Chapter 4: Chemical Bonding and Molecular Geometry

Ionic Bonding

Ionic bonding involves the transfer of electrons from one atom (typically a metal) to another (typically a nonmetal), resulting in the formation of oppositely charged ions that are held together by electrostatic attraction.

• Predicting Charges of Monatomic Ions: The charge of a main-group element's ion can often be predicted from its group number in the periodic table.

  • Group 1: +1 (e.g., Na+)

  • Group 2: +2 (e.g., Mg2+)

  • Group 13: +3 (e.g., Al3+)

  • Group 15: -3 (e.g., N3−)

  • Group 16: -2 (e.g., O2−)

  • Group 17: -1 (e.g., Cl−)

• Electron Configurations of Ions: Ions achieve noble gas configurations by losing or gaining electrons. For example, Na (1s22s22p63s1) loses one electron to become Na+ (1s22s22p6).

Ionic vs. Covalent Bonding

• Ionic Bonding: Involves electron transfer from metal to nonmetal, forming cations and anions.

• Covalent Bonding: Involves sharing of electrons between nonmetals.

• Key Difference: Ionic bonds result from electrostatic attraction; covalent bonds result from shared electron pairs.

Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract electrons in a chemical bond.

• Periodic Trend: Electronegativity increases across a period and decreases down a group.

• Ranking Bonds by Polarity: The greater the difference in electronegativity, the more polar the bond.

• Bond Types:

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

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

  • Ionic: Electrons transferred (e.g., NaCl).

• Partial Charges and Dipole Direction: Use δ+ and δ− to indicate partial charges; arrow points toward the more electronegative atom.

Covalent Bond Order

• Single Bond: One shared pair of electrons (bond order = 1).

• Double Bond: Two shared pairs (bond order = 2).

• Triple Bond: Three shared pairs (bond order = 3).

Chemical Nomenclature

• Naming Ionic Compounds:

  • Monatomic ions: Name of metal + name of nonmetal with -ide (e.g., sodium chloride).

  • Metals with variable charge: Indicate charge with Roman numerals (e.g., iron(III) chloride).

  • Polyatomic ions: Use standard names (e.g., sodium sulfate for Na2SO4).

• Naming Binary Molecular Compounds: Use prefixes (mono-, di-, tri-, etc.) to indicate number of atoms (e.g., CO2: carbon dioxide).

• Naming Acids:

  • Binary acids: "hydro-" + root + "-ic acid" (e.g., HCl: hydrochloric acid).

  • Oxyacids: Based on polyatomic ion (e.g., H2SO4: sulfuric acid).

Lewis Symbols and Structures

• Lewis Symbols: Dots represent valence electrons around atomic symbols.

• Lewis Structures: Show bonding and lone pairs in molecules and polyatomic ions.

  • Count total valence electrons.

  • Arrange atoms, connect with single bonds, complete octets, assign lone pairs.

Exceptions to the Octet Rule

• Odd-Electron Species: Molecules with an odd number of electrons (e.g., NO).

• Incomplete Octets: Atoms like B and Be may have fewer than 8 electrons.

• Expanded Octets: Atoms in period 3 or beyond can have more than 8 electrons (e.g., SF6).

Formal Charge

• Definition: Formal charge helps determine the most reasonable Lewis structure.

• Calculation:

  • Formula:

• Choosing Structures: The best Lewis structure minimizes formal charges and places negative charges on the most electronegative atoms.

Resonance

• Resonance Structures: Multiple valid Lewis structures for a molecule or ion (e.g., O3).

• Resonance Hybrid: The actual structure is an average of resonance forms, affecting bond order and length.

VSEPR Theory

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

• Electron-Pair Geometry: Arrangement of all electron domains (bonding and lone pairs).

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

• Common Geometries and Ideal Bond Angles:

  • Linear: 180°

  • Trigonal planar: 120°

  • Bent: <120° or <109.5°

  • Tetrahedral: 109.5°

  • Trigonal pyramidal: <109.5°

  • Trigonal bipyramidal: 90°, 120°

  • Seesaw, T-shaped, Octahedral, Square pyramidal, Square planar: various angles

Molecular Polarity

• Bond Dipoles: Arise from differences in electronegativity.

• Net Dipole Moment: Vector sum of bond dipoles; determines if molecule is polar or nonpolar.

• Using Geometry: Symmetrical molecules (e.g., CO2) may be nonpolar even with polar bonds.

Chapter 5: Advanced Theories of Bonding

Valence Bond Theory

Valence bond theory describes covalent bond formation as the overlap of atomic orbitals from two atoms.

• σ (Sigma) Bonds: Formed by head-on overlap; all single bonds are σ bonds.

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

Hybrid Atomic Orbitals

• sp Hybridization: Linear geometry (180°); e.g., BeCl2.

• sp2 Hybridization: Trigonal planar geometry (120°); e.g., BF3.

• sp3 Hybridization: Tetrahedral geometry (109.5°); e.g., CH4.

• Determining Hybridization: Count electron domains around the atom.

Multiple Bonds: σ and π Bonds

• Single Bond: 1 σ bond

• Double Bond: 1 σ + 1 π bond

• Triple Bond: 1 σ + 2 π bonds

Molecular Orbital (MO) Theory

• Linear Combination of Atomic Orbitals: Atomic orbitals combine to form molecular orbitals (MOs).

• Bonding vs. Antibonding MOs:

  • Bonding (σ, π): Lower energy, electrons stabilize the molecule.

  • Antibonding (σ*, π*): Higher energy, electrons destabilize the molecule.

• MO Energy-Level Diagrams: Used for homonuclear diatomic molecules (H2 to Ne2).

• Filling MO Diagrams: Fill from lowest to highest energy, obeying the Pauli exclusion principle and Hund's rule.

• Bond Order:

  • Formula:

  • Higher bond order = stronger, shorter bond; bond order of 0 = unstable molecule.

• Paramagnetism vs. Diamagnetism:

  • Paramagnetic: Unpaired electrons (e.g., O2 is paramagnetic).

  • Diamagnetic: All electrons paired.

Common Ions: Names, Formulas, and Charges

The following table lists common ions, their formulas, and charges. Knowing these is essential for nomenclature and predicting compound formulas.

Example: Naming an Ionic Compound

• Na2SO4: sodium sulfate

• FeCl3: iron(III) chloride

• NH4NO3: ammonium nitrate

Additional info:

• Practice drawing Lewis structures and predicting molecular shapes using VSEPR for common molecules and ions.

• Be familiar with the relationship between bond order, bond length, and bond strength.

• Understand how to use formal charge and resonance to determine the most stable structure.