4/24/25 CCP Chemistry Lecture (Unit 4.2, Chap 11)

Lewis Symbols and Valence Electrons

• Lewis Symbols: Represent the valence electrons of an element using dots around its chemical symbol.
• Valence Electrons: The outermost electrons involved in bonding.
  • For elements in the A groups, the number of valence electrons equals the group number.
  • Example: Aluminum (Al) has 3 valence electrons as it is in group 3.
• Drawing Lewis Symbols:
  • Place dots on four sides of the element's symbol.
  • Start with single unpaired electrons until all four sides have been filled.
  • Do not pair electrons until necessary.

Examples of Lewis Structures

• Aluminum Example:

  • Aluminum's Lewis symbol will show three unpaired dots on three of its four sides.

• Sulfur Example:

  • Sulfur has 6 valence electrons.
  • First four electrons are placed unpaired on different sides; the remaining two are paired.
  • Hence, it shows two lone pairs and two unpaired electrons.

Bond Types and Structure Construction

• Types of Bonds:

  • Single Bond: 1 line = 2 electrons
  • Double Bond: 2 lines = 4 electrons
  • Triple Bond: 3 lines = 6 electrons

• Lewis Structure Construction (AXN):

  • Count total valence electrons.
  • Example:
  • For a molecule with formula AXN, determine the valence count.
  • Arsenic in group 5 and Fluorine in group 7, leads to a total of 26 valence electrons.
  • Hook outer atoms (X) to the central atom (A) with single bonds.
  • Start placing remaining electrons as lone pairs on outer atoms.
  • Ensure outer atoms achieve octets.

Lone Pairs and Expanded Valences

• Lone Pairs: Unpaired electrons that are not involved in bonding.

• Exceptions in Octet Rule:

  • Atoms like Beryllium can be stable with fewer than 8 electrons.
  • Boron with 6 valence electrons is another exception.
  • Odd-electron molecules can also feature less than 8 at the central atom.

• Expanded Valence Shells:

  • Certain elements can hold more than 8 electrons due to their position on the periodic table (magic number is 13).
  • Example: Chlorine and Selenium can exhibit expanded valences.

Formal Charge Calculation

• Formal Charge Formula:
  • $Formal \, Charge = (Valence \, Electrons) - (Lone \, Pair \, Electrons) - (0.5 imes Shared \, Electrons)$
  • Example Calculation for Nitrogen:
  • Starts with 5 valence electrons, determine if it has lone pairs or shared in bonds.

Major Structure Determination

• Identify formal charges, preferring structures with more zero charges.
• If tie occurs, check for electronegativity: Negative charge should be on the more electronegative atom.
• Determining major structures helps in understanding true bond character.

Condensed Structures and Bonding

• Condensed Structure:
  • Represents connections without complex details about bonds.
  • Organic chemists favor structural formulas over molecular ones because they provide more insight about connectivity.

Bond Enthalpy and Reaction Energy

• Bond Enthalpy: Energy required to break a bond.
• Understanding reactions requires identifying what bonds are being broken/formed.
• Net change in energy can be calculated using bond enthalpies to reveal if the reaction is exothermic or endothermic.
• Example: For a reaction involving Carbon, Hydrogen and Oxygen, consider bond changes to find enthalpy.
  • Metric calculations can simplify analyzing full structures.

Final Notes

• Review of key concepts and reactions in chemistry should flow into the next chapter.
• Reinforcement of Lewis structures, bonding types, and how to apply these concepts in real calculations.
• Continuation of study to ensure comfort with the upcoming material and test preparation.