A Lewis structure is the working drawing of organic chemistry. It shows three things: which atoms are joined to which, where every valence electron sits, and which atoms carry a charge. Every valence electron appears either in a bond (a pair shared by two atoms) or in a lone pair (a pair kept by one atom). From those three things you can predict a molecule's shape, polarity, acidity and reactivity. Most mistakes in a mechanism trace back to a Lewis structure drawn carelessly at the start.
A Lewis structure is flat. It records connections and electrons, not the 3D shape; Molecular geometry reads the shape off a finished structure. Every structure on this page ends with the formula from Formal charge: FC = valence electrons − lone-pair electrons − bonds.
The method you will actually use: the valence rule
Take ethanol, CH₃CH₂OH. Its heavy atoms (every atom except hydrogen) form a chain: C, C, O. None of them sits in the middle of the others. General chemistry teaches a recipe that starts by picking a central atom, and a chain has none. So organic chemists use a different method, one built for chains.
Look at what each atom in ethanol does. Each carbon makes four bonds and keeps no lone pairs. The oxygen makes two bonds and keeps two lone pairs. Each hydrogen makes one bond. These numbers are not special to ethanol: every neutral carbon, nitrogen, oxygen and halogen has a fixed number of bonds and lone pairs. Call that pair of numbers the atom's quota.
The table lists the quotas. It is the octet counting from Atomic structure, written out atom by atom: the shared pairs plus the kept pairs fill the outer shell.
| Atom | Valence e⁻ | Bonds | Lone pairs | Why |
|---|---|---|---|---|
| H | 1 | 1 | 0 | 1 shared pair fills its first shell (2) |
| C | 4 | 4 | 0 | 4 shared pairs = 8 electrons |
| N | 5 | 3 | 1 | 3 shared + 1 kept = 8 |
| O | 6 | 2 | 2 | 2 shared + 2 kept = 8 |
| F, Cl, Br, I | 7 | 1 | 3 | 1 shared + 3 kept = 8 |
Skeleton. Join the heavy atoms in the order the formula writes them: C–C–O. Put each hydrogen on the atom written just before it: three on the first carbon, two on the second, one on the oxygen.
Bonds. Count each atom's bonds against its quota. Carbon 1: three H plus carbon 2 makes 4. ✓ Carbon 2: two H, one C and one O makes 4. ✓ Oxygen: one C and one H makes 2. ✓ Each hydrogen: 1. ✓ Single bonds meet every quota, so no double bond is needed.
Lone pairs. Only the oxygen's quota includes lone pairs. Two bonds means two lone pairs. The carbons get none. Draw every lone pair, in every structure: leaving them off is how charges get lost later.
Charges. Every atom is at its quota, so every formal charge is zero. The molecule is neutral. ✓ As a last check, count electrons: the drawing holds 8 bonds and 2 lone pairs, which is 20 electrons, and the atoms bring 2(4) + 6(1) + 6 = 20 valence electrons. ✓
The ethanol example used four steps. They work for any organic molecule:
- Skeleton. Join the heavy atoms in the order the formula gives, and hang each hydrogen on the atom it follows.
- Bonds. Give every atom its quota of bonds. Where single bonds leave two neighboring atoms short, add a double or triple bond between them.
- Lone pairs. Add lone pairs until every atom reaches its quota.
- Charges. Work out each formal charge. An atom that departs from its quota carries a charge; write it on that atom. Then check that the electrons drawn match the valence electrons the atoms bring.
Reading a condensed formula
CH₃CH₂OH is a condensed formula: it lists the atoms in bonding order and leaves most bonds out. Until skeletal drawings arrive in Drawing Molecules & Moving Electrons, it is the only compact notation you have. Three conventions cover nearly all of it.
Hydrogens follow the atom they are on. In CH₃CH₂OH, the first carbon carries three hydrogens, the second carries two, and the O carries one. Read left to right, each heavy atom bonds to the heavy atom before it.
Parentheses mean branches. The subscript outside a bracket counts copies of the whole group. At the start of a formula, the bracketed group hangs off the atom that follows it: in (CH₃)₂CHOH, one CH carbon carries two methyl (CH₃) groups and an OH, and in (CH₃)₃COH three methyls hang off a carbon that has no hydrogen left.
Anywhere else, a bracket hangs its group on the atom before it. CH₃CH(OH)CH₃ puts the OH on the middle carbon, so it is the same molecule as (CH₃)₂CHOH. C(CH₃)₄ is one carbon carrying four methyls, with nothing after the bracket at all.
Multiple bonds are sometimes written. Double and triple bonds between carbons, and to a nitrogen, are often spelled out: CH₂=CH₂, CH₃C≡N. Everything else joins by single bonds, with one exception below.
One pattern needs care: a carbon followed by two O's, written COO, as in CH₃COOH or CH₃COOCH₃. The two O's are not a chain of oxygens. The first O takes a double bond to the carbon, and the second O continues the chain. A CHO at the end of a formula, as in CH₃CHO, also hides a C=O: the O ends the chain, so it bonds only to that carbon. A single O that continues the chain, as in (CH₃)₃COH, is an ordinary single bond. When in doubt, count bonds: each carbon must reach four. Functional groups uses these patterns most.
When an atom leaves its quota: charges
Skeleton. C–N, with three H on the carbon and two on the nitrogen.
Bonds. Carbon: 3 H + 1 N = 4. ✓ Nitrogen: 1 C + 2 H = 3. ✓
Lone pairs. Nitrogen's quota is 3 bonds and 1 lone pair, so it gets one lone pair.
Charges. All zero.
Why the lone pair matters. It drives the chemistry of amines. It makes methylamine a base, because an H⁺ can bond to it. It makes methylamine a nucleophile, an electron-pair donor that bonds to an electron-poor atom. It also sets the shape: the three bonds point down from nitrogen like the legs of a low pyramid, and the lone pair sits on top.
Now add H⁺. The H⁺ bonds to that lone pair. Nitrogen now has 4 bonds and no lone pair, so FC = 5 − 0 − 4 = +1. Methylammonium, CH₃NH₃⁺, carries its + on the nitrogen. Nothing else in the drawing changes.
Methylammonium shows the general rule. An atom at its quota is neutral. An atom off its quota carries a charge, and the direction is fixed.
- A bond in place of a lone pair gives +1. Oxygen with 3 bonds and 1 lone pair; nitrogen with 4 bonds and none.
- A lone pair in place of a bond gives −1. Oxygen with 1 bond and 3 lone pairs; carbon with 3 bonds and 1 lone pair.
- A missing bond with no lone pair to replace it gives +1. Carbon with 3 bonds and nothing else is a carbocation.
Skeleton. CH₃–C–O–H, with the other O on the middle carbon. By the condensed-formula rule, the two O's are not a chain.
Bonds. The methyl carbon: 3 H + 1 C = 4. ✓ With single bonds only, the middle carbon has 3 bonds (CH₃, O, OH), one short. The O without an H has 1 bond, also one short. A double bond between them fixes both: the middle carbon reaches 4 ✓ and that oxygen reaches 2 ✓.
Lone pairs. The OH oxygen has 2 bonds, so 2 lone pairs. The C=O oxygen also has 2 bonds (a double bond counts twice), so 2 lone pairs.
Charges. Every atom is at its quota; all zero.
Now remove H⁺. The H leaves without electrons, so the O–H bond pair stays on the oxygen. That oxygen now has 1 bond and 3 lone pairs: FC = 6 − 6 − 1 = −1. The ion is acetate, CH₃COO⁻. ✓
Skeleton. C–O, with three H on the carbon and none on the oxygen.
Bonds. Carbon: 3 H + 1 O = 4. ✓ Oxygen: 1 bond, one short of its quota of 2.
Lone pairs and charge. One bond short means one extra lone pair and a −1, so the oxygen gets 3 lone pairs and a −1. Check: 6 − 6 − 1 = −1. ✓ The sum over the ion is −1, matching its charge.
Where to write it. The charge goes on the oxygen in the drawing, not floating beside the formula.
The general-chemistry recipe, for molecules with a real center
Some small molecules do have one atom in the middle. In CO₂, the carbon sits between two oxygens. In nitrate, NO₃⁻, the nitrogen sits among three oxygens. For these, the general-chemistry recipe works well. It aims at the octet rule from Atomic structure: most atoms are most stable with 8 electrons in their outer shell, and hydrogen with 2.
Count. C (4) + O (6) + O (6) = 16 valence electrons.
Connect. O–C–O with two single bonds uses 4, leaving 12.
Fill the outer atoms. Three lone pairs on each oxygen use all 12, leaving 0.
Check the center. Carbon has only 2 bonds: 4 electrons, four short of an octet.
Fix it. Move one lone pair from each oxygen into a second bond, giving O=C=O. Carbon now has 4 bonds (8 electrons). Each oxygen has 2 bonds plus 2 lone pairs (8 electrons).
Charges. Carbon: 4 − 0 − 4 = 0. Each oxygen: 6 − 4 − 2 = 0. The sum is 0, as a neutral molecule needs. ✓
The CO₂ example is the whole recipe. In general form:
- Count total valence electrons. Add up each atom's group-number valence count. Then adjust for charge: add one electron per negative charge, subtract one per positive charge.
- Connect the atoms with single bonds. The least electronegative atom (never hydrogen) goes in the middle. Each bond uses 2 of your electrons.
- Give the outer atoms full octets with lone pairs, starting with the most electronegative.
- Put any electrons still left over on the central atom.
- If the central atom is still short of an octet, turn a lone pair on a neighbor into a second shared pair: a double bond. This gives the central atom more electrons without changing the total, which is why step 1 must be right.
Then work out the formal charges and check that they add up to the overall charge. Formal charge explains why that sum is a real check.
Count. N (5) + 3 × O (6) = 23, plus 1 for the negative charge = 24.
Connect. Three N–O single bonds use 6, leaving 18.
Fill the outer atoms. 18 is exactly three lone pairs for each oxygen. All 24 are placed.
Check the center. Nitrogen has 3 bonds: 6 electrons, two short.
Fix it. Move one lone pair from one oxygen into a double bond. Nitrogen now has 4 bonds: FC = 5 − 0 − 4 = +1. The doubly bonded oxygen: 6 − 4 − 2 = 0. Each singly bonded oxygen: 6 − 6 − 1 = −1. Sum: +1 + 0 − 1 − 1 = −1. ✓
Which oxygen? Any of the three. The choice is arbitrary, so three equally good structures exist. In real nitrate, all three N–O bonds are measurably identical. That is not a defect in the method. It is the molecule telling you it has resonance, the subject of Drawing Molecules & Moving Electrons.
Using formal charge as the tiebreaker
CO₂ has a second structure that also gives every atom an octet: one single bond and one triple bond, O–C≡O. The single-bonded oxygen keeps 3 lone pairs, so its FC is 6 − 6 − 1 = −1. The triple-bonded oxygen keeps 1 lone pair, so its FC is 6 − 2 − 3 = +1. Same atoms, same octets, but two charges where O=C=O has none.
That is the tiebreaker. When more than one structure satisfies the octet rule, the realistic one keeps formal charges closest to zero. If a charge cannot be avoided, the negative charge goes on the most electronegative atom. Formal charge sets out the full ranking.
Three real exceptions to the octet rule
Incomplete octets. Boron brings only 3 valence electrons. In BF₃ it makes three bonds and has no lone pair: 6 electrons. That is the correct structure, not an error. Boron and beryllium routinely stop short of eight.
Boron's empty 2p orbital makes BF₃ a powerful electrophile: it forms bonds by accepting an electron pair. Acids & Bases calls an electron-pair acceptor a Lewis acid, and BF₃ is a strong one. Carbocations belong in the same family.
Expanded octets: why period 2 cannot. This is the exception the rest of the chapter points back to. The n = 2 shell contains exactly four orbitals, one 2s and three 2p, and four orbitals hold eight electrons. A ninth electron on carbon, nitrogen, oxygen or fluorine would have nowhere to sit. A period-2 atom is also small, so it has no room to crowd more than four neighbors around it. This is why "I gave carbon five bonds" is the most common structural error in a first organic course.
So the safe statement is this: carbon, nitrogen, oxygen and fluorine never hold more than eight electrons. With the ordinary two-electron bonds this course draws, carbon has at most four bonds, and so does nitrogen. A few exotic ions, such as CH₅⁺, do attach five atoms to one carbon. They share one electron pair among three atoms, so the carbon still holds eight. You will not draw them in this course.
Expanded octets: period 3 and below can. PCl₅ puts 10 electrons on phosphorus, and SF₆ puts 12 on sulfur. Both are real, stable compounds.
The 10 and 12 are counts from the drawing. Phosphorus and sulfur cannot use extra orbitals to hold them either: their empty 3d orbitals are too high in energy to take much part. Their bonds have strong ionic character: the bonding electrons sit mostly on the more electronegative atom, part-way to an ionic bond. So much of that electron density really sits on the chlorines and fluorines. Period-3 atoms are also larger, so five or six neighbors fit around them without crowding.
Older textbooks explained this with the empty 3d orbitals: they were said to mix into the bonding, which is where the labels "sp³d" and "sp³d²" came from. Modern computation has retired that explanation, because real d-orbital participation turns out to be minimal. The shapes and the practical rule are unaffected: sulfur and phosphorus can be drawn past an octet.
That ionic character can change the drawing itself. Sulfuric acid, H₂SO₄, is usually drawn with two S=O double bonds: every formal charge is zero, and sulfur holds 12 electrons. It can also be drawn with four single bonds: sulfur +2, two oxygens −1, and every atom at an octet. Bonding calculations favor the second picture. Most courses and exams still use the first, and so does this course, but you should recognize both.
The same choice comes up for two sulfur groups in organic molecules. A sulfoxide has a sulfur bonded to two carbons and one oxygen; a sulfone has a sulfur bonded to two carbons and two oxygens. Both are usually drawn with S=O bonds, which put 10 and 12 electrons on sulfur. Drawn with S⁺–O⁻ bonds instead, sulfur keeps an octet.
Odd-electron species. Nitric oxide, NO, has 5 + 6 = 11 valence electrons. An odd total cannot be split into pairs, so one electron stays unpaired and some atom falls short of an octet. Any species with an unpaired electron is a free radical; the methyl radical, CH₃•, is the organic example. Radicals are real and reactive, and they get their own treatment when radical mechanisms come up.
Four structures that are wrong
Each of these is drawn often. The valence rule catches each one in seconds: compare what is drawn with the atom's quota.
- A carbon with five bonds, or four bonds and a lone pair. Either way, that is ten electrons on a period-2 atom. The legal options for carbon in this course are four bonds (neutral), three bonds and a lone pair (a carbanion, a carbon with a negative charge: −1), three bonds and nothing else (a carbocation, +1), and three bonds with one unpaired electron (a radical).
- A nitrogen with four bonds and no charge. The +1 from the charge rule above is missing. A four-bonded nitrogen always carries a +: in ammonium, in a protonated amine, and in a nitro group (–NO₂, a nitrogen bonded to two oxygens).
- An oxygen with three bonds and no charge. Again the +1 is missing: draw it or redraw the bond. This slip is common when an oxygen lone pair forms a new bond and the charge is forgotten.
- A hydrogen with two bonds. Hydrogen's only shell holds 2 electrons, which is one shared pair. An ordinary organic structure never has an H bridging two atoms; an H drawn between two atoms is a slip of the pen.
What carries forward
Everything. Molecular geometry is read off a completed Lewis structure. Resonance is the observation that some molecules need more than one. Arrow pushing is the practice of converting one Lewis structure into another legally, and every mechanism in the course is a sequence of Lewis structures. If one section of Foundations is worth over-practicing, it is this one.