Electrons: electronic energy levels, ionisation energies, atomic orbitals, extranuclear structure

Welcome to Atomic Structure: The World of Electrons!

Hello there! Today, we are diving into the most exciting part of the atom: the electrons. If the nucleus is the heart of the atom, the electrons are the personality. They determine how an atom behaves, who it "hangs out" with (bonding), and how much energy it carries.

Don’t worry if some of these terms like "orbitals" or "subshells" sound like science fiction at first. We will break them down piece by piece. By the end of these notes, you’ll be able to map out exactly where electrons live and understand the energy needed to "kidnap" them from an atom!

1. Where do Electrons Live? (Atomic Orbitals)

In lower secondary, you learned that electrons move in "shells." In H2 Chemistry, we look closer. Think of an atom like a hotel:

• Principal Quantum Number (\(n\)): This is the floor of the hotel. The higher the number, the further it is from the nucleus and the higher the energy.
• Subshells (\(s, p, d\)): These are the types of rooms on each floor.
• Orbitals: These are the beds in the rooms. Each orbital can hold a maximum of two electrons, but only if they have opposite spins (like two people sleeping head-to-toe!).

Types of Orbitals and Their Shapes

You need to know the shapes of these "rooms" because electrons don't move in perfect circles; they move in probability clouds.

• s orbitals: These are spherical (like a ball). There is one \(s\) orbital in every principal shell.
• p orbitals: These are dumbbell-shaped. They come in sets of three: \(p_x\), \(p_y\), and \(p_z\), arranged along the axes.
• d orbitals: These are more complex (mostly clover-leaf shaped). There are five \(d\) orbitals in a set.

Did you know? An "orbital" isn't a physical container. It’s just a region of space where there is a 95% chance of finding an electron!

The Energy Ladder

Electrons are lazy—they always want to be in the lowest energy room available. Usually, the order is:
\(1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p\)

Quick Review & Common Mistake: Notice that the \(4s\) subshell is actually lower in energy than the \(3d\) subshell when it's empty. So, we fill \(4s\) before \(3d\). Think of it like a hotel where the room on the 4th floor is slightly cheaper than the luxury suite on the 3rd floor!

Key Takeaway: Electrons occupy orbitals of the lowest energy first. Principal shells (\(n\)) contain subshells (\(s, p, d\)), and each orbital holds 2 electrons max.

2. Writing Electronic Configurations

To show where electrons are, we use a code. For example, Magnesium (12 electrons) is: \(1s^2 2s^2 2p^6 3s^2\).

The Three Golden Rules

1. Aufbau Principle: Always fill the lowest energy level first.
2. Pauli Exclusion Principle: Two electrons in the same orbital must have opposite spins.
3. Hund’s Rule: Electrons will occupy orbitals singly first before pairing up.
Analogy: On a bus, people usually sit in empty rows before sitting next to a stranger. Electrons do the same thing!

Special Cases: Chromium (\(Z=24\)) and Copper (\(Z=29\))

These two are "rebels" because a half-filled or fully-filled \(d\) subshell is extra stable.
• Cr: \([Ar] 3d^5 4s^1\) (NOT \(3d^4 4s^2\))
• Cu: \([Ar] 3d^{10} 4s^1\) (NOT \(3d^9 4s^2\))

Making Ions

When an atom becomes a positive ion, it loses electrons.
CRITICAL RULE: For transition metals (like Iron or Copper), electrons are removed from the \(4s\) subshell BEFORE the \(3d\) subshell. Even though \(4s\) filled first, it also empties first because once \(3d\) is occupied, \(4s\) becomes higher in energy.

Key Takeaway: Always follow the filling order, but watch out for the stable \(d^5\) and \(d^{10}\) configurations in Cr and Cu. For ions, "First in (\(4s\)), first out!"

3. Ionisation Energy (I.E.)

First Ionisation Energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous \(1+\) ions.

Equation: \(X(g) \rightarrow X^+(g) + e^-\)

Factors Affecting I.E. (The "Tug-of-War")

Think of the nucleus as a magnet and the electron as a paperclip. How hard is it to pull the paperclip away?
1. Nuclear Charge (\(+\)): More protons = stronger magnet = higher I.E.
2. Distance (\(-\)): Further away = weaker attraction = lower I.E.
3. Shielding (\(-\)): Inner shells of electrons act like a "shield," blocking the pull of the nucleus = lower I.E.

Trends in the Periodic Table

• Across a Period: I.E. generally increases. Protons increase (stronger magnet) while shielding stays roughly the same.
• Down a Group: I.E. decreases. New shells are added, so the electron is much further away and has more shielding.

Small Dips (The "Exceptions")

Sometimes the trend isn't a straight line. For example, in Period 3:
• Mg to Al dip: Al's outer electron is in a \(3p\) orbital, which is higher in energy and further from the nucleus than Mg's \(3s\). It’s easier to remove!
• P to S dip: In Sulfur, two electrons are paired in the same \(3p\) orbital. They repel each other, making it easier for one to be kicked out.

Key Takeaway: I.E. measures how tightly the nucleus holds onto electrons. It increases across a period and decreases down a group, with small dips due to subshell energies and electron repulsion.

4. Successive Ionisation Energies

This is the energy needed to remove the 1st, then the 2nd, then the 3rd electron from the same atom.

How to Read the Data

Successive I.E.s always increase because you are removing a negative electron from an increasingly positive ion. However, look for a HUGE JUMP in the numbers.

Example: An element has these I.E. values:
1st: 578 kJ/mol
2nd: 1,817 kJ/mol
3rd: 2,745 kJ/mol
--- BIG JUMP ---
4th: 11,578 kJ/mol

Step-by-step Deduction:
1. The jump happens after the 3rd electron.
2. This means there were 3 electrons in the outer shell.
3. The 4th electron was taken from an inner shell closer to the nucleus (much harder!).
4. Therefore, the element is in Group 13.

Quick Review Box:
• Small jump? You are likely moving to a different subshell (e.g., from \(p\) to \(s\)).
• Big jump? You have moved to a new principal quantum shell (closer to the nucleus).

Key Takeaway: Successive I.E. data is like a "map" of the atom’s shells. The number of electrons removed before the first big jump equals the group number of the element.

Summary Checklist

• Can you draw the shapes of \(s\) and \(p\) orbitals?
• Do you remember to fill \(4s\) before \(3d\)?
• Can you explain why I.E. increases across a period using "nuclear charge" and "shielding"?
• Can you identify an element's group from a list of ionisation energies?

Don't worry if this feels like a lot to memorize. Just keep practicing the electronic configurations—they are the foundation for everything else in Chemistry!