Lecture_8-Bonding__Lewis_Structures

Lecture Overview

• Lecture 8: Bonding and Lewis Structures

Upcoming Due Dates

• Last REACT orientation meeting: February 13, 6 PM in 319 Greg Hall

• Quiz 2 due: February 20, 8 AM (Covers lectures/assignments 6-9)

• OWL HW 4 due: February 18, 8 AM

• Exam 1: February 26, 7-8:30 PM

General Bonding Characteristics of Representative Elements

• Elements strive for the lowest energy electron arrangement (most stable).

• Noble gas electron configurations (ns²np⁶) represent stable arrangements.

• Elements from Groups 1A-8A bond to achieve noble gas configurations through:

  • Ionic Bonds: Transfer of valence electrons (metal + nonmetal)

  • Covalent Bonds: Sharing of valence electrons (nonmetal + nonmetal)

Bond Types for Representative Elements (Groups 1A-8A)

Ionic Bonds

• Involves transfer of valence electrons to achieve noble gas configuration.

  • Example: Na + Cl → NaCl

  • Metals lose valence electrons, nonmetals gain.

Stabilization of Ionic Bonds

• Stabilization occurs through the attraction of oppositely charged ions (Coulomb’s Law).

• Lewis Structures: Illustrate arrangement of valence electrons in ionic compounds, aiming for noble gas configurations per ion.

Energetics of Forming Ionic Bonds

Energy Costs

• Example:

  • Mg → Mg²⁺ + 2e⁻ (DH = 2200 kJ)

  • O + 2e⁻ → O²⁻ (DH = 900 kJ)

  • Total Energy Cost: DH = 3100 kJ

Energy Savings

• Example:

  • Mg²⁺ + O²⁻ → MgO (DH = -3900 kJ)

  • Total DH = 3100 kJ - 3900 kJ = -800 kJ

  • Ionic bonds form as this yields a lower energy state (negative DH).

Higher Charges in Ionic Bonds

• Although ionic bonds like [Mg³⁺][O³⁻] would have larger energy of attraction due to higher charges, they do not form because:

  • Energy costs outweigh energy savings (DHoverall > 0). An exothermic reaction (DHoverall << 0) is required for bond formation.

Covalent Bonds

Formation of Covalent Bonds

• Involves sharing valence electrons to achieve the octet rule (H follows duet rule).

  • Example: H₂ shares electrons to achieve He configuration.

  • Example: F₂ shares electrons to achieve Ne configuration (8 valence electrons).

Stabilization of Covalent Bonds

• Covalent bonds are stabilized by the attraction between bonding electrons and the nuclei of both atoms involved.

Types of Covalent Bonds

Nonpolar Covalent Bonds

• Occur with equal sharing of electrons (e.g., H₂, O₂, F₂).

Polar Covalent Bonds

• Occur when electrons are shared unequally (e.g., HF, CO).

Bond Dipoles

• Bond dipoles indicate polar covalent bonds, featuring a partial positive and partial negative end.

• Represented with an arrow pointing from the positive to the negative end.

Electronegativity (EN) Trends

• Electronegativity refers to an atom's tendency to attract bonding electrons.

  • F: Most electronegative element; EN values increase towards F in the periodic table.

  • Hydrogen: Has an EN value between B and C, identical to P.

Predicting Bond Dipoles

• The direction of bond dipoles always points towards the atom with higher electronegativity.

Bond Type Classifications

• Classification based on electronegativity differences:

  • ΔEN = 0: Covalent bonds

  • Intermediate ΔEN: Polar covalent

  • Large ΔEN: Ionic bonds

Bond Type Assumptions

• Metal + nonmetal = Ionic bond (large ΔEN)

• Nonmetal + nonmetal = Covalent bond (smaller ΔEN)

• Identical nonmetals = Pure covalent bond (nonpolar)

• Different nonmetals generally produce polar covalent bonds unless specified exceptions (e.g., B-H, C-H, P-H are nonpolar).

Lewis Structure Basics

Steps in Drawing Lewis Structures

1. Count valence electrons.

2. Draw skeletal structure (place central atom).

3. Arrange valence electrons adhering to the octet rule (duet for H).

4. Utilize trial and error if needed.

Common Molecules

• Examples of more challenging Lewis structures with calculations for valence electrons:

  • H₂O: 2(1) + 6 = 8 valence electrons

  • NCl₃: 5 + 3(7) = 26 valence electrons

  • SO₄²⁻: 6 + 4(6) + 2 = 32 valence electrons

Exceptions to the Octet Rule

• Examples include BCl₃, NO₂, PCl₅, XeF₄.

Conclusion

• Always follow the octet rule whenever possible. Exceptions occur only when absolutely necessary. Understand these concepts for forming Lewis structures effectively.