Chapter 4: Molecular Compounds – Structured Study Notes

Chapter 4: Molecular Compounds

4.1 Covalent Bonds

Covalent bonds are fundamental to molecular compounds, formed by the sharing of electrons between atoms. This section introduces the nature of covalent bonds and distinguishes them from ionic bonds.

• Covalent Bond: A chemical bond formed when two atoms share one or more pairs of electrons.

• Molecule: A group of atoms held together by covalent bonds.

• Octet Rule: Main group elements react to achieve eight valence electrons (or two for hydrogen), resulting in a noble gas configuration.

• Electron Sharing: Nonmetals achieve an octet by sharing electrons in covalent bonds.

• Bond Formation Example: Hydrogen atoms (each with one 1s electron) combine to form a hydrogen molecule (H2) by sharing their electrons, resulting in a shared electron pair between the nuclei.

• Bond Length: The optimal distance between nuclei where attractive and repulsive forces balance.

• Diatomic Elements: Seven elements exist as diatomic molecules: nitrogen, oxygen, hydrogen, fluorine, chlorine, bromine, and iodine.

4.2 Covalent Bonds and the Periodic Table

The periodic table helps predict the number of covalent bonds an atom will form to achieve an octet. Molecular compounds consist of molecules, not ions.

• Bonding Patterns: Main group elements typically form bonds to complete their octet.

• Exceptions: Boron forms only three bonds (six electrons), and elements in the third row and below can have expanded octets due to vacant d orbitals.

4.3 Multiple Covalent Bonds

Some molecules require double or triple bonds to satisfy the octet rule. Multiple bonds involve sharing more than one pair of electrons between atoms.

• Single Bond: One electron pair shared ($\text{A--B}$).

• Double Bond: Two electron pairs shared ($\text{A==B}$).

• Triple Bond: Three electron pairs shared ($\text{A\equiv B}$).

• Common Elements: Carbon, nitrogen, and oxygen frequently form multiple bonds.

4.4 Coordinate Covalent Bonds

A coordinate covalent bond forms when both electrons in the shared pair are donated by the same atom. Once formed, it is indistinguishable from other covalent bonds.

• Unusual Patterns: Coordinate bonds can result in atoms like nitrogen with four bonds or oxygen with three bonds.

4.5 Characteristics of Molecular Compounds

Molecular compounds differ from ionic compounds in structure and properties.

• Ionic Compounds: High melting and boiling points due to strong electrostatic attractions.

• Molecular Compounds: Lower melting and boiling points; molecules are neutral and lack strong electrostatic attractions.

Property	Ionic Compounds	Molecular Compounds
Melting/Boiling Point	High	Low
Structure	Crystal lattice	Discrete molecules
Electrical Conductivity	Conducts when molten or dissolved	Does not conduct
Bonding	Electrostatic attraction	Electron sharing

4.6 Molecular Formulas and Lewis Structures

Molecular formulas indicate the types and numbers of atoms in a molecule. Lewis structures show both atom connections and lone pairs.

• Molecular Formula: Shows exact numbers and types of atoms (e.g., H2O).

• Structural Formula: Shows connections using lines for bonds.

• Lewis Structure: Shows bonds and lone-pair electrons.

4.7 Drawing Lewis Structures

Lewis structures are drawn using a systematic method based on the octet rule.

• Step 1: Count total valence electrons.

• Step 2: Draw lines between connected atoms (bond pairs).

• Step 3: Add lone pairs to complete octets (except H).

• Step 4: Place remaining electrons on the central atom.

• Step 5: If the central atom lacks an octet, form multiple bonds using lone pairs from neighboring atoms.

4.8 The Shapes of Molecules

The VSEPR model predicts molecular geometry by considering electron charge clouds around atoms.

• Step 1: Draw the Lewis structure and identify the central atom.

• Step 2: Count charge clouds (bonds + lone pairs).

• Step 3: Predict shape by maximizing distance between charge clouds.

• Bond Angles: Linear (180°), Trigonal planar (120°), Tetrahedral (109.5°).

• Lone Pairs: Cause bond angles to decrease slightly due to stronger repulsion.

4.9 Polar Covalent Bonds and Electronegativity

Bond polarity arises when electrons are shared unequally. Electronegativity measures an atom's ability to attract electrons in a bond.

• Polar Covalent Bond: Electrons are attracted more strongly by one atom.

• Electronegativity: Highest for fluorine (4.0); decreases leftward and downward in the periodic table.

• Bond Classification:

  • Nonpolar covalent: Small electronegativity difference ($\Delta EN < 0.5$)

  • Polar covalent: Moderate difference ($0.5 < \Delta EN < 1.7$)

  • Ionic: Large difference ($\Delta EN > 1.7$)

• Dipole: Partial charges at each end of a polar bond.

4.10 Polar Molecules

Molecular polarity depends on both bond polarities and molecular geometry.

• Polar Molecule: Electrons are attracted more strongly to one part of the molecule.

• Symmetry: Symmetrical molecules with polar bonds can be nonpolar overall.

4.11 Naming Binary Molecular Compounds

Binary molecular compounds are named using prefixes to indicate the number of atoms and modifying the second element with the –ide suffix.

• Formula Writing: Lesser electronegative element first; metals before nonmetals.

• Naming Steps:

  1. Name the first element, using a prefix if needed.

  2. Name the second element, modify with –ide, and add numerical prefixes.

• Prefixes: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), etc.

• Example: N2O3 is dinitrogen trioxide; GeCl4 is germanium tetrachloride.

Summary Table: Numerical Prefixes

Number	Prefix
1	mono-
2	di-
3	tri-
4	tetra-
5	penta-
6	hexa-
7	hepta-
8	octa-
9	nona-
10	deca-

Additional info: These notes expand on brief points from the original slides, providing definitions, examples, and context for GOB Chemistry students. All equations are formatted in LaTeX as required.