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 DNA holds its shape. When you understand how and why atoms connect, you can predict properties like melting points, conductivity, and solubility without memorizing endless data tables. These concepts show up 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 a lot 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 is whether electrons are transferred, shared, or pooled.

Ionic Bonds

Ionic bonds form when a metal transfers one or more electrons to a nonmetal. The atom that loses electrons becomes a positively charged cation, and the atom that gains electrons becomes a negatively charged anion. These oppositely charged ions attract each other through strong electrostatic forces.

• $Na$ loses one electron to become $Na^+$; $Cl$ gains that electron to become $Cl^-$. The attraction between them forms $NaCl$.
• Ions arrange into a repeating crystalline lattice structure, which is why ionic compounds tend to have high melting and boiling points. Breaking that lattice takes a lot of energy.
• Ionic compounds conduct electricity only when dissolved or melted. In solid form, the ions are locked in place. Once dissolved in water or melted, the ions are free to move and carry charge.

Covalent Bonds

Covalent bonds form when two nonmetals share electron pairs rather than transferring them. This happens because neither atom has a low enough electronegativity to simply give up electrons.

• Atoms can share one, two, or three pairs of electrons, forming single, double, or triple bonds respectively. More shared pairs means a stronger and shorter bond.
• Polarity depends on the electronegativity difference between the bonded atoms. If the atoms pull on the shared electrons unequally, you get a polar covalent bond (like in $H_2O$, where oxygen pulls harder than hydrogen). If they pull equally, you get a nonpolar covalent bond (like in $O_2$, where both atoms are identical).

Metallic Bonds

In metals, valence electrons aren't held by any single atom. Instead, they're delocalized across the entire structure, forming what's often called a "sea of electrons" surrounding positive metal ion cores.

• Electrical conductivity: the delocalized electrons flow freely, carrying charge through the metal.
• Malleability and ductility: when you hammer a metal, layers of ions can slide past each other without breaking bonds, because the electron sea adjusts. Compare this to ionic crystals, where shifting the lattice puts like-charged ions next to each other, causing the crystal to shatter.
• Luster: free electrons absorb incoming light and re-emit it, giving metals their characteristic shine.
• Bond strength increases with more delocalized electrons and smaller ion size. That's why tungsten (many valence electrons, relatively small ions) has an extremely high melting point.

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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 they dramatically affect physical properties like boiling point, surface tension, and solubility.

Hydrogen Bonds

Hydrogen bonds are a special, stronger type of dipole-dipole attraction. They occur when a hydrogen atom bonded to a highly electronegative atom ($N$, $O$, or $F$) is attracted to a lone pair on another electronegative atom in a nearby molecule.

• Water's unusually high boiling point (100°C for such a small molecule) comes from extensive hydrogen bonding between $H_2O$ molecules.
• Ice floats because hydrogen bonds hold water molecules in an open, hexagonal lattice when frozen, making ice less dense than liquid water. This is unusual; most solids are denser than their liquid form.
• In biology, hydrogen bonds hold together DNA's double helix (between base pairs) and help maintain protein shapes. They're strong enough to provide structure but weak enough to be broken and reformed as needed.

Van der Waals Forces

Van der Waals forces arise from temporary, momentary charge imbalances in molecules. Even in nonpolar molecules, electrons are constantly moving, and at any given instant the electron distribution can be slightly uneven, creating a temporary dipole. That temporary dipole can then induce a dipole in a neighboring molecule, creating a brief attraction.

• London dispersion forces are the most basic type and are present in all molecules and atoms, polar or nonpolar. For nonpolar substances, they're the only intermolecular force at work.
• Dipole-dipole interactions occur between polar molecules that have permanent partial charges. These are stronger than London dispersion forces alone.
• Strength increases with molecular size. Larger molecules have more electrons, which means stronger temporary dipoles. This is why propane ($C_3H_8$, boiling point $-42°C$) boils at a higher temperature than methane ($CH_4$, boiling point $-161°C$), even though both are nonpolar.

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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?