lewis structures

Lewis Structure Tutorial on Formal Charge

Introduction

  • Continuation of Lewis structure tutorial from the previous video.
  • Focus on the worksheet used in the first video, particularly the blank last column that discusses formal charge.

Formal Charge Definition

  • Formal charge: A hypothetical charge on an atom within a molecular ion.
    • Defined as the charge an atom would have if the bonding electrons were shared equally among atoms.

Importance of Formal Charge

  • Used to differentiate between multiple Lewis structures for a given molecular ion.
    • Important in the context of resonance where more than one structure can be drawn.
  • Helps in selecting the best Lewis structure by applying the formal charge concept.

Calculating Formal Charge

  • Calculation of formal charge: {
    • Formula:
      extFormalCharge=extValenceElectronsofFreeAtom−extUnsharedValenceElectrons+12imesextSharedValenceElectronsext{Formal Charge} = ext{Valence Electrons of Free Atom} - ext{Unshared Valence Electrons} + \frac{1}{2} imes ext{Shared Valence Electrons}
    • Components:
    • Valence Electrons in free atom: Determine the total available valence electrons for the atom from the periodic table.
    • Unshared valence electrons: Number of lone pairs on the atom.
    • Shared valence electrons: Number of electrons involved in bonds summed across the atom's bonded atoms.

Example 1: Sulfur Difluoride (SF2)

  • Calculation for Sulfur (S):
    • Valence electrons = 6 (from periodic table).
    • Unshared valence electrons (lone pairs) = 4.
    • Shared valence electrons (bonding) = 4 (2 for each of the 2 S-F bonds).
    • Formal charge calculation:
      extFormalChargeonS=6−4+12imes4=0ext{Formal Charge on S} = 6 - 4 + \frac{1}{2} imes 4 = 0
  • Calculation for Fluorine (F) (same for both F atoms):
    • Valence electrons = 7.
    • Unshared valence electrons (lone pairs) = 6.
    • Shared valence electrons (bonding) = 2 (from the S-F bond).
    • Formal charge calculation:
      extFormalChargeonF=7−6+12imes2=0ext{Formal Charge on F} = 7 - 6 + \frac{1}{2} imes 2 = 0
  • Summary of Results:
    • Formal charge on Sulfur = 0; Formal charge on Fluorine = 0.
    • Check: The sum of the formal charges equals 0, matching the charge on the neutral molecule.

Example 2: Formaldehyde (CH2O)

  • Calculation for Carbon (C):
    • Valence = 4;
    • Unshared valence electrons = 0.
    • Shared valence electrons = 8 (from the double bond with oxygen and two bonds with hydrogen).
    • Formal charge:
      extFormalChargeonC=4−0+12imes8=0ext{Formal Charge on C} = 4 - 0 + \frac{1}{2} imes 8 = 0
  • Calculation for Oxygen (O):
    • Valence = 6;
    • Unshared valence electrons = 4 (2 lone pairs).
    • Shared valence electrons = 4 (double bond with carbon).
    • Formal charge:
      extFormalChargeonO=6−4+12imes4=0ext{Formal Charge on O} = 6 - 4 + \frac{1}{2} imes 4 = 0
  • Calculation for Hydrogen (H):
    • Valence = 1;
    • Unshared valence electrons = 0;
    • Shared valence electrons = 2 (bonding with carbon).
    • Formal charge:
      extFormalChargeonH=1−0+12imes2=0ext{Formal Charge on H} = 1 - 0 + \frac{1}{2} imes 2 = 0
  • Summary of Results:
    • Formal charges: Carbon = 0; Oxygen = 0; Hydrogens = 0.
    • Check: Sum of formal charges adds up to 0 for the neutral molecule.

Example 3: Cyanide Ion (CN−)

  • Calculation for Carbon (C):
    • Valence = 4;
    • Unshared valence electrons = 2;
    • Shared valence electrons = 6 (triple bond with nitrogen).
    • Formal charge:
      extFormalChargeonC=4−2+12imes6=−1ext{Formal Charge on C} = 4 - 2 + \frac{1}{2} imes 6 = -1
  • Calculation for Nitrogen (N):
    • Valence = 5;
    • Unshared valence electrons = 2;
    • Shared valence electrons = 6 (triple bond with carbon).
    • Formal charge:
      extFormalChargeonN=5−2+12imes6=0ext{Formal Charge on N} = 5 - 2 + \frac{1}{2} imes 6 = 0
  • Summary of Results:
    • Formal charges: Carbon = -1; Nitrogen = 0.
    • Check: The sum equals -1, which corresponds to the charge on the cyanide ion.

Example 4: Comparing Two Lewis Structures

  1. First Molecule:
    • Nitrogen (N): Valence = 5; Unshared = 2; Shared = 6.
    • Formal charge:
      extFormalChargeonN=5−2+12imes6=0ext{Formal Charge on N} = 5 - 2 + \frac{1}{2} imes 6 = 0
    • Oxygen (O): Valence = 6; Unshared = 4; Shared = 2.
    • Formal charge:
      extFormalChargeonO=6−4+12imes2=0ext{Formal Charge on O} = 6 - 4 + \frac{1}{2} imes 2 = 0
    • Chlorine (Cl): Valence = 7; Unshared = 6; Shared = 2.
    • Formal charge:
      extFormalChargeonCl=7−6+12imes2=0ext{Formal Charge on Cl} = 7 - 6 + \frac{1}{2} imes 2 = 0
  2. Second Molecule:
    • Nitrogen: Same as first molecule (charge = 0).
    • Oxygen: Valence = 6; Unshared = 6; Shared = 2.
    • Formal charge:
      extFormalChargeonO=6−6+12imes2=−1ext{Formal Charge on O} = 6 - 6 + \frac{1}{2} imes 2 = -1
    • Chlorine: Valence = 7; Unshared = 4; Shared = 2.
    • Formal charge:
      extFormalChargeonCl=7−4+12imes2=+1ext{Formal Charge on Cl} = 7 - 4 + \frac{1}{2} imes 2 = +1
  • Summary of Results:
    • First molecule formal charges = 0, 0, 0.
    • Second molecule formal charges = 0, -1, +1.
    • Both structures yield zero net formal charge.
    • Best structure based on the smallest formal charge is the first structure.

Final Criteria for Choosing Lewis Structures

  • Ensure that:
    • The sum of formal charges equals the overall charge of the ion or molecule.
    • Choose structures with the smallest possible formal charges.
    • Ideally, negative formal charges should reside on the more electronegative atoms.

Conclusion

  • Formal charge calculations are critical for choosing the best Lewis structure representations.
  • Engage in assessments and practice calculations to solidify understanding and prepare for further discussions in class regarding resonance.