Chemical Bonding: Molecular Compounds and Lewis Structures

Chemical Bonding: Molecular Compounds and Lewis Structures

Lecture Objectives

This section introduces the key learning goals for understanding molecular compounds and their representations. By the end of this topic, students should be able to:

• Write the name and formula for molecular compounds.

• Draw Lewis structures for molecular compounds with single and multiple bonds.

Molecular Compounds: Sharing Electrons

Introduction to Molecular Compounds

Molecular compounds are formed when atoms of two or more nonmetals share electrons, resulting in the formation of covalent bonds. This sharing allows each atom to achieve a stable electron configuration, often resembling that 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 discrete group of atoms held together by covalent bonds.

• There are many more molecular compounds than ionic compounds.

Example: Water (H2O) is a molecular compound where each hydrogen atom shares an electron with oxygen.

Common Molecular Compounds and Their Uses

The following table lists several common molecular compounds and their commercial uses:

Formula	Name	Commercial Uses
CO2	Carbon dioxide	Fire extinguishers; dry ice; propellant in aerosols; carbonation of beverages
CS2	Carbon disulfide	Manufacture of rayon
N2O	Dinitrogen oxide	Inhalation anesthetic, "laughing gas"
NO	Nitrogen oxide	Stabilizer, biochemical messenger in cells
SO2	Sulfur dioxide	Preserving fruits, vegetables; disinfectant in breweries; bleaching textiles
SF6	Sulfur hexafluoride	Electrical circuits
SO3	Sulfur trioxide	Manufacture of explosives

Naming Molecular Compounds

Rules for Naming

Molecular compounds are named using a systematic approach to indicate the number and type of atoms present:

• Name the first nonmetal by its element name.

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

• Add prefixes to indicate the number of atoms (subscripts) for each element.

Prefixes for Number of Atoms:

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

Example: CO2 is named carbon dioxide; N2O is dinitrogen oxide.

Writing Formulas for Molecular Compounds

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

• Write the chemical symbols in the order of the elements in the name.

• Write any prefixes as subscripts to indicate the number of atoms.

Example: Dinitrogen trioxide is N2O3; sulfur hexafluoride is SF6.

Classifying Compounds: Ionic vs. Molecular

How to Distinguish Between Ionic and Molecular Compounds

• Ionic Compounds: The first element in the formula or name is a metal (or the polyatomic ion NH4+).

• Molecular Compounds: The first element in the formula is a nonmetal.

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

Summary: Naming Flowchart

The following flowchart summarizes the process for naming chemical compounds:

• If the compound contains a metal (or ammonium), it is ionic. Name the cation and anion accordingly.

• If the compound contains only nonmetals, it is molecular. Use prefixes to indicate the number of atoms and add "ide" to the second element.

Additional info: The flowchart visually distinguishes between the steps for naming ionic and molecular compounds, including the use of Roman numerals for transition metals and the use of prefixes for molecular compounds.

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 octets (except for hydrogen, which achieves a duet).

• Valence electrons are shown as dots around the chemical symbols.

• 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.

Example: The Lewis structure for F2 shows each fluorine atom sharing one pair of electrons, with three lone pairs on each atom.

Steps for Drawing Lewis Structures

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

2. Determine 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. Use the remaining electrons to complete the octets of the outer atoms, then the central atom.

Example: For PCl3, phosphorus is the central atom, and each chlorine is attached with a single bond. Remaining electrons complete the octets.

Multiple Bonds in Lewis Structures

Some molecules require double or triple bonds to satisfy the octet rule for all atoms. This occurs when there are not enough electrons to complete octets with only single bonds.

• Double bond: Two pairs of electrons are shared between two atoms.

• Triple bond: Three pairs of electrons are shared between two atoms.

• Atoms of nitrogen (N), oxygen (O), and carbon (C) are most likely to form multiple bonds.

Example: In CO2, carbon forms two double bonds with oxygen to complete its octet:

$O = C = O$

Additional info: If the central atom does not have a complete octet after assigning single bonds and lone pairs, convert lone pairs from surrounding atoms into additional bonds as needed.