Chemical Bonding I: The Lewis Model – Electronegativity, Bond Polarity, Resonance, Formal Charge, and Bond Energies

Chemical Bonding I: The Lewis Model

Introduction

This study guide covers the foundational concepts of chemical bonding as described by the Lewis model, focusing on electronegativity, bond polarity, resonance, formal charge, exceptions to the octet rule, and trends in bond energies and bond lengths. These topics are essential for understanding molecular structure and reactivity in general chemistry.

Electronegativity

Definition and Trends

• Electronegativity is the ability of an atom to attract bonding electrons to itself.

• Electronegativity increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Fluorine is the most electronegative element; Francium is the least.

• Noble gases typically do not have assigned electronegativity values.

• Electronegativity is inversely related to atomic size: larger atoms have less ability to attract electrons.

Electronegativity Ordering Example

• Arrange the elements P, Na, N, Al in order of decreasing electronegativity: N > P > Al > Na

Electronegativity Difference and Bond Type

Bond Polarity and Classification

• The degree of polarity in a chemical bond depends on the electronegativity difference ($\Delta EN$) between bonded atoms.

• If $\Delta EN = 0$, the bond is pure covalent (nonpolar, equal sharing).

• If $\Delta EN$ is between 0.1 and 0.4, the bond is covalent.

• If $\Delta EN$ is between 0.4 and 1.9, the bond is polar covalent (unequal sharing).

• If $\Delta EN \geq 2.0$, the bond is ionic (electrons transferred).

The Continuum of Bond Types

• Pure (nonpolar) covalent bond: Electrons shared equally.

• Polar covalent bond: Electrons shared unequally, partial charges develop ($\delta^+$ and $\delta^-$).

• Ionic bond: Electrons transferred, resulting in full charges.

Bond Dipole Moments

Definition and Measurement

• Dipole moment ($\mu$) is a measure of bond polarity.

• A dipole is a bond in a molecule with a $\delta^+$ and $\delta^-$ end.

• Dipole moment is directly proportional to the size of the partial charges and the distance between them.

• Measured in Debyes (D).

• Generally, the more electrons two atoms share and the larger the atoms, the larger the dipole moment.

Dipole Moment Table

Molecule	$\Delta EN$	Dipole Moment (D)
Cl2	0	0
ClF	1.0	0.88
HF	1.9	1.82
LiF	3.0	6.33

Percent Ionic Character

Definition and Application

• Percent ionic character is the percentage of a bond's measured dipole moment compared to what it would be if the electrons were completely transferred.

• Indicates the degree to which the electron is transferred in a bond.

• No bond is 100% ionic.

• Bonds with >50% ionic character are classified as ionic bonds.

Percent Ionic Character Table (Inferred from Graph)

Compound	Electronegativity Difference	Percent Ionic Character
HF	1.9	~40%
LiF	3.0	~90%
KBr	~2.0	~70%
HCl	~0.9	~20%
NaCl	~2.1	~75%
Additional info:	Values are approximate, inferred from the graph.

Writing Lewis Structures of Molecules

Steps for Drawing Lewis Structures

1. Write the correct skeletal structure for the molecule.

   • Hydrogen atoms are always terminal.

   • More electronegative atoms are placed in terminal positions.

2. Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom.

   • For polyatomic ions, adjust for charge: add 1 electron for each negative charge, subtract 1 for each positive charge.

3. Distribute electrons among the atoms, giving octets (or duets for hydrogen) to as many atoms as possible.

   • Place 2 electrons between every 2 atoms.

   • Distribute remaining electrons as lone pairs starting with terminal atoms.

4. If any atoms lack an octet, form double or triple bonds as necessary.

5. Move lone pairs from terminal atoms into bonding regions with the central atom as needed.

Resonance

Concept and Importance

• The Lewis model localizes electrons as lone pairs or bonding pairs.

• Some molecules exhibit delocalization of electrons, known as resonance.

• Delocalization of charge helps stabilize the molecule.

Resonance Structures

• Resonance structures are two or more Lewis structures with the same skeletal formula but different electron arrangements.

• The actual molecule is a combination of the resonance forms, called a resonance hybrid.

• Look for multiple bonds or lone pairs that can be rearranged.

• Example: Ozone ($O_3$) can be represented by two resonance structures:

  • $\mathrm{O=O-O} \leftrightarrow \mathrm{O-O=O}$

Formal Charge

Definition and Calculation

• Formal charge is a fictitious charge assigned to each atom in a Lewis structure to help distinguish among competing structures.

• Calculated as:

  • $\text{FC} = \text{valence electrons} - [\text{nonbonding electrons} + \frac{1}{2} \text{bonding electrons}]$

• Sum of all formal charges in a molecule must be zero; in an ion, it equals the ion's charge.

Rules for Evaluating Resonance Structures

1. The sum of all formal charges in a neutral molecule must be zero.

2. The sum of all formal charges in an ion must equal the charge of the ion.

3. Small (or zero) formal charges on individual atoms are better than large ones.

4. When formal charge cannot be avoided, negative formal charge should reside on the most electronegative atom.

Exceptions to the Octet Rule

Types of Exceptions

• Odd number electron species (e.g., NO):

  • Will have one unpaired electron; called free radicals, which are very reactive.

• Incomplete octets:

  • Most important for boron, which forms compounds with only 6 electrons around it.

• Expanded octets:

  • Elements in the third row and beyond can have 12 or sometimes 14 electrons.

  • Occurs when empty d orbitals are available.

Example Questions

• Free Radical: Which molecule would you expect to be a free radical? Answer: NO (N2O)

• Expanded Octet: Which molecule could have an expanded octet? Answer: H3PO4

Bond Energy and Bond Length

Bond Energies

• Bond energy (or bond enthalpy) is the amount of energy required to break one mole of a bond in the gas phase.

• Chemical reactions involve breaking bonds in reactants and forming new bonds in products.

• The overall enthalpy change for a reaction can be calculated as the sum of the enthalpy changes associated with breaking and forming bonds.

Trends in Bond Energies

• The more electrons two atoms share, the stronger (higher bond energy) the covalent bond.

• Bond enthalpy increases as bond order increases:

  • Single bonds < double bonds < triple bonds

• The shorter the covalent bond, the stronger the bond (for similar types of bonds).

• Bonds get weaker down a column and stronger across a period.

Bond Lengths

• Bond length is the distance between the nuclei of bonded atoms.

• Actual bond length depends on surrounding atoms; average bond length is used for similar bonds across many compounds.

Trends in Bond Lengths

• The more electrons two atoms share, the shorter the covalent bond.

• Bond length decreases from left to right across a period (C–C > C–N > C–O).

• Bond length increases down a column (F–F < Cl–Cl < Br–Br).

• As bonds get longer, they also get weaker.

• Bond length decreases with increasing bond order:

  • Single bond > double bond > triple bond

Self-Assessment Table

Topic	Confident	Kind of Get It	Not a Clue
Electronegativity Difference/Bond Type
Resonance
Exceptions
Bond Length

Use this table to reflect on your understanding of the material and identify areas for further study.