Short Answer

The valence electron configuration of an element determines how easily it gains, loses, or shares electrons in chemical reactions.

• Elements with one or two valence electrons (like alkali and alkaline earth metals) tend to lose electrons easily, making them highly reactive.
• Elements with almost full valence shells (like halogens) readily gain electrons, also showing high reactivity.
• Elements with stable noble gas configurations are mostly inert.

Detailed Explanation :

Valence Electrons and Chemical Reactivity

• Valence electrons are the outermost electrons that participate in chemical bonding.
• The number of valence electrons and their configuration determine whether an element behaves as a metal or non-metal, and how it reacts with other elements.

1. Metals and Reactivity

• Metals generally have 1–3 valence electrons.
• Their reactivity increases when they can easily lose electrons to achieve a stable configuration.
• Example:
  • Sodium (Na): 1 valence electron → loses 1 e⁻ → forms Na⁺ → highly reactive.
  • Calcium (Ca): 2 valence electrons → loses 2 e⁻ → forms Ca²⁺ → reactive, but less than Na.
• Trend across the period:
  • Metals on the left side of the periodic table are more reactive because they have fewer valence electrons and lower ionization energy.
• Trend down the group:
  • Reactivity increases due to larger atomic size and weaker attraction of valence electrons to the nucleus.

2. Non-Metals and Reactivity

• Non-metals have 5–7 valence electrons.
• They are reactive because they gain electrons to complete their octet.
• Example:
  • Fluorine (F): 7 valence electrons → gains 1 e⁻ → forms F⁻ → extremely reactive.
  • Oxygen (O): 6 valence electrons → gains 2 e⁻ → forms O²⁻ → reactive.
• Trend across the period:
  • Reactivity increases from left to right for non-metals because elements get closer to achieving a full valence shell.
• Trend down the group:
  • Reactivity decreases because valence electrons are farther from the nucleus → weaker attraction → harder to gain electrons.

3. Noble Gases and Inertness

• Noble gases have full valence shells (ns² np⁶).
• Their stable electron configuration makes them chemically inert.
• They rarely react because no tendency exists to lose or gain electrons.

4. Role of Electron Configuration in Predicting Reactions

1. Octet Rule:
   • Atoms tend to gain, lose, or share electrons to achieve a full octet.
   • Determines whether elements act as oxidizing or reducing agents.
2. Periodic Trends:
   • Reactivity in metals: Decreases across a period, increases down a group.
   • Reactivity in non-metals: Increases across a period, decreases down a group.
3. Transition Elements:
   • Valence electrons in (n-1)d and ns orbitals allow multiple oxidation states.
   • Reactivity depends on which d- and s-electrons participate in bonding.
4. Predicting Chemical Behavior:
   • By examining valence configuration, we can determine:
     • Which ions an element forms
     • Type of bond likely to form (ionic or covalent)
     • Relative reactivity in chemical reactions

Conclusion

The valence electron configuration is the key factor determining an element’s reactivity.

• Elements with few valence electrons react by losing electrons (metals).
• Elements with almost full valence shells react by gaining electrons (non-metals).
• Noble gases, with complete valence shells, are mostly inert.

Understanding the valence configuration helps predict chemical reactivity, bond formation, and periodic trends across the periodic table.