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Chemical Bonding
Ionic, Covalent, and Metallic Bonds
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Chemical bonds form when atoms interact in ways that lower their total energy - almost always by achieving a more stable electron configuration. The three main bond types are ionic (electron transfer between metals and nonmetals), covalent (electron sharing between nonmetals), and metallic (electrons delocalized across a sea of metal cations). Each type produces distinct physical and chemical properties.
Electronegativity difference determines bond type: a large difference (ΔEN > 1.7) produces ionic character; a small difference (ΔEN < 0.5) means nonpolar covalent; intermediate values produce polar covalent bonds. Real bonds exist on a spectrum from pure covalent to pure ionic, not in discrete categories.
Key Facts
- Ionic bonds form between metals and nonmetals; electrons are transferred. Produce crystalline solids with high melting points.
- Covalent bonds form between nonmetals; electrons are shared. Produce molecules with lower melting points.
- Metallic bonds form between metal atoms; electrons are delocalized. Produce conductive, malleable solids.
- Polar covalent: unequal sharing (one atom pulls harder). Creates a partial negative charge (δ-) on the more electronegative atom.
- Lewis structures show valence electrons as dots around atoms and bond pairs as lines.
- VSEPR theory predicts molecular shape based on repulsion of electron pairs.
Vocabulary
- Ionic bond
- A bond formed by the electrostatic attraction between oppositely charged ions (one atom transfers electrons to another).
- Covalent bond
- A bond formed by the sharing of electron pairs between atoms.
- Metallic bond
- A bond in metals formed by a 'sea' of delocalized electrons surrounding positively charged metal cations.
- Electronegativity
- Measure of an atom's ability to attract bonding electrons; drives polarity.
- Polar covalent bond
- A covalent bond in which electrons are shared unequally, creating partial charges on each atom.
Common Mistakes to Avoid
- Thinking ionic compounds are always soluble in water. Many ionic compounds are insoluble - solubility depends on specific lattice and hydration energies.
- Drawing Lewis structures with electrons for inner-shell electrons. Only valence electrons are shown in Lewis structures.
- Assuming all covalent compounds don't conduct electricity. Some polar molecules in solution ionize and conduct. The rule is for pure covalent molecular substances.
- Confusing bond polarity with molecular polarity. A molecule can have polar bonds but be nonpolar overall if the dipoles cancel by symmetry (e.g. CO₂ is linear and nonpolar).
Practice Questions
- 1 Predict the type of bond (ionic, polar covalent, or nonpolar covalent) between: Na and Cl, H and O, C and H.
- 2 Draw the Lewis structure for H₂O. How many bonding pairs and lone pairs does the oxygen atom have?
- 3 Explain why NaCl has a much higher melting point than H₂O using bond type reasoning.