Chemistry | Electron configurations

Welcome to Structure 1.3: Electron Configurations!

Hello future chemists! This chapter, Electron Configurations, might seem abstract, but it is one of the most foundational concepts in all of Chemistry. Why? Because the way an atom arranges its electrons determines virtually everything about how it behaves, who it bonds with, and what kind of substances it forms.

Think of the electrons as the "personality" of the atom. We are going to learn the rules governing where these electrons live, how they fill up their available spaces, and how to write down their addresses. Don't worry if this seems tricky at first; we will use lots of analogies to break it down!

Part 1: Energy Levels and Sub-levels (The Address System)

The Principal Energy Level (Shells)

We know from the previous chapter that electrons exist outside the nucleus. The first step in describing their arrangement uses the Principal Quantum Number, represented by the letter n.

• n tells us the main energy level, or shell, the electron occupies.
• The larger the value of n (\(n = 1, 2, 3, 4, ...\)), the further the electrons are, on average, from the nucleus, and the higher their energy.

Analogy: Think of n as the "floor" of a building. The first floor (\(n=1\)) is the lowest energy and closest to the ground (nucleus). The 4th floor (\(n=4\)) is much higher energy.

The Sub-levels (The Apartments)

Inside each principal energy level (shell), there are smaller energy groups called sub-levels or subshells. These are named using the letters s, p, d, and f.

The number of sub-levels available within a shell is equal to n.

• For \(n=1\), only the s sub-level exists.
• For \(n=2\), s and p sub-levels exist.
• For \(n=3\), s, p, and d sub-levels exist.
• For \(n=4\), s, p, d, and f sub-levels exist.

Electron Capacities of Sub-levels

Each sub-level contains a specific number of tiny regions of space called orbitals, and each orbital can hold a maximum of two electrons.

Sub-level Type	Number of Orbitals	Maximum Electrons
s (spherical shape)	1 orbital	2 electrons
p (dumbbell shape)	3 orbitals	6 electrons
d (complex shape)	5 orbitals	10 electrons
f (very complex shape)	7 orbitals	14 electrons

Key Takeaway: Electrons fill up the spaces in the order s < p < d < f, and the total capacity of an energy level \(n\) is \(2n^2\).

Part 2: The Rules for Filling Orbitals

Electrons are organized by three fundamental rules. These rules ensure that electrons occupy the most stable (lowest energy) configuration possible.

Rule 1: The Aufbau Principle (The "Building Up" Rule)

The German word Aufbau means "building up."

The principle states that electrons will always fill the lowest available energy levels first before moving to higher levels.

While you might assume the order is simply \(1s, 2s, 2p, 3s, 3p, 3d, 4s...\), the energy levels sometimes overlap, making the true filling order slightly complicated after the third shell.

• Crucial Order Alert: \(4s\) fills before \(3d\). (It is lower in energy!)

Memory Aid: The Diagonal Rule

Write the sub-levels in columns and follow the diagonal arrows to get the correct filling order:
\(1s\)
\(2s \ 2p\)
\(3s \ 3p \ 3d\)
\(4s \ 4p \ 4d \ 4f\)
... and so on.

The order is: \(1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, ...\)

Rule 2: The Pauli Exclusion Principle

This rule relates to electron "spin." Electrons in the same orbital must have opposite spins.

• Principle: No two electrons in the same atom can have the exact same set of quantum numbers (or, more simply, an orbital can hold a maximum of two electrons, and they must spin in opposite directions).

Rule 3: Hund's Rule (The "Bus Seat" Rule)

Hund's rule applies when you have multiple orbitals with the exact same energy (like the three p-orbitals or the five d-orbitals). These are called degenerate orbitals.

• Principle: Every orbital in a sub-level must be singly occupied with electrons having parallel spins (all facing the same way) before any orbital is doubly occupied.

Analogy: Think about getting onto an empty bus. If you have three seats in a row (p-orbitals), you don't sit right next to a stranger if you can help it! You spread out, one person per seat, until every seat has one person. Only then, if more people get on, do you start pairing up. Electrons behave the same way!

Part 3: Writing Electron Configurations

Spectroscopic (spdf) Notation (SL & HL)

This is the standard, concise way of writing the electron address.

It follows the format: (Principal Quantum Number) (Sub-level Letter)^{(Number of Electrons)}

Example: Oxygen (O)
Oxygen has 8 electrons (its atomic number is 8). We fill them in order:

1. Fill \(1s\): \(1s^2\) (2 electrons used, 6 remaining)
2. Fill \(2s\): \(2s^2\) (4 electrons used, 4 remaining)
3. Fill \(2p\): \(2p^4\) (4 electrons remaining, and the p-sublevel can hold up to 6)

The full configuration is: \(1s^2 2s^2 2p^4\)

Noble Gas Abbreviation (The Shortcut)

For large atoms (HL students, this is especially useful!), writing out the whole configuration can be tedious. We use the nearest preceding Noble Gas (Group 18 element) to simplify.

Noble gases have full outer shells, making their configurations very stable.

Example: Potassium (K)
Potassium has 19 electrons.

Full configuration: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^1\)

The nearest preceding noble gas is Argon (Ar), which has 18 electrons. Argon's configuration is \(1s^2 2s^2 2p^6 3s^2 3p^6\).

Abbreviated configuration: \([Ar] 4s^1\)

Common Mistake Alert: \(4s\) versus \(3d\)

Remember, the \(4s\) sub-level is filled *before* the \(3d\) sub-level because \(4s\) is lower in energy.

\(1s^2 ... 3p^6 \ 4s^2 \ 3d^{10} \ 4p^6 ...\)

However, when transition metals lose electrons to form ions, they always lose the electrons from the highest principal quantum number (\(n\)) first. This means they lose the 4s electrons before the 3d electrons, even though they were filled first.

Did you know? The slight difference in energy between \(4s\) and \(3d\) levels is why transition metals can form ions with multiple different charges (e.g., Fe²⁺ and Fe³⁺).

Part 4: Valence Electrons (The Chemical Fingerprint)

What are Valence Electrons?

The most important part of the electron configuration for chemists is the number of valence electrons.

• Valence Electrons are the electrons found in the atom's outermost occupied principal energy level (the highest value of n).
• These are the electrons involved in bonding and chemical reactions.
• The group number (1 to 18) on the Periodic Table (for Groups 1, 2, and 13-18) often tells you the number of valence electrons.

Example Recap: Oxygen (O)
Configuration: \(1s^2 2s^2 2p^4\)

The highest principal energy level is \(n=2\).
The total number of electrons in the \(n=2\) level is \(2\) (from 2s) + \(4\) (from 2p) = 6 valence electrons.