Chemistry: Atoms First 2e · Chemical Bonding and Molecular Geometry
Lewis Symbols and Structures
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A Lewis symbol Element symbol surrounded by dots representing valence electrons Full entry → (electron-dot symbol) is a shorthand drawing of an atom showing only its valence electrons Electrons in the outermost occupied shell of an atom Full entry → — the outermost electrons that participate in bonding. The element symbol stands for the nucleus plus all inner (core) electrons; dots around it represent valence electrons. A Lewis structure Two-dimensional drawing showing connectivity, bonds, and lone pairs in a molecule or ion Full entry → extends the idea to molecules and ions: it shows which atoms are connected, which electron pairs are shared (bonding pairs), and which pairs belong to a single atom (lone pairs). Proposed by Gilbert N. Lewis in 1916, it remains chemistry's first tool for predicting how atoms combine.
Why this matters
Lewis structures turn a periodic table into a predictive tool: a quick drawing reveals how many bonds an atom tends to form, whether two elements will form an ionic or covalent compound, and where lone pairs sit. Later topics — VSEPR geometry, formal charge, resonance, and molecular orbital theory — all start from a Lewis structure. In the real world, water's bent shape, the hydrogen bonds holding DNA strands together, and the way a drug fits a receptor are all consequences of electron pairs mapped on paper this way.
The college version
Core Concepts
Valence electrons and the octet rule
Valence electrons are the electrons in an atom's highest occupied principal energy level. For representative (main-group) elements, the number of valence electrons equals the group number: 1 for group 1, 2 for group 2, 3 for group 13, up to 8 for group 18. The octet rule Tendency of main-group atoms to surround themselves with eight valence electrons Full entry → states that main-group atoms tend to share, gain, or lose electrons until surrounded by eight valence electrons — the stable configuration of the nearest noble gas. Two qualifications matter: hydrogen and helium need only two electrons (a "duet"), and elements in period 3 and beyond (P, S, Cl) can hold more than eight when needed. The octet rule is a strong tendency, not a law.
Lewis symbols for atoms and ions
To write a Lewis symbol, place dots around the four sides of the element symbol (top, right, bottom, left), one per side before any side gets a second. Chlorine (group 17, seven valence electrons) gets one dot on each side plus a pair on the fourth. For ions, add or remove dots to reflect the charge: a sodium cation is written Na⁺ with no dots (it lost its one valence electron), while a chloride anion is written [Cl]⁻ with eight dots in brackets, showing the octet belongs to the ion as a whole.
The Lewis structure algorithm
For covalent molecules and polyatomic ions, use this sequence:
- Count valence electrons. Sum group numbers; add one electron per negative charge, subtract one per positive charge.
- Choose a skeleton. Place the least electronegative atom in the center (hydrogen is never central) and connect the others with single bonds.
- Distribute remaining electrons. Place lone pairs on the outer (terminal) atoms first, aiming for octets (duets for H).
- Form multiple bonds if needed. If the central atom still lacks an octet, convert a terminal lone pair Electron pair owned by a single atom in a structure Full entry → into a shared pair (double or triple bond).
Ionic vs. covalent representations
In ionic compounds, electrons transfer, so the picture shows separate ions: the cation loses dots, the anion gains them, each in brackets with its charge. In a covalent compound, electrons are shared, so one structure shows every atom connected by bonds with lone pairs around it. The same counting rules work for both; the difference is where electrons end up.
Multiple bonds and common exceptions
When the central atom cannot reach an octet with single bonds alone, the structure uses double or triple bonds — carbon dioxide (O=C=O) and nitrogen (N≡N) are everyday examples. Some molecules legitimately break the rule: BeCl₂ leaves beryllium with four electrons, BF₃ leaves boron with six, and SF₆ and PCl₅ expand beyond eight using d orbitals in period 3 and beyond. Recognize these patterns so you do not force an octet the drawing does not require.
How It Works / Step-by-Step Process
- Decide whether the compound is ionic or covalent.
- Count total valence electrons, adjusting for charge.
- Assemble a skeleton: least electronegative atom central; H and halogens terminal.
- Connect atoms with single bonds (2 electrons per bond).
- Distribute leftover electrons as lone pairs on terminal atoms, then the central atom, until octets are satisfied.
- If the central atom is short of an octet, move terminal lone pairs into the bond region to form double or triple bonds.
- For ions, enclose the structure in brackets with the charge outside.
Common Confusions
| Do Not Confuse | With | The Difference |
|---|---|---|
| Total valence electrons | Dots in a finished structure | The total is your starting budget; the finished structure shows the same electrons as bonds and lone pairs |
| The octet rule | An absolute law | It is a strong tendency with well-known exceptions (H, Be, B, expanded octets) |
| A double bond | Two lone pairs | A double bond is two shared pairs between the same two atoms; lone pairs are unshared |
| The element's group number | The number of dots for any element | Group number gives valence electrons only for main-group elements; transition metals follow different patterns |
| Electronegativity (who pulls electrons) | Formal charge (bookkeeping) | Electronegativity is real electron attraction; formal charge is bookkeeping covered in the next topic |
| Ion charge | Charge of an individual atom in a covalent structure | The ion's charge is the total; per-atom charges are formal charges, a different concept |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Lewis symbols are like sticker maps of an atom: the element's letter is the atom, and the stickers around it are the outer electrons that do the bonding. Lewis structures are the same idea for whole molecules — they show which atoms hold hands (share electrons) and which electrons hang around on one atom. The octet rule is the rule of eight: most atoms feel happiest with eight outer electrons and will share or trade stickers to get there.
Worked example
Example 1: Lewis structure of carbon dioxide (CO₂)
Step 1 — count electrons with unit analysis:
ne = 1 C atom × 4 e-C atom + 2 O atoms × 6 e-O atom = 4 e- + 12 e- = 16 e-
Step 2 — skeleton. Carbon is less electronegative than oxygen, so C is central: O–C–O. Two single bonds use 4 electrons, leaving 12.
Step 3 — lone pairs. Give each O three lone pairs (12 electrons). Oxygen atoms now have octets, but carbon has only 4 electrons.
Step 4 — multiple bonds. Move one lone pair from each oxygen into a shared pair, giving O=C=O. Every atom now has an octet, and all 16 electrons are accounted for: 4 in bonds plus 12 as lone pairs.
Example 2: Lewis structure of the ammonium ion (NH₄⁺)
Count electrons for the ion, remembering the +1 charge removes one electron:
ne = 5 e- + 4 × 1 e- - 1 e- = 8 e-
Nitrogen is central (H cannot be central). Four N–H single bonds use all 8 electrons, leaving zero lone pairs; every H has its duet and N has an octet. The structure is written [NH₄]⁺ because the charge belongs to the whole ion — which also explains why ammonium salts dissolve readily in water.
Key takeaways
- Number of dots in a Lewis symbol = group number for main-group elements (1–8).
- Octet rule exceptions: H/He (2 electrons), Be and B (4 and 6), period-3+ elements (expanded octets).
- Sequence: total valence electrons → skeleton (least electronegative central, H never central) → lone pairs on outer atoms → multiple bonds if needed.
- Add one electron per negative charge, subtract one per positive charge, when counting ions.
- A double bond is two shared pairs (4 electrons); a triple bond is three shared pairs (6 electrons).
- Draw multiple bonds only to complete the central atom's octet — never as a shortcut.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
How many valence electrons does sulfur (group 16) have, and how many dots does its Lewis symbol show?
Show answer
Six; six dots (one on each side, then paired).
Why is hydrogen never placed in the center of a Lewis structure?
Show answer
Hydrogen can hold only two electrons (a duet), so it can form only one bond and cannot connect multiple atoms.
How many electrons get added when counting for a polyatomic ion with a 2− charge?
Show answer
Two electrons — add one per negative charge.
When does a Lewis structure require a multiple bond rather than lone pairs?
Show answer
When the central atom still lacks an octet after all lone pairs are placed on terminal atoms.
Name two elements that commonly exceed the octet rule and one that commonly has fewer than eight.
Show answer
Expanded octets: P and S (or Cl); fewer than eight: Be and B.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- valence electrons
- Electrons in the outermost occupied shell of an atom
- Lewis symbol
- Element symbol surrounded by dots representing valence electrons
- Lewis structure
- Two-dimensional drawing showing connectivity, bonds, and lone pairs in a molecule or ion
- octet rule
- Tendency of main-group atoms to surround themselves with eight valence electrons
- lone pair
- Electron pair owned by a single atom in a structure
- bonding pair
- Electron pair shared between two atoms in a covalent bond
Sources & references
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