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. The core skill here is explaining how atoms interact, what forces govern molecular behavior, and why different substances have dramatically different physical properties.

The key concepts are electronegativity differences, electron sharing vs. transfer, and intermolecular vs. intramolecular forces. Don't fall into the trap of memorizing bond types as isolated definitions. 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.

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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

Covalent bonds form when two atoms (typically nonmetals) have similar electronegativities, so neither can pull the electrons away entirely. Instead, they share electron pairs.

• Nonpolar covalent bonds occur when the sharing is roughly equal (electronegativity difference $\approx 0$), as in $\text{H}_2$ or $\text{Cl}_2$. Polar covalent bonds occur when the sharing is unequal (electronegativity difference roughly 0.4–1.7), creating partial charges, as in $\text{H-Cl}$.
• Bond multiplicity matters: single, double, or triple bonds form depending on how many electron pairs are shared. More shared pairs means a shorter, stronger bond. For example, the triple bond in $\text{N}_2$ (bond energy ~945 kJ/mol) is much stronger than the single bond in $\text{F}_2$ (~155 kJ/mol).
• Molecular geometry results from covalent bonding patterns and lone pairs, which in turn determines polarity and reactivity.

Ionic Bonds

Ionic bonds form when there's a large electronegativity difference (typically >1.7) between two atoms, resulting in complete electron transfer rather than sharing.

• Metals lose electrons to become cations; nonmetals gain electrons to become anions. For example, in $\text{NaCl}$, sodium transfers one electron to chlorine.
• The electrostatic attraction between oppositely charged ions produces crystalline lattice structures with characteristically high melting points ($\text{NaCl}$ melts at 801°C).
• Ionic compounds conduct electricity when dissolved or molten because the ions become mobile. In solid form, the ions are locked in place and can't carry charge. This is a key testable distinction from covalent compounds.

Coordinate Covalent Bonds

Sometimes called dative bonds, these are a special case of covalent bonding where one atom donates both electrons in the shared pair. A Lewis base (the donor) shares its lone pair with a Lewis acid (the acceptor).

• A classic example is $\text{NH}_4^+$: ammonia ($\text{NH}_3$) donates its lone pair to $\text{H}^+$, which has an empty orbital. Once formed, the coordinate covalent bond is indistinguishable from any other covalent bond in the molecule.
• Metal coordination compounds (like $[\text{Cu(NH}_3)_4]^{2+}$) rely heavily on these bonds, with ligands donating electron pairs to a central metal ion. This is important in catalysis and biochemistry.

Metallic Bonds

In metals, valence electrons aren't bound to individual atoms. Instead, a delocalized "sea" of electrons surrounds positively charged metal nuclei (cations arranged in a lattice).

• This model explains electrical and thermal conductivity: mobile electrons carry charge and kinetic energy efficiently through the lattice.
• Malleability and ductility result because metal layers can slide past each other without breaking bonds. The electron sea simply adjusts to the new arrangement, unlike an ionic lattice where displacing ions brings like charges together and shatters the crystal.

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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. When you boil water, you're overcoming intermolecular forces (the molecules separate), not breaking covalent bonds (the $\text{O-H}$ bonds within each molecule stay intact).

Hydrogen Bonds

A hydrogen bond is a special, unusually strong dipole-dipole interaction. It occurs when hydrogen is covalently bonded to a highly electronegative atom ($\text{O}$, $\text{N}$, or $\text{F}$), which pulls electron density away from hydrogen, leaving it with a strong partial positive charge ($\delta^+$). This exposed proton then attracts a lone pair on a nearby electronegative atom.

• Hydrogen bonds are typically 5–30 kJ/mol, much weaker than covalent bonds (~200–800 kJ/mol) but significantly stronger than other intermolecular forces.
• They explain water's anomalous properties: high boiling point (100°C for such a small molecule), high surface tension, and the fact that ice floats (the hydrogen bonding network in ice creates an open lattice less dense than liquid water).
• Biological structures depend on them: DNA base pairing (A-T has 2 hydrogen bonds, G-C has 3) and protein folding both rely on precise hydrogen bonding patterns.

Dipole-Dipole Interactions

These occur between polar molecules with permanent dipoles. The positive end ($\delta^+$) of one molecule attracts the negative end ($\delta^-$) of another.

• They require molecules with asymmetric charge distribution due to electronegativity differences and molecular geometry. A molecule like $\text{CO}_2$ is nonpolar despite having polar bonds, because its linear shape makes the dipoles cancel.
• Their strength is intermediate between hydrogen bonds and London dispersion forces, and they directly raise boiling points of polar compounds relative to nonpolar ones of similar size.

London Dispersion Forces

Even in nonpolar molecules, electrons are constantly moving. At any instant, the electron cloud can be slightly asymmetric, creating a temporary (instantaneous) dipole. This temporary dipole induces a dipole in a neighboring molecule, and the two attract each other briefly.

• London dispersion forces are present in ALL molecules, but they're the only intermolecular force in nonpolar substances like noble gases and hydrocarbons.
• Strength scales with polarizability: larger molecules with more electrons have bigger, more easily distorted electron clouds, producing stronger London forces. This is why boiling points increase down the noble gases (He: -269°C, Ne: -246°C, Ar: -186°C, Kr: -152°C) and why longer-chain alkanes have higher boiling points than shorter ones.

Van der Waals Forces

This is an umbrella term encompassing both dipole-dipole interactions and London dispersion forces. It's not a separate force type. You'll sometimes see it used loosely to mean just London forces, but strictly it covers all non-covalent, non-ionic intermolecular attractions (excluding hydrogen bonds in some conventions, though definitions vary by textbook).

• Van der Waals forces govern deviations from ideal gas behavior, particularly at high pressures and low temperatures where molecules are close together and these forces become significant.
• In large biological macromolecules like proteins, the cumulative effect of many weak Van der Waals interactions across a large surface area can provide substantial stabilization.

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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.