Chemistry 2e · Chemical Bonding and Molecular Geometry
Lewis Symbols and Structures
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Before discussing a molecule's shape, polarity, or reactivity, chemists need a picture of how its atoms connect. Lewis symbols show an atom's valence electrons as dots around its symbol; Lewis structures extend that idea to whole molecules, with bonds as lines and lone pairs as dots. Writing a Lewis structure Drawing of a molecule showing bonds (lines) and lone pairs (dots) Full entry → is a systematic counting exercise: every electron is accounted for, and the final drawing must use exactly the available valence electrons. This topic covers that procedure, the octet rule Tendency of atoms to surround themselves with eight valence electrons Full entry → that usually governs it, and the exceptions (hydrogen's duet, boron's sextet, expanded octets, odd-electron radicals).
Why this matters
The Lewis structure is the single most reused tool in introductory chemistry. Every later topic in this chapter — resonance, formal charges, molecular geometry, polarity — starts from a correct Lewis structure, as do Chapter 8's bonding theories. Lewis structures also explain real behavior — why water is bent, why CO₂ is linear, why ammonia is a base — and counting right now prevents later errors.
The college version
Core Concepts
Lewis symbols for atoms
A Lewis symbol Element symbol with dots for valence electrons Full entry → is the element's symbol surrounded by dots, one per valence electron Electron in the outermost shell, available for bonding Full entry →. For main-group elements, that number equals the group number: sodium (group 1) has one dot, chlorine (group 17) seven. Dots are placed one per side first, then paired, because electrons in the same orbital repel; phosphorus (group 15) shows one dot on each of four sides plus a fifth pairing with one. The unpaired dots are the electrons available for bonding.
The octet rule and its exceptions
The octet rule states that atoms tend to gain, lose, or share electrons until they have eight valence electrons — a noble-gas configuration. Hydrogen is the famous exception, satisfied with two (a duet, like helium). Boron often stops at six, forming compounds like BF₃ with only a sextet. Elements in period 3 and beyond (P, S, Cl) can exceed eight using empty d orbitals, giving expanded octets such as PCl₅ and SF₆. A few species, like NO and NO₂, have an odd total of electrons and must leave one unpaired; these are radicals.
A reliable procedure for writing structures
For any neutral molecule or ion:
- Count valence electrons. Sum all atoms' valence electrons; for an ion, add one per negative charge and subtract one per positive charge.
- Choose the skeleton. Least electronegative atom central (never hydrogen), atoms connected by single bonds.
- Distribute remaining electrons as lone pairs, starting with outer atoms, until each has an octet (duet for H).
- Check the central atom. If it has fewer than eight, move outer lone pairs into bonds as double or triple bonds.
- Verify the count. The drawing must use exactly the electrons counted in step 1.
Multiple bonds appear when the count demands them
Carbon dioxide has 4 + 2(6) = 16 valence electrons. Connecting O–C–O with single bonds and filling outer octets uses all 16 electrons, leaving carbon with only four — short of an octet. Moving one lone pair Pair of valence electrons not involved in bonding Full entry → from each oxygen into the bond gives O=C=O, and every atom has an octet. Multiple bonds are not a choice; they are required whenever single bonds leave the central atom electron-poor. The same logic produces the double bonds in SO₃ and the delocalized structures of CO₃²⁻ and NO₃⁻.
How It Works / Step-by-Step Process
- Count the total valence electrons (adjust for ionic charge).
- Arrange the skeleton: least electronegative atom central, hydrogens outside.
- Connect atoms with single bonds (two electrons each).
- Give outer atoms lone pairs until they reach octets.
- Move lone pairs into the central atom's bonding region if it still lacks an octet, forming double or triple bonds.
- Check the final electron count and each atom's octet.
Common Confusions
| Do not confuse | With | Difference |
|---|---|---|
| Central atom = most electronegative atom | Central atom = least electronegative atom | The least electronegative atom goes in the center; F is almost always terminal |
| The octet rule always holding | A rule with exceptions | H duet, B sextet, expanded octets, radicals all break strict octets |
| Lone pairs being unimportant | Lone pairs driving shape and reactivity | Lone pairs occupy space and act as electron donors (bases) |
| Ignoring charge when counting electrons | Adjusting the count for ions | CO₃²⁻ has 4 + 18 + 2 = 24 electrons; skipping the +2 for a negative charge is a classic error |
| "Any atom can be central" | Hydrogen is always terminal | H forms one bond only, so it sits outside |
| Placing dots randomly around a symbol | Placing one dot per side before pairing | "One per side first" reflects electron repulsion and is standard |
Eli explains
The same idea, in plain words
Explain it like I’m 10
A Lewis structure is a dot-picture of a molecule: each atom is its letter with dots for outer electrons, and shared dots become lines — the bonds. The picture shows how many bonds each atom wants and where lone pairs sit.
Worked example
Example 1: Carbon dioxide, CO₂
Count electrons: carbon contributes 4, each oxygen 6:
4 + 2(6) = 16 valence electrons
Place carbon in the center: O–C–O. Two single bonds use 4 electrons, leaving 12 — three lone pairs per oxygen. Now each oxygen has an octet, but carbon has only 4; move one lone pair from each oxygen into the bond:
O=C=O
Each oxygen retains two lone pairs plus the double bond (8 electrons), and carbon has two double bonds (8 electrons). The count is still 16, and every atom obeys the octet rule — predicting the experimentally observed 180° bond angle.
Example 2: Ammonia, NH₃, and water, H₂O
Ammonia: nitrogen contributes 5, each hydrogen contributes 1:
5 + 3(1) = 8 valence electrons
Nitrogen is central, connected to three hydrogens by single bonds (6 electrons used), leaving one lone pair on nitrogen. Every atom is satisfied: each H has 2, and N has 2 + 2 + 2 + 2 = 8. That lone pair makes ammonia a base — it accepts a proton, as in Topic 2.
Water: oxygen contributes 6, each hydrogen contributes 1:
6 + 2(1) = 8 valence electrons
Two O–H single bonds use 4 electrons, leaving two lone pairs on oxygen. Those lone pairs are why water's bond angle is about 104.5° rather than 180° — they take up space and compress the H–O–H angle.
Example 3: Sulfur hexafluoride, SF₆ — an expanded octet
Count electrons:
6 + 6(7) = 48 valence electrons
Sulfur is central, bonded to six fluorines by single bonds (12 electrons used), leaving 36 electrons as three lone pairs per fluorine. Sulfur itself now holds 12 electrons in its bonding region — an expanded octet — possible because period-3 sulfur has empty 3d orbitals. A strict octet structure is impossible: six fluorines require at least six bonds from sulfur, meaning at least 12 electrons around it.
Key takeaways
- Dots in a Lewis symbol = valence electrons; for main-group elements that is the group number.
- Octet rule: atoms tend toward eight valence electrons; H and He need only two.
- Counting is everything: sum valence electrons, then add 1 per negative charge or subtract 1 per positive charge.
- Least electronegative atom goes in the center; hydrogen is always terminal.
- Fill outer octets first; only then move lone pairs in to form multiple bonds for an electron-poor central atom.
- Exceptions: H duet, B sextet (BF₃), expanded octets for period 3+ (PCl₅, SF₆), radicals (NO, NO₂).
- Always verify the total electron count matches the drawing.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
How many valence electrons does phosphorus have? How many dots in its Lewis symbol?
Show answer
Phosphorus is in group 15: five valence electrons, five dots in its Lewis symbol (one per side, then a fifth paired).
Write the Lewis structure of methane, CH₄, and state how many lone pairs sit on carbon.
Show answer
Carbon shares one pair with each of four hydrogens: four C–H single bonds, no lone pairs on carbon; 8 valence electrons total.
Why must SF₆ have an expanded octet Central atom holding more than eight electrons via d orbitals Full entry → around sulfur?
Show answer
Six fluorines require six S–F bonds, placing at least 12 electrons around sulfur; period-3 sulfur can exceed an octet via empty 3d orbitals.
How many total valence electrons does the carbonate ion, CO₃²⁻, have?
Show answer
Carbon contributes 4, each oxygen 6, and the 2− charge adds 2: 4 + 3(6) + 2 = 24 electrons.
In the Lewis structure of HCN, how many bonds does carbon form?
Show answer
Carbon forms four bonds: a single bond to hydrogen and a triple bond to nitrogen (C≡N) — a full octet.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- Lewis symbol
- Element symbol with dots for valence electrons
- Lewis structure
- Drawing of a molecule showing bonds (lines) and lone pairs (dots)
- valence electron
- Electron in the outermost shell, available for bonding
- octet rule
- Tendency of atoms to surround themselves with eight valence electrons
- lone pair
- Pair of valence electrons not involved in bonding
- bonding pair
- Pair of electrons shared between two atoms
- expanded octet
- Central atom holding more than eight electrons via d orbitals
- radical
- Molecule or ion with an unpaired electron
Sources & references
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