Chapter 10: Chemical Bonding – Lewis Model, Ionic and Covalent Bonds, and Molecular Structure

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Chemical Bonding: Introduction

Bonding Models and AIDS Drugs

Chemical bonding theories are essential for understanding how atoms combine to form molecules, which in turn determines the properties and functions of substances, including drugs. For example, the structure of HIV-protease was determined using bonding models, enabling the design of protease inhibitors such as Indinavir to treat HIV/AIDS. These models help predict how drug molecules interact with biological targets.

Bonding theories predict how atoms bond to form compounds and explain molecular shapes and properties.
Drug design uses these models to simulate interactions between molecules and biological targets.
Protease inhibitors have been developed to treat HIV and other viral diseases.

Bonding Theories: The Lewis Model

Overview of the Lewis Model

The Lewis model is a foundational theory in chemistry for representing how atoms bond. It uses dot structures to show valence electrons and helps predict molecular stability and shape.

Lewis structures (dot structures) represent valence electrons as dots around element symbols.
The model predicts which atom combinations form stable molecules.
It is widely used for quick predictions, despite the existence of more advanced models.

Representing Valence Electrons with Dots

Lewis Structures for Main-Group Elements

Valence electrons are shown as dots around the element symbol. The number of valence electrons equals the group number for main-group elements (except helium).

Each dot = one valence electron.
Maximum two dots per side; fill singly before pairing.
Atoms with eight valence electrons (octet) are especially stable (e.g., Neon).
Helium is an exception, stable with two electrons (duet).

Example: Phosphorus (Group 5A) has 5 valence electrons: P with five dots.

Example: Oxygen electron configuration: (6 valence electrons).

Lewis Model of Chemical Bonding

Ionic and Covalent Bonds

Chemical bonds form when atoms share or transfer electrons to achieve stable configurations.

Ionic bond: Electrons are transferred (metal to nonmetal).
Covalent bond: Electrons are shared (between nonmetals).
Octet rule: Atoms tend to have eight electrons in their valence shell (except H and He, which require two).

Lewis Structures for Ionic Compounds

Electrons Transferred

When metals bond with nonmetals, electrons are transferred, forming cations and anions that attract to form ionic compounds.

Electron dots move from metal to nonmetal in the Lewis model.
Anion Lewis structures are shown in brackets with the charge.

Example: Potassium and chlorine:

K: • Cl: •• •• ••
K transfers electron to Cl:
Compound formed: KCl

Example: Magnesium and oxygen:

Mg: •• O: •• •• •• ••
Mg loses 2 electrons (), O gains 2 ():
Compound formed: MgO

Example: Sodium and sulfur:

Na: • S: •• •• •• ••
Two Na atoms needed for one S atom:
Compound formed: Na2S

Lewis Structures for Covalent Compounds

Electrons Shared

Nonmetals bond by sharing electrons, forming molecular compounds. Shared electrons allow atoms to achieve octets (or duets for hydrogen).

Bonding pair electrons: Shared between atoms (represented by dashes).
Lone pair electrons: Belong to one atom only.

Example: Water (H2O):

H: • O: •• •• ••
Each H shares electrons with O, forming two single bonds; O achieves octet, H achieves duet.
Lewis structure: H—O—H

Diatomic Molecules

Halogens (e.g., Cl2) and hydrogen (H2) form diatomic molecules to achieve stable configurations.
Lewis structures show shared pairs for each atom.

Double and Triple Bonds

Some molecules require double or triple bonds to satisfy the octet rule.

Double bond: Two pairs of electrons shared (e.g., O2).
Triple bond: Three pairs of electrons shared (e.g., N2).
Double and triple bonds are shorter and stronger than single bonds.

Steps to Write Lewis Structures for Covalent Compounds

Write the correct skeletal structure (central atom, terminal atoms).
Calculate total valence electrons (sum group numbers; adjust for ion charge).
Distribute electrons to give octets/duets (start with bonding pairs, then lone pairs).
If any atom lacks an octet, form double or triple bonds by moving lone pairs into bonding regions.

Example: CCl4 and CO2 structures follow these steps.

Polyatomic Ions and Resonance

Lewis Structures for Polyatomic Ions

Adjust electron count for ion charge (add for negative, subtract for positive).
Show structure in brackets with charge.
Example: Cyanide ion (CN-): 10 valence electrons, triple bond between C and N.

Resonance Structures

Some molecules (e.g., SO2) have more than one valid Lewis structure. The true structure is an average (resonance hybrid).

Resonance structures are shown with double-headed arrows.
Bonds in resonance hybrids are intermediate in length and strength.

Exceptions to the Octet Rule

Molecules with odd numbers of electrons (e.g., NO) cannot satisfy the octet rule for all atoms.
Boron compounds (e.g., BF3) often have only six electrons around B.
Some molecules (e.g., SF6) have expanded octets (more than eight electrons).

Predicting Molecular Shapes: VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory predicts molecular geometry based on the repulsion between electron groups (lone pairs, single, double, or triple bonds) around the central atom.

Electron groups arrange themselves to minimize repulsion.
Common geometries: linear (180°), trigonal planar (120°), tetrahedral (109.5°).
Lone pairs affect molecular geometry (e.g., NH3 is trigonal pyramidal, H2O is bent).

Steps to Predict Geometry Using VSEPR

Draw the Lewis structure.
Count electron groups around the central atom.
Count bonding groups and lone pairs.
Refer to geometry tables to determine electron and molecular geometry.

Table: Electron and Molecular Geometries

Electron Groups	Bonding Groups	Lone Pairs	Electron Geometry	Molecular Geometry	Example
2	2	0	Linear	Linear	CO2
3	3	0	Trigonal planar	Trigonal planar	H2CO
4	4	0	Tetrahedral	Tetrahedral	CH4
4	3	1	Tetrahedral	Trigonal pyramidal	NH3
4	2	2	Tetrahedral	Bent	H2O

Electronegativity and Polarity

Electronegativity

Electronegativity is the ability of an atom to attract electrons in a covalent bond. It is measured on a scale from 0.7 to 4.0 (Pauling scale), with fluorine being the most electronegative (4.0).

Electronegativity increases across a period and decreases down a group.
Unequal sharing of electrons leads to partial charges (dipole moment).

Bond Polarity

Nonpolar covalent bond: Electrons shared equally (e.g., Cl2).
Polar covalent bond: Electrons shared unequally (e.g., HCl).
Ionic bond: Electrons transferred (e.g., NaCl).

Electronegativity Difference	Bond Type	Example
0–0.4	Nonpolar covalent	Cl2
0.4–2.0	Polar covalent	HCl
>2.0	Ionic	NaCl

Polar Molecules

A molecule is polar if its polar bonds do not cancel out, resulting in a net dipole moment (e.g., H2O).
Linear molecules with identical polar bonds (e.g., CO2) are nonpolar because dipoles cancel.
Geometry affects overall polarity.

Applications: Polarity in Everyday Life

Why Oil and Water Don’t Mix

Water is polar; oil is nonpolar.
Polar molecules attract each other, excluding nonpolar molecules.

How Soap Works

Soap molecules have a polar head (attracts water) and a nonpolar tail (attracts grease/oil).
Soap allows water and grease to mix, enabling cleaning.

Review and Learning Objectives

Write Lewis structures for elements, ionic, and covalent compounds.
Draw resonance structures.
Predict molecular shapes using VSEPR theory.
Determine molecular polarity.