Electron structure - Chemistry A - H032 - Cambridge OCR AS Level

Welcome to the World of Electrons!

In your previous science classes, you likely learned that electrons orbit the nucleus in simple circles. While that was a great starting point, the truth is much more interesting! In this chapter, we are going to explore electron structure. Understanding where electrons live and how they move is the "secret code" to chemistry—it explains why some elements are explosive, why some are lazy (noble gases), and how atoms stick together to form everything in the universe.

Don't worry if this seems a bit abstract at first. We’ll use plenty of analogies to make these invisible particles feel much more real!

1. Energy Levels: The Atom's "Floors"

Electrons aren't just floating around randomly; they exist in specific shells (also called principal energy levels). You can think of an atom like a hotel where the nucleus is the ground floor, and the electrons are guests staying in rooms on the floors above.

The Rules of the Hotel:

Each shell has a maximum number of electrons it can hold. We use the formula \(2n^2\) (where \(n\) is the shell number) to figure this out:

Shell 1 (n=1): Holds up to 2 electrons \( (2 \times 1^2) \)
Shell 2 (n=2): Holds up to 8 electrons \( (2 \times 2^2) \)
Shell 3 (n=3): Holds up to 18 electrons \( (2 \times 3^2) \)
Shell 4 (n=4): Holds up to 32 electrons \( (2 \times 4^2) \)

Quick Review: As the shell number increases, the distance from the nucleus increases, and the energy level of the electrons also increases.

Key Takeaway: Shells are the main energy levels, numbered \(n = 1, 2, 3, 4\), and they get larger and hold more electrons as you move away from the nucleus.

2. Atomic Orbitals and Sub-shells

If shells are the "floors" of our hotel, sub-shells are the different types of hallways, and orbitals are the actual rooms.

What is an Atomic Orbital?

An atomic orbital is a region around the nucleus that can hold up to two electrons. Crucially, these two electrons must have opposite spins (imagine them spinning in different directions so they don't repel each other too much).

Types of Sub-shells and Their Shapes

There are three main types of sub-shells you need to know for AS Level:

s-sub-shell: Contains 1 orbital (holds 2 electrons total). Shape: Spherical (like a ball).
p-sub-shell: Contains 3 orbitals (holds 6 electrons total). Shape: Dumbbell-shaped.
d-sub-shell: Contains 5 orbitals (holds 10 electrons total). These have more complex shapes!

Did you know? Even though a p-sub-shell has three orbitals (called \(p_x, p_y,\) and \(p_z\)), they all have the same energy level. We call orbitals with the same energy degenerate.

Key Takeaway: Orbitals are "clouds" where electrons live. An s-orbital is a sphere; a p-orbital is a dumbbell. Every single orbital, no matter the type, can hold a maximum of two electrons.

3. How Orbitals are Filled

Electrons are a bit like people boarding a bus—they follow specific rules to decide where to sit. There are three main "rules" for filling orbitals:

Rule 1: The "Lowest Energy First" Rule (Aufbau Principle)

Electrons always fill the orbital with the lowest energy first. It’s easier to stay on the bottom floor than to climb the stairs!

The Energy Sequence: \(1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \rightarrow 4p\)

Common Mistake Alert! Notice that the 4s sub-shell actually has a lower energy than the 3d sub-shell. Therefore, 4s fills up before 3d starts. Think of 4s as a "shortcut" that electrons take because it's slightly easier to get into.

Rule 2: The "Bus Seat" Rule (Hund’s Rule)

If there are multiple orbitals of the same energy (like the three rooms in a p-sub-shell), electrons will occupy them singly first before they start pairing up. Just like on a bus, you’d usually sit in an empty row before sitting next to a stranger!

Rule 3: Opposite Spins

When two electrons finally do pair up in an orbital, they must have opposite spins. In diagrams, we represent this using one arrow pointing up and one pointing down: \( \uparrow\downarrow \).

Summary Table: Sub-shell Capacity

s: 1 orbital = 2 electrons max
p: 3 orbitals = 6 electrons max
d: 5 orbitals = 10 electrons max

4. Writing Electron Configurations

We use a standard notation to show where all the electrons in an atom are. This is called the electron configuration.

Example: Oxygen (Atomic Number 8)
1. Start at the bottom: \(1s\) holds 2 electrons \(\rightarrow 1s^2\)
2. Next is \(2s\): holds 2 electrons \(\rightarrow 2s^2\)
3. Next is \(2p\): we have 4 electrons left (8 total - 4 used) \(\rightarrow 2p^4\)
Full configuration: \(1s^2 2s^2 2p^4\)

Step-by-Step for Atoms up to Z=36 (Krypton):
1. Find the atomic number (this is the number of electrons).
2. Fill sub-shells in order: \(1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^2, 3d^{10}, 4p^6\).
3. Stop when you run out of electrons!

A Note on Ions:
When an atom becomes an ion, it loses or gains electrons:
- Negative ions (Anions): Add electrons (e.g., \(F^-\) has 10 electrons instead of 9).
- Positive ions (Cations): Remove electrons (e.g., \(Mg^{2+}\) has 10 electrons instead of 12).
Memory Aid: Pawsitive ions are Positive (like a cat!).

Crucial Rule for Cations: When removing electrons for 4s and 3d elements, electrons are lost from the 4s sub-shell first, even though it was filled first. They are the "last in, first out" in terms of physical distance from the nucleus.

Quick Review Box:
- Shell 1: \(1s\)
- Shell 2: \(2s, 2p\)
- Shell 3: \(3s, 3p, 3d\)
- Shell 4: \(4s, 4p, 4d, 4f\)
- The filling order trick: 4s comes before 3d!

Key Takeaway: Electron configuration is a map of where electrons are. Use the periodic table as a guide: Groups 1-2 are the s-block, Groups 13-18 are the p-block, and the transition metals are the d-block.

5. Summary of Electron Structure

Shells are the main energy levels (\(n=1, 2, 3, 4\)).
Sub-shells (\(s, p, d\)) exist within shells.
Orbitals are specific regions holding 2 electrons with opposite spins.
Shapes: \(s\) is a sphere, \(p\) is a dumbbell.
Order: Fill \(1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p\).
Ions: Remember to add or subtract electrons from the total number before writing the configuration.

You've got this! Practice writing the configurations for the first 20 elements, and the pattern will start to feel like second nature.

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