AP Chem Periodic Trends to Know

What You Need to Know

Periodic trends are the predictable patterns in atomic/ionic properties across the periodic table. AP Chem loves them because they test whether you can connect electron configuration + Coulombic attraction to real chemical behavior (bonding, reactivity, acidity, etc.).

The big idea (the “why” behind every trend)

Most periodic trends can be explained by two competing effects:

• Effective nuclear charge: Z_{\text{eff}} = Z - S

  • Z = nuclear charge (protons)
  • S = shielding by inner electrons
  • Higher Z_{\text{eff}} means valence electrons feel more pull toward the nucleus.

• Distance (principal energy level): As n increases (going down a group), valence electrons are farther from the nucleus and more shielded.

Underlying physics (qualitative):

• Coulombic attraction increases when charges are larger and distance is smaller: F \propto \frac{q_1 q_2}{r^2}
• Stability of attraction increases as particles get closer / charges increase: U \propto -\frac{q_1 q_2}{r}

If you can say “Across a period, Z_{\text{eff}} increases; down a group, n and shielding increase,” you can justify almost every ranking question.

Trends you’re expected to know cold

• Atomic radius
• Ionic radius (including isoelectronic series)
• Ionization energy (especially 1st IE + key exceptions)
• Electron affinity (general trend + exceptions)
• Electronegativity
• Related “applications” trends: metallic character, reactivity, bond polarity, oxide acidity/basicity

Step-by-Step Breakdown

Use this quick procedure for any periodic trend comparison/ranking problem.

A. Ranking atomic size (atomic radius)

1. Locate elements (same period or same group?).
2. If same period: size decreases left → right (higher Z_{\text{eff}} pulls electrons in).
3. If same group: size increases top → bottom (higher n dominates).
4. If diagonal / mixed, decide which effect dominates:
   • Going down usually increases size strongly.
   • Going right usually decreases size moderately.

B. Ranking ionic size (ionic radius)

1. Identify whether each ion is a cation or anion.
   • Cations are smaller than their neutral atoms.
   • Anions are larger than their neutral atoms.
2. Check for an isoelectronic set (same number of electrons).
   • If isoelectronic: more protons ⇒ smaller ion.
3. If not isoelectronic, fall back to: higher n (down group) tends to mean larger radius, and higher positive charge tends to shrink.

C. Ranking ionization energy (IE)

1. Make sure you know which IE: 1st, 2nd, 3rd…
2. For 1st IE:
   • Increases left → right (higher Z_{\text{eff}})
   • Decreases top → bottom (higher n and shielding)
3. Check for classic exceptions (see below).
4. For successive IE:
   • Expect a huge jump when you start removing a core electron (after all valence electrons are gone).

D. Ranking electron affinity (EA)

1. Remember the sign convention: EA is energy change when an atom gains an electron.
   • More favorable additions are typically more negative (more energy released).
2. Trend: generally becomes more negative left → right, less negative down a group.
3. Check exceptions (Groups 2, 15, 18).

E. Ranking electronegativity (EN) and bond polarity

1. EN increases left → right and decreases top → bottom.
2. Larger EN difference \Delta EN ⇒ more polar bond.
3. Highest EN element: F.

Mini worked “decision” examples

• Isoelectronic ions: O^{2-}, F^-, Na^+, Mg^{2+} all have 10 e^-. More protons means smaller, so: Mg^{2+} < Na^+ < F^- < O^{2-} (smallest → largest).
• Successive IE jump: If an element shows a massive jump between IE_2 and IE_3, it likely has 2 valence electrons (after 2 removals, you hit core).

Key Formulas, Rules & Facts

Core drivers (use to justify trends)

Directional trends (the “arrows”)

Atomic radius: what to remember

• Largest atoms: bottom-left (Cs/Fr region).
• Smallest atoms: top-right (He is smallest by radius definition issues; among bonding atoms, think F/Ne region).

Ionic radius: key rules

• Cations < neutral atom (lost e^-, sometimes lost an entire shell).
• Anions > neutral atom (added e^- increases repulsion).
• Isoelectronic series (same e^- count): radius decreases as protons increase.

Ionization energy (IE): key facts + the two classic exceptions

General trend: up and to the right.

Exception 1: Group 13 vs Group 2 (B < Be, Al < Mg)

• Removing from p orbital (Group 13) is easier than from filled s orbital (Group 2).
• So IE drops from Group 2 to Group 13 in the same period.

Exception 2: Group 16 vs Group 15 (O < N, S < P)

• Group 15 has half-filled p (more stable).
• Group 16 has one paired p electron → extra repulsion → easier to remove.

Successive IE rule:

• Big jump occurs after removing all valence electrons.
• Use jumps to infer valence electron count (and likely group).

Electron affinity (EA): what AP expects

• Halogens tend to have very negative EA (they “want” an electron).
• EA becomes less negative going down a group (larger atoms, added e^- farther away).

Commonly tested exceptions:

• Group 18 (noble gases): EA is near 0 or positive (adding e^- goes to a new shell).
• Group 2 (alkaline earths): relatively small/less negative EA (would add e^- to higher-energy p subshell).
• Group 15: less negative EA than you might expect (half-filled p is stable).

Subtle point that shows up: Cl has a slightly more negative EA than F because F is so small that e^-–e^- repulsions are stronger in its compact 2p orbital.

Electronegativity (EN)

• Increases toward F (highest).
• Decreases down a group.
• Used for:
  • bond polarity (bigger \Delta EN ⇒ more polar)
  • predicting partial charges and intermolecular forces trends.

Reactivity trends (quick, AP-style)

Oxides: acidity/basicity trend (periodic trend application)

Across a period, oxides generally go:

• basic (metal oxides) → amphoteric → acidic (nonmetal oxides)

Reason: increasing EN makes bonds more covalent and favors formation of oxyacids / acidic behavior.

Examples you should recognize:

• Na_2O: strongly basic
• Al_2O_3: often treated as amphoteric
• SO_3, CO_2: acidic oxides

Examples & Applications

Example 1: Rank atomic radius

Rank (largest → smallest): Na, Mg, Al, Si (all Period 3).

• Across Period 3, Z_{\text{eff}} increases → radius decreases.
• Answer: Na > Mg > Al > Si.

Example 2: Isoelectronic ionic radii

Rank (smallest → largest): O^{2-}, F^-, Na^+, Mg^{2+}.

• All have 10 e^-.
• More protons pulls same electron cloud in tighter.
• Proton counts: O (8) < F (9) < Na (11) < Mg (12).
• Answer (smallest → largest): Mg^{2+} < Na^+ < F^- < O^{2-}.

Example 3: First ionization energy with exceptions

Which has the higher 1st IE: Al or Mg?

• Naively “to the right is higher,” but here you cross from Group 2 to Group 13.
• Mg has a filled 3s subshell; Al’s electron removed is a 3p electron (higher energy, easier to remove).
• Answer: Mg has higher 1st IE than Al.

Which has the higher 1st IE: P or S?

• Across a period would suggest S higher, but Group 16 has paired p electron repulsion.
• Answer: P has higher 1st IE than S.

Example 4: Bond polarity + EN trend

Which bond is more polar: H–Cl or H–Br?

• EN decreases down the halogens: Cl is more EN than Br.
• Larger \Delta EN for H–Cl.
• Answer: H–Cl is more polar.

How it shows up on AP: they might tie this to stronger dipole-dipole forces, boiling point, or acid strength trends (e.g., HCl vs HBr vs HI).

Common Mistakes & Traps

1. Mixing up “atomic radius” vs “ionic radius”
   Students apply atomic trends to ions directly. Ions change size dramatically when electrons are gained/lost. Fix: first decide cation/anion and check isoelectronic sets.

2. Forgetting isoelectronic logic (protons win)
   Students rank ions by position on the periodic table instead of proton count. In an isoelectronic series, more protons always means smaller radius. Fix: count electrons; if equal, compare Z.

3. Assuming ionization energy always increases left → right with no exceptions
   The Be/B and N/O style exceptions are AP favorites. Fix: check subshell: removing from a p electron is easier; paired p electrons add repulsion.

4. Misreading electron affinity sign
   Many forget “more negative” usually means “more favorable.” Fix: interpret EA qualitatively: “How badly does it want an electron?” (halogens want it most).

5. Thinking shielding increases across a period like it does down a group
   Across a period, added electrons go to the same shell; shielding doesn’t increase much, so Z_{\text{eff}} rises. **Fix**: across = same n; down = new shell.

6. Overgeneralizing halogen EA: assuming F is always most negative
   F is extremely EN, but EA can be slightly less negative than Cl due to electron-electron repulsions in small 2p. Fix: remember the classic Cl vs F nuance.

7. Confusing electronegativity with electron affinity
   EN is about attraction for shared electrons in a bond; EA is an energy change for adding an electron to a gaseous atom. Fix: EN = bonds; EA = isolated atom gaining e^-.

8. Using “top-right” rules on noble gases without thinking
   Noble gases don’t “want” electrons (EA near 0) and rarely form bonds (EN often not assigned). Fix: treat Group 18 as special.

Memory Aids & Quick Tricks

Quick Review Checklist

• [ ] You can explain trends using Z_{\text{eff}} and shielding/n.
• [ ] Atomic radius: decreases across, increases down.
• [ ] Ionic radius: cations smaller, anions larger; isoelectronic: more protons → smaller.
• [ ] 1st IE: increases across, decreases down, with Group 13 < Group 2 and Group 16 < Group 15 exceptions.
• [ ] Successive IE: identify the big jump to find valence electron count.
• [ ] EA: generally more negative across, less negative down; know Groups 2, 15, 18 exceptions and Cl vs F nuance.
• [ ] EN: increases toward F; use \Delta EN to compare bond polarity.
• [ ] Metals most reactive down-left; nonmetals most reactive up-right (excluding noble gases).
• [ ] Oxides trend across a period: basic → amphoteric → acidic.

You’ve got this—if you justify with Z_{\text{eff}} and electron configuration, the trends questions become predictable.