4.2 Covalent Bonding | Chemistry: Atoms First 2e

Learning Objectives

Describe the formation of covalent bonds
Define electronegativity and assess the polarity of covalent bonds

Core Concepts & Principles

Unlike ionic bonds formed by complete electron transfer between metals and nonmetals, covalent bonds result when two atoms share a pair of valence electrons. This occurs primarily between nonmetal atoms that have similar ionization energies and electron affinities.

Physical Properties of Covalent Compounds

Because covalent molecules are electrically neutral, the intermolecular attractions between them are much weaker than the electrostatic forces between full ions in ionic crystals. Consequently, covalent compounds generally feature:

Much lower melting and boiling points (often existing as liquids or gases at room temperature).
Softer solid states.
Poor electrical conductivity in all states and little to no solubility in water.

Formation and Energy of Covalent Bonds

When two hydrogen atoms approach each other, their $1s$ valence orbitals overlap, and their single electrons interact with both nuclei.

Potential Energy: As the atoms get closer, potential energy decreases due to attractive forces until reaching a minimum at the optimal bond length. If pushed closer still, nuclear-nuclear repulsion causes the potential energy to spike.
Thermodynamics: Breaking chemical bonds is an endothermic process (requires energy input), while forming chemical bonds is an exothermic process (releases energy). For example: $\text{H}_2(g) \longrightarrow 2\text{H}(g) \quad \Delta H = +436\text{ kJ/mol}$ $2\text{H}(g) \longrightarrow \text{H}_2(g) \quad \Delta H = -436\text{ kJ/mol}$

Bond Energy & Stability Rule

Bond formation releases energy (exothermic, $-\Delta H$ ), creating a stable system at the lowest potential energy point (bond length). Bond breaking absorbs energy (endothermic, $+\Delta H$ ).

Pure vs. Polar Covalent Bonds

The sharing of electrons is not always equal:

Pure covalent bond: Electrons are shared equally between identical atoms (e.g., $\text{H}_2$ , $\text{Cl}_2$ ) with a zero electronegativity difference.
Polar covalent bond: Electrons are shared unequally between different atoms because one atom pulls harder on the electron density than the other. This creates partial positive ( $\delta^+$ ) and partial negative ( $\delta^-$ ) charges (e.g., $\text{H}^{\delta+}-\text{Cl}^{\delta-}$ ).

Electronegativity

Electronegativity measures an atom's tendency to attract shared electrons toward itself in a chemical bond.

Periodic Trends: Electronegativity increases up and to the right on the periodic table. Fluorine is the most electronegative element ( $\text{EN} = 4.0$ ). Group 1 metals have the lowest values.
Electronegativity vs. Electron Affinity: Electron affinity is a measurable physical quantity ( $\text{kJ/mol}$ ) of an isolated gas-phase atom gaining an electron. Electronegativity is a calculated, dimensionless relative scale (ranging from 0 to 4) pioneered by Pauling to describe bonded atoms.

Bond Polarity & Electronegativity Difference ($\Delta EN$)

The absolute electronegativity difference ( $\Delta EN$ ) between two bonded atoms determines bond type:

$\Delta EN \approx 0$ : Pure (nonpolar) covalent
$\Delta EN \approx 0.4 - 1.4$ : Polar covalent
$\Delta EN \ge 2.0$ : Ionic character predominates

Problem-Solving Routines & Methods

How to Determine Bond Polarity and Partial Charges

1
Identify the two bonded atoms in the molecule or amino acid side chain.
2
Look up the electronegativity (EN) values for both atoms using the periodic table.
3
Calculate the absolute electronegativity difference: $\Delta EN = |EN_1 - EN_2|$ .
4
Assign partial charges: the atom with the higher EN receives $\delta^-$ , and the lower EN receives $\delta^+$ .

Pro-Tip: Always verify atom types: bonds between two nonmetals are generally covalent, while metal-nonmetal pairs are typically ionic regardless of intermediate EN differences.

Electronegativity Difference Formula

\Delta EN = |EN_1 - EN_2|

Quantifies the polarity of a bond based on the electronegativities of the two participating atoms.

Variables & Constants

\Delta EN

=absolute electronegativity difference;

EN_1

=electronegativity of atom 1;

EN_2

=electronegativity of atom 2

Practice & Concept Checks

Concept Check

Why do covalent compounds generally have much lower melting and boiling points than ionic compounds?

Concept Check

If a covalent bond has an electronegativity difference (

\Delta EN

) of 1.4 (such as an

\text{O}-\text{H}

bond), which atom carries the partial negative (

\delta^-

) charge and why?

Key Terms & Vocabulary

covalent bondsBonding

Chemical bonds formed by the mutual attraction of atoms for a shared pair of valence electrons.

Example: The H–H bond in H₂

bond lengthMolecular Structure

The internuclear distance between two bonded atoms where the lowest potential energy of the system is achieved.

Example: Optimal internuclear distance in H₂ molecule

pure covalent bondBonding

A covalent bond in which bonding electrons are shared equally between identical atoms with zero electronegativity difference.

Example: Cl–Cl bond in Cl₂

polar covalent bondBonding

A covalent bond where electrons are shared unequally, producing an uneven electron density distribution and partial charges.

Example: H–Cl with Hδ⁺ and Clδ⁻

electronegativityPeriodic Properties

A dimensionless measure of an atom's tendency to attract shared electrons or electron density toward itself in a bond.

Example: Fluorine is the most electronegative element (EN = 4.0)

PaulingHistorical Figures

Linus Pauling, the pioneering chemist who established the first relative electronegativity scale and foundational bonding theories.

Example: Pauling electronegativity values