Molecular Compounds, Lewis Structures, Bond Polarity, and Intermolecular Forces

Molecular Compounds: Sharing Electrons

Covalent Bonds and Molecular Compounds

Molecular compounds are formed when atoms of two or more nonmetals share electrons, resulting in covalent bonds. The shared valence electrons allow each atom to achieve stability, often reaching an octet configuration. A molecule is a discrete group of atoms in a definite proportion.

• Covalent bond: A chemical bond formed by the sharing of electron pairs between atoms.

• Molecule: A group of atoms bonded together in a specific ratio.

• Example: Water (H2O) is a molecular compound formed by sharing electrons between hydrogen and oxygen.

Naming Molecular Compounds

The names of molecular compounds require prefixes because multiple compounds can be formed from the same two nonmetals. The first nonmetal is named by its element name, and the second nonmetal is named using the first syllable followed by "ide." Prefixes indicate the number of atoms present.

• Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-

• Example: NCl3 is named nitrogen trichloride.

• Rule: If the subscript is one for the first element, no prefix is used.

Writing Formulas from Names

To write the formula from the name of a molecular compound, use the element symbols and add subscripts according to the prefixes.

• Example: Diphosphorus pentoxide is written as P2O5.

Identifying Ionic vs. Covalent Compounds

A compound is usually ionic if the first element is a metal or a polyatomic ion, and covalent if the first element is a nonmetal.

• Ionic example: Potassium oxide (K2O)

• Covalent example: Dinitrogen oxide (N2O)

Lewis Structures for Molecules and Polyatomic Ions

Drawing Lewis Structures

Lewis structures represent molecules by showing the arrangement of valence electrons to achieve octets. Shared electrons (bonding pairs) are shown as dots or lines between atoms, while lone pairs are placed on the outside.

• Bonding pairs: Shared electrons between atoms.

• Lone pairs: Nonbonding electrons on atoms.

• Example: The Lewis structure for fluorine (F2) shows each atom sharing one electron to form a covalent bond.

Steps for Drawing Lewis Structures

1. Determine the arrangement of atoms (central atom is usually the least numerous).

2. Calculate the total number of valence electrons.

3. Attach each bonded atom to the central atom with a pair of electrons.

4. Use remaining electrons to complete octets.

5. If there are too many electrons, create double or triple bonds as needed.

Example: Drawing the Lewis structure for carbonate ion (CO32−) involves these steps:

Exceptions to the Octet Rule

Some elements do not follow the octet rule. Hydrogen requires only two electrons, while elements like phosphorus, sulfur, chlorine, bromine, and iodine can have expanded octets (10 or 12 electrons).

• Example: Sulfur hexafluoride (SF6) has 12 valence electrons around sulfur.

Electronegativity and Bond Polarity

Electronegativity

Electronegativity is the ability of an atom to attract shared electrons in a chemical bond. It increases across a period and decreases down a group. Nonmetals have high electronegativity, with fluorine being the highest.

• Electronegativity difference: Used to predict bond polarity.

Bond Polarity

The difference in electronegativity between atoms determines whether a bond is nonpolar covalent, polar covalent, or ionic.

• Nonpolar covalent: Electrons are shared equally (difference 0.0–0.4).

• Polar covalent: Electrons are shared unequally (difference 0.5–1.8).

• Ionic: Electrons are transferred (difference >1.8).

Example: HCl is a polar molecule because the dipole does not cancel out.

Shapes and Polarity of Molecules

VSEPR Theory and Molecular Shapes

Valence Shell Electron-Pair Repulsion (VSEPR) theory states that electron groups are arranged as far apart as possible around the central atom to minimize repulsion. The shape of a molecule depends on the number of atoms and lone pairs attached to the central atom.

• Linear: Two electron groups (180° bond angle).

• Trigonal planar: Three electron groups (120° bond angle).

• Bent: Two bonds and one or two lone pairs (less than 120° or 109.5°).

• Tetrahedral: Four bonds (109.5° bond angle).

• Trigonal pyramidal: Three bonds and one lone pair (109.5° bond angle).

Polarity of Molecules

A molecule is polar if it contains polar bonds in an unsymmetrical arrangement, causing dipoles that do not cancel. Nonpolar molecules have symmetrical shapes and identical polar bonds.

• Nonpolar shapes: Linear, trigonal planar, tetrahedral with identical bonds.

• Polar shapes: Bent, trigonal pyramidal, or any shape with non-identical bonds.

Intermolecular Forces in Compounds

Types of Intermolecular Forces

Intermolecular forces are attractions between molecules that affect physical properties like melting points.

• Hydrogen bonds: Strongest force, occurs between H bonded to F, O, or N and a lone pair on F, O, or N.

• Dipole–dipole attractions: Occur between polar molecules.

• Dispersion forces: Weak attractions between nonpolar molecules due to temporary dipoles.

• Ionic bonds: Strongest overall, occur between ions in ionic compounds.

Intermolecular Forces and Melting Points

The strength of intermolecular forces determines the melting points of substances:

• Dispersion forces: Lowest melting points.

• Dipole–dipole: Moderate melting points.

• Hydrogen bonds: Higher melting points.

• Ionic bonds: Highest melting points.

Summary Table: Bond Types and Intermolecular Forces

Bond Type	Electronegativity Difference	Example
Nonpolar Covalent	0.0–0.4	Cl2, H2
Polar Covalent	0.5–1.8	HCl, H2O
Ionic	>1.8	NaCl, KCl

Intermolecular Force	Relative Strength	Example
Dispersion Forces	Weakest	CH4, O2
Dipole–Dipole	Moderate	SO2, HCl
Hydrogen Bonds	Strong	H2O, NH3
Ionic Bonds	Strongest	NaCl, KCl

Additional info: Academic context was added to clarify naming rules, Lewis structure steps, exceptions to the octet rule, and intermolecular force comparisons.