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 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 in a crystal lattice 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 drives ionic bonding. Metals lose valence electrons to form cations, and nonmetals gain electrons to form anions. The oppositely charged ions attract each other through Coulombic (electrostatic) interactions, arranging themselves into a repeating crystal lattice.

• The lattice structure gives ionic compounds high melting and boiling points because you have to overcome many strong ion-ion attractions, not just one bond
• Ionic compounds conduct electricity only when dissolved or molten because the ions must be free to move and carry charge. In the solid lattice, ions are locked in place.
• Lattice energy (a measure of bond strength in ionic compounds) increases with higher ion charges and smaller ionic radii. That's why $MgO$ (2+ and 2- charges) has a much higher melting point than $NaCl$ (1+ and 1-)

Covalent Bonding

Electron sharing between nonmetals is the basis of covalent bonding. Atoms share one or more pairs of electrons to achieve stable electron configurations (typically an octet).

• Single, double, or triple bonds form depending on how many electron pairs are shared (1, 2, or 3 respectively). Bond strength increases and bond length decreases as you go from single to triple.
• Covalent compounds typically have lower melting points than ionic compounds because the molecules themselves are held together by weaker intermolecular forces. The covalent bonds within each molecule are strong, but it's the forces between molecules that you overcome when melting or boiling.

Polar Covalent Bonding

Unequal electron sharing occurs when bonded atoms have different electronegativities. The more electronegative atom pulls electron density toward itself, creating partial charges ($\delta^+$ on the less electronegative atom, $\delta^-$ on the more electronegative one).

• The dipole moment of a bond reflects the magnitude of this charge separation. Greater electronegativity difference means a more polar bond.
• Molecular polarity depends on both bond polarity and molecular geometry. $CO_2$ has polar bonds, but the linear shape causes the dipoles to cancel, making the molecule nonpolar overall.
• Polar molecules like $H_2O$ dissolve other polar substances and ionic compounds ("like dissolves like").

Metallic Bonding

In metals, valence electrons aren't held by individual atoms. Instead, they form a "sea" of delocalized electrons shared among all the metal cations in the structure.

• This model explains electrical and thermal conductivity: delocalized electrons flow freely through the metal when a voltage is applied
• Malleability and ductility make sense too: metal cations can slide past each other without breaking the bond because the electron sea adjusts around them. In an ionic solid, shifting ions would put like charges next to each other, causing the lattice to shatter.
• Bond strength varies with the number of valence electrons and the size of the metal cation. More valence electrons and smaller cations mean stronger metallic bonds and higher melting points (compare tungsten at 3422°C to sodium at 98°C).

Coordinate Covalent Bonding

A coordinate covalent bond forms when one atom donates both electrons in the shared pair. The donor is a Lewis base (has a lone pair), and the acceptor is a Lewis acid (has an empty orbital).

• This is how complex ions form, such as $[Cu(NH_3)_4]^{2+}$, where each ammonia molecule donates a lone pair to the copper ion
• The hydronium ion $H_3O^+$ also contains a coordinate covalent bond: water's oxygen donates a lone pair to $H^+$
• 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 along the internuclear axis, formed by s-s, s-p, or p-p overlap
• Free rotation is allowed around sigma bonds because the electron density is cylindrically 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 creates electron density above and below the internuclear axis, not along it
• Restricts rotation because rotating around the bond axis would break the parallel alignment of the p orbitals, destroying the bond
• Found only 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. The general rule: stronger intermolecular forces = higher boiling points.

London Dispersion Forces (LDFs)

Temporary dipoles from random electron fluctuations cause these forces. At any instant, electrons in a molecule may be unevenly distributed, creating a momentary dipole that induces a dipole in a neighboring molecule.

• Present in ALL molecules and atoms. For nonpolar substances like $Cl_2$, noble gases, or $CH_4$, LDFs are the only intermolecular force.
• Strength increases with molar mass and surface area. More electrons means larger, more frequent temporary dipoles. This is why $I_2$ (large, many electrons) is a solid at room temperature while $F_2$ (small, few electrons) is a gas.
• Molecular shape matters too: long, straight molecules have more surface contact than compact, spherical ones, so they experience stronger LDFs.

Dipole-Dipole Interactions

Permanent dipoles attract each other. The $\delta^+$ end of one polar molecule is attracted to the $\delta^-$ end of another.

• For molecules of similar size, dipole-dipole forces make polar compounds have higher boiling points than nonpolar ones
• Greater dipole moment = stronger interactions
• However, a large nonpolar molecule (with strong LDFs) can have a higher boiling point than a small polar molecule. Size can outweigh polarity.

Hydrogen Bonding

This is a special, unusually strong type of dipole-dipole interaction. It occurs when H is bonded directly to N, O, or F, which are small, highly electronegative atoms with lone pairs.

• The combination of hydrogen's tiny size and the large electronegativity difference creates an exceptionally strong dipole. The exposed $\delta^+$ hydrogen interacts strongly with a lone pair on a neighboring molecule's N, O, or F.
• Explains water's anomalous properties: high boiling point (100°C vs. -60°C for $H_2S$), high specific heat, and the fact that ice floats (hydrogen bonds hold water molecules in an open, less dense lattice structure when frozen)
• Essential for biological structures: DNA base pairing, protein folding, and enzyme-substrate recognition all rely on hydrogen bonds

Van der Waals Forces

This is an umbrella term that includes both London dispersion forces and dipole-dipole interactions. Some sources also include hydrogen bonding under this label, though that varies by textbook.

• Van der Waals forces collectively determine the physical properties of molecular substances: boiling points, surface tension, viscosity
• All van der Waals forces are weaker than any intramolecular bond, which is why molecular compounds generally have lower melting points than ionic or network covalent 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.