Also found in: Dictionary, Thesaurus, Medical, Wikipedia.
adhesion and cohesion
The tendency of atoms or molecules to coalesce into extended condensed states. This tendency is practically universal. In all but exceptional cases, condensation occurs if the temperature is sufficiently low; at higher temperatures, the thermal motions of the constituents increase, and eventually the solid assumes gaseous form. The cohesive energy is the work required to separate the condensed phase into its constituents or, equivalently, the amount by which the energy of the condensed state is lower than that of the isolated constituents. The science of cohesion is the study of the physical origins and manifestations of the forces causing cohesion, as well as those opposing it. It is thus closely related to the science of chemical bonding in molecules, which treats small collections of atoms rather than extended systems. See Intermolecular forces
The origin and magnitude of the attractive forces depend on the chemical nature of the constituent atoms or molecules. Strong attractive interactions are usually associated with constituents having valence electron shells which are partly filled or open; if the valence electron shells are completely filled or closed, the interactions are weaker.
For open-shell constituents, as the atoms approach, the electron energy levels on different atoms begin to interact, forming a complex of energy levels in the solid. Some of these are below the atomic energy levels and some above. Since the atomic shells are partly filled, the lower energy levels in the solid are filled, but at least some of the higher levels are empty. Thus the average energy of the occupied levels in the solid is lower than that in the isolated atoms, resulting in an attractive force. Bonding in open-shell systems can be approximately divided into three categories, although most cases involve a combination. See Band theory of solids, Fermi-Dirac statistics, Solid-state physics, Valence band
1. Covalent bonding. This type of bonding is most similar to the molecular bond. The electron energy levels in the solid are split into a lower and a higher portion, with the states in the lower one filled and the higher one empty. Covalent bonds are strongly directional, with electron charge accumulating around the bond centers. Materials bonded in this fashion typically form structures with low coordination numbers, prototypical materials elements in group 14 of the periodic table, the insulator carbon, and the semiconductors silicon and germanium. See Semiconductor
2. Metallic bonding. In this case, there is no split between the lower and higher states of the electrons in the solid; rather, they occupy levels from the bottom up to a cutoff point known as the Fermi level. For example, in transition metals, the electron states in the solid derived from the atomic d orbitals form a complex which is gradually filled with increasing atomic number. The bulk of the cohesive energy is due to this complex. The metallic bond is less directional than the covalent bond, with a more uniform distribution of electronic charge. Metals usually form closely packed structures. See Fermi surface, Free-electron theory of metals
3. Ionic bonding. This occurs in compounds having at least two distinct types of atoms. One or more of the species of atoms (the cations) have only a small number of electrons in their valence shells, whereas at least one species (the anions) has a nearly filled valence shell. As the atoms approach each other, electrons drop from the cation valence states into holes in the anion valence shell, forming a closed-shell configuration in the solid. The different types of atoms in the solid have net charges; a strong attractive force results from the interaction between unlike charges. For example, in sodium chloride (NaCl), the sodium atoms acquire positive charges, and the chlorine atoms acquire negative charges. The closest interatomic separations in the solid are between sodium and chlorine, so that the attractive electrostatic interactions outweigh the repulsive ones. See Ionic crystals
In closed-shell constituents, the above effects are greatly reduced because the atomic or molecular shells are basically inert. The constituents retain their separate identities in the solid environment. If the constituents are atomic, as in rare-gas solids, the cohesion is due to the van der Waals forces. The positions of the electrons in an atom fluctuate over time, and at any given time their distribution is far from spherical. This gives rise to fluctuating long-ranged electric fields, which average zero over time, but can still have appreciable effects on neighboring atoms. The electrons on these atoms move in the direction of the force exerted by the electric field. The net result is that the interactions between unlike charges (electrons and nuclei) are increased in the solid, whereas the interactions between like charges are reduced. Thus the solid has a lower energy than the isolated atoms.
In solids made up of molecules, there are additional electrostatic interactions due to the nonspherical components of the molecular charge density. These interactions are strongest if the molecules are polar. This means that the center of the positive charge on the molecule is at a different point in space from that of the negative charge. Polar molecules, such as water (H2O), form structures in which the positive charge on a molecule is close to the negative charges of its neighbors. For nonpolar molecules, the electrostatic interactions are usually weaker than the van der Waals forces. The nonspherical interactions in such cases are often so weak that the molecules can rotate freely at elevated temperatures, while the solid is still held together by the van der Waals forces.
The repulsive forces in the condensed phase are a dramatic illustration of the combined action of two quantum-mechanical principles, the exclusion principle and the uncertainty principle.
The exclusion principle states that the quantum-mechanical wave function for the electrons in the solid must be antisymmetric under the interchange of the coordinates of any two electrons. Consequently, two electrons of the same spin are forbidden from being very close to each other. See Exclusion principle
The uncertainty principle states that if the motion of an electron is confined, its kinetic energy must rise, resulting in a repulsive force opposing the confinement. The kinetic energy due to the confinement is roughly inversely proportional to the square of the radius of the region of confinement. According to the exclusion principle, the motion of an electron in a solid is partially confined because it is forbidden from closely approaching other electrons of the same spin. Thus the uncertainty principle in turn implies a repulsive force. See Uncertainty principle
cohesionsee SOCIAL COHESION.
the binding of molecules, atoms, or ions of physical solids under the influence of forces of attraction, such as the forces of intermolecular interactions, hydrogen bonds, and/or chemical bonds. The forces determine the aggregate of physical and physicochemical properties of matter: state of aggregation, volatility, solubility, mechanical properties, and so on. The intensity of intermolecular and interatomic interactions (and, consequently, of the forces of cohesion) diminishes sharply with distance.
Cohesion is stronger in solids and liquids—that is, in condensed states, where the distances between molecules, atoms, or ions are small (of the order of a few angstroms). In gases, the average distance between molecules is large compared with their sizes; therefore, cohesion in them is insignificant. The cohesion energy density is a measure of the intensity of intermolecular interactions. It is equivalent to the work required to separate mutually attracting molecules or atoms to an infinitely large distance from each other, which corresponds in practice to evaporation or sublimation of the substance.
L. A. SHITS