Chemical Bonds, Electronegativity, and Covalent Compounds: Study Notes for General Chemistry

Chemical Bonds and the Octet Rule

Introduction to Chemical Bonding

Chemical bonds are the attractive forces that hold atoms together in compounds. Atoms form bonds to achieve a stable electron configuration, often resembling the nearest noble gas (the octet rule). There are three primary types of chemical bonds: ionic, covalent, and metallic.

• Ionic bonds: Formed by the transfer of electrons from a metal to a nonmetal, resulting in the formation of cations and anions.

• Covalent bonds: Formed by the sharing of electrons between two nonmetals.

• Metallic bonds: Involve the pooling of electrons among metal atoms.

Electronegativity and Covalent Bonds

Electronegativity: Definition and Trends

Electronegativity (EN) is the tendency of an atom to attract electrons in a chemical bond. It is a key factor in determining bond polarity and the type of bond formed between atoms.

• Electronegativity increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Fluorine (F) is the most electronegative element; noble gases generally do not have EN values because they rarely form bonds.

Bond Polarity and Dipole Moments

The difference in electronegativity (ΔEN) between two atoms determines the bond type:

• Nonpolar covalent: ΔEN ≤ 0.4 (equal sharing of electrons)

• Polar covalent: 0.4 < ΔEN ≤ 2.0 (unequal sharing, partial charges form)

• Ionic: ΔEN > 2.0 (complete transfer of electrons)

A dipole moment (μ) measures the separation of charge in a molecule and is calculated as:

$\mu = Qr$ where Q = charge, r = separation distance

Lewis Structures and the Octet Rule

Drawing Lewis Structures

Lewis structures represent the arrangement of valence electrons in a molecule. The steps to construct a Lewis structure are:

1. Find the total number of valence electrons by summing the group numbers of all atoms.

2. Place the least electronegative atom in the center (except hydrogen).

3. Connect atoms with single bonds (pairs of electrons), then complete the octets of surrounding atoms.

4. Place any remaining electrons on the central atom. If the central atom lacks an octet, form double or triple bonds as needed (except for Be, B, F, and metals).

5. Calculate formal charges to determine the most reasonable structure.

Octet Rule and Its Exceptions

• Incomplete octet: Some elements (e.g., B, Be) can have fewer than 8 electrons.

• Odd-electron species: Molecules with an odd number of electrons (e.g., NO).

• Expanded octet: Elements in period 3 or higher can have more than 8 electrons due to available d orbitals.

Single, Double, and Triple Bonds

• Single bond: One shared pair of electrons (2 electrons).

• Double bond: Two shared pairs (4 electrons).

• Triple bond: Three shared pairs (6 electrons).

Resonance Structures

Some molecules can be represented by two or more valid Lewis structures, called resonance structures. The actual structure is a resonance hybrid, with delocalized electrons.

• Resonance is indicated by a double-headed arrow between structures.

• Bond lengths and strengths in resonance hybrids are intermediate between single and double bonds.

Formal Charge

Formal charge helps identify the most stable Lewis structure. It is calculated as:

$\text{Formal charge} = \text{valence electrons} - \text{unshared electrons} - \frac{1}{2}(\text{shared electrons})$

• Structures with the smallest formal charges (closest to zero) are preferred.

• Negative formal charges should reside on the most electronegative atoms.

Naming Covalent Compounds

Rules for Naming Binary Covalent Compounds

Binary covalent compounds are formed between two nonmetals. The naming rules are:

• Name the less electronegative element first.

• Name the second element with the root and the suffix -ide.

• Use numerical prefixes to indicate the number of atoms (except mono- for the first element).

Examples:

• CO2: carbon dioxide

• Cl2O7: dichlorine heptoxide

• CO: carbon monoxide

Special Names for Common Covalent Compounds

Compound	Common Name	Comments
H2O	Water	No other name is accepted for this compound.
NH3	Ammonia	No other name is accepted for this compound.
N2O	Nitrous oxide	Dinitrogen oxide and dinitrogen monoxide are allowed, but not generally used.
NO	Nitric oxide	Nitrogen oxide and nitrogen monoxide are allowed, but not generally used.

Naming Hydrocarbons

Hydrocarbons are compounds containing only carbon and hydrogen. The simplest type, alkanes, are named with a root and the suffix -ane.

Recognizing Ionic and Molecular Compounds

Distinguishing Features

• Ionic compounds: Contain a metal and a nonmetal, or include polyatomic ions like NH4+.

• Molecular (covalent) compounds: Contain only nonmetals.

Examples:

• Ionic: CaO, Fe(NO3)3, NaC2H3O2, NH4Cl

• Molecular: CO2, H2SO4, CH3OH

Comparison Table: Ionic vs. Covalent Compounds

Property	Ionic Compounds	Covalent Compounds
Type of elements	Metal and nonmetal	Nonmetals only
Bonding	Ionic (attraction between anions and cations)	Covalent (sharing of electrons)
Representative unit	Formula unit	Molecule
Physical state at room temperature	Solid	Gas, liquid, or solid
Water solubility	Usually high	Variable
Melting and boiling points	Generally high	Generally low
Electrical conductivity	Good when molten or in solution	Poor
State when dissolved in water	Separates into ions	Remains molecules

Summary

This guide covers the essential concepts of chemical bonding, electronegativity, Lewis structures, naming covalent compounds, and distinguishing between ionic and molecular compounds. Mastery of these topics is fundamental for success in general chemistry.