Ординатура / Офтальмология / Английские материалы / Biochemistry of the Eye 2nd edition_Whikehart_2003

2• Biochemistry of the Eye

polar nature to the molecule and makes it attractive to oppositely charged ions on the appropriate sides of the water molecule. Other polar molecules, such as ethanol (ethyl alcohol), are similarly attracted toward water molecules on the appropriate oppositely partially charged sides. In biological fluids, such as cell cytoplasm, this attractive property of water helps large molecules (e.g., proteins) to maintain both their presence (soluble state) and functional conformation (shape). Furthermore, water acts as a containment medium for the many smaller molecules (e.g., ions and sugars) whose presence is also required for biological functions.

Solubility of small and large molecules in water and biological fluids is, therefore, explained partly in terms of the ability of water to associate with charged atoms by the use of its partially charged regions (polar interactions). This form of solvation is shown in Figure 1–2 for sodium chloride (common table salt). The partially, positively charged hydrogen atoms of water tend to surround and pull negatively charged chloride ions away from salt crystals while the partially, negatively charged oxygen atoms of water tend to surround and pull positively charged sodium ions away from the same crystals. A second form of solvation is also explained in terms of the ability of water to associate with other polar molecules by means of hydrogen bonding.

WATER MOLECULE

δ+

δ+

Water and Ocular Fluids • 3

HYDROGEN BONDS

These are weak bonds that form between a hydrogen (whose electrons have been delocalized or pulled away by an electronegative element such as oxygen) and a second electronegative atom (commonly oxygen or nitrogen). Hydrogen bonds are not unique to water, but water constantly forms hydrogen bonds with its nearest neighbor water molecules both in the liquid and solid states (Figure 1–3). During solvation, water may also form hydrogen bonds with polar portions of proteins as seen in Figure 1–4. As shown, a water molecule forms a hydrogen bond with an amino group (R-NH2) on a protein. Hydrogen bonds are individually quite weak (approximately 14 kJ/mol) (Table 1–1), but their aggregate strength, especially in proteins may be very strong. This explains how water can be a very powerful solvent for large molecules.

Weak Electrolytes and Buffers in Water

Cells and tissues (with some notable exceptions such as cells that line the stomach) are bathed in aqueous solutions that must maintain a consistent near neutral concentration of hydrogen ions. The hydrogen ion concentration in these cells is usually at 1 × 10–7.4 M and is necessary to sustain protein structure and function (i.e., life itself). This concentration is conveniently expressed as its negative logarithm value of 7.4 and is the

Figure 1–3

Hydrogen bonds between water molecules. Water molecules form “temporary” hydrogen bonds, that is, bonds that are constantly made and broken. When water freezes, the bonds become permanent giving water a more ordered structure and it expands.

HYDROGEN BONDS

4• Biochemistry of the Eye

HYDROGEN

BOND

WATER

MOLECULE

T A B L E 1 – 1 STABILIZATION ENERGIES FOR WEAK AND STRONG

MOLECULAR BONDS

1A kilojoule (kJ) is 1000 × the amount of energy necessary to accelerate a 1 kg mass through a distance of 1 meter per sec2. One kJ is equal to .239 kilocalories. The values are given as the amounts necessary per molecule (mol) to break one bond. See Appendix.

(Data from Weast, 1986; Daniels, Alberty, 1961; Lehninger, Nelson, Cox, 1993.)

most common pH value for biological fluids both inside and outside of cells. The term “pH” was first explained and used by the Scandinavian biochemist Sørensen in 1909 (Voet, Voet, 1995). The pH is controlled primarily by weak electrolytes present in both ocular and nonocular biological, aqueous solutions. These weak electrolytes are acids that only partially ionize (form ions) in solution. Due to their partial ionization, they can either take up or donate hydrogen ions to control solution pH.

An example of a weak electrolyte is acetic acid, the acid found in household vinegar.

In order to understand how buffers function, it is necessary to grasp the concept of how weak electrolyte ionization can be determined. This is done by comparing the concentrations of the ionized to the nonionized forms to obtain a ratio known as the dissociation constant. Therefore, one uses the expression:

Water and Ocular Fluids • 5

Ka = [H+] [A–] / [HA]

where Ka is the dissociation constant; [H+] is the molar hydrogen ion concentration; [A–] is the molar anion concentration; and [HA] is the molar concentration of the undissociated (nonionized) acid. Using the example of acetic acid, it has been found that the dissociation constant (Ka) for this acid is: 1.74 × 10–5. The negative logarithm or pKa of that value = 4.76. Using the negative logarithm of the Ka is convenient to determine the pH of a solution of this acid.

Now, if you are getting a little bored with this discussion, go open a bottle of vinegar and sniff it. The characteristic odor is acetate anion from the acetic acid and the sour taste (take a little taste) is from the hydrogen ions dissociated in the vinegar. Although the amount dissociated is only 0.000009 M from the total of 0.5 M present in the vinegar, it is potent to one’s senses! It is easily realized then that not much is needed to produce a pH effect in biological systems including the eye.

Let us return to the pKa value for acetic acid (4.76). This value is equivalent to the pH value of a solution of acetic acid when it is 50% ionized. Its pH can be changed, but only with difficulty, by adding acid or base to the weak electrolyte (i.e., the acetic acid/acetate mixture). This is a desirable property since, by resisting pH changes, the weak electrolyte acts as a buffer to pH change and does so by ionizing to a greater extent whenever base enters the solution and by ionizing to a lesser extent whenever acid enters the solution (Figure 1–5). The capacity of the buffer is the extent to which it will absorb acid or base and depends on its concentration in solution. The range of the buffer is the pH limit (increased or decreased) that will be buffered and this depends on the pKa of the buffer to act as the mid-range value. The range will extend approximately 1 pH unit above and below the pKa of the buffer as shown in Figure 1–6. Buffering is a very useful and necessary property of all biological fluids both ocular and nonocular.

THE HENDERSON-HASSELBALCH EQUATION

This equation is useful in the determination of pH, pKa, and the relative concentrations of ionized and nonionized components of a buffer. It is important in the preparation of buffers to be used in clinical and research laboratories and in the formulation of drugs and drug vehicles that require buffers on topical installation of a drug (such as ocular drugs, which are placed on the corneal surface). The equation is:

pH = pKa + log [salt] / [acid]

where pH is the negative logarithm of the hydrogen ion concentration, pKa is the negative logarithm of the dissociation constant of the weak electrolyte, and log [A–]/[HA] is the logarithm of the ratio of the ionized anion to the nonionized acid of the weak electrolyte. The derivation of the equation may be found in a number of contemporary biochemical texts (Lehninger, Nelson, Cox, 1993). Examples of problems involving the Henderson-Hasselbalch equation are found at the end of this chapter.

Water and Ocular Fluids • 7

This system is more complex in that the carbon dioxide (CO2) can be removed with the expired air and, therefore, one’s breathing rate can influence the HCO3–/H2CO3 ratio and, conclusively, the extracellular pH of blood and all ocular fluids. As a consequence of this system, an elevated breathing rate would tend to raise the pH by shifting both arrows to the left as more CO2 is removed from the body while the converse would hold true. Removing CO2 converts HCO3– to H2CO3 and then to CO2. The H+ is incorporated into water as CO2 is formed. One can produce respiratory alkalosis just by hyperventilating (forced rapid breathing). The pH in one’s own blood may be made to rise to as high as 7.9!

Protein buffers, it turns out, are the most important buffers both inside and outside of cells (Mathews, van Holde, 1990). This is due to the presence of the myriad number of acidic and basic groups, which are present on protein amino acid components, and the fact that there are so many proteins present both inside and outside of cells. An example of an ionizable group on a protein is the glutamate ion:

R1 – NH – CH – CO – R2

CH2 – CH2 – COO–

Its acidic properties are balanced by basic groups such as that found on the amino acid lysine on the same protein. (See Chapter 2 for further information on the amino acid components of proteins.) In this manner, the pH buffering from a single protein is quite sophisticated. Moreover, the ionizable groups on proteins have altered pK values such that prediction of the exact buffering tendencies and capacities of proteins are impossible without experimental evidence.

Ocular Fluids

Ocular fluids that are found in common with nonocular regions of the body are cellular cytoplasm, interstitial fluid, and blood. Ocular fluids that are exclusive to the eye are aqueous fluid, the vitreous, and precorneal tears. All of these fluids make use of the particular physicalchemical properties of water to carry out their functions. No further mention will be made of cellular cytoplasm and the interstitial fluids as these subjects are covered in basic texts in cell biology and biochemistry. Their roles are identical in the eye.

BLOOD IN THE OCULAR GLOBE

Blood is a complex mixture of cellular and biochemical components that can be separated by centrifugation. The ocular physiological functions of blood include the nourishment and ridding of waste components of ocular cells; a source for the generation of the intraocular pressure of the eye; and a source for the formation of aqueous and vitreous fluids as well as the homeostasis of retinal functions.

The usual pH of blood is 7.4, but it can vary from 7.33 to 7.45 under normal circumstances (Tietz, 1976). Gases dissolved or carried in the blood include oxygen, nitrogen, and carbon dioxide. The partial