Bicarbonate-Mediated CO2 Formation on Both Sides of Photosystem II.

The effect of bicarbonate (HCO3-) on photosystem II (PSII) activity was discovered in the 1950s, but only recently have its molecular mechanisms begun to be clarified. Two chemical mechanisms have been proposed. One is for the electron-donor side, in which mobile HCO3- enhances and possibly regulates water oxidation by acting as proton acceptor, after which it dissociates into CO2 and H2O. The other is for the electron-acceptor side, in which (i) reduction of the QA quinone leads to the release of HCO3- from its binding site on the non-heme iron and (ii) the Em potential of the QA/QA•- couple increases when HCO3- dissociates. This suggested a protective/regulatory role of HCO3- as it is known that increasing the Em of QA decreases the extent of back-reaction-associated photodamage. Here we demonstrate, using plant thylakoids, that time-resolved membrane-inlet mass spectrometry together with 13C isotope labeling of HCO3- allows donor- and acceptor-side formation of CO2 by PSII to be demonstrated and distinguished, which opens the door for future studies of the importance of both mechanisms under in vivo conditions.

Photosynthetic organisms need atmospheric carbon dioxide (CO2) as the terminal electron acceptor to store the captured energy of sunlight as energy-rich carbohydrates.[1,2] Cyanobacteria, algae, and higher plants also require CO2 in solution in the form of bicarbonate ions (HCO3–), for the optimal function of photosystem II (PSII), the enzyme that catalyzes light-induced reduction of quinone and oxidation of water to molecular oxygen and protons.[3−6] The discovery of the “bicarbonate effect” on PSII activity in 1958[7] triggered a long-running debate about its role(s).[8,9] Two sites of interaction of HCO3– with PSII have been considered: one on the electron-donor side of PSII, where water oxidation takes place, and the other on the electron-acceptor side, associated with quinone reduction (Figure A).

Figure 1

Redox-active cofactors and bicarbonate interactions in PSII. (A) Arrangement of redox-active cofactors and sites where bicarbonate (HCO3–) plays functional roles within PSII. One HCO3– molecule binds, in the dark, with high affinity to the Fe2+ between quinones QA and QB on the acceptor side of PSII, and this is resolved in the crystal structure;[13] mobile HCO3– molecules act as proton acceptors on the water-oxidizing side. Their locations are unknown but suggested to be in the proton exit channels[18] and thus have been placed there to illustrate that model. (B) Nonlabeling measuring conditions. Light-induced formation of CO2 from HCO3– may occur as 12CO2 on both sides of PSII. Thus, all released CO2 should be detected by TR-MIMS at m/z 44. (C) Labeling of the medium with H13CO3–. 13CO2 formation is expected (m/z 45) from light-induced water splitting leading to the protonation of mobile H13CO3– (Figure and Figure S1), while the release of slowly exchangeable (in darkness with a low CO2/HCO3– content) H12CO3– from its binding site at the Fe2+ would generate 12CO2 (m/z 44) (Figure S2). The structure of PSII was generated using Protein Data Bank entry 3ARC.[13]

Figure 2

Online TR-MIMS measurements of the production of O2 and CO2 by spinach thylakoids. Dark-adapted thylakoids (0.5 mg of Chl mL–1) were illuminated with 100 (panels A, C, and E) or 500 (panels B, D, and F) xenon flashes (at 2 Hz) in the absence (A– traces) or presence (A+ traces) of artificial electron acceptors {0.25 mM PPBQ and 0.5 mM K3[Fe(CN)6]}. Traces of O2 (A and B) and 12CO2 (C and D) evolution were measured simultaneously at m/z 32 and 44, respectively, after the addition of 1 mM NaH12CO3 and subsequent sample degassing for 20–30 min to reach a stable baseline. Traces of 13CO2 (E and F) evolution were obtained at m/z 45 after the addition of 1 mM NaH13CO3 and subsequent degassing of thylakoid preparations inside the MIMS cell for 20–30 min. The measurements were performed in MES medium (3 mM MES at pH 6.39) at 20 °C. The arrows indicate the start and end of the train of flashes. Zero levels are offset for the sake of clarity of presentation. In all panels, the means (solid traces) of two to three repeats and standard errors (shaded areas) are presented.

It is now clear that HCO3– is not a tightly bound component of the oxygen-evolving complex and its catalytic site, the Mn4CaO5 cluster.[10−13] However, it was proposed that mobile, easily exchangeable HCO3– ions can stimulate water splitting by accepting protons produced by water splitting.[14−17] This action of HCO3– was proven experimentally by the direct detection of light-induced formation of O2 and CO2 by PSII using time-resolved membrane-inlet mass spectrometry (TR-MIMS).[18] This study also showed that the O2 activity of PSII in HCO3–-depleted media is reversibly reduced by ∼20%, likely via both donor- and acceptor-side (see below) effects of HCO3–. Moreover, a recent PSII mutagenesis study supports the role of HCO3– in the proton-egress pathway.[19] These observations led to the proposal that mobile HCO3– may contribute to a feedback mechanism that may adapt the availability of electrons for CO2 reduction.[18]

On the acceptor side of PSII, HCO3– removal slows electron transfer through the quinone electron acceptors QA and QB.[20] This was subsequently rationalized by the discovery that HCO3– is a ligand of the non-heme iron (Fe2+) (Figure A).[21−23] It was also suggested that HCO3– ions facilitate the protonation of QB.[24−26]

Given its reported binding constant and its concentration in the stroma, one HCO3– was thought to be permanently bound to the Fe2+.[27,28] However, some earlier quantitative assays determined less than one HCO3– per PSII.[11,29,30] Our recent TR-MIMS study, carried out in the dark under air-saturated conditions, revealed exactly one HCO3– per PSII.[31] This unexpected variation was clarified recently, when it was shown that formation of QA•– results in a weakening of HCO3– binding, which can lead to a release of HCO3– that in turn shifts the Em of the QA/QA•– couple by +74 mV.[32] The release of HCO3– was evidenced by the typical slowing of electron transfer from QA•– to QB.[32] These results were interpreted as providing a mechanism by which the PSII activity is regulated by the ambient HCO3–, and hence by CO2, and by the redox state of QA.[32] In this case, the upshift in the Em of the QA/QA•– couple upon the release of HCO3– increases the energy gap between QA•– and pheophytin (Pheo)D1, disfavoring the back-reaction route that is known to generate the radical-pair recombination chlorophyll (Chl) triplet state[33] and hence singlet oxygen.[34] The acceptor-side redox tuning mechanism is thus not only regulatory but also protective.[32] The aspect of protection is crucial, as simply slowing PSII electron transfer when CO2, the terminal acceptor, is limiting is likely to result in an increased level of photodamage unless a protective mechanism is also triggered.

The acceptor-side regulation may be modulated by the binding of other carboxylic acids.[32] This proposal was supported by reports that acetate in the growth medium appears to displace HCO3– from the acceptor side of PSII in Chlamydomonas.[35] A similar effect was seen in a photorespiration mutant of Arabidopsis in which high concentrations of glycolate accumulated.[36] However, despite the observations mentioned above, no direct experimental evidence for light-induced evolution of CO2 from the non-heme Fe2+ has been reported.

In the study presented here, we refined the TR-MIMS experiments to probe selectively for the light-induced production of CO2 at the acceptor and donor sides of PSII (Figure B). For this, we employed H13CO3– labeling of the thylakoid suspension (Figure C) and modulated the oxidation state of the acceptor-side quinones by using varying concentrations of exogenous electron acceptors. The MIMS signals obtained clearly demonstrate that both the HCO3– bound at the Fe2+ on the PSII acceptor side and mobile HCO3– molecules protonated within PSII during water splitting contribute to light-induced CO2 formation by PSII.

Materials and Methods

Chemicals and Reagents

NaH13CO3 (99% 13C), NaH12CO3 (>99.7%), 2-phenyl-p-benzoquinone (PPBQ, >95%), and potassium ferricyanide, K3[Fe(CN)6] (>99.99%), were purchased from Sigma-Aldrich. All [12C/13C]bicarbonate stock solutions were prepared shortly before the experiments in deionized and filtered (Milli-Q) water depleted of inorganic carbon (C). Depletion of C in water was carried out as described earlier by intensive flushing with nitrogen in septum-sealed vials for 20–30 min.[16] To avoid contamination with atmospheric CO2, the C-depleted water was added to the weighed batches of NaH13CO3/NaH12CO3 inside a glovebox (OMNI-Lab System, VAC, Hawthorne, CA) under a N2 atmosphere. The resulting stock solutions of NaH13CO3 and NaH12CO3 (30 mM each) were kept in septum-sealed vials until they were used. PPBQ (50 mM) and K3[Fe(CN)6] (100 mM) stock solutions were freshly prepared in DMSO (>99.9%) and in Milli-Q water, respectively.

Sample Preparations

Isolated thylakoids were prepared from fresh leaves of Spinacia oleracea as described previously.[37,38] After being isolated, the thylakoids were frozen in liquid N2 in small aliquots (in sucrose buffer containing 400 mM sucrose, 5 mM CaCl2, 5 mM MgCl2, 15 mM NaCl, and 50 mM MES/NaOH adjusted to pH 6.0 and at [Chl] = ∼2 mg mL–1) and stored at −80 °C until they were used. Control rates of O2 evolution for our thylakoid preparations were ∼180–200 μmol of O2 (mg of Chl)−1 h–1 [as measured by a Clark-type electrode at 20 °C using continuous saturating light (∼1500 μmol of photons m–2 s–1) in the presence of 0.25 mM PPBQ and 0.5 mM K3[Fe(CN)6] as artificial electron acceptors]. Shortly before the measurements, the samples were thawed in the dark on ice, washed once in MES medium (containing 15 mM NaCl, 5 mM CaCl2, 5 mM MgCl2, and 3 mM MES/NaOH adjusted to pH 6.39), and diluted to the desired Chl concentrations (see below). This MES medium was used in all of the MIMS experiments described herein. A sufficiently low concentration of MES in this medium was needed to minimize the competition of MES with HCO3– in proton removal from the water-splitting site of PSII while still allowing it to buffer the pH of the medium.[18]

Online TR-MIMS Assays and H13CO3— Labeling

Our TR-MIMS setup[39,40] consisted of an isotope ratio mass spectrometer (Finnigan DELTAplusXP, Thermo, Bremen, Germany) connected via a cold trap (dry ice/ethanol) to a 150 μL in-house-constructed membrane-inlet cell described previously.[18] The sample volume of the cell was isolated from the vacuum (3 × 10–8 mbar) of the mass spectrometer by a gas-permeable silicon membrane (25 μm thick; type MEM-213, MemPro, Troy, NY) that was supported by a porous Teflon disc (Ø 1 cm; Small Parts Inc., Miami Lakes, FL). Thylakoids diluted with MES medium (see above) were injected into the MIMS cell to a final concentration of 0.5 mg of Chl mL–1. If not stated otherwise, the sample suspension also contained 0.25 mM PPBQ and 0.5 mM K3[Fe(CN)6] as an exogenous electron-acceptor system. All of these manipulations were performed under dim green light. During the assays, the MIMS cell was thermostated at 20 °C and the sample suspension was constantly stirred at high speed (1000 rpm) with a magnetic stir bar. After the sample suspension had been loaded into the MIMS cell, the samples were thoroughly degassed in the dark for ∼40 min by the vacuum pump of the mass spectrometer. After degassing, the bulk medium of the sample suspension was labeled with 13C by injection of 5 μL of NaH13CO3 into the MIMS cell to a final concentration of 1 mM. For control experiments, the same amount of NaH12CO3 was injected into the MIMS cell. All transfers and injections of bicarbonate solutions were performed with gastight Hamilton syringes that had been preflushed with nitrogen. After addition of the [13C/12C]bicarbonate solutions, the sample suspensions were incubated and stirred for approximately 20–30 min. This time was enough to reach stable and identical baseline values for 12CO2/13CO2 and to equilibrate the 13C label between the remaining C species in the aqueous fraction of the sample. Light-induced evolution of gases (O2 and 13CO2/12CO2) was monitored upon illumination of the thylakoids with a train of 100 or 500 short (∼5 μs full width at half-maximum) saturating flashes (2 Hz) given by a xenon flash lamp (model FX-4400, Excelitas Technologies Illumination, Inc., Salem, MA). O2 and nonlabeled CO2 evolved by thylakoids were detected simultaneously at m/z 32 and 44, respectively, with the same sensitivity of Faraday Cups (Figure B). In the 13C labeling experiments (Figure C), 13CO2 evolution was monitored at m/z 45 using the same Faraday Cup as for detection of 12CO2m/z 44 signals (Figures S1 and S2). This allowed us to obtain 13CO2-MIMS signals with the same sensitivity and selectivity as for the 12CO2-MIMS signals.

Results

Light-induced O2 [m/z 32 (Figure A,B)] and 12CO2 [m/z 44 (Figure C,D)] evolution by spinach thylakoids was monitored using TR-MIMS in nonlabeled bulk medium. The experiments were performed in the presence or absence (labeled A+ or A–, respectively) of the artificial electron-acceptor system (0.25 mM PPBQ and 0.5 mM potassium ferricyanide). Gas evolution was monitored upon illumination with 100 (Figure A,C) and 500 saturating xenon flashes (Figure B,D).

In the absence of exogenous electron acceptors, the total O2 yield (gray traces in panels A and B of Figure ) was equally small irrespective of the application of 100 or 500 flashes. This is simply a reflection of the limited number of enzyme turnovers possible with the small intrinsic electron-acceptor pool.

As expected, the level of O2 production was much higher in the presence of electron acceptors (black lines in panels A and B of Figure ). However, Figure B (black trace) shows that the O2 evolution reached a plateau at ∼300 flashes and decreased thereafter. Because the concentration of the artificial electron acceptors was adjusted to be just enough for 500 flashes (for details, see the Supporting Information), we propose that the decline in O2 production is due to the limited acceptor pool size and possibly some photodamage.

For 100 flashes in the presence of the electron acceptors (Figure A, black trace), the trace looks different compared to the trace obtained in the absence of the acceptors (Figure A, gray trace). Here, the O2 yield continued to increase for ∼7 s in the dark after the flash train ended, reaching a maximum before falling back to the original level ∼200 s after the flash train. The increase in the rate of release of O2 immediately after the end of the flash train reflects the instrument response time, which is also observed for all traces as an apparent lag in O2 evolution at the onset of illumination. The decrease in the O2 concentration after the flash train is due to the pervaporation of the dissolved O2 through the silicon membrane into the high vacuum of the mass spectrometer.[40]

Panels C and D of Figure show the kinetics of 12CO2 evolution when monitored simultaneously with O2 concentration. This experiment is sensitive to 12CO2 released by both mechanisms: (i) from the donor side, where mobile H12CO3– reacts with protons generated by water oxidation to form CO2 and H2O,[18] and (ii) from the acceptor side, where the bound H12CO3– is released from its site at the non-heme Fe2+ during illumination,[31] after which it may also dissociate into CO2 and H2O after protonation. In contrast to O2 evolution (Figure A,B), the 12CO2 evolution traces (Figure C,D) were much less dependent on the presence of the exogenous electron acceptors; in fact, a similar quantity of CO2 was formed with and without the added electron-acceptor system. Nevertheless, the presence of the electron acceptors did result in CO2 release occurring more rapidly than in the absence of acceptors, and this was the case whether 100 or 500 flashes were used. This more rapid increase in the level of CO2 when electron acceptors were present (Figure C,D, dark blue traces) appears to correspond to the rapid increase in the level of O2 evolution (Figure A,B, black traces), and this part of CO2 release can thus be attributed to the donor-side mechanism (see also below).

The kinetics for 12CO2 evolution in the absence of an exogenous electron acceptors (Figure C,D, bright blue traces) is noticeably slower than the diffusional step for O2 detection (Figure A, black trace), and we thus associate this slower phase with CO2 formation at the acceptor side. The slow kinetics for the onset of CO2 production could reflect (i) the loss of long-lived QA•– under these conditions due to charge recombination with the donor side, (ii) the slow release of HCO3– when QA•– is present,[32] or (iii) the slow conversion of the released HCO3– to CO2 at this pH (see below). In comparison to O2 evolution, the CO2 release measured in the absence of electron acceptors also continued for much longer after the end of the flash train (compare panels C and D of Figure with panels A and B of Figure ). In Figure C, after the 100 flashes were given, at 50 s during the time course, the amount of CO2 continues to increase in the dark for a further ∼50 s. The lengthy phase of CO2 emission in the dark after the cessation of illumination could reflect the second and (more likely) the third of the three options discussed above.

When 500 flashes were given (Figure D), the amount of CO2 released in the sample lacking the exogenous electron acceptors increased with flash number, but the slope decreases toward the end of the flash train. Given the absence of any CO2 originating from the donor side at this time under these conditions, due to the lack of O2 production, the 12CO2 released (Figure D, bright blue trace) can be assigned as arising purely from the gradual release of the bound acceptor-side bicarbonate, which appears to be nearly complete at this time point.

The two kinetic components in the data in panels C and D of Figure thus demonstrate the presence of two sources for CO2 formation, which we assign to the donor side and acceptor side, respectively. This is in line with our previous MIMS measurements showing CO2 production on the donor side, where HCO3– was thought to act as a proton acceptor during water oxidation, and our recent electrochemical and fluorescence data providing strong evidence that indicated the release of HCO3– from the non-heme iron when QA•– was present for a prolonged period of time.[18,32] The data in Figures and 3 also show that acceptor-side CO2 formation is dominant under the experimental conditions used, which entail a low availability of HCO3– in the medium, thus largely suppressing CO2 production at the donor side. The low HCO3– concentration additionally promotes the release of HCO3– from the acceptor side.

Figure 3

Resulting TR-MIMS CO2 signals upon subtraction of m/z 45 curves (Figure E,F) from corresponding m/z 44 traces (Figure C,D) obtained after illumination with (A) 100 and (B) 500 Xe flashes. For other conditions, see Figure and Materials and Methods.

To confirm this assignment further and to selectively record the evolution of CO2 from the electron-donor side, we added H13CO3– to the bulk medium. This approach is based on the assumption that H12CO3– bound to the non-heme iron does not significantly exchange with H13CO3– within the 20–30 min incubation time after its addition to the medium. This assumption appears to be reasonable given that (i) HCO3– binds in the dark tight enough to the non-heme iron to allow only slow dissociation into the medium, even at low external HCO3– concentrations, and (ii) the binding to empty sites at the non-heme iron competes with the consumption by the mass spectrometer.[31,32] In contrast, HCO3– acting as a proton acceptor at the donor side can be easily removed and thus exchangeable.[18]

Light-induced 13CO2, monitored at m/z 45 (Figure E,F), occurred with kinetics and relative amplitudes that closely resembled those of O2 evolution under the various conditions, albeit at much smaller absolute amplitudes. While the present setup does not allow for calibration of the signals, we estimate the 13CO2/O2 ratio to be on the order of 0.3–1.3% on the basis of the relative signal amplitudes and amplification factors (this estimate neglects differences in ionization efficiencies).[18] The excellent kinetic match between the O2 and 13CO2 release data thus provides strong support for the assumptions made above and the assignment of the fast phase of CO2 evolution observed in panels C and D of Figure in the presence of electron acceptors to mobile bicarbonate acting as a proton acceptor during water splitting.[18]

Taking into account the fact that m/z 44 and 45 signals (panels C and D and panels E and F, respectively, of Figure ) were recorded with identical sensitivity and selectivity (because both were monitored by the same Faraday cup), it is clear that the donor-side CO2 release is slower than acceptor-side release, as already suggested above by the unlabeled data (Figure C,D). To obtain pure kinetics for acceptor-side CO2 formation, we subtracted the corresponding curves in Figure C–F (Figure ). In principle, this difference signal may reflect one HCO3– per PSII reaction center after 500 flashes where the acceptor-side release appears to be nearly complete (see above). However, under the conditions of the TR-MIMS experiment with the degassing procedures used, we estimate the occupancy of bicarbonate at the non-heme iron site to be 0.3–0.5 HCO3– per PSII.[11]

Surprisingly, acceptor-side CO2 formation is nearly identical for samples with and without the acceptor during the flash train, as seen from the same initial slopes for CO2 formation in Figure . However, after the train of 100 flashes, there is marked difference comparing the traces obtained with and without the acceptor. This difference is less marked in the data from the experiments using 500 flashes, reflecting the nearly complete release of HCO3– during the flash train under both experimental conditions.

Discussion

The results of this study provide strong evidence of two separate mechanisms of light-induced CO2 formation by PSII, one involving mobile (bulk) HCO3– and one firmly bound (slowly exchangeable) HCO3–. We assign the slower kinetics to CO2 formation at the electron-acceptor side of PSII, where HCO3– bound to the Fe2+ is released into the medium under illumination, and in the absence of added electron acceptors also for extended times in the dark after the end of the illumination. While some open questions remain (see below), the data overall support the idea that this CO2 formation is due to the reduction of the endogenous electron acceptors of PSII, in agreement with the earlier correlation of formation of QA•– with bicarbonate release.[32] We assign the faster phase of CO2 formation to the electron-donor side, where mobile HCO3– accepts protons produced by the Mn4CaO5 cluster during light-induced water splitting, presumably within the channels of PSII, and then decomposes into H2O and CO2.[18]

The donor-side release of CO2 correlates closely with O2 evolution in terms of both its rate and its relative extent. The absolute extent of CO2 release is small (≈1% of the O2 yield) in this work because of the low level of free bicarbonate left in the experimental buffer after the extensive degassing that was required to allow the experiments to be performed. The CO2 yield is further reduced by the competition for protons between HCO3– and MES molecules in the medium.[18] Experiments in intact cells or chloroplasts will be required to determine the magnitude and functional importance of the donor-side CO2 evolution under in vivo conditions.

The CO2 attributed to the HCO3– released from the acceptor side of PSII was observed when thylakoids were illuminated, supporting the hypothesis that reduced quinones decrease the binding affinity of HCO3– at the non-heme iron. However, we did not find the expected clear difference during the flash train with and without the added electron-acceptor mix. However, a clear difference was seen in the dark after the series of 100 flashes. In the absence of the electron-acceptor system, the CO2 concentration continued to increase in the dark after the flash train, while when the acceptor system was present, the post-illumination increase in the level of CO2 was much less marked. When the acceptor system was absent, 100 flashes was more than enough to reduce QA (by one electron), QB (by two electrons), and the plastoquinone (PQ) pool (a capacity of 14 electrons when taken as seven PQs). The significant quantity of CO2 released in the dark after the flash train could reflect (i) the presence of QA•–, which is known to trigger HCO3– release,[32] and (ii) the presence of other reduced forms of quinone, QB•–/QBH2, and the fully reduced PQH2 pool, all of which can equilibrate with QA•– and could result in bicarbonate dissociation. It is also possible that bicarbonate binding may be weakened by the physical presence of reduced quinones in the QB site (i.e., QB•– or QBH2). We note that an additional complexity arises from the fact that both PPBQ and K3[Fe(CN)6] can oxidize the non-heme Fe2+ to Fe3+ in some of the centers. While K3[Fe(CN)6] may do so in a fraction of the centers during the dark time before the flash train, PPBQ is known to oxidize the non-heme iron on odd-numbered flashes by a mechanism known as reductant-induced oxidation.[41] However, the presence of the K3[Fe(CN)6] is likely to compete for the semiquinone that oxidizes the Fe2+. While we might expect Fe3+ to bind bicarbonate more tightly than Fe2+, this is likely affected by the proton release that accompanies the oxidation of the iron, and a bidentate/monodentate ligation difference could also affect the binding properties upon the Fe2+ to Fe3+ oxidation step.[22] This aspect will require more detailed attention in the future.

The absence of a post-100 flash CO2 release with an added electron acceptor could be due to the remaining PPBQ/K3[Fe(CN)6] acceptor system removing any electrons from the intrinsic acceptors. This would mean that the binding of the bicarbonate would revert to being tight due to the absence of reduced intrinsic quinone acceptors.

The unexpected observation that bicarbonate (in the form of CO2) is also released upon illumination when the exogenous electron acceptors are present could at least partially be due to the low concentration of free bicarbonate ions in solution under the conditions needed for the experiments. The low concentration of bicarbonate would decrease the rate at which bicarbonate would bind to the non-heme iron. Consequently, the rate of bicarbonate dissociation could become competitive with the electron transfer kinetics of the forward reaction, QA•– to QB, and/or the back-reactions, QB•– to QA and QBH2 to QA. This view is supported by MIMS experiments at a higher level of residual HCO3– in the medium (see Figure S3), in which acceptor-side CO2 formation was not observed after the end of the flash train in samples containing no added acceptor. Further experiments combining TR-MIMS with simultaneous fluorescence measurements are required to fully resolve the details of the light-induced release of HCO3– from the electron-acceptor side.

Conclusions

The direct measurement of light-induced CO2 formation at the acceptor side of PSII reported here using TR-MIMS complements the less direct demonstrations of the release of HCO3– from the non-heme Fe reported previously, i.e., (i) the light-induced shift in the redox potential of the QA/QA•– couple to a value typical of bicarbonate-depleted PSII and (ii) the light-induced shift in the kinetics of QA•– oxidation to a slower rate typical of bicarbonate-depleted PSII and its recovery by re-addition of bicarbonate.[32] In addition, our data demonstrate a second kinetic phase of release of CO2 that closely correlates with O2 evolution and thus is assigned to HCO3– acting as a proton acceptor during water oxidation.

The donor- and acceptor-side effects may both be rationalized in terms of regulatory mechanisms in which the terminal electron acceptor, CO2, in the form of HCO3–, influences PSII function (Figure ).[18,32] Decreased concentrations of CO2 will result in decreased concentrations of HCO3– ions, which will have an immediate effect on the donor-side function if no other proton acceptors are present. On the other hand, this will initially have little effect on the acceptor side as the bicarbonate is strongly bound under normal functional conditions. However, once the electron-acceptor side becomes reduced, as it would when the PQ pool is reduced when CO2 fixation becomes limiting (Figure ), the affinity for bicarbonate would decrease, leading to its release. The release of HCO3– slows QBH2/PQ exchange and also results in a positive shift in the Em of the QA/QA•– couple, and this increases the energy gap between QA and PheoD1, disfavoring the Chl triplet-mediated, singlet O2-generating, back-reaction route that gives rise to photodamage.[32,34]

Figure 4

Possible regulation sites of oxygenic photosynthesis by inorganic carbon. Abbreviations of the components involved in electron transfer: PQ, plastoquinone; PQH2, plastoquinol; PC, plastocyanin; Fd, ferredoxin; FNR, ferredoxin-NADP reductase.

Although indications have been reported that bicarbonate effects are present in plants and cells of green algae,[35,36,42−44] further experiments with simultaneous TR-MIMS and fluorescence assays are required to determine the mechanistic aspects and the importance of HCO3–/CO2-mediated feedback regulation at the donor and acceptor sides of PSII in vivo.