Consider the different ways atomic-scale particles attract other atomic-scale particles and how strengths of those attractions affect macroscopic properties. Noble-gas atoms have only weak London dispersion forces (LDFs) between them, leading to very low boiling points.
Metals consist of positive ions surrounded by valence electrons that are not associated with any specific positive ion (that is, with any specific metal-atom core). Attractions between metal atoms involve positive ions and electrons and therefore are much larger than attractions between noble-gas atoms. Metallic bonding becomes stronger as the number of valence electrons in the metal atom increases. Metal atoms are attracted strongly enough that most metals are solids at room temperature.
When an atom with low attraction for electrons (a metal atom with low ionization energy) approaches an atom with greater attractions for electrons (a nonmetal with large negative electron affinity), electron density can transfer from the metal atom to the nonmetal atom to form ions. Coulomb’s law attractions between ions are large and depend on the charges of the ions and the distance between the ions. This results in formation of ionic crystal lattices that require significant increase in temperature for melting or boiling.
When two nonmetal atoms approach, the typical result is a covalent chemical bond, although there are cases such as He2 where there are enough electrons to fill antibonding as well as bonding molecular orbitals and give a bond order of zero. The characteristics of covalent bonds depend on properties of the bonded atoms, such as size and number of electrons. Covalent bonds connect atoms to form molecules. The strengths of bonds in molecules are typically as large as or larger than the strengths of metallic bonds and ionic bonds. Thus, when a molecular substance melts or boils, atoms remain bonded and the atomic-scale particles in the liquid or gas phase are molecules.
For example, when NaCl melts, Na+ ions and Cl‾ ions break from the crystal lattice and move freely around each other. Breaking the ionic bonds requires significant energy, and hence NaCl has a high melting point of 801 °C. In contrast, when methane (CH4) melts, individual methane (CH4) molecules stay intact, but they can move freely around other methane molecules. The attractions between methane molecules must be partially overcome, but those attractions are not as strong as covalent bonds or ionic bonds. Hence, the melting point of methane is much lower than for NaCl. Methane melts at −182 °C.
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