Periodic Table Trends

Why This Matters

Periodic trends aren't just patterns to memorize—they're the foundation for predicting how atoms behave in chemical reactions, how bonds form, and why certain elements are more reactive than others. On the AP Chemistry exam, you're being tested on your ability to explain these trends using core concepts like effective nuclear charge, electron shielding, and atomic structure. Whether you're analyzing why sodium reacts violently with water while potassium reacts even more violently, or predicting which compound will have a more polar bond, these trends give you the tools to reason through problems.

The key insight is that almost every periodic trend traces back to the same fundamental forces: the attraction between positive protons and negative electrons, modified by distance and shielding. Once you understand how effective nuclear charge ($Z_{eff}$) changes across periods and down groups, the rest falls into place. Don't just memorize that electronegativity increases across a period—know why it happens and how it connects to ionization energy, atomic radius, and bonding behavior. That's what earns you points on FRQs.

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Trends Driven by Effective Nuclear Charge

The effective nuclear charge ($Z_{eff}$) is the net positive charge felt by valence electrons after accounting for shielding by inner electrons. As you move across a period, protons are added to the nucleus while electrons enter the same shell, so shielding stays roughly constant but nuclear charge increases. This single concept explains most horizontal trends.

Effective Nuclear Charge

• $Z_{eff}$ increases across a period—more protons pull valence electrons closer, shrinking atomic size and increasing attraction
• Shielding remains approximately constant across a period because electrons add to the same principal energy level
• Calculated using Slater's rules or approximated as $Z_{eff} \approx Z - S$, where $S$ is the shielding constant

Atomic Radius

• Decreases across a period due to increasing $Z_{eff}$ pulling electrons closer to the nucleus
• Increases down a group as new electron shells are added, placing valence electrons farther from the nucleus
• Directly affects ionization energy and electronegativity—smaller atoms hold electrons more tightly

Electronegativity

• Increases across a period because higher $Z_{eff}$ strengthens the nucleus's pull on bonding electrons
• Decreases down a group as greater atomic radius and shielding weaken electron attraction
• Fluorine has the highest electronegativity (3.98)—critical for predicting bond polarity and molecular geometry

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Trends Involving Electron Removal and Addition

These trends describe how easily atoms gain or lose electrons—fundamental for understanding ion formation, redox reactions, and reactivity patterns. The energy required to remove an electron depends on how tightly the nucleus holds it, while the energy released when adding an electron depends on how favorable the resulting configuration is.

Ionization Energy

• Increases across a period as higher $Z_{eff}$ and smaller atomic radius make electrons harder to remove
• Decreases down a group because valence electrons are farther from the nucleus and more shielded
• Successive ionization energies increase dramatically when removing core electrons—watch for large jumps indicating a new shell

Electron Affinity

• Generally becomes more negative (exothermic) across a period—atoms closer to a full octet release more energy when gaining electrons
• Less predictable down a group due to increased electron-electron repulsion in smaller orbitals
• Noble gases have positive electron affinities—their full valence shells make adding electrons unfavorable

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Trends in Ionic Size

When atoms gain or lose electrons, their radii change significantly. Cations shrink because removing electrons reduces electron-electron repulsion and often eliminates an entire shell, while anions expand because added electrons increase repulsion without adding protons.

Ionic Radius

• Cations are smaller than their parent atoms—losing electrons decreases repulsion and often removes the outermost shell entirely
• Anions are larger than their parent atoms—gaining electrons increases electron-electron repulsion while nuclear charge stays constant
• Isoelectronic species (same electron count) decrease in size as nuclear charge increases: $O^{2-} > F^- > Na^+ > Mg^{2+}$

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Trends in Metallic and Nonmetallic Behavior

The periodic table's staircase line separates metals from nonmetals, but this boundary reflects gradual changes in atomic properties. Metallic character correlates with the tendency to lose electrons, while nonmetallic character correlates with the tendency to gain electrons.

Metallic Character

• Increases down a group as larger atomic radius and lower ionization energy make electron loss easier
• Decreases across a period as increasing $Z_{eff}$ makes atoms hold electrons more tightly
• Metals are reducing agents—they donate electrons in redox reactions, which connects to standard reduction potentials in electrochemistry

Reactivity

• Alkali metals increase in reactivity down the group—lower ionization energy means easier electron loss (think Li < Na < K < Rb < Cs)
• Halogens decrease in reactivity down the group—smaller atoms like fluorine attract electrons more strongly than larger iodine
• Reactivity patterns are opposite for metals and nonmetals—metals want to lose electrons while nonmetals want to gain them

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Trends in Electron Configuration

Electron configuration determines which electrons are available for bonding and how atoms interact. The periodic table is organized by electron configuration—each block (s, p, d, f) corresponds to the orbital being filled.

Electron Configuration

• Follows Aufbau principle, Hund's rule, and Pauli exclusion—electrons fill lowest-energy orbitals first, maximize unpaired electrons, and have opposite spins when paired
• Valence electron count determines group chemistry—all alkali metals have $ns^1$, all halogens have $ns^2np^5$
• Anomalies occur in transition metals—chromium is $[Ar]3d^54s^1$ (not $3d^44s^2$) due to half-filled subshell stability

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Physical Property Trends

Melting and boiling points depend on the type and strength of bonding within a substance. These trends are less predictable than electronic trends because they depend on bonding type—metallic, covalent network, or molecular.

Melting and Boiling Points

• Metals generally have high melting points due to strong metallic bonding from delocalized electrons
• Peaks occur at covalent network solids like carbon (graphite/diamond) and silicon in Period 3
• Molecular substances (like $S_8$ and $Cl_2$) have low melting points because only weak intermolecular forces must be overcome

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Quick Reference Table

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Self-Check Questions

1. Why does atomic radius decrease across a period while ionic radius of anions increases compared to neutral atoms? Explain using $Z_{eff}$ and electron repulsion.

2. Arrange the following isoelectronic species in order of increasing ionic radius: $K^+$, $Ca^{2+}$, $Cl^-$, $S^{2-}$. Justify your ranking.

3. Compare and contrast the reactivity trends of Group 1 metals and Group 17 nonmetals. Why do they trend in opposite directions down their respective groups?

4. An element has successive ionization energies (in kJ/mol) of 578, 1817, 2745, 11,578, and 14,831. Identify the group this element belongs to and explain your reasoning.

5. If an FRQ asks you to predict which bond is more polar—$H-F$ or $H-I$—what periodic trend would you use, and how would you explain the difference in terms of atomic structure?