Ionic compounds are held together by electrostatic attractions between oppositely charged ions. Coulomb’s law provides a simple, powerful way to compare how strongly different cation–anion pairs attract based on charge and distance.

Core idea: electrostatic attraction controls interaction strength

Ionic interaction strength refers to how strongly a cation and anion attract each other due to opposite charges. In AP Chemistry, you mainly use Coulomb’s law qualitatively to compare relative strengths, not to compute exact forces.

Key quantities you compare

• Ionic charge: e.g., $+1$, $+2$, $-1$, $-2$

• Interionic distance: approximated by the sum of ionic radii (smaller ions sit closer)

Larger charge magnitude and shorter distance both produce a stronger attraction.

Coulomb’s law applied to ions

In the point-charge model, ions are treated as charged particles whose attraction depends on charge and separation.

This diagram illustrates Coulombic interactions between two point charges separated by a distance $r$, showing both repulsion for like charges and attraction for unlike charges. The equal-and-opposite force vectors emphasize that the interaction strength depends on charge and separation, while the direction depends on the signs of the charges. Source

This radius-ratio diagram models ions as hard spheres and shows how the cation-to-anion size ratio controls whether a given coordination environment is geometrically stable. It helps you connect “smaller ionic radii” to a shorter center-to-center separation $r$, which (via Coulomb’s law) increases the magnitude of electrostatic attraction. Source

Ionic radius is not a directly “visible” boundary; it is a model-based size that helps you reason about distances in ionic solids.

For AP comparisons, focus on proportionality: $F \propto \dfrac{|q_1 q_2|}{r^2}$.

How charge affects attraction (qualitative comparisons)

Increasing the magnitude of ionic charge increases attraction strongly because the product $|q_1 q_2|$ increases.

• If distance is similar, a $2+$ / $2-$ pair attracts more strongly than a $1+$ / $1-$ pair.

• Comparing products:

  • $|(+1)(-1)| = 1$

  • $|(+2)(-1)| = 2$

  • $|(+2)(-2)| = 4$ So, with comparable distances, attraction strength tends to follow these charge products.

What to say in words (expected AP phrasing)

• “Larger ionic charges strengthen interactions because the electrostatic attraction increases with $|q_1 q_2|$.”

How ionic size (distance) affects attraction

Distance matters even more sharply because it appears as $r^2$ in the denominator.

• Smaller ionic radii → smaller $r$ → much stronger attraction.

• For ions with the same charges, the pair involving smaller ions typically has stronger attraction because the ions can approach more closely.

What counts as “smaller ions” in typical AP comparisons

• Higher-charge cations (e.g., $Al^{3+}$ vs $Na^+$) are often smaller due to stronger attraction to electrons.

• Within a group, ions generally get larger down the group, increasing $r$ and weakening attractions (when charges are the same).

What to say in words (expected AP phrasing)

• “Smaller ionic radii (shorter distances) increase interaction strength because $F$ increases as $r$ decreases.”

Using Coulomb’s law to rank interaction strength

When asked to compare two ionic interactions, use a consistent checklist:

• Compare charge product $|q_1 q_2|$ first (bigger product → stronger).

• If charge products are equal or similar, compare distance using ionic radii (smaller $r$ → stronger).

• If one pair has larger charges but also much larger ions, explicitly state both factors and decide which dominates (charge changes often have a large effect, but very large distance can weaken attraction).

Limits of the model (what not to overclaim)

Coulomb’s law assumes point charges and ignores electron-cloud shape and short-range repulsions. It is best used for relative comparisons of ionic attraction, not exact numerical prediction in real crystals or solutions.