Atomic Structure and Electron Configuration — Quick Reference

Atomic weight is the weighted average of naturally occurring isotopes, reported in amu (AMU).
Atomic number Z equals the number of protons; neutrons may vary, yielding different isotopes of the same element.
Neutral atoms have the same number of electrons as protons; electrons determine reactivity.
Example: chlorine isotopes
- Isotopes: Cl-35 (mass = $34.97\ \mathrm{amu}$ , abundance = $0.7578$ ) and Cl-37 (mass = $36.97\ \mathrm{amu}$ , abundance = $0.2422$ ).
- Weighted average (atomic weight):
  $A_{avg} = 34.97(0.7578) + 36.97(0.2422) \approx 35.45\ \mathrm{amu}$
- Atomic weight shown on the periodic table is this weighted average, rounded to significant figures.
Practical note: to compute an unknown element’s atomic weight you would need the masses and their experimental abundances.

Nucleus contains protons (Z) and neutrons; electrons reside outside in shells.
Shells are labeled by principal quantum number n: n = 1, 2, 3, … (first, second, third shells, etc.).
Subshells within each shell: s, p, d, f (lowest to highest energy: s < p < d < f).
Orbitals and capacities:
- s subshell: 1 orbital → 2 electrons
- p subshell: 3 orbitals → 6 electrons
- d subshell: 5 orbitals → 10 electrons
- f subshell: 7 orbitals → 14 electrons
Maximum electrons per shell (by sum of subshell capacities):
- n = 1: 2
- n = 2: 8
- n = 3: 18
- n = 4: 32
Electron distance from nucleus corresponds to energy: lower energy (closer) electrons are more strongly attracted to the nucleus.

Aufbau principle: electrons fill lowest-energy orbitals first (start with 1s).
Energy ordering can cause detours (e.g., 4s can fill before 3d in many cases).
Periods reflect shell filling patterns:
- 1st period: 2 elements (1s shell)
- 2nd period: 8 elements (2s and 2p)
- 3rd period: 8 elements (3s and 3p, with 3d starting later in the period diagram)
- 4th period: 18 elements (4s, 3d, 4p, etc.)
Concept: the periodic table is organized to mirror shell/subshell filling, not to require memorizing all numbers.

Electron configuration describes where electrons reside: lowest energy to highest.
Notation examples (ascending order of energy):
- Hydrogen: $1s^1$
- Helium: $1s^2$
- Lithium: $1s^2 2s^1$
- Carbon: $1s^2 2s^2 2p^2$
- Neon: $1s^2 2s^2 2p^6$
Hund’s rule (rule for equal-energy orbitals): electrons fill singly in degenerate orbitals before pairing.
- Example: in the 2p subshell, electrons occupy separate orbitals before pairing.
Box/orbital notation vs line notation:
- Box notation (orbital diagram) shows boxes for orbitals and arrows for electrons.
- Line notation uses the same subshell labels with superscripts, e.g., $1s^2\ 2s^2\ 2p^6$ .
Practice check:
- 2p orbital holds $2$ electrons total (in two degenerate orbitals you place one in each before pairing).
- 3d subshell holds $10$ electrons (5 orbitals × 2).

Abbreviate inner-shell configurations using the nearest noble gas from the previous period in brackets, then add the remaining electrons.
Examples:
- Carbon: $\text{C} = [\text{He}]\ 2s^2\ 2p^2$
- Iodine (example of a large element): $[\text{Xe}]\ 4f^{14}\ 5d^{10}\ 6s^2\ 6p^5$
Noble gases include: He, Ne, Ar, Kr, Xe, Rn, Og.
Three ways to write configurations:
- Box notation, e.g., two arrows per orbital
- Full line notation, e.g., $1s^2\ 2s^2\ 2p^6$
- Noble gas notation, e.g., $[\text{Ne}]\ 3s^2\ 3p^5$

Two electrons occupy any orbital (paired with opposite spins).
2p has 2 electrons max per orbital; with 3 degenerate p orbitals, total 6 electrons in 2p.
3d subshell has 5 orbitals; max 10 electrons.
The periodic table is a visual guide to electron-shell filling, not just memorization.
When given a configuration, you can infer the element by summing the superscripts (total electrons = atomic number for a neutral atom).