Chapter 4: Molecules and Compounds

Polar and Nonpolar Covalent Bonds; Electronegativity

Chemical bonds can be classified based on the sharing or transfer of electrons between atoms. Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond.

• Nonpolar covalent bond: Electrons are shared equally between atoms (usually identical or similar electronegativities).

• Polar covalent bond: Electrons are shared unequally due to a difference in electronegativity between atoms.

• Electronegativity trend: Increases across a period (left to right) and decreases down a group in the periodic table.

Example: In H2O, the O–H bonds are polar because oxygen is more electronegative than hydrogen.

Lewis Structures of Molecular Compounds and Polyatomic Ions

Lewis structures represent the arrangement of valence electrons among atoms in a molecule or ion.

• Count the total number of valence electrons (add or subtract electrons for ions).

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

• Distribute electrons to form single bonds, then complete octets (or duets for H).

Example: For CO32−, sum valence electrons: C (4) + 3×O (3×6) + 2 (for charge) = 24 electrons.

Assigning Formal Charges

Formal charge helps determine the most stable Lewis structure.

• Formula:

• The best structure has formal charges closest to zero and negative charges on the most electronegative atoms.

Exceptions to the Octet Rule

Some molecules do not follow the octet rule:

• Incomplete octet: e.g., BeCl2, BF3

• Expanded octet: Elements in period 3 or higher (e.g., SF6, PCl5)

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

Bond Energy, Bond Strength, and Bond Length

Bond energy is the energy required to break one mole of a bond in a molecule. Bond strength is related to bond energy, and bond length is the distance between nuclei of bonded atoms.

• Stronger bonds have higher bond energies and shorter bond lengths.

• Triple bonds > double bonds > single bonds (in strength and energy).

Example: (triple bond) is stronger and shorter than (double bond).

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shape of molecules based on electron group repulsions around a central atom.

• Electron groups: Bonding pairs and lone pairs around the central atom.

• Common geometries:

  • Linear (2 groups, 180°)

  • Trigonal planar (3 groups, 120°)

  • Tetrahedral (4 groups, 109.5°)

  • Trigonal bipyramidal (5 groups, 90°, 120°)

  • Octahedral (6 groups, 90°)

  • Bent, trigonal pyramidal (due to lone pairs)

Example: NH3 is trigonal pyramidal due to one lone pair on nitrogen.

Dipole Moment, Bond Polarity, and Molecular Polarity

Dipole moment is a measure of molecular polarity, resulting from unequal sharing of electrons.

• Bond polarity depends on electronegativity difference.

• Molecular polarity depends on both bond polarity and molecular geometry.

Example: CO2 has polar bonds but is nonpolar overall due to its linear geometry.

Chapter 5: Chemical Bonding I

Valence Bond Theory and Hybridization

Valence bond theory describes covalent bond formation as the overlap of atomic orbitals. Hybridization is the mixing of atomic orbitals to form new, equivalent hybrid orbitals.

• Types of hybridization:

  • sp (linear, 2 groups)

  • sp2 (trigonal planar, 3 groups)

  • sp3 (tetrahedral, 4 groups)

  • sp3d (trigonal bipyramidal, 5 groups)

  • sp3d2 (octahedral, 6 groups)

Example: In methane (CH4), carbon is sp3 hybridized.

Sigma (σ) and Pi (π) Bonding

Sigma (σ) bonds are formed by head-on overlap of orbitals; pi (π) bonds are formed by side-to-side overlap.

• Single bonds: 1 σ bond

• Double bonds: 1 σ + 1 π bond

• Triple bonds: 1 σ + 2 π bonds

Formation of Ions and Ionic Bonds

Ionic bonds form when electrons are transferred from a metal to a nonmetal, creating cations and anions.

• Metals lose electrons to form cations; nonmetals gain electrons to form anions.

• The resulting electrostatic attraction forms the ionic bond.

Naming Ionic Compounds

Rules for naming ionic compounds:

• Name the cation first, then the anion.

• For transition metals, indicate the charge with Roman numerals (e.g., FeCl2: iron(II) chloride).

• For polyatomic ions, use the ion's name (e.g., NaNO3: sodium nitrate).

Molecular Formula and Empirical Formula

Molecular formula shows the actual number of atoms of each element in a molecule. Empirical formula shows the simplest whole-number ratio of atoms.

• Example: Glucose (C6H12O6) has the empirical formula CH2O.

Ionic Compounds of Transition Metals and Their Naming

Transition metals can form ions with different charges. The charge is specified in the compound's name.

• Example: CuO is copper(II) oxide; Cu2O is copper(I) oxide.

Polyatomic Ions and Naming Compounds with Polyatomic Ions

Common polyatomic ions include:

When naming, use the name of the cation followed by the polyatomic ion (e.g., CaCO3: calcium carbonate).

Hydrates and Naming Hydrates

Hydrates are ionic compounds with water molecules incorporated into their structure.

• Name the ionic compound, then add a prefix for the number of water molecules and the word 'hydrate' (e.g., CuSO4·5H2O: copper(II) sulfate pentahydrate).

Lattice Energy

Lattice energy is the energy released when gaseous ions form an ionic solid. It is a measure of the strength of the ionic bond.

• Increases with higher charge and smaller ionic radius.

• Formula: (where and are ion charges, is the distance between ions, is a constant)

Lewis Dot Symbols

Lewis dot symbols represent valence electrons as dots around the element symbol.

• Used to predict bonding and the formation of ions.

Formation of Covalent Bonds; Bonding and Nonbonding Electrons; Diatomic Molecules

Covalent bonds form when atoms share electrons. Bonding electrons are shared between atoms; nonbonding electrons (lone pairs) are not shared.

• Diatomic molecules: Molecules composed of two atoms (e.g., O2, N2).

Naming Covalent Compounds and Acids

For covalent (molecular) compounds:

• Use prefixes to indicate the number of each atom (e.g., CO2: carbon dioxide).

• For acids:

  • If the anion ends in -ide: hydro-...-ic acid (e.g., HCl: hydrochloric acid)

  • If the anion ends in -ate: ...-ic acid (e.g., H2SO4: sulfuric acid)

  • If the anion ends in -ite: ...-ous acid (e.g., H2SO3: sulfurous acid)

Properties of Ionic and Covalent Compounds