Bonding & Lewis Structures

Introduction

This chapter covers the fundamental concepts of chemical bonding, focusing on electronegativity, bond types, Lewis dot structures, covalent bonding, multiple bonds, and the construction of Lewis structures. These topics are essential for understanding molecular structure and reactivity in GOB Chemistry.

Electronegativity

Definition and Periodic Trends

• Electronegativity is the ability of an atom to attract electrons in a chemical bond.

• It is a periodic property that varies across the periodic table:

  • Electronegativity decreases down a group (from top to bottom).

  • Electronegativity increases across a period (from left to right).

• The most electronegative element is fluorine (F).

Electronegativity Values Table

The following table summarizes electronegativity values for selected elements (Pauling scale):

Additional info: Electronegativity values are used to predict bond polarity and type.

Bond Types

Classification by Electronegativity Difference

The type of bond formed between atoms depends on the difference in their electronegativities:

• Nonpolar Covalent Bond: Equal sharing of electrons. EN Difference: 0.0 – 0.4

• Polar Covalent Bond: Unequal sharing of electrons. EN Difference: 0.5 – 1.8

• Ionic Bond: Transfer of electrons. EN Difference: 1.9 or greater

• Metallic Bond: Metals share all valence electrons as a "sea of electrons". EN Difference: low (0.0 – 1.0)

Bond Type Examples Table

Lewis Dot Structure

Representing Valence Electrons

• Lewis dot structures represent valence electrons as dots around the chemical symbol of an element.

• The number of valence electrons for main group elements equals the group number.

Examples

Covalent Bond

Formation and Octet Rule

• When two nonmetal atoms share a pair of electrons, a covalent bond is formed.

• Atoms share electrons to complete an octet (8 electrons) in their valence shell.

• Each shared pair of electrons constitutes one covalent bond.

Example: Chlorine Molecule

Two chlorine atoms share a pair of electrons:

• Each Cl now has 8 valence electrons.

Covalent Bonds (cont.)

Valence Shell Completion

• Atoms share enough electrons to fill their valence shell (usually 8 electrons for main group elements).

• Hydrogen is an exception, requiring only 2 electrons.

Valence Electrons Table

Multiple Bonds

Single, Double, and Triple Bonds

• More than one pair of electrons shared between atoms produces multiple bonds:

  • Single bond: 1 pair of electrons (e.g., Cl–Cl)

  • Double bond: 2 pairs of electrons (e.g., O=O)

  • Triple bond: 3 pairs of electrons (e.g., N≡N)

Bond Representation

• Single bond:

• Double bond:

• Triple bond:

Lewis Structures

Steps for Assigning Lewis Structures

1. Determine the total number of valence electrons contributed from all atoms in the molecule.

2. For ions, add 1 electron for each negative charge and subtract 1 electron for each positive charge.

3. Calculate the total number of bonds formed in the molecule (usually ½ the total number of bonds expected from all atoms).

4. Arrange atoms symmetrically around the least electronegative atom and connect with bonds (single lines for single bonds).

5. Add electrons to outlying atoms to satisfy the octet rule.

6. Place remaining electrons on the central atom.

7. If the central atom has fewer than 8 electrons, form multiple bonds as needed.

8. Check that the total number of electrons represented matches the number of valence electrons (including ion charge).

Example: Lewis Structure for Water ()

• Oxygen: 6 valence electrons

• Hydrogen: 1 valence electron each

• Total: electrons

• Structure: with two lone pairs on O

Polyatomic Ions

Lewis Structures for Polyatomic Ions

• Polyatomic ions contain covalent bonds and have a net charge.

• The charge must be accounted for when counting total electrons.

• Example: (Ammonium ion)

  • Nitrogen: 5 valence electrons

  • Hydrogen: 1 valence electron each

  • Subtract 1 electron for the positive charge

  • Total: electrons

Molecular Structure and VSEPR Theory

Electron Pair Repulsion and Molecular Shapes

• Molecular structure is determined by the arrangement of valence electrons (bonding and non-bonding pairs).

• Electron pairs repel each other and arrange themselves as far apart as possible (Valence Shell Electron Pair Repulsion, VSEPR theory).

• Common shapes include:

  • Tetrahedral: 4 areas of electron density, bond angle ≈ 109.5°

  • Trigonal pyramidal: 3 bonds and 1 lone pair, bond angle ≈ 107°

  • Bent: 2 bonds and 2 lone pairs, bond angle ≈ 104.5°

Example: Water () Shape

• Bent shape due to two lone pairs on oxygen

• Bond angle ≈ 104.5°

Molecular Polarity

Polar and Nonpolar Molecules

• Molecules with polar covalent bonds can be either polar or nonpolar overall.

• If bond dipoles cancel each other (symmetrical arrangement), the molecule is nonpolar (e.g., ).

• If bond dipoles do not cancel (asymmetrical arrangement), the molecule is polar (e.g., ).

• Requirements for a polar molecule:

  • Presence of polar covalent bonds

  • Asymmetric electron distribution

Example Table: Molecular Polarity

Additional info: Molecular polarity affects physical properties such as boiling point, solubility, and intermolecular forces.