Types of Molecular Bonds

Why This Matters

Molecular bonds are the foundation of everything you'll encounter in molecular physics—they determine why water behaves differently than methane, why metals conduct electricity, and why DNA holds its shape. You're being tested on your ability to explain how atoms interact, what forces govern molecular behavior, and why different substances exhibit dramatically different physical properties. The key concepts here include electronegativity differences, electron sharing vs. transfer, and intermolecular vs. intramolecular forces.

Don't fall into the trap of memorizing bond types as isolated definitions. Instead, focus on the underlying mechanisms: What's happening to the electrons? How does this affect macroscopic properties like melting point, conductivity, or solubility? When you understand the why, you can predict behavior for molecules you've never seen before—and that's exactly what FRQs will ask you to do.

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Intramolecular Bonds: What Holds Atoms Together

These are the strong bonds that form molecules themselves. Intramolecular forces involve electron sharing or transfer between atoms and require significant energy to break.

Covalent Bonds

• Electron sharing between atoms—occurs when two atoms (typically nonmetals) have similar electronegativities and neither can "win" the electrons outright
• Bond multiplicity matters: single, double, or triple bonds form depending on how many electron pairs are shared, directly affecting bond strength and length
• Molecular geometry results from covalent bonding patterns, determining properties like polarity and reactivity

Ionic Bonds

• Complete electron transfer creates oppositely charged ions—metals lose electrons (cations), nonmetals gain them (anions)
• Electrostatic attraction between ions produces crystalline lattice structures with characteristically high melting and boiling points
• Conductivity when dissolved or molten because ions become mobile—a key testable property distinguishing ionic from covalent compounds

Coordinate Covalent Bonds

• One atom donates both electrons—also called dative bonds, these form when a Lewis base shares its lone pair with a Lewis acid
• Essential for complex ions like $\text{NH}_4^+$ where ammonia donates electrons to $\text{H}^+$
• Metal coordination compounds rely on these bonds, influencing reactivity in catalysis and biochemistry

Metallic Bonds

• Delocalized "sea" of electrons surrounds positively charged metal nuclei, creating a unique bonding model
• Explains electrical and thermal conductivity—mobile electrons carry charge and energy efficiently through the metal lattice
• Malleability and ductility result because metal layers can slide past each other without breaking bonds—the electron sea simply adjusts

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Intermolecular Forces: What Holds Molecules Together

These weaker forces act between molecules rather than within them. Intermolecular forces determine bulk properties like boiling point, viscosity, and solubility without breaking molecular identity.

Hydrogen Bonds

• Special dipole-dipole interaction occurs when hydrogen bonds to highly electronegative atoms ($\text{O}$, $\text{N}$, or $\text{F}$), creating a strong partial positive charge on hydrogen
• Explains water's anomalous properties—high boiling point, surface tension, and ice floating all result from extensive hydrogen bonding networks
• Critical for biological structure—DNA base pairing and protein folding depend on precise hydrogen bonding patterns

Dipole-Dipole Interactions

• Permanent dipoles attract—the positive end of one polar molecule aligns with the negative end of another
• Requires polar molecules with asymmetric charge distribution due to electronegativity differences within the molecule
• Intermediate strength between hydrogen bonds (stronger) and London dispersion forces (weaker), directly affecting boiling points of polar compounds

London Dispersion Forces

• Temporary dipoles from electron fluctuation—even nonpolar molecules experience momentary asymmetric electron distribution
• Present in ALL molecules but are the only intermolecular force in nonpolar substances like noble gases and hydrocarbons
• Strength scales with polarizability—larger molecules with more electrons have stronger London forces, explaining why larger alkanes have higher boiling points

Van der Waals Forces

• Umbrella term encompassing both dipole-dipole interactions and London dispersion forces—not a separate force type
• Governs gas behavior and deviations from ideal gas law, particularly at high pressures and low temperatures
• Cumulative effect in large molecules stabilizes biological macromolecules like proteins through numerous weak interactions

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Quick Reference Table

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Self-Check Questions

1. Which two bond types both involve electrostatic attraction but differ in whether electrons are localized or delocalized? How does this difference explain their conductivity behavior?

2. A molecule has a high boiling point but doesn't conduct electricity in any state. What type of bonding is most likely responsible, and what intermolecular forces might be present?

3. Compare and contrast hydrogen bonds and standard dipole-dipole interactions. Why does water have such an unusually high boiling point compared to hydrogen sulfide?

4. London dispersion forces are present in all molecules, yet we often ignore them when analyzing polar substances. Under what circumstances would London forces become the dominant factor in determining physical properties?

5. An FRQ presents two compounds: $\text{NaCl}$ and $\text{CCl}_4$. Both contain chlorine, but one dissolves in water while the other doesn't. Using bond type and intermolecular force concepts, explain this difference.