Hameroff Was Talking About Living Water Forty Years Ago...,

hameroff  |  Biomolecules have evolved and flourished in aqueous environments, and basic interactions among biomolecules and their pervasive hosts, water molecules, are extremely important. The properties of intracellular water are controversial. Many authors believe that more than 90 percent of intracellular water is in the “bulk” phase-water as it exists in the oceans (Cooke and Kuntz, 1974; Schwan and Foster, 1977; Fung and McGaughy, 1979). This traditional view is challenged

by others who feel that none of the water in living cells is bulk (Troshin, 1966, Cope 1976, Negendank and Karreman, 1979). A middle position is assumed by those who feel that about half of “living” water is bulk and the other half “ordered” (Hinke, 1970; Clegg, 1976; Clegg, 1979; Horowitz and Paine, 1979). This group emphasizes the importance of “ordered” water to cellular structure and function.

Many techniques have been used to study this issue, but the results still require a great deal of interpretation. Nuclear magnetic resonance (NMR), neutron diffraction, heat capacity measurement, and diffusion studies are all inconclusive. Water appears to exist in both ordered and aqueous forms within cells. The critical issue is the relation between intracellular surfaces and water. Surfaces of all kinds are known to perturb adjacent water, but within cells it is unknown precisely how far from the surfaces ordering may extend. We know the surface area of the microtrabecular lattice and other cytoskeleton components is extensive (billions of square nanometers per cell) and that about one fifth of cell interiors consist of these components. Biologist James Clegg (1981) has extensively reviewed these issues. He concludes that intracellular water exists in three phases. 1) “Bound water” is involved in primary hydration, being within one or two layers from a biomolecular surface. 2) “Vicinal water” is ordered, but not directly bound to structures except other water molecules. This altered water is thought to extend 8 to 9 layers of water molecules from surfaces, a distance of about 3 nanometers. Garlid (1976, 1979) has shown that vicinal water has distinct solvent properties which differ from bulk water. Thus “borders” exist between water phases which partition solute molecules. 3) “Bulk water” extends beyond 3 nanometers from cytoskeletal surfaces (Figure 6.4).

Drost-Hansen (1973) described cooperative processes and phase transitions among vicinal water molecules. Clegg points out the potential implications of vicinal water on the function of enzymes which had previously been considered “soluble.” Rather than floating freely in an aqueous soup, a host of intracellular enzymes appear instead to be bound to the MTL surface within the vicinal water phase. Significant advantages appear evident to such an arrangement: a sequence of enzymes which perform a sequence of reactions on a substrate would be much more efficient if bound on a surface in the appropriate order. Requirements for diffusion of the substrate, the most time consuming step in enzymatic processes, would be minimal. Clegg presents extensive examples of associations of cytoplasmic enzymes which appear to be attached to and regulated by, the MTL. These vicinal water multi-enzyme complexes may indeed be part of a cytoskeletal information processing system. Clegg conjectures that dynamic conformational activities within the cytoskeleton/MTL can selectively excite enzymes to their active states.

The polymerization of cytoskeletal polymers and other biomolecules appears to flow upstream against the tide of order proceeding to disorder which is decreed by the second law of thermodynamics. This apparent second law felony is explained by the activities of the water molecules involved (Gutfreund, 1972). Even in bulk aqueous solution, water molecules are somewhat ordered, in that each water molecule can form up to 4 hydrogen bonds with other water molecules. Motion of the water molecules (unless frozen) and reversible breaking and reforming of these hydrogen bonds maintain the far miliar liquid nature of bulk water. Outer surfaces of biomolecules form more stable hydrogen bonding with water, “ordering” the water surrounding them. This results in a decrease in entropy (increased order) and increase in free energy: factors which would strongly inhibit the solubility of biomolecules if not for the effects of hydrophobic interactions. Hydrophobic groups (for example amino acids whose side groups are non-polar, that is they have no charge-like polar groups to form hydrogen bonds in water) tend to combine, or coalesce for two main reasons: Van der Waals forces and exclusion of water. 

Combination of hydrophobic groups “liberate” ordered water into free water, resulting in increased entropy and decreased free energy, factors which tend to drive reactions. The magnitude of the favorable free energy change for the combination of hydrophobic groups depends on their size and how well they fit together “sterically.” A snug fit between groups will exclude more water from hydrophobic regions than will loose fits. Consequently, specific biological reactions can rely on hydrophobic interactions. Forma, tion of tertiary and quaternary protein structure (including the assembly of microtubules and other cytoskeletal polymers) are largely regulated by hydrophobic interactions, and by the effect of hydrophobic regions on the energies of other bonding. A well studied example of the assembly of protein subunits into a complex structure being accompanied by an increase in entropy (decrease in order) is the crystallization of the tobacco mosaic virus. When the virus assembles from its subunits, an increase in entropy occurs due to exclusion of water from the virus surface. Similar events promote the assembly of microtubules and other cytoskeletal elements The attractive forces which bind hydrophobic groups are distinctly different from other types of chemical bonds such as covalent bonds and ionic bonds. These forces are called Van der Waals forces after the Dutch chemist who described them in 1873. At that time, it had been experimentally observed that gas molecules failed to follow behavior predicted by the “ideal gas laws” regarding pressure, temperature and volume relationships. Van der Waals attributed this deviation to the volume occupied by the gas molecules and by attractive forces among the gas molecules. These same attractive forces are vital to the assembly of organic crystals, including protein assemblies. They consist of dipole-dipole attraction, “induction effect,” and London dispersion forces. These hydrophobic Van der Waals forces are subtly vital to the assembly and function of important biomolecules.

Dipole-dipole attractions occur among molecules with permanent dipole moments. Only specific orientations are favored: alignments in which attractive, low energy arrangements occur as opposed to repulsive, high energy orientations. A net attraction between two polar molecules can result if their dipoles are properly configured. The “induction” effect occurs when a permanent dipole in one molecule can polarize electrons in a nearby molecule. The second molecule’s electrons are distorted so that their interaction with the dipole of the first molecule is attractive. The magnitude of the induced dipole attraction force was shown by Debye in 1920 to depend on the molecules’ dipole moments and their polarizability. Defined as the dipole moment induced by a standard field, polarizability also depends on the molecules’ orientation relative to that field. Subunits of protein assemblies like the tobacco mosaic virus have been shown to have high degrees of polarizability. London dispersion forces explain why all molecules, even those without intrinsic dipoles, attract each other. The effect was recognized by F. London in 1930 and depends on quantum mechanical motion of electrons. Electrons in atoms without permanent dipole moments (and “shared” electrons in molecules) have, on the average, a zero dipole, however “instantaneous dipoles” can be recognized. Instantaneous dipoles can induce dipoles in neighboring polarizable atoms or molecules. The strength of London forces is proportional to the square of the polarizability and inversely to the sixth power of the separation. Thus London forces can be significant only when two or more atoms or molecules are very close together (Barrow, 1966). Lindsay (1987) has observed that water and ions ordered on surfaces of biological macromolecules may have “correlated fluctuations” analogous to London forces among electrons. Although individually tenuous, these and other forces are the collective “glue” of dynamic living systems.