Covalent compounds are substances formed when atoms share valence electrons, known as a covalent bond, typically occurring between two nonmetal atoms. This electron sharing is the defining feature of these compounds. The nature of this bond dictates a distinct set of physical and chemical characteristics.
The Fundamental Structure: Discrete Molecules
The properties of most covalent compounds stem directly from their molecular structure. Unlike ionic compounds, which form extended crystalline lattices, covalent compounds exist as distinct, small molecules (e.g., \(\text{H}_2\text{O}\) or \(\text{CH}_4\)). The atoms within each molecule are held together by strong intramolecular covalent bonds that require a large amount of energy to break.
These individual molecules are attracted to each other by much weaker intermolecular forces (IMFs), collectively known as van der Waals forces. These forces include London dispersion forces, which are transient attractions present in all molecules, and dipole-dipole interactions, which occur between polar molecules with a permanent separation of charge.
A particularly strong type of dipole-dipole interaction is the hydrogen bond, which forms when a hydrogen atom bonded to a highly electronegative atom (nitrogen, oxygen, or fluorine) is attracted to a lone pair of electrons on a neighboring molecule. These intermolecular forces are significantly weaker than the covalent bonds holding the atoms together. This difference in strength determines the bulk physical properties, as physical changes only require overcoming the weak intermolecular forces.
Low Energy Requirements and Physical State
The relative weakness of the intermolecular forces directly impacts the energy needed to change the compound’s physical state. Covalent compounds exhibit low melting points (MP) and low boiling points (BP). Only a small amount of thermal energy is required to break the weak attractions between molecules, allowing them to transition between states.
Consequently, many covalent compounds are found as gases or liquids at standard room temperature. For instance, oxygen (\(\text{O}_2\)) and carbon dioxide (\(\text{CO}_2\)) are gases, while water and alcohol are liquids. Covalent compounds that exist as solids at room temperature, such as paraffin wax, tend to be soft and easily melted. The melting points for these simple molecular solids are usually below \(400^\circ\text{C}\).
Electrical Behavior and Interaction with Solvents
Most covalent compounds cannot conduct electricity. Covalent molecules are electrically neutral and lack mobile charged particles, such as free ions or delocalized electrons, in any state. Since electrical current relies on the movement of these charged carriers, covalent compounds are classified as non-conductors.
The interaction of a covalent compound with a solvent is governed by the principle of “like dissolves like,” which relates to molecular polarity. Polarity depends on the distribution of electrons within the molecule; a polar molecule has an uneven electron distribution, creating a partial positive and partial negative end. Polar covalent compounds (e.g., sugar or ethanol) readily dissolve in polar solvents like water because the attractive forces between the different molecules are strong enough to separate the solute molecules.
Conversely, nonpolar covalent compounds, which share electrons evenly, dissolve best in nonpolar solvents. Nonpolar substances (e.g., oil, fats, and hexane) mix with each other but do not dissolve in polar solvents like water. This explains why oil and water do not mix, as the cohesive forces within the highly polar water molecules are much stronger than the weak attractive forces they would form with nonpolar oil molecules.
The Exception: Covalent Network Solids
While most covalent substances form discrete molecules, a few exist as a structural exception known as covalent network solids. In these materials, such as diamond, quartz (\(\text{SiO}_2\)), and silicon carbide, the atoms are not organized into small, separate molecules. Instead, every atom is connected to its neighbors by a continuous series of strong covalent bonds extending throughout the entire crystal structure.
This continuous network effectively makes the entire solid one giant molecule. Because melting the solid requires breaking these numerous, strong covalent bonds, network solids exhibit extremely high melting points, often exceeding \(1000^\circ\text{C}\). This structural arrangement also gives them exceptional hardness and makes them generally insoluble in almost all solvents.