General Principles

Group number relates to the number of valence electrons:

• s-block: Group number = number of valence electrons
• p-block: Number of valence electrons = Group number - 10
• Elements tend to achieve stable electron configurations (like noble gases)

Ion Formation Trends

• 1-3 valence electrons: Tend to lose electrons, form +1, +2, +3 ions
• 5-7 valence electrons: Tend to gain electrons, form -3, -2, -1 ions
• 4 valence electrons: Can form +4 or -4 ions

Group 1 (Alkali Metals)

• One valence electron
• Form +1 ions (monovalent cations)
• Examples: Li⁺, Na⁺, K⁺

Group 2 (Alkaline Earth Metals)

• Two valence electrons
• Form +2 ions (divalent cations)
• Examples: Be²⁺, Mg²⁺, Ca²⁺

Group 17 (Halogens)

• Seven valence electrons
• Gain one electron to form -1 ions (monovalent anions)
• Examples: F^-, Cl^-, Br^-

Group 16 (Chalcogens)

• Six valence electrons
• Gain two electrons to form -2 ions (divalent anions)
• Examples: O^2-, S^2-

Group 18 (Noble Gases)

Full valence electron shells (except He). Chemically stable, do not typically form ions.

Determining Element Position from Electronic Configuration

• Write electronic configuration
• Identify valence shell configuration
• Period number: Coefficient of s or p sub-shell
• Group number: Total electrons in valence shell (+ 10 for p-block)

Example

Nitrogen (N): 1s² 2s² 2p³
Period: 2, Group: 15 (5 + 10)

Oxygen (O): 1s² 2s² 2p⁴
Period: 2, Group: 16 (6 + 10)