Chemical Bonding I: Lewis Theory – Study Notes and Practice

Chemical Bonding I: The Lewis Model

Introduction to Lewis Theory

The Lewis model is a foundational theory in chemistry for understanding how atoms bond to form molecules. It uses simple diagrams to represent valence electrons and predict molecular structure, bond order, and stability.

• Lewis Structures: Diagrams that show the arrangement of valence electrons around atoms in a molecule.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a noble gas configuration (8 valence electrons).

• Bonding Pairs: Shared pairs of electrons between atoms (represented as lines).

• Lone Pairs: Non-bonding pairs of electrons localized on a single atom (represented as dots).

• Formal Charge: A bookkeeping tool to determine the most stable Lewis structure.

Drawing Lewis Structures

Lewis structures are drawn by following a systematic process:

1. Count the total number of valence electrons for all atoms in the molecule or ion.

2. Arrange the atoms, usually with the least electronegative atom in the center (except hydrogen).

3. Connect atoms with single bonds (lines), then distribute remaining electrons as lone pairs to complete octets.

4. If necessary, form double or triple bonds to satisfy the octet rule.

5. Assign formal charges to check for the most stable structure.

Example: The Lewis structure for CO2 is drawn as O=C=O, with each oxygen atom having two lone pairs.

Resonance Structures

Some molecules cannot be represented by a single Lewis structure. Resonance structures are multiple valid Lewis structures for the same molecule, differing only in the placement of electrons.

• Resonance Hybrid: The actual structure is a hybrid of all resonance forms.

• Example: The carbonate ion, CO32−, has three resonance structures, each with one C=O double bond and two C–O single bonds.

Bond Order and Bond Strength

Bond order refers to the number of chemical bonds between a pair of atoms. Higher bond order generally means stronger and shorter bonds.

• Single Bond: Bond order = 1

• Double Bond: Bond order = 2

• Triple Bond: Bond order = 3

• Bond Strength: Increases with bond order; triple bonds are stronger than double, which are stronger than single bonds.

• Bond Length: Decreases with bond order; triple bonds are shortest.

Example: In N2, the bond order is 3 (triple bond), making it very strong and short.

Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract electrons in a chemical bond. Differences in electronegativity lead to bond polarity.

• Nonpolar Covalent Bond: Electrons are shared equally (e.g., Cl2).

• Polar Covalent Bond: Electrons are shared unequally (e.g., HCl).

• Ionic Bond: Electrons are transferred from one atom to another (e.g., NaCl).

Electronegativity Trend: Increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

Lattice Energy

Lattice energy is the energy required to separate one mole of an ionic solid into its gaseous ions. It is a measure of the strength of the ionic bonds in a crystal lattice.

• Factors Affecting Lattice Energy:

  • Charge of ions: Higher charges lead to higher lattice energy.

  • Size of ions: Smaller ions lead to higher lattice energy.

• Example: MgO has a much higher lattice energy than NaCl due to higher charges and smaller ionic radii.

Equation:

$\text{Lattice Energy} \propto \dfrac{Q_1 Q_2}{r}$

where $Q_1$ and $Q_2$ are the charges on the ions, and $r$ is the distance between ion centers.

Bond Energy and Enthalpy Calculations

Bond energy is the energy required to break one mole of a bond in the gas phase. It is used to estimate the enthalpy change ($\Delta H$) of reactions.

• Bond Dissociation Energy: The energy needed to break a specific bond in a molecule.

• Estimating Reaction Enthalpy: Use bond energies to estimate $\Delta H$ for a reaction:

$\Delta H_{\text{rxn}} = \sum \text{(Bond energies of bonds broken)} - \sum \text{(Bond energies of bonds formed)}$

Example: Calculating $\Delta H$ for the reaction $\mathrm{CH_4 + 2Cl_2 \rightarrow CH_2Cl_2 + 2HCl}$ using bond energies.

Formal Charge

Formal charge helps determine the most stable Lewis structure by assigning charges to atoms based on electron ownership.

• Formula:

$\text{Formal Charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \dfrac{1}{2} \times \text{Bonding electrons}$

Example: In OCN−, formal charges help identify the most stable resonance structure.

Practice with Lewis Structures

• Draw Lewis structures for molecules and ions such as Cl2, Br2, NO3−, SO2, NH4+, and CO32−.

• Identify the best Lewis structure based on octet completion and formal charges.

• Recognize resonance structures and delocalized electrons.

Summary Table: Bond Order, Bond Length, and Bond Strength

Bond Type	Bond Order	Bond Length	Bond Strength
Single	1	Longest	Weakest
Double	2	Intermediate	Intermediate
Triple	3	Shortest	Strongest

Key Terms and Definitions

• Electronegativity: The ability of an atom to attract electrons in a chemical bond.

• Lattice Energy: The energy required to separate an ionic solid into gaseous ions.

• Bond Energy: The energy required to break a bond in a molecule.

• Formal Charge: A calculated charge on an atom in a Lewis structure.

• Resonance: The concept that some molecules are best described by two or more valid Lewis structures.

Applications and Examples

• Predicting molecular stability using formal charges and resonance.

• Ranking compounds by lattice energy, bond strength, or bond length.

• Estimating reaction enthalpy using bond energies.

• Determining the polarity of bonds and molecules based on electronegativity differences.

Additional info: These notes expand on the practice questions by providing academic context, definitions, and examples relevant to Lewis theory, bond energies, lattice energy, and electronegativity, as covered in a typical General Chemistry curriculum.