The fundamental nature of the chemical bond within a substance is the single most critical factor determining its macroscopic physical properties. The strength of the forces holding the constituent particles together – whether these are ions in a lattice or individual molecules – directly influences how much energy is required to overcome these forces, thereby dictating properties like melting point, boiling point, hardness, and electrical conductivity.

• High Melting and Boiling Points: This is perhaps the most defining characteristic. The strong, omnidirectional electrostatic forces of attraction holding the cations and anions together in a rigid, extended crystal lattice require a tremendous amount of thermal energy to overcome. Consequently, ionic compounds possess very high melting and boiling points, meaning they are typically solids at room temperature. For instance, magnesium oxide melts at an astonishing 2852^\circ\text{C}.
• Hard and Brittle Solids: The strong, fixed electrostatic attractions make ionic compounds very hard. However, their ordered structure also makes them brittle. If a mechanical force is applied that causes one layer of ions to shift relative to another, similarly charged ions will align, leading to strong electrostatic repulsion that can cause the crystal to cleave or shatter.
• Electrical Conductivity: Ionic compounds are excellent electrical conductors when in the molten (liquid) state or when dissolved in water (aqueous solution). In these states, the individual ions become mobile and are free to move throughout the substance, acting as charge carriers to conduct electricity. Conversely, ionic compounds are poor or non-conductors of electricity in their solid state because their ions are held rigidly within the crystal lattice and are not free to move.
• Solubility in Water: Many ionic compounds exhibit good solubility in water and other polar solvents. Water molecules are polar (they have distinct positive and negative ends) and can interact strongly with the charged ions. These water molecules surround and effectively pull the ions away from the crystal lattice, dispersing them into the solution.
• Low Volatility: Due to their strong interionic forces and high boiling points, ionic compounds have very low volatility, meaning they do not readily turn into a gas.

• Low Melting and Boiling Points: This is a key distinguishing feature. Within each covalent molecule, the atoms are held together by strong covalent bonds. However, the forces between separate, individual molecules (known as intermolecular forces, IMFs) are significantly weaker than the intramolecular covalent bonds or the interionic forces in ionic compounds.
• Poor Electrical Conductivity: Covalent compounds generally do not conduct electricity in any state (solid, liquid, or gas). This is because they are composed of neutral molecules, which do not possess free-moving charged particles (ions) or delocalized electrons that can carry an electrical current.
• Variable Solubility: The solubility of covalent compounds in water is highly variable and depends on the polarity of the molecule. Small, polar covalent molecules (like sugar, ethanol, or ammonia) can dissolve in water because they can form attractive interactions with polar water molecules. However, non-polar covalent molecules (like oil, methane, or waxes) are generally insoluble in water ("like dissolves like" principle).
• High Volatility: Due to the weak intermolecular forces and low boiling points, many covalent compounds are volatile, meaning they readily evaporate and turn into a gas at room temperature.
• Soft Solids or Gaseous/Liquid States at Room Temperature: Reflecting their low melting and boiling points, many covalent compounds are soft solids, liquids, or gases under ambient conditions.