Recording-2025-01-21T16:12:38.301Z

Drawing Electron Dot Structures

• Valence Electrons: Determine the number of valence electrons by looking at the family number of each atom.

  • Example: Sulfur has 6 valence electrons.

  • Oxygen also has 6 valence electrons.

  • For two Chlorine atoms: Total is 26 electrons (6 from sulfur, 6 from oxygen, and 6 from each chlorine, 2x6=12).

Constructing the Bond Skeleton

• Begin by drawing a bond skeleton.

  • Arrange atoms with the first listed atom in the center (e.g., O, Cl, Cl).

  • Outline may vary, but symmetry is preferred.

  • Starting with 26 total electrons, using 6 for the skeleton leaves 20 remaining electrons.

Adding Electrons and Checking Octets

• Placing Electrons: Use remaining electrons as dots to complete the octets.

  • Count out 20 electrons to ensure all atoms (except hydrogen) have an octet.

  • Extra electrons go on the central atom if there are any after fulfilling octet.

• Octet Rule Confirmation:

  • Confirm that all relevant atoms satisfy octet rule, remembering hydrogen can only have 2 electrons.

  • Verify the total hasn't exceeded the available valence electrons, i.e., 26.

Resonance Structures

• Lewis structures can sometimes be represented in different ways (resonance structures).

  • It doesn’t significantly affect Lewis structure whether double, single, or triple bonds are used, as long as total valence and octet rule are satisfied.

Molecular Geometry and Bond Angles

• Identify center atom (A) and exterior atoms (X).

  • For example, a structure with sulfur as A and three exterior atoms (X) such as two chlorines and one oxygen can be noted as AX3E (one lone pair on the central atom).

• VSEPR Theory: Identify the geometry based on lone pairs around the central atom, yielding a pyramidal shape for AX3E.

  • Bond angle = 109.5°.

Dipole Moment and Polarity

• Determine if the molecule is polar or nonpolar based on:

  1. Electronegativity Difference: A significant difference (0.3 or greater) in electronegativity between bonded atoms indicates polarity.

     • Example calculations:

       • Oxygen: 3.5, Sulfur: 2.5, Chlorine: 3.0.

  2. Molecular Geometry: Asymmetric charge distribution leads to polarity. Even distributions are nonpolar.

Example Analysis

• Considering sulfur-chlorine and sulfur-oxygen bonds:

  • Draw arrows to represent electron pulls from chlorine to sulfur due to differences in electronegativity, noting directionality (positive/negative ends).

  • Check electrons drawn for each bond.

• Since there’s unequal electron distribution due to lack of symmetry, conclude that the molecule has a dipole moment, indicating it is polar.

Practice and Application

• Assign each molecule for practice:

  • Draw Lewis structure and VSEPR structure, noting geometry and bond angles.

  • Use electronegativity chart to evaluate bonds and draw dipole arrows as necessary.

  • Determine net polarity based on symmetry and electronegativity differences.

Conclusion and Next Steps

• Complete the exercises with peers in class to solidify understanding before the upcoming lesson.

• Resources available for those needing assistance, including a recap on Lewis structure rules.