Chapter 9: Chemical Bonding I – Lewis Structures and Bond Energies

Chapter 9: Chemical Bonding I – Lewis Structures

Learning Objectives

• Draw Lewis structures for simple compounds and polyatomic ions.

• Distinguish between ionic, polar covalent, and covalent bonding.

• Use the Born-Haber cycle to predict lattice energy for ionic compounds.

• Define electronegativity and use its value to assess bond polarity and metallic/nonmetallic character.

• Describe the relationship between electronegativity difference and percent ionic character of a bond.

• Apply the octet rule and formal charge concepts to determine the most plausible Lewis structure.

• Recognize and draw resonance structures.

• Draw Lewis structures for odd-electron and electron-deficient species.

• Identify expanded octets and hypervalent compounds.

• Use bond energies to compute enthalpy changes of reactions.

• Predict molecular polarity and dipole moments using electronegativity and molecular shape.

Types of Chemical Bonds

Ionic, Covalent, and Metallic Bonds

Chemical bonds are the forces that hold atoms together in compounds. The three main types are:

• Ionic Bond: Electrons are transferred from a metal to a non-metal, forming cations and anions. Example: $\text{Na} + \text{Cl} \rightarrow \text{Na}^+ + \text{Cl}^-$ (NaCl).

• Covalent Bond: Electrons are shared between non-metals. Example: $\text{H}_2\text{O}$, where O shares electrons with two H atoms.

• Metallic Bond: Valence electrons are delocalized over many metal nuclei, forming a 'sea of electrons' (e.g., in solid metals).

Representing Valence Electrons: Lewis Structures

Lewis Dot Symbols

Lewis structures use dots to represent valence electrons around an element's symbol. Only valence electrons are shown; inner shell electrons and electron spin are not depicted.

• Example: Si ([Ne]3s23p2) has 4 valence electrons: Si with 4 dots.

• Example: N ([He]2s22p3) has 5 valence electrons: N with 5 dots.

Octet Rule

Achieving Full Valence Shells

Atoms tend to form bonds to achieve a full octet (8 electrons) in their valence shell, similar to noble gases. Hydrogen is an exception, achieving a duet (2 electrons).

• Ionic Example: K transfers an electron to Cl, forming $\text{K}^+$ and $\text{Cl}^-$, both with full octets.

• Covalent Example: In $\text{H}_2\text{O}$, O shares electrons with two H atoms to complete its octet.

Drawing Lewis Structures for Molecular Compounds

Single, Double, and Triple Bonds

Covalent bonds form when atoms share electrons to achieve octets. A single bond is a shared pair (2 electrons), a double bond is two pairs (4 electrons), and a triple bond is three pairs (6 electrons).

• Single Bond: $\text{H}—\text{O}—\text{H}$ in water.

• Double Bond: $\text{O}=\text{O}$ in O2.

• Triple Bond: $\text{N}\equiv\text{N}$ in N2.

• Lone pairs: Electrons not involved in bonding; do not draw with a dash.

General Steps for Drawing Lewis Structures

1. Calculate total valence electrons in the molecule.

2. Draw single bonds between each pair of bonded atoms.

3. Calculate remaining valence electrons after accounting for bonds.

4. Distribute remaining electrons as lone pairs to terminal atoms first, then to the central atom.

5. If any atom lacks an octet, form double or triple bonds as needed by converting lone pairs to bonding pairs.

Valence electron counts for atoms are determined from the periodic table:

Group Number	Valence Electrons
1	1
2	2
13	3
14	4
15	5
16	6
17	7
18	8

Examples

• Monochloroamine (NH2Cl): N is central, 14 valence electrons, single bonds, lone pairs on Cl and N.

• Carbon Dioxide (CO2): C is central, 16 valence electrons, double bonds to each O, lone pairs on O.

Lewis Structures for Polyatomic Ions

Charged Species

Polyatomic ions are covalently bonded groups of atoms with an overall charge. The procedure is similar to neutral molecules, but add or subtract electrons to account for the charge. Enclose the structure in square brackets with the charge indicated.

• Example: ClO- has 14 valence electrons (7 from Cl, 6 from O, plus 1 for the negative charge).

• Example: NH4+ has 8 valence electrons (5 from N, 4 from H, minus 1 for the positive charge).

Lewis Structures for Ionic Compounds

Electron Transfer and Octet Achievement

Ionic compounds form by transferring electrons from metals to non-metals, resulting in ions with full octets. Example: Sodium sulfide (Na2S) – two Na atoms each lose one electron to S, forming Na+ and S2-.

The Ionic Bonding Model

Lattice Energy and Born-Haber Cycle

Formation of ionic solids is highly exothermic due to lattice energy, the energy released when gaseous ions form a solid lattice. The Born-Haber cycle uses Hess's law to calculate lattice energy from known enthalpy changes:

• Step 1: Sublimation of metal (e.g., Na(s) → Na(g))

• Step 2: Dissociation of non-metal molecule (e.g., ½ Cl2(g) → Cl(g))

• Step 3: Ionization of metal atom (e.g., Na(g) → Na+(g) + e-)

• Step 4: Electron affinity of non-metal (e.g., Cl(g) + e- → Cl-(g))

• Step 5: Formation of solid lattice (e.g., Na+(g) + Cl-(g) → NaCl(s))

The overall lattice energy is calculated by summing the enthalpy changes:

$\Delta H_{\text{lattice}} = \Delta H_1 + \Delta H_2 + \Delta H_3 + \Delta H_4 - \Delta H_5$

Trends in Lattice Energies

Factors Affecting Lattice Energy

• Ionic radius: As ions get larger, the distance between them increases, and lattice energy becomes less exothermic.

• Charge magnitude: Higher charges on ions result in much more exothermic lattice energies (Coulomb's law).

Compound	Lattice Energy (kJ/mol)
LiCl	-834
NaCl	-788
KCl	-701
CsCl	-657

For compounds with higher charges:

Compound	Lattice Energy (kJ/mol)
NaF	-910
CaO	-3414

Ionic Bonding: Models and Reality

Properties of Ionic Solids

• Exist as crystalline lattices of cations and anions.

• High melting points due to strong electrostatic attractions.

• Do not conduct electricity in solid state; conduct when dissolved in water (ions become mobile).

Covalent Bond Energies

Bond Energy and Enthalpy Calculations

Bond energy is the energy required to homolytically break 1 mol of bonds in the gas phase. Homolytic cleavage means electrons are divided equally between atoms.

• Bond energies are always positive (energy input required).

• Example: $\text{Cl}_2(g) \rightarrow 2\text{Cl}(g)$, $\Delta H = 243 \text{ kJ/mol}$

• H–Cl bond is stronger than Cl–Cl bond.

• Tabulated bond energies are averages from many compounds.

Using Bond Energies to Estimate Reaction Enthalpy

To estimate $\Delta H$ for a reaction:

• Sum the energies of bonds broken (positive values).

• Sum the energies of bonds formed (negative values).

• Overall: $\Delta H = \sum \Delta H_{\text{bonds broken}} + \sum \Delta H_{\text{bonds formed}}$

Example:

• CH4 + 2H2O → 4H2 + CO2

• Bonds broken: 4 C–H, 4 O–H

• Bonds formed: 4 H–H, 2 C=O

$\Delta H = \sum \Delta H_{\text{bonds broken}} + \sum \Delta H_{\text{bonds formed}}$

A reaction is exothermic ($-\Delta H$) when weaker bonds break and stronger bonds form; endothermic ($+\Delta H$) when strong bonds break and weak bonds form.

Additional info:

• Resonance structures, formal charge, incomplete octets, and expanded octets are covered in the full chapter but not shown in the provided images. These are essential for advanced Lewis structure analysis.