Electronegativity itself does not directly predict the colour of a compound, but it can influence factors that affect colour. The colour of a compound is primarily determined by the wavelengths of light it absorbs and transmits. This absorption is related to the electronic structure of the compound, specifically the energy gap between different electron orbitals. In compounds where there is a significant electronegativity difference between bonded atoms, such as in transition metal complexes, the polarisation of bonds can affect the energy levels of d-orbitals. The splitting of these d-orbitals in a polarised environment can lead to specific wavelengths of light being absorbed, which in turn determines the colour of the compound. However, predicting the exact colour requires a detailed understanding of electronic transitions and molecular orbitals, which are influenced by various factors including electronegativity, molecular geometry, and the nature of the chemical bonds.

Electronegativity differences within molecules significantly affect their melting and boiling points. Substances with polar molecules, where there is a notable difference in electronegativity between the bonded atoms, generally have higher melting and boiling points compared to nonpolar molecules. This is due to the stronger intermolecular forces present in polar molecules, such as dipole-dipole interactions and hydrogen bonding. For example, water, with a significant electronegativity difference between oxygen and hydrogen, exhibits strong hydrogen bonding, leading to relatively high boiling and melting points. In contrast, nonpolar molecules, where the electronegativity difference is minimal, primarily exhibit weaker London dispersion forces. These forces increase with molecular size but are generally weaker than dipole-dipole or hydrogen bonding, resulting in lower melting and boiling points for nonpolar substances. Additionally, in ionic compounds, where the electronegativity difference is so significant that electrons are transferred, the resulting ionic bonds lead to very high melting and boiling points due to the strong electrostatic attraction between the ions.

Electronegativity significantly influences the acidity of a molecule, particularly in organic compounds like carboxylic acids. The acidity of a molecule is determined by the ease with which it donates a proton (H⁺ ion). In molecules with acidic hydrogen atoms, if these hydrogen atoms are attached to atoms with high electronegativity, the bond becomes polar. For instance, in carboxylic acids, the oxygen atoms attached to the hydrogen of the -OH group are highly electronegative, drawing electron density towards themselves and weakening the O-H bond. This weakening makes the hydrogen atom more prone to ionisation, enhancing the molecule's acidity. Additionally, electronegative atoms within the molecule, especially those near the acidic hydrogen, can stabilise the negative charge of the conjugate base formed after deprotonation, further increasing acidity. This stabilisation occurs through inductive effects, where the electronegative atoms pull electron density through sigma bonds, reducing the electron density on the conjugate base and making it more stable. Thus, the presence and position of electronegative atoms are crucial in determining a molecule's ability to donate protons and hence its acidity.

Electronegativity plays a significant role in determining the solubility of compounds in water. Water is a polar solvent, with a high dielectric constant due to its molecular polarity, influenced by the difference in electronegativity between hydrogen and oxygen atoms. When a compound is introduced into water, the polar water molecules interact with the compound's ions or molecules. For ionic compounds, the polar water molecules effectively weaken the electrostatic forces between the ions due to their partial charges, aiding in the dissolution process. For covalent compounds, those with polar bonds (where there's a significant difference in electronegativity between the bonded atoms) tend to dissolve better in water as they can form hydrogen bonds or dipole-dipole interactions with water molecules. Nonpolar covalent compounds, which lack significant differences in electronegativity, are generally insoluble in water as they cannot interact effectively with the polar water molecules. This principle underlies the common adage in chemistry: "like dissolves like," indicating that the solubility of a substance is related to its polarity, which in turn is influenced by the electronegativities of its constituent atoms.

Electronegativity significantly influences the reactivity of molecules. In a chemical reaction, the formation and breaking of bonds involve the redistribution of electrons between atoms. In molecules where there is a significant difference in electronegativity between the bonded atoms, the resulting polar bonds create regions of partial positive and negative charges. These polarised regions can be more susceptible to attack by nucleophiles (electron-rich species) or electrophiles (electron-deficient species). For example, in polar molecules like aldehydes and ketones, the carbon atom of the carbonyl group is partially positive due to the higher electronegativity of the oxygen atom. This makes it a site susceptible to nucleophilic attack, a fundamental concept in organic chemistry. Conversely, molecules with nonpolar bonds are less reactive in polar reactions but can undergo reactions driven by factors like steric strain or the presence of radical species. Therefore, understanding the distribution of electronegativity within a molecule can provide valuable insights into its potential reaction pathways and reactivity.