Covalent Compounds: Structure, Bonding, and Properties

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Covalent Compounds: Structure, Bonding, and Properties

Introduction to Covalent Bonding

Covalent bonds are a fundamental type of chemical bond where two atoms share electrons to achieve a stable electron configuration. This sharing allows each atom to attain the electron arrangement of the nearest noble gas, resulting in the formation of discrete molecules.

Covalent bond: A chemical bond formed by the sharing of one or more pairs of electrons between two atoms.
Molecule: A group of atoms held together by covalent bonds, representing the smallest unit of a covalent compound.

Formation of a covalent bond between two hydrogen atoms

Diatomic Elements

Certain elements naturally exist as diatomic molecules, meaning they are found in nature as pairs of atoms bonded together by covalent bonds. These elements are essential in both chemical reactions and biological processes.

The seven diatomic elements are: H2, N2, O2, F2, Cl2, Br2, I2.

Table of diatomic elements Periodic table highlighting diatomic elements

Nomenclature of Covalent Compounds

Naming Covalent Compounds

Covalent compounds are named using a systematic approach that indicates the number and type of atoms present. Prefixes are used to denote the number of each atom, and the second element's name ends with the suffix "-ide." The prefix "mono-" is usually omitted for the first element.

First nonmetal: Use the element name.
Second nonmetal: Use the root of the element name + "-ide".
Prefixes indicate the number of atoms (see table below).

Prefixes for naming covalent compounds

Examples:

CO: carbon monoxide
NO2: nitrogen dioxide
N2O4: dinitrogen tetroxide

Writing Formulas for Covalent Compounds

To write the formula from a compound's name, use the element symbols in the order given and apply the appropriate prefixes as subscripts.

Example: Sulfur hexafluoride → SF6
Example: Dinitrogen trioxide → N2O3

Lewis Structures and Resonance

Lewis Structures

Lewis structures are diagrams that show the arrangement of valence electrons among atoms in a molecule. They help predict molecular shape, bond order, and the presence of lone pairs.

Each bond represents two shared electrons.
Lone pairs are non-bonding pairs of electrons localized on a single atom.
Atoms (except hydrogen) strive for an octet (eight electrons) in their valence shell.

Lewis structure of HF showing lone pairs and shared electrons

Resonance Structures

Some molecules cannot be adequately represented by a single Lewis structure. Resonance structures are two or more valid Lewis structures that differ only in the position of electrons, not atoms. The true structure is a resonance hybrid, which stabilizes the molecule by delocalizing electrons.

Resonance structures have the same arrangement of atoms but different arrangements of electrons.
The actual molecule is a hybrid of all resonance forms.

Resonance structures of HCO3-

Molecular Shape and VSEPR Theory

VSEPR Theory

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the three-dimensional shape of molecules based on the repulsion between electron groups (bonding and lone pairs) around a central atom. The arrangement minimizes repulsion, resulting in specific molecular geometries.

Electron groups arrange themselves as far apart as possible.
The number of electron groups determines the molecular geometry.

Basic molecular shapes: linear, trigonal planar, tetrahedral

Common Molecular Geometries

Linear: 2 groups, 180° bond angle (e.g., CO2)
Trigonal planar: 3 groups, 120° bond angle (e.g., BF3)
Tetrahedral: 4 groups, 109.5° bond angle (e.g., CH4)
Trigonal pyramidal: 3 atoms + 1 lone pair (e.g., NH3), bond angle ~107°
Bent: 2 atoms + 1 or 2 lone pairs (e.g., H2O, SO2), bond angle ~105° or 120°

Linear geometry example HCN linear geometry CO2 linear geometry Trigonal planar geometry BF3 trigonal planar geometry H2CO trigonal planar geometry SO2 bent geometry SO2 bent geometry 120 degrees Tetrahedral geometry Tetrahedral carbon Tetrahedral bonds in the plane, in front, and behind Trigonal pyramidal geometry (NH3) Bent geometry (H2O)

Electronegativity and Bond Polarity

Electronegativity

Electronegativity is a measure of an atom's ability to attract 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.

Electronegativity trend in the periodic table

Bond Polarity

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

Nonpolar covalent: Electrons are shared equally (ΔEN < 0.5).
Polar covalent: Electrons are shared unequally (0.5 ≤ ΔEN ≤ 1.9).
Ionic: Electrons are transferred (ΔEN > 1.9).

Bond polarity and electronegativity difference Dipole in a polar covalent bond Bond type and electron sharing Table of electronegativity difference and bond type

Polarity of Molecules

Determining Molecular Polarity

The overall polarity of a molecule depends on both the polarity of its bonds and its molecular geometry. If the dipoles cancel due to symmetry, the molecule is nonpolar; if they reinforce, the molecule is polar.

Linear, trigonal planar, and tetrahedral molecules with identical surrounding atoms are usually nonpolar.
Bent and trigonal pyramidal molecules are usually polar.

NH3 is a polar molecule BF3 is a nonpolar molecule CO2 is a nonpolar molecule CCl4 is a nonpolar molecule H2O is a polar molecule CH3Cl is a polar molecule Summary of molecular polarity examples

Hybridization

Hybridization and Molecular Geometry

Hybridization explains the observed shapes of molecules by describing the mixing of atomic orbitals to form new, equivalent hybrid orbitals. The type of hybridization corresponds to the number of electron groups around the central atom.

sp: 2 groups, linear geometry
sp2: 3 groups, trigonal planar geometry
sp3: 4 groups, tetrahedral geometry

Hybridization and observed geometries Hybridization and electronic structure Carbon ground-state electron configuration Hybridization in methane Formation of four equivalent bonds in methane Hybrid orbitals Shape of hybrid orbitals Basic hybridization geometries Hybridization geometries: trigonal planar, bent, tetrahedral, trigonal pyramidal, nonlinear Hybridization geometries: seesaw, T-shaped, linear, square pyramidal, square planar Determining hybridization Examples of hybridization

Intermolecular Forces

Types of Intermolecular Forces (IMFs)

Intermolecular forces are attractions between molecules that influence physical properties such as boiling and melting points. They are generally weaker than intramolecular (covalent or ionic) bonds.

London Dispersion Forces (LDF): Present in all molecules; weakest IMF; arise from temporary dipoles.
Dipole-Dipole Forces: Present in polar molecules; due to permanent dipoles.
Hydrogen Bonding: Special dipole-dipole interaction involving H bonded to O, N, or F; stronger than regular dipole-dipole forces.
Ion-Dipole Forces: Attraction between an ion and a polar molecule; strongest IMF.

Periodic Trends

Effective Nuclear Charge (Zeff)

The effective nuclear charge is the net positive charge experienced by valence electrons, accounting for shielding by inner electrons. It increases across a period and decreases down a group.

Formula:

Atomic Size

Atomic size (radius) increases down a group and decreases across a period due to increasing nuclear charge pulling electrons closer.

Ionization Energy (IE)

Ionization energy is the energy required to remove an electron from a neutral atom. It decreases down a group and increases across a period.

Electron Affinity (EA)

Electron affinity is the energy released when an atom gains an electron. It generally becomes more negative (greater) across a period.

Metallic Character

Metallic character increases down a group and decreases across a period. Elements with high metallic character lose electrons easily.

*Additional info: This guide covers the essential concepts of covalent bonding, molecular structure, polarity, hybridization, intermolecular forces, and periodic trends, providing a comprehensive foundation for further study in general chemistry.*