Valence Electrons

Chemistry is almost entirely the behaviour of an atom's outermost electrons. Iron has twenty-six electrons and eighteen of them might as well not exist for most purposes: they are locked into filled inner shells, held so tightly that no ordinary reaction disturbs them.

The ones that are left over — the valence electrons — determine what an element bonds to, how many bonds it makes, what charge its ions carry and where it sits in the periodic table. Counting them correctly is the prerequisite for nearly everything else, and for main-group elements it takes about two seconds once you know the rule.

Why the inner electrons are out of play

Two things put core electrons beyond reach. They sit closer to the nucleus, so they feel a much stronger attraction, and they are shielded from disturbance by the electrons outside them.

The evidence is blunt and quantitative. Removing sodium's single outer electron costs 496 kJ per mole. Removing the next one — the first from the filled shell beneath — costs 4,562 kJ, more than nine times as much. No chemical reaction releases energy on that scale, so sodium's second electron is simply unavailable, and sodium is a +1 metal in every compound it forms. That enormous step, and where it falls for each element, is the subject of ionisation energy.

The definition follows: valence electrons are those in the outermost occupied shell — the one with the highest principal quantum number. Everything beneath is the core, usually written as the preceding noble gas in square brackets.

Group number, minus ten

For main-group elements the count is readable straight off the table.

• Groups 1 and 2: the count is the group number, so hydrogen and sodium have one, magnesium and calcium two.
• Groups 13 to 18: subtract ten. Group 13 has three, group 15 five, group 17 seven, group 18 eight.

Phosphorus is in group 15, so five, and its configuration confirms it: [Ne] 3s² 3p³, with 2 + 3 = 5 electrons in shell 3. Chlorine is group 17: [Ne] 3s² 3p⁵, so seven.

If your course uses the older numbering with Roman numerals, the A-group number is the valence electron count directly, which is the one genuine advantage that system retains.

Two elements need care. Helium has two valence electrons, not eight, yet it belongs with the noble gases because two is a full shell for shell 1. Hydrogen has one valence electron and sits above the alkali metals without being anything like them.

Dots around a symbol

A Lewis dot symbol writes the element's symbol with one dot per valence electron, distributed around four imaginary sides. Dots go singly onto each side before any side takes a second — the same preference for staying unpaired that governs how orbitals fill.

• Carbon, four valence electrons: four single dots, one per side.
• Oxygen, six: two sides paired, two sides with single dots. Those two singles are exactly the two bonds oxygen typically forms.
• Neon, eight: all four sides paired, nothing available, no ordinary chemistry.

The notation earns its keep because the number of unpaired dots usually predicts the number of bonds. Nitrogen's five electrons give one pair and three singles, and nitrogen forms three bonds in ammonia while keeping the pair to itself as a lone pair.

Eight is just two plus six

The octet rule says main-group atoms tend toward eight valence electrons. Stated like that it sounds like numerology, and it is nothing of the sort.

For any shell, the accessible orbitals in ordinary chemistry are one s orbital and three p orbitals. The s holds two electrons and the three p orbitals hold six. Two plus six is eight, and an atom with eight valence electrons has filled every orbital available at that energy — the configuration s²p⁶ that every noble gas from neon onwards has.

Hydrogen, helium, lithium and beryllium play by a duet rule instead, because shell 1 has no p orbitals at all and two electrons already fill it.

The octet is a strong tendency, not a law, and a substantial minority of stable molecules violate it — boron trifluoride settles for six, sulfur hexafluoride carries twelve, and nitrogen monoxide has an odd number of electrons and so cannot possibly reach eight. Those cases are grouped and explained on the covalent bonding page, where the Lewis structures that expose them are drawn.

Iron has two valence electrons. Iron has eight.

Both statements appear in reputable textbooks, and the contradiction is real rather than an error, because two different definitions are in circulation and neither is wrong.

Iron's configuration is [Ar] 3d⁶ 4s².

• By the outermost-shell definition, only the 4s electrons qualify, since shell 4 is the highest occupied. Iron has two valence electrons.
• By the available-for-bonding definition, which counts the outer shell plus any incomplete inner subshell, the six 3d electrons count too. Iron has eight.

The second definition is the more useful one for the d-block, because it predicts what the elements actually do. Manganese has seven electrons outside its argon core, and manganese reaches +7 in permanganate. Chromium has six, and chromium reaches +6 in chromate. Osmium and ruthenium have eight and both form tetroxides in which they are +8. Scandium has three, forms Sc³⁺, and stops there because there is nothing left.

The practical advice is to state which definition you are using whenever the answer matters. The rule that reads a count off the group number applies to groups 1, 2 and 13–18 only; groups 3–12 do not obey it, and applying it to iron would suggest eight from "group 8 minus nothing", which gets the right answer for the wrong reason and fails immediately for copper.

A related generalisation exists for transition metals: the 18-electron rule, which holds that many stable complexes have eighteen electrons around the metal. Eighteen is 2 + 6 + 10 — a filled s, p and d set — so it is the d-block's version of the octet, and it is honoured about as reliably, which is to say usefully but not universally.

Three words that are not synonyms

Three terms overlap in ordinary speech and mean genuinely different things.

• Valence electrons is a count of electrons. Sulfur has six.
• Valency, or combining capacity, is the number of bonds an atom forms. Sulfur's is usually two, as in hydrogen sulfide, though it can be four or six.
• Oxidation state is a signed bookkeeping number that assumes bonds are fully ionic. Sulfur ranges from −2 to +6 across its compounds.

Sulfur therefore has six valence electrons, a valency of two in hydrogen sulfide, and an oxidation state of −2 in the same molecule. All three numbers describe the same atom in the same compound, and confusing them produces answers that are wrong in ways that are hard to spot. The signed one has its own conventions, set out under oxidation states.

Counting for a whole ion, not just an atom

Before drawing any structure you need the total valence electron count for the species, and the charge changes it.

Add up the valence electrons of every atom, then add one for each unit of negative charge and subtract one for each unit of positive charge. The sign is the opposite of what the charge looks like, because a negative ion has gained electrons.

• Carbonate, CO₃²⁻: carbon contributes 4, each oxygen 6, and the 2− charge adds 2. 4 + 18 + 2 = 24 electrons.
• Sulfate, SO₄²⁻: 6 + 24 + 2 = 32.
• Ammonium, NH₄⁺: 5 + 4 = 9, minus 1 for the positive charge = 8.

The last of those is worth pausing on. Ammonium's eight electrons form four bonds around nitrogen and nothing else — no lone pairs left — which is why the ion is a neat tetrahedron and why its charge is best thought of as belonging to the whole ion rather than sitting on any one atom.