The water molecule (H2O) is an ideal solvent for physiologic reactions. H2O has a dipole moment where oxygen slightly pulls away electrons from the hydrogen atoms and creates a charge separation that makes the molecule polar. This allows water to dissolve a variety of charged atoms and molecules. It also allows the H2O molecule to interact with other H2O molecules via hydrogen bonding. The resulting hydrogen bond network in water allows for several key properties relevant to physiology: (1) water has a high surface tension, (2) water has a high heat of vaporization and heat capacity, and (3) water has a high dielectric constant. In layperson’s terms, H2O is an excellent biologic fluid that serves as a solute; it provides optimal heat transfer and conduction of current.
- Solvent
- Water Influences the Structure of Biomolecules
- Excellent Nucleophile for Metabolic Reactions
- Acts as Both Acid and Base
Water has the ability to solvate a wide range of organic and inorganic molecules. Its ability is derived from (1) water's dipolar structure and (2) exceptional capacity for forming hydrogen bonds.
1. Water Molecules Form Dipoles
A molecule with electrical charge distributed asymmetrically about its structure is referred to as a dipole. Water's strong dipole is responsible for its high dielectric constant. As described quantitatively by Coulomb's law, the strength of interaction F between oppositely charged particles is inversely proportionate to the dielectric constant ε of the surrounding medium. The dielectric constant for a vacuum is unity and for water, 78.5. Water therefore greatly decreases the force of attraction between charged and polar species relative to water-free environments with lower dielectric constants. Its strong dipole and high dielectric constant enable water to dissolve large quantities of charged compounds such as salts.
2. Water Molecules Form Hydrogen Bonds
A partially unshielded hydrogen nucleus covalently bound to an electron-withdrawing oxygen or nitrogen atom can interact with an unshared electron pair on another oxygen or nitrogen atom to form a hydrogen bond. Since water molecules contain both of these features, hydrogen bonding favors the self-association of water molecules into ordered arrays.
Left: Association of two dipolar water molecules by a hydrogen bond (dotted line). Right: Hydrogen-bonded cluster of four water molecules. Note that water can serve simultaneously both as a hydrogen donor and as a hydrogen acceptor.
Hydrogen bonding profoundly influences the physical properties of water and accounts for its exceptionally high viscosity, surface tension, and boiling point. On average, each molecule in liquid water associates through hydrogen bonds with 3.5 others.
Hydrogen bonding enables water to dissolve many organic biomolecules that contain functional groups which can participate in hydrogen bonding. The oxygen atoms of aldehydes, ketones, and amides, for example, provide lone pairs of electrons that can serve as hydrogen acceptors. Alcohols, carboxylic acids, and amines can serve both as hydrogen acceptors and as donors of unshielded hydrogen atoms for formation of hydrogen bonds.
Covalent and Noncovalent Bonds Stabilize Biologic Molecules
The covalent bond is the strongest force that holds molecules together.
Noncovalent forces, while of lesser magnitude, make significant contributions to the structure, stability, and functional competence of macromolecules in living cells. These forces, which can be either attractive or repulsive, involve interactions both within the biomolecule and between it and the water that forms the principal component of the surrounding environment.
Biomolecules Fold to Position Polar & Charged Groups on Their Surfaces
Most biomolecules are amphipathic; that is, they possess regions rich in charged or polar functional groups as well as regions with hydrophobic character. Proteins tend to fold with the R-groups of amino acids with hydrophobic side chains in the interior. Amino acids with charged or polar amino acid side chains (eg, arginine, glutamate, serine) generally are present on the surface in contact with water. A similar pattern prevails in a phospholipid bilayer, where the charged “head groups” of phosphatidyl serine or phosphatidyl ethanolamine contact water while their hydrophobic fatty acyl side chains cluster together, excluding water. This pattern maximizes the opportunities for the formation of energetically favorable charge–dipole, dipole–dipole, and hydrogen bonding interactions between polar groups on the biomolecule and water. It also minimizes energetically unfavorable contacts between water and hydrophobic groups.
Hydrophobic Interactions
Hydrophobic interaction refers to the tendency of nonpolar compounds to self-associate in an aqueous environment. This self-association is driven neither by mutual attraction nor by what are sometimes incorrectly referred to as “hydrophobic bonds.” Self-association minimizes the disruption of energetically favorable interactions between the surrounding water molecules.
While the hydrogens of nonpolar groups such as the methylene groups of hydrocarbons do not form hydrogen bonds, they do affect the structure of the water that surrounds them. Water molecules adjacent to a hydrophobic group are restricted in the number of orientations ( degrees of freedom) that permit them to participate in the maximum number of energetically favorable hydrogen bonds. Maximal formation of multiple hydrogen bonds, which maximizes enthalpy, can be maintained only by increasing the order of the adjacent water molecules, with an accompanying decrease in entropy.
It follows from the second law of thermodynamics that the optimal free energy of a hydrocarbon–water mixture is a function of both maximal enthalpy (from hydrogen bonding) and minimum entropy (maximum degrees of freedom). Thus, nonpolar molecules tend to form droplets that minimize exposed surface area and reduce the number of water molecules whose motional freedom becomes restricted. Similarly, in the aqueous environment of the living cell the hydrophobic portions of biopolymers tend to be buried inside the structure of the molecule, or within a lipid bilayer, minimizing contact with water.
Electrostatic Interactions
Interactions between charged groups help shape biomolecular structure. Electrostatic interactions between oppositely charged groups within or between biomolecules are termed salt bridges. Salt bridges are comparable in strength to hydrogen bonds but act over larger distances. They therefore often facilitate the binding of charged molecules and ions to proteins and nucleic acids.
van der Waals Forces
van der Waals forces arise from attractions between transient dipoles generated by the rapid movement of electrons of all neutral atoms. Significantly weaker than hydrogen bonds but potentially extremely numerous, van der Waals forces decrease as the sixth power of the distance separating atoms. Thus, they act over very short distances, typically 2–4 Å.
The strength of van der Waals interactions varies with the distance, R, between interacting species. The force of interaction between interacting species increases with decreasing distance until they are separated by the van der Waals contact distance (see arrow marked A). Repulsion due to interaction between the electrons of each atom or molecule then supervenes. While individual van der Waals interactions are extremely weak, the cumulative effect is nevertheless substantial for macromolecules such as DNA and proteins with many atoms in close contact.
Multiple Forces Stabilize Biomolecules
The DNA double helix illustrates the contribution of multiple forces to the structure of biomolecules. While each individual DNA strand is held together by covalent bonds, the two strands of the helix are held together exclusively by noncovalent interactions such as hydrogen bonds between nucleotide bases (Watson–Crick base pairing) and van der Waals interactions between the stacked purine and pyrimidine bases. The double helix presents the charged phosphate groups and polar hydroxyl groups from the ribose sugars of the DNA backbone to water while burying the relatively hydrophobic nucleotide bases inside. The extended backbone maximizes the distance between negatively charged phosphates, minimizing unfavorable electrostatic interactions.
Metabolic reactions often involve the attack by lone pairs of electrons residing on electron-rich molecules termed nucleophiles upon electron-poor atoms called electrophiles.
Water, whose two lone pairs of sp3 electrons bear a partial negative charge, is an excellent nucleophile.
Nucleophilic attack by water typically results in the cleavage of the amide, glycoside, or ester bonds that hold biopolymers together. This process is termed hydrolysis. Conversely, when monomer units are joined together to form biopolymers such as proteins or glycogen, water is a product, for example, during the formation of a peptide bond between two amino acids:
While hydrolysis is a thermodynamically favored reaction, the amide and phosphoester bonds of polypeptides and oligonucleotides are stable in the aqueous environment of the cell. This seemingly paradoxic behavior reflects the fact that the thermodynamics governing the equilibrium of a reaction do not determine the rate at which it will proceed. In the cell, protein catalysts called enzymes accelerate the rate of hydrolytic reactions when needed. Proteases catalyze the hydrolysis of proteins into their component amino acids, while nucleases catalyze the hydrolysis of the phosphoester bonds in DNA and RNA. Careful control of the activities of these enzymes is required to ensure that they act only on appropriate target molecules at appropriate times.
Many Metabolic Reactions Involve Group Transfer
Many of the enzymic reactions responsible for synthesis and breakdown of biomolecules involve the transfer of a chemical group G from a donor D to an acceptor A to form an acceptor group complex, A—G:
The hydrolysis and phosphorolysis of glycogen, for example, involve the transfer of glucosyl groups to water or to orthophosphate. The equilibrium constant for the hydrolysis of covalent bonds strongly favors the formation of split products. Conversely, in many cases the group transfer reactions responsible for the biosynthesis of macromolecules involve the thermodynamically unfavored formation of covalent bonds. Enzyme catalysts play a critical role in surmounting these barriers by virtue of their capacity to directly link two normally separate reactions together. By marrying an energetically unfavorable group transfer reaction with a thermodynamically favorable reaction, such as the hydrolysis of ATP, a new coupled reaction can be generated whose net overall change in free energy favors biopolymer synthesis.
Given the nucleophilic character of water and its high concentration in cells, why are biopolymers such as proteins and DNA relatively stable? And how can synthesis of biopolymers occur in an aqueous, seemingly prohydrolytic, environment? Central to both questions are the properties of enzymes. In the absence of enzymic catalysis, even reactions that are highly favored thermodynamically do not necessarily take place rapidly. Precise and differential control of enzyme activity and the sequestration of enzymes in specific organelles determine under what physiologic conditions a given biopolymer will be synthesized or degraded. Newly synthesized biopolymers are not immediately hydrolyzed because the active sites of biosynthetic enzymes sequester substrates in an environment from which water can be excluded.
Since water can act both as an acid and as a base, its ionization may be represented as an intermolecular proton transfer that forms a hydronium ion (H3O+) and a hydroxide ion (OH−):
The transferred proton is actually associated with a cluster of water molecules. Protons exist in solution not only as H3O+, but also as multimers such as H5O2+ and H7O3+. The proton is nevertheless routinely represented as H+, even though it is in fact highly hydrated.
Since hydronium and hydroxide ions continuously recombine to form water molecules, an individual hydrogen or oxygen cannot be stated to be present as an ion or as part of a water molecule. At one instant it is an ion; an instant later it is part of a water molecule. Individual ions or molecules are therefore not considered. We refer instead to the probability that at any instant in time a given hydrogen will be present as an ion or as part of a water molecule. Since 1 g of water contains 3.46 × 1022 molecules, the ionization of water can be described statistically. To state that the probability that a hydrogen exists as an ion is 0.01 means that at any given moment in time, a hydrogen atom has 1 chance in 100 of being an ion and 99 chances out of 100 of being part of a water molecule. The actual probability of a hydrogen atom in pure water existing as a hydrogen ion is approximately 1.8 × 10−9. The probability of its being part of a water molecule thus is almost unity. Stated another way, for every hydrogen ion or hydroxide ion in pure water, there are 1.8 billion or 1.8 × 109 water molecules. Hydrogen ions and hydroxide ions nevertheless contribute significantly to the properties of water.