The melting point of a substance is the temperature at which it transitions from a solid state to a liquid state. This phase change requires an input of energy to disrupt the forces holding the solid structure together. Comparing the two major categories of elements, the general principle is straightforward: metals possess significantly higher melting points than nonmetals. This difference is directly attributable to the fundamental disparity in the atomic-level forces that hold these two classes of materials together.
The Role of Metallic Bonding in High Melting Points
The high temperatures required to melt most metals are a direct consequence of metallic bonding. In a solid metal, the valence electrons are delocalized, forming a mobile “sea of electrons” that extends throughout the crystal lattice. The remaining metal atoms become positive ions, or cations, which are held in fixed positions within this electron cloud.
The metallic bond is defined by the strong electrostatic attraction between these fixed positive ion cores and the surrounding electron sea. This force is substantial and uniformly distributed in three dimensions, creating a cohesive and robust structure. For example, melting copper (1085°C) or iron (1538°C) requires a large amount of thermal energy.
This energy is necessary to overcome the powerful electrostatic attractions and allow the positive ions to break free from their ordered lattice positions. Because the metallic bond is strong and pervasive throughout the solid, the temperature must be raised considerably before the structure collapses into the liquid state. Transition metals, such as tungsten, often exhibit high melting points because they contribute a greater number of electrons to the delocalized sea, intensifying the attractive forces.
The strength of the metallic bond is also influenced by the size and charge of the metal ions. Smaller ions with a higher positive charge density create a stronger pull on the electron sea, resulting in higher melting points. This strong bonding mechanism ensures that most common metals remain solid until heated to several hundred degrees Celsius or more.
Molecular Structure and Low Melting Points in Nonmetals
Most nonmetals exhibit low melting points due to the nature of their bonding forces. Many nonmetals, such as oxygen (\(O_2\)), iodine (\(I_2\)), and sulfur (\(S_8\)), exist as discrete, small molecules. Within these molecules, the atoms are held together by strong covalent bonds, known as intramolecular forces.
When these nonmetal molecules form a solid, they are held next to each other by much weaker forces called intermolecular forces. These forces, which include London Dispersion forces, are temporary attractions arising from fleeting electron movements. These weak attractions are the only forces that must be overcome for the substance to melt.
Melting a nonmetal does not require breaking the strong covalent bonds that hold the individual molecules together. Only a small input of thermal energy is needed to disrupt the weak intermolecular forces that keep the molecules arranged in a solid state. This minimal energy requirement explains why many nonmetals are gases or liquids at room temperature, or are solids that melt well below 100°C, such as iodine at 114°C.
Notable Exceptions to the General Rule
While the distinction between high-melting-point metals and low-melting-point nonmetals is a reliable rule, significant exceptions exist in both categories. The most prominent nonmetal exceptions are the network covalent solids, such as diamond and silicon. In these materials, atoms are linked by strong covalent bonds into a continuous, three-dimensional lattice, rather than being composed of discrete molecules.
To melt diamond, which has a melting point exceeding 3500°C, every single covalent bond in the structure must be broken. This process demands an immense amount of energy, giving these nonmetals melting points comparable to the most refractory metals. The entire solid effectively acts as one giant molecule, fundamentally changing the melting process compared to simple molecular nonmetals.
On the other side of the periodic table, some metals exhibit unusually low melting points. Mercury is the most famous example, existing as a liquid at room temperature with a melting point of \(-39^\circ\text{C}\). Other metals like Cesium (\(28^\circ\text{C}\)) and Gallium (\(30^\circ\text{C}\)) are also liquid or near-liquid at or slightly above room temperature.
This low melting behavior is often seen in the alkali metals. It is due to a less robust metallic bond caused by the larger atomic size of these elements. Having only a single valence electron participating in the electron sea results in weaker overall electrostatic attraction, requiring less thermal energy to overcome the bond.