Lodish H. - Molecular Cell Biology (5ed, Freeman, 2003) (794361), страница 21
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In some cases,these binding reactions are independent, with their own distinct Kd values that are constant. In other cases, binding ofa molecule at one site on a macromolecule can change thethree-dimensional shape of a distant site, thus altering thebinding interactions at that distant site. This is an importantmechanism by which one molecule can alter (regulate) theactivity of a second molecule (e.g., a protein) by changing itscapacity to interact with a third molecule. We examine thisregulatory mechanism in more detail in Chapter 3.Biological Fluids Have Characteristic pH ValuesThe concept of chemical equilibrium also applies to the binding of one molecule to another.
Many important cellularprocesses depend on such binding “reactions,” which involvethe making and breaking of various noncovalent interactionsrather than covalent bonds, as discussed above. A commonexample is the binding of a ligand (e.g., the hormone insulinor adrenaline) to its receptor on the surface of a cell, triggering a biological response. Another example is the binding ofa protein to a specific sequence of base pairs in a moleculeof DNA, which frequently causes the expression of a nearbygene to increase or decrease (Chapter 11).
If the equilibriumconstant for a binding reaction is known, the intracellularstability of the resulting complex can be predicted. To illustrate the general approach for determining the concentrationof noncovalently associated complexes, we will calculate theextent to which a protein is bound to DNA in a cell.Most commonly, binding reactions are described in termsof the dissociation constant Kd, which is the reciprocal of theequilibrium constant.
For the binding reaction P DPD, where PD is the specific complex of a protein (P) andDNA (D), the dissociation constant is given byKd 47(2-4)Typical reactions in which a protein binds to a specific DNAsequence have a Kd of 1010M, where M symbolizes molarity, or moles per liter (mol/L). To relate the magnitude of thisdissociation constant to the intracellular ratio of bound tounbound DNA, let’s consider the simple example of a bacterial cell having a volume of 1.5 1015 L and containingThe solvent inside cells and in all extracellular fluids is water.An important characteristic of any aqueous solution is theconcentration of positively charged hydrogen ions (H) andnegatively charged hydroxyl ions (OH). Because these ionsare the dissociation products of H2O, they are constituentsof all living systems, and they are liberated by many reactionsthat take place between organic molecules within cells.When a water molecule dissociates, one of its polarHOO bonds breaks.
The resulting hydrogen ion, often referred to as a proton, has a short lifetime as a free particleand quickly combines with a water molecule to form a hydronium ion (H3O). For convenience, however, we refer tothe concentration of hydrogen ions in a solution, [H], eventhough this really represents the concentration of hydroniumions, [H3O]. Dissociation of H2O generates one OH ionalong with each H. The dissociation of water is a reversiblereaction,H2 OH OHAt 25 ºC, [H][OH] 1014 M2, so that in pure water,[H] [OH] 107 M.The concentration of hydrogen ions in a solution is expressed conventionally as its pH, defined as the negative logof the hydrogen ion concentration. The pH of pure water at25 ºC is 7:pH log[H ] log117 log[H ]107It is important to keep in mind that a 1 unit difference inpH represents a tenfold difference in the concentration of48CHAPTER 2 • Chemical Foundationsprotons.
On the pH scale, 7.0 is considered neutral: pH values below 7.0 indicate acidic solutions (higher [H]), andvalues above 7.0 indicate basic (alkaline) solutions. For instance, gastric juice, which is rich in hydrochloric acid (HCl),has a pH of about 1. Its [H] is roughly a millionfold greaterthan that of cytoplasm with a pH of about 7.Although the cytosol of cells normally has a pH of about7.2, the pH is much lower (about 4.5) in the interior of lysosomes, one type of organelle in eukaryotic cells. The manydegradative enzymes within lysosomes function optimally inan acidic environment, whereas their action is inhibited inthe near neutral environment of the cytoplasm.
This illustrates that maintenance of a specific pH is imperative forproper functioning of some cellular structures. On the otherhand, dramatic shifts in cellular pH may play an importantrole in controlling cellular activity. For example, the pH ofthe cytoplasm of an unfertilized sea urchin egg is 6.6. Within1 minute of fertilization, however, the pH rises to 7.2; that is,the H concentration decreases to about one-fourth its original value, a change that is necessary for subsequent growthand division of the egg.The equilibrium constant for this reaction, denoted Ka (subscript a for “acid”), is defined as Ka [H][A]/[HA].
Taking the logarithm of both sides and rearranging the resultyields a very useful relation between the equilibrium constantand pH:pH pKa log[A][HA](2-5)where pKa equals –log Ka.From this expression, commonly known as the HendersonHasselbalch equation, it can be seen that the pKa of any acidis equal to the pH at which half the molecules are dissociatedand half are neutral (undissociated). This is because when pKa pH, then log ([A]/[HA]) 0, and therefore [A] [HA].The Henderson-Hasselbalch equation allows us to calculate thedegree of dissociation of an acid if both the pH of the solutionand the pKa of the acid are known. Experimentally, by measuring the [A] and [HA] as a function of the solution’s pH, onecan calculate the pKa of the acid and thus the equilibriumconstant Ka for the dissociation reaction.Hydrogen Ions Are Releasedby Acids and Taken Up by BasesBuffers Maintain the pH of Intracellularand Extracellular FluidsIn general, an acid is any molecule, ion, or chemical groupthat tends to release a hydrogen ion (H), such as hydrochloric acid (HCl) and the carboxyl group (OCOOH),which tends to dissociate to form the negatively charged carboxylate ion (OCOO).
Likewise, a base is any molecule,ion, or chemical group that readily combines with a H,such as the hydroxyl ion (OH), ammonia (NH3), whichforms an ammonium ion (NH4), and the amino group(ONH2).When acid is added to an aqueous solution, the [H] increases (the pH goes down). Conversely, when a base isadded to a solution, the [H] decreases (the pH goes up). Because [H][OH] 1014M2, any increase in [H] is coupled with a decrease in [OH], and vice versa.Many biological molecules contain both acidic and basicgroups.
For example, in neutral solutions (pH 7.0), aminoacids exist predominantly in the doubly ionized form inwhich the carboxyl group has lost a proton and the aminogroup has accepted one:A growing cell must maintain a constant pH in the cytoplasm of about 7.2–7.4 despite the metabolic production ofmany acids, such as lactic acid and carbon dioxide; the latter reacts with water to form carbonic acid (H2CO3).
Cellshave a reservoir of weak bases and weak acids, calledbuffers, which ensure that the cell’s pH remains relativelyconstant despite small fluctuations in the amounts of H orOH being generated by metabolism or by the uptake or secretion of molecules and ions by the cell. Buffers do this by“soaking up” excess H or OH when these ions are addedto the cell or are produced by metabolism.If additional acid (or base) is added to a solution thatcontains a buffer at its pKa value (a 1:1 mixture of HA andA), the pH of the solution changes, but it changes less thanit would if the buffer had not been present.
This is becauseprotons released by the added acid are taken up by the ionized form of the buffer (A); likewise, hydroxyl ions generated by the addition of base are neutralized by protonsreleased by the undissociated buffer (HA). The capacity of asubstance to release hydrogen ions or take them up dependspartly on the extent to which the substance has already takenup or released protons, which in turn depends on the pH ofthe solution.
The ability of a buffer to minimize changes inpH, its buffering capacity, depends on the relationship between its pKa value and the pH, which is expressed by theHenderson-Hasselbalch equation.The titration curve for acetic acid shown in Figure 2-22illustrates the effect of pH on the fraction of molecules inthe un-ionized (HA) and ionized forms (A). At one pH unitbelow the pKa of an acid, 91 percent of the molecules are inthe HA form; at one pH unit above the pKa, 91 percent areNH3HCCOORwhere R represents the side chain.
Such a molecule, containing an equal number of positive and negative ions, is calleda zwitterion. Zwitterions, having no net charge, are neutral.At extreme pH values, only one of these two ionizablegroups of an amino acid will be charged.The dissociation reaction for an acid (or acid group in alarger molecule) HA can be written as HAH A .2.3 • Chemical Equilibrium8CH3COOHCH3COO − + H +6pHpK a = 4.754200.20.40.60.81.0Fraction of dissociated CH3COOHAdded OH−▲ FIGURE 2-22 The titration curve of acetic acid(CH3COOH). The pKa for the dissociation of acetic acid tohydrogen and acetate ions is 4.75. At this pH, half the acidmolecules are dissociated.
Because pH is measured on a logarithmicscale, the solution changes from 91 percent CH3COOH atpH 3.75 to 9 percent CH3COOH at pH 5.75. The acid has maximumbuffering capacity in this pH range.49in the A form. At pH values more than one unit above orbelow the pKa, the buffering capacity of weak acids andbases declines rapidly. In other words, the addition of thesame number of moles of acid to a solution containing a mixture of HA and A that is at a pH near the pKa will cause lessof a pH change than it would if the HA and A were notpresent or if the pH were far from the pKa value.All biological systems contain one or more buffers.
Phosphate ions, the ionized forms of phosphoric acid, are present in considerable quantities in cells and are an importantfactor in maintaining, or buffering, the pH of the cytoplasm.Phosphoric acid (H3PO4) has three protons that are capableof dissociating, but they do not dissociate simultaneously.Loss of each proton can be described by a discrete dissociation reaction and pKa as shown in Figure 2-23. The titrationcurve for phosphoric acid shows that the pKa for the dissociation of the second proton is 7.2. Thus at pH 7.2, about 50percent of cellular phosphate is H2PO4 and about 50 percent is HPO42 according to the Henderson-Hasselbalchequation. For this reason, phosphate is an excellent bufferat pH values around 7.2, the approximate pH of the cytoplasm of cells, and at pH 7.4, the pH of human blood.KEY CONCEPTS OF SECTION 2.3Chemical EquilibriumA chemical reaction is at equilibrium when the rate ofthe forward reaction is equal to the rate of the reverse reaction (no net change in the concentration of the reactantsor products).■14pKa = 12.7HPO 42−12The equilibrium constant Keq of a reaction reflects theratio of products to reactants at equilibrium and thus is ameasure of the extent of the reaction and the relative stabilities of the reactants and products.■PO 43− + H+10pH8pKa = 7.2H2PO4−HPO 42− + H+6■ For any reaction, the equilibrium constant Keq equalsthe ratio of the forward rate constant to the reverse rateconstant (kf /kr).4pKa = 2.12■ The Keq depends on the temperature, pressure, andchemical properties of the reactants and products, but isindependent of the reaction rate and of the initial concentrations of reactants and products.H3PO4H2PO 4− + H+0Added OH−▲ FIGURE 2-23 The titration curve of phosphoric acid(H3PO4).
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