Chemical Bonding, Molecular Structure, and Nomenclature: Study Guide

Chemical Bonding and Molecular Structure

Bond Polarity and Dipole Moments

Chemical bonds can be classified by the distribution of electron density between the bonded atoms. Understanding bond polarity is essential for predicting molecular properties and reactivity.

• Bond Polarity: A covalent bond is polar if the two atoms have different electronegativities, causing unequal sharing of electrons.

• Bond Dipole: The shift in electron density creates a dipole moment, represented by an arrow pointing toward the more electronegative atom.

• Molecular Polarity: The overall polarity of a molecule depends on both the polarity of individual bonds and the molecular geometry.

• Dipole Moment: A vector quantity representing the separation of charge in a molecule.

Example: In H2O, the O-H bonds are polar, and the bent geometry results in a net molecular dipole moment.

Valence Bond Theory and Hybridization

Valence bond theory explains how atomic orbitals combine to form chemical bonds. Hybridization describes the mixing of atomic orbitals to create new, equivalent hybrid orbitals for bonding.

• Valence Bond Theory: Bonds form when atomic orbitals overlap.

• Hybridization: The process of mixing atomic orbitals (s, p, d) to form hybrid orbitals (sp, sp2, sp3, etc.).

• Geometry Connection: The type of hybridization determines the geometry around the central atom.

Example: Methane (CH4) has sp3 hybridization, resulting in a tetrahedral geometry.

Sigma (σ) and Pi (π) Bonds

Covalent bonds are classified as sigma or pi bonds based on the type of orbital overlap.

• Sigma (σ) Bond: Formed by head-on overlap of orbitals; all single bonds are sigma bonds.

• Pi (π) Bond: Formed by side-on overlap of p orbitals; present in double and triple bonds.

Example: In ethene (C2H4), the C=C bond consists of one sigma and one pi bond.

Molecular Orbital (MO) Theory

MO theory describes the combination of atomic orbitals to form molecular orbitals, which can be bonding or antibonding.

• Bonding Orbital: Lower in energy, increases electron density between nuclei.

• Antibonding Orbital: Higher in energy, decreases electron density between nuclei.

• Bond Order: Indicates the strength and stability of a bond.

Bond Order Formula:

$\text{Bond Order} = \frac{(\text{Number of electrons in bonding orbitals}) - (\text{Number of electrons in antibonding orbitals})}{2}$

Example: O2 has a bond order of 2 and is paramagnetic due to unpaired electrons in its MO diagram.

Steric Number, Electron-Pair, and Molecular Geometries

The steric number helps determine the geometry around a central atom using VSEPR theory.

• Steric Number: The sum of bonded atoms and lone pairs on the central atom.

• Electron-Pair Geometry: Arrangement of regions of electron density (bonds and lone pairs).

• Molecular Geometry: Arrangement of atoms (ignoring lone pairs).

Common Steric Numbers and Geometries:

Steric Number	Electron-Pair Geometry	Ideal Bond Angles
2	Linear	180°
3	Trigonal Planar	120°
4	Tetrahedral	109.5°
5	Trigonal Bipyramidal	90°, 120°
6	Octahedral	90°

Effect of Lone Pairs: Lone pairs decrease bond angles relative to the ideal geometry.

Lewis Structures and Resonance

Counting Valence Electrons

Valence electrons are the outermost electrons involved in bonding. Accurate counting is essential for drawing Lewis structures.

• Atoms: Use group number for main-group elements.

• Ions: Add electrons for negative charge, subtract for positive charge.

Example: NO3- has 5 (N) + 3×6 (O) + 1 (charge) = 24 valence electrons.

Octet Rule and Extended Octet

The octet rule states that atoms tend to have eight electrons in their valence shell. Some elements can have expanded octets.

• Octet Rule: Applies to C, N, O, F, and most main-group elements.

• Extended Octet: Elements in period 3 or higher (e.g., P, S, Cl) can have more than 8 electrons.

Formal Charge and Resonance

Formal charge helps evaluate the stability of Lewis structures. Resonance structures represent delocalized bonding.

• Formal Charge Formula:

$\text{Formal Charge} = (\text{Valence electrons}) - (\text{Nonbonding electrons}) - \frac{1}{2}(\text{Bonding electrons})$

• Resonance: When more than one valid Lewis structure exists, the actual structure is a hybrid.

Example: The carbonate ion (CO32-) has three equivalent resonance forms.

Bond Order, Bond Length, and Bond Energy

Bond order correlates with bond length and bond energy.

• Bond Order: Single (1), double (2), triple (3).

• Bond Length: Decreases as bond order increases.

• Bond Energy: Increases as bond order increases.

Energy Changes: Breaking bonds is endothermic (absorbs energy); forming bonds is exothermic (releases energy).

Chemical Nomenclature

Classification of Compounds

Compounds are classified as ionic or molecular, which determines their naming conventions.

• Ionic Compounds: Composed of metals and nonmetals or polyatomic ions.

• Molecular Compounds: Composed of nonmetals only.

Naming Elements and Compounds

• Monoatomic Elements: Elements that exist as single atoms (e.g., Na, Fe).

• Diatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2.

• Binary Molecular Compounds: Use prefixes (mono-, di-, tri-, etc.) to indicate the number of each atom.

• Common Ions: Memorize names and formulas of common monoatomic and polyatomic ions (e.g., Na+, SO42-).

• Ionic Compounds: Name cation first, then anion; include Roman numerals for transition metals if necessary.

• Acids: Binary acids (e.g., HCl) and polyatomic acids (e.g., H2SO4).

Example Table: Common Polyatomic Ions

Name	Formula	Charge
Sulfate	SO4	2-
Nitrate	NO3	1-
Phosphate	PO4	3-
Ammonium	NH4	1+
Hydroxide	OH	1-

Example: Na2SO4 is named sodium sulfate; CO2 is carbon dioxide.

Additional info: This guide integrates key concepts from chemical bonding, molecular structure, and nomenclature, providing a foundation for understanding molecular properties and chemical reactivity.