Chemical Bonding I: The Lewis Model – Study Notes and Practice Guide

Chemical Bonding I: The Lewis Model

Introduction to Lewis Structures

The Lewis model is a foundational concept in chemistry for representing the arrangement of valence electrons in molecules and ions. Lewis structures help predict molecular shape, bond order, and reactivity by illustrating how atoms share or transfer electrons to achieve stable electron configurations.

• Lewis Structure: A diagram showing the bonding between atoms and the lone pairs of electrons in a molecule.

• Valence Electrons: Electrons in the outermost shell of an atom, involved in chemical bonding.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (except H, which seeks two).

Example: The Lewis structure for Cl2 is shown as Cl–Cl, with each Cl atom surrounded by three lone pairs.

Drawing Lewis Structures

• Count total valence electrons for all atoms in the molecule or ion.

• Arrange atoms with the least electronegative atom (except H) in the center.

• Connect atoms with single bonds (each bond = 2 electrons).

• Distribute remaining electrons as lone pairs to complete octets.

• Form double or triple bonds if necessary to satisfy the octet rule.

Example: For CO2, the Lewis structure is O=C=O, with each O atom having two lone pairs.

Bond Types and Strengths

• Single Bond: Two electrons shared between two atoms (e.g., H–H).

• Double Bond: Four electrons shared (e.g., O=O).

• Triple Bond: Six electrons shared (e.g., N≡N).

• Bond Strength: Triple bonds > Double bonds > Single bonds.

• Bond Length: Single bonds are longest; triple bonds are shortest.

Example: In N2, the triple bond makes the molecule very stable and short.

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: Delocalization of electrons across multiple atoms, increasing stability.

• Actual structure is a hybrid of all resonance forms.

Example: The carbonate ion (CO32–) has three resonance structures, each with a different C=O double bond location.

Formal Charge

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

• Formula: $\text{Formal charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2}(\text{Bonding electrons})$

• Structures with formal charges closest to zero are preferred.

Example: In OCN–, the best Lewis structure minimizes formal charges on each atom.

Lattice Energy

Lattice energy is the energy required to separate one mole of an ionic solid into gaseous ions. It reflects the strength of ionic bonds in a compound.

• Trends: Lattice energy increases with higher ionic charge and smaller ionic radius.

• Order: MgO > NaF > KCl (due to charge and size differences).

Equation: $\text{Lattice Energy} \propto \frac{q_1 q_2}{r}$, where $q_1$ and $q_2$ are ion charges, $r$ is the distance between ions.

Bond Energy and Enthalpy Calculations

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

• Formula: $\Delta H = \sum \text{(Bond energies of bonds broken)} - \sum \text{(Bond energies of bonds formed)}$

• Breaking bonds requires energy (endothermic); forming bonds releases energy (exothermic).

Example: For the reaction $\text{PCl}_3 + \text{Cl}_2 \rightarrow \text{PCl}_5$, use bond energies to estimate $\Delta H$.

Electronegativity and Bond Polarity

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

• Trend: Increases across a period, decreases down a group.

• Bond Polarity: Nonpolar (equal sharing), polar (unequal sharing), ionic (electron transfer).

Example: In H2O, O is more electronegative than H, making the O–H bonds polar.

Dipole Moment

The dipole moment measures the separation of positive and negative charges in a molecule, indicating molecular polarity.

• Formula: $\mu = q \times r$ (charge × distance)

• Molecules with polar bonds and asymmetric shapes have nonzero dipole moments.

Example: H2O has a large dipole moment due to its bent shape and polar bonds.

Bond Order and Strength

• Bond Order: Number of chemical bonds between a pair of atoms (single = 1, double = 2, triple = 3).

• Higher bond order means stronger and shorter bonds.

Example: In O2, the bond order is 2 (double bond).

Summary Table: Bond Types and Properties

Bond Type	Number of Shared Electrons	Bond Order	Relative Strength	Relative Length
Single	2	1	Weakest	Longest
Double	4	2	Intermediate	Intermediate
Triple	6	3	Strongest	Shortest

Practice Applications

• Predict the Lewis structure for various molecules and ions (e.g., CH2Cl2, NO3–, SO42–).

• Determine the number of bonding and lone pairs in molecules (e.g., H2O, NH3).

• Rank compounds by lattice energy, bond strength, or dipole moment using periodic trends and molecular structure.

• Estimate reaction enthalpy using bond energies.

Key Definitions

• Bond Energy: Energy required to break a bond.

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

• Electronegativity: Atom's ability to attract electrons in a bond.

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

• Resonance: Multiple valid Lewis structures for a molecule.

Additional info: These notes synthesize the main concepts and skills assessed in the provided question set, expanding on Lewis structures, resonance, bond energies, lattice energy, and related periodic trends as covered in a General Chemistry course.