Topic 1.7 Notes – Periodic Trends

1. How the Periodic Table Reflects Electron Structure

The periodic table is arranged by increasing atomic number (Z).
$Z =$ number of protons = nuclear charge.

Periods (rows)

• Same number of occupied energy levels (shells).
• Moving left → right:
  • Protons increase.
  • Electrons are added to the same principal energy level.

So across a period, the valence electrons stay in the same shell with similar core shielding, but nuclear charge increases, so they are pulled in closer.

Groups (columns)

• Same number of valence electrons.
• Similar valence electron configurations → similar chemistry.
• Moving down:
  • A new shell is added each step.

Here’s a standard periodic table labeled by periods and groups:

Periodic table showing groups (1-18) and periods

The repeating pattern of valence configurations explains why properties repeat.
Completely filled or half-filled subshells (like $p^{3}$ or $p^{6}$) add extra stability. That’s why a few ionization energy and electron affinity values don’t follow the smooth trend perfectly.

2. Coulomb’s Law, Shielding, and Effective Nuclear Charge

All periodic trends come from electrostatic attraction.

Coulomb’s Law (qualitative version)

Attraction increases with:

• Higher charge (more protons)
• Shorter distance between charges

Shielding

• Inner (core) electrons repel valence electrons.
• They block some of the nucleus’s pull.

Effective Nuclear Charge ($Z_{\text{eff}}$)

This is the net positive charge felt by valence electrons.

• Across a period → $Z_{\text{eff}}$ increases
  (more protons, similar shielding)
• Down a group → roughly similar $Z_{\text{eff}}$ for valence electrons
  (more protons but also more shielding)

Shortcut logic

• Across → stronger pull.
• Down → greater distance + more shielding → weaker pull.

When you explain a trend using charge + distance, that’s AP-level reasoning.

3. The Four Major Periodic Trends

Let’s connect each trend to electron structure.

A. Atomic Radius

Distance from nucleus to outermost electrons.

• Across → decreases
  • $Z_{\text{eff}}$ increases → electrons pulled closer.
• Down → increases
  • More shells → greater distance and shielding.

Largest atoms are bottom left.
Smallest are top right.

B. Ionic Radius

Compare ion vs. neutral atom.

Cations (positive ions)

• Formed by losing electrons.
• Smaller than parent atom
  • Less repulsion.
  • Sometimes entire valence shell removed.

Example: $\ce{Na} \rightarrow \ce{Na^{+}}$ loses its 3rd shell.

Anions (negative ions)

• Formed by gaining electrons.
• Larger than parent atom
  • More electron-electron repulsion.

Isoelectronic Series

Same number of electrons, different protons.

More protons pulling on the same electrons → smaller radius.

Isoelectronic comparisons show up often in multiple choice.

C. Ionization Energy (IE)

Energy required to remove an electron from a gaseous atom.

• Across → increases
  • Smaller radius + stronger pull.
• Down → decreases
  • Electrons farther away.

Successive Ionization Energies

Each one is larger than the previous.

If data looks like this:

IE₁ = 600
IE₂ = 1200
IE₃ = 7500

The huge jump after IE₂ means the third electron removed was a core electron.
So the atom had 2 valence electrons.

That jump question appears constantly on quizzes and FRQs.

Small Exceptions

• Group 13 slightly lower than Group 2 (removing a p electron vs s).
• Group 16 slightly lower than Group 15 (paired electron repulsion).

You don’t need to write exception electron configurations on the AP exam.

D. Electron Affinity (EA)

Energy change when an atom gains an electron.

• Across → more negative
  • Stronger attraction for added electron.
• Down → less negative
  • More distance and shielding.

Halogens are very negative.
Chlorine is slightly more negative than fluorine because fluorine is so small that electron-electron repulsion matters more.

E. Electronegativity (EN)

Ability of an atom in a bond to attract shared electrons.

• Across → increases
• Down → decreases
• Highest: Fluorine

Same reasoning as IE and EA. Small radius + high $Z_{\text{eff}}$.

4. Using Periodicity to Predict Properties

If you don’t remember an exact value, use position.

Across a period:

• Radius ↓
• IE ↑
• EA more negative
• EN ↑

Down a group:

• Radius ↑
• IE ↓
• EA less negative
• EN ↓

Bottom left elements are large, low IE, low EN, metallic.
Top right elements are small, high IE, high EN, nonmetallic.