Photosynthesis Research 19:225-236 (1988) © Kluwer Academic Publishers, Dordrechl - Printed in the Netherlands
Minireview
Active transport, ion movements, and pH changes
I. The chemistry of pH changes N O R M A N E. G O O D Department of Botany and Plant Pathology, Michigan State University, East Lansing, MI 48824, U.S.A. Received 1 March 1988; accepted in revised form 15 June 1988 Abstract. The transport of substances across cell membranes may be the most fundamental activity of living things. When the substance transported is any ion there can be a change in the concentration of hydrogen ions on the two sides of the membrane. These hydrogen ion concentration changes are not caused by fluxes of hydrogen ions although fluxes of hydrogen ions may sometimes be involved. The reason for the apparent contradiction is quite simple. All aqueous systems are subject to two constraints: (1) to maintain the charge balance, the sum of the cationic charges must equal the sum of the anionic charges and (2) the product of the molar concentration of H + and the molar concentration of O H - , established and maintained by the association and the dissociation of water, remains always at 10 14. As a consequence the concentrations of H + and OH are determined uniquely by differences between the concentrations of the other cations and anions, with [H +] and [OH-] being dependent variables. Hydrogen ions and hydroxyl ions can be produced or consumed in local reactions whereas any strong ions such as Cl-, Mg 2+, or K + can be neither produced nor consumed in biological reactions. Further consequences of these truisms are outlined here in terms of the chemistry of the kinds of reactions which can lead to pH changes.
Introduction
The assumption that changes in pH can be equated to hydrogen ion fluxes can lead to serious misconceptions. This paper details the nature of the various kinds of reactions which really do change pH, especially in biological systems. The next paper then expands on implications of the arguments regarding possible mechanisms of ATP synthesis and reexamines experimental data. The counter-intuitive contention that pH changes cannot be the direct consequence of hydrogen ion movements is based on solid chemical dicta. All aqueous systems are subject to two constraints. First, in order to maintain essential electrical neutrality in the bulk of the liquid C÷ + H÷ =
A - + OH
226 where C + represents all of the cationic charges other than H +, and A represents all of the anionic charges other than O H - . The distinction between the pair C + and A - on the one hand and the pair H + and O H - on the other hand is profoundly important and it is regrettable that our language has not given us terms to express the difference. In general, C + and A- are independent variables. In contrast H + and O H - are dependent variables. That is because H + is involved in a host of reversible protonation reactions, not least the protonation of O H - to make H20, carried out in a vast 55 M proton source. The second constraint on aqueous media is that the product of the molar concentration of O H - and of the molar concentration of H + remains constant at 10 -14, with the reaction H + + O H - ~ H20 proceeding in such a way as to assure this constancy. Thus the charge balance equation dictates that an excess of C + is tantamount to an excess of O H - while the water dissociation equation dictates that an increase in O H - is associated with a decrease in H + . With these constraints in mind, let us consider the meaning o f p H changes and how one can go about changing pH. Only when the diversity of the mechanisms which alter the concentration of H + has been appreciated can we begin to seek the real meaning of the Mitchell-Jagendorf electron transport-dependent pH changes in suspensions of mitochondria or chloroplast lamellae, the pH changes commonly but perhaps inappropriately attributed to hydrogen ion "pumps". How differences in hydrogen ion activities arise in aqueous media In preparing the following list I have attempted to cover most types of reactions although I cannot guarantee that the list is complete. Some processes, such as the oxidation or reduction of metal electrodes by an electric current, are unlikely to have much relevance to biological systems but most of the processes of pH change described below probably do have biological counterparts. A cautionary note may be needed here. The equation C + + H + = A- + O H - is exact except near the surfaces of high charge density. The charged membranes of most biological systems are an example but probably do not have sufficient electrical capacitance to account for significant deviations from this crucial equation in the bulk of the surrounding aqueous phase. In any event, the measurements of pH that we make are almost always at a great distance from such charged membranes. Therefore the pH values measured by electrodes in biological suspensions reflect the equation exactly.
227
a) Addition of HCI to K+ X - (titration o f a buffer by a strong acid) Reaction: H + + C1- + K + + X -
~
K + + C1- + HX
In this case the concentration of H + is determined by the equilibrium constant of the association-dissociation reaction H + + X - ~- HX and therefore [H + ] = Kd [HX]/[X- ] where K~ is the dissociation constant. Thus changes in [H + ] are complex functions of the amount HC1 (not H + ] added and the amount of K + X - originally present. It should be noted that it is the input of the chloride ion in HC1 just as much as it is the input of hydrogen ion that changes the pH. It should also be noted that the relationship between increases in [H + ] and the amount of HC1 added is ever-changing as K + X - is titrated. Obviously H X can represent a large reservoir of hydrogen ions available to replace those disappearing for any reason. When we are titrating a base such as K + O H - or a buffering salt such as K + X - , neither the change in [H +] nor the energy stored in the accumulation of H + is proportional to the flux of "protons".
b) Introduction of H + through membranes exclusively permeable to hydrogen ions. p H meters and the glass electrode It is interesting to consider the working of a pH meter. The glass electrode depends on the fact that the special glass membrane is, in effect, appreciably permeable to hydrogen ions but almost impermeable to all other ions. Therefore hydrogen ions, unaccompanied by counter ions, tend to diffuse across the membrane down any activity gradient which may exist and in this case we are indeed dealing with a phenomenon equivalent to a flux of protons. The result is the development of a transmembrane electric charge just sufficient to stop net diffusion of H + . That is to say, the work capacity of the charge on the extra electrons on the negative side of the membrane is just equal and opposite to the work potential of the hydrogen ions "expanding" from the higher concentration to the lower concentration. But we cannot measure the work potential of electrons without using some of them to do work, albeit a very small number sometimes. How then can we use hydrogen ions diffusing down a concentration gradient if these same hydrogen ions cannot change their own concentration by moving, as our charge-balance equation implies? (Clearly we cannot hope to measure the work done in dissipating a concentration difference if the difference is not being dissipated.) This problem is only a problem if we insist on looking at a partial reaction. When we consider the other reactions involved in completing the circuit of the meter, we find that the flux of hydrogen ions is
228 indeed changing the pH but only because of the electrode reactions which produce and consume silver ions. As the uncompensated hydrogen ions flow down their concentration gradient across the glass barrier, silver ions are reduced to silver metal on the "downhill" side and silver metal is oxidized to silver ions on the "uphill" side. The production of silver ions from silver metal on the "uphill" side ties up an amount of chloride exactly balancing the amount of hydrogen ion leaving that side, while the reduction of silver ions on the "downhill" side releases an amount of chloride exactly balancing the number of H ÷ ions arriving from the "uphill" side. Thus changes in hydrogen ion concentration do occur, changes driven by the hydrogen ion concentration difference. But these changes in hydrogen ion concentration are balanced by changes in chloride ion (driven it is true by H + activity differences) and therefore C + and H + continues to equal A - + O H - . I have belaboured this instance because it seems at first glance to refute my contention that moving hydrogen ions without moving (or creating) other ions cannot change the pH. However this case is simply an example of the errors we fall into when we confuse partial reactions with real-world overall reactions. Partial reactions "float" with no reference point, they cannot be balanced, and by themselves they have no meaningful energy relations. Here the flux of hydrogen ions drives the change in pH by driving the reduction of silver ions and the release of chloride ions. Therefore the charge balance equation is in no way violated. This example, while instructive, probably has little relevance to biology because the parts of the reaction which involve the oxidation and reduction of silver and precipitation of almost insoluble silver chloride are in no way biological.
c) The production of cations and anions by oxidation and reduction: electron donors and hydrogen donors If we oxidize one metallic ion by another metallic ion, there is no change in the balance of C + and A and therefore no change in pH is to be expected. However if we oxidize a metal to a cation, or oxidize a cation to the next higher valence state without producing a compensating anion, there is an increase in the number of cationic charges and a corresponding drop in the concentration of H + . By the same reasoning, reductions such as the conversion of chlorine gas to chloride ion can cause an increase in the concentration of H + . Let it be noted that the production or consumption of H + in these instances come entirely from the dissociation or association of water or of protonated acids in a reaction which balances the positive and negative charges in the medium and need not be due to hydrogen ion or hydroxyl ion movement.
229 If we oxidize a hydrogen donor H2 X to non-ionic products there is, of course, no pH change. If, however, we oxidize H2X with ferric salts: 2(Fe 3+ + 3A 3-) + H2X --+ 2(Fe 2+ + 2A ) + X + 2H + + 2Athere is a net decrease in the number of cationic charges and the pH falls. In the latter reaction we usually think of the production of hydrogen ions as coming from the oxidation of the hydrogen atoms in H2X and at first thought it seems logical to attribute the new hydrogen ions to the oxidation of H to H + , but in thinking thus we are considering a half-reaction which is incapable of changing pH by itself. In fact, the reason for the acidification is better described in terms of the decrease in C + which accompanies the disappearance of cationic charges when Fe 3+ becomes Fe 2+ . When we fit the loss of a cationic charge into the equation C + + H + = A - + O H - it is easier to see the real nature of the pH change. This distinction may be important because much has been made of the undoubted importance of "hydrogen donors" and so called proton pumps in phosphorylationcoupled electron transport.
d) Electrophoresis If one passes an electric current through an aqueous medium, the current is carried entirely by the charges on the dissolved ions when these ions migrate in the electric field. Thus cations, including H + , migrate toward the negative electrode (the cathode) and anions, including O H - , migrate toward the positive electrode (the anode). Since the C + ions accumulate toward the negative pole and the A - ions accumulate toward the positive pole, one would expect the concentration of hydrogen ions to decrease toward the negative pole even though H + is migrating into that region. That is exactly what happens. The situation is complicated, however, by the electrode reactions which introduce or remove ions. Nevertheless it is clear that the migrations of H + and O H - have no effect on the pH of water since pure water remains at pH 7.0 forever, no matter how long it is subjected to electrophoresis. If C + + H + = A - + O H - , H + must equal O H - as long as C + and A are zero; moreover [H +] and [ O H - ] must both be 10 -7 since their product is always 10 -14 e) Changes in the conformation of complex protonated molecules which alter a dissociation constant Complicated molecules such as proteins often have a number of very different conformations, so that the environments of their protonatable amino and carboxyl groups can change. These changes in the environments of the
23O atoms in turn change the ease or difficulty of protonation. For instance, suppose the following: HX ~
H + + X - ; pKa
=
5.0
now allow a conformation change in convert H X to HXo, HXo ~
H + + Xo; pKa
=
4.0
Thus, if we started with an equimolar mix of H X and K + X - , the pH would be 5.0 but if the conformation change occurred we would have almost equimolar HX o and K ÷ X o , and the pH would approximate 4.0. Note that there need be no import or export of H ÷ associated with the change in pH, the additional H + coming from the protonated acid group. This kind of reaction must play a very important role in biological systems, if only because proteins are so numerous and so subject to rearrangements. The problem is complicated by the fact that protonation of proteins is itself one of the major factors causing conformational changes.
f) Equilibrium of pH across membranes Many phenomena associated with photophosphorylation involve reversible pH changes. As will be discussed more fully below, electron transport frequently causes the medium suspending lamellar vesicles to become more alkaline while at the same time the inner parts of the system become more acid. Within a few seconds of the beginning of electron transport a steadystate ion redistribution is achieved, presumably because the influx and egress of ions balance. When electron transport ceases, the pH of the medium falls again and equilibrium is restored at the original pH and with approximately the original distribution of ions. Let us now consider the ion fluxes associated with re-equilibrium since these fluxes are also among the many processes changing pH. In the next section we will consider how transmembrane redox reactions might lead to pH changes. To illustrate the principles involved in pH equilibration consider a vesicle bounded by a membrane, with inner and outer aqueous phases. Now for simplicity consider the situation where only KC1 and water are present, the concentrations of K ÷ and C1- being identical on both sides of the membrane. The pH will then be exactly 7.0 on both sides and a stable equilibrium will exist. Now suppose that a disequilibrium is introduced, say by adding some K + outside then in the form of K O H or by removing some CI- outside with silver oxide. In either case, the pH outside will rise relative to the pH inside and equilibration processes can begin. But the disequilibrium tending
231 to dissipate, although it involves differences in H + concentration, is not caused by changes in H + and cannot be dispelled by movements of H + ; it resides in differences between K + and C1-. Ions of C1- or K + must move to restore transmembrane equilibrium and equilibrium can only be restored when the concentrations of K + are identical on the two sides of the membrane and when the concentration of C1 are identical on the two sides of the membrane. To achieve that end several ionic species must move, with kinetics too complex to consider here. Many calculations have been made in terms of the permeabilities of chloroplast membranes to hydrogen ions in order to account for "basal" electron transport rates and in order to account for the decay constants of the reversible change in the medium pH. However, in view of the considerations above, it is not evident that permeability to hydrogen ions need be a factor limiting the pH equilibration. In fact, if the membranes in the model system described here were infinitely permeable to hydrogen ions but utterly impermeable to K + and C1 , the pH would never equilibrate; even with infinite permeability to H + , there would be no "proton leakage". The first few hydrogen ions traversing the membrane would build up an electric potential just sufficient to prevent the flow of any more H + (which flow would not of itself change pH anyway). A membrane infinitely permeable to K + but utterly impermeable to other ions would never cause the pH to equilibrate either, because the first few K + ions moving would also build up a membrane potential just sufficient to prevent the flow of more K + . Clearly transmembrane ion equilibrations must always involve the movement of at least two ionic species, in this case perhaps the exchange of K + and H +, (although we must not overlook the role of magnesium ions and the binding of M f + to proteins in any consideration of actual chloroplast reactions (Dilley and Vernon 1965)). Furthermore, permeation of the membrane by the least permeant of the ions will always be rate-determining in the process restoring pH equilibrium. Attributing transmembrane pH equilibration and non-phosphorylating electron transport solely to the permeability of thylakoid membranes to hydrogen ions involves an assumption, the implausible assumption that the membranes are orders of magnitude more permeable to the counterions moving (including M f + ) than they are to hydrogen ions.
g) Changes in pH associated with transmembrane redox reactions: biological "proton pumps" Peter Mitchell's great contribution to our understanding of photophosphorylation and oxidative phosphorylation lies in recognition of the fact that electron transport can occur across membranes in such a way as to cause ionic disequilibria and pH changes. As an example he proposed a
232 transmembrane "proton conducting loop composed of a hydrogen carrier R/RH 2and an electron carrier loop M/M E+", both traversing the membrane and working together (Mitchell 1966). According to this hypothesis, illustrated in Fig. l, the reduced form of the hydrogen carrier RH 2 is oxidized on one side of the membrane by the oxidized form of the electron carrier M 2÷. The reaction depicted in Fig. 1, RH 2 + M z+ ~
R + M + 2H +,
releases two hydrogen ions on that side of the membrane. The oxidized hydrogen carrier (R) and the reduced electron carrier (M) then return to the other side of the membrane where they are regenerated as RH2 and M 2÷. This mechanism and analogous mechanisms have become very popular, not only to account for ATP synthesis but also to account for many active transport systems, all thought to be driven by the above described "proton pump". Unfortunately describing the process as a "proton pump" implies primacy of hydrogen ion movements when no such movements are actually postulated in the Mitchell model. In the model an electron acceptor oxidizes a hydrogen donor so that H ÷ replaces one of the positive charges on a PHASE L
PHASE R
MEMBRANE
f
s.
2H+,
Minireview
Active transport, ion movements, and pH changes
I. The chemistry of pH changes N O R M A N E. G O O D Department of Botany and Plant Pathology, Michigan State University, East Lansing, MI 48824, U.S.A. Received 1 March 1988; accepted in revised form 15 June 1988 Abstract. The transport of substances across cell membranes may be the most fundamental activity of living things. When the substance transported is any ion there can be a change in the concentration of hydrogen ions on the two sides of the membrane. These hydrogen ion concentration changes are not caused by fluxes of hydrogen ions although fluxes of hydrogen ions may sometimes be involved. The reason for the apparent contradiction is quite simple. All aqueous systems are subject to two constraints: (1) to maintain the charge balance, the sum of the cationic charges must equal the sum of the anionic charges and (2) the product of the molar concentration of H + and the molar concentration of O H - , established and maintained by the association and the dissociation of water, remains always at 10 14. As a consequence the concentrations of H + and OH are determined uniquely by differences between the concentrations of the other cations and anions, with [H +] and [OH-] being dependent variables. Hydrogen ions and hydroxyl ions can be produced or consumed in local reactions whereas any strong ions such as Cl-, Mg 2+, or K + can be neither produced nor consumed in biological reactions. Further consequences of these truisms are outlined here in terms of the chemistry of the kinds of reactions which can lead to pH changes.
Introduction
The assumption that changes in pH can be equated to hydrogen ion fluxes can lead to serious misconceptions. This paper details the nature of the various kinds of reactions which really do change pH, especially in biological systems. The next paper then expands on implications of the arguments regarding possible mechanisms of ATP synthesis and reexamines experimental data. The counter-intuitive contention that pH changes cannot be the direct consequence of hydrogen ion movements is based on solid chemical dicta. All aqueous systems are subject to two constraints. First, in order to maintain essential electrical neutrality in the bulk of the liquid C÷ + H÷ =
A - + OH
226 where C + represents all of the cationic charges other than H +, and A represents all of the anionic charges other than O H - . The distinction between the pair C + and A - on the one hand and the pair H + and O H - on the other hand is profoundly important and it is regrettable that our language has not given us terms to express the difference. In general, C + and A- are independent variables. In contrast H + and O H - are dependent variables. That is because H + is involved in a host of reversible protonation reactions, not least the protonation of O H - to make H20, carried out in a vast 55 M proton source. The second constraint on aqueous media is that the product of the molar concentration of O H - and of the molar concentration of H + remains constant at 10 -14, with the reaction H + + O H - ~ H20 proceeding in such a way as to assure this constancy. Thus the charge balance equation dictates that an excess of C + is tantamount to an excess of O H - while the water dissociation equation dictates that an increase in O H - is associated with a decrease in H + . With these constraints in mind, let us consider the meaning o f p H changes and how one can go about changing pH. Only when the diversity of the mechanisms which alter the concentration of H + has been appreciated can we begin to seek the real meaning of the Mitchell-Jagendorf electron transport-dependent pH changes in suspensions of mitochondria or chloroplast lamellae, the pH changes commonly but perhaps inappropriately attributed to hydrogen ion "pumps". How differences in hydrogen ion activities arise in aqueous media In preparing the following list I have attempted to cover most types of reactions although I cannot guarantee that the list is complete. Some processes, such as the oxidation or reduction of metal electrodes by an electric current, are unlikely to have much relevance to biological systems but most of the processes of pH change described below probably do have biological counterparts. A cautionary note may be needed here. The equation C + + H + = A- + O H - is exact except near the surfaces of high charge density. The charged membranes of most biological systems are an example but probably do not have sufficient electrical capacitance to account for significant deviations from this crucial equation in the bulk of the surrounding aqueous phase. In any event, the measurements of pH that we make are almost always at a great distance from such charged membranes. Therefore the pH values measured by electrodes in biological suspensions reflect the equation exactly.
227
a) Addition of HCI to K+ X - (titration o f a buffer by a strong acid) Reaction: H + + C1- + K + + X -
~
K + + C1- + HX
In this case the concentration of H + is determined by the equilibrium constant of the association-dissociation reaction H + + X - ~- HX and therefore [H + ] = Kd [HX]/[X- ] where K~ is the dissociation constant. Thus changes in [H + ] are complex functions of the amount HC1 (not H + ] added and the amount of K + X - originally present. It should be noted that it is the input of the chloride ion in HC1 just as much as it is the input of hydrogen ion that changes the pH. It should also be noted that the relationship between increases in [H + ] and the amount of HC1 added is ever-changing as K + X - is titrated. Obviously H X can represent a large reservoir of hydrogen ions available to replace those disappearing for any reason. When we are titrating a base such as K + O H - or a buffering salt such as K + X - , neither the change in [H +] nor the energy stored in the accumulation of H + is proportional to the flux of "protons".
b) Introduction of H + through membranes exclusively permeable to hydrogen ions. p H meters and the glass electrode It is interesting to consider the working of a pH meter. The glass electrode depends on the fact that the special glass membrane is, in effect, appreciably permeable to hydrogen ions but almost impermeable to all other ions. Therefore hydrogen ions, unaccompanied by counter ions, tend to diffuse across the membrane down any activity gradient which may exist and in this case we are indeed dealing with a phenomenon equivalent to a flux of protons. The result is the development of a transmembrane electric charge just sufficient to stop net diffusion of H + . That is to say, the work capacity of the charge on the extra electrons on the negative side of the membrane is just equal and opposite to the work potential of the hydrogen ions "expanding" from the higher concentration to the lower concentration. But we cannot measure the work potential of electrons without using some of them to do work, albeit a very small number sometimes. How then can we use hydrogen ions diffusing down a concentration gradient if these same hydrogen ions cannot change their own concentration by moving, as our charge-balance equation implies? (Clearly we cannot hope to measure the work done in dissipating a concentration difference if the difference is not being dissipated.) This problem is only a problem if we insist on looking at a partial reaction. When we consider the other reactions involved in completing the circuit of the meter, we find that the flux of hydrogen ions is
228 indeed changing the pH but only because of the electrode reactions which produce and consume silver ions. As the uncompensated hydrogen ions flow down their concentration gradient across the glass barrier, silver ions are reduced to silver metal on the "downhill" side and silver metal is oxidized to silver ions on the "uphill" side. The production of silver ions from silver metal on the "uphill" side ties up an amount of chloride exactly balancing the amount of hydrogen ion leaving that side, while the reduction of silver ions on the "downhill" side releases an amount of chloride exactly balancing the number of H ÷ ions arriving from the "uphill" side. Thus changes in hydrogen ion concentration do occur, changes driven by the hydrogen ion concentration difference. But these changes in hydrogen ion concentration are balanced by changes in chloride ion (driven it is true by H + activity differences) and therefore C + and H + continues to equal A - + O H - . I have belaboured this instance because it seems at first glance to refute my contention that moving hydrogen ions without moving (or creating) other ions cannot change the pH. However this case is simply an example of the errors we fall into when we confuse partial reactions with real-world overall reactions. Partial reactions "float" with no reference point, they cannot be balanced, and by themselves they have no meaningful energy relations. Here the flux of hydrogen ions drives the change in pH by driving the reduction of silver ions and the release of chloride ions. Therefore the charge balance equation is in no way violated. This example, while instructive, probably has little relevance to biology because the parts of the reaction which involve the oxidation and reduction of silver and precipitation of almost insoluble silver chloride are in no way biological.
c) The production of cations and anions by oxidation and reduction: electron donors and hydrogen donors If we oxidize one metallic ion by another metallic ion, there is no change in the balance of C + and A and therefore no change in pH is to be expected. However if we oxidize a metal to a cation, or oxidize a cation to the next higher valence state without producing a compensating anion, there is an increase in the number of cationic charges and a corresponding drop in the concentration of H + . By the same reasoning, reductions such as the conversion of chlorine gas to chloride ion can cause an increase in the concentration of H + . Let it be noted that the production or consumption of H + in these instances come entirely from the dissociation or association of water or of protonated acids in a reaction which balances the positive and negative charges in the medium and need not be due to hydrogen ion or hydroxyl ion movement.
229 If we oxidize a hydrogen donor H2 X to non-ionic products there is, of course, no pH change. If, however, we oxidize H2X with ferric salts: 2(Fe 3+ + 3A 3-) + H2X --+ 2(Fe 2+ + 2A ) + X + 2H + + 2Athere is a net decrease in the number of cationic charges and the pH falls. In the latter reaction we usually think of the production of hydrogen ions as coming from the oxidation of the hydrogen atoms in H2X and at first thought it seems logical to attribute the new hydrogen ions to the oxidation of H to H + , but in thinking thus we are considering a half-reaction which is incapable of changing pH by itself. In fact, the reason for the acidification is better described in terms of the decrease in C + which accompanies the disappearance of cationic charges when Fe 3+ becomes Fe 2+ . When we fit the loss of a cationic charge into the equation C + + H + = A - + O H - it is easier to see the real nature of the pH change. This distinction may be important because much has been made of the undoubted importance of "hydrogen donors" and so called proton pumps in phosphorylationcoupled electron transport.
d) Electrophoresis If one passes an electric current through an aqueous medium, the current is carried entirely by the charges on the dissolved ions when these ions migrate in the electric field. Thus cations, including H + , migrate toward the negative electrode (the cathode) and anions, including O H - , migrate toward the positive electrode (the anode). Since the C + ions accumulate toward the negative pole and the A - ions accumulate toward the positive pole, one would expect the concentration of hydrogen ions to decrease toward the negative pole even though H + is migrating into that region. That is exactly what happens. The situation is complicated, however, by the electrode reactions which introduce or remove ions. Nevertheless it is clear that the migrations of H + and O H - have no effect on the pH of water since pure water remains at pH 7.0 forever, no matter how long it is subjected to electrophoresis. If C + + H + = A - + O H - , H + must equal O H - as long as C + and A are zero; moreover [H +] and [ O H - ] must both be 10 -7 since their product is always 10 -14 e) Changes in the conformation of complex protonated molecules which alter a dissociation constant Complicated molecules such as proteins often have a number of very different conformations, so that the environments of their protonatable amino and carboxyl groups can change. These changes in the environments of the
23O atoms in turn change the ease or difficulty of protonation. For instance, suppose the following: HX ~
H + + X - ; pKa
=
5.0
now allow a conformation change in convert H X to HXo, HXo ~
H + + Xo; pKa
=
4.0
Thus, if we started with an equimolar mix of H X and K + X - , the pH would be 5.0 but if the conformation change occurred we would have almost equimolar HX o and K ÷ X o , and the pH would approximate 4.0. Note that there need be no import or export of H ÷ associated with the change in pH, the additional H + coming from the protonated acid group. This kind of reaction must play a very important role in biological systems, if only because proteins are so numerous and so subject to rearrangements. The problem is complicated by the fact that protonation of proteins is itself one of the major factors causing conformational changes.
f) Equilibrium of pH across membranes Many phenomena associated with photophosphorylation involve reversible pH changes. As will be discussed more fully below, electron transport frequently causes the medium suspending lamellar vesicles to become more alkaline while at the same time the inner parts of the system become more acid. Within a few seconds of the beginning of electron transport a steadystate ion redistribution is achieved, presumably because the influx and egress of ions balance. When electron transport ceases, the pH of the medium falls again and equilibrium is restored at the original pH and with approximately the original distribution of ions. Let us now consider the ion fluxes associated with re-equilibrium since these fluxes are also among the many processes changing pH. In the next section we will consider how transmembrane redox reactions might lead to pH changes. To illustrate the principles involved in pH equilibration consider a vesicle bounded by a membrane, with inner and outer aqueous phases. Now for simplicity consider the situation where only KC1 and water are present, the concentrations of K ÷ and C1- being identical on both sides of the membrane. The pH will then be exactly 7.0 on both sides and a stable equilibrium will exist. Now suppose that a disequilibrium is introduced, say by adding some K + outside then in the form of K O H or by removing some CI- outside with silver oxide. In either case, the pH outside will rise relative to the pH inside and equilibration processes can begin. But the disequilibrium tending
231 to dissipate, although it involves differences in H + concentration, is not caused by changes in H + and cannot be dispelled by movements of H + ; it resides in differences between K + and C1-. Ions of C1- or K + must move to restore transmembrane equilibrium and equilibrium can only be restored when the concentrations of K + are identical on the two sides of the membrane and when the concentration of C1 are identical on the two sides of the membrane. To achieve that end several ionic species must move, with kinetics too complex to consider here. Many calculations have been made in terms of the permeabilities of chloroplast membranes to hydrogen ions in order to account for "basal" electron transport rates and in order to account for the decay constants of the reversible change in the medium pH. However, in view of the considerations above, it is not evident that permeability to hydrogen ions need be a factor limiting the pH equilibration. In fact, if the membranes in the model system described here were infinitely permeable to hydrogen ions but utterly impermeable to K + and C1 , the pH would never equilibrate; even with infinite permeability to H + , there would be no "proton leakage". The first few hydrogen ions traversing the membrane would build up an electric potential just sufficient to prevent the flow of any more H + (which flow would not of itself change pH anyway). A membrane infinitely permeable to K + but utterly impermeable to other ions would never cause the pH to equilibrate either, because the first few K + ions moving would also build up a membrane potential just sufficient to prevent the flow of more K + . Clearly transmembrane ion equilibrations must always involve the movement of at least two ionic species, in this case perhaps the exchange of K + and H +, (although we must not overlook the role of magnesium ions and the binding of M f + to proteins in any consideration of actual chloroplast reactions (Dilley and Vernon 1965)). Furthermore, permeation of the membrane by the least permeant of the ions will always be rate-determining in the process restoring pH equilibrium. Attributing transmembrane pH equilibration and non-phosphorylating electron transport solely to the permeability of thylakoid membranes to hydrogen ions involves an assumption, the implausible assumption that the membranes are orders of magnitude more permeable to the counterions moving (including M f + ) than they are to hydrogen ions.
g) Changes in pH associated with transmembrane redox reactions: biological "proton pumps" Peter Mitchell's great contribution to our understanding of photophosphorylation and oxidative phosphorylation lies in recognition of the fact that electron transport can occur across membranes in such a way as to cause ionic disequilibria and pH changes. As an example he proposed a
232 transmembrane "proton conducting loop composed of a hydrogen carrier R/RH 2and an electron carrier loop M/M E+", both traversing the membrane and working together (Mitchell 1966). According to this hypothesis, illustrated in Fig. l, the reduced form of the hydrogen carrier RH 2 is oxidized on one side of the membrane by the oxidized form of the electron carrier M 2÷. The reaction depicted in Fig. 1, RH 2 + M z+ ~
R + M + 2H +,
releases two hydrogen ions on that side of the membrane. The oxidized hydrogen carrier (R) and the reduced electron carrier (M) then return to the other side of the membrane where they are regenerated as RH2 and M 2÷. This mechanism and analogous mechanisms have become very popular, not only to account for ATP synthesis but also to account for many active transport systems, all thought to be driven by the above described "proton pump". Unfortunately describing the process as a "proton pump" implies primacy of hydrogen ion movements when no such movements are actually postulated in the Mitchell model. In the model an electron acceptor oxidizes a hydrogen donor so that H ÷ replaces one of the positive charges on a PHASE L
PHASE R
MEMBRANE
f
s.
2H+,
