Electronegativity, Bond Polarity, and Lewis Structures: Study Guide

Electronegativity and Bond Polarity

Definition and Trends

Electronegativity is a measure of how strongly an atom attracts electrons in a chemical bond. It is a dimensionless quantity, with higher values indicating a stronger attraction for electrons.

• Trend in the Periodic Table: Electronegativity increases across a period (left to right) and decreases down a group (top to bottom).

• Most Electronegative Element: Fluorine (F) is the most electronegative element.

Example: In a bond between hydrogen (H) and fluorine (F), F attracts electrons more strongly, resulting in a polar bond.

Bond Polarity and Dipole Moments

Bonds can be classified based on the difference in electronegativity (ΔEN) between the bonded atoms:

Electronegativity Difference (ΔEN)	Bond Type	Example
Small (0–0.4)	Pure (nonpolar) covalent	Cl2
Intermediate (0.4–2.0)	Polar covalent	HCl
Large (>2.0)	Ionic	NaCl

• Pure Covalent Bond: Electrons are shared equally (e.g., Cl2).

• Polar Covalent Bond: Electrons are shared unequally, creating partial charges (δ+ and δ−), as in H–F or H–Cl.

• Ionic Bond: Electrons are transferred, resulting in full charges (e.g., Na+ and Cl−).

Dipole Moment: A polar bond creates a dipole moment, represented by an arrow pointing toward the more electronegative atom.

Examples of Bond Types

• Br and Br: Pure covalent (ΔEN ≈ 0)

• C and Cl: Polar covalent (ΔEN ≈ 0.5)

• C and S: Pure covalent (ΔEN ≈ 0.03)

• Sr and O: Ionic (ΔEN ≈ 2.4)

Lewis Structures

Definition and Purpose

A Lewis structure represents the arrangement of atoms, bonds, and valence electrons in a molecule or ion. It helps visualize bonding and predict molecular geometry.

Steps to Draw a Lewis Structure

1. Find the total number of valence electrons.

2. Draw the skeleton structure using single bonds.

3. Assign remaining valence electrons as lone pairs to complete octets.

4. Make multiple bonds if all octets are not filled.

Example: For SO2, two valid Lewis structures can be drawn with the double bond on alternate sides. The actual structure is a resonance hybrid.

Resonance Structures

When more than one valid Lewis structure can be drawn for a molecule, the actual structure is an average (hybrid) of these forms, called resonance structures. Resonance is indicated by a double-headed arrow between structures.

Formal Charge

Formal charge is a bookkeeping tool to determine the most stable Lewis structure. It is calculated as:

$\text{Formal charge} = V - N - \frac{B}{2}$

• V: Number of valence electrons for the atom

• N: Number of non-bonding (lone pair) electrons

• B: Number of bonding (shared) electrons

The best Lewis structure minimizes formal charges and places negative charges on the most electronegative atoms.

Examples of Lewis Structures

• PH3: Central P atom with three single bonds to H, one lone pair on P.

• SCl2: Central S atom with two single bonds to Cl, two lone pairs on S.

• NF3: Central N atom with three single bonds to F, one lone pair on N.

• SF2: Central S atom with two single bonds to F, two lone pairs on S.

• HCOOH (Formic Acid): Central C atom double bonded to O, single bonded to OH and H.

Odd-Electron Species, Incomplete Octets, and Expanded Octets

Exceptions to the Octet Rule

• Odd-Electron Species: Molecules with an odd number of electrons (e.g., NO, NO2).

• Incomplete Octets: Atoms with fewer than 8 electrons (e.g., B in BCl3).

• Expanded Octets: Atoms in period 3 or below can have more than 8 electrons (e.g., PF5, SF4, ClF5).

Examples

• BCl3: Boron has only 6 electrons around it.

• PF5: Phosphorus has 10 electrons around it.

• SF4: Sulfur has 10 electrons around it.

• ClF5: Chlorine has 12 electrons around it.

Practice Problems and Applications

Bond Type Classification

• Given pairs of atoms, determine if the bond is pure covalent, polar covalent, or ionic using electronegativity differences.

Lewis Structure Drawing

• Draw Lewis structures for a variety of molecules and ions, including resonance forms and formal charges.

• Examples: CO, BrF, PH3, SCl2, NF3, HBr, SF2, HCOOH, CH2O, C2Cl4, N2O, ClO−, SeO2, CO32−, NO2−, ClO3−, ClO4−, SO42−, PF5, I3−, SF4, ClF5, AsF6−, Cl3PO, IF5.

Resonance and Formal Charge Tables

Tables are used to organize formal charge calculations for resonance structures. For example, the formal charge of NO2− and ClO3− can be summarized as follows:

Atom	Valence Electrons	Lone Pair Electrons	Bonding Electrons	Formal Charge
N	5	2	6	0
O1	6	6	2	-1
O2	6	4	4	0

Additional info: The best Lewis structure minimizes formal charges and distributes negative charges on the most electronegative atoms.

Summary Table: Bond Types by Electronegativity Difference

ΔEN	Bond Type	Example
0–0.4	Pure Covalent	Cl2
0.4–2.0	Polar Covalent	HCl
>2.0	Ionic	NaCl