Chemical Bonding I: The Lewis Model – Ionic and Covalent Bonding, Lattice Energy, Electronegativity, and Lewis Structures

Chapter 10: Chemical Bonding I – The Lewis Model

Introduction to the Lewis Model

The Lewis Model is a foundational concept in chemical bonding, representing valence electrons as dots and using Lewis electron-dot structures to depict molecules. This model helps explain how atoms bond to achieve stable electron configurations.

• Valence electrons are the outermost electrons involved in bonding.

• Lewis structures use dots to represent valence electrons around atomic symbols.

• Atoms bond to achieve a stable electron configuration, often an octet (eight electrons).

Types of Chemical Bonds

Chemical bonds are classified based on how electrons are shared or transferred between atoms.

• Ionic Bond: Formed when a metal transfers one or more electrons to a nonmetal, resulting in cations and anions held together by electrostatic attraction.

• Covalent Bond: Formed when two nonmetals share electrons to achieve stable configurations.

Example: Sodium (Na) and Chlorine (Cl) form NaCl by electron transfer.

Ionic Bonding and Electron Transfer

Ionic bonding involves the transfer of electrons from a metal to a nonmetal, resulting in the formation of ions.

• Na (sodium) transfers its valence electron to Cl (chlorine):

$\text{Na} + \text{Cl} \rightarrow \text{Na}^+ + \text{Cl}^-$

• Both ions achieve an octet configuration.

• Formation of ionic compounds is usually exothermic.

Lattice Energy

Lattice energy is the energy associated with the formation of a crystalline ionic solid from gaseous ions. It is a measure of the strength of the ionic bonds in a solid.

• Lattice energy is usually negative (exothermic).

• Calculated using the Born-Haber Cycle, which involves several steps:

Example Calculation for NaCl: Step 1: Sublimation of Na(s) to Na(g) Step 2: Ionization of Na(g) to Na+(g) Step 3: Dissociation of Cl2(g) to Cl(g) Step 4: Electron affinity of Cl(g) to Cl-(g) Step 5: Formation of NaCl(s) from Na+(g) and Cl-(g)

$\Delta H_{\text{lattice}} = \Delta H_{\text{f}} - (\text{sum of other enthalpy changes})$

• Lattice energy depends on the magnitude of charges (directly) and the distance between ions (inversely).

Electrostatic Potential Formula: $E = \dfrac{1}{4\pi\varepsilon_0} \dfrac{q_1 q_2}{r}$ Where $q_1$ and $q_2$ are the charges, and $r$ is the distance between them.

Trends in Lattice Energies

Lattice energy varies with ion size and ion charge.

• Ion Size: Smaller ions result in higher (more negative) lattice energies.

• Ion Charge: Higher charges result in higher (more negative) lattice energies.

Example Table: Comparison of Lattice Energies

Compound	Ion Size	Ion Charge	Lattice Energy (kJ/mol)
LiF	Small	+1/-1	-1036
NaI	Large	+1/-1	-682
MgO	Small	+2/-2	-3795
KI	Large	+1/-1	-632

Key Points:

• Greater distance between ions = less negative lattice energy.

• Greater charge magnitude = more negative lattice energy.

Lewis Structures for Covalent Bonds

Lewis structures are used to represent covalent bonds, showing how atoms share electrons to achieve octet configurations.

• Single Covalent Bond: Atoms share one pair of electrons (e.g., H2, F2).

• Double Covalent Bond: Atoms share two pairs of electrons (e.g., O2, CO2).

• Triple Covalent Bond: Atoms share three pairs of electrons (e.g., N2).

• Lone pairs are nonbonding electrons shown as pairs of dots.

Example: Water (H2O) has two bonding pairs and two lone pairs on oxygen.

Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract electrons in a chemical bond. Bond polarity arises when electrons are not equally shared.

• Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges (e.g., H–F).

• Nonpolar Covalent Bond: Electrons are shared equally (e.g., H2).

• Electronegativity increases across a period and decreases down a group.

• Fluorine (F) is the most electronegative element.

Writing Lewis Structures for Molecular Compounds

Lewis structures for molecules are drawn using a systematic approach:

1. Write the correct skeletal structure for the molecule.

2. Calculate the total number of valence electrons.

3. Distribute electrons among atoms to satisfy the octet rule (or duet for hydrogen).

4. Assign lone pairs as needed.

5. Use double or triple bonds if necessary to satisfy octet rule.

Example: Drawing the Lewis structure for CO2:

• Total valence electrons: 16

• Central atom: C

• Each O forms a double bond with C

Resonance Structures

Some molecules have more than one valid Lewis structure. These are called resonance structures, and the actual molecule is a resonance hybrid.

• Resonance structures differ only in the placement of electrons, not atoms.

• The true structure is an average of all resonance forms.

Example: Ozone (O3) has two resonance structures.

Summary Table: Bond Types and Electron Sharing

Bond Type	Electron Sharing	Example
Ionic	Transfer	NaCl
Covalent (Single)	1 pair shared	H2
Covalent (Double)	2 pairs shared	O2
Covalent (Triple)	3 pairs shared	N2

Additional info:

• Lewis structures are essential for predicting molecular geometry and reactivity.

• Lattice energy calculations are important for understanding the stability of ionic solids.

• Electronegativity differences help classify bonds as ionic, polar covalent, or nonpolar covalent.