Types of Chemical Bonds

Why This Matters

Chemical bonds are the foundation of everything you'll study in chemistry—from why ice floats to how your DNA holds its shape. When you understand how and why atoms connect, you unlock the ability to predict properties like melting points, conductivity, and solubility without memorizing endless data tables. These concepts appear repeatedly in questions about compound behavior, molecular structure, and physical properties.

You're being tested on your ability to distinguish between electron transfer, electron sharing, and intermolecular attractions—and to connect each bonding type to observable properties. Don't just memorize that ionic compounds have high melting points; know why (strong electrostatic forces between ions require lots of energy to overcome). Each bond type illustrates a different principle of atomic interaction, and that's what exam questions are really after.

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Intramolecular Bonds: Holding Atoms Together

These are the strong bonds that form within molecules and compounds, directly connecting atoms to create new substances. The key distinction here is whether electrons are transferred, shared, or pooled.

Ionic Bonds

• Electron transfer between metals and nonmetals—one atom loses electrons, another gains them, creating oppositely charged ions ($Na^+$ and $Cl^-$, for example)
• Strong electrostatic attraction between ions results in crystalline lattice structures with high melting and boiling points
• Conduct electricity only when dissolved or melted—ions must be free to move, which doesn't happen in solid form

Covalent Bonds

• Electron sharing between nonmetals—atoms with similar electronegativities share pairs of electrons rather than transferring them
• Single, double, or triple bonds form depending on how many electron pairs are shared (more sharing = stronger, shorter bonds)
• Polarity varies based on electronegativity difference—unequal sharing creates polar molecules (like $H_2O$), while equal sharing creates nonpolar molecules (like $O_2$)

Metallic Bonds

• "Sea of electrons" model—valence electrons are delocalized and flow freely among positive metal ion cores
• Explains metal properties like electrical conductivity (electrons move freely), malleability (layers slide without breaking bonds), and luster (electrons absorb and re-emit light)
• Bond strength depends on electron count and ion size—more delocalized electrons and smaller ions create stronger metallic bonds

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Intermolecular Forces: Attractions Between Molecules

These are weaker forces that act between separate molecules, not within them. They don't form new compounds but dramatically affect physical properties like boiling point and solubility.

Hydrogen Bonds

• Special dipole-dipole attraction—occurs when hydrogen bonded to $N$, $O$, or $F$ is attracted to another electronegative atom on a nearby molecule
• Explains water's unusual properties—high boiling point, surface tension, and why ice floats (hydrogen bonds create an open lattice structure in solid water)
• Critical for biological structures—holds DNA's double helix together and maintains protein shapes

Van der Waals Forces

• Temporary dipoles from electron fluctuations—even nonpolar molecules experience momentary charge imbalances that create weak attractions
• Includes London dispersion forces (present in all molecules), dipole-dipole interactions, and induced dipole forces
• Strength increases with molecular size—larger molecules have more electrons, creating stronger temporary dipoles (explains why larger hydrocarbons have higher boiling points)

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Quick Reference Table

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Self-Check Questions

1. Which two bond types both involve electrons being held by multiple atoms, but differ in whether the electrons are localized or delocalized?

2. A compound has a high melting point and conducts electricity when dissolved in water but not as a solid. What type of bonding does it have, and why does conductivity depend on its state?

3. Compare and contrast hydrogen bonds and covalent bonds—how do they differ in strength, location (inter- vs. intramolecular), and the role they play in water's properties?

4. Why do larger nonpolar molecules generally have higher boiling points than smaller ones, even though neither has permanent dipoles?

5. If an exam question asks you to explain why metals are malleable but ionic crystals shatter when struck, which bonding concepts would you use in your response?