Chapter 4: Molecular Compounds – Structure, Bonding, and Properties

• Chapter 4: Molecular Compounds

Metals, Nonmetals, and Metalloids

Classification of Elements

Elements are classified as metals, nonmetals, or metalloids based on their properties and position in the periodic table. This classification is important for predicting the type of bonding that will occur in compounds.

• Metals: Good conductors of heat and electricity, malleable, ductile, and tend to lose electrons to form cations.

• Nonmetals: Poor conductors, brittle, and tend to gain electrons to form anions or share electrons in covalent bonds.

• Metalloids: Exhibit properties intermediate between metals and nonmetals.

Why Do Compounds Form?

The Octet Rule and Noble Gas Configuration

The driving force for compound formation is the attainment of a noble gas electronic configuration (usually eight valence electrons, known as the octet rule). Atoms exchange or share electrons during chemical reactions to achieve this stable configuration.

• Octet Rule: Atoms tend to gain, lose, or share electrons to have eight electrons in their valence shell.

• Stability: Noble gases are stable due to their full valence shells.

Types of Bonding Models

Ionic vs. Covalent Bonding

• Ionic Bonding: Occurs between a metal and a nonmetal. Involves a complete transfer of electrons from the metal to the nonmetal, forming cations and anions.

• Covalent Bonding: Occurs between two nonmetals. Involves sharing of electron pairs so that both atoms achieve an octet.

Properties of Ionic and Molecular Compounds

• Ionic Compounds: Usually solid at room temperature, high melting and boiling points, conduct electricity when molten or dissolved in water.

• Molecular (Covalent) Compounds: Can be solids, liquids, or gases; usually have low melting and boiling points; do not conduct electricity in aqueous solution; often soluble in nonpolar solvents.

Covalent Bonding

Formation and Types of Covalent Bonds

Covalent bonds form when two nonmetal atoms share electrons to achieve a noble gas configuration. The shared electrons are located in the space between the two nuclei.

• Single Bond: Sharing of one pair of electrons (e.g., H2).

• Double Bond: Sharing of two pairs of electrons (e.g., O2, C2H4).

• Triple Bond: Sharing of three pairs of electrons (e.g., N2).

Example: In a hydrogen molecule (H2), each hydrogen atom shares its single electron, resulting in a shared pair (single bond) and a stable configuration.

Lewis Dot Structures

Lewis Dot Symbols for Main Group Elements

Lewis dot symbols represent the valence electrons of an atom as dots around the element symbol. For main group elements (Periods 2 and 3), the number of dots corresponds to the group number.

Steps for Drawing Lewis Structures (Octet Rule)

1. Decide which atom is the central atom (usually the least electronegative, often listed first in the formula).

2. Count the total number of valence electrons (sum of group numbers for each atom, adjusted for charge in ions).

3. Connect the central atom to peripheral atoms with single bonds (each bond = 2 electrons).

4. Distribute remaining electrons to complete octets on peripheral atoms, then on the central atom.

5. If the central atom lacks an octet, form double or triple bonds as needed by sharing lone pairs from peripheral atoms.

Lewis Structures for Polyatomic Ions

• For anions: Add electrons equal to the negative charge.

• For cations: Subtract electrons equal to the positive charge.

Example: For NO3-, add one electron to the total count for the negative charge.

Isoelectronic Molecules

Molecules or ions with the same number of electrons and similar structures are called isoelectronic. For example, CO2 is isoelectronic with CS2.

Shapes of Molecules: VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory states that electron pairs (bonding and lone pairs) around a central atom arrange themselves to minimize repulsion, determining the molecule's shape.

Charge Clouds

• A charge cloud is any region of electron density (single, double, triple bond, or lone pair).

Electronic Geometries

Number of Charge Clouds	Electronic Geometry
2	Linear
3	Trigonal Planar
4	Tetrahedral

Significant Bond Angles

• Linear: 180°

• Trigonal Planar: 120°

• Tetrahedral: 109.5°

Bond Polarity and Electronegativity

Bond Polarity

Bond polarity arises when two atoms in a covalent bond have different electronegativities, causing unequal sharing of electrons and creating a dipole moment.

• Nonpolar Covalent Bond: Equal sharing of electrons (e.g., H2).

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

• Ionic Bond: Complete transfer of electrons (e.g., NaCl).

Electronegativity

Electronegativity is the ability of an atom to attract shared electrons in a covalent bond. The Pauling scale assigns values from 0 to 4 (fluorine is the most electronegative).

• Electronegativity increases across a period (left to right) and decreases down a group.

Predicting Bond Type by Electronegativity Difference

Electronegativity Difference (ΔEN)	Bond Type
ΔEN > 2.0	Ionic
0.4 < ΔEN < 2.0	Polar Covalent
ΔEN ≤ 0.4	Nonpolar Covalent

Molecular Polarity

Determining Molecular Polarity

A molecule is polar if it has a net dipole moment, resulting from the vector sum of all bond dipoles and the molecular geometry. If the bond dipoles cancel, the molecule is nonpolar.

• Both the presence of polar bonds and the shape of the molecule must be considered.

• Linear, trigonal planar, and tetrahedral molecules with identical peripheral atoms are usually nonpolar.

• Angular (bent) and trigonal pyramidal molecules are usually polar.

Examples of Molecular Polarity

Formula	Molecular Geometry	Polar/Nonpolar
CO2	Linear	Nonpolar
H2O	Bent (Angular)	Polar
BF3	Trigonal Planar	Nonpolar
NH3	Trigonal Pyramidal	Polar
CH4	Tetrahedral	Nonpolar
CHCl3	Tetrahedral	Polar

Naming Binary Covalent Compounds

Rules for Naming

• Name the element closest to the metals (leftmost) in the periodic table first.

• Name the second element as the root plus "-ide".

• Use Greek prefixes to indicate the number of each atom (except "mono-" is usually omitted for the first element).

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

Example: NI3 is named nitrogen triiodide.