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The properties of covalent compounds can be largely attributed to the presence of van der Waals forces, which are weak attractions between molecules or individual atoms. Despite their relatively weak nature, these forces play a critical role in the physical properties of covalent compounds.

Van der Waals forces include London dispersion forces, dipole-dipole interactions, and hydrogen bonds. London dispersion forces occur between all atoms and molecules, resulting from temporary fluctuations in electron distribution which induce a temporary dipole. Dipole-dipole interactions occur in polar molecules where there is a permanent separation of charge. Hydrogen bonds, which are a type of dipole-dipole attraction, occur when hydrogen is bound to a highly electronegative atom such as oxygen, nitrogen, or fluorine.

These interactions account for the low melting and boiling points of covalent compounds since a minimal amount of energy is required to break them apart. Understanding this principle helps explain why substances like ice, which is held together by hydrogen bonds, melt at relatively low temperatures compared to other solid materials.

The electrical conductivity of a substance is a measure of its ability to allow the flow of an electric current. In the context of covalent compounds, poor electrical conductivity is a notable property. This property stems from the nature of covalent bonding, where electrons are shared between atoms, forming a stable configuration that localizes the electrons.

Unlike metals, where electrons can move freely as a 'sea of electrons', or ionic compounds which contain mobile ions, covalent compounds typically have no charge carriers to support an electric current. The solid forms of these compounds generally lack the delocalized electrons required for conductivity, leading to their characterization as insulators or semiconductors, depending on the purity and molecular structure of the substance. Addressing common misconceptions, not all covalent substances are non-conductive; for example, a diamond (a form of carbon with a covalent network structure) can exhibit semiconductor properties under specific conditions.

The molecular structure of covalent compounds is determined by the specific ways in which atoms bond together through covalent bonds. These bonds arise from the sharing of electrons between nonmetal atoms, which have similar tendencies to attract electrons. The molecular structures can range from simple, such as the diatomic oxygen molecule (O2), to complex, like the large biomolecules found in living organisms.

Most covalent compounds exist as discrete molecules with definite shapes and sizes. The specific geometry of a molecule—such as linear, bent, tetrahedral, or trigonal planar—is defined by the Valence Shell Electron Pair Repulsion (VSEPR) theory, which posits that electron pairs around a central atom will arrange themselves to minimize repulsion.

The brittleness or softness of covalent solids is influenced by the manner in which these molecules pack together. In solid forms, they are often organized in a lattice structure held by van der Waals forces. Because these forces are weaker than ionic or metallic bonds, covalent compounds can be easily deformed or shattered when subjected to stress.