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 by 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 to indicate the number of atoms of each element, as multiple compounds can be formed from the same two nonmetals.

• The first nonmetal is named by its element name.

• The second nonmetal is named using the first syllable of its name followed by ide.

• Prefixes are used to indicate the number of atoms (e.g., mono-, di-, tri-, tetra-, penta-).

• Example: NCl3 is named nitrogen trichloride.

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 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 give octets. Shared electrons (bonding pairs) are shown as dots or lines between atoms, while lone pairs are placed on the outside.

• Central atom is usually the atom with the least number of atoms in the formula.

• Hydrogen forms only one bond (needs two electrons).

• Some elements (B, P, S, Cl, Br, I) can have expanded octets.

Stepwise Example: Drawing the Lewis Structure for a Polyatomic Ion

• Step 1: Arrange atoms (central atom is C).

• Step 2: Count total valence electrons.

• Step 3: Attach bonded atoms to the central atom with pairs of electrons.

• Step 4: Complete octets with remaining electrons.

• Step 5: If there are too many electrons, create multiple bonds (double or triple).

• Step 6: For polyatomic ions, place brackets and indicate the charge.

Double and Triple Bonds

Double bonds occur when atoms share two pairs of electrons, and triple bonds occur when atoms share three pairs. These are formed when there are not enough electrons to complete octets with single bonds.

• Example: Carbon dioxide (CO2) has two double bonds.

Exceptions to the Octet Rule

Hydrogen requires only two electrons. Some nonmetals (P, S, Cl, Br, I) can have more than eight electrons (expanded octets).

• 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; metals have low.

• Highest electronegativity: Fluorine

Bond Polarity

The difference in electronegativity between atoms determines bond polarity.

• Nonpolar covalent bond: Equal sharing, difference 0.0–0.4

• Polar covalent bond: Unequal sharing, difference 0.5–1.8

• Ionic bond: Electron transfer, difference >1.8

Dipoles and Bond Polarity

A polar covalent bond has a dipole, with partial positive and negative charges. The dipole is indicated by the Greek letter delta and an arrow pointing from positive to negative.

Shapes and Polarity of Molecules

VSEPR Theory

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 depends on the number of atoms and lone pairs attached to the central atom.

• Linear: Two electron groups, bond angle 180°

• Trigonal planar: Three electron groups, bond angle 120°

• Bent: Two bonds and one or two lone pairs, bond angle <120° or <109.5°

• Tetrahedral: Four bonds, bond angle 109.5°

• Trigonal pyramidal: Three bonds and one lone pair, bond angle <109.5°

Polarity of Molecules

Molecular polarity depends on both bond polarity and molecular shape. Nonpolar molecules have symmetrical shapes (linear, trigonal planar, tetrahedral with identical bonds). Polar molecules have unsymmetrical shapes (bent, pyramidal) or non-identical bonds.

• Example: Water (H2O) is polar due to its bent shape.

Intermolecular Forces in Compounds

Types of Intermolecular Forces

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

• Hydrogen bonds: Strongest, occur 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: Weakest, occur between nonpolar molecules due to temporary dipoles.

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

Intermolecular Forces and Melting Points

Melting points are higher for substances with stronger intermolecular forces. Ionic compounds have the highest melting points, followed by hydrogen bonds, dipole–dipole, and dispersion forces.

Comparison of Bonding and Intermolecular Forces

Type of Force	Relative Strength	Example
Ionic Bonds	Very Strong	KCl
Hydrogen Bonds	Strong	H2O
Dipole–Dipole	Moderate	SO2
Dispersion Forces	Weak	CO2

Summary: Variations in Bonding

The difference in electronegativity values for two atoms predicts the type of chemical bond:

• 0–0.4: Nonpolar covalent

• 0.5–1.8: Polar covalent

• >1.8: Ionic

Bond type and molecular shape together determine the physical and chemical properties of compounds.