Chapter 4: Molecular Compounds – Structure, Bonding, and Properties

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Chapter 4: Molecular Compounds

4.1 Covalent Bonds

Covalent bonds are chemical bonds formed by the sharing of electrons between two nonmetal atoms. This sharing allows each atom to achieve a stable electron configuration, typically an octet (eight valence electrons), similar to noble gases. Covalent bonds are distinct from ionic bonds, which involve the transfer of electrons from one atom to another.

Molecule: A group of atoms held together by covalent bonds, forming a discrete unit.
Octet Rule: Main group elements tend to form bonds until they are surrounded by eight valence electrons (two for hydrogen).
Bond Formation Example: Two hydrogen atoms (each with one electron) share their electrons to form a hydrogen molecule (H2), achieving the electron configuration of helium.

Formation of a water molecule from hydrogen and oxygen atoms Formation of a hydrogen molecule by sharing electrons

The covalent bond results from the balance of attractive and repulsive forces between the nuclei and electron clouds of the bonded atoms. The optimal distance between nuclei is called the bond length.

Attractive and repulsive forces in a covalent bond

Chlorine forms a diatomic molecule (Cl2) by the overlap of 3p orbitals, and there are seven diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2.

Covalent bonding in chlorine Diatomic elements in the periodic table

4.2 Covalent Bonds and the Periodic Table

Molecular compounds consist of molecules, not ions. The number of covalent bonds an atom forms is generally determined by its group in the periodic table, as atoms seek to achieve an octet configuration.

Examples of covalent bonding in water, ammonia, and methane Periodic table showing typical number of covalent bonds for main group elements

Exceptions to the Octet Rule: Boron forms only three bonds (six electrons), and elements in the third period and below can have expanded octets due to available d orbitals.

Boron trifluoride, an octet exception Phosphorus pentachloride, an expanded octet example

Worked Examples: Predicting Molecular Existence and Formulas

Atoms form the number of covalent bonds needed to achieve a complete octet (or duet for hydrogen).
Electron-dot symbols can be used to visualize the sharing of electrons in covalent bonds.

Worked example: Lewis structures for various molecules Electron-dot symbols for H and F forming HF

4.3 Multiple Covalent Bonds

Some molecules require double or triple bonds to satisfy the octet rule. A single bond involves one shared pair of electrons, a double bond involves two pairs, and a triple bond involves three pairs.

Carbon, nitrogen, and oxygen commonly form multiple bonds.
Double bonds are found in molecules like O2 and CO2; triple bonds are found in N2 and HCN.

Unstable Lewis structures for CO2 and N2 without multiple bonds Double and triple bonds in CO2 and N2 1-Butene structure Identifying the double bond in 1-butene

4.4 Coordinate Covalent Bonds

A coordinate covalent bond forms when both electrons in a shared pair come from the same atom. Once formed, these bonds are indistinguishable from regular covalent bonds.

Typical vs. coordinate covalent bond formation Formation of the ammonium ion via coordinate covalent bond Formation of BF4- via coordinate covalent bond

4.5 Characteristics of Molecular Compounds

Molecular compounds differ from ionic compounds in structure, composition, and properties. The following table summarizes these differences:

Ionic Compounds	Molecular Compounds
Smallest components are ions (e.g., Na+, Cl-)	Smallest components are molecules (e.g., CO2, H2O)
Usually metal + nonmetal	Usually nonmetal + nonmetal
Crystalline solids	Gases, liquids, or low-melting solids
High melting/boiling points	Low melting/boiling points
Conduct electricity when molten/dissolved	Do not conduct electricity
Many are water soluble	Few are water soluble
Not soluble in organic liquids	Many are soluble in organic liquids

Ionic vs. molecular compounds

4.6 Molecular Formulas and Lewis Structures

A molecular formula shows the number and types of atoms in a molecule. A structural formula shows how atoms are connected, and a Lewis structure also shows lone pairs of electrons.

Lewis structures for water, ammonia, and methane

4.7 Drawing Lewis Structures

Lewis structures are drawn using a systematic method:

Count total valence electrons (add for negative ions, subtract for positive ions).
Connect atoms with single bonds (lines).
Add lone pairs to complete octets (except H).
Place remaining electrons on the central atom.
If the central atom lacks an octet, form multiple bonds as needed.

Ball-and-stick and space-filling models of molecules Common bonding patterns for C, N, O, halogens, and H Condensed and ball-and-stick representations of molecules Lewis structure for hydrogen cyanide (HCN) Stepwise construction of vinyl chloride Lewis structure Adding lone pairs to vinyl chloride Lewis structure Final Lewis structure for vinyl chloride with double bond

4.8 The Shapes of Molecules

The Valence-Shell Electron-Pair Repulsion (VSEPR) model predicts molecular geometry by considering the repulsion between electron charge clouds (bonding and lone pairs) around a central atom. The arrangement of these clouds determines the shape and bond angles of the molecule.

Examples of molecular shapes Table of molecular geometries for 2, 3, and 4 charge clouds

Linear: 2 charge clouds, 180° bond angle
Trigonal planar: 3 charge clouds, 120° bond angle
Bent: 2 bonds + 1 or 2 lone pairs
Tetrahedral: 4 charge clouds, 109.5° bond angle
Pyramidal: 3 bonds + 1 lone pair, ~107° bond angle

Linear geometry Trigonal planar geometry Bent geometry Tetrahedral geometry Effect of lone pairs on bond angles Molecular geometry in larger molecules Hydronium ion geometry (pyramidal) Acetaldehyde geometry: tetrahedral and trigonal planar

4.9 Polar Covalent Bonds and Electronegativity

When atoms in a covalent bond have different electronegativities, the electrons are shared unequally, resulting in a polar covalent bond. Electronegativity is a measure of an atom's ability to attract electrons in a bond.

Electronegativity increases across a period and decreases down a group.
The difference in electronegativity determines bond type:

Electronegativity Difference	Bond Type
0 – 0.4	Nonpolar covalent
0.5 – 1.9	Polar covalent
2.0 and above	Ionic

Polarity in HCl molecule Electronegativity values in the periodic table Comparing bond polarities Bond dipole illustration

4.10 Polar Molecules

A molecule is polar if it has a net dipole moment, meaning electrons are drawn more to one part of the molecule. Molecular polarity depends on both the polarity of individual bonds and the overall shape of the molecule.

Polarity in water and chloromethane Nonpolar molecules with polar bonds Polarity in hydrogen cyanide (HCN) Polarity in vinyl chloride

4.11 Naming Binary Molecular Compounds

Binary molecular compounds are named using prefixes to indicate the number of each atom present. The less electronegative element is named first, followed by the more electronegative element with an -ide suffix.

Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-
The prefix mono- is omitted for the first element unless needed for clarity.

Naming binary molecular compounds Examples of binary molecular compound names

Concept Map: Covalent and Ionic Bonds

Concept map of intramolecular forces