shells is ‘n’, where n is the number of the shell.
n = 1 is the first shell, n = 2 is the second, & so on.
What is a Sub-Shell?
Each Principal Quantum Shell contains a number of sub-divisions.
Each sub-division is a SUB-SHELL (or SUB-LEVEL).
A sub-shell is a group of orbitals of the same shape & energy level.
For example, the 1s sub-shell contains 1 ‘s’ orbital. The 2p sub-shell contains 3 ‘p’ orbitals.
What are Orbitals?
A region of space around the nucleus which has a high probability of containing electrons.
Due to quantum physics weirdness (which we don’t have to know about yet), it’s not possible to know both the momentum & location of an electron.
Thus, orbitals show where electrons are MOST LIKELY to be at any given time.
Orbitals come in 4 different types, each with a different shape & energy level:
s, p, d, f (but we only learn about s, p & d for A-Levels).
Extra Question: why those letters?
Types of Orbitals
|Shapes||Occur in shells..||
Number of Orbitals per Shell
From the table, you can see that:
- The ‘s’ orbital is spherical in shape, & has only 1 possible orientation.
Thus, only 1 s orbital can exist in each shell.
- The ‘p’ orbital is dumb-bell-shaped, & has 3 possible orientations.
Thus, 3 p orbitals can exist in each shell.
- The ‘d’ orbital has various shapes, & has 5 possible orientations.
Thus, 5 p orbitals can exist in each shell.
Differentiating between Orientations of the Same Orbital Type
Let’s take the ‘p’ orbital for example.
Imagine a 3-dimensional Cartesian Plane around the orbital.
The ‘p’ orbital can be in 1 of 3 orientations, all 90 degrees from each other:
- “upright”: lying across the z-axis
- lying across the x-axis
- lying across the y-axis
Thus, we name these orbitals as:
Each orbital can hold a maximum of 2 electrons.
However, they do not have to be filled (an orbital can hold 0, 1 or 2 electrons, but never more than 2).
When writing out electron configurations of any atom, you have to write:
- the Principal Quantum Number (shell) (written as 1,2,3,…)
- the Sub-Shell (which is a group of orbitals of the same type) (written as s, p, or d)
- the Number of electrons in the sub-shell (written as a superscript number)
An atom with an electron configuration of:
1s2 2s2 2p6 3s1
2 electrons in the s-orbital of the 1st shell
2 electrons in the s-orbital of the 2nd shell
6 electrons in the p-orbitals of the 3rd shell
1 electron in the s-orbital of the 2nd shell
There’s a specific order to filling orbitals.
To know how to fill these orbitals, we will have to learn 3 basic rules..
3 Rules of Filling Orbitals & Shells:
- Aufbau’s Principle
- Pauli Exclusion Principle
- Hund’s Rule
Let’s examine what each rule means:
“Electrons enter the LOWEST available energy level”.
Always start filling up the orbitals at a lower energy level (sub-shell) first.
When these orbitals are full, only then can you fill the orbitals of the next level.
Do NOT fill an orbital of a higher energy before the previous sub-shell is full.
The sequence of sub-shells are (in ascending energy level): s, p, d.
This means that you must fill the ‘s’ orbital of a principal shell before filling the ‘p’ orbitals of the same shell, & lastly fill the ‘d’ orbitals of that shell.
Only then can you move on to fill the ‘s’ orbital of the NEXT shell.
If you recall, each shell can hold a maximum of 1 s, 3 p & 5 d orbitals.
Also, each orbital can hold a maximum of 2 electrons.
So, you have to put in 2 electrons in the (one & only) s orbital, then you can fill the 3 p orbitals with 2 electrons each, making a total of 6 electrons in the p sub-shell, etc.
Also recall that as the shell sizes increase, the number of sub-shells (& thus orbitals) inside them increases.
- the 1st Principal Shell only contains 1 s orbital
- the 2nd Principal Shell contains 1 s & 3 p orbitals
- the 3rd Principal Shell contains 1 s, 3 p, & 5 d orbitals
Here’s a diagram to summarise:
If you notice, now it’s clear why we learned previously that electrons fill shells in the order of 184.108.40.206.
2 = 2 electrons in the sub-shell 1s
8 = 2 electrons in the sub-shell 2s + 6 electrons in the sub-shell 2p
18 = 2 electrons in the sub-shell 3s + 6 electrons in the sub-shell 3p + 10 electrons in the sub-shell 3d
& so on…
HOWEVER, as with most rules, there are a few details that may seem confusing.
Look at this part:
As you can see, although the 3d orbital is in the 3rd shell, the 4s orbital in the 4th shell must be filled in FIRST because the 4s orbital has a lower energy level than 3d.
Shortcut: Cross-Out Method
A shortcut to conveniently know the order of filling up sub-shells that takes into account counter-intuitive things like these is the cross-out method.
Start filling the first row (1s) & cross out the sub-shells diagonally as you fill them with electrons.
Follow the direction of the arrow to know the next sub-shell to fill, until you reach the 1st column again.
Hund’s Rule of Maximum Multiplicity
“When in orbitals of equal energy, electrons will try to remain unpaired”.
Now we know that you have to fill sub-shells with lower energies first.
But how about the orbitals in the same sub-shell?
For example, when there are 3 p orbitals in the same sub-shell, do we fill the px orbital first, then move on to the py?
Each orbital can hold a maximum of 2 electrons; so do we put 2 electrons in one p orbital before putting the next in another p orbital?
Hund’s Rule states that electrons will remain UNPAIRED if possible.
This means that if you have 2 electrons to put into the p sub-shell, they will fill DIFFERENT orbitals to avoid each other as much as possible.
Only when all 3 p orbitals have 1 electron, then can the 2nd electron come to fill a p orbital.
The same rule applies to any sub-shell with multiple orbital: p, d, but NOT s (the s sub-shell only has 1 s orbital).
Pauli Exclusion Principle
“2 electrons in the same orbital must have opposite spin”.
A little piece of info before we tackle this one:
Electrons have an intrinsic property (just like mass, charge) called SPIN.
Just like how charge can be NEGATIVE or POSITIVE, spin can be UP or DOWN.
It’s not really angular momentum like we’re used to, but we don’t have to know what exactly spin is for A-Levels.
Now that you know spin has two opposite values, we can apply the Pauli Exclusion Principle.
All this means is that when 2 electrons fill an orbital, 1 must have an UP SPIN, while the other must have a DOWN SPIN.
This fact doesn’t matter if we’re just writing out the electron configuration of an atom
(for example: )
but it DOES matter if we show the spin of each electron by drawing….
Electrons in Boxes
(used to show both the CONFIGURATION & SPIN of each electron in the atom)
Each orbital is drawn as a box, which can be filled with 2 electrons.
The SPIN of the electrons is drawn as an arrow (Up Arrow = Spin Up, Down Arrow = Spin Down)
Here’s an example:
All 3 rules apply when drawing electrons in boxes, so be sure to be aware of all of them when filling electrons of an atom.
Here’s a handy gif that shows the electron-in-box diagrams for a few elements in ascending proton number:
(From this website)
These 3 rules above aren’t explicitly stated to be necessary to describe/memorise in A-Levels, but they are important & handy when you need to write out electron configurations (which IS a requirement!).
In summary, we learned:
- a more accurate representation of atomic structure
- how to write electron configurations for atoms
You can practice these skills by playing this short game:
In the next post, we’ll be covering IONISATION ENERGY.