Kössel-Lewis Approach

Why does sodium hand over an electron to chlorine, while two chlorine atoms simply share a pair instead? In 1916, working independently, Walther Kössel and Gilbert N. Lewis answered this with the earliest electronic picture of bonding: atoms combine because each one is chasing the stable, noble-gas-like arrangement of eight electrons in its outer shell — the octet rule. Kössel pictured this stability reached by outright electron transfer, producing ions held together by electrostatic attraction (an ionic bond); Lewis pictured it reached by two atoms sharing a pair of electrons between them (a covalent bond). Between them they gave chemistry its first working vocabulary for a bond: dot symbols for valence electrons, dot-and-dash structures for whole molecules, and formal charge, a bookkeeping tool to judge which of several candidate structures is most believable. This article builds that vocabulary, shows where the octet rule breaks down, and works a full Lewis structure — the resonance-bearing nitrate ion — so you can draw and defend a structure under exam conditions.

The core insight

Atoms attain a stable octet of valence electrons either by transferring them (ionic bond, Kössel) or by sharing them in pairs (covalent bond, Lewis), and a Lewis structure's plausibility is checked using formal charge, FC=V−N−12B\text{FC} = V - N - \tfrac{1}{2}B.

The core idea

The octet rule. Noble gases (He, Ne, Ar, ...) are chemically inert because their outermost shell is completely filled — He with 2 electrons, the rest with 8. Kössel and Lewis proposed that every other atom "wants" this same filled-shell stability, and reacts by either giving away, taking, or sharing electrons until its valence shell looks like the nearest noble gas's. This single idea is the seed from which ionic bonding, covalent bonding, and eventually molecular shape all grow.

Lewis dot symbols. A representative (main-group) element's valence electrons are shown as dots placed around its symbol, one dot per electron, up to four sides before pairing begins. The dot count equals the group number for groups 1, 2, and 13-18 (in the older 1-8 notation): sodium has 1 dot, oxygen has 6, chlorine has 7. This dot symbol is the atom's "bonding budget" — it tells you at a glance how many electrons it can lose, gain, or share.

Lewis structures. For a whole molecule or ion, every atom's symbol is drawn together with its bonding pairs (shared electrons between two nuclei, usually drawn as a short line, one line per pair) and its lone pairs (electrons that belong to just one atom, drawn as dots). A correctly drawn Lewis structure lets you count electrons directly around each atom and check its octet (or duet, for hydrogen) is satisfied.

Formal charge. When more than one Lewis structure can be drawn for the same molecule, formal charge is the tool used to decide which is the best representation of reality. It compares how many electrons an atom "owns" in the structure against how many it would own as a free, neutral atom:

FC=V−N−12B\text{FC} = V - N - \frac{1}{2}B

where: VV is the number of valence electrons the free atom would normally have, NN is the number of non-bonding (lone-pair) electrons it carries in the structure, and BB is the number of bonding (shared) electrons around it, so that 12B\tfrac{1}{2}B is the electrons the atom is credited with from each shared pair. The sum of formal charges over every atom in the structure must always equal the overall charge of the species — this is a built-in arithmetic check, not a coincidence. Among competing structures, the one with formal charges closest to zero, and with any negative charge sitting on the more electronegative atom, is preferred.

Figure
Two Lewis structures side by side showing sodium losing one electron to chlorine to form an ionic bond and two chlorine atoms sharing a pair of electrons to form a covalent bond
Two routes to an octet — transfer versus sharing ionic (Kössel): NaCl Na Cl e− Na⁺ Cl⁻ electron transferred, then electrostatic attraction covalent (Lewis): Cl2 Cl Cl one shared pair, one bonding line each side
Derivation AFormal charge from the electron-counting rule

The goal is to derive the formal-charge formula from a single, plain electron-counting rule so that its origin is obvious rather than memorised: an atom in a Lewis structure is credited with all of its own lone-pair electrons in full, and with exactly half of every bonding pair it takes part in, because a shared pair is by definition owned equally by the two bonded atoms.

Step 1. Define the atom's "assigned" electron count in the structure. If an atom has NN non-bonding (lone-pair) electrons drawn directly on it, and it takes part in bonds carrying a total of BB bonding electrons around it, then by the equal-sharing rule its assigned share of those bonding electrons is 12B\tfrac{1}{2}B. Its total assigned electron count is therefore N+12BN + \tfrac{1}{2}B.

Step 2. Compare this assigned count to what the atom would have as a free, neutral, isolated atom. That free-atom count is simply its number of valence electrons, VV, obtainable from its group number. Any difference between what it "should" have as a free atom and what it is "assigned" in the structure is charge that has effectively moved onto or off that atom in the bonding process.

Step 3. Write the difference as (free-atom electrons) minus (assigned electrons in the structure): FC=V−(N+12B)=V−N−12B\text{FC} = V - \left(N + \tfrac{1}{2}B\right) = V - N - \tfrac{1}{2}B. A positive result means the atom is assigned fewer electrons than it would normally have (it looks like it lost some), a negative result means it is assigned more (it looks like it gained some), and zero means the structure assigns it exactly its free-atom count.

Step 4. Note the built-in consistency check: since every electron in the structure is either a lone-pair electron on one atom or a bonding electron shared between exactly two atoms, summing V−N−12BV - N - \tfrac{1}{2}B over every atom must reproduce the species' overall charge. This is why formal charges of a correctly drawn structure always add up to the actual charge — a direct consequence of the formula, not an extra rule to memorise.

Worked example

Draw the Lewis structure of the nitrate ion, NO3−\text{NO}_3^-, and verify its formal charges.

Step 1. Count total valence electrons: nitrogen contributes 5, three oxygens contribute 3×6=183\times6=18, and the ion's −1-1 charge contributes 1 more electron. Total =24=24 electrons =12=12 pairs.

Step 2. Place N centrally (nitrogen forms the most bonds among these atoms), single-bonded to each of the three O atoms: 3 bonds use 6 electrons (3 pairs), leaving 24−6=1824-6=18 electrons (9 pairs) to distribute.

Step 3. Give each terminal O three lone pairs, which uses exactly the remaining 9 pairs. At this point N has only 3 bonds and no lone pair of its own — one electron short of an octet, so this structure cannot be final.

Step 4. Fix the shortfall by converting one N–O single bond into a double bond: move one lone pair from that oxygen into a second N–O bonding pair. N now has 4 bonds (octet satisfied); one oxygen is doubly bonded with 2 lone pairs, the other two oxygens remain singly bonded with 3 lone pairs each.

Step 5. Compute formal charges with FC=V−N−12B\text{FC}=V-N-\tfrac12B. For N (4 bonds, 0 lone pairs): 5−0−12(8)=+15-0-\tfrac12(8)=+1. For the doubly bonded O (2 bonds, 2 lone pairs): 6−4−12(4)=06-4-\tfrac12(4)=0. For each singly bonded O (1 bond, 3 lone pairs): 6−6−12(2)=−16-6-\tfrac12(2)=-1, and there are two such oxygens. Sum: (+1)+0+(−1)+(−1)=−1(+1)+0+(-1)+(-1)=-1, which correctly matches the ion's overall −1-1 charge. The Lewis structure of NO3−\text{NO}_3^- has N with formal charge +1+1, one O at 00, and two O atoms at −1-1 each — since the double bond could equally sit on any of the three oxygens, NO3−\text{NO}_3^- is a resonance hybrid of three equivalent structures.

Competitive stretch

Applying the same recipe to carbon dioxide (4+2×6=164+2\times6=16 electrons, 8 pairs) forces two C=O double bonds rather than one, because a single-bonded skeleton leaves carbon with only 4 electrons around it (nowhere near an octet) while oxygen already has a full octet with just one bond and three lone pairs. The final structure, O=C=O\text{O}=\text{C}=\text{O}, gives every atom a formal charge of exactly zero — a useful cross-check that a symmetric, neutral molecule should carry no formal charge at all when drawn correctly.

Common traps

  • Forgetting to add or remove electrons for the charge on a polyatomic ion before drawing the skeleton — the single most common arithmetic slip in this topic.
  • Treating formal charge as the actual charge on an atom; it is a bookkeeping estimate for comparing candidate structures, not a measured or "real" charge.
  • Assuming every atom needs exactly eight electrons — hydrogen needs only 2 (a duet), while B and Be are stable with fewer than 8, and P, S can exceed 8 in period-3 compounds.
  • Placing the wrong atom at the centre — the central atom is usually the one that can form the most bonds (often the least electronegative non-hydrogen atom in the formula), not necessarily the atom written first.

Ask the tutor: Not sure why nitrate needs a double bond, or want your own formal-charge arithmetic checked step by step — ask the AI tutor to walk through your Lewis structure.

Why does nitrogen need a double bond in nitrate but not in ammonia?

It comes down to how many electron pairs are available once the skeleton is drawn. In NH3\text{NH}_3, nitrogen's three single bonds to hydrogen plus one lone pair already use up all 8 electrons around it — no shortfall, so no double bond is needed. In NO3−\text{NO}_3^-, after placing three N–O single bonds and filling every oxygen's octet with lone pairs, nitrogen is left with only 3 bonds and 0 lone pairs, one electron pair short of an octet. The only way to fix that shortfall without changing the total electron count is to convert one single bond into a double bond, exactly what Step 4 of the worked example does.

If formal charge is not the real charge, why do we bother calculating it?

Formal charge is a comparison tool, not a measurement. A real molecule can often be drawn with two or more different, equally electron-counted Lewis structures, and formal charge gives a consistent rule for picking the most chemically reasonable one: prefer structures where formal charges are closest to zero, and where any negative formal charge lands on the more electronegative atom. It is analogous to a spreadsheet check — it tells you which drawing is more believable, while the true distribution of charge in the molecule is smoother and is better described later by resonance and, eventually, quantum-mechanical bonding theories.

Does the octet rule mean atoms literally count their electrons and stop at eight?

No — it is a useful bookkeeping pattern, not a law of physics that atoms obey deliberately. The real reason a filled outer shell is stable is quantum mechanical: electrons fill shells and subshells in an order governed by their energies, and a completely filled valence shell corresponds to an especially low-energy, hard-to-disturb arrangement. The octet rule is simply the observation that, for most second-period elements, that low-energy arrangement happens to hold eight electrons. That is also exactly why the rule breaks for elements where this coincidence does not hold, such as boron (stable with six) or sulfur (able to exceed eight using its available d-orbitals).

Before this

Where this leads