Periodic Trends Study Pack

Kibin's free study pack on Periodic Trends includes a 3-section study guide, 8 quiz questions, 10 flashcards, and 1 open-ended Explain review question. Sign up free to track your progress toward mastery, plus upload your own notes and recordings to create personalized study packs organized by course.

Last updated May 21, 2026

Topic mastery0%
Save progress →

Periodic Trends Study Guide

Trace the patterns that govern atomic behavior across the periodic table, from atomic radius and ionization energy to electronegativity and metallic character. This pack unpacks how effective nuclear charge and electron shielding drive nearly every major trend, including tricky exceptions in electron affinity at half-filled and fully filled subshells — giving you the conceptual foundation to predict and explain element properties with confidence.

Key Takeaways

  • •Atomic radius generally decreases across a period (left to right) due to increasing nuclear charge pulling electrons closer, and increases down a group as additional electron shells are added.
  • •Ionization energy—the energy required to remove an electron from a neutral atom—follows the opposite trend from atomic radius: it increases across a period and decreases down a group.
  • •Electron affinity, the energy change when a neutral atom gains an electron, tends to become more negative (more favorable) across a period, with notable exceptions at elements that have fully filled or half-filled subshells.
  • •Electronegativity, which measures an atom's pull on shared electrons in a bond, increases across a period and decreases down a group, making fluorine the most electronegative element.
  • •Effective nuclear charge (Zeff) is the net positive charge experienced by valence electrons after accounting for shielding by inner electrons, and it is the underlying driver of most periodic trends.
  • •Metallic character decreases across a period and increases down a group, reflecting the ease with which atoms lose valence electrons.

Effective Nuclear Charge: The Engine Behind Periodic Trends

Most periodic trends can be traced back to one concept: how strongly the nucleus attracts the outermost electrons. Understanding effective nuclear charge and electron shielding explains why properties shift predictably across the periodic table.

Effective Nuclear Charge (Zeff)

  • •Zeff is calculated as the actual nuclear charge (atomic number, Z) minus the shielding constant (S): Zeff = Z − S.
  • •Inner-shell electrons partially cancel the nuclear attraction felt by valence electrons — a phenomenon called electron shielding or screening.
  • •Electrons in the same shell shield each other less effectively than inner-shell electrons do, so adding a proton across a period raises Zeff significantly.

Trends in Zeff Across the Table

  • •Moving left to right across a period, each additional proton increases Z while shielding stays nearly constant, so Zeff rises steadily.
  • •Moving down a group, each new period adds a complete inner shell that shields the valence electrons, keeping Zeff roughly constant even as Z grows substantially.
  • •The net effect is that valence electrons in period 3 sodium (Na) feel a much weaker nuclear pull than those in period 3 chlorine (Cl), despite both being in the same row.

Atomic Radius: Size Patterns Across Periods and Groups

Atomic radius describes the effective size of an atom and is typically measured as half the distance between two identical bonded atoms (covalent radius) or half the distance between adjacent atoms in a metallic lattice (metallic radius). Size changes predictably in response to Zeff and the number of occupied electron shells.

Decrease Across a Period

  • •As Zeff increases from left to right, the nucleus pulls all valence electrons inward more forcefully, shrinking the electron cloud.
  • •Example: lithium (Li, period 2) has a covalent radius of approximately 128 pm, while neon (Ne, period 2) has an atomic radius of about 58 pm.
  • •The number of electron shells does not change within a period, so the only variable is the strength of nuclear attraction.

Increase Down a Group

  • •Each successive period adds a new principal energy level (n = 1, 2, 3…), placing valence electrons in orbitals that are physically farther from the nucleus.
  • •Inner shells also add shielding, further reducing the effective pull on outer electrons.
  • •Example: fluorine (F, period 2) has a covalent radius of about 64 pm, while iodine (I, period 5) has a covalent radius of about 133 pm, despite both being in Group 17.

Ionic Radius Compared to Atomic Radius

  • •Cations (positive ions) are smaller than their parent atoms because removing electrons reduces electron–electron repulsion and increases Zeff per remaining electron.
  • •Anions (negative ions) are larger than their parent atoms because added electrons increase repulsion among valence electrons, expanding the cloud.

About this Study Pack

Created by Kibin to help students review key concepts, prepare for exams, and study more effectively. This Study Pack was checked for accuracy and curriculum alignment using authoritative educational sources. See sources below.

Sources

More in AP Chemistry

See all topics →

Browse other courses

See all courses →