How electrons are arranged

You're coming in knowing that an atom has a nucleus and electrons in fuzzy 3D regions around it. This module is about the structure of those regions: why they have the energies and shapes they do, and what that arrangement determines about an atom's behaviour. If you internalise this material, the rest of chemistry will feel like consequences rather than rules.

Energy levels — the staircase, not the ramp

Established Electrons in an atom can only have certain specific energies, not arbitrary ones. This is quantization, and it is one of the foundational discoveries of quantum mechanics. Think of it as a staircase rather than a ramp — you can stand on step 1 or step 2, but not at step 1.5.

These allowed energies are organised into shells, labelled n = 1, 2, 3, 4 and so on. Shell 1, closest to the nucleus and lowest energy, holds up to 2 electrons. Shell 2 holds up to 8. Shell 3 holds up to 18, though it gets complicated. Shell 4 holds up to 32.

Why these specific numbers? They come from the mathematics of quantum mechanics — specifically, the number of allowed orbital shapes at each energy level. We're about to see this directly.

Subshells and orbitals — the real picture

Each shell is divided into subshells, named with letters: s, p, d, f. (The letters are leftovers from spectroscopy — sharp, principal, diffuse, fundamental. The historical naming doesn't matter; the structure does.)

Each subshell contains orbitals, and each orbital can hold exactly 2 electrons (with opposite "spins," which we set aside for now). The shapes:

Subshell	Orbitals	Max electrons	Shape
s	1	2	sphere
p	3	6	dumbbell (3 orientations: x, y, z)
d	5	10	complex (cloverleaf and others)
f	7	14	very complex

So shell 1 contains only the 1s orbital, holding 2 electrons. Shell 2 contains 2s and 2p, holding 2 + 6 = 8 electrons. Shell 3 contains 3s, 3p, and 3d, holding 2 + 6 + 10 = 18. The numbers from earlier are consequences, not rules.

This is the key insight: orbital shapes are 3D regions with specific geometries. The s orbital is a sphere. The three p orbitals are dumbbells pointing along the x, y, and z axes. These shapes are real, and they determine how atoms can bond — because bonds form where electron orbitals overlap.

How electrons fill orbitals

Established Electrons fill orbitals from lowest energy upward. This is the Aufbau principle (German for "building up"). Within a subshell, they spread out across orbitals before pairing up — electrons repel each other and prefer their own orbital before sharing one. The filling order:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s …

Notice that 4s fills before 3d. This is because 4s actually has slightly lower energy than 3d, despite being in a higher shell. Energy levels start overlapping at higher shells, which is why the periodic table has those funny block layouts.

Writing electron configurations

The notation is [shell][subshell][number of electrons]. A few worked examples:

Hydrogen, 1 electron: 1s¹
Helium, 2 electrons: 1s² — shell 1 full
Lithium, 3 electrons: 1s² 2s¹
Carbon, 6 electrons: 1s² 2s² 2p²
Oxygen, 8 electrons: 1s² 2s² 2p⁴
Neon, 10 electrons: 1s² 2s² 2p⁶ — shell 2 full

Valence electrons — the only ones that matter for chemistry

The electrons in the outermost shell are valence electrons. These are the only ones that participate in chemical bonding. Inner-shell electrons are "core" electrons and are essentially inert spectators.

This single fact is why the periodic table works. Elements in the same column have the same number of valence electrons and therefore behave chemically similarly.

Hydrogen has 1 valence electron. Carbon has 4 (the 2s² 2p² in shell 2). Oxygen has 6. Neon has 8 (a full shell). Sodium has 1 (the lonely 3s¹).

The octet rule — why atoms bond at all

Here is the principle that explains essentially all of chemistry:

Atoms are most stable when their outer shell is full.

Established A full outer shell is usually 8 electrons — hence "octet rule." Hydrogen and helium are exceptions, needing only 2 because shell 1 maxes out at 2.

Atoms with full outer shells already (helium, neon, argon — the noble gases) are extremely unreactive. They don't bond, don't form compounds, just float around as single atoms.

Every other atom is, in a sense, "trying" to get to a full outer shell. This is teleological language and atoms don't actually want anything — but the energy landscape pushes them toward configurations with full outer shells, the same way water "wants" to flow downhill.

There are three ways to achieve this: lose electrons (if you have just a few valence electrons, dump them), gain electrons (if you're close to full, grab some), or share electrons (split the difference with another atom). These three strategies correspond exactly to the three main types of chemical bonds, which is where M-Chem-03 picks up.