Covalent Bonds

Definition and Basic Principles

Concept

Covalent bonds: chemical bonds formed by the sharing of electron pairs between atoms. Unlike ionic bonds, no electron transfer occurs. Result: stable molecule with shared valence electrons.

Historical Context

Introduced by G.N. Lewis (1916) via electron dot structures. Concept expanded quantum mechanically by Heitler and London (1927) using wavefunction overlap.

Atoms Involved

Usually nonmetals with similar electronegativities. Examples: H2, O2, N2, CO2, CH4.

Octet Rule

Atoms share electrons to achieve noble gas configuration (eight electrons in valence shell). Exceptions exist (e.g., H with duet, expanded octet in third period).

Types of Covalent Bonds

Single Bonds

One shared electron pair (two electrons). Example: H–H in hydrogen molecule.

Double Bonds

Two shared pairs (four electrons). Example: O=O in oxygen molecule.

Triple Bonds

Three shared pairs (six electrons). Example: N≡N in nitrogen molecule.

Coordinate (Dative) Bonds

One atom donates both electrons forming the bond. Example: NH4+ ion formation.

Polar and Nonpolar Covalent Bonds

Polar: unequal sharing due to electronegativity difference. Nonpolar: equal sharing.

Electron Sharing Mechanism

Orbital Overlap

Bonding orbitals formed by constructive interference of atomic orbitals. Overlap magnitude correlates with bond strength.

Sigma (σ) Bonds

Head-on overlap along internuclear axis. Strongest covalent bond type.

Pi (π) Bonds

Side-on overlap of p orbitals above and below the axis. Weaker than sigma bonds.

Electron Pair Sharing

Electron pairs localize between nuclei, reducing potential energy and increasing stability.

Bond Formation Process

Energy Changes

Bond formation releases energy (exothermic). Bond dissociation requires energy input (endothermic).

Potential Energy Diagram

Atoms approach: potential energy decreases. Minimum energy corresponds to equilibrium bond length.

Quantum Mechanical Model

Wavefunctions combine to form molecular orbitals. Bonding orbitals lower energy; antibonding raise energy.

Valence Bond Theory

Hybridization explains molecular shape and bond angles.

Properties of Covalent Bonds

Bond Strength

Measured by bond dissociation energy. Stronger bonds resist breaking.

Bond Length

Distance between nuclei at minimum potential energy. Inversely related to bond strength.

Electrical Conductivity

Covalent compounds typically poor conductors; electrons localized.

Solubility

Polar covalent compounds soluble in polar solvents; nonpolar soluble in nonpolar solvents.

Bond Polarity and Electronegativity

Electronegativity Concept

Atom’s ability to attract electrons in bond. Scale: Pauling scale most common.

Nonpolar Covalent Bonds

Electronegativity difference (ΔEN) < 0.4. Equal electron sharing.

Polar Covalent Bonds

ΔEN between 0.4 and 1.7. Partial charges develop: δ+ and δ−.

Ionic Character

ΔEN > 1.7 indicates ionic bond; continuum exists between covalent and ionic.

Dipole Moment

Quantitative measure of bond polarity. Vector quantity: magnitude and direction.

Bond Energy and Strength

Definition

Energy required to break one mole of bonds in gaseous state.

Factors Affecting Bond Energy

Bond order, bond length, atomic size, electronegativity.

Typical Bond Energies

Single bonds: 150-400 kJ/mol, double: 400-700 kJ/mol, triple: 700-1100 kJ/mol.

Bond Strength and Reactivity

Stronger bonds less reactive; weaker bonds more reactive.

Energy Diagram Example

Bond type Energy (kJ/mol)Single ~350Double ~610Triple ~840

Molecular Geometry and VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR)

Electron pairs repel to minimize energy, determining molecular shape.

Common Geometries

Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

Effect of Lone Pairs

Lone pairs occupy space, distort bond angles.

Hybridization

sp, sp2, sp3 hybrid orbitals correlate with geometry.

Example: Methane

CH4: tetrahedral, bond angle ~109.5°, sp3 hybridization.

Bond Length and Factors Affecting It

Definition

Distance between nuclei of bonded atoms at minimum potential energy.

Influencing Factors

Atomic radii, bond order, bond polarity, hybridization.

Bond Order Effect

Higher bond order → shorter bond length.

Atomic Size Effect

Larger atoms → longer bonds.

Table: Selected Bond Lengths

Multiple Bonds: Double and Triple Bonds

Structure

Double bonds: one sigma + one pi bond. Triple bonds: one sigma + two pi bonds.

Bond Strength

Multiple bonds stronger and shorter than single bonds.

Reactivity

Multiple bonds more reactive due to pi bond exposure.

Examples

Alkenes (C=C), alkynes (C≡C), O2 (double bond), N2 (triple bond).

Hybridization

Double bonds: sp2 hybridization; triple bonds: sp hybridization.

Applications of Covalent Bonding

Organic Chemistry

Basis for carbon compounds, pharmaceuticals, polymers.

Materials Science

Design of covalent network solids: diamond, graphene, silicones.

Biochemistry

Protein folding, DNA base pairing, enzyme-substrate complexes.

Industrial Chemistry

Synthesis of fuels, plastics, dyes, and agrochemicals.

Nanotechnology

Molecular machines, self-assembly via covalent linkages.

Experimental Techniques for Studying Covalent Bonds

X-ray Crystallography

Determines bond lengths, angles, and molecular structure.

Infrared (IR) Spectroscopy

Identifies bond types via vibrational frequencies.

Nuclear Magnetic Resonance (NMR)

Probes electronic environment around atoms.

Electron Microscopy

Visualizes molecular assemblies and bond arrangements.

Computational Chemistry

Quantum calculations predict bond energies, structures.

References

• Pauling, L. The Nature of the Chemical Bond. Cornell University Press, 1960, pp. 1-600.
• Atkins, P., de Paula, J. Physical Chemistry. 10th ed., Oxford University Press, 2014, pp. 200-230.
• Levine, I.N. Quantum Chemistry. 7th ed., Pearson, 2014, pp. 150-190.
• Housecroft, C.E., Sharpe, A.G. Inorganic Chemistry. 4th ed., Pearson, 2012, pp. 130-160.
• Silverstein, R.M., Webster, F.X., Kiemle, D.J. Spectrometric Identification of Organic Compounds. 7th ed., Wiley, 2005, pp. 50-90.