Welcome to the World of Transition Metals!

In this chapter, we are going to explore the elements that live in the middle of your Periodic Table. You might know them as the metals that give us beautiful colored gemstones, powerful magnets, and the materials used to build skyscrapers. For your International AS Level (XCH11), we are going to focus on the "blueprints" of these atoms—their electronic configurations.

Don’t worry if the letters and numbers (like \(3d\) and \(4s\)) seem a bit like a secret code right now. We’re going to break that code together step-by-step! By the end of these notes, you’ll be able to write the electronic configuration for any element up to Krypton (\(Z=36\)).

1. Finding the d-block

The Periodic Table is divided into "neighborhoods" called blocks. These blocks are named after the sub-shell that is being filled with electrons.

  • s-block: Groups 1 and 2.
  • p-block: Groups 3 to 0 (or 13 to 18).
  • d-block: The big rectangular block in the middle!

The Transition Metals we study are located in the d-block. Specifically, we look at the first row, which starts with Scandium (\(Sc\)) and ends with Zinc (\(Zn\)).

Did you know? Even though they are in the middle, these elements are the reason why rubies are red and emeralds are green! It’s all down to how their electrons are arranged.

2. The Rules of the House (Electronic Configuration)

To understand how electrons fit into an atom, think of the atom as a hotel. Electrons are the guests, and they always want the cheapest room (the lowest energy level) first.

The Order of Filling

Usually, we fill shells in order: \(1s\), then \(2s\), then \(2p\), and so on. However, there is a tiny "glitch" in the system when we get to the transition metals. The \(4s\) sub-shell is actually slightly lower in energy than the \(3d\) sub-shell.

The Golden Rule: You must fill the \(4s\) room before you start putting electrons into the \(3d\) suite.

The sequence goes: \(...3p \rightarrow 4s \rightarrow 3d\)

Maximum Capacity

Remember how many "guests" each sub-shell can hold:

  • s sub-shell: 1 orbital (max 2 electrons)
  • p sub-shell: 3 orbitals (max 6 electrons)
  • d sub-shell: 5 orbitals (max 10 electrons)

Quick Review Box:
Always fill \(4s\) before \(3d\) because \(4s\) is lower in energy!

3. Two "Rebellious" Elements: Chromium and Copper

Most elements follow the rules perfectly. For example, Iron (\(Fe\), \(Z=26\)) is \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6\).

However, Chromium (\(Cr\)) and Copper (\(Cu\)) like to be different. They "steal" one electron from the \(4s\) sub-shell to help stabilize their \(3d\) sub-shell.

Chromium (\(Cr\), \(Z=24\))

Expected: \(...4s^2 3d^4\)
Actual: \(...4s^1 3d^5\)
Why? An atom is extra stable if the \(3d\) sub-shell is exactly half-full. It’s like having one person in every seat of a 5-seater car—it’s symmetrical and happy!

Copper (\(Cu\), \(Z=29\))

Expected: \(...4s^2 3d^9\)
Actual: \(...4s^1 3d^{10}\)
Why? An atom is even more stable if the \(3d\) sub-shell is completely full.

Common Mistake to Avoid: Don't forget these two! Examiners love to ask about \(Cr\) and \(Cu\) because they are the exceptions to the "fill \(4s\) first" rule.

4. Making Ions: The "Last In, First Out" Rule

This is the part that trips up many students, so let's use an analogy. Imagine the \(4s\) sub-shell is the front door of the atom, and the \(3d\) sub-shell is the living room.

When electrons enter, they go through the door (\(4s\)) and then sit in the living room (\(3d\)).
When electrons leave (to form a positive ion), they must leave through the door (\(4s\)) first!

Key Rule: When forming transition metal ions, always remove the \(4s\) electrons before the \(3d\) electrons.

Example: Iron(II) Ion (\(Fe^{2+}\))

1. Start with the atom: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6\)
2. Remove 2 electrons (the \(2+\) charge means 2 electrons left).
3. Take them from \(4s\) first!
Result: \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^6\) (The \(4s\) is now empty).

Key Takeaway: For any transition metal ion in the first row, you will usually see an empty \(4s\) sub-shell in the final answer.

5. Using "Electron-in-Boxes" Notation

Sometimes the syllabus asks you to draw the electrons. We use boxes to represent orbitals and arrows to represent electrons.

  • Each box can hold two arrows.
  • The arrows must point in opposite directions (this represents "opposite spin").
  • Hund’s Rule: In the \(d\) sub-shell, electrons will occupy their own empty box (orbital) before they start pairing up.

Analogy: It’s like people getting on a bus. Everyone takes an empty double-seat for themselves first. They only sit next to a stranger if all the rows already have one person in them!

Summary Checklist

✓ The d-block: The middle section of the table where \(d\) sub-shells are filled.
✓ Filling Order: \(4s\) fills before \(3d\).
✓ Exceptions: Chromium is \(4s^1 3d^5\) and Copper is \(4s^1 3d^{10}\).
✓ Ion Formation: Always remove \(4s\) electrons first.
✓ Boxes and Arrows: Fill orbitals singly before pairing, and use opposite arrows.

Don't worry if this seems tricky at first! Electronic configuration is a bit like learning a new musical instrument—once you practice the scales (the rules), the songs (the atoms) become much easier to play. Keep practicing with your Periodic Table!