Molecules with covalent bonds are generally poor conductors of electricity because they lack free-moving charged particles, which are essential for electrical conductivity. In covalent compounds, electrons are shared between atoms and are localized in specific bonds, meaning they cannot move freely throughout the substance. Additionally, covalent compounds do not have an overall electric charge. For instance, in a molecule like methane (CH₄), the electrons are tightly held in covalent bonds between carbon and hydrogen atoms and cannot move to conduct electricity. Even in polar covalent molecules like HCl, where there is an uneven distribution of electrons, the lack of free ions or electrons in the solid or liquid state prevents these substances from conducting electricity. Electrical conductivity in substances requires mobile ions or free electrons, which is characteristic of metallic and ionic compounds, not covalent compounds.

Covalent bonds contribute to the low melting and boiling points of molecular compounds primarily due to the nature of the interactions between molecules. In covalently bonded substances, the strong bonds exist within the molecules themselves (intramolecular forces), but the forces between different molecules (intermolecular forces) are relatively weak. For example, in a substance like iodine (I₂), the I-I covalent bond within each molecule is strong, but the forces between different I₂ molecules are weak van der Waals forces. When melting or boiling a covalent compound, it's these weak intermolecular forces that are overcome, not the strong covalent bonds within the molecules. This is why substances with simple molecular structures, where covalent bonds dominate, generally have lower melting and boiling points compared to ionic or metallic compounds, where the forces between the basic units are much stronger.

Yes, covalent bonds can form between atoms of the same element, and this is commonly observed in diatomic molecules. Diatomic molecules consist of two atoms of the same element bonded together covalently. A classic example is oxygen (O₂), where two oxygen atoms share two pairs of electrons, forming a double covalent bond. This sharing allows each oxygen atom to achieve a stable electronic configuration with eight valence electrons. Other examples include nitrogen (N₂), which forms a triple bond between two nitrogen atoms, and hydrogen (H₂), chlorine (Cl₂), fluorine (F₂), bromine (Br₂), and iodine (I₂), all of which form single covalent bonds. These diatomic molecules are essential components of the atmosphere and are fundamental in various biological and chemical processes.

Lone pairs, or non-bonding pairs, are pairs of valence electrons that are not shared between atoms in a molecule. They play a significant role in determining the shape and polarity of molecules. For example, in the water molecule (H₂O), the oxygen atom has two lone pairs. According to VSEPR (Valence Shell Electron Pair Repulsion) theory, electron pairs around a central atom, including bonding and lone pairs, repel each other and orient themselves as far apart as possible. In H₂O, the two lone pairs on oxygen push the hydrogen atoms closer together, resulting in a bent molecular shape rather than a linear one. This shape contributes to water's polarity. Similarly, in ammonia (NH₃), the nitrogen atom has one lone pair, which repels the bonding pairs, giving ammonia a trigonal pyramidal shape. The presence of lone pairs can significantly reduce the bond angles in molecules, affecting their physical and chemical properties.

Electronegativity refers to an atom's ability to attract and hold onto electrons. In covalent bonding, the difference in electronegativity between the two atoms influences the bond's character. When two atoms have similar electronegativities, they equally share electrons, resulting in a nonpolar covalent bond. A classic example is the Cl₂ molecule, where both chlorine atoms have the same electronegativity, leading to equal sharing of the electron pair. However, when there's a significant difference in electronegativity, as seen in the HCl molecule, the more electronegative atom (chlorine, in this case) attracts the shared electron pair more strongly. This uneven sharing creates a polar covalent bond, where the electron density is higher around the more electronegative atom, giving it a partial negative charge, while the less electronegative atom gains a partial positive charge. The degree of polarity in a covalent bond is crucial as it affects the molecule's physical properties, such as boiling and melting points, solubility, and reactivity.