aqueous hydride transfer 
 is an essential process in enzymatic catalysis 
 in nature , emerging fuel synthesis schemes for alternative 
 energy , and biphasic catalysis in chemical industry . hydrogenase 
 enzymes , for example , can produce ( or split ) hydrogen 
 with exceptional rates via transition metal hydride intermediates . long - sought synthetic mimics that catalyze hydrogen 
 evolution in aqueous solutions remain a pressing challenge . further , in the petrochemical industry , an aqueous - phase rh hydride 
 produces 800  000 tons / year n - butyraldehyde 
 for plastics . understanding and predicting 
 the reactivity of transition metal hydrides in water will continue 
 to gain importance as alternative feedstocks such as biomass and co2 , which can be reduced using hydride complexes , become increasingly 
 prevalent in the synthesis of chemicals and fuels . thermochemical studies of metal hydrides provide 
 a foundation for 
 rational design of catalysts and for mechanistic studies of 2e proton - coupled electron transfer ( pcet ) reactions . in acetonitrile , dubois and rakowski dubois pioneered the determination 
 of hydricity and its use as a powerful tool for reaction development 
 in organic solvents . in one stunning example , a decade of thermochemistry - guided 
 catalyst design culminated in a hydrogen evolution electrocatalyst 
 that can operate faster than hydrogenase enzymes . scheme 
 illustrating the hydricity of reference complex [ cp*ir(bpy - coo)(h ) ] ( 2h ) and thermochemical cycles that establish 
 aqueous hydricity of ir and ru hydrides . in light of the benefits that thermochemical understanding 
 could 
 have in the development of hydride - mediated catalysis in water , the 
 aqueous hydricity of metal hydrides has been , until recently , surprisingly 
 unexplored . creutz and chou s 
 seminal early efforts relied on experimentally challenging approach - to - equilibrium 
 kinetics under co2 . recently , the groups 
 of yang and berben each reported the hydricity of one metal hydride 
 based on the thermodynamics of h2 cleavage reactions . our interest in a family of cp*ir - based catalysts ( cp * is pentamethylcyclopentadienyl ) 
 motivated us to develop a general and expedient method for aqueous 
 hydricity determination in this series of complexes . preliminary 
 studies on the parent complex [ cp*ir(bpy)(h ) ] ( 1h ; bpy is 2,2-bipyridine ) were stymied by 
 the water - insolubility of the conjugate base cp*ir(bpy ) ( 1 ) , so we charted a course utilizing 
 carboxylate groups on the ligand to confer water solubility . pka thermochemical cycle in water ( figure 1a ) and then map the relative 
 hydricity of other complexes based on hydride transfer equilibria 
 ( figure 1b ) . the potential - pka thermochemical cycle has been used extensively 
 in acetonitrile , but has not been successfully 
 applied in water . the strategy depicted in figure 1 has enabled the construction 
 of an extensive , self - consistent 
 aqueous hydricity scale . the broad range of ir and ru hydricity values 
 reveals how the polar , protic aqueous environment impacts hydride 
 transfer thermodynamics . substantial shifts in the hydricity values 
 are observed relative to acetonitrile , with electronic changes to 
 the supporting ligand correlated strongly to the hammett parameter 
 p. a dramatic impact of water was also observed 
 in the primary coordination sphere : a variety of suitable ligands 
 present in aqueous media can bind the ir or ru centers after hydride 
 transfer , shifting the effective hydricity substantially . describing 
 the complexities of hydride transfer in water allows interpretation 
 of previously reported catalytic reactions and predictions that can 
 guide improvements in the hydrogenation of carboxylic acids , the disproportionation of formic acid to methanol , and other metal hydride - mediated reactions such 
 as h2 evolution and co2 reduction . the first 
  reference  hydride investigated was [ cp*ir(bpy - coo)(h ) ] ( 2h ; bpy - x = 4,4-x - bpy ) , with 
 carboxylate groups installed on the bipyridine ligand to confer good 
 water solubility over a wide ph range , independent of metal ligation 
 or oxidation state . the reduction potential 
 of [ cp*ir(bpy - cooh)(cl)][cl ] ( 2cl ) was initially assessed 
 using cyclic voltammetry ( cv ) in 1 m naoh . 
 under these conditions , the chloride is displaced by hydroxide to 
 form [ cp*ir(bpy - coo)(oh ) ] ( 2oh ) based 
 on nmr and ms data , and a 2e reduction forms the 
 freely diffusing species [ cp*ir(bpy - coo ) ] ( 2 ) . unfortunately , a large peak - to - peak separation was observed 
 between the reduction of 2oh and the oxidation of 2 ( figure s4 ) . this electrochemical 
 irreversibility , attributed to slow electron transfer or slow ligand 
 dissociation , prevented the use of cv to determine e1/2 . ( a ) spectral changes 
 of a ph 14 solution of [ cp*ir(bpy - coo ) ( oh ) ] ( 2oh ) as the solution potential is decreased 
 by electrolysis to form [ cp*ir(bpy - coo ) ] ( 2 ) . ( b ) absorbance at 620 nm stepping in the negative potential 
 direction ( red dots ) , the positive potential direction ( blue dots ) , 
 and the fit to the nernst equation ( dot - dashed line ) giving e = 0.60 v. the lack 
 of hysteresis indicates that equilibrium was established . ( c ) absorbance 
 at 570 nm of a ph titration of [ cp*ir(bpy - coo)(h ) ] ( 2h ) forming 2 ( red dots ) and the fit 
 to the henderson  hasselbalch equation ( dot - dashed line ) giving 
 pka = 12.4 . biochemists have developed an electrochemical technique suitable 
 for quantifying reduction potentials that are hampered by slow kinetics : redox potentiometry . solutions 
 varying the relative concentrations of 2oh and 2 were prepared by partial electrolysis of a ph 14 solution 
 of 2oh ( figure s7 ) . between 
 each stage of the electrolysis , the solution was allowed to reach 
 equilibrium ( as judged by a constant open circuit potential ) and the 
 concentrations of the ir species were determined by uv  vis 
 ( figure 2a ) . this method 
 provided e = 0.60 
 v for the reduction of 2oh to 2 at ph 14 
 ( figure 2b ) . as thermodynamic 
 constants for proton reduction are determined at the standard state 
 of ph 0 , this reduction was extrapolated to ph 0 
 by applying a 29.5 mv per ph unit shift ( 2e reduction 
 with loss of hydroxide ) , giving e = 0.19 v. redox potentiometry is seldom used in organometallic 
 chemistry , but this method was essential to overcoming the slow 
 kinetics that prevented the straightforward measurement of thermodynamic 
 values . with a reduction potential in hand , hydricity could 
 be determined 
 if paired with the metal hydride pka value . spectrophotometric titrations established the pka of the carboxylic acid groups in [ cp*ir(bpy - cooh)(oh2 ) ] and [ cp*ir(bpy - cooh)(h ) ] as 1.9 
 and 2.7 , respectively . the acidity of 2h was then measured 
 spectrophotometrically by addition of base to a yellow - orange solution 
 of 2h to produce a deep purple solution of 2 ( figure 2c ) , providing 
 pka(2h ) = 12.4 . the relatively 
 acidic carboxylic acid groups provide a doubly anionic supporting 
 ligand at ph 7 and ensure that hydride donation will not be coupled 
 to protonation changes at the ligand.1234 the hydricity of reference complex 2h was established 
 by combining the pka of 
 the metal hydride ( eq 1 ) , the oxidation potential of the conjugate base ( eq 2 ) , and the free energy of proton 
 reduction to hydride ( eq 3 , 34.2 kcalmol ) . this thermochemical cycle provides gh(oh ) = 42.4 kcalmol ( eq 4 ) , employing the 
 conventional standard state of ph 0 . the hydricity gh(oh ) is the free energy 
 of hydride transfer from 2h with formation of the hydroxo 
 complex 2oh . hydroxide 
 binding is involved in the experimentally measured reduction potential , 
 so thermochemistry involving this ligand is obtained directly . the 
 free energy of hydride transfer from 2h with formation 
 of the aquo complex 2oh2 can also be determined 
 by taking into account the pka of 2oh2 ( eq 5 , pka = 7.6 by spectrophotometric 
 titration ) : gh(oh2 ) = 32.0 kcalmol.5 an unusual situation arises 
 when taking into account the metal 
 aquo acidity : there are two different hydricity values 
 for 2h , gh(oh ) and gh(oh2 ) . formal hydride transfer initially results 
 in a 16e complex with a vacant coordination site , 
 and this hydride dissociation process ( gh in scheme 1 ) is most commonly associated with hydricity . in many 
 cases , however , the coordinatively unsaturated complex rapidly binds 
 a ligand ( e.g. , solvent or a counterion ) during the net hydride transfer 
 process . in organic solvents , solvation of the metal center after 
 hydride transfer is commonly ignored in the thermochemistry : the activity 
 of the solvent is taken as unity . water inherently contains hydroxide 
 ions capable of binding the metal center , leading to a distinct ( and 
 ph dependent ) hydricity value . the obtained thermodynamic values 
 gh(oh ) 
 and gh(oh2 ) include the formal hydricity and the binding affinity for the incoming ligand ( scheme 1 ) . a similar situation 
 arises for acidities when , following proton loss , aggregation through 
 hydrogen - bonding interactions ( e.g. , homoconjugation ) influences effective 
 acidity . to distinguish the different 
 effective hydricity values that couple hydride transfer and ligand 
 association , aqueous catalysis is typically carried out in the presence 
 of various 
 buffers and salts , and these species can also alter hydricity through 
 metal ligation . to better understand the role of incoming ligands , 
 we explored the effect of phosphate and chloride on hydricity . effective 
 hydricity values were determined by measuring the free energy of ligand 
 exchange with 2oh2 ( figure 3a ) and adding that thermodynamic value to 
 gh(oh2 ) . the relative free energy of chloride substitution was determined 
 by nmr titration of nacl into a pd 7 solution of 2oh2 , goh2cl = 4.4 kcalmol ( eq 6 ) . because ligand exchange is slow 
 on the nmr time scale , the concentrations of the iridium species could 
 be determined directly . the hydricity of 2h to form the 
 chloride product is thus gh(cl ) = 27.6 kcalmol ( figure 3a).67 the phosphate buffer presents both h2po4 and hpo4 ligands 
 at 
 ph 7 , either of which can bind ir(iii ) . phosphate binding is apparent by nmr spectroscopy in ph 7 phosphate 
 buffer , but rapid proton exchange prevents precise identification 
 of the ligand protonation state . the relative binding affinity of 
 the phosphate mixture ( goh2pi = 1.9 kcalmol at ph 7 , eq 7 ) provides gh(pi ) = 30.1 kcalmol . this hydricity is strictly accurate only at ph 7 , where the measurement 
 was made for the specific h2po4/hpo4 mixture which pi represents . the concentrations 
 of h2po4 and hpo4 will change based on the solution ph , however , 
 which could impact hydride transfer . ( a ) summary of thermochemical values of 
 [ cp*ir(bpy - coo)(h ) ] ( 2h ) . free energies 
 ( kcalmol ) and reduction potentials ( v vs 
 nhe ) are cited at 
 the standard state of ph 0 , 1 m reagents , and 1 atm gases , except 
 for gh(pi ) that refers to ph 7 . ( b ) summary of the 
 ph dependence of gh(y ) with the h2o / h2 and co2/hco2 couples . complex 2h is substantially more hydridic in 
 water 
 ( smaller gh(oh2 ) value ) than in acetonitrile , consistent with prior 
 studies . the large 
 differences in hydricity as a function of the ligands present in aqueous 
 solution , however , were previously unexplored and suggest that water 
 plays a role in hydride transfer reactions beyond simply providing 
 a high polarity medium . transition metal hydride transfer can be described 
 by a manifold of hydricity values comprised of the 
 heterolytic m the effective hydricity , gh(y ) , is expected to 
 be experimentally relevant 
 to catalysis . hydride transfer reactions for d hydrides 
 during catalysis will involve ligand association , so understanding 
 the overall thermodynamics of that process is vital . for example , in a typical ph 7 phosphate buffer solution 
 used in photoelectrocatalytic h2 evolution , hydride 2h reacts with water to 
 release h2 and generate an equilibrium mixture of ir(iii ) 
 chloride , aquo , and phosphate complexes  representing three 
 different h2 release pathways with three different hydricity 
 values . in water for one , 
 the 
 h2o / h2 potential will shift to lower values 
 as ph increases ( 1.36 kcalmolph ) , indicating that as protons become scarcer , stronger 
 hydrides are required to evolve h2 . yet while h2 is shifting , gh(cl ) and gh(oh2 ) remain constant across the accessible ph range , 
 altering net h2 release thermodynamics . on the other hand , 
 gh(oh ) is influenced 
 by ph as the concentration of ligand available for binding changes 
 with ph . at ph 0 , hydroxide ligation is unfavorable , leading to gh(oh ) > gh(oh2 ) ; 
 while chemical intuition might suggest that hydride transfer to form 
 an aquo complex would be a less favorable than hydride transfer to 
 form a complex with the more basic hydroxide ion , the extremely low 
 concentration of hydroxide at ph 0 leads to unfavorable energetics . as the solution ph increases , however , formation of the hydroxide 
 complex will become more favorable , and the value gh(oh ) will shift smoothly . though the differences caused by incoming ligands 
 in the aqueous 
 medium are striking , their impact is best assessed by comparison to 
 the effect of changing the metal center and supporting ligands . modification 
 of the structure of the hydride is the most common route to tune hydricity , 
 and these synthetic strategies are typically assumed to have a greater 
 influence than solvation of the product . to make these comparisons , 
 we sought to explore a wider range of metal complexes and began by 
 determining the hydricity of another soluble  reference hydride 3h hails from a family of ( arene)ru(diimine ) 
 catalysts that carry out aqueous transfer hydrogenation , water splitting , 
 and co2 reduction . the reduction potential of [ ( cymene)ru(bpy - coo)(oh ) ] ( 3oh ) between ph 8 and 12 was measured by cv . the quasi - reversible 3oh/3 couple ( ep = 60 mv ) shifted 26 mv per ph unit , close to the ideal value of 
 29.5 mv expected for a 1oh/2e process ( figure s20 ) . extrapolating the 
 trend in e1/2 to ph 0 provided the standard 
 reduction potential e = 0.30 
 v. spectrophotometric titrations 
 provided the acidity of the hydride 3h , pka = 11.8 . 
 from the pka and e , 
 gh(oh ) 
 = 36.5 kcalmol can be determined . ( all the relevant pka and gh values for this system are collected in table s1 . ) chloride 
 association free energy , goh2cl = 2.9 kcalmol , 
 was significantly smaller than that of the ir complex . taken together , 
 the hydricity to form the chloride was determined to be gh(cl ) = 23.1 kcalmol . having established two well - defined reference 
 hydricity values , 
 we set out to determine the hydricity of related hydrides , including 
 the parent bpy complexes . to probe hydride transfer equilibria between 
 ir and ru hydrides , however , chloride counterions were sought to increase water solubility 
 ( the previously reported pf6 and cf3so3 salts of 1h 
 were insoluble above 2 mm in water ) and 
 to reduce speciation . electrochemical and chemical synthetic 
 methods were developed to 
 provide rapid access to a wide range of water - soluble metal hydrides . 
 in a representative controlled potential electrolysis , the chloride 
 salt of 1cl was converted to > 20 mm of 1h in 0.1 m ph 7 napi . if the ph and electrolysis 
 potential were appropriately controlled to facilitate a reduction protonation 
 sequence , the electrolysis method was quite general , as detailed in 
 the supporting information ( p. s20 ) . chemical 
 syntheses were also carried out , as needed , according to a newly developed 
 procedure . for example , reduction of the chloride salt of 1cl by nabh4 in 1 m naoh resulted in precipitation of purple 1 in nearly quantitative yield . dropwise addition of hclet2o to a stirring solution of 1 in et2o prompted precipitation of the golden yellow chloride salt of hydride 1h . this procedure is also generally applicable , except when 
 the metal hydride can not be deprotonated in water or the conjugate 
 base does not precipitate from water ( see supporting information p. s20 for full details ) . with a collection 
 of hydride complexes ( see figure 4 for numbering scheme ) , relative hydricity 
 could be determined by mixing a hydride donor and a hydride acceptor 
 and allowing the system to reach an equilibrium distribution of both 
 hydrides and acceptors . the concentration of each species was determined 
 by nmr , and the equilibrium constant provided the difference in hydricity 
 ( gh ) between the two complexes , according to figure 1b.figure 4 depicts the relative 
 hydricity of each hydride complex , with each reaction representing 
 a hydride  chloride exchange . in a representative hydride equilibration , a solution of 2cl in pd 7 0.1 m napi ( produced 
 electrochemically in 84% yield , with 16% unreacted 2cl ) 
 was mixed with 1cl . after the reaction was allowed to 
 reach equilibrium , the concentrations of 1h , 1cl , 2h , and 2cl were measured by h nmr spectroscopy . the equilibrium constant , keq = 0.35 , provided gh = 0.6 kcalmol ( eq 8) and established 
 the hydricity of 1h in a single experiment : gh(cl ) = 27.0 kcalmol . it is noteworthy that equilibration was established 
 in < 15 min , and though our present focus is on thermodynamic hydricity , this contrasts with the frequently kinetically slow hydride 
 transfer reactions reported in acetonitrile.8 relative hydricity values of ir and ru complexes ( blue ) . aqueous hydricity scale of the complexes we 
 report along with those 
 previously reported in the literature . y represents the incoming ligand 
 such that the top scale shows gh(cl ) and the bottom scale shows gh(oh2 ) . tspp = tetra(p - sulfonatophenyl)porphyrin ; tmps = 
 tetrakis(3,5-disulfonatomesityl)porphyrin ; tpy = terpyridine ; dhmpe 
 = 1,2-bis(dihydroxymethylphosphino)ethane . a series of hydride transfer equilibrium experiments 
 established 
 the relative hydricity scale of figure 4 . equilibrium could be established from either direction 
 to give gh values that were identical within experimental uncertainty 
 ( 0.1 kcalmol , see supporting information p. s36 for full experimental details ) . hydricity values were determined from these relative hydricities by 
 comparison to the gh(cl ) of reference 2h for ir complexes and reference 3h for ru complexes , and the scale is self - consistent within 
 the 1 kcalmol estimated uncertainty 
 of the measurements . the gh(oh2 ) for all complexes was determined by measuring the 
 aquo - chloride relative association energy of each of these species 
 ( table 1 ) . our values are also 
 consistent with one of the few other well - defined 
 hydricity values available in the literature : the hydricity of [ ( c6me6)ru(bpy)(h ) ] ( 5h ) with 
 formation of 5oh2 was reported by creutz , 
 gh(oh2 ) = 22.2 kcalmol , which we independently determined to be gh(oh2 ) = 22.9 kcalmol . in figure 5 , our 
 continuum of hydricity values is contextualized against previously 
 reported hydricity values ( gh(oh2 ) ) for transition metal hydrides 
 and substrates relevant to alternative energy pursuits ( h and co2 ) . the two parallel scales illustrate 
 the role of the ligand bound to the product and the influence of changes 
 to the supporting ligands or metal center . in general , the hydricity 
 values electron - donating groups promote hydride transfer , as evidenced 
 by a strong correlation between gh(cl ) and the hammett parameter p ( figure 6 ) . the ir and ru catalysts investigated 
 herein become better hydride donors ( lower gh ) with increasing electron density . 
 hydricity is moderated by electronic effects : increasing electron 
 density increases pka and shifts e more negative which raise and lower hydricity , respectively . in fact , other systems have been found to be more influenced by the 
 ligand bite angle than by electronics . the ease with which each ligand can stabilize increased electron 
 density is reflected in electronic spectroscopy : hydricity is correlated 
 to the metal - to - ligand charge - transfer band around 400 nm that is 
 present in each of the ir hydride complexes ( figure s42 ) . interestingly , the activity of aqueous 
 hydrogen evolution catalysis involving cp*ir - based catalysts also 
 correlates with electron - donating ability of the bipyridine ligand , suggesting that perhaps the increase in rate 
 is due to an increase in the hydricity of the metal hydride intermediate . electronic changes to the bipyridine 
 ligands affect the acidity of the metal hydride more 
 dramatically than the 
 hydricity . the hydricity difference between methoxy - substituted 7h ( pka > 14 ) and methylester - substituted 8h ( pka  5 , estimated 
 from cv , see supporting information p. 
 s4 ) is only 2.4 kcalmol , while the acidity 
 difference between these complexes spans 9 orders of magnitude 
 ( 12 kcalmol ) . ligand effects 
 on hydricity were more pronounced when changes were 
 made to the arene rings . cymene complex 4h and hexamethylbenzene complex 5h displayed 
 a 3 kcalmol difference in hydricity 
 that is larger than observed for bpy ligand modifications , but of 
 a similar magnitude to the effect of chloride ligation . these differences 
 warrant further studies into possible steric effects in these thermodynamic 
 hydricity values . the emerging picture of aqueous hydricity 
 tunable by both ligands and the medium could impact 
 catalysis . electrocatalytic 
 hydrogen evolution in water is usually carried out with ph - stabilizing 
 buffer bases , and water splitting schemes 
 that employ salt water must wrestle with an abundance of chloride , which would lead to a 5 kcalmol difference in the hydricity of ir catalysts . the hydricity trends 
 in figure 5 also predict 
 the ph at which h2 evolution will occur , as a function 
 of the ligand electronics and the presence of incoming ligands in 
 solution . all of the complexes investigated , for example , are predicted 
 to produce h2 at ph 0 ( gh < 34.2 kcalmol ) , but at ph 10 only ru complex 5h is thermodynamically 
 capable of forming h2 ( and only at high chloride concentration ) . 
 under basic conditions , the ability of a hydride to reduce 
 co2 to formate at 
 ph 0 can also be predicted by inspection of figure 5 . species more hydridic than formate ( gh < 24.1 kcalmol ) are thermodynamically capable of co2 reduction . an intriguing prediction arises from figure 5 : co2 reduction by hydride transfer 
 from ( cymene)ru complexes 3h and 4h should 
 be unfavorable in unbuffered water and favorable only when chloride 
 anion is present . the less hydritic hydrides would require increased 
 co2 pressure to enable hydride transfer to co2 . in a prior report of co2 hydrogenation , the ir hydride 1h underwent slow , rate - limiting hydride transfer to co2 , while the ru hydride 5h transferred hydride 
 sufficiently quickly that h2 cleavage became rate - limiting . our studies show that 5h is more 
 hydridic than the parent ir complex 1h ; the hydricity 
 scale correctly predicts that 5h will more readily hydrogenate 
 co2 ( and less readily cleave h2 ) . comparisons across a range of well - known catalytic intermediates 
 were enabled by an electrochemical technique well suited to the complications 
 of water and by new synthetic routes to water - soluble hydrides . thermodynamic 
 hydricity in water is not only influenced by the supporting ligands , 
 but also by the ligating species present in aqueous media . rather 
 than a single value defined in terms of the hydride donor , a continuum 
 of hydricity values should be considered . being cognizant of the resulting 
 product after hydride transfer makes direct comparisons between catalysts 
 and conditions possible . the synthetic chemist instinctively tunes catalysts through ligand 
 modifications , but tuning the medium itself can also effect changes 
 in hydricity . the present findings will guide further thermodynamic 
 studies of pcet events in water and guide aqueous catalyst development . all solutions containing metal 
 hydride species were protected from ambient light to prevent excited 
 state reactions . commercial hplc - grade 
 water was used as a solvent , and organic solvents were dried and degassed 
 with argon using a pure process technology solvent system . electrochemical experiments were performed on a pine wavenow potentiostat 
 or pine wavedriver bipotentiostat controlled by aftermath software . solution ph was recorded using an orionstar a111 ph meter with a beckman - coulter , 
 hanna , or hach isfet ph probe . vis spectra were obtained 
 using an ocean optics usb2000 + spectrometer with a dt - mini-2gs deuterium / tungsten 
 halogen light source controlled by oceanview software . h 
 nmr spectra were referenced to the residual solvent signals ( or dioxane 
 or naots as an internal standard in d2o ) . spectra were processed using the mestrenova software suite 
 from mestrelab research s. l. the solution acidity in nmr experiments 
 is reported as pd , obtained by addition of + 0.4 to the reading of 
 a ph electrode that was calibrated using h2o standards . esi - ms were obtained on a thermo scientific 
 ltq ft - icr ms with 
 samples introduced either through direct infusion or by lc . inductively 
 coupled plasma - mass spectrometry ( icp - ms , agilent technologies 7500x 
 series ) was employed to determine the precise ir and ru concentrations 
 in uv  vis samples ( for molar extinction coefficient determination ) , 
 with the aid of a calibration curve for 10500 ppb ir and ru . electrochemical experiments were carried 
 out with carbon working electrodes , platinum wire counter electrodes , 
 and ag / agcl ( 3 m nacl ) reference electrode in a small glass tube fitted 
 with a vycor glass frit . solutions were thoroughly degassed by sparging 
 with nitrogen for at least 15 min before beginning an experiment . all potentials are reported relative to nhe , with values obtained 
 by adding 0.21 v to the experimentally observed potential vs ag / agcl . cyclic voltammetry experiments were carried 
 out with a glassy carbon working electrode ( polished with 0.05 m 
 alumina powder between scans ) in an undivided cell . controlled potential 
 electrolysis experiments were carried out with reticulated vitreous 
 carbon ( rvc ) as the working electrode separated from the counter electrode 
 and reference electrodes by a fine frit in an h - cell . potentiometric 
 experiments were performed in a custom - made three - compartment 
 cell divided by fine frits and with a 10 mm  10 mm pyrex glass 
 cuvette affixed to the central working electrode chamber ( figure s1 ) . the solution was stirred at the base 
 of the cuvette and by a slow bubble of n2 through the length 
 of the cuvette to ensure sufficient mixing near the electrode . an 
 rvc electrode was used as the working electrode for both the electrolysis 
 and open circuit potential experiments . reduction and oxidation of 
 the analyte was achieved via short periods of electrolysis , and after 
 each pulse of current , sufficient time was allowed for the solution 
 components to come into equilibrium ( typically 510 min ) as 
 judged by an unchanging open circuit potential over 30 s. after equilibrium 
 was established , uv the complexes [ cp*ir(bpy)(cl)][cl ] ( 1cl ) , [ ( cymene)ru(bpy)(cl)][cl ] 
 ( 4cl ) , [ ( c6me6)ru(bpy)(cl)][cl ] 
 ( 5cl ) , [ cp*ir(bpy - me)(cl)][cl ] 
 ( 6cl ) , and [ cp*ir(bpy - ome)(cl)][cl ] ( 7cl ) 
 were prepared following the method of dadci et al . [ cp*ir(cl)2]2 , [ cp*ir(bpy - cooh)(cl)][cl ] , and [ cp*ir(bpy)(h)][otf ] were prepared following literature procedures . [ cp*ir(bpy)(oh2)][so4 ] ( 1oh2 ) , [ cp*ir(bpy - cooh)(oh2)][otf]2 ( 2oh2 ) , and 
 [ ( cymene)ru(bpy - cooh)(oh2)][otf]2 ( 3oh2 ) were prepared 
 following literature procedures with the appropriate silver salt . electrolysis 
 of 2cl in napi , na2so4 , or naoh electrolytes ( depending on the desired use of the product ) 
 past the first reduction feature (   1.0 v ) resulted 
 in conversion to reduced products , consistent with previously reported 
 spectroscopic and electrochemical properties . the form of these products ( either 2h or 2 ) was highly dependent on solution ph , giving 2h at 
 neutral ph , 2 at high ph , and a mixture in between . to 
 confirm the identities of these reduced products , 2cl 
 ( 9.3 mg , 0.014 mmol ) was reduced by excess nabh4 ( 3.7 mg , 
 0.98 mmol ) by stirring for 30 min in meoh . 2h : h nmr ( 600 mhz , d2o + dioxane ) 
  8.69 ( d , j = 5.9 hz , 2h ) , 8.55 ( s , 2h ) , 7.72 
 ( d , j = 5.9 hz , 2h ) , 1.77 ( s , 15h ) , 11.90 
 ( s , 1h ) . abs , max ( ph 7 0.1 m napi ) = 428 nm ( 3700 m cm ) . 2 : h nmr ( 400 mhz , d2o + dioxane ) 
  8.81 ( d , j = 7.0 hz , 1h ) , 8.42 ( s , 1h ) , 6.82 
 ( dd , j = 6.9 , 2.0 hz , 1h ) . 1.80 ( s , 15 h ) . abs , max ( 1 m naoh ) = 292 nm ( 22  000 m cm ) , 364 nm ( 10  800 m cm ) , 535 nm ( 23  000 m cm ) . under 
 nitrogen , [ ( cymene)rucl2]2 ( 50.3 mg , 0.082 mmol ) 
 and bpy - cooh ( 40.3 mg , 0.165 mmol ) were allowed to stir in 8 ml dmf 
 at 60 c for 3 h. after filtering the solution in air to remove 
 unreacted ligand , the dmf was removed in vacuo . the resulting film 
 was dissolved in meoh , and yellow 3cl ( 83.4 mg , 92% yield ) 
 precipitated from solution on addition of ether . the hydride 3h and reduced complex 3 were prepared according 
 to the procedures for 2h and 2 . 3h : h nmr ( 600 mhz , d2o + dioxane )  
 8.87 ( d , j = 5.8 hz , 2h ) , 8.53 ( d , j = 1.5 hz , 2h ) , 7.64 ( dd , j = 5.9 , 1.6 hz , 2h ) , 
 5.70 ( d , j = 6.2 hz , 2h ) , 5.51 ( d , j = 6.1 hz , 2h ) , 2.63 ( sept , j = 6.9 hz , 1h ) , 2.24 
 ( s , 3h ) , 1.07 ( d , j = 6.9 hz , 6h ) , 6.21 ( s , 
 1h ) . abs , max ( ph 7 0.1 m napi ) = 295 nm ( 20  800 m cm ) , 434 nm ( 5500 m cm ) . 3 : h nmr ( 600 mhz , d2o + dioxane )  
 8.82 ( d , j = 6.8 hz , 2h ) , 8.37 ( s , 2h ) , 7.12 ( d , j = 6.7 hz , 2h ) , 4.84 ( d , j = 5.5 hz , 2h ) , 
 4.81 ( d , j = 5.7 hz , 2h ) , 2.32 ( sept , j = 7.7 hz , 1h ) , 1.85 ( s , 3h ) , 0.86 ( d , j = 6.7 hz , 
 6h ) . abs , max ( 1 m naoh ) = 298 nm ( 22  600 
 m cm ) , 374 nm ( 13  500 
 m cm ) , 506 nm ( 17  600 
 m cm ) , 610 nm ( 14  400 
 m cm ) . in a nitrogen filled glovebox , 
 [ cp*ir(bpy)(cl)][cl ] ( 15.5 mg , 0.028 mmol ) and excess nabh4 ( 8.5 mg , 0.223 mmol ) were allowed to stir in 2 ml of 1 m naoh . dark 
 purple solids quickly formed . after letting stir for 4 h , the solid was filtered off , washed 3 with water , collected in benzene , 
 and evaporated to dryness , yielding 1 ( 13.3 mg , 98% yield ) . 
 the h nmr spectrum of 1 prepared in this 
 way matched previously reported data . to a stirring 
 solution of 1 ( 13.3 mg , 0.028 mmol ) in ether , a dilute 
 solution of hclet2o ( 40 mm ) was added dropwise until 
 a change from a dark purple solution to bright yellow solids was observed . typically 11.5 equiv of hcl was added with the excess acid 
 immediately pumped off after completion of the addition . samples of 
 hydride prepared in this way typically contained small amounts ( < 5% ) 
 of [ cp*ir(bpy)(cl)][cl ] ( formed by protonation of hydride releasing 
 h2 ) , and the h nmr spectrum is consistent with 
 previously reported [ cp*ir(bpy)(h ) ] . deep purple 4 was prepared in quantitative yield from 
 the chloride salt 
 of 4cl , according to the procedure used in the synthesis 
 of 1 . the chloride 
 salt of 4h was prepared from 4 , according 
 to the procedure used in the synthesis of the chloride salt of 1h . h nmr ( 600 mhz , d2o + dioxane ) 
  8.78 ( d , j = 4.1 hz , 2h ) , 8.09 ( d , j = 8.2 hz , 2h ) , 7.87 ( t , j = 7.8 hz , 2h ) , 
 7.32 ( dd , j = 7.3 , 5.8 hz , 2h ) , 5.56 ( d , j = 6.0 hz , 2h ) , 5.36 ( d , j = 6.0 hz , 2h ) , 
 2.53 ( sept , j = 7.1 hz , 1h ) , 2.10 ( s , 3h ) , 0.99 ( d , j = 6.9 hz , 6h ) , 6.32 ( s , 1h ) . deep purple 5 was prepared in quantitative yield from 
 the chloride salt of 5cl , according to the procedure 
 used in the synthesis of 1 . the chloride salt of 5h was prepared from 5 , according to the procedure used in the synthesis of the 
 chloride salt of 1h . the h nmr spectrum is 
 consistent with the reported spectrum for the triflate salt of 5h in water . in a nitrogen filled 
 glovebox , [ cp*ir(bpy - me)(cl)][cl ] ( 5.8 mg , 0.010 mmol ) and excess 
 nabh4 ( 5.1 mg , 0.135 mmol ) were allowed to stir in 2 ml 
 of 5 m naoh , and a dark violet solid quickly formed . after letting 
 stir for 4 h , the solid was extracted into c6h6 , dried over mgso4 , and evaporated to dryness , yielding 6 in quantitative yield . h nmr ( 600 mhz , c6d6 )  8.91 ( d , j = 6.8 hz , 
 2h ) , 7.44 ( s , 2h ) , 6.06 ( dd , j = 6.8 , 2.1 hz , 2h ) , 
 2.02 ( s , 6h ) , 1.82 ( s , 15h ) . c nmr ( 151 mhz , c6d6 )  148.20 , 141.06 , 127.59 , 122.82 , 117.92 , 83.32 , 
 21.44 , 10.24 . the bright 
 yellow chloride salt of 6h was prepared from 6 , according to the procedure used in the synthesis of the chloride 
 salt of 1h . h nmr ( 600 mhz , d2o + dioxane )  8.55 ( d , j = 5.8 hz , 2h ) , 7.91 
 ( s , 2h ) , 7.32 ( d , j = 6.0 hz , 2h ) , 2.51 ( s , 6h ) , 
 1.75 ( s , 15h ) , 11.51 ( s , 1h ) . abs , max ( ph 
 7 0.1 m napi ) = 394 nm ( 2900 mcm ) . controlled 
 potential electrolysis of [ cp*ir(bpy - ome)(cl)][cl ] at 1.0 
 v in 0.1 m pd 7 napi resulted in clean 
 formation of [ cp*ir(bpy - ome)(h ) ] , consistent with previously 
 reported spectroscopic and electrochemical properties.h nmr ( 600 mhz , d2o + 
 dioxane )  8.49 ( d , j = 6.5 hz , 2h ) , 7.44 ( s , 
 2h ) , 7.06 ( dd , j = 6.6 , 2.8 hz , 2h ) , 3.98 ( s , 6h ) , 
 1.75 ( s , 15h ) , 11.23 ( s , 1h ) . in a nitrogen filled 
 glovebox , 8.1 mg ( 0.012 mmol ) [ cp*ir(bpy - coome)(cl)][cl ] and 4.7 g 
 ( 0.069 mmol ) nao2ch were stirred in 2 ml ph 7 0.1 m napi . while stirring for 4 h , a royal purple 
 solid precipitated from solution . the solution was filtered , and the 
 solids were washed 3 with water , collected by dissolving in 
 benzene , and evaporated under vacuum to yield 8 ( 3.4 
 mg , 47% yield ) . h nmr ( 600 mhz , c6d6 )  8.78 ( d , j = 7.0 hz , 1h ) , 8.69 ( d , j = 2.1 hz , 1h ) , 7.20 ( dd , j = 7.0 , 2.0 
 hz , 1h ) , 3.52 ( s , 2h ) , 1.55 ( s , 6h ) . c6d6 )  166.76 , 147.22 , 142.02 , 125.72 , 120.11 , 115.04 , 
 85.57 , 51.76 , 9.68 . the scarlet 
 chloride salt of 8h was prepared from 8 , 
 according to the procedure used in the synthesis of the chloride salt 
 of 1h . h nmr ( 600 mhz , d2o + dioxane ) 
  8.89 ( d , j = 5.9 hz , 2h ) , 8.69 ( s , 2h ) , 7.92 
 ( d , j = 5.9 hz , 2h ) , 4.10 ( s , 6h ) , 1.80 ( s , 15h ) , 
 12.28 ( s , 1h ) . c nmr ( 151 mhz , d2o ) 
  165.82 , 155.95 , 153.00 , 138.06 , 126.69 , 123.19 , 92.70 , 54.39 , 
 8.94 . abs , max ( ph 3 0.1 m napi ) = 388 nm ( 4200 mcm ) , 451 nm ( 4300 mcm ) , 481 nm ( 4400 mcm ) . in a typical equilibration experiment 
 to determine relative hydricity according to figure 1b , 19.3 mg of 2cl was dissolved 
 by sonication in 2 ml of pd 7 0.1 napi , added to the working electrode compartment of an h - cell , and degassed 
 for 15 min . the solution was electrolyzed 
 at 1.0 v for 6 h , transferred to a n2 purged bomb 
 flask , and brought into a glovebox . different volumes of the electrolyzed 
 solution ( 100 , 200 , and 300 l ) were added to three samples 
 containing 3.5 mmol 1cl and dioxane , and the total volume 
 was brought to 500 l . samples were monitored by h nmr , and equilibrium of the experimental samples was quickly achieved ; 
 though the samples were monitored over 25 h by h nmr , 
 equilibrium ( keq = 0.35 ) was established 
 by the first time point , giving gh = 0.6  0.1 kcalmol . alternatively , following protonation with hclet2o , the solid hydride was extracted into the nmr solvent ( either 
 pd 7 0.1 m napi or pd 4.3 20 mm naoac 
 with dioxane internal standards ) , filtered to remove any residual 
 cp*ir(bpy - x ) or ( arene)ru(bpy - x ) , and combined with a hydride acceptor . 2.0 mg of [ cp*ir(bpy)(h)][cl ] ( 0.004 mmol ) was dissolved 
 in 490 l of pd 7 0.1 m napi with 
 10 l of 0.5 m dioxane as an internal standard . after confirming 
 the purity of the hydride sample by h nmr , 2.1 mg [ cp*ir(bpy - me ) 
 ( cl)][cl ] ( 0.004 mmol ) for each species , 
 a series solutions of a known concentration of chloride were prepared 
 in pd 7 napi and monitored by nmr to ensure 
 that the aquo , phosphate , and chloride species were in equilibrium . for example , in air , a 5.5 mm solution of [ cp*ir(bpy - cooh)(oh2)][otf]2 in 50 mm pd 7 napi with a dioxane internal standard was split between six samples 
 each containing dry nacl to produce final solutions with [ cl ] from 0 to 18 mm . the samples were monitored by h nmr 
 over 24 h to ensure that equilibrium had been established between 
 [ cp*ir(bpy - coo)(oh2 ) ] , [ cl ] and [ cp*ir(bpy - coo)(cl ) ] . the initial [ cl ] left from the halide abstraction with agotf was fit 
 by minimizing the variance of g of the 0 mm 
 nacl added sample with that of the remaining five samples . the free 
 energy of the ligand exchange was found to be 4.4  0.2 
 kcalmol . the relative aquo - phosphate association 
 free energy was determined similarly with solutions of increasing 
 total [ pi ] at pd 7 .
aqueous 
 hydride transfer is a fundamental step in emerging alternative 
 energy transformations such as h2 evolution and co2 reduction . 
 hydricity ,  the hydride donor ability 
 of a species , is a key metric for understanding transition metal hydride 
 reactivity , but comprehensive studies of aqueous hydricity are scarce . 
 an extensive and self - consistent aqueous hydricity scale is constructed 
 for a family of ru and ir hydrides that are key intermediates in aqueous 
 catalysis . 
 a reference hydricity is determined using redox potentiometry 
 and spectrophotometric titration for a particularly water - soluble 
 species . 
 then , relative hydricity values for a range of species are 
 measured using hydride transfer equilibria , taking advantage of expedient 
 new synthetic procedures for ru and ir hydrides . 
 this large collection 
 of hydricity values provides the most comprehensive picture so far 
 of how ligands impact hydricity in water . 
 strikingly , we also find 
 that hydricity can be viewed as a continuum in water : 
 the free energy of hydride transfer changes with ph , buffer composition , 
 and salts present in solution .