BIO202 Handouts Lectures 1-45

Water, pH and Acid-base balance

Water is the most abundant substance in living systems, making up 70% or more of the weight of most organisms. The water molecule and its ionization products, H and OH, profoundly influence the structure, self assembly, and properties of all cellular components, including proteins, nucleic acids, and lipids. Hydrogen bonds between water molecules provide the cohesive forces that make water a liquid at room temperature and that favor the extreme ordering of molecules that is typical of crystalline water (ice). Polar biomolecules dissolve readily in water because they can replace water-water interactions with more energetically favorable water-solute interactions. Water has a higher melting point, boiling point, and heat of vaporization than most other common solvents. Each hydrogen atom of a water molecule shares an electron pair with the central oxygen atom. The HOH bond angle is 104.5.

The oxygen nucleus attracts electrons more strongly than does the hydrogen nucleus (a proton); that is, oxygen is more electronegative. The sharing of electrons between H and O is therefore unequal; the electrons are more often in the vicinity of the oxygen atom than of the hydrogen. The result of this unequal electron sharing is two electric dipoles in the water molecule, one along each of the HO bonds; each hydrogen bears a partial positive charge ( +) and the oxygen atom bears a partial negative charge. As a result, there is an electrostatic attraction between the oxygen atom of one water molecule and the hydrogen of another called a hydrogen bond. Hydrogen bonds account for the relatively high melting point of water, because much thermal energy is required to break a sufficient proportion of hydrogen bonds to destabilize the crystal lattice of ice. When ice melts or water evaporates, heat is taken up by the system: H2O (solid)  H2O (liquid) ∆H = +5.9 kj/mol H2O (liquid)  H2O (gas) ∆H = +44.0 kj/mol Hydrogen atoms covalently bonded to carbon atoms do not participate in hydrogen bonding, because carbon is only slightly more electronegative than hydrogen and thus the COH bond is only very weakly polar. Uncharged but polar biomolecules such as sugars dissolve readily in water because of the stabilizing effect of hydrogen bonds. Alcohols, aldehydes, ketones, and compounds containing N−H bonds all form hydrogen bonds with water molecules and tend to be soluble in water. The attraction between the partial electric charges is greatest when the three atoms involved (in this case O, H, and O) lie in a straight line. Water is a polar solvent. It readily dissolves most biomolecules, which are generally charged or polar compounds; compounds that dissolve easily in water are hydrophilic (Greek, ―water-loving‖). In contrast, nonpolar solvents such as chloroform and benzene are poor solvents for polar biomolecules but easily dissolve those that are hydrophobic—nonpolar molecules such as lipids and waxes. Gases CO2, O2, and N2 are nonpolar. In CO2, each C=O bond is polar, but the two dipoles are oppositely directed and cancel each other. Note that molecules dissolve far better even at low temperatures than do nonpolar molecules at relatively high temperatures.

Amphipathic:

Compounds contain regions that are polar (or charged) and regions that are nonpolar. Hydrophilic region interacts favorably with the solvent and tends to dissolve, but the nonpolar, hydrophobic region tends to avoid contact with the water. The nonpolar regions of the molecules cluster together. These stable structures of amphipathic compounds in water, called micelles. When two uncharged atoms are brought very close together, electrons around one nucleus may create a transient electric dipole, which induces a transient, opposite electric dipole in the nearby atom. The two dipoles weakly attract each other, bringing the two nuclei closer. These weak attractions are called van der Waals interactions. As the two nuclei draw closer together, their electron clouds begin to repel each other. At the point where the van der Waals attraction exactly balances this repulsive force, the nuclei are said to be in van der Waals contact. Each atom has a characteristic van der Waals radius, a measure of how close that atom will allow another to approach. When two atoms are joined covalently, the atomic radii at the point of bonding are less than the van der Waals radii, because the joined atoms are pulled together by the shared electron pair. Solutes of all kinds alter certain physical properties of the solvent, water: its vapor pressure, boiling point, melting point (freezing point), and osmotic pressure. These are called colligative (―tied together‖) properties, because the effect of solutes on all four properties has the same basis: the concentration of water is lower in solutions than in pure water. Independent of the chemical properties of the solute; it depends only on the number of solute particles (molecules, ions) in a given amount of water. A compound such as NaCl, which dissociates in solution, has twice the effect on osmotic pressure, for example, as does an equal number of moles of a nondissociating solute such as glucose. When two different aqueous solutions are separated by a semipermeable membrane (one that allows the passage of water but not solute molecules), water molecules diffusing from the