Periodic Trends

Organization of the Periodic Table

The periodic table is organized by:

• Periods (rows): Horizontal rows numbered 1-7
• Groups (columns): Vertical columns numbered 1-18
• Blocks: s, p, d, f based on electron configuration

Key Regions

• Metals: Left and center (conduct electricity, malleable, lustrous)
• Nonmetals: Right side (poor conductors, brittle as solids)
• Metalloids: Staircase line (properties between metals and nonmetals)

Four Major Periodic Trends

1. Atomic Radius

Definition: Distance from nucleus to outermost electron (half the distance between nuclei of bonded atoms)

Trends:

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

Explanation:

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

| Property | Across Period (→) | Down Group (↓) | |----------|-------------------|----------------| | Atomic Radius | Decreases | Increases | | Ionization Energy | Increases | Decreases | | Electron Affinity | More negative | Less negative | | Electronegativity | Increases | Decreases | | Metallic Character | Decreases | Increases |

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