Across a period:

• Same energy level, but increasing nuclear charge
• More protons pull electrons closer
• Effective nuclear charge increases

Down a group:

• Each period adds a new energy level (shell)
• Electrons are farther from nucleus
• Shielding from inner electrons

Example: Na > Mg > Al > Si > P > S > Cl (Period 3)

2. Ionization Energy (IE)

Definition: Energy required to remove an electron from a gaseous atom

$X(g) + \text{energy} \rightarrow X^+(g) + e^-$

Trends:

• Across a period (left → right): Increases
• Down a group (top → bottom): Decreases

Explanation:

Across a period:

• Smaller atomic radius
• Stronger attraction between nucleus and electrons
• Harder to remove electron

Down a group:

• Larger atomic radius
• Electrons farther from nucleus
• Easier to remove electron

Exceptions:

• Slight decrease from Group 15 to 16 (half-filled stability)
• Slight decrease from Group 2 to 13 (new sublevel)

Example: He has highest IE; Cs has lowest IE

3. Electron Affinity (EA)

Definition: Energy change when an electron is added to a gaseous atom

$X(g) + e^- \rightarrow X^-(g) + \text{energy}$

Convention: Negative EA means energy is released (favorable)

Trends:

• Across a period (left → right): Generally becomes more negative (more favorable)
• Down a group (top → bottom): Generally becomes less negative

Explanation:

Across a period:

• Smaller atoms can accommodate extra electron more easily
• Higher effective nuclear charge attracts electron

Down a group:

• Larger atoms have more diffuse electron cloud
• Added electron is farther from nucleus

Exceptions:

• Noble gases have positive EA (stable configuration)
• Group 2 and 15 have less negative EA (filled/half-filled stability)

Example: Cl has most negative EA (excluding noble gases)

4. Electronegativity

Definition: Ability of an atom to attract electrons in a chemical bond

Pauling Scale: 0.7 (Cs) to 4.0 (F)

Trends:

• Across a period (left → right): Increases
• Down a group (top → bottom): Decreases

Explanation:

Similar to ionization energy:

• Smaller atoms attract bonding electrons more strongly
• Higher nuclear charge increases pull on electrons

Example: F (4.0) > O (3.5) > N (3.0) > C (2.5)

Note: Noble gases are not assigned electronegativity values (don't form bonds)

Ionic Radius Trends

Cations (Positive Ions)

• Always smaller than parent atom
• Lost electrons → less electron-electron repulsion
• Same nuclear charge pulls fewer electrons

Example: Na (186 pm) > Na⁺ (102 pm)

Anions (Negative Ions)

• Always larger than parent atom
• Gained electrons → more electron-electron repulsion
• Same nuclear charge pulls more electrons

Example: Cl (99 pm) < Cl⁻ (181 pm)

Isoelectronic Series

Atoms/ions with same number of electrons:

Example: O²⁻, F⁻, Ne, Na⁺, Mg²⁺ (all have 10 electrons)

Trend: As nuclear charge increases, radius decreases $O^{2-} > F^- > Ne > Na^+ > Mg^{2+}$

More protons → stronger pull → smaller radius

Successive Ionization Energies

First ionization energy (IE₁): Remove first electron Second ionization energy (IE₂): Remove second electron Third ionization energy (IE₃): Remove third electron, etc.

Trend: IE₁ < IE₂ < IE₃ < ...

Why: Each successive electron is removed from a more positive ion

Big Jump: Large increase when removing electron from inner shell

Example: Magnesium (Mg)

• IE₁ = 738 kJ/mol (remove 3s¹)
• IE₂ = 1451 kJ/mol (remove 3s¹)
• IE₃ = 7733 kJ/mol (HUGE jump - removing from filled 2p⁶)

This confirms Mg forms Mg²⁺, not Mg³⁺

Summary of Trends

Applications

1. Predicting reactivity: Most reactive metals (lower left), most reactive nonmetals (upper right)
2. Bond polarity: Difference in electronegativity determines bond type
3. Ion formation: Elements lose/gain electrons to achieve noble gas configuration
4. Chemical behavior: Trends explain why elements in same group have similar properties

Step 3: Apply the trend

Going down the group (increasing period):

• O is smallest (Period 2)
• S is next (Period 3)
• Se is larger (Period 4)
• Te is largest (Period 5)

Answer: O < S < Se < Te

Explanation: Each successive element has one more electron shell, placing the outermost electrons farther from the nucleus despite increasing nuclear charge. The shielding effect from inner electrons outweighs the increased nuclear charge.

Verification:

• All in same group ✓
• Order follows period numbers ✓
• Radius increases down group ✓