Imagine an electron as a tiny magnet trapped in orbit around a positively charged nucleus, held there purely by electrostatic attraction. To rip that electron away completely — so it becomes a free, gaseous electron with no more pull from the nucleus — you have to do work against that attraction. That "work" is the ionisation energy.

Formally: the ionisation energy (IE) of an element is the energy required to remove one mole of electrons from one mole of gaseous atoms of that element, to form one mole of gaseous 1+ ions. It's measured under standard conditions (298 K, 101 kPa) and given in kJ mol⁻¹.

Notice three "one mole" conditions packed into that definition — this precision matters in exam answers. Also notice the state symbols: everything must be gaseous. If atoms were in a solid or liquid, neighbouring particles would interfere with the measurement, so chemists always define IE for isolated gas-phase particles.

Nothing stops you from removing a second electron once the first is gone — you just now have to pull it away from a positively charged ion instead of a neutral atom, which is harder (more attraction, less repulsion between remaining electrons). This gives the second ionisation energy, third, and so on.

Notice the huge jump between IE₂ (1145) and IE₃ (4912) for calcium. That's not random — it's the whole basis of how we deduce electron configuration from data, covered in the next section.

Whether an ionisation energy is "big" or "small" for a given electron comes down to a tug-of-war between four physical factors. Think of it like judging how hard it would be to pull a fridge magnet off a fridge — it depends on how strong the fridge's magnetism is, how far away the magnet is, whether something is blocking the pull, and whether another magnet is pushing it away.

This is one of the most exam-tested skills in this chapter, so let's go slowly. Chemists plot log(ionisation energy) against the number of electrons removed. Using logs compresses the huge range of values (hundreds to tens of thousands of kJ mol⁻¹) onto one readable graph.

The logic, step by step: as you strip electrons away one at a time, you're pulling from a progressively more positive ion each time, so successive IEs always increase. But they don't increase smoothly — they jump. A small jump happens when you finish one subshell and start the next (e.g. moving from 3p into 3s within the same shell). A big jump happens when you finish an entire shell and drop down into a shell much closer to the nucleus — that electron is far more strongly attracted (factor: distance ↓, shielding ↓, charge felt ↑), so it takes dramatically more energy to remove.

This means: if you count how many electrons were removed before the biggest jump, that number tells you how many electrons were in the outer (valence) shell — which tells you the group number.

Worked walkthrough — Sodium, Magnesium, Aluminium

Sodium: huge jump between IE₁ (494) and IE₂ (4560) — nearly 10×. Only one electron came off before that massive jump, so sodium has 1 outer electron → Group 1. Config: 1s² 2s² 2p⁶ 3s¹.

Magnesium: jump is between IE₂ (1450) and IE₃ (7740) — over 5×. Two electrons came off before the jump → Group 2. Config: 1s² 2s² 2p⁶ 3s².

Aluminium: jump is between IE₃ (2740) and IE₄ (11600) — over 4×. Three electrons came off before the jump → Group 3. Config: 1s² 2s² 2p⁶ 3s² 3p¹.

Calcium's configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s². Let's trace exactly why each successive IE changes the way it does — this kind of reasoning is exactly what examiners want to see written out.

Think of an atom like a huge apartment building. The building has floors (shells, numbered by principal quantum number n = 1, 2, 3...). Each floor is divided into room types (subshells: s, p, d, f). And each room type contains a fixed number of individual rooms (orbitals), and each room can only fit two people (electrons) — and importantly, those two people must have opposite "spins," like two magnets facing opposite ways so they can coexist peacefully in a cramped space.

You've already seen this in action above, but here's the reusable method as a clean checklist you can apply to any successive-IE table:

1. Scan through the list of IE values and find the single biggest percentage jump between consecutive values.
2. Count how many electrons were removed before that jump — this equals the number of electrons in the outer (valence) shell.
3. That number = the group number of the element (for main-group elements).
4. Smaller jumps within the same run (before the big one) hint at subshell changes, e.g. s → p.

Electrons behave like tiny spinning charges, and that spin creates a miniature magnetic field — picture each electron as a microscopic bar magnet with a North and South pole, spinning either clockwise or anticlockwise.

Electrons with the same spin repel each other (this is spin-pair repulsion). So when several empty orbitals of equal energy are available within a subshell, electrons will spread out — one to each empty orbital, all spinning the same way — before any of them pair up. Only once every orbital in that subshell has one electron will additional electrons start pairing up (with opposite spin, to minimise repulsion). This behaviour is called Hund's Rule.

Why does this matter energetically? Two electrons crammed into the same orbital repel each other more than the energy "cost" of the alternative — but jumping up to a completely different, higher-energy empty orbital costs even more energy than that repulsion. So electrons take the middle path: spread out within a subshell first (avoiding repulsion cheaply), then pair up within that same low-energy subshell rather than jumping to a pricier one.

This is simply a visual way of drawing out an electron configuration. Each box = one orbital. Arrows show spin (↑ and ↓ for opposite spins). Boxes are drawn left-to-right in order of increasing energy.

An s orbital is shaped like a solid sphere centred on the nucleus — imagine a fuzzy ball of probability where the electron is likely to be found, with the nucleus sitting right at the centre. Every shell has exactly one s orbital, and as the shell number increases, the sphere simply gets bigger (a 3s orbital is a larger sphere than a 2s orbital, which is larger than 1s).

A p orbital looks like two balloons tied together at the nucleus, pointing in opposite directions along one axis — chemists call this a "dumbbell" shape. Every shell from n=2 upward has three p orbitals, oriented perpendicular to one another along the x, y, and z axes — labelled p_x, p_y, and p_z. (The first shell, n=1, has no p orbital at all — only s.)

Just like s orbitals, p orbital lobes get larger and longer as the shell number increases — a 3p orbital's lobes stretch further from the nucleus than a 2p orbital's.

So the second shell (n=2) contains two subshells (2s and 2p), and those subshells together contain four orbitals (one 2s + three 2p), which together can hold a maximum of 8 electrons (2 + 6) — which is exactly why the second row of the Periodic Table has 8 elements!

Electrons always fill the lowest-energy available space first (this is called the Aufbau Principle, though your syllabus may not name it explicitly). Two rules govern energy order:

• Shells increase in energy as n increases: n=1 < n=2 < n=3...
• Within a shell, subshells increase in energy: s < p < d < f

Combine these and you'd expect 3d to fill before 4s (since "3" is a lower shell number than "4"). But here's the twist: the 3d subshell actually sits at slightly higher energy than the 4s subshell, so 4s fills first — this is the single most important exception in this whole chapter, and it explains a lot of odd-looking behaviour in the d-block.

The full configuration writes out every single subshell from 1s upward. The shorthand configuration replaces the "core" electrons (everything up to the nearest preceding noble gas) with that noble gas's symbol in square brackets, then only writes out the extra electrons beyond it.

Important convention: even though the 4s subshell fills before 3d, the full configuration is conventionally still written in numerical shell order (3d before 4s) once both are occupied — e.g. gallium is written 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p¹, not with 4s before 3d.

Based on the normal filling rules, you'd predict:

• Chromium (Cr, 24 electrons): [Ar] 3d⁴ 4s²
• Copper (Cu, 29 electrons): [Ar] 3d⁹ 4s²

But experimentally, that's not what happens. The real configurations are:

One electron "jumps" from 4s into 3d in both cases. Why? Because a half-full d subshell (d⁵, one electron in each of the 5 d orbitals) or a completely full d subshell (d¹⁰) is unusually energetically stable — more stable than the "expected" arrangement. The atom effectively trades a small amount of energy to promote one electron, in exchange for a bigger stability payoff from the symmetric, evenly-spread d arrangement.

The Periodic Table isn't arranged randomly — its entire shape is a direct map of electron configuration. Elements are grouped into blocks based on which subshell their highest-energy (valence) electron sits in:

This connects directly back to reactivity. Elements in the same group share the same number of outer electrons, so they behave similarly. Elements in the same period share the same number of occupied shells.

What Examiners Are Actually Looking For

• Precise definitions using the exact phrasing: "one mole of electrons," "one mole of gaseous atoms/ions," "gaseous ions."
• Correctly balanced equations with state symbols for every ionisation energy question.
• Clear reasoning that links jumps in successive IE data to shell/subshell changes — don't just state the answer, explain the "why."
• Correct use of terminology: "shielding," "nuclear charge," "spin-pair repulsion," "orbital" vs "subshell" vs "shell" used precisely (these are NOT interchangeable words).
• Full electron configurations written with correct superscripts and in the conventional numerical order.