Chemical Bonding Types

Why This Matters

Chemical bonding is the foundation of everything you'll study in chemistry—from predicting molecular shapes to explaining why ice floats or why metals conduct electricity. On the AP exam, you're being tested on your ability to explain why substances behave the way they do, and that always traces back to how their atoms are connected. Understanding bonding types helps you predict physical properties (melting points, conductivity, solubility), molecular geometry, and reactivity patterns.

The key concepts here include electron transfer vs. sharing, electronegativity differences, intermolecular vs. intramolecular forces, and the relationship between bond strength and physical properties. Don't just memorize that ionic compounds have high melting points—know that it's because breaking the electrostatic attractions between ions requires enormous energy. Every bonding type on this list illustrates a principle about how electron behavior determines macroscopic properties.

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Intramolecular Bonds: The Strong Connections

These are the bonds that hold atoms together within molecules or compounds. Breaking intramolecular bonds requires significant energy and constitutes a chemical change. The type of intramolecular bond depends on whether electrons are transferred, shared equally, or shared unequally.

Ionic Bonding

• Electron transfer between metals and nonmetals—metals lose electrons to form cations, nonmetals gain electrons to form anions, creating oppositely charged ions
• Strong electrostatic attraction between ions results in crystal lattice structures with high melting and boiling points
• Conducts electricity only when dissolved or molten because ions must be mobile to carry charge

Covalent Bonding

• Electron sharing between nonmetals—atoms share one or more pairs of electrons to achieve stable electron configurations
• Single, double, or triple bonds form depending on how many electron pairs are shared (1, 2, or 3 respectively)
• Lower melting points than ionic compounds because molecules are held together by weaker intermolecular forces, not the covalent bonds themselves

Polar Covalent Bonding

• Unequal electron sharing occurs when atoms have different electronegativities, creating partial charges ($\delta^+$ and $\delta^-$)
• Dipole moment results from the asymmetric charge distribution—the greater the electronegativity difference, the more polar the bond
• Determines solubility and molecular interactions; polar molecules like $H_2O$ dissolve other polar substances ("like dissolves like")

Metallic Bonding

• Sea of delocalized electrons—valence electrons are shared among all metal atoms rather than localized between specific pairs
• Explains metal properties including electrical conductivity, malleability, and ductility (electrons flow; atoms slide past each other)
• Bond strength varies with the number of valence electrons and ionic radius—more electrons and smaller ions mean stronger bonds and higher melting points

Coordinate Covalent Bonding

• One atom donates both electrons—a Lewis base (electron donor) shares a lone pair with a Lewis acid (electron acceptor)
• Forms complex ions like $[Cu(NH_3)_4]^{2+}$ where ammonia donates electron pairs to copper
• Critical in biochemistry—hemoglobin uses coordinate bonds between iron and oxygen for oxygen transport

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Sigma and Pi Bonds: Orbital Overlap Matters

Understanding how orbitals overlap explains bond strength, rotation, and reactivity. Sigma bonds form the backbone of molecules; pi bonds add rigidity and reactivity.

Sigma Bonds

• Head-on orbital overlap—formed by direct overlap of s-s, s-p, or p-p orbitals along the internuclear axis
• Free rotation is allowed around sigma bonds because electron density is symmetric around the bond axis
• Always the first bond formed between two atoms; every single bond is a sigma bond

Pi Bonds

• Side-by-side p orbital overlap—electron density lies above and below the internuclear axis, not along it
• Restricts rotation because rotating would break the orbital overlap and the bond itself
• Found in multiple bonds—a double bond = 1$\sigma$ + 1$\pi$; a triple bond = 1$\sigma$ + 2$\pi$

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Intermolecular Forces: The Weaker Attractions

These forces act between molecules, not within them. They're much weaker than intramolecular bonds but determine physical properties like boiling point, viscosity, and solubility. Stronger intermolecular forces = higher boiling points.

London Dispersion Forces

• Temporary dipoles from electron fluctuations—even nonpolar molecules experience momentary uneven electron distribution
• Present in ALL molecules—this is the only intermolecular force for nonpolar substances like $Cl_2$ or noble gases
• Strength increases with size—more electrons and larger surface area mean stronger dispersion forces (why $I_2$ is solid but $F_2$ is gas)

Dipole-Dipole Interactions

• Attraction between permanent dipoles—the $\delta^+$ end of one polar molecule attracts the $\delta^-$ end of another
• Stronger than London forces for molecules of similar size, contributing to higher boiling points for polar compounds
• Depends on molecular polarity—greater dipole moment means stronger interactions

Hydrogen Bonding

• Special dipole-dipole interaction—occurs when H is bonded to N, O, or F (highly electronegative atoms with lone pairs)
• Explains water's anomalous properties—high boiling point, high specific heat, ice floating (hydrogen bonds create open lattice structure)
• Essential for biological structures—DNA base pairing, protein folding, and enzyme function all depend on hydrogen bonds

Van der Waals Forces

• Umbrella term for weak IMFs—includes both London dispersion forces and dipole-dipole interactions (sometimes hydrogen bonding)
• Collectively determine physical properties of molecular substances—boiling points, surface tension, viscosity
• Weaker than any intramolecular bond—this is why molecular compounds have lower melting points than ionic or metallic substances

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Quick Reference Table

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Self-Check Questions

1. Which two bonding types both involve electron sharing but differ in how the electrons are contributed? What distinguishes them?

2. A substance 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 state?

3. Compare and contrast London dispersion forces and dipole-dipole interactions. Under what circumstances might a nonpolar molecule have stronger intermolecular forces than a polar molecule?

4. Why does water have a much higher boiling point than hydrogen sulfide ($H_2S$), even though sulfur is larger than oxygen? Which specific intermolecular force explains this?

5. A double bond consists of which types of bonds? Explain why molecules with double bonds cannot rotate freely around that bond, and describe one consequence of this restriction that might appear on an FRQ.