General Chemistry: Chapter 2 - The Periodic Table

2.1 Levels of Organization

Groups

• Vertical columns in the periodic table
• Elements share the same valence shell configuration
  • Group 1: 1 electron in valence shell
  • Group 2: 2 electrons in valence shell

Periods

• Horizontal rows in the periodic table
• Indicate energy levels of valence electrons
  • Period 1: n=1 energy level (e.g., Hydrogen)
  • Period 2: n=2 energy level (e.g., Lithium, Beryllium)

Types of Elements

• Representative Elements: Groups 1, 2, 13-18 (predictable behaviors based on their positions)
  • Example: Sodium and Potassium react similarly with Chlorine
• Non-Representative Elements: Transition metals, lanthanides, and actinides (less predictable behaviors)
  • Example: Transition metals have variable valence electrons

2.2 Characteristics of Element Groups

Metals (yellow on periodic table)

• Properties: Lustrous, malleable, ductile, good conductors of electricity
• Oxidation States: Multiple stable states (up to +7)
• Electron Sea Model: Electrons are shared among atoms, allowing flexibility
• Low Electronegativity: Prefer to lose electrons (form cations)

Non-Metals (blue on periodic table)

• Properties: High ionization energy, high electron affinity, high electronegativity, brittle, poor conductors
• Bonding: Form covalent bonds, rigid structures

Metalloids (green on periodic table)

• Locations: Elements touching the staircase (except Aluminum)
• Properties: Mixed characteristics of metals and non-metals

2.3 Periodic Trends

Effective Nuclear Charge (Z_eff)

• Attraction felt by valence electrons
• Increases from left to right in a period
• Examples:
  • Lithium: Z_eff = +1
  • Beryllium: Z_eff = +2

Noble Gas Configuration

• Noble gases have filled valence shells and are very stable
• Other elements aim to achieve this configuration

Atomic Radii

• Decreases from left to right (increased Z_eff pulls electrons closer)
• Increases from top to bottom (additional energy levels increase size)

Ionic Radii

• Cations: Smaller than their atoms (less repulsion among fewer electrons)
• Anions: Larger than their atoms (more repulsion among more electrons)

Ionization Energy

• Energy required to remove an electron
• Increases from left to right (higher Z_eff makes it harder to remove electrons)
• Decreases from top to bottom (valence electrons are farther from nucleus)

Electron Affinity

• Energy released when an electron is added
• Increases from left to right (elements closer to noble gas configuration release more energy)

Electronegativity

• Tendency to attract shared electrons in a bond
• Increases from left to right
• Example:
  • Sodium Chloride (NaCl): Ionic bond due to large difference in electronegativity
  • Carbon-Chlorine bond: Covalent bond with unequal sharing

2.4 Chemistry of Element Groups

Alkali Metals (Group 1)

• Low density, large atomic radius
• Easily lose one electron to achieve noble gas configuration
• Very reactive, rarely found in elemental form

Alkaline Earth Metals (Group 2)

• Form divalent cations (e.g., Ca²⁺, Be²⁺)
• Also reactive and not found in elemental form

Chalcogens (Group 16)

• Six valence electrons, small atomic radius
• Include important biological elements like Oxygen and Sulfur

Halogens (Group 17)

• Very electronegative, high Z_eff
• Tend to exist as diatomic molecules or in ionic form

Noble Gases (Group 18)

• Inert, very stable
• Low boiling points, gases at room temperature

Transition Metals

• Multiple oxidation states (e.g., Mn can be +2, +3, +4, +6, +7)
• Low electron affinity, do not form anions

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End of Chapter 2 Summary