Chapter 4 Notes: Covalent Bonding, Lewis Structures, & Molecular Geometry

4.1 Introduction to Covalent Bonding

• Definition of a covalent bond
  • Two-electron bond resulting from the sharing of valence electrons between two atoms.
  • Gives rise to discrete units called molecules.
• Lone (non-bonded) electron pairs
  • Pairs of valence electrons that are not shared.
• Octet/duet objective
  • $\text{H}$ attains $2\,e^-$ (duet).
  • Main-group elements (Groups 1A–8A, except H) share electrons until each atom possesses an octet (8 e⁻) in its valence shell.
• Periodic-table perspective
  • Covalent bonds form between two non-metals or a metalloid and a non-metal.
  • Hydrogen = 1 bond.
  • For atoms with 1–3 valence e⁻, the number of covalent bonds equals the number of valence e⁻.
  • For atoms with ≥ 4 valence e⁻, use
  $\text{predicted number of bonds}=8-\text{number of valence }e^-$
• General bonding rule (excl. H)
  $\text{number of bonds}+\text{number of lone pairs}=4$
• Human-body connection
  • Many heart-related biomolecules (e.g., nitric oxide, carbon monoxide in low amounts) are covalent compounds.

4.2 Lewis Structures

• Molecular formula vs. Lewis structure
  • Molecular formula: identities + counts only.
  • Lewis: connectivity and placement of all bonding & non-bonding valence e⁻.
• General rules for drawing Lewis structures
  1. Draw only valence electrons.
  2. Give every main-group atom an octet (H gets 2 e⁻).
  3. Place H & halogens at the periphery (only 1 bond each).
• Stepwise procedure
  1. Arrange atoms likely bonded; use common bonding patterns (C=4, N=3, O=2, halogens=1).
  2. Count total valence e⁻: sum of group numbers.
  3. Connect atoms with single bonds (2 e⁻ per bond). Distribute remaining e⁻ as lone pairs to satisfy octets/duets, starting with peripheral atoms.
  4. If octets are incomplete after all valence e⁻ are placed: convert lone-pair(s) on adjacent atoms into additional bonding pair(s) → double or triple bonds.
• Multiple bonds
  • Double bond = 4 e⁻ in two shared pairs.
  • Triple bond = 6 e⁻ in three shared pairs.
• Worked example: $\mathrm{C<em>2H</em>4}$
  • 12 valence e⁻.
  • After placing single bonds, each C lacks an octet → convert one lone pair to a second C–C bond → ethene structure with a C=C double bond, all atoms octet/duet satisfied.

4.3 Exceptions to the Octet Rule

• Electron-deficient atoms (Group 3A)
  • $\text{B}$, $\text{Al}$ often stable with only $6\,e^-$.
  • Example: $\mathrm{BF_3}$ displays 6 valence e⁻ on B.
• Expanded octets (3rd period & beyond)
  • Elements such as $\text{P}$, $\text{S}$ can utilize empty $d$-orbitals to accommodate > 8 e⁻ (e.g., $\mathrm{PF<em>5}$, $\mathrm{SF</em>6}$).
• Hydrogen
  • Never exceeds 2 e⁻.

4.4 Resonance

• Polyatomic ions: electron accounting
  • Add 1 e⁻ per negative charge; subtract 1 e⁻ per positive charge.
• Resonance definition
  • Two or more valid Lewis structures with identical atomic positions but different e-distribution.
  • Neither structure is the “true” one; actual species is a resonance hybrid.
• Stabilization
  • Delocalization of e⁻ lowers energy by spreading charge and multiple-bond character over a larger region.
• Environmental focus – Ozone ($\mathrm{O_3}$)
  • Exists as two resonance forms.
  • Formed in the stratosphere (reaction of O₂ with O atoms).
  • Protects Earth by absorbing high-energy UV radiation.
  • Graphical representation (altitude vs. temperature) shows protective ozone layer between ≈ 15–35 km.

4.5 Naming Covalent Compounds

• Two-step protocol
  1. Name 1st non-metal by its element name; 2nd non-metal gets “-ide.”
  2. Apply prefixes to denote number of atoms (Table 4.1).
     • Prefix list: 1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa-, 7 hepta-, 8 octa-, 9 nona-, 10 deca-.
     • Omit “mono-” from first element but retain for second (e.g., CO = carbon monoxide).
     • Drop the first vowel when two vowels clash (monooxide → monoxide).
• Examples
  • $\mathrm{NO<em>2}$ → nitrogen dioxide. • $\mathrm{N</em>2O_4}$ → dinitrogen tetroxide.

4.6 Molecular Shape (VSEPR Theory)

• Concept
  • Geometry determined by number of electron groups (atoms + lone pairs) around a central atom.
  • Most stable arrangement separates groups as far apart as possible.
• Two groups
  • Shape: linear; bond angle $\approx 180^\circ$.
  • Example: $\mathrm{CO_2},\,HCN$.
• Three groups
  • Shape: trigonal planar; angle $\approx 120^\circ$.
  • Examples: $\mathrm{BF<em>3},\,H</em>2CO$.
• Four groups
  • All atoms (0 lone pairs): tetrahedral; angle $109.5^\circ$ (e.g., $\mathrm{CH<em>4}$). • 1 lone pair: trigonal pyramidal; angle $\approx 107^\circ$ (e.g., $\mathrm{NH</em>3}$).
  • 2 lone pairs: bent; angle $\approx 105^\circ$ (e.g., $\mathrm{H_2O}$).
• Table 4.2 summary
  • Consolidates group count, lone pairs, shape, and sample molecules (see above).
• 3-D drawing conventions (text alt)
  • Solid lines = in-plane; wedge = toward viewer; dashed = behind plane.

4.7 Electronegativity & Bond Polarity

• Electronegativity (EN): relative attraction an atom exerts on shared e⁻.
  • Higher EN → stronger pull.
• Bond classification based on EN difference (ΔEN)
  • $\Delta EN<0.5$ → non-polar covalent (equal sharing). • $0.5\le\Delta EN\le1.9$ → polar covalent (unequal sharing, bond dipole). • $\Delta EN>1.9$ → ionic (electron transfer).
• Dipole notation
  • Arrow points toward more electronegative atom; cross marks δ⁺ end.
  • Example: C–O bond has δ⁻ on O, δ⁺ on C.

4.8 Polarity of Molecules

• Criteria for molecular polarity
  1. Presence/absence of polar bonds.
  2. Spatial cancellation or reinforcement of bond dipoles via molecular geometry.
• Non-polar molecules
  • Either contain no polar bonds or individual dipoles cancel (vector sum ≈ 0).
  • Example: $\mathrm{BF_3}$ (three identical 120° dipoles cancel).
• Polar molecules
  • Contain ≥ 1 polar bond and dipoles do not cancel.
  • Examples & rationales (image alternatives):
  – $\mathrm{H<em>2CO}$: trigonal planar; C=O dipole not cancelled → net dipole along C=O. – $\mathrm{NH</em>3}$: trigonal pyramidal; three N–H dipoles reinforce toward N.
  – $\mathrm{CH<em>2Cl</em>2}$: tetrahedral; C–Cl dipoles dominate → net toward Cl atoms.

Supplemental Bonding Patterns (Text Alternative)

• Typical valence patterns
  • H: 1 bond, 0 lone pairs.
  • C: 4 bonds, 0 lone pairs.
  • N: 3 bonds, 1 lone pair.
  • O: 2 bonds, 2 lone pairs.
  • Halogens (F, Cl, Br, I): 1 bond, 3 lone pairs.

Ethical / Practical / Real-World Connections

• Medical relevance: Accurate knowledge of covalent compounds underpins understanding of cardiac drugs (e.g., nitroglycerin → NO release) and toxicants (e.g., CO binding to hemoglobin).
• Environmental relevance: Resonance-stabilized ozone layer critical in filtering UV; anthropogenic CFCs jeopardize this natural shield.