Atomic Structure & Electron Configuration

Every atom is built from three particles. Protons and neutrons sit packed together in a tiny, dense nucleus at the center of the atom. Electrons move around that nucleus in a much larger cloud of space called the electron cloud, organized into orbitals. Almost all of an atom's mass is in the nucleus, but almost all of its volume is empty space occupied by the electron cloud.

Two numbers describe an atom's identity. The atomic number (Z) is the number of protons — it defines which element the atom is, and every atom of a given element has the same Z. The mass number (A) is the total count of protons plus neutrons in the nucleus (electrons are too light to count meaningfully toward mass).

Isotope notation writes the mass number as a superscript before the element symbol, with the atomic number as a subscript before it: 126C is read "carbon-12," meaning Z = 6 (6 protons, since carbon's atomic number is always 6) and A = 12 (so 12 − 6 = 6 neutrons). The same information is written in words as carbon-12.

Isotopes are atoms of the same element (same Z, same number of protons) that have different numbers of neutrons — so they have different mass numbers. Because isotopes have identical electron counts and configurations, they behave almost identically in chemical reactions. What differs is mass and nuclear stability, which is why isotopes matter in physics, medicine, and dating rock or fossil samples but rarely change an element's chemistry.

Electrons don't orbit the nucleus in neat circles like planets. Instead, each electron occupies an orbital — a region of space where it has a high probability of being found. Orbitals are grouped first by energy level (or shell, given the number n = 1, 2, 3…) and then by subshell type (s, p, d, f), each with a different shape and a different maximum number of electrons it can hold.

Each individual orbital, regardless of shape, holds a maximum of two electrons. A subshell's total capacity is just orbitals × 2 — that's why p (3 orbitals) holds 6 and d (5 orbitals) holds 10. Shell n = 1 only has an s subshell; n = 2 has s and p; n = 3 has s, p, and d; n = 4 and up have s, p, d, and f.

Three rules together determine how electrons fill an atom's orbitals from the ground up.

The Aufbau filling order does NOT simply go 1s, 2s, 2p, 3s, 3p, 3d, 4s… — the 4s subshell is actually lower in energy than 3d, so it fills first. The standard order, memorized by following the diagonal arrows on an Aufbau diagram, is:

A full electron configuration lists every subshell an atom's electrons occupy, in fill order, with the electron count in that subshell written as a superscript. The superscripts for one atom must add up to that atom's total number of electrons (Z, for a neutral atom).

Check yourself by adding the superscripts: for iron, 2 + 2 + 6 + 2 + 6 + 2 + 6 = 26, matching iron's atomic number. That running total is the fastest way to catch a dropped or extra electron.

Full configurations get long for larger atoms, so chemists abbreviate the inner, already-filled shells using the symbol of the noble gas that has that same configuration, in brackets, followed by only the remaining outer electrons. Pick the noble gas with the largest atomic number that is still smaller than the element you're writing.

1. Write the full electron configuration for phosphorus (P, Z = 15). 2. Silver-107 and silver-109 are both isotopes of silver (Z = 47). How many neutrons does each have? 3. Write the noble-gas shorthand configuration for potassium (K, Z = 19). Using elements-electron-config.csv: 4. For each element, write the full electron configuration in Aufbau order. Check that the superscripts add up to the atomic number. 5. For each element, write the noble-gas shorthand configuration. 6. For each element, give the number of valence electrons (electrons in the highest occupied energy level).