Molecular Compounds and Covalent Bonds

Overview of Molecular Compounds

Molecular compounds are chemical compounds composed of molecules held together by covalent bonds. Unlike ionic compounds, which are formed from ions, molecular compounds consist of nonmetal atoms sharing electrons.

• Molecular compounds are typically formed between two or more nonmetal atoms.

• They are held together by covalent bonds, which involve the sharing of electrons.

• These compounds are often found as low-melting solids, liquids, or gases and generally do not conduct electricity.

Comparison: Molecular vs. Ionic Compounds

Molecular and ionic compounds differ in their composition, structure, and physical properties.

Covalent Bonds

Nature and Formation of Covalent Bonds

A covalent bond is formed when two atoms share one or more pairs of electrons. This type of bonding allows atoms to achieve a stable electron configuration, often resembling that of noble gases.

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

• Main group elements typically achieve an octet (eight valence electrons) by sharing electrons, except hydrogen, which achieves a duet (two electrons).

• Example: In a water molecule (H2O), two hydrogen atoms and one oxygen atom share electrons to form covalent bonds.

Forces in Covalent Bond Formation

When atoms approach each other, both attractive and repulsive forces are present:

• Repulsive: Nucleus-nucleus and electron-electron repulsions.

• Attractive: Nucleus-electron attractions.

• A covalent bond forms when the attractive forces outweigh the repulsive forces, resulting in a stable molecule.

Bond Length and Bond Energy

The bond length is the optimal distance between the nuclei of two bonded atoms where the potential energy is minimized. The bond energy is the energy required to break the bond.

The Octet Rule and Exceptions

The Octet Rule

The octet rule states that atoms tend to form bonds until they are surrounded by eight valence electrons, achieving a noble gas configuration.

• Hydrogen achieves a duet (2 electrons).

• Carbon, nitrogen, oxygen, and fluorine typically obey the octet rule.

Exceptions to the Octet Rule

• Incomplete octet: Elements like boron may have fewer than 8 electrons.

• Expanded octet: Elements in period 3 or below (e.g., phosphorus, sulfur) can have more than 8 electrons.

• Odd-electron species (radicals): Molecules with an odd number of electrons.

Multiple Covalent Bonds

Single, Double, and Triple Bonds

• Single bond: Sharing of one electron pair ($2$ electrons).

• Double bond: Sharing of two electron pairs ($4$ electrons).

• Triple bond: Sharing of three electron pairs ($6$ electrons).

Examples:

• Oxygen molecule (O2): Double bond.

• Nitrogen molecule (N2): Triple bond.

• Ethylene (C2H4): Double bond between carbons.

• Acetylene (C2H2): Triple bond between carbons.

Coordinate Covalent Bonds

A coordinate covalent bond forms when both electrons in a shared pair come from the same atom. This is common in complex ions and coordination compounds.

• Example: The ammonium ion (NH4+) contains a coordinate covalent bond formed when ammonia (NH3) donates a lone pair to a proton (H+).

Representing Molecular Compounds

Molecular and Structural Formulas

• Molecular formula: Shows the number and type of atoms (e.g., H2O).

• Structural formula: Shows how atoms are connected using lines for bonds.

• Condensed formula: A compact version omitting some bonds (e.g., CH3CH2OH).

• Ball-and-stick model: Atoms as spheres, bonds as sticks.

• Space-filling model: Atoms as overlapping spheres, showing the molecule's shape.

Lewis Structures

Lewis structures show the arrangement of atoms, bonds, and lone pairs of electrons in a molecule.

• Lone pairs are non-bonding pairs of electrons shown as dots.

• Bonding pairs are shown as lines between atoms.

Common bonding patterns:

• H: 1 bond

• C: 4 bonds

• N: 3 bonds, 1 lone pair

• O: 2 bonds, 2 lone pairs

• Halogens: 1 bond, 3 lone pairs

Steps for Drawing Lewis Structures

1. Arrange atoms (least electronegative in the center; H and F are always terminal).

2. Count total valence electrons (adjust for ions).

3. Draw single bonds between central and surrounding atoms.

4. Complete octets of surrounding atoms with lone pairs.

5. Place remaining electrons on the central atom.

6. If the central atom lacks an octet, form multiple bonds as needed.

Molecular Shape and the VSEPR Model

Valence-Shell Electron-Pair Repulsion (VSEPR) Theory

The VSEPR model predicts molecular shapes based on the repulsion between electron pairs (bonding and lone pairs) around a central atom.

1. Draw the Lewis structure and identify the central atom.

2. Count the number of electron charge clouds (bonding and lone pairs) around the central atom.

3. Predict the shape by arranging charge clouds as far apart as possible.

Bond Polarity and Electronegativity

Electronegativity

Electronegativity is the ability of an atom to attract electrons in a covalent bond. It increases across a period and decreases down a group in the periodic table. Fluorine is the most electronegative element (value = 4.0).

Polar and Nonpolar Covalent Bonds

• Nonpolar covalent bond: Electrons are shared equally (e.g., H2, Cl2).

• Polar covalent bond: Electrons are shared unequally due to differences in electronegativity (e.g., HCl).

• Ionic bond: Electrons are transferred (difference in electronegativity ≥ 2.0).

Electronegativity difference guidelines:

• 0.0 – 0.4: Nonpolar covalent

• 0.5 – 1.9: Polar covalent

• ≥ 2.0: Ionic

Molecular Polarity

A molecule is polar if it contains polar bonds arranged asymmetrically, resulting in a net dipole moment. Symmetrical molecules (e.g., CO2, CCl4) may have polar bonds but are nonpolar overall due to cancellation of dipoles.

• Polarity affects melting/boiling points and solubility.

• Dipole moments are represented by arrows pointing toward the negative end.

Naming Binary Molecular Compounds

Nomenclature Rules

• The less electronegative element is named first.

• Prefixes indicate the number of each atom (mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-).

• The prefix mono- is omitted for the first element.

• The second element is named with the suffix -ide.

Examples:

• CO2: Carbon dioxide

• N2O5: Dinitrogen pentoxide

• SF6: Sulfur hexafluoride

Summary

• Covalent bonds involve the sharing of electrons between nonmetal atoms.

• Molecular compounds have distinct properties compared to ionic compounds, including lower melting/boiling points and lack of electrical conductivity.

• The octet rule guides bonding, but exceptions exist.

• Lewis structures and the VSEPR model help predict molecular structure and shape.

• Electronegativity differences determine bond polarity, which in turn affects molecular polarity and physical properties.