SR

8/20 Electronegativity and Bonding – Comprehensive Notes

Electronegativity and Bond Types

  • Electronegativity is the pull an individual atom exerts on its valence (outermost) electrons, not on all electrons. It matters most for the valence shell.
  • A bond always occurs between two atoms. To determine the bond type, compare the difference in electronegativity between the two atoms.
  • The bigger the difference in electronegativity, the more likely the bond is to be ionic; a small to moderate difference leads to polar covalent bonds; essentially no difference leads to nonpolar covalent bonds.
  • The goal is to understand why bonds form the way they do, not just memorize definitions like ionic (I give) vs covalent (I share).

Sodium example: electron configuration and ion formation

  • Sodium has atomic number Z = 11, so it has 11 protons and, in a neutral atom, 11 electrons.
  • Electron shell occupancies:
    • 1st shell max occupancy: 2 electrons
    • 2nd shell max occupancy: 8 electrons
    • Therefore Na has: 2, 8, 1 electrons across shells (2 in the first shell, 8 in the second, and 1 in the valence shell).
  • The valence electron (the lone electron in the outermost shell) is unstable and can be lost to stabilize the atom.
  • Sodium tends to lose that one valence electron rather than gain seven, because its electronegativity is not very high.
  • Loss of the valence electron results in the sodium cation: ext{Na}
    ightarrow ext{Na}^{+} + e^{-}
  • After losing the electron, Na now has 11 protons and 10 electrons, giving a net positive charge (cation).
  • The electron is gained by a more electronegative atom (often a halogen in group 17): the rightmost non-noble gas elements in group 17 are highly electronegative and tend to gain electrons to complete their valence shells.
  • Halogens (e.g., chlorine) have typically one electron short of a full valence shell and will gain an electron:
    • Chlorine gains one electron: ext{Cl} + e^{-}
      ightarrow ext{Cl}^{-}
  • Chloride now has 17 protons and 18 electrons, giving a negative charge. This creates the ionic pair Na^+ and Cl^-.
  • The attraction between the oppositely charged ions (Na^+ and Cl^-) forms the ionic bond.
  • Important note: ionic bonds are very strong in the solid state, but they can break (dissolve) when placed in water or other solvents that can stabilize the ions.

Dissolution vs dissociation in water

  • Dissolve: a solute becomes surrounded by solvent molecules; the solute is dispersed within the solvent to form a solution. In this context, water acts as the solvent and surrounds the ions.
    • Example: NaCl dissolves in water because water molecules orient themselves around Na^+ and Cl^- ions (hydration).
  • Dissociate: the ionic compound splits into its constituent ions in solution. For NaCl in water, the solid dissociates into Na^+ and Cl^- ions:
    • NaCl(s) → Na^+(aq) + Cl^-(aq)
  • Glucose, by contrast, does not have ionic bonds (it is covalently bonded) and does not dissociate into ions in water.
  • Summary:
    • Dissolve = the solute is surrounded by solvent; can lead to a solution, but not necessarily to ions.
    • Dissociate = the compound splits into ions in solution; occurs for substances with ionic bonds.

No difference in electronegativity: nonpolar covalent bonds

  • When two atoms have no difference in electronegativity, electrons are shared equally; this forms a nonpolar covalent bond.
  • Examples where atoms are the same or have nearly identical electronegativities:
    • H–H, O–O, N–N (single bonds), also possible for double bonds (O=O, N=N) and triple bonds (N≡N).
  • Rule of thumb: equal electronegativities → no net charge on either atom; both remain neutral.
  • An exception often discussed is carbon in C–H bonds: carbon and hydrogen have similar electronegativities, so the C–H bond is often treated as nonpolar covalent, even though the atoms are different.
  • Visualizing the tug-of-war: neither atom gains nor loses electrons; electrons remain with their original atoms.

Polar covalent bonds: small to moderate electronegativity differences

  • When there is a small to medium difference in electronegativity, electrons are pulled more toward one atom but not fully transferred; this creates a polar covalent bond with partial charges.
  • Four key polar covalent bonds (ordered by the strength of the differential, i.e., how strongly electrons are pulled):
    • O–H, N–H, O–C, N–C
  • Rationale: Oxygen and Nitrogen are more electronegative than Hydrogen and Carbon, so O and N pull electrons more strongly.
  • In biochemistry and many contexts (ignoring fluorine for a moment), oxygen is the most electronegative element among common biochemistry-relevant elements, so it tends to attract electrons from H or C.
  • Water as a primary example:
    • Water molecule: ext{H}_2 ext{O} consists of two O–H bonds.
    • Oxygen exerts a stronger pull on the shared electrons than hydrogen, but not to the extent of completely removing electrons from hydrogen.
    • This results in partial charges:
    • Oxygen becomes partial negative: ext{O}^{ ext{δ}-}
    • Each hydrogen becomes partial positive: ext{H}^{ ext{δ}+}
    • The polarity of the O–H bonds makes water a polar molecule.
  • Key takeaway for polarity:
    • Polar covalent bonds have partial charges, leading to dipole moments and intermolecular interactions such as hydrogen bonding and dipole-dipole interactions.
    • Nonpolar covalent bonds have no significant partial charges and weaker dipole interactions.

Recap: connections to broader concepts

  • Electronegativity differences explain bond formation, bond strength, and chemical behavior (ionic vs covalent vs polar covalent).
  • Ionic bonds arise from complete electron transfer and electrostatic attraction between ions; they are strong in solids but dissociate in polar solvents like water.
  • Covalent bonds result from electron sharing; the degree of sharing determines polarity (nonpolar vs polar).
  • Molecular polarity influences solubility, boiling/melting points, and interaction with other molecules (e.g., water as a universal solvent due to its polarity).

Transition to next topic

  • The transcript transitions from intramolecular bonds (within molecules and atoms) to intermolecular bonds (between molecules). Intermolecular interactions include forces like hydrogen bonding, dipole-dipole interactions, and London dispersion forces, which govern properties such as boiling points and solubility of compounds in water.

Real-world relevance and quick references

  • Ionic compounds in water: dissolution and dissociation underlie most biological salt transport and many physical processes (salts in physiology, ocean chemistry).
  • Polar covalent bonds explain why water is a good solvent for many ionic and polar substances and why many biological macromolecules fold and interact the way they do.
  • Nonpolar covalent substances (like O2, N2, and H2) are typically insoluble in water, reflecting a lack of favorable dipole interactions with water.
  • Notation reminders:
    • Ion formation example: ext{Na}
      ightarrow ext{Na}^{+} + e^{-}; ext{Cl} + e^{-}
      ightarrow ext{Cl}^{-}
    • Ionic compound in water: ext{NaCl(s)}
      ightarrow ext{Na}^{+} (aq) + ext{Cl}^{-} (aq)
    • Water formula: ext{H}_{2} ext{O}
    • Partial charges in H2O: ext{O}^{ ext{δ}-}, ext{H}^{ ext{δ}+}