Electronegativity, Bond Polarity, and Lewis Structures: Study Guide

Study Guide - Smart Notes

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Electronegativity and Bond Polarity

Definition and Trends

Electronegativity is a measure of how strongly an atom attracts electrons in a chemical bond. The most electronegative element is fluorine (F). Electronegativity increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

Trend: Increases from left to right across a period and decreases from top to bottom in a group.
Application: Used to predict bond polarity and the type of bond formed between atoms.

Example: Fluorine is the most electronegative element, so it strongly attracts electrons in bonds.

Bond Polarity and Types

The difference in electronegativity between two atoms determines the type of bond:

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

Pure Covalent Bond: Electrons are shared equally (ΔEN ≈ 0).
Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges (δ+ and δ−).
Ionic Bond: Electrons are transferred from one atom to another, forming ions.

Example: In H–F, the bond is polar covalent with F being δ− and H being δ+.

Bond Dipole and Molecular Polarity

A bond dipole is represented by an arrow pointing toward the more electronegative atom. In an electric field, polar molecules align with the field due to their dipole moments.

Dipole Moment: A measure of bond polarity, shown as an arrow (→) or δ+ and δ− notation.
Example: H–F aligns in an electric field with F toward the negative plate.

Lewis Structures

Definition and Purpose

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

Key Features: Shows all valence electrons as dots, bonds as lines, and lone pairs.
Application: Used to determine resonance, formal charge, and molecular shape.

Steps to Draw a Lewis Structure

Find the total number of valence electrons.
Draw the skeleton structure using single bonds.
Assign remaining valence electrons as lone pairs to complete octets.
Make multiple bonds if all octets are not filled.

Example: For H2O, O is the central atom with two lone pairs and two single bonds to H.

Resonance Structures

Some molecules can be represented by more than one valid Lewis structure, called resonance structures. The actual molecule is a hybrid of these forms.

Notation: Resonance structures are connected by a double-headed arrow (↔).
Example: SO2 has two resonance structures with the double bond on different oxygens.

Formal Charge

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

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.

Example: In H–F, both H and F have a formal charge of 0.

Practice: Lewis Structures for Selected Molecules and Ions

Examples

PH3: P is central, three single bonds to H, one lone pair on P.
SCl2: S is central, two single bonds to Cl, two lone pairs on S.
NF3: N is central, three single bonds to F, one lone pair on N.
SF2: S is central, two single bonds to F, two lone pairs on S.
HCOOH (Formic Acid): C is central, double bond to O, single bond to OH, single bond to H.
ClO− (Hypochlorite ion): Cl single bonded to O, three lone pairs on each, negative charge on O.

Resonance and Formal Charge Examples

CO32− (Carbonate ion): Three resonance structures, each with a double bond to a different O. Formal charge on each O is -1 or 0 depending on bonding.
NO2− (Nitrite ion): Two resonance structures, formal charge on N and O calculated for each.
SO42− (Sulfate ion): Four resonance structures, all S–O bonds equivalent.

Odd-Electron Species, Incomplete Octets, and Expanded Octets

Special Cases in Lewis Structures

Odd-Electron Species: Molecules with an odd number of electrons (e.g., NO) cannot satisfy the octet rule for all atoms.
Incomplete Octets: Some atoms (e.g., B in BCl3) are stable with fewer than eight electrons.
Expanded Octets: Atoms in period 3 or below (e.g., P, S, Cl) can have more than eight electrons (e.g., PF5, SF4, ClF5).

Summary Table: Bond Types and Electronegativity Difference

Bond Type	Electronegativity Difference (ΔEN)	Electron Distribution
Pure Covalent	0–0.4	Electrons shared equally
Polar Covalent	0.4–2.0	Electrons shared unequally
Ionic	>2.0	Electrons transferred

Key Equations

Formal Charge:

Additional info:

Lewis structures are foundational for predicting molecular geometry (VSEPR theory), bond polarity, and reactivity.
Resonance structures do not represent real, rapidly interconverting forms, but rather the delocalization of electrons in the molecule.