Lecture on Chemical Bonding

Topics Covered

• Ionization Energy
• Electron Affinity
• Electronegativity

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Ionization Energy (IE)

• Definition: The energy required to remove an electron from an isolated gaseous atom.

  • Also known as Ionization Enthalpy or Ionization Potential.
  • IE is always positive as energy is absorbed.

• Factors Affecting IE:

  • Atomic Radius: Smaller radius increases IE due to stronger attraction between nucleus and electron.
  • Nuclear Charge: Higher charge increases IE as it attracts the electrons more strongly.
  • Shielding Effect: More inner electrons reduce IE by shielding outer electrons from the nucleus.
  • Electron Configuration: Half-filled and fully-filled orbitals are more stable, thus increasing IE.

• Periodic Trends:

  • Across a period, IE generally increases due to decreasing atomic radius and increasing nuclear charge.
  • Down a group, IE generally decreases due to increasing atomic radius and shielding effect.

Electron Affinity (EA)

• Definition: The energy change when an electron is added to a neutral atom to form a negative ion.

  • First EA is usually exothermic (energy is released), while second EA can be endothermic (energy is absorbed).

• Factors Affecting EA:

  • Atomic Radius: Smaller radius increases EA because the added electron is closer to the nucleus.
  • Electron Configuration: Atoms with stable configurations (e.g., noble gases) have low EA values.

• Periodic Trends:

  • Generally increases across a period due to decreasing atomic radius.
  • Decreases down a group due to increasing atomic size and shielding effect.

Electronegativity (EN)

• Definition: The tendency of an atom to attract shared electrons in a covalent bond.

  • Not a thermodynamic quantity; no units.
  • Pauling Scale is commonly used to measure EN.

• Factors Affecting EN:

  • Atomic Radius: Smaller size usually means higher EN.
  • Nuclear Charge: Greater charge tends to increase EN.
  • Electron Configuration: Atoms with nearly full valence shells have higher EN.

• Periodic Trends:

  • EN generally increases across a period as atomic size decreases.
  • EN generally decreases down a group as atomic size increases.

Applications & Implications

• Ionization Energy:

  • Determines metallic character: Metals have low IE.
  • Influences reactivity: Lower IE often means higher reactivity for metals.
  • Stability: High IE indicates difficulty in removing an electron, thus more stability.

• Electron Affinity:

  • Nonmetals typically have high EA, leading to more negative ions.
  • Periodic Table position can predict EA behavior.

• Electronegativity:

  • Affects bond polarity; large EN difference leads to ionic bonds.
  • Influences molecular shape and properties.

Key Observations

• Metals are usually found on the left side of the Periodic Table, characterized by low IE, EA, and EN.
• Nonmetals are on the right side, with high IE, EA, and EN.
• Noble gases have high IE and low EA as they are stable and unreactive.

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Miscellaneous Points

• Understanding these concepts helps explain chemical reactivity and bonding.
• Ionization energy and electron affinity are crucial for predicting reactions and compounds' stability.
• Electronegativity is essential in determining the type and strength of chemical bonds.

Conclusion

• The lecture emphasized the importance of these properties in understanding and predicting chemical behavior.
• Encouragement to study chemistry with an open mind, recognizing it as a solution-oriented science.

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"Chemistry is not difficult; it provides solutions to every science problem."