Chemical Bonding: Molecular and Lewis Structures

Introduction to Molecular Compounds

Molecular compounds are formed when atoms of two or more nonmetals share electrons, resulting in covalent bonds. These compounds are distinct from ionic compounds, which involve the transfer of electrons between metals and nonmetals. The sharing of valence electrons allows nonmetal atoms to achieve stable electron configurations, often resembling those of noble gases.

• Covalent bond: A chemical bond formed by the sharing of one or more pairs of electrons between two nonmetal atoms.

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

• Molecular compounds are more numerous than ionic compounds due to the variety of possible combinations among nonmetals.

Example: Water (H2O), carbon dioxide (CO2), and ammonia (NH3) are common molecular compounds.

Naming Molecular Compounds

Naming molecular compounds follows a systematic approach to indicate the types and numbers of atoms present in a molecule. Prefixes are used to denote the number of each type of atom.

• Step 1: Name the first nonmetal by its element name.

• Step 2: Name the second nonmetal by using the first syllable of its element name followed by "-ide".

• Step 3: Add prefixes to indicate the number of atoms (subscripts) for both elements.

Prefixes for Number of Atoms:

Example:

• CO: carbon monoxide

• N2O: dinitrogen monoxide

• PF3: phosphorus trifluoride

• CCl4: carbon tetrachloride

Writing Formulas for Molecular Compounds

To write the formula from the name of a molecular compound:

• Step 1: Write the chemical symbols in the order of the elements in the name.

• Step 2: Use the prefixes to determine the subscripts for each element.

Example:

• Phosphorus pentachloride: PCl5

• Dinitrogen trioxide: N2O3

• Sulfur hexafluoride: SF6

Distinguishing Ionic and Molecular Compounds

It is important to distinguish between ionic and molecular compounds based on their composition:

• Ionic Compounds: The first element is a metal or the polyatomic ion NH4+.

• Molecular Compounds: The first element is a nonmetal.

Example: K2O (potassium oxide) is ionic; N2O (dinitrogen oxide) is molecular.

Summary: Naming Flowchart

The following flowchart summarizes the process for naming chemical compounds, distinguishing between ionic and molecular compounds and the rules for each.

Lewis Structures for Molecules

Introduction to Lewis Structures

Lewis structures are diagrams that represent the arrangement of valence electrons among atoms in a molecule. They help visualize how atoms share electrons to achieve stable electron configurations (octets), except for hydrogen, which achieves a duet.

• Bonding pairs (shared electrons) are shown as two dots or a single line between atoms.

• Lone pairs (non-bonding electrons) are placed on the outside of atoms.

Drawing Lewis Structures: Steps

To draw a Lewis structure for a molecule:

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

2. Calculate the total number of valence electrons for all atoms.

3. Attach each bonded atom to the central atom with a pair of electrons (single bond).

4. Distribute remaining electrons to complete octets (or duets for hydrogen).

Example: Lewis structures for CH4, NH3, and H2O.

Example: Drawing Lewis Structure for PCl3

Phosphorus trichloride (PCl3) is a molecular compound with phosphorus as the central atom. Each chlorine atom forms a single bond with phosphorus, and lone pairs are placed to complete the octets.

Diatomic Molecules

Diatomic molecules consist of two atoms of the same element bonded together, such as O2, N2, or F2. These molecules are often used as examples for illustrating covalent bonding and Lewis structures.

Double and Triple Bonds

Some molecules require double or triple bonds to satisfy the octet rule for all atoms. Double bonds involve the sharing of two pairs of electrons, while triple bonds involve three pairs. Atoms such as carbon, nitrogen, oxygen, and sulfur commonly form multiple bonds.

• Double bond: Two pairs of electrons are shared between two atoms (e.g., O2, CO2).

• Triple bond: Three pairs of electrons are shared (e.g., N2).

• Multiple bonds form when there are not enough valence electrons to complete octets with single bonds alone.

Example: The Lewis structure for carbon dioxide (CO2) shows double bonds between carbon and each oxygen atom.

Appendix: Commercial Uses of Common Molecular Compounds

References: Timberlake, K. (2018). Chemistry: An Introduction to General, Organic, and Biological Chemistry (13th ed.). Pearson Education.