Chemical Bond Types

Why This Matters

Chemical bonding is the foundation of chemistry itself—it explains why matter behaves the way it does, from why salt dissolves in water to why DNA holds genetic information. When you're tested on bonding, you're really being tested on your understanding of electronegativity, electron behavior, and energy stability. These concepts show up everywhere: in predicting molecular properties, explaining phase changes, and understanding biological structures.

Don't just memorize that ionic bonds involve "electron transfer"—know why that transfer happens (electronegativity differences) and what it means for the compound's properties. The same principle applies to every bond type: the mechanism determines the properties. Master that connection, and you'll handle any question they throw at you.

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Intramolecular Bonds: The Strong Stuff

These are the bonds that hold atoms together within a molecule or compound. Breaking them requires significant energy, which is why they determine fundamental properties like melting points and chemical stability.

Ionic Bonds

• Electron transfer between atoms—occurs when electronegativity difference is large (typically > 1.7), creating oppositely charged ions that attract
• Metals + nonmetals form ionic compounds; metals lose electrons to become cations, nonmetals gain electrons to become anions
• High melting points and electrical conductivity when dissolved—the strong electrostatic attraction requires lots of energy to overcome, but mobile ions carry charge in solution

Covalent Bonds

• Electron sharing between atoms—occurs when electronegativity values are similar, allowing atoms to achieve stable electron configurations together
• Single, double, and triple bonds reflect how many electron pairs are shared; more sharing means shorter, stronger bonds
• Lower melting points and no conductivity—molecules are neutral with no free charges, so they don't conduct electricity in any phase

Metallic Bonds

• Delocalized "sea of electrons"—metal atoms release valence electrons into a shared pool that flows freely around positive metal ions
• Conductivity, malleability, and luster all result from electron mobility; electrons carry current, allow layers to slide, and interact with light
• Variable melting points—depends on number of valence electrons and ion size; more electrons in the "sea" generally means stronger bonding

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Intermolecular Forces: The Weak (But Essential) Stuff

These forces act between molecules rather than within them. They're much weaker than intramolecular bonds, but they determine crucial physical properties like boiling point, viscosity, and solubility.

Hydrogen Bonds

• Attraction between H and electronegative atoms (F, O, N)—hydrogen bonded to one electronegative atom is attracted to a lone pair on another nearby electronegative atom
• Explains water's unique properties—high boiling point, surface tension, and ice floating all result from extensive hydrogen bonding networks
• Critical for biological molecules—holds DNA strands together and maintains protein secondary structure (alpha helices and beta sheets)

Van der Waals Forces

• Temporary dipoles from electron movement—even nonpolar molecules experience momentary charge imbalances that induce attractions
• Three types: London dispersion, dipole-dipole, and dipole-induced dipole—London forces exist in all molecules; the others require permanent dipoles
• Strength increases with molecular size—larger molecules have more electrons, creating stronger temporary dipoles; this explains why larger hydrocarbons have higher boiling points

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Bond Strength Hierarchy

Understanding relative strength is essential for predicting properties and answering comparison questions.

The Strength Spectrum

• Covalent and ionic bonds are strongest—breaking them requires chemical reactions; triple covalent bonds (like in $N_2$) are especially strong
• Metallic bond strength varies—tungsten has extremely strong metallic bonds (high melting point), while mercury's are weak (liquid at room temperature)
• Intermolecular forces are weakest—hydrogen bonds are strongest among them, followed by dipole-dipole, then London dispersion forces

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Self-Check Questions

1. Which two bond types both result in high melting points but differ in their electrical conductivity as solids? Explain why.

2. A molecule has a boiling point much higher than expected for its molecular weight. Which intermolecular force is most likely responsible, and what structural feature would you look for?

3. Compare and contrast how electrons behave in covalent bonds versus metallic bonds. How does this difference explain why metals conduct electricity but most covalent compounds don't?

4. If an FRQ asks you to explain why ice floats on liquid water, which bond type is central to your answer? What specific property of that bond type causes this phenomenon?

5. Arrange these in order of strength: hydrogen bond, ionic bond, London dispersion force, triple covalent bond. For each, give one example substance where that interaction is the dominant force.