Back/Chemistry: Atoms First 2e
Section 4.410 Key Terms

Lewis Symbols and Structures

Learning Objectives
  • Write Lewis symbols for neutral atoms and monatomic ions
  • Draw Lewis structures depicting covalent bonding in molecules and polyatomic ions
  • Identify and explain exceptions to the octet rule, including odd-electron, electron-deficient, and hypervalent species

Core Concepts & Principles

Lewis symbols and structures provide a straightforward visual method for tracking valence electrons and understanding chemical bonding.

  • Lewis Symbols: An elemental symbol surrounded by dots representing its valence electrons. Monatomic cations show fewer dots (lost electrons), while anions show gained dots filling their outer shell.
  • Lewis Structures: Drawings that illustrate covalent bonding in molecules and polyatomic ions. Shared electron pairs form covalent bonds (represented by dashes or pairs of dots), while unshared valence electrons are represented as lone pairs.
  • Bond Multiplicity: Atoms can share one, two, or three pairs of electrons, forming single bonds, double bonds, or triple bonds, respectively.
The Octet Rule & Bonding
  • Main-group elements tend to form enough bonds to acquire eight valence electrons (octet rule), achieving a stable noble gas configuration.
  • Hydrogen is a primary exception, requiring only two electrons to fill its valence shell (a duet).
  • Period 3 and higher elements can accommodate more than eight electrons due to empty valence d-orbitals, forming hypervalent molecules.

Problem-Solving Routines & Methods

Drawing Lewis Structures
  1. 1
    Calculate the total valence electron pool: sum the valence electrons of all constituent atoms. For anions, add one electron per negative charge; for cations, subtract one per positive charge.
  2. 2
    Draw the skeleton structure: place the least electronegative atom in the center (hydrogen and fluorine are never central). Connect atoms with single bonds.
  3. 3
    Distribute remaining electrons as lone pairs on terminal atoms to complete their octets (or duets for hydrogen).
  4. 4
    Place any leftover electrons on the central atom as lone pairs.
  5. 5
    Form multiple bonds: if the central atom lacks an octet, move lone pairs from terminal atoms to form double or triple bonds until octets are satisfied where possible.
Pro-Tip: Always verify that your final structure uses the exact total number of valence electrons calculated in Step 1 and does not violate element bonding capacities.
Valence Electron Pool Calculation
Total e−=∑(valence e−)±charge\text{Total } e^- = \sum (\text{valence } e^-) \pm \text{charge}

Determines the total number of valence electrons available for distribution in a Lewis structure.

Variables & Constants
∑(valence e−)\sum (\text{valence } e^-)=sum of valence electrons from neutral atoms;
++=add 1 for each negative charge;
−-=subtract 1 for each positive charge

Practice & Concept Checks

Concept Check
Why does boron in boron trifluoride (BF3BF_3) often remain electron-deficient rather than forming a double bond to satisfy the octet rule?
Concept Check
How many total valence electrons are used to draw the Lewis structure for the cyanide ion (CN−CN^-)?

Key Terms & Vocabulary

Lewis symbolAtomic Structure
An elemental symbol surrounded by dots that represents the valence electron configuration of an atom or monatomic ion.
Example: A sodium atom represented with one dot.
Lewis structuresMolecular Bonding
Diagrams that show the bonding between atoms in a molecule or polyatomic ion, using lines for covalent bonds and dots for lone pairs.
Example: H−O¨−HH-\ddot{O}-H for water.
lone pairsMolecular Bonding
Pairs of valence electrons that are not involved in covalent bond formation.
Example: Three pairs of unshared electrons on each chlorine atom in Cl2Cl_2.
single bondCovalent Bonding
A covalent bond formed by the sharing of a single pair of electrons between two atoms.
Example: The H−ClH-Cl bond in hydrogen chloride.
octet ruleBonding Rules
The tendency of main-group atoms to form enough covalent bonds to achieve eight valence electrons in their outer shell.
Example: Carbon forming four bonds in CH4CH_4.
double bondMultiple Bonding
A covalent bond formed by the sharing of two pairs of electrons between two atoms.
Example: The C=OC=O bond in formaldehyde (CH2OCH_2O).
triple bondMultiple Bonding
A covalent bond formed by the sharing of three pairs of electrons between two atoms.
Example: The C≡OC \equiv O bond in carbon monoxide (COCO).
SmalleyHistory & Discovery
Richard Smalley, a pioneer in fullerene chemistry who co-discovered buckminsterfullerene (C60C_{60}) and was honored as a father of nanotechnology.
Example: Co-recipient of the 1996 Nobel Prize in Chemistry.
free radicalsExceptions to Octet Rule
Molecules that contain an odd number of valence electrons and possess an unpaired electron.
Example: Nitric oxide (NONO).
hypervalent moleculesExceptions to Octet Rule
Molecules with a central atom from period 3 or higher that holds more than eight valence electrons using empty d-orbitals.
Example: PCl5PCl_5 and SF6SF_6.
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4.3 Chemical Nomenclature
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4.5 Formal Charges and Resonance