Chemistry · Unit 2: Molecular and Ionic Compound Structure and Properties · 14 min read · Updated 2026-05-11
Types of chemical bonds — AP Chemistry
AP Chemistry · Unit 2: Molecular and Ionic Compound Structure and Properties · 14 min read
1. Ionic Bonding★★☆☆☆⏱ 5 min
The strength of ionic attraction follows Coulomb's law:
F = k \frac{|q_1 q_2|}{r^2}
where $q_1$ and $q_2$ are the charges of the two ions, $r$ is the distance between the centers of the ions, and $k$ is Coulomb's constant. The AP-accepted threshold for predicting ionic bonding is an electronegativity difference $\Delta \text{EN} > 1.7-2.0$. Ionic compounds form crystalline lattice structures, are hard, brittle, and conduct electricity only when molten or dissolved (not as solids, since ions are fixed in place).
Exam tip: AP will never ask you to memorize electronegativity values — they will always provide any EN values you need for classification. Focus on applying rules, not memorizing numbers.
2. Covalent Bonding (Polar and Nonpolar)★★☆☆☆⏱ 4 min
Covalent bonding forms when two atoms (almost always nonmetals with similar high electronegativities) share pairs of valence electrons to achieve a stable full valence shell. The attraction arises from the shared electron pair being pulled toward the positively charged nuclei of both bonded atoms. Covalent bonds are split into two subclasses based on electronegativity difference:
Nonpolar covalent: $\Delta \text{EN} < 0.5$ — electron density is shared nearly equally between the two atoms, so no permanent partial charge separation
Polar covalent: $0.5 \leq \Delta \text{EN} \leq 2.0$ — the more electronegative atom pulls the shared electron pair closer, creating a partial negative charge ($\delta^-$) on the more electronegative atom and a partial positive charge ($\delta^+$) on the less electronegative atom, forming a bond dipole
A larger $\Delta \text{EN}$ within the covalent range always means a more polar bond.
3. Metallic Bonding★★★☆☆⏱ 5 min
Metals have low ionization energy, so their valence electrons are not tightly held to individual atoms, leading to delocalization. Unlike ionic and covalent bonds (which bind specific pairs of atoms), metallic bonding is non-directional, which explains why metals are malleable (can be hammered into sheets) and ductile (can be drawn into wires): layers of cations can slide past each other without breaking the bonding network. Delocalized mobile electrons also make metallic solids good conductors of electricity and heat in both solid and liquid states. Metallic bonding occurs between atoms of the same pure metal or between different metal atoms in alloys.
Exam tip: AP FRQs always require you to link the microscopic bonding structure to the macroscopic property. Don’t just write "metals conduct electricity" — explicitly mention delocalized mobile electrons to earn full credit.
4. AP Style Practice Problems★★★☆☆⏱ 6 min
Common Pitfalls
Why: Students memorize the "metal + nonmetal = ionic" rule and forget that high oxidation state metals have significant covalent character
Why: The topics are taught back-to-back, so students mix up bond-level vs molecular-level properties
Why: Students forget that interionic distance is in the denominator of the force equation, so they flip the relationship
Why: Students overthink the small electronegativity difference and forget the AP classification threshold
Why: Students oversimplify the structure of metal atoms