CHEM10007 · Fundamentals Of Chemistry
Chemical Bonding, Lewis Structures & Shape
This chapter explains why atoms join and what shape the result takes. You distinguish ionic, covalent and metallic bonding, learn the charges of common monatomic and polyatomic ions, and name ionic and covalent compounds. The core skill is drawing Lewis (electron-dot) structures — counting valence electrons, satisfying the octet rule, placing lone pairs and showing multiple bonds, formal charge and resonance — which feeds directly into polarity and molecular shape in later weeks.
What this chapter covers
- 01Ionic bonding: metal + non-metal, electron transfer, ionic lattices
- 02Covalent bonding: non-metal + non-metal, electron sharing; molecules vs covalent lattices
- 03Metallic bonding: a 'sea' of delocalised electrons
- 04Common ions: monatomic charges from group position; polyatomic ions (NO3−, SO42−, CO32−, PO43−, NH4+, OH−)
- 05Naming: ionic compounds (Roman numerals for variable-charge metals) and covalent compounds (Greek prefixes)
- 06Lewis structures: count valence electrons, octet rule, lone vs bonding pairs, multiple bonds
- 07Formal charge and resonance; first-year octet exceptions
- 08Molecular shape from electron domains (introduced lightly as a supporting visual)
Lewis structure and valence-electron count
- 1 mark — total valence electron countCount valence electrons: 5 (N) + 3 × 6 (O) + 1 (for the 1− charge) = 24 electrons.
- 1 mark — correct skeletonPlace N as the central atom with three N–O bonds; that uses 6 electrons, leaving 18 for lone pairs.
- 1 mark — octets satisfied and charge shownComplete octets on the oxygens; to give N an octet, form one N=O double bond, and enclose the whole structure in brackets with an overall 1− charge.
- 1 mark — resonance explainedThe double bond can sit on any of the three equivalent oxygens, so NO3− is described by three resonance structures; the real ion is the average, with three identical N–O bonds.
Key terms
- Ionic bond
- An electrostatic attraction between oppositely charged ions formed by electron transfer from a metal to a non-metal, giving a 3-D lattice.
- Covalent bond
- A bond formed by shared electron pairs between non-metal atoms, producing discrete molecules or, less often, covalent lattices.
- Lewis (electron-dot) structure
- A diagram showing all valence electrons as bonding pairs and lone pairs, used to predict bonding, charge distribution and shape.
- Formal charge
- The hypothetical charge on an atom if bonding electrons were shared equally; the best Lewis structure minimises formal charges. Calculated as valence electrons − lone-pair electrons − ½ bonding electrons.
- Resonance
- When two or more valid Lewis structures differ only in electron placement, the true structure is a weighted average (hybrid) of them, e.g. the three equivalent forms of NO3−.
Chemical Bonding, Lewis Structures & Shape FAQ
How do I count valence electrons for an ion?
Sum the valence electrons of every atom, then add one electron for each unit of negative charge or subtract one for each unit of positive charge. For NO3− that is 5 + 18 + 1 = 24.
When do I use Roman numerals in a name?
Use them for metals with more than one common charge (transition metals and some main-group metals), e.g. iron(II) vs iron(III). Fixed-charge metals like Na or Mg never take a Roman numeral; covalent compounds use Greek prefixes instead.
What is the difference between a polar bond and a polar molecule?
A polar bond comes from an electronegativity difference between two atoms. A molecule is polar only if the vector sum of its bond dipoles, given its geometry, is non-zero — symmetric molecules such as CO2 have polar bonds but are non-polar overall.
Exam move
Memorise the common polyatomic ions and their charges early — they appear in naming, precipitation and equilibrium questions all semester. Practise Lewis structures as a fixed routine (count → skeleton → octets → multiple bonds → formal charge/resonance) so you never freeze on an unfamiliar species. Keep molecular shape light here, but note which geometries cancel bond dipoles, because that connects directly to the intermolecular-forces chapter.
