Periodic Trends — AP Chemistry
1. Core Concept: Effective Nuclear Charge ★★☆☆☆ ⏱ 3 min
All periodic trends originate from differences in effective nuclear charge, the net positive attraction valence electrons experience from the nucleus, after accounting for shielding by inner core electrons.
Z_{eff} = Z - S
Trends in $Z_{eff}$ are consistent: moving left to right across a period, $Z$ increases by 1 per element while the number of core electrons (and thus $S$) stays constant, so $Z_{eff}$ increases steadily. Moving down a group, $Z$ increases but $S$ increases by a full new shell of electrons, so $Z_{eff}$ is nearly constant with only a very small increase for most groups.
Exam tip: On AP FRQs, you will not earn a justification point just saying "this follows the periodic trend". Always explicitly connect the property to $Z_{eff}$ and shielding to get full credit.
2. Atomic and Ionic Radius ★★★☆☆ ⏱ 4 min
The trend in atomic radius directly follows $Z_{eff}$ and principal quantum number $n$: moving left to right across a period, increasing $Z_{eff}$ pulls valence electrons closer to the nucleus, so atomic radius decreases. Moving down a group, valence electrons occupy higher energy levels with a larger average distance from the nucleus, and this effect dominates over the small $Z_{eff}$ increase, so atomic radius increases down a group.
For ionic radius, additional rules apply: Cations are smaller than their parent neutral atom (lose an entire valence shell, higher $Z_{eff}$ per electron). Anions are larger than their parent neutral atom (add electrons to the valence shell, increasing electron-electron repulsion with no change in $Z$). For isoelectronic ions (ions with the same number of electrons), ionic radius decreases with increasing atomic number, because higher $Z$ gives higher $Z_{eff}$ that pulls the same number of electrons closer.
Exam tip: Always check for isoelectronic ions first when ranking radii. Students often default to the general atomic radius trend and miss that same-electron-count ions follow a different rule.
3. First Ionization Energy ★★★☆☆ ⏱ 4 min
X(g) \rightarrow X^+(g) + e^- \quad \Delta H = IE_1
The general trend follows $Z_{eff}$ and radius: $IE_1$ generally increases left to right across a period (higher $Z_{eff}$ holds electrons tighter, so more energy is needed to remove one) and decreases down a group (valence electrons are farther from the nucleus and easier to remove).
AP Chemistry regularly tests two common exceptions to the general trend: (1) Group 13 elements have lower $IE_1$ than Group 2 elements in the same period, because the electron removed from Group 13 is in a higher-energy p orbital, compared to the lower-energy s orbital valence electrons of Group 2. (2) Group 16 elements have lower $IE_1$ than Group 15 elements in the same period, because Group 15 has a stable half-filled p subshell, while Group 16 has one paired p electron that experiences extra electron-electron repulsion, making it easier to remove.
Exam tip: Always write the valence electron configuration when justifying an ionization energy exception. AP graders explicitly require this connection to electron configuration to award the point.
4. Electron Affinity and Electronegativity ★★★★☆ ⏱ 3 min
X(g) + e^- \rightarrow X^-(g) \quad \Delta H = EA
The general trend for EA is that it becomes more negative (higher affinity) moving left to right across a period, and less negative moving down a group. Common exceptions mirror ionization energy: noble gases have positive EA (adding an electron requires placing it in a new higher energy shell, so the process is unfavorable), Group 2 has less negative EA than Group 1, and Group 15 has less negative EA than Group 14.
Electronegativity follows the same general trend as $IE_1$: EN increases left to right across a period, and decreases down a group. Fluorine is the most electronegative element on the periodic table, and francium is the least. A critical distinction: electron affinity describes isolated gaseous atoms gaining an electron, while electronegativity describes attraction for bonding electrons in a molecule.
Exam tip: Do not confuse electron affinity with electronegativity: mixing these definitions costs points on FRQs.
Common Pitfalls
Why: Students confuse adding valence electrons to the same shell with adding new shells; across a period, shielding is constant so higher $Z_{eff}$ offsets any added electron repulsion
Why: Students assume more protons mean larger size, forgetting the number of electrons is identical
Why: Students think removing a second electron is easier after the first, but the second electron is removed from a positively charged ion that holds electrons tighter
Why: Students memorize the general trend but forget the common orbital-based exceptions that are frequently tested
Why: Students mix up the sign convention: more negative means more energy is released when adding an electron, so higher affinity
Why: Students confuse increasing $Z$ with $Z_{eff}$ for valence electrons; larger distance from the nucleus dominates over the small $Z_{eff}$ increase