Molecular Compounds: Covalent Bonding, Naming, and Lewis Structures

6.5 Molecular Compounds: Sharing Electrons

Introduction to Molecular Compounds

Molecular compounds are substances formed when two or more nonmetal atoms share electrons, resulting in covalent bonds. The names of these compounds require prefixes because multiple compounds can be formed from the same two nonmetals, each with different proportions of atoms.

• Molecular compounds are discrete groups of atoms in definite proportions.

• Examples: Carbon dioxide ($CO_2$) and carbon monoxide ($CO$) are both formed from carbon and oxygen but differ in the number of oxygen atoms.

Covalent Bonds

Formation and Properties

• Covalent bonds form when atoms of two or more nonmetals share electrons to achieve stability (noble gas configuration).

• Valence electrons are shared between nonmetal atoms.

• A molecule is a discrete group of atoms held together by covalent bonds.

Names and Formulas of Molecular Compounds

Naming Rules

• The first nonmetal in the formula is named by its element name.

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

• When a subscript indicates two or more atoms of an element, a prefix is used in front of its name.

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

• When the vowels o and o or a and o appear together, the first vowel is omitted (e.g., "monoxide" not "monooxide").

Examples:

• $CO_2$ = carbon dioxide

• $CO$ = carbon monoxide

• $NO$ = nitrogen monoxide

Naming Example: $NCl_3$

1. Name the first nonmetal by its element name: nitrogen.

2. Name the second nonmetal using the first syllable plus "ide": chloride.

3. Add prefixes to indicate the number of atoms: "tri" for three chlorines. No prefix for one nitrogen.

Result: $NCl_3$ is named nitrogen trichloride.

Practice: Name $B_2O_3$

• "di" for two boron atoms, "tri" for three oxygen atoms: diboron trioxide.

Writing Formulas from Names

• Use the element symbols in the order given by the name.

• Use the prefix to determine the subscript for each element.

Example: Diphosphorus pentoxide

1. "di" = 2 phosphorus atoms ($P_2$)

2. "penta" = 5 oxygen atoms ($O_5$)

Formula: $P_2O_5$

Compounds: Ionic or Covalent?

• Ionic compounds contain a metal or the polyatomic ion $NH_4^+$ as the first element (e.g., $K_2O$ is potassium oxide, ionic).

• Covalent compounds contain only nonmetals (e.g., $N_2O$ is dinitrogen oxide, covalent).

6.6 Lewis Structures for Molecules and Polyatomic Ions

Introduction to Lewis Structures

A Lewis structure represents a molecule by showing the arrangement of valence electrons among atoms, ensuring that each atom achieves a stable electron configuration (usually an octet).

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

• Lone pairs (nonbonding electrons) are placed on the outside of atoms.

The Simplest Molecule: $H_2$

• Forms as two hydrogen atoms move closer and share electrons, forming a covalent bond.

• Each hydrogen achieves the electron configuration of helium (2 electrons).

• The molecule is more stable than individual atoms.

Lewis Structures: Examples

• Ammonia ($NH_3$): Nitrogen shares three electrons with three hydrogens, with one lone pair on nitrogen.

• Fluorine ($F_2$): Each fluorine shares one electron, forming a single bond and completing the octet for both.

Elements That Exist as Diatomic Molecules

Diatomic Molecule	Name
$H_2$	Hydrogen
$N_2$	Nitrogen
$O_2$	Oxygen
$F_2$	Fluorine
$Cl_2$	Chlorine
$Br_2$	Bromine
$I_2$	Iodine

Bonding Patterns of Nonmetals

Group	Element	Typical Number of Bonds
1A (1)	H	1 bond
3A (13)	B	3 bonds
4A (14)	C, Si	4 bonds
5A (15)	N, P	3 bonds
6A (16)	O, S	2 bonds
7A (17)	F, Cl, Br, I	1 bond

Drawing Lewis Structures: Steps

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

2. Count the total number of valence electrons (sum from 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: $PCl_3$ (Phosphorus Trichloride)

• P is the central atom; three Cl atoms are bonded to it.

• Total valence electrons: $5 + (3 \times 7) = 26$ electrons.

• Three single bonds use 6 electrons; the remaining 20 electrons are placed as lone pairs to complete octets.

Double and Triple Bonds

• Double bond: Two pairs of electrons are shared (e.g., $O_2$, $CO_2$).

• Triple bond: Three pairs of electrons are shared (e.g., $N_2$, $HCN$).

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

Exceptions to the Octet Rule

• Hydrogen requires only two electrons (duet rule).

• Some nonmetals (e.g., P, S, Cl, Br, I) can have expanded octets (10 or 12 valence electrons).

• Sulfur hexafluoride ($SF_6$) is an example where sulfur has 12 valence electrons (6 bonds).

Additional info: Lewis structures are foundational for understanding molecular geometry, polarity, and reactivity in chemistry. Mastery of these concepts is essential for further study in chemical bonding and molecular structure.