Bond Type and Electronegativity Difference

Classification of Chemical Bonds by Electronegativity Difference

• Electronegativity (EN) is a dimensionless number that quantifies how strongly an atom attracts a shared pair of electrons.
• Pauline scale is the standard; ranges roughly 0.7\,(\text{Cs}) to 4.0\,(\text{F}).
• Notation: ENX = electronegativity of element X. • Bond polarity is governed by the magnitude of the difference \Delta EN between two bonded atoms: \Delta EN = |ENA - EN_B|

• Three canonical bond types (thresholds memorized for most general‐chemistry courses):

Non-polar covalent bond
• \Delta EN < 0.5 ("essentially equal sharing" of electrons).
• No permanent dipole moment; electron cloud is symmetrically distributed.
• Typical examples: \text{H–H},\ \text{F–F},\ \text{C–H} (although C–H is often treated as weakly polar in higher‐level courses, it falls under the <0.5 guideline here).
• Macroscopically yields substances with low dielectric constants (e.g., \text{O}_2, hydrocarbons).
• Intermolecular forces dominated by London (dispersion) interactions.
Polar covalent bond
• 0.5 \le \Delta EN \le 2.1 (partial electron displacement toward more electronegative atom).
• Leads to bond dipoles; molecule may be overall polar if dipole vectors do not cancel.
• Representative pairs highlighted in transcript:
– \text{F–H}:\ \Delta EN \approx 4.0-2.1 = 1.9
– \text{F–S}:\ \Delta EN \approx 4.0-2.5 = 1.5
– \text{F–P}:\ \Delta EN \approx 4.0-2.1 = 1.9
• Consequences: higher boiling points than non-polar analogues, miscibility with water, ability to engage in dipole–dipole & hydrogen bonding (when \text{H} is attached to \text{N},\text{O},\text{F}).
Ionic bond
• \Delta EN > 2.1 (approximate full electron transfer from electropositive to electronegative species).
• Formation of a lattice of discrete cations and anions rather than individual molecules.
• Classical example: \text{NaCl}:\ EN{\text{Na}}\,\approx 0.9,\ EN{\text{Cl}}\,\approx 3.0,\ \Delta EN\,\approx 2.1.
• Properties: high melting/boiling points, brittleness, electrical conductivity in molten/aqueous states.

Conceptual Rationale Behind Cut-Off Values

• Thresholds (0.5 & 2.1) are empirical rules of thumb useful in introductory settings; actual borderline behaviour is continuous, not discrete.

• Bond character can be expressed as % ionic character:\%\,\text{ionic} \approx \left(1 - e^{-0.25(\Delta EN)^2}\right)\times 100 (Pauling’s empirical formula).

• At \Delta EN \approx 2.1, % ionic character surpasses ~50 %, justifying the qualitative jump from polar covalent to ionic.

Practical / Real-World Relevance

• Solubility rules correlate with bond polarity ("like dissolves like"): ionic & highly polar compounds dissolve in polar solvents (water); non-polar molecules dissolve in hexane, benzene, etc.

• Biomolecular function: polar covalent bonds create partial charges that enable hydrogen bonding → critical for protein secondary structure and DNA base pairing.

• Material science: tailoring bond polarity in polymers controls dielectric behaviour, thermal resistance, and mechanical flexibility.

Periodic Trends Driving \Delta EN

• Electronegativity generally increases left → right across a period and decreases top → bottom in a group.

• This explains why bonds with fluorine (highest EN) often fall into the polar/ionic regimes and why Cs–F is one of the most ionic bonds known (\Delta EN \approx 4.0 - 0.7 = 3.3).

Caveats & Advanced Notes

• Metallic bonding: not captured by \Delta EN tables; involves a delocalised "sea" of electrons.

• Polyatomic ions & resonance (e.g., \text{SO}_4^{2-}) display partial bond orders; individual S–O bonds have intermediate character that simple \Delta EN rules cannot fully describe.

• C–Cl (\Delta EN \approx 0.5) sits on the non-polar/polar cusp; textbooks may classify it either way depending on context.

• Bond polarity ≠ molecular polarity; e.g., \text{CO}_2 has polar C–O bonds (\Delta EN \approx 1.0) but is non-polar overall due to linear geometry and dipole cancellation.