Chapter 3 – Structure & Properties of Ionic and Covalent Compounds

3.1 Chemical Bonding

• Chemical bond = attractive force linking two atoms
  • Overcomes nucleus–nucleus repulsion
  • Arises from interactions of valence electrons

Lewis Symbols (Dot Structures)

• Represent atom with elemental symbol + dots for valence e⁻
• Construction rules
  • Four sides around symbol; max 2 dots per side ⇒ $8$ total
  • Place one dot per side (↑,→,↓,←) before pairing
  • Unpaired dot = potential bonding electron
• Enables easy use of octet rule (atoms seek $8$ e⁻ in valence shell)

Principal Bond Types

• Ionic bond
  • Complete e⁻ transfer: metal → non-metal
  • Electrostatic attraction between cation (+) and anion (–)
• Covalent bond
  • e⁻ pair shared between atoms (usually non-metals)
• Many real bonds fall on an ionic ↔ covalent continuum

Ionic Bonding Details

• Representative-element metals form cations (low IE & EA)
• Non-metals form anions (high IE & EA)
• Both ions adopt nearest noble-gas electron configuration (octet rule)
• Example: $\text{Na} \rightarrow \text{Na}^+ + e^-$ (Neon core)
  $\text{Cl}+e^- \rightarrow \text{Cl}^-$ (Argon core)
• Crystal lattice: repeated Na⁺/Cl⁻ ion pairs; cubic habit visible macroscopically

Covalent Bonding Details

• Forms between atoms with similar ability to gain/lose e⁻
• Diatomic elements (H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂) = purely covalent
• Example H₂ formation
  • Each H shares its 1 e⁻ ⇒ both achieve He configuration (2 e⁻)
  • Shared pair = covalent bond

Polar Covalent Bonds & Electronegativity (EN)

• Polar covalent = unequal e⁻ sharing
• Electronegativity (EN): tendency to attract bonding e⁻
  • Increases up & right across periodic table
  • Difference $\Delta EN$ dictates bond type
  • $\Delta EN>1.9$ ≈ ionic
  • $0.5<\Delta EN<1.9$ polar covalent
  • $\Delta EN<0.5$ non-polar covalent
• Example: $\Delta EN_{\text{H–F}}=4.0-2.1=1.9$ ⇒ more polar than $\text{H–Cl}\;(\Delta EN=0.9)$

3.2 Nomenclature – Naming & Writing Formulas

Basics

• Formula: symbols + numeric subscripts (implied 1 if absent)
• Nomenclature systems
  • Ionic compounds
  • Covalent (molecular) compounds

Ionic Compounds

• Smallest whole-number ratio of ions reported
• To write formula:
  1. Determine ionic charges (Group number cues)
  2. Combine to achieve charge neutrality
• Example predictions
  • Na⁺ & O²⁻ ⇒ $\text{Na}_2\text{O}$
  • Al³⁺ & O²⁻ ⇒ $\text{Al}<em>2\text{O}</em>3$

Naming Rules

• Binary ionic: cation name + anion stem + “-ide”
  • $\text{NaCl} \rightarrow$ sodium chloride
• Variable-charge metals (transition metals): Stock system
  • Roman numeral gives cation charge; e.g. $\text{FeCl}_3$ = iron(III) chloride
  • Common system: -ous (lower) vs -ic (higher): Fe²⁺ = ferrous, Fe³⁺ = ferric

Monatomic Ions (selected)

• Cations: $\text{H}^+,\;\text{Li}^+,\;\text{Na}^+,\;\text{Mg}^{2+},\;\text{Al}^{3+},\;\text{Ag}^+$
• Anions: $\text{F}^-,\;\text{O}^{2-},\;\text{N}^{3-},\;\text{P}^{3-}$

Polyatomic Ions (selected)

• Positive: $\text{NH}<em>4^+$ (ammonium), $\text{H}</em>3\text{O}^+$ (hydronium)
• Negative: $\text{SO}<em>4^{2-}$ (sulfate), $\text{NO}</em>3^-$ (nitrate), $\text{OH}^-$ (hydroxide), $\text{CH}_3\text{COO}^-$ (acetate)

Sample Naming Practice

• $\text{NH}_4\text{Cl}$ → ammonium chloride
• $\text{Fe}(\text{NO}<em>3)</em>3$ → iron(III) nitrate
• $\text{Ca(OH)}_2$ → calcium hydroxide

Writing Ionic Formula from Name

• Barium chloride ⇒ Ba²⁺ + Cl⁻ gives $\text{BaCl}_2$

Covalent (Molecular) Compounds

• Typically non-metal + non-metal
• Prefixes indicate atom counts
  • mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-
• Omit “mono-” on first element; last element stem + “-ide”
  • $\text{CO}$ = carbon monoxide (not monocarbon…)
  • $\text{N}<em>2\text{O}</em>5$ = dinitrogen pentoxide
• Formula writing uses prefixes directly
  • nitrogen trifluoride ⇒ $\text{NF}_3$
• Common names: $\text{H}<em>2\text{O}$ (water), $\text{NH}</em>3$ (ammonia), $\text{C}<em>2\text{H}</em>5\text{OH}$ (ethanol), $\text{C}<em>6\text{H}</em>{12}\text{O}_6$ (glucose)

3.3 Physical Properties: Ionic vs Covalent

• Physical state @ RT
  • Ionic → usually solid, crystalline
  • Covalent → solid/liquid/gas; crystalline or amorphous
• Melting/boiling points
  • Ionic: very high (many hundreds °C) due to strong electrostatic lattice
  • Covalent: lower (varies with IM forces)
• Conductivity in water
  • Ionic solids dissociate → electrolytes (conduct)
  • Covalent compounds generally remain molecules → nonelectrolytes

3.4 Lewis Structures – Guidelines & Practice

1. Draw skeletal structure (least EN atom central; H & halogens terminal)
2. Count total valence e⁻ (adjust for ionic charge)
3. Connect with single bonds; complete octets of terminals, then central
4. If central lacks octet, convert lone pairs into multiple bonds
5. Verify octets & e⁻ count

Examples

• CO₂: O=C=O with two C=O double bonds (16 e⁻ total)
• $\text{CO}_3^{2-}$
  • 24 valence e⁻
  • Resonance: three equivalent structures, each with one C=O double bond
• Exceptions
  • Incomplete octet: BeH₂ (4 e⁻ around Be)
  • Odd e⁻: NO (11 e⁻ total)
  • Expanded octet: PF₅ (10 e⁻ around P)

Bond Order & Energies

• Single < Double < Triple in energy; reverse order for bond length
  • $E<em>{\text{triple}}>E</em>{\text{double}}>E_{\text{single}}$
  • $\text{length}<em>{\text{single}}>\text{length}</em>{\text{double}}>\text{length}_{\text{triple}}$

Molecular Geometry – VSEPR Theory

• Electron groups (bonded atoms + lone pairs) arrange to minimize repulsion
• Basic shapes & ideal angles (central atom):
  • 2 groups ⇒ Linear, $180^\circ$ (e.g., CO₂)
  • 3 groups, 0 lone ⇒ Trigonal planar, $120^\circ$ (SO₃)
  • 3 groups, 1 lone ⇒ Bent, $<120^\circ$ (SO₂)
  • 4 groups, 0 lone ⇒ Tetrahedral, $109.5^\circ$ (CH₄)
  • 4 groups, 1 lone ⇒ Trigonal pyramidal, $\approx107^\circ$ (NH₃)
  • 4 groups, 2 lone ⇒ Bent, $\approx104.5^\circ$ (H₂O)
• Exceptions revisited
  • BeH₂: linear despite incomplete octet
  • BF₃: trigonal planar with 6 e⁻ on B

Molecular Polarity

• Molecule polar if centers of + and – charge differ
• Procedure
  1. Draw Lewis & geometry
  2. Mark bond dipoles (arrow toward more EN atom)
  3. Assess vector cancellation
• Heuristics
  • No lone pair + identical terminal atoms ⇒ non-polar (e.g., CO₂)
  • One lone pair ⇒ usually polar (e.g., NH₃, SO₂)
  • Multiple lone pairs ⇒ usually polar (H₂O)
• Practice outcomes
  • HF, H₂O polar; CH₄, O₂ non-polar

3.5 Intermolecular Forces (IMF) & Macroscopic Properties

• Intramolecular = bonds inside molecule; intermolecular = between molecules
• Types (in covalent compounds)
  • Dipole–dipole (polar molecules)
  • Hydrogen bonding: $\text{H}···\text{F, O, N}$ strong special dipole
  • London (dispersion) forces (all molecules; dominant in non-polar)
• Solubility rule: “like dissolves like”
  • Polar dissolves in polar; non-polar in non-polar
  • Example: NH₃ dissolves in H₂O via hydrogen bonding; oil (non-polar) immiscible with water
• Boiling/Melting trends
  1. Stronger IMF ⇒ higher $m.p.$, $b.p.$
  2. Greater molecular mass ⇒ higher $m.p.$, $b.p.$ (if IMF similar)
  3. Polarity increases $m.p./b.p.$ for similar masses
• Sample data (℃)
  • $\text{H}_2\text{O}: m.p.=0,\;b.p.=100$ (polar, H-bond)
  • $\text{N}_2: m.p.=-210,\;b.p.=-196$ (non-polar)
  • $\text{NaCl}: m.p.=801, b.p.=1413$ (ionic lattice)

Practice & Applications

• Predict ionic vs covalent, polarity, geometry, formula writing, naming, boiling point ordering
• Dimethyl ether (CH₃–O–CH₃)
  • Two tetrahedral CH₃ groups; O center geometry similar to water ⇒ bent region producing net molecular dipole

Key Numerical/Statistical References

• $\Delta EN$ classification thresholds: $0.5, 1.9$
• Ideal bond angles: $180^\circ, 120^\circ, 109.5^\circ, 107^\circ, 104.5^\circ$
• Sample lattice melting points: $\text{NaCl }801^\circ\text{C},\;\text{KBr }730^\circ\text{C}$

Ethical & Practical Implications

• Understanding bonding guides material design (pharmaceuticals, polymers)
• Knowledge of ion formation critical for environmental toxicity (heavy-metal cation behavior)
• Naming conventions ensure global, unambiguous chemical communication, preventing laboratory or industrial errors