Bonding and Molecular Structure: Study Notes for General Chemistry

Chapter 4: Bonding

Ionic and Covalent Bonds

Chemical bonding is a fundamental concept in chemistry, describing how atoms combine to form compounds. The two main types of chemical bonds are ionic bonds and covalent bonds.

• Ionic Bond: Formed when electrons are transferred from one atom to another, typically between metals and nonmetals. Ionic compounds consist of positive and negative ions held together by electrostatic forces.

• Covalent Bond: Formed when two atoms share one or more pairs of electrons, usually between nonmetals.

• Examples: NaCl (ionic), H2O (covalent)

Bond Strength and Lattice Energy

The strength of a bond and the stability of ionic compounds are influenced by lattice energy and bond order.

• Lattice Energy: The energy required to separate one mole of an ionic solid into its gaseous ions. Higher lattice energy indicates stronger ionic bonds.

• Bond Order: The number of shared electron pairs between two atoms (single, double, triple bonds). Higher bond order means stronger and shorter bonds.

• Formula (Coulomb's Law):

$E = \dfrac{kQ_1Q_2}{r}$

• Where E is the electrostatic energy, Q1 and Q2 are the charges, r is the distance between ions, and k is a constant.

Periodic Trends: Ionic Radii

Ionic radii refer to the size of ions. Cations are generally smaller than their parent atoms, while anions are larger.

• Cations: Lose electrons, resulting in a smaller radius.

• Anions: Gain electrons, resulting in a larger radius.

• Trend: Ionic radius increases down a group and decreases across a period for cations and anions.

Electronegativity and Polarity

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

• Polar Bonds: Electrons are shared unequally due to differences in electronegativity.

• Nonpolar Bonds: Electrons are shared equally.

• Example: HCl is polar, Cl2 is nonpolar.

Dipole Moments and Molecular Polarity

The dipole moment measures the separation of positive and negative charges in a molecule. It is influenced by bond polarity and molecular geometry.

• Formula:

$\mu = q \times r$

• Where q is the charge and r is the distance between charges.

• Example: Comparing HF, HCl, HBr, HI: Dipole moment decreases as electronegativity difference decreases.

Lewis Structures and Resonance

Lewis structures represent the arrangement of electrons in a molecule. Resonance occurs when more than one valid Lewis structure can be drawn for a molecule.

• Steps to Draw Lewis Structures:

  1. Count total valence electrons.

  2. Arrange atoms and connect with single bonds.

  3. Distribute remaining electrons to complete octets.

  4. Check for resonance structures.

• Exceptions to the Octet Rule: Some atoms (e.g., H, B, P, S) may have fewer or more than eight electrons.

• Resonance: Delocalization of electrons across multiple structures (e.g., O3, NO3-).

Formal Charge and Structure Stability

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

• Formula:

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

• Structures with formal charges closest to zero are generally more stable.

Naming Compounds

Correctly naming compounds is essential in chemistry. Different rules apply to ionic, molecular, hydrates, and oxyanions.

• Ionic Compounds: Name cation first, then anion. Use Roman numerals for transition metals.

• Molecular Compounds: Use prefixes (mono-, di-, tri-) to indicate the number of atoms.

• Hydrates: Add "hydrate" with appropriate prefix (e.g., CuSO4·5H2O is copper(II) sulfate pentahydrate).

• Oxyanions: Use suffixes (-ate, -ite) and prefixes (hypo-, per-) to indicate oxygen content.

Chapter 5: Advanced Theory - Hybrid Orbitals, MO Theory

Valence Bond Theory and Hybridization

Valence Bond Theory explains how atomic orbitals combine to form chemical bonds. Hybridization is the mixing of atomic orbitals to form new, equivalent hybrid orbitals.

• Why Hybrid Orbitals? They allow for the observed molecular geometries and bond angles.

• Types of Hybridization:

  • sp (linear geometry)

  • sp2 (trigonal planar geometry)

  • sp3 (tetrahedral geometry)

• Example: CH4 uses sp3 hybridization.

VSEPR Theory and Molecular Geometry

VSEPR (Valence Shell Electron Pair Repulsion) Theory predicts the shapes of molecules based on electron pair repulsion.

• Five Basic Geometries:

  • Linear (180°)

  • Trigonal planar (120°)

  • Tetrahedral (109.5°)

  • Trigonal bipyramidal (90°, 120°)

  • Octahedral (90°)

• Molecular Geometry vs. Shape: Geometry considers both bonds and lone pairs; shape refers to the arrangement of atoms only.

• Bond Angles: Lone pairs can reduce bond angles below ideal values (e.g., H2O bond angle is 104.5°, not 109.5°).

• Polarity: Shape and bond polarity determine if a molecule is polar or nonpolar.

Hybrid Orbitals and Bond Types

Hybrid orbitals explain the formation of sigma and pi bonds in molecules.

• Sigma (σ) Bond: Formed by direct overlap of orbitals along the bond axis.

• Pi (π) Bond: Formed by sideways overlap of p orbitals.

• Double and Triple Bonds: Double bonds have one sigma and one pi bond; triple bonds have one sigma and two pi bonds.

Molecular Orbital (MO) Theory

Molecular Orbital Theory describes electrons in molecules as occupying molecular orbitals that are spread over the entire molecule.

• Molecular Orbitals: Formed by the combination of atomic orbitals.

• Bonding and Antibonding Orbitals: Bonding orbitals stabilize the molecule; antibonding orbitals destabilize it.

• Example: O2 has both bonding and antibonding molecular orbitals, explaining its paramagnetism.

Summary Table: Types of Compounds and Naming Rules

The following table summarizes the main types of compounds and their naming conventions:

Type of Compound	Naming Rule	Example
Ionic (with transition metals)	Cation (Roman numeral) + Anion	FeCl3: Iron(III) chloride
Molecular	Prefix + Name of first element + Prefix + Name of second element	CO2: Carbon dioxide
Hydrates	Compound name + Prefix + hydrate	CuSO4·5H2O: Copper(II) sulfate pentahydrate
Oxyanions	Prefix/Suffix based on oxygen content	NO3-: Nitrate; NO2-: Nitrite

Additional info: Some context and examples have been expanded for clarity and completeness.