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.

• Example: Na (1s22s22p63s1) → Na+ (1s22s22p6)

Ionic vs. Covalent Bonding

Chemical bonds can be classified based on how electrons are distributed between atoms.

• Ionic Bonds: Complete transfer of electrons from metal to nonmetal.

• Covalent Bonds: Sharing of electrons between nonmetals.

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

Electronegativity and Bond Polarity

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

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

• Ranking Bonds by Polarity: Greater difference in electronegativity leads to more polar bonds.

• Bond Types:

  • Nonpolar Covalent: Equal sharing (e.g., Cl2)

  • Polar Covalent: Unequal sharing (e.g., HCl)

  • Ionic: Electron transfer (e.g., NaCl)

• Partial Charges and Dipole Direction: Indicated by δ+ (partial positive) and δ− (partial negative); dipole arrow points toward the more electronegative atom.

Covalent Bond Order

Bond order refers to the number of shared electron pairs between two atoms.

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

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

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

• Example: O2 has a double bond; N2 has a triple bond.

Chemical Nomenclature

Naming compounds follows systematic rules based on composition and structure.

• Naming Ionic Compounds:

  • Main group metal + nonmetal: Name cation, then anion (e.g., NaCl: sodium chloride)

  • Metals with variable charge: Indicate charge with Roman numerals (e.g., FeCl2: iron(II) chloride)

  • Polyatomic ions: Use standard names (e.g., NaNO3: sodium nitrate)

• 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 represent valence electrons as dots around element symbols. Lewis structures show bonding and lone pairs in molecules and ions.

• Drawing Lewis Structures:

  1. Count total valence electrons.

  2. Arrange atoms (central atom usually least electronegative).

  3. Connect atoms with single bonds.

  4. Distribute remaining electrons as lone pairs to complete octets.

  5. Form double/triple bonds if needed to satisfy octet rule.

• Exceptions to the Octet Rule:

  • Odd-electron species (e.g., NO)

  • Incomplete octets (e.g., B, Be)

  • Expanded octets (period 3 and beyond, e.g., SF6)

Formal Charge

Formal charge helps determine the most reasonable Lewis structure.

• Calculation:

• Choose structures with formal charges closest to zero and negative charges on more electronegative atoms.

Resonance

Some molecules have multiple valid Lewis structures (resonance forms). The actual structure is a resonance hybrid.

• Drawing Resonance Structures: Move electrons, not atoms, to show alternative bonding arrangements.

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

• Example: Ozone (O3) has two resonance structures, bond order = 1.5.

VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes 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

Molecular polarity depends on bond dipoles and molecular geometry.

• Bond Dipoles: Arise from differences in electronegativity.

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

• Example: CO2 (linear, nonpolar); H2O (bent, polar)

Chapter 5: Advanced Theories of Bonding

Valence Bond Theory

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

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

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

• Example: Ethene (C2H4): C=C double bond consists of one σ and one π bond.

Hybrid Atomic Orbitals

Hybridization explains molecular shapes by combining atomic orbitals into new hybrid orbitals.

• Types of Hybridization:

  • sp: linear (180°)

  • sp2: trigonal planar (120°)

  • sp3: tetrahedral (109.5°)

• Determining Hybridization: Count electron domains around atom (lone pairs + bonds).

• Example: CH4 (methane): C is sp3 hybridized.

Multiple Bonds: σ and π Bonds

Double and triple bonds consist of one σ bond and one or two π bonds, respectively.

• Counting Bonds:

  • Single bond: 1 σ

  • Double bond: 1 σ + 1 π

  • Triple bond: 1 σ + 2 π

Molecular Orbital (MO) Theory

MO theory describes bonding by combining atomic orbitals to form molecular orbitals that extend over the entire molecule.

• Bonding vs. Antibonding Orbitals:

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

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

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

• Filling MO Diagrams: Add electrons according to the Aufbau principle and Pauli exclusion principle.

• Bond Order:

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

• Paramagnetism vs. Diamagnetism:

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

  • Diamagnetic: All electrons paired

Common Ions: Names, Formulas, and Charges

The following table lists common ions that students should know by name, formula, and charge.

Additional info: Students should be able to write names and formulas for these ions, recognize their charges, and use them in naming and writing formulas for ionic compounds.