6.5 Molecular Compounds: Sharing Electrons

Introduction to Molecular Compounds

Molecular compounds are formed when two or more nonmetal atoms 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 electrons allows each atom to achieve a stable electron configuration, often resembling that of the nearest noble gas.

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

• Molecule: A discrete group of atoms held together by covalent bonds in a definite proportion.

Naming Molecular Compounds

The names of molecular compounds require prefixes to indicate the number of each type of atom present, as several different 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 (except 'mono' is never used for the first atom).

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

Example: CO2 is carbon dioxide; CO is carbon monoxide.

Flowchart for Naming Compounds

To determine whether a compound is ionic or molecular and how to name it, use the following flowchart:

6.6 Lewis Structures for Molecules and Polyatomic Ions

Lewis Structures

A Lewis structure represents the arrangement of valence electrons among atoms in a molecule. Shared pairs (bonding pairs) are shown as lines or pairs of dots, while lone pairs (nonbonding pairs) are shown as dots on individual atoms.

• Atoms tend to achieve an octet (eight valence electrons) for stability (the octet rule).

• Hydrogen is an exception, requiring only two electrons.

Steps for Drawing Lewis Structures:

1. Sum the valence electrons for all atoms (adjust for ions).

2. Arrange atoms (most metallic or least electronegative in the center; H is always terminal).

3. Connect atoms with single bonds.

4. Complete octets for outer atoms, then central atom.

5. If needed, form double or triple bonds to complete octets.

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

Diatomic Molecules

Certain elements exist naturally as diatomic molecules (two atoms bonded together): H2, N2, O2, F2, Cl2, Br2, I2.

Bond Types: Single, Double, and Triple Bonds

• Single bond: One pair of shared electrons.

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

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

Exceptions to the Octet Rule

• Hydrogen requires only two electrons.

• Elements like P, S, Cl, Br, and I can have expanded octets (10 or 12 electrons).

6.7 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 in the periodic table. Fluorine is the most electronegative element.

Bond Polarity

The difference in electronegativity between two atoms determines the type of bond:

6.8 Shapes and Polarity of Molecules

VSEPR Theory and Molecular Geometry

Valence Shell Electron-Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on the repulsion between electron groups around a central atom. Electron groups are arranged as far apart as possible to minimize repulsion.

• Linear: 2 electron groups, 180° bond angle (e.g., CO2).

• Trigonal planar: 3 electron groups, 120° bond angle (e.g., H2CO).

• Bent: 3 electron groups (1 lone pair), 120° bond angle (e.g., SO2); or 4 electron groups (2 lone pairs), 109° bond angle (e.g., H2O).

• Tetrahedral: 4 electron groups, 109° bond angle (e.g., CH4).

• Trigonal pyramidal: 4 electron groups (1 lone pair), 109° bond angle (e.g., NH3).

Examples of Molecular Shapes

• CH4 (Methane): Tetrahedral shape, 109° bond angles.

• NH3 (Ammonia): Trigonal pyramidal shape, 109° bond angles.

• H2O (Water): Bent shape, 109° bond angles.

Polarity of Molecules

A molecule is nonpolar if it contains only nonpolar bonds or if the dipoles in polar bonds cancel due to symmetry (e.g., CO2, CCl4). A molecule is polar if it contains polar bonds that do not cancel, resulting in a partial negative and partial positive end (e.g., H2O, NH3).

Summary Table: Molecular Shape and Polarity

Key Points:

• Use prefixes to name molecular compounds and indicate the number of each atom.

• Draw Lewis structures to visualize bonding and lone pairs.

• Apply VSEPR theory to predict molecular shapes and bond angles.

• Determine bond polarity using electronegativity differences.

• Assess molecular polarity by considering both bond polarity and molecular geometry.