Before you can read a single Lewis structure, you need a small toolbox of ideas. The parent note quietly assumes every one of them. Below we build each from absolute zero, in the order they depend on each other — no symbol appears before it is earned.
Look at the figure. The grey dot in the middle is the nucleus. The blue rings are shells. The first ring (closest to the nucleus) can hold only 2 electrons. The second ring can hold 8. Electrons fill inner rings first, then spill outward — like filling the bottom row of seats before the next row.
Why do we care about rings at all? Because chemistry is decided almost entirely by the electrons in the outermost occupied ring. The inner ones are buried and unavailable.
In the figure above, the electrons on the outer blue ring (drawn orange) are the valence electrons. The buried inner electrons (grey) never take part.
Why single these out? When two atoms meet, only their outer surfaces touch. Inner electrons are shielded by distance and by the outer ones. So bonding = a story about valence electrons only. This is why the parent note counts "group-valence electrons" and nothing else.
Reading the table: for groups 13–18, subtract 10 from the group number to get valence electrons (14 → 4, 16 → 6, 17 → 7). This is exactly why the parent writes "C is group 14 → 4 valence; each O is group 16 → 6."
See Periodic trends for why the table is shaped this way.
The figure shows why "8" is special. The valence shell is built from orbitals: one s orbital (holds 2 electrons) plus three p orbitals (hold 6 electrons). Fill them all:
2+6=8.
Why the superscripts?s2 means "the s orbital contains 2 electrons"; p6 means "the three p orbitals contain 6 electrons total." The little raised number counts electrons in that type of room. This is the notation behind the parent's phrase "s2p6."
Why 2 for hydrogen and helium? They only have the first shell, which has just the tiny s room — one room, 2 seats. So their "full" is 2, called a duet, not 8.
Look at the figure. Around each symbol are up to 8 positions (top, bottom, left, right — two seats each). We place dots one per side first, then pair them up. Oxygen (6 valence) ends up with 2 lone pairs and 2 single dots; nitrogen (5) with 1 lone pair and 3 single dots.
Why this matters for counting later: the formal-charge formula treats lone-pair electrons and bonding electrons differently, so you must be able to tell them apart on sight.
The figure shows two hydrogen atoms. Each has 1 electron and wants 2. Alone, each is short. Sharing their two electrons in the middle, both now count 2 — both reach the duet. That shared pair is a single bond, drawn as one line: H−H.
Single bond = 1 shared pair = 2 e⁻ = one line.
Double bond = 2 shared pairs = 4 e⁻ = two lines (=).
Triple bond = 3 shared pairs = 6 e⁻ = three lines (≡).
Why the parent divides shared electrons by 2: each bond is a pair, so bonds=Nshared/2.
This is exactly the "charge correction" in the parent's Step A. Example: CO32− has a 2− charge, so we add 2 electrons to the count. Get this backwards and every later step (bonds, formal charge) goes wrong.
The parent uses FC=V−L−21B without slowing down. Here is every letter.
Why halve B? A bonding pair is shared, so each atom is credited only half of it — that is the 21B. Lone pairs belong entirely to the atom, so they are subtracted in full.
Worked micro-example (single-bonded O in CO32−):V=6, it has 3 lone pairs so L=6, one single bond so B=2.
FC=6−6−21(2)=6−6−1=−1.
That negative charge sits on oxygen (very electronegative) — perfectly reasonable.
NO has 5+6=11 electrons — odd — so one electron is always unpaired. You literally cannot pair everything, so a perfect octet everywhere is impossible. Recognising odd totals is a prerequisite for the odd-electron exception.
