doi.org/10.26434/chemrxiv.13114019.v1

Temperature and Solvent Effects on H2 Splitting and Hydricity:Ramifications on CO2 Hydrogenation by a Rhenium Pincer CatalystJenny Hu, Quinton J. Bruch, Alexander Miller

Submitted date: 19/10/2020 • Posted date: 20/10/2020Licence: CC BY-NC-ND 4.0Citation information: Hu, Jenny; Bruch, Quinton J.; Miller, Alexander (2020): Temperature and Solvent Effectson H2 Splitting and Hydricity: Ramifications on CO2 Hydrogenation by a Rhenium Pincer Catalyst. ChemRxiv.Preprint. https://doi.org/10.26434/chemrxiv.13114019.v1

The catalytic hydrogenation of carbon dioxide holds immense promise for applications in sustainable fuelsynthesis and hydrogen storage. Mechanistic studies that connect thermodynamic parameters with thekinetics of catalysis can provide new understanding and guide predictive design of improved catalysts.Reported here are thermochemical and kinetic analyses of a new pincer-ligated rhenium complex(tBuPOCOP)Re(CO)2(tBuPOCOP = 2,6-bis(di-tert-butylphosphinito)phenyl) that catalyzes CO2 hydrogenationto formate with faster rates at lower temperature. Because the catalyst follows the prototypical “outer sphere”hydrogenation mechanism, comprehensive studies of temperature and solvent effects on the H2splitting andhydride transfer steps are expected to be relevant to many other catalysts. Strikingly large entropy associatedwith cleavage of H2 results in a strong temperature dependence on the concentration of[(tBuPOCOP)Re(CO)2H]– present during catalysis, which is further impacted by changing the solvent fromtoluene to tetrahydrofuran to acetonitrile. New methods for determining the hydricity of metal hydrides andformate at temperatures other than 298 K were developed, providing insight into how temperature caninfluence the favorability of hydride transfer during catalysis. These thermochemical insights guided theselection of conditions for CO2 hydrogenation to formate with high activity (up to 364 h–1 at 1 atm or 3330 h–1

at 20 atm of 1:1 H2 CO2). In cases where hydride transfer is the highest individual kinetic barrier, entropiccontributions to outer sphere H2 splitting lead to a unique temperature dependence: catalytic activityincreases as temperature decreases in tetrahydrofuran (200-fold increase upon cooling from 50 to 0 °C) andtoluene (4-fold increase upon cooling from 100 to 50 °C). Ramifications on catalyst structure-functionrelationships are discussed, including comparisons between “outer sphere” mechanisms and metal–ligandcooperation mechanisms.

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Temperature and Solvent Effects on H2 Splitting and Hydricity: Ramifications on CO2 Hydrogenation by a Rhenium Pincer Catalyst JennyHu,‡QuintonJ.Bruch,‡andAlexanderJ.M.Miller*

DepartmentofChemistry,UniversityofNorthCarolinaatChapelHill,ChapelHill,NorthCarolina27599–3290,UnitedStates

ABSTRACT: The catalytic hydrogenation of carbon dioxide holds immense promise for applications in sustainable fuelsynthesisandhydrogenstorage.Mechanisticstudiesthatconnectthermodynamicparameterswiththekineticsofcatalysiscanprovidenewunderstandingandguidepredictivedesignofimprovedcatalysts.Reportedherearethermochemicalandkinetic analyses of a new pincer-ligated rhenium complex (tBuPOCOP)Re(CO)2 (tBuPOCOP = 2,6-bis(di-tert-butylphosphinito)phenyl)thatcatalyzesCO2hydrogenationtoformatewithfasterratesatlowertemperature.Becausethecatalyst follows the prototypical “outer sphere” hydrogenation mechanism, comprehensive studies of temperature andsolventeffectsontheH2splittingandhydridetransferstepsareexpectedtoberelevanttomanyothercatalysts.Strikinglylarge entropy associated with cleavage of H2 results in a strong temperature dependence on the concentration of[(tBuPOCOP)Re(CO)2H]– present during catalysis, which is further impacted by changing the solvent from toluene totetrahydrofurantoacetonitrile.Newmethodsfordeterminingthehydricityofmetalhydridesandformateattemperaturesotherthan298Kweredeveloped,providinginsightintohowtemperaturecaninfluencethefavorabilityofhydridetransferduringcatalysis.ThesethermochemicalinsightsguidedtheselectionofconditionsforCO2hydrogenationtoformatewithhighactivity(upto364h–1at1atmor3330h–1at20atmof1:1H2CO2). Incaseswherehydridetransfer isthehighestindividual kinetic barrier,entropic contributions to outersphereH2 splitting lead to a unique temperaturedependence:catalyticactivityincreasesastemperaturedecreasesintetrahydrofuran(200-foldincreaseuponcoolingfrom50to0°C)andtoluene (4-fold increase upon cooling from 100 to 50 °C). Ramifications on catalyst structure-function relationships arediscussed,includingcomparisonsbetween“outersphere”mechanismsandmetal–ligandcooperationmechanisms.

1.INTRODUCTION

Hydrogenation reactions are ubiquitous in chemistry, withindustrial applications in pharmaceutical synthesis,commoditychemicalsmanufacturing,andfuelgeneration.ThehydrogenationofCO2toformateorformicacidhasattractedparticularly intense interest inthecontextofcarbonfixationand liquid fuel synthesis.1 Efficient CO2 hydrogenation oftenrequires elevated temperatures, even though entropicpenalties render the overall reactionmore unfavorablewithincreasing temperature (Scheme 1A, see SI forthermochemicaldetails).This couldbedue toahighkineticbarrier under ambient conditions; however, a distinctpossibility is that an individual step in the cycle becomesthermodynamically unfavorable at low temperature.Mechanistic studies thatconnectkinetic and thermodynamicanalysis could elucidate the factors that control thetemperaturedependenceofCO2hydrogenation.Twogeneralmechanisms,showninScheme1B,arefrequentlyproposedinleadinghydrogenationcatalysts.1–8ThemodeofH2splitting is distinct in these two mechanisms: the “outersphere” mechanism utilizes an external base to produce amonohydrideinaformallytermolecularreaction(sometimesinvolvingtwosequentialbimolecularsteps).3,4Conversely, inthemetal–ligand cooperationmechanisma basic site on theligandisproposedtoassistintheheterolyticcleavageofH2toformamonohydrideandaprotonatedligand.5,7,8Enthalpyandentropyofreactionparametershavebeendeterminedforafewcatalysts that split H2 via the metal-ligand cooperativemechanism.9–11 However, to our knowledge no analogoustemperature-dependentthermochemicaldata isavailable for

catalyststhatfollowanoutersphereH2splittingmechanism,precludingcomparisonsbetweenmechanisms.Scheme1.ThermodynamicanalysisofCO2hydrogenation.

AfterH2heterolysis, thesubsequentC–Hbond-formingstepsare often also different between the two mechanisms ofScheme1B.Inmetal–ligandcooperativepathways,concertedhydrideandprotontransfertoCO2isofteninvoked,althoughCO2bindingandmigratoryinsertionorothermechanismsarealsopossible.5,7 In theouter spherepathway, theH2-derived

LnM

LnM H

HBase–

H2LnM

HH

+

LnMOH

CH

O

CO2

CO2

[HBase][HCO2] [HBase][HCO2]

H2(g) + CO2(g)

[HBase][HCO2](aq)HCO2H(aq) + Base(aq)

∆G100ºC = 8.8 kcal·mol–1∆G25ºC = 5.4 kcal·mol–1

[LnM H]– LnM + H–

CO2 + H– HCO2–

∆GH–,LnMH–= ∆HºH– – T∆SºH––∆GH–,HCO2– = –∆HºH– + T∆SºH–

[LnM H]– LnM + HCO2–+ CO2 ∆∆GH–

A. Hydrogenation of CO2 to formic acid or formate

B. Two mechanisms of CO2 hydrogenation

C. Relating hydricity (∆GH–) to hydride ion transfer

∆G = –2.3RT(∆pKa)

Hydride ion transferΔΔGH–

metal–ligand cooperation

outer sphere

Base

H2 HeterolysisΔG1

H2Base

HCO2H(aq)

hydrideundergoeshydrideiontransfertofreeCO2,generatingformate ion.Whenformatedoesnotbindtothecatalyst, thethermodynamicsofhydridetransfercanbedirectlyrelatedtothedifferenceinhydricity(ΔG°H–)betweenthemetalhydrideintermediate and the free formate ion (Scheme 1C).12Thermodynamichydricityvalueshavebeendeterminedforawide variety of transition metal hydrides in MeCN.12,13However,MeCNisrarelyusedinCO2hydrogenationcatalysis.Thefirsthydricityvaluesintetrahydrofuran(THF),asolventcommonly used in hydrogenation studies, have only veryrecently appeared.10,13,14 Irrespective of solvent, so farΔG°H–values have only been reported at the standard statetemperatureof25˚C,wherefewcatalystsoperate.Toconnectthetemperature-dependentthermochemistryofH2splitting and hydride transfer steps to overall catalytic CO2hydrogenationactivity,we setouttoprepareanewcatalystthatwould beamenable to detailed thermochemical studiesacrossarangeoftemperaturesandsolvents.Wehypothesizedthat a pincer rhenium carbonyl hydride anion would be apromising catalyst candidate, based on several trends inthermodynamic hydricity.13 First, early metal hydrides areoftenmorehydridicthanlatemetalhydrides(strongerhydridedonorswithlower∆G°H–values).12,13Second,anionichydridecomplexes tend tobemorehydridic thanneutralor cationichydridecomplexes.12Third,complexeswithahydridesittingacrossfromastrongtranseffectligand(COinthiscase)canbemore potent hydride donors.15–17 Rhenium complexes,includingpincercomplexes,cancatalyzethehydrogenationoforganic carbonyls;18–23 we are not aware of examples of RecatalystsforCO2hydrogenation.16Wealsoanticipatedanoutersphere hydrogenation mechanism based on the octahedralgeometrywith18valenceelectronsandthenegativechargeonthecomplex.Herein we report a new rhenium pincer catalyst for CO2hydrogenationthatprovidesaplatformforunderstandinghowtemperature and solvent affect formate synthesis by outersphere mechanisms. Thermodynamic studies of thetemperature dependence of H2 heterolysis in a variety ofsolventsrevealstrikingentropiccontributionstothehydrideformation step. Insight into the formate-producing hydridetransferstepcomefromnewmethodsforthedeterminationoftemperature-dependent thermodynamic hydricity. Kineticstudiesofcatalyticactivityrevealedconditionswherethenewcatalystproducesformateathigherratesatlowertemperaturesin THF and toluene. Mechanistic models that connect thethermodynamic and kinetic factors that enable improvedreactivityatlowtemperatureareintroduced,andimplicationsondesigningfuturecatalystsystemsareconsidered.

2.RESULTSANDDISCUSSIONSynthesis of RheniumCarbonyl Complexes. A rhenium(I)carbonyl complex supported by the anionic pincer ligandtBuPOCOP (2,6-bis(di-tert-butylphosphonito)phenyl) wasprepared by refluxing tBuPOCOP-H with Re(CO)5Cl andtriethylamine in chlorobenzene for 16 h (Scheme 2).Analyticallypure,colorlesscrystalsof(tBuPOCOP)Re(CO)3(1)wereobtaineduponcoolingasaturatedpentanesolutionto–30˚C.Thestructureof1wasascertainedthroughNMRandIRspectroscopy (nCO 2023,1923, and 1902cm–1) anda single-crystal X-ray diffraction study (Figure 1). The structuralmetricsandvibrationalspectraaresimilartorelatedpincerRecarbonylcomplexes.24–28

Scheme 2. Synthesis of tricarbonyl complex(tBuPOCOP)Re(CO)3(1).

In order to open a coordination site for H2 activation andhydride formation, carbonyl ligand removal methods wereexplored.Heatingwhitesolid1at200°Cundervacuumfor5hled to sublimation of a dark orange-brown solid that wasisolated and found by NMR spectroscopy to feature a newspecieswitha31Presonanceshifteddownfieldbyover20ppmrelativeto1(whichwasstillpresentasa27%impurity).ThisnewspeciesshowedjusttwostretchesinIRspectra(nCO1910,1840cm–1), consistentwith formationof (tBuPOCOP)Re(CO)2(2).25,29Asingle-crystalX-raydiffractionstudyconfirmedtheformula as the anticipated dicarbonyl complex in a squarepyramidalgeometry(Figure1).Evidencefor2havingstrongerp-backbonding than 1 comes from crystallographic data(shorter Re–C bonds in 2) and IR data (lower-energy COstretchesfor2).Thesolid-statestructureof2haspseudo-CSmolecularsymmetry,whilethesamecomplexexhibitspseudo-C2v symmetry in solution, perhaps reflecting a rapidisomerizationorotherstructuralfluxionality.25

Figure1.Structuralrepresentationsof(tBuPOCOP)Re(CO)3(1,top)and(tBuPOCOP)Re(CO)2(2,bottom)fromsinglecrystalx-ray diffraction (ellipsoids are set at 50%, hydrogen atomsomittedforclarity).SeeSISectionXforcrystallographicdetails.Isolationoffive-coordinate2inhighpuritywasessentialforsubsequentstudiesofH2heterolysis.Therefore,analternative,two-step decarbonylation strategy was designed based oninitialhydrideformationfollowedbyprotonolysistoafford2.We hypothesized that hydride transfer to complex1 wouldresultinaformylcomplexthatwouldreleaseCOandundergoa-migration to form a hydride complex (although othermechanismsarepossible).30–33Treatingthehydridewithacid

intheabsenceofCOwouldthenreleaseH2andform2.Heatingtricarbonyl1with5equivLiHBEt3(1MinTHF)at70 CinTHFledtoformationofarheniumhydridespecies(Scheme3,SIFigureS11).After removal of free CO via freeze-pump-thawcycling,additionofHCletherateledtovigorousH2evolutionasthe colorless solution turned dark brown. Dicarbonyl2wasisolatedin>99%puritybythismethod,basedonmultinuclearNMRspectroscopyandelementalanalysis.Scheme 3. Synthesis of dicarbonyl complex(tBuPOCOP)Re(CO)2 (2) and hydride complex[(tBuPOCOP)Re(CO)2H]–(3).

Attempts to isolate the hydride intermediate provedchallenging, as decomposition was always observed duringworkupduetoinstabilitytovacuumandextremesensitivitytotracemoisture.We found itmost efficient to treat2with 1equiv of NaHBEt3 (1 M in toluene) in toluene, followed byprecipitation with pentane. The isolated solid displayed ahydride 1H resonanceat –6.13ppmwith strong JPH coupling(25.8Hz) in freshlydriedTHF-d8.ThemoleculeexhibitedCSsymmetry, as reflected in twodistinct tert-butyl resonances,indicatingacisdicarbonylgeometry(furtherbuttressedbyIRspectra showing two carbonyl stretches,nCO 1873and 1754

cm–1) andallowing assignment as cis-[(tBuPOCOP)Re(CO)2H]–(3,Scheme3).TemperatureandSolventEffectsontheFreeEnergyofH2

Heterolysis, ΔG1. Five-coordinate complex2 and hydride3were considered likely catalytic intermediates that couldplausibly be interconverted by anH2 splitting reactionwithexogeneous base. Three organic solvents were selected forreactivity and thermochemistry studies: toluene, THF, andMeCN.These solvents not only spana range of polarity anddonorability,butTHFandtoluenearecommonlyemployedincatalysisandwerethesolventsusedinthermodynamicstudiesofH2splittingbymetal–ligandcooperativesystems.10,11WhileMeCNisnotoftenemployedinhydrogenationcatalysis,itisaclassicsolventforthermochemicalstudies.12,34,35Solutions of dicarbonyl 2 and various organic bases werepreparedineachofthethreesolventsandplacedunder1atmH2toprobeforhydrogencleavagereactivity.InthepresenceofNEt3, no reaction was observed in toluene, THF, or MeCN(though spectroscopic evidence for MeCN binding can beobserved,seeSIFigureS6).Infact,noevidenceforH2bindingwasobservedevenwhensolutionsof2intoluene-d8under1atmH2were cooled from298K to198KandmonitoredbyNMRspectroscopy.ThisindicatesthatH2coordinationto2issubstantiallyendergonic(SIFigureS21-S22).WehypothesizedthatastrongerbasethanNEt3mightbetterfacilitate outer sphere H2 splitting, even though a stabledihydrogen complexwas not observed.12,13 Indeed, evidenceforheterolyticH2cleavageandformationof3wasobservedwith stronger bases (Figure 2A). In MeCN-d3, DBU (1,8-diazabicyclo[5.4.0]undec-7-ene) proved to be the ideal base,producingequilibriummixturesof2and3under1atmH2(seeSI Section VI for more details). In tetrahydrofuran-d8, H2splitting with TBD (1,5,7-triazabicyclo[4.4.0]dec-5-ene)producedasimilarequilibriummixture.Intoluene-d8at25˚C,Verkade’s base (Vk’s, 2,8,9-triisobutyl-2,5,8,9-tetraaza-1-phosphabicyclo[3.3.3]undecane) furnished completeconversion of brown2 to colorless3; however, the solutioncolor turned brown upon heating, indicating access to thedesiredequilibriumathighertemperatures.

Figure2.A)Enthalpy,entropy,andfreeenergyvaluesforH2splitting.aB)van’tHoffplotsconstructedusingK1determinedbyvariable-temperatureNMRintoluene(bluesquares,Vk’s),MeCN(redcircles,DBU),andTHF(blacktriangles,TBD).C)Molefractionof3(cReH)intoluene(bluesolidline,Vk’s),MeCN(reddashedline,DBU),andTHF(blackdottedline,TBD)ata1:1ratioofM:Band1atmofH2(SeeSISectionVIII).aAveragedover2trials,uncertaintybasedonstandarddeviation.bFreeenergyvaluesextrapolatedfromΔΗ˚1andΔS˚1.

Having identified appropriate conditions for equilibrium H2cleavage,thetemperaturedependenceofhydrideformationineachsolventwasinvestigated.Teflon-sealedNMRtubeswerechargedwithdicarbonylcomplex2andtheappropriatebaseinthesolventofinterest,filledwith1atmH2,andtheequilibriumconstantsweredeterminedbyvariable-temperature(VT)NMRspectroscopy.ThetemperaturerangeswereselectedbasedontheabilitytodetectallspeciesbyNMRspectroscopyandbythesolventboilingpoints.Figure2showsthevan’tHoffplots(Figure2B)resultingfromthevariabletemperaturestudies,fromwhichtheenthalpyofreaction(ΔΗ˚1)andentropyofreaction(ΔS˚1)forH2heterolysiswereobtainedineachsolvent(Figure2A).EachH2splittingreactionisexothermic,ΔΗ˚1between–17and–27kcal·mol–1.TheΔS˚1valuesarestrikinglylargeandnegative,reflectinganentropicpenaltyassociatedwith conversionof threeneutralspecies to two solvated ions. The reaction becomes moreentropically unfavorable as the solvent polarity decreases,suggestingthattheionsaremoretightlypaired(morehighlyordered)inthelesspolarsolvents.The enthalpy and entropy terms for this outer sphere H2splittingreactioncanbecomparedtoothersystemsthatcleaveH2.Outer sphereH2 splittingby themaingroup “frustrated”Lewis pair B(C6F5)3/P(tBu)3 has ΔΗ˚1 = –31.4 kcal·mol–1 inbromobenzene.36 This value is similar to the enthalpy of H2splitting by2 and Vk’s in toluene (ΔΗ˚1 = –26.7 kcal·mol–1).Metal-ligandcooperativeexamples include1,2-additionofH2acrossFe–NorFe–Bbonds(ΔΗ˚1≈–9kcal·mol–1intoluene11orbenzene9)and1,3-additionofH2in(PNP*)RuH(CO)(ΔΗ˚1=–17.4 kcal·mol–1 in THF; PNP* = 2-(tBu2PCH2)-6-(tBu2PCH)-C5H3N).10TheentropyofH2splittingby2isfarmoreunfavorablethananycomparablesystemsreportedintheliterature.The∆S˚1of2variesfrom–63.3cal·mol–1·K–1inMeCNto–76.0cal·mol–1·K–1 in toluene. For comparison, the entropy of metal-ligandcooperative1,3-additionofH2isreported(ΔS˚1=–45cal·mol–1·K–1 in THF10) and the 1,2-addition examples are lessunfavorable (ΔS˚1 = –10 cal·mol–1·K–1 in toluene11 or –28cal·mol–1·K–1 benzene9). Even the termolecular “frustrated”Lewispair,forwhichonlyacomputationally-derivedestimateisavailable,onlyhasanentropicpenaltyof–56cal·mol–1·K–1inbromobenzene.12The highly unfavorable entropy term associated with H2splittingby2leadstoastrongtemperaturedependenceintheformation of hydride complex3. Furthermore, because∆H˚1and ∆S˚1 vary across the three solvents studied, distincttemperature effects can be expected for each solvent. Tovisualize the influence of temperature and solvent on theconcentrationofhydride3 present in solution, thevaluesofΔΗ˚1 and ΔS˚1 were used (assuming temperatureindependence) tocalculate themole fractionof3 (cReH) as afunctionoftemperatureat1atmofH2anda1:1ratioofM:B(seeSISectionVIIIforderivation).AsshowninFigure2Candobserved experimentally in VTNMR specotroscopic studies,hydride3isdominantatlowtemperature,whiledicarbonyl2dominatesathightemperature.Thisindicatesthatifthekineticbarriersremainsurmountable,catalysismayproceedfasteratlower temperatures. Further changes in reaction conditionssuchasbaseidentity(pKa),ratioofbase:metalconcentrations,

orH2pressureallimpactthecReHplot(seeSIsectionVIIIformoredetails).Solvent Effects on Thermodynamic Hydricity, ΔG˚H–. Wenext turned our attention to understanding thethermodynamics of the hydride transfer step. As shown inScheme 1C above, the favorability of hydride transfer isdictatedbythedifferenceinhydricitybetweenformateandthehydride donor of interest. The hydricity of3was thereforeexperimentally determined. The H2 heterolysis method ofScheme4isbuilton∆G1,sothethermodynamicmeasurementsoftheprecedingsectioncouldbeuseddirectlyinconstructingathermochemicalcycleforhydricity.Scheme4.DeterminationofhydricitythroughheterolysisofH2.12

MeCNwasexaminedfirsttoenablevaluablecomparisonswithother hydride complexes.12,37,38 Using ∆G˚1 (˚ denotes thestandardstatetemperatureof298K)fromFigure2A,thefreeenergyofprotonation,∆G˚2,andtheH2heterolysisconstantinMeCN,∆G˚H2(76.0kcal·mol–1),12 thethermochemicalcycleofScheme4wasusedtodeterminethehydricityof3inMeCNat25˚C(∆G˚H–=40.6±1.0kcal·mol–1,Table1).Table1.Hydricity(∆G˚H–)of3inTHFandMeCNat298K.

Solvent Base(pKa)39 ΔG˚H–(kcal·mol–1)

THF TBD(21.0) 37.6±1.0MeCN DBU(24.3) 40.6±1.0

Rheniumhydride3isapotenthydridedonorinMeCN:hydridetransferfrom3toCO2isthermodynamicallyfavoredbyca.3kcal·mol–1 (∆G˚H– ca. 44 kcal·mol–1 for HCO2–).12 Althoughhydricity studies of other rhenium hydride complexes arelacking,3isca.5-7kcal·mol–1morehydridicthantheneutralMntricarbonylhydrideswithsubstitutedbipyridineligands.40Hydride3issimilarlyhydridictotheanionicgroup6hydride[W(CO)5H]–.12Thehydricityof3wasalsodeterminedinTHF.Using∆G˚1fromFigure2A,alongwiththerecentlyreportedvalueof∆G˚H2(68.7kcal·mol–1),14 the hydricity in THF at 25 °Cwas determined(∆G˚H–=37.6±1.0kcal·mol–1,Table1).HydricitiesinTHFhavenotbeenreporteduntilrecently.10,13ComparingthehydricityofhydridesinTHF(seeTableS1intheSI),complex3ismorehydridicthanneutral(PNP)Ru(H)2(CO)(∆G°H– inTHF=44.6kcal·mol–1; PNP = 2,6-bis(tBu2PCH2)-C5H3N), and falls in therangeofanionicbimetalliccobalthydrides.10,13Onthebasisofan estimated hydricity of formate in THF and the knowncatalytic activity of anionic bimetallic cobalt hydrides,10,13,41complex 3 should also be sufficiently hydridic to produceformatefromhydridetransfertoCO2inTHF.Toconfirmthishypothesis,3wasgeneratedinsituinTHF-d8byadditionofLiHBEt3to2andsubsequentlyplacedunderanatmosphereofCO2 (1atm).Completeconversionof3 to thefive-coordinate complex2,accompaniedby the formationoffreeformate(d8.22)wasobservedby1HNMRspectroscopy(seeSIFigureS19-S20).Thelackofformatebindingsimplifies

the situation, as the thermodynamicsof ligandbindingneednotbe considered.Considering thathydride complex3 isanoctahedral complexwith avalenceelectron count of 18andthat formate does not coordinate after hydride transfer, anouter sphere hydride ion transfer is highly likely.A seminalstudyoffac-Re(bpy)(CO)3Hwasalsoconsistentwithanouterspherepathway.42The lack of a defined pKa scale in toluene and availablethermochemical data to determine ∆G˚H2 precluded thedeterminationofhydricityintoluene.TemperatureEffectsonThermodynamicHydricity,ΔGH–.To date, thermodynamic hydricity values have only beendeterminedat298K.Determiningthetemperature-dependenthydricity (∆GH–) requires knowledge of the temperaturedependence of outer sphere H2 splitting, the acidity of theexternal acid/base pair, and the heterolysis of free H2. Wedevelopedathermochemicalmethodologythatsumsthefreeenergyofeachreaction inScheme4atagiventemperature,focusing on MeCN solvent based on the availability ofthermodynamicparameters.ThefreeenergyofH2heterolysis,∆G1,wasavailablefrom293to320KfromVTNMRequilibriumstudies(Figure2).ThepKaofDBUwasassumedtobetemperature-independent,basedonstudiesofnitrogenheterocyclesinMeCN.43,44AlthoughthepKachangeislikelynegligible,∆G2decreasesby3kcal·mol–1asthetemperature increases over the studied range (SI Table S3).The temperature-dependent free energy of H2 heterolysis,∆GH2,wasestimatedbasedonthefreeenergyofH+reductiontoH–inwateratagiventemperature(basedontemperature-dependent reduction potential data),45 coupled with the iontransferfreeenergyofH+andH–fromwatertoMeCNatthesame given temperature (see SI Section VII for derivation).∆GH2 increasesasa functionof temperature(seeSITableS3andSIFigureS30).Summing∆G1,∆G2,and∆GH2atagiventemperatureenabledthedetermination of the hydricity of complex 2 at thattemperature.Hydride3becomesmorehydridic(smaller∆GH–values) with increasing temperature over the experimentalrange of 293 to 320 K in MeCN (Figure 3, red triangles).Enthalpy and free energymeasurements of organic hydridedonors and a few transition hydrides at 298 K predict adecreasein∆GH–withincreasingtemperature.12,46,47Thenewthermochemicalmethodologyenabled the first experimentalvalidation of this prediction and provide a quantitativeestimateofthetemperaturedependenceofhydricity.To gain insight into how temperature influences thethermodynamicsofhydridetransferfrom3toCO2,whichisakeystepinCO2hydrogenation,thetemperaturedependenceofthehydricityofformate isneeded.Wethereforeadditionallydeveloped a thermochemical cycle to determine ∆GH– forformate in MeCN (see SI Section VII for derivation). Thehydricityofformatewasfoundtobesignificantlylesssensitivetotemperature(Figure3,bluesquares).Figure3showsthathydridetransferfrom3toCO2isthermodynamicallyfavorableoverthefulltemperaturerangestudiedinMeCN.Assumingalineartemperaturedependence,thedatacanbeextrapolatedto predict that this favorability will be maintained above –12°C.HydridetransfertoCO2becomesmoreexergonicasthetemperatureincreasesinMeCN.

Figure 3.Plot of ∆GH– of3 (red triangles) and HCO2– (bluesquares)aswellas∆∆GH– forhydridetransfer from3toCO2(purplearrows).Hydrogenation of CO2 to Formate: Thermochemistry-Guided Development of Catalytic Conditions. The H2splitting and hydride transfer reactivity observed in thepreceding thermochemical studies constitute the individualstepsofCO2hydrogenationcatalysisviaan“outersphere”typemechanism (Scheme 1 above). The two-step outer spheremechanism would facilitate direct connections betweenthermodynamicparameters and thekineticsof catalysis.WethereforecarriedoutcatalyticCO2hydrogenationstudiesoverawiderangeoftemperaturesinthreesolvents.CatalyticCO2hydrogenation activity was assessed using the turnoverfrequency(TOF,definedasmolesofformatedividedbymolesofcatalystandthereactiontime),basedoninitialrates(<10%conversion) of CO2 hydrogenation under 1 atm 1:1 H2:CO2.Conditionsalignedwiththethermodynamicstudieswheneverpossible to facilitate comparisons with experimentalthermodynamic data. Additionally, using the weakest basespossibleminimizesexcessdrivingforceinthereaction;akeyparameterinliquidfuelsynthesisorH2storage.1,6,48In THF, dicarbonyl complex 2 and 100 equivalents oftBuP1(pyrr)3 (pyrr = pyrrolidinyl, pKa = 20.3 in THF49) wereallowedtoreactwithH2andCO2at0°C,25°C,and50°C.ThebasetBuP1(pyrr)3waschosenasamoresolublealternativetoTBD with a similar pKa value. As shown in Figure 4A, thecatalyticactivityexhibitsastrikingtemperaturedependence:thereactionbecomesca.200-fold fasterasthetemperature isdecreased(TOF=24±6h–1at0˚Cand0.13±0.06h–1at50˚C).Thealmostcompletelossofcatalyticactivityat50°Cisinlinewiththepredictedlowconcentrationof3(cReH=0.10at50˚C;Figure4B)resultingfromuphillH2splittingthermodynamics(DG1=5.4kcal·mol–1).At0°C,whereahighconcentrationof3ispredicted(cReH=0.76)andH2splittingisalmostergoneutral(DG1 = 0.8 kcal·mol–1), catalysis can proceed— the kineticbarriers remain surmountable at this temperature. Catalystactivity varied linearlywith the change in favorability of H2

O

O

PtBu2

PtBu2

Re COCO

H

3

+ CO2

O

O

PtBu2

PtBu2

Re COCO

+ [HCO2]–∆∆GH–

splitting,DG1,asshowninFig4C,suggestingthatH2splittingisinfluencing the rate limiting step (see below for furtherdetails). While faster net hydride ion transfer at lowertemperatureshasbeenobservedfororganohydridedonors,weare not aware of any such examples in catalytic CO2hydrogenation.50,51If H2 splitting is involved in the turnover-limiting step(s), astrongerbaseshouldacceleratethereaction.UsingVk’sbase(estimatedpKaof26.6inTHF)52inTHFwouldshiftDG1at50˚C from+5.4kcal·mol–1(with tBuP1(pyrr)3)to–3.9kcal·mol–1,increasing cReH from 0.1 to ca. 1.0. As predicted, CO2hydrogenationat50˚CinTHFproceededfasterwithVk’sbasewhencomparedwithtBuP1(pyrr)3(increasedfrom0.13±0.06h–1to324±8h–1).In MeCN containing 100 equiv DBU, catalyst 2 exhibits theoppositetemperaturedependence.Thereactionisslowat0˚C(TOF = 9 ± 1 h–1), and the rate increases slightly as thetemperatureincreasesto50 C(TOF=29±2h–1,Figure4).WeproposethatH2splittingbecomestheturnover-limitingstepinMeCN, due to solvent binding (vide supra) that inhibits H2binding. Consistent with slow H2 association, the use of astrongerbase(Vk’s,pKa=33.553;8.2pKaunitsstrongerthanDBU)didnotchangetheinitialTOFat50˚C(29±2h–1withDBUversus32±8h–1withVk’s).Correlationswith∆∆GH–areconsistentwithpre-equilibriumcontributionstorate,butweproposethattheTOFtemperaturedependence isdominatedbyH2binding(seebelowforfurtherdetails).Intoluenecontaining100equivVk’sbase,catalyst2operateswithhigherTOFvaluesincomparisontoTHForMeCNatalltemperatures examined. This broadly correlates withpredictions, given that H2 splitting was most favorable intolueneandthehighestconcentrationofhydride3ispredictedin this solvent.However,adistinct temperaturedependenceprofile was observed in toluene. A maximum in activity isobservedat50 ˚C(TOF=364±18h–1),withslowerratesatboth lower and higher temperatures (Figure 4). DuringinvestigationsofthecatalyststabilitybyNMRspectroscopy,nodecompositionwasobservedafter24hat25°Candonly13%conversiontotricarbonyl1after11hat50°C(SIFigureS43).

ThissuggeststhatcatalystdecompositionisunlikelyaffectingtheinitialTOFreactivity(collectedwithin5minutes inmostcases).The initial rates TOF of CO2 hydrogenation in toluene atambienttemperatureof245h–1atonly0.5atmeachofH2andCO2 compares favorably with other CO2-to-formatehydrogenationcatalysts(TableS7-8intheSI).Increasingfrom1atmto20atm1:1CO2:H2leadstoa14-foldincreaseinrate,TOF=3330±340h–1(seeSISectionIX).Runningcatalysistofullconversionat25 Cand1atmof1:1H2:CO2intoluenewith0.1mol%2and1,000equivVk’sresultedinaturnovernumber(TON)of1023±27(102±3%yield).ThefinalTOF(227±6h–1)waswithinerroroftheTOFvaluefrominitialrates(245±30h–1), consistent with sustained activity of 2 over extendedreactiontimesandrulingoutanysignificantproductinhibition.AGeneralModel forUnderstanding theRamificationsofTemperature on Outer Sphere CO2 Hydrogenation

Reactions. Studying both the thermodynamics of individualstepsandthekineticsofcatalysisinmultiplesolventsallowsusto develop a general mechanistic understanding of howtemperatureaffectsoutersphereCO2hydrogenationcatalysts.In this section, we introduce representative reactioncoordinate diagrams (RCDs) and discuss the influence oftemperatureontheactivationbarriersandindividualreactionfree energies.The insights should begenerally applicable toothercatalyststhatfollowanouterspheremechanism.Figure6showstwoRCDsforoutersphereCO2hydrogenationtoformate.Onthebasisofourreactivityandthermodynamicstudies,weadoptamodelwhere∆G1isendergonicand∆∆GH–isexergonic.Theoverallreaction,∆Grxn,mustbeexergonicinorder to proceed to completion. In this experimentallyinformed model, there are two limiting regimes: when thehighest single-step barrierheight is associatedwith hydrideiontransfer(∆G‡HIT),orwhenitisassociatedwithH2splitting(∆G‡H2).Figure6Ashowsthecaseofalargerhydrideiontransfer(HIT)barrier,∆G‡HIT.With∆G1>0kcal·mol–1,TOFwilldependonboth ∆G1 and ∆G‡HIT.54 As the temperature decreases, thebarrier height ∆G‡HIT will increase, as almost universally

Figure4.A)InitialratesofCO2hydrogenationtoformateinMeCN(redcheckeredbars,withDBU),THF(blackdiagonalbars,withtBuP1(pyrr)3)andtoluene(bluesolidbars,withVk’s)using2,100equivbase,and1:1mixofH2:CO2at1atm.aB)Molefractionof3(cReH)intoluene(bluesolidline,Vk’s),MeCN(reddashedline,DBU),andTHF(blackdottedline,tBuP1(pyrr)3ata1:100ratioofM:Band0.5atmofH2.C)PlotofinitialratesTOFversusDG1intoluene(bluesquaresusingVk’s),MeCN(redcirclesusingDBU),andTHF(black triangles for tBuP1(pyrr)3;black star forVk’s).DG1wasdeterminedusing thermodynamicparameters inFigure2AandcorrectingforbasepKaandstoichiometryandH2pressure.

observed for elementary reactions.Because outer sphereH2splitting to form a metal hydride features a large negativeentropy of reaction, ∆G1 will decrease dramatically withdecreasingtemperature.Forcatalyst2inTHF(andintolueneabove50˚C),thedecreasein∆G1islargerthantheincreasein∆G‡HIT,andthustheTOFincreaseswithdecreasingtemperature.Knowing that higher temperatures will not always lead tofastercatalysisisabroadlyimportantfinding—onethatmayruncountertoexpectations,butisreadilyexplainedbythepre-equilibriummodeloftheRCD.Figure6BshowsthecaseofalargerH2splittingbarrier,∆G‡H2.If ∆∆GH– is large and negative, the TOFwill depend only on∆G‡H2 and decreasing temperature will decrease the rate. If∆∆GH– is small, however, this thermodynamicparameterwillalsoaffecttheequilibriumconcentrationofMduringturnoverand thus influenceTOF.For catalyst2 inMeCN, equilibriumsolventbindingtothecatalystisproposedtoresultinalarge∆G‡H2barrier(likelycomposedofbothsolventdissociationandH2 splitting) as shown in Figure 6B. Strong correlationsbetweenTOFof2and∆∆GH–suggestapre-equilibriumregimein MeCN. The TOF depends on ∆∆GH– and ∆G‡H2, but themagnitude and temperature dependence of ∆G‡H2 dominatesuch that TOF decreases with decreasing temperature.QuantitativeRCDsinMeCNsupportthismodel(FigureS45intheSI).

Figure 6. Idealized reaction coordinate diagrams for outer-sphere CO2 hydrogenation highlighting limiting cases of: (A)high-barrier hydride ion transfer and (B) high-barrier H2splitting.

Forfuturecatalysisstudies,simplifiedqualitativeRCDssuchasthose inFigure6canenablepredictionsabouttheexpectedtemperaturedependenceofthehydrogenationTOF.Infact,theobservation of inverse temperature effects (higher TOF atlower temperature) can be taken as a strong mechanisticindicator:onlypre-equilibriumH2splittingfollowedbyhigh-barrier hydride transfer is expected to follow such atemperature dependence in outer sphere hydrogenationreactions.Knowledgeof theRCDcanalsoenablepredictionsaboutwhenchangingH2pressure,CO2pressure,orbasepKamightinfluenceTOF.ImplicationsofThermochemicalDataonCatalystDesign.Inthissection,wecomparetheouterspheremechanismwiththe metal–ligand cooperation mechanism as they relate tocatalystdesignprinciplesinCO2hydrogenation.ThemoststrikingdifferencesareintheentropyofH2splitting.The entropy associated with metal–ligand cooperative H2heterolysis has been determined for two hydrogenationcatalysts.The1,2-additionofH2to(iPrPNP)FeH(CO)intoluenehasa smallnegativeentropy term(ΔS˚1=–9.7 cal·mol–1·K–1,ΔH˚1 = –7.8 kcal/mol; iPrPNP = N(CH2CH2PiPr2)2).11 The 1,3-additionofH2to(PNP*)RuH(CO)inTHFhasasomewhatlargernegativeentropyterm(ΔS˚1=–45cal·mol–1·K–1,ΔH˚1=–17.4kcal/mol).10Thenewlyreportedthermodynamicdatarevealsamuch largerentropicpenalty forouter sphereH2 splitting,∆S˚1 of –63.3 to–76.0 cal·mol–1·K–1.Theentropicdifferencesare reflected in distinct temperature-dependent speciationprofiles forcatalysts followingthetwodifferentmechanisms(Figure7).ThechangeincMHasafunctionoftemperatureismuchmoregradualformetal–ligandcooperativeH2splittingthan for termolecular outer sphere H2 heterolysis. Thissuggests that catalysts following metal–ligand cooperationmechanismswillbelesslikelytoshowrateenhancementswithdecreasing temperature, as thebarrierheights could changemore than ∆G1 as a function of temperature. Furthermore,while the temperature-dependence of metal–ligandcooperative H2 splitting is dictated by the structure of thecatalyst, the temperature-dependence of outer sphere H2splittingwilldependontheboththecatalyststructureandthechoiceofbase.

Figure7.Molefractionofhydride(cMH)at1:1M:Bratioand1atmofH2for3intoluene(redsolid,Vk’s),3inTHF(bluesolid,TBD), (H-iPrPNP)FeH2(CO) in toluene (red dashed), and(PNP)Ru(H)2(CO)inTHF(bluedashed).

ΔG‡H2

ΔG‡HIT

ΔΔGH–

lower T

ΔG‡H2

ΔG‡HIT

ΔΔGH–

[HBase][MH] + CO2

M +H2 + CO2 + Base

lower T

lower T

lower T

lower TA. Higher TOF at lower TTOF ∝ ΔG1 + ΔG‡

HIT

B. Lower TOF at lower TTOF ∝ ΔG‡

H2 (∆∆GH– large and negative)

TOF ∝ ΔG‡H2 + ΔΔGH– (∆∆GH– small)

ΔGrxn

[HBase][MH] + CO2

M + [HBase][HCO2]

M +H2 + CO2 + Base

ΔG1

lower TΔG1

lower T

ΔGrxn

M + [HBase][HCO2]

lower T

The distinct temperature-dependent thermochemistry ofcatalyststhatoperateviaouterspheremechanismsandthosethat operate via metal–ligand cooperation has not beenrecognizedpreviously.The starkdifferences in temperature-dependent behavior suggest may help chemists identifyappropriate catalysts to suit the target application. Forexample, an outer sphere catalyst that undergoes speciationchangesoveraverynarrowtemperaturerange(tunablebythechoice of base) might be ideal for thermally reversible H2storage.1,6,48

3.CONCLUSIONS

AnewrheniumcatalystforCO2hydrogenationtoformateviaanouterspheremechanismwasthesubjectofadetailedstudyconnectingthethermodynamicsofindividualH2splittingandhydridetransferstepswiththekineticsofcatalysis.The coordinatively unsaturated rhenium complex 2heterolytically cleaves H2 with an exogenous base to formhydride complex3. The first experimental determination ofenthalpyandentropyofanoutersphereH2splittingreactionofthiskindrevealedanextremelylargeandnegativeentropyof reaction that varies systematically across three organicsolvents.This finding is expected tobe representativeof theentropy parameters for themany other catalysts known tooperateviaH2splittingwithexternalbases.TheotherstepintheouterspheremechanismishydrideiontransferfromthemetalhydrideintermediatetoCO2,producingthe formate anion. The first hydricitymeasurements of anyrheniumhydriderevealedastronghydridedonorcapableofreducing CO2 to formate. Further, we have introduced newthermochemicalmethodologytoenablethedetermination of temperature-dependent hydricity values. Thepresentrheniumhydridecomplexbecomesmorehydridicathigher temperatures. Coupled to the distinct temperaturedependenceofthehydricityofformate,wefoundthathydridetransfer from the rhenium complex to CO2 becomingincreasinglyfavorablewithincreasingtemperature.By connecting thermodynamic parameters with catalyticactivity, broadly applicable reaction coordinate diagrammodelscouldbeconstructedthatexplainthetemperature-andsolvent-dependentreactivity.Thehighlyunusualtemperatureeffectwherebyhydrogenationacceleratesasthetemperaturedecreases is attributed to the large entropy associated withtermolecular H2 splitting in a pre-equilibrium precedinghydride ion transfer. This analysis not only shows howthermochemicalstudiescanassistincatalystdevelopment,butsheds light on the salient differences between termolecularsystems that undergo base-assisted H2 heterolysis, andbimolecularsystemsthatfeaturemetal–ligandcooperation.SupportingInformationExperimentaldetailsandcharacterizationdata(PDF)Crystallographicdata(CIF)

Corresponding Author *A.J.M.M.Email:ajmm@email.unc.edu

Notes ‡Theseco-authorscontributedequally.

ACKNOWLEDGMENT Thisworkwas supportedby theU.S.DepartmentofEnergy,OfficeofScience,OfficeofBasicEnergySciences,underAwardNo. DE-SC0014255. J.H was supported by the J. ThurmanFreezescholarshipfund.Q.J.B.acknowledgessupportfromtheNSFGraduateResearchFellowshipProgram(DGE-1650116)andtheUNCDissertationCompletionFellowshipprogram.TheauthorsthankAndrewCampandMarcterHorstforassistancewithNMRspectroscopyexperiments.ThemassspectrometryworkwassupportedbytheNationalScienceFoundationunderGrant No. CHE-1726291. The NMR spectroscopy work wassupportedbytheNationalScienceFoundationunderGrantNo.CHE-1828183.REFERENCES(1) Sordakis,K.;Tang,C.;Vogt,L.K.;Junge,H.;Dyson,P.J.;

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S1

Supporting Information

Temperature and Solvent Effects on H2 Splitting and Hydricity: Ramifications on CO2

Hydrogenation by a Rhenium Pincer Catalyst

Jenny Hu,‡ Quinton J. Bruch,‡ Alexander J. M. Miller*

Department of Chemistry, University of North Carolina at Chapel Hill, Chapel Hill, North Carolina 27599–3290, United States

*Corresponding Author Email Address: A.J.M.M.: ajmm@email.unc.edu

S2

Table of Contents

I. General Considerations S3

II. Synthetic Procedures S5

III. NMR Spectra of New Compounds S8

IV. Infrared Spectra of New Compounds S14

V. Hydricity Determination S17

VI. Thermodynamics of H2 Heterolysis S22

VII. Hydricity as a Function of Temperature S29

VIII. Impact of H2 Splitting Thermodynamics on cReH S37

IX. Catalytic Hydrogenation of CO2 S48

X. Crystallographic Details S57

XI. References S66

S3

I. General Considerations

All manipulations were carried out under an N2 atmosphere using standard glovebox and

Schlenk techniques. Under standard glovebox operating conditions pentane, diethyl ether (Et2O),

benzene, toluene, and tetrahydrofuran (THF) were used without purging, such that traces of those

solvents were present in the atmosphere and in the solvent bottles. 1H, 31P{1H}, and 13C{1H} NMR

spectra were recorded on 400 or 600 MHz spectrometers at 298 K unless otherwise specified in

the University of North Carolina at Chapel Hill Department of Chemistry NMR Core Laboratory.

NMR solvents were purchased from Cambridge Isotopes Laboratories, Inc. Benzene-d6 (C6D6),

acetonitrile-d3 (CD3CN), and toluene-d8 (tol-d8) were freeze−pump−thaw degassed three times

before drying by passage through a small column of activated alumina and storage over 3 Å

molecular sieves. Tetrahydrofuran-d8 (THF-d8) was purchased in ampoules and dried by passage

over two small columns of activated alumina. THF-d8 was passed over an additional short column

of activated alumina immediately before use. 1H and 13C chemical shifts are reported relative to

residual protio solvent resonances.1 31P chemical shifts are absolutely referenced to the 1H solvent

residual and reported versus phosphoric acid (0 ppm).2 All other reagents were commercially

available and used without further purification. High-resolution mass spectrometry (HRMS) was

performed in the University of North Carolina at Chapel Hill Department of Chemistry Mass

Spectrometry Core Laboratory on a Q Exactive HF-X (ThermoFisher, Bremen, Germany) mass

spectrometer. Samples (in CH2Cl2) were introduced via a microelectrospray source at a flow rate

of 3 µL/min. Xcalibur (ThermoFisher, Bremen, Germany) was used to analyze the data. Molecular

formula assignments were determined with Molecular Formula Calculator (v 1.2.3). Infrared

spectroscopy was carried out with a Thermo Scientific Nicolet iS5 FT-IR equipped with an iD1

Transmission Accessory (Thermo Scientific) for solution measurements in a demountable liquid

S4

cell with CaF2 windows (0.05 mm pathlength) (Pike Technologies Inc.). Single-crystal X-ray

diffraction (XRD) data were collected in the University of North Carolina at Chapel Hill

Department of Chemistry X-Ray Core Laboratory on a Bruker APEX-II CCD diffractometer at

100 K with Cu Kα radiation (l = 1.54175 Å). See Section X for crystallographic details. Elemental

analysis was performed at Robertson Microlit Laboratories (Ledgewood, NJ). High pressure

catalysis was performed with a Parr Series 5000 Multiple Reactor System operated by a Parr 4871

Process Controller. Each of the six reactors was individually pressure and temperature controlled

and was monitored by computer software SpecView 32.

S5

II. Synthetic Procedures

Synthesis of tBuPOCOP-H. The synthesis was adapted from a literature procedure.3 In a N2-filled

glovebox, a 20 mL scintillation vial was charged with resorcinol (200 mg, 1.8 mmol), 1,8-

diazabicyclo[5.4.0]undec-7-ene (DBU, 542 µL, 3.6 mmol), and 5 mL of THF. The milky

suspension was stirred while adding PtBu2Cl (692 µL, 3.6 mmol), resulting in immediate formation

of large chunks of a white insoluble material. The suspension was stirred for 24 h, then filtered

over a glass frit to remove white solids, which were washed with 3 ×3 mL portions of THF before

the filtrate was collected and dried under vacuum. The crude material was extracted with pentane

(3 × 1 mL) and filtered, and the solvent was removed from the filtrate under vacuum to provide

tBuPOCOP-H as a colorless oil that solidified overnight (400 mg, 55% yield, 99% purity by 1H

NMR spectroscopy). 1H and 31P NMR features matched with previously reported data.

Synthesis of (tBuPOCOP)Re(CO)3 (1). In a N2-filled glovebox, tBuPOCOP-H (202.6 mg, 0.507

mmol) was dissolved in 5 mL of chlorobenzene and Re(CO)5(Cl) (208.4 mg, 0.576 mmol) was

suspended in 5 mL of chlorobenzene. Both were transferred to a Schlenk flask and Et3N (0.157

mL, 1.13 mmol) was added. The reaction mixture was refluxed for 16 h, during which time the

Re(CO)5(Cl) solid dissolved, yielding a pale-yellow solution. The solution was dried under

vacuum without heating, leaving a sticky off-white solid. In the glovebox, the solid was extracted

with pentane (2 × 5 mL) and filtered. The pale-yellow filtrate was dried under vacuum, giving a

light-yellow microcrystalline solid. The solid was dissolved in minimal pentane and placed in the

freezer overnight, resulting in precipitation of a crystalline solid. The mother liquor was decanted,

the volume was reduced by half, and placed back in the freezer for a second recrystallization. The

crystals resulting from both crops were combined and dried under vacuum to give 204.0 mg of

S6

colorless 1 (60% yield, >99% pure by multinuclear NMR spectroscopy and elemental analysis).

Single crystal X-ray diffraction studies were run on crystals isolated from the mother liquor prior

to drying. 1H NMR (600 MHz, C6D6): δ 6.93 (t, 7.9 Hz, 1H, p-CH), 6.80 (d, 7.8 Hz, 2H, m-CH),

1.29 (m, 36H, C(CH3)). 31P{1H} NMR (162 MHz, C6D6): δ 188.94 (s). 13C{1H} (151 MHz, C6D6):

δ 200.83 (t, 8.8 Hz, Re(CO)), 200.14 (s, Re(CO)), 167.41 (t, 6.5 Hz, aryl-C), 130.77 (t, 7.3 Hz,

aryl-C), 127.65 (s, aryl-C), 105.71 (t, 4.8 Hz, aryl-C), 43.03 (t, 10.4 Hz, C(CH3)), 29.33 (t, 2.9 Hz,

C(CH3)). FTIR (THF, cm–1): nCO 2023, 1923, 1902. Anal. Calcd for C25H39O5P2Re: C, 44.97; H,

5.89; N, 0.00. Found: C, 45.06; H, 5.72; N, <0.02.

Synthesis of (tBuPOCOP)Re(CO)2 (2). In a N2-filled glovebox, 1 (49.5 mg, 0.074 mmol) was

dissolved in 5 mL of THF and lithium triethylborohydride (371 µL of a 1 M solution in THF, 0.371

mmol) was added by syringe. The solution was transferred to a Teflon-sealed pressure vessel and

heated at 70 oC overnight. The excess CO was then removed through five freeze-pump-thaw cycles.

In the glovebox, the clear solution was stirred while adding HCl etherate (290 µL of a 2 M solution

in diethyl ether, 0.580 mmol). The solution immediately turned dark brown and effervesced. After

stirring for 10 minutes, the solution was dried under vacuum. The product was extracted with

pentane (3 × 2 mL), filtered, and dried under vacuum, yielding 32.6 mg of a dark brown solid (69%

yield, >99% by multinuclear NMR spectroscopy and elemental analysis). NMR spectra and

elemental analysis are reported for the isolated powder. Single crystal X-ray diffraction data was

collected on crystals grown via slow evaporation of a pentane solution. 1H NMR (600 MHz, C6D6):

δ 7.13 (t, 7.9 Hz, 1H, p-CH), 6.88 (d, 7.9 Hz, 2H, m-CH), 1.21 (m, 36H, C(CH3)). 31P{1H} NMR

(243 MHz, C6D6): δ 210.76 (s). 13C{1H} (151 MHz, C6D6): δ 211.65 (t, 4.0 Hz, Re(CO)), 172.72

(t, 8.5 Hz, aryl-C), 166.98 (t, 7.6 Hz, aryl-C), 134.93 (s, aryl-C), 105.13 (t, 5.1 Hz, aryl-C), 41.98

S7

(t, 10.1 Hz, C(CH3)), 28.33 (t, 5.1 Hz, C(CH3)). FTIR (THF, cm–1): nCO 1910, 1840. FTIR (MeCN,

cm–1): nCO 1912, 1834. FTIR (toluene, cm–1): nCO 1913, 1840. Anal. Calcd for C24H39O4P2Re: C,

45.06; H, 6.15; N, 0.00. Found: C, 45.34; H, 6.11; N, <0.10.

Synthesis of [Na][(tBuPOCOP)Re(CO)2H] (3). In a N2-filled glovebox, 2 (20 mg, 0.031 mmol)

was dissolved in 5 mL toluene. Sodium triethylborohydride (34.5 µL of a 1 M solution in toluene,

0.0345 mmol) was added followed by 15 mL of pentane and allowed to sit for 16 h, during which

time a white precipitate formed and the solution color changed from dark brown to pale yellow.

The solution was decanted and the remaining white solids were triturated with pentane (3 × 3 mL).

Afterwards, the remaining white solid was dried under vacuum for one minute and collected (9.8

mg, 47% yield, >99% pure by multinuclear NMR spectroscopy). Complex 3 was observed to react

readily with trace moisture and decomposed upon prolonged exposure to vacuum, so elemental

analysis could not be performed. 1H NMR (400 MHz, THF-d8): δ 6.37 (t, 7.6 Hz, 1H, p-CH), δ

6.16 (d, 7.7 Hz, 2H, m-CH), δ 1.40 (m, 18H, C(CH3)), δ 1.30 (m, 18H, C(CH3)), δ –6.13 (t, 25.8

Hz, 1H, ReH). 31P{1H} NMR (161.92 MHz, THF-d8): δ 201.37 (s). 13C{1H} (151 MHz, THF-d8):

δ 211.83 (s, Re(CO)), 210.86 (t, 6.0 Hz, Re(CO)), 167.07 (d, 8.8 Hz, aryl-C), 142.96 (t, 8.5 Hz,

aryl-C), 121.82 (m, aryl-C), 102.36 (t, 4.8 Hz, aryl-C), 42.23 (t, 7.8 Hz, C(CH3)), 41.42 (t, 10.1

Hz, C(CH3)), 30.17 (t, 3.5 Hz, C(CH3)), 29.58 (t, 2.7 Hz, C(CH3)). FTIR (THF, cm–1): nCO 1873,

1754, 1737. HRMS (ESI–) m/z [(tBuPOCOP)Re(CO)2(H)]– calcd for C24H40O4P2Re 641.19593;

found 641.19936.

S8

III. NMR Spectra of New Compounds

Figure S1. 1H NMR spectrum of (tBuPOCOP)Re(CO)3 (1) (600 MHz, C6D6).

Figure S2. 31P{1H} NMR spectrum of (tBuPOCOP)Re(CO)3 (1) (162 MHz, C6D6).

S9

Figure S3. 13C{1H} NMR spectrum of (tBuPOCOP)Re(CO)3 (1) (150 MHz, C6D6).

Figure S4. 1H NMR spectrum of (tBuPOCOP)Re(CO)2 (2) (600 MHz, C6D6).

S10

Figure S5. 31P{1H} NMR spectrum of (tBuPOCOP)Re(CO)2 (2) (243 MHz, C6D6).

Figure S6. 31P{1H} NMR spectrum of (tBuPOCOP)Re(CO)2 (2) in toluene-d8 (top), CD3CN (middle), and THF-d8 (bottom), and under N2 (151 MHz).

S11

Figure S7. 13C{1H} NMR spectrum of (tBuPOCOP)Re(CO)2 (2) (150 MHz, C6D6).

Figure S8. 1H NMR spectrum of [Na][(tBuPOCOP)Re(CO)2(H)] (3) (400 MHz, THF-d8).

tol tol

S12

Figure S9. 31P{1H} NMR spectrum of [Na][(tBuPOCOP)Re(CO)2(H)] (3) (162 MHz, THF-d8).

Figure S10. 13C{1H} NMR spectrum of [Na][(tBuPOCOP)Re(CO)2(H)] (3) (150 MHz, THF-d8).

THF-d8 THF-d8

tol

S13

Figure S11. 31P{1H} NMR stack of in situ conversion of 1 to 3 using 5 equivalents of LiHBEt3 (1 M solution in THF). Stack shows 31P NMR spectra before (top) and after (bottom) heating at 70 ˚C (400 MHz, THF-d8).

t = 0min

t = 4 days

S14

IV. Infrared Spectra of New Compounds

Figure S12. FTIR spectrum of (tBuPOCOP)Re(CO)3 (1) in THF.

Figure S13. FTIR spectrum of (tBuPOCOP)Re(CO)2 (2) in THF.

S15

Figure S14. FTIR spectrum of (tBuPOCOP)Re(CO)2 (2) in MeCN.

Figure S15. FTIR spectrum of (tBuPOCOP)Re(CO)2 (2) in toluene.

S16

Figure S16. FTIR spectrum of [Na][(tBuPOCOP)Re(CO)2(H)] (3, black) in THF. Due to the high water-sensitivity of 3, some hydrolysis to 2 was observed. The FTIR spectrum of 2 (red) is included for comparison.

S17

V. Hydricity Determination Preparation of samples for hydricity measurements. In a N2-filled glovebox, 2 (5 mg, 0.008

mmol) and the selected base were dissolved in 0.5 mL deuterated solvent in a Teflon sealed NMR

tube. In CD3CN, 6 or 13 equiv of DBU was used; in in THF-d8, 1 or 13 equiv TBD (1,5,7-

triazabicyclo[4.4.0]dec-5-ene) was used. Mesitylene in either CD3CN or THF-d8 (10 µL of a 90

mM stock solution; 0.0009 mmol) was injected as an internal standard. The tube was then

subjected to three freeze-pump-thaw cycles before being backfilled with 1 atm of H2 gas at 298 K.

The sample was monitored by 1H NMR spectroscopy every 24 h until equilibrium was reached.

Similar Keq values were obtained for each of the two base concentrations utilized, confirming that

a true equilibrium of H2 splitting was measured.

Calculation of hydricity. Concentrations of 2, 3, and base were determined by integrating against

the mesitylene internal standard and an equilibrium constant was determined. The hydricity was

calculated using the equations below.4 Note that Eq S1 is the reverse of the H2 splitting reaction,

which we define as ∆Gº1. Uncertainties presented here are based on the standard deviation of

experimentally determined Keq values; in the main text the hydricity values are reported with an

uncertainty of ± 1 kcal·mol–1 due to underlying assumptions.4

MH– + HB+ ⇄ H2 + M + B 1.364log(Keq) = –∆G˚1 ; 𝐾%& =[)*+][*-.]/01[)][-]

Eq. S1

H+ + B ⇄ HB+ –1.364pKa = ∆G˚2 Eq. S2 H2 ⇄ H+ + H– ∆G˚H2 Eq. S3

MH– ⇄ M + H– ∆G˚H– Eq. S4 In MeCN, the average equilibrium constant, Keq, was determined to be 0.026 ± 0.009 atm–1, the

pKa of DBU in MeCN is 24.33,5 and the H2 splitting constant, DG°H2 is 76.0.4 Hydricity is thus:

D𝐺°*+ = 1.364 ∗ log=𝐾%&> − 1.364 ∗ 𝑝𝐾A + D𝐺°*1 D𝐺°*+ = 1.364 ∗ log(0.026) − 1.364 ∗ 24.33 + 76.0

D𝐺°*+ = 40.6 ± 0.2 kcal·mol–1

S18

Figure S17. 1H NMR spectrum of equilibrium between 2, 3, and DBU under 1 atm of H2 at 298 K (400 MHz, CD3CN). Calculated ΔG˚H– = 40.6 ± 0.2 kcal·mol–1 in MeCN.

mesitylene

DBU

3 2

H2

S19

Figure S18. 1H NMR of equilibrium between 2, 3, and TBD under 1 atm of H2 at 298 K (400 MHz, THF-d8). Calculated ΔG˚H– = 37.6 ± 0.8 kcal·mol–1 in THF.

Hydride transfer from 3 to CO2. In a N2-filled glovebox, 3 was generated in situ by dissolving

2 (11 mg, 0.017 mmol) in THF-d8 and adding LiHBEt3 (15.5 µL of a 1 M solution in THF, 0.0155

mmol, 0.9 equiv). Upon addition of LiHBEt3, the solution color immediately changed from brown

to pale orange, indicative of consumption of 2 and formation of 3. The solution was transferred to

a Teflon-sealed NMR tube and placed under 1 atm of CO2. Upon mixing the CO2 headspace and

the solution, the solution color immediately changed from pale orange to brown. 1H NMR

spectroscopy of the solution showed a prominent singlet at 8.22 ppm, indicative of formate

formation. 31P{1H} NMR spectroscopy showed no evidence for 3, instead only showing 2 and <3 %

of 1, indicating that hydride transfer from 3 to CO2 proceeds to completion and that formate

binding does not occur, as no new species were observed.

TBD

3 2 H2

mesitylene

S20

Figure S19. 1H NMR spectrum after addition of CO2 to in situ generated of 3 (400 MHz, THF-d8).

Figure S20. 31P{1H} NMR spectrum after addition of CO2 to in situ generated of 3 (162 MHz, THF-d8).

THF

THF

2

2

2 HCO2–

2

S21

Table S1. Comparison of known transition metal hydride hydricities in THF.6

Hydride Donor ∆GºH–(THF) (kcal/mol) Reference

54.4 7

40.5 8

35.4 8

44.6 9

37.6 This work

S22

VI. Thermodynamics of H2 Heterolysis

Sample preparation for H2 binding. In an N2-filled glovebox, 2 was dissolved in toluene-d8 and

transferred to a Teflon-sealed NMR tube. The tube was subjected to three freeze-pump-thaw cycles

before being backfilled with 1 atm of H2 gas at 298 K and allowed to equilibrate for at least 24 h

before starting variable temperature NMR studies.

Variable Temperature NMR Spectroscopy for H2 Binding. Quantitative 1H NMR spectra were

collected in 20 °C increments spanning a temperature range of 198 to 298 K. The sample was

allowed to equilibrate for 4 minutes at temperature before each spectrum was collected. A 2.0 Hz

line broadening was applied to all spectra. Over the course of the experiment, the observed pseudo-

C2v symmetry of 2 by 1H NMR spectroscopy did not change. Similarly, the chemical shift of 2 by

31P{1H} NMR spectroscopy did not change as a function of temperature.

Figure S21. 1H NMR spectroscopy of 2 under 1 atm of H2 in toluene at temperatures increasing from –75 °C to 25 °C (500 MHz, toluene-d8).

tol tol H2

2 2

T = –75 °C

T = –55 °C

T = –35 °C

T = –15 °C

T = 5 °C

T = 25 °C

S23

Figure S22. 31P{1H} NMR spectroscopy of 2 under 1 atm of H2 in toluene at temperatures increasing from –75 °C to 25 °C (202 MHz, toluene-d8).

Sample preparation for H2 Heterolysis. Samples in THF and MeCN were prepared in the same

manner as in Section V. Samples in toluene-d8 were prepared analogously, except that 1 or 4 equiv

Verkade’s base (Vk’s, 2,8,9-triisobutyl-2,5,8,9-tetraaza-1-phosphabicyclo[3.3.3]undecane) was

used.

Variable Temperature NMR Spectroscopy for H2 Heterolysis. Quantitative 1H NMR spectra

were collected in 10 °C increments, with the sample allowed to equilibrate for 4 minutes before

collection of each spectrum. Once a maximum or minimum temperature was reached, the probe

temperature was increased (or decreased) by 5 oC, before collecting in 10 °C increments. This

ensured that no temperature-dependent hysteresis was observed between warming and cooling the

2

T = –75 °C

T = –55 °C

T = –35 °C

T = –15 °C

T = 5 °C

T = 25 °C

S24

probe. Two samples were investigated in each solvent and the average enthalpy and entropy are

reported. A 2.0 Hz line broadening was applied to all spectra.

In MeCN the temperature range studied was 293 to 333 K. Over this temperature range,

both the equilibrium concentration of 3 changed and the 1H NMR resonances of 3 broadened

linearly as a function of temperature. At temperatures above 320 K, the hydride resonance of 3

was too broad to observe. The 1H NMR resonances of conjugate hydride acceptor 2 remained sharp,

suggesting that the broadness of 3 stems from interactions with DBU or the conjugate acid

[HDBU]+. When aromatic 1H NMR resonances of 3 were used to determine the equilibrium

concentrations, the entropy and enthalpy of H2 splitting were within error of those derived from

hydride integration and reported in the main text.

In THF the temperature range studied was 273 to 323 K, while in toluene a temperature

range of 333 to 373 K was studied. In both THF and toluene the 1H NMR resonances of 3 were

sharp at all temperatures examined. At temperatures below 333 K in toluene, only hydride species

3 is observed and only dicarbonyl species 2 was observed in THF at temperatures above 323 K.

During variations in temperature, the pressure in the headspace of the Teflon-sealed NMR

tube was calculated to vary approximately 0.1 atm per 30 °C according to the ideal gas law. To

ensure the experimental impact of changing H2 pressure over the course of the variable temperature

NMR studies was minimal, solutions of 2 and TBD (15 equivalents) in THF-d8 were filled with

H2 at 0 and 50 ˚C. Keq was measured via 1H NMR spectroscopy at 0, 25, and 45 ˚C and in all cases,

Keq was within error of values determined during previous runs (Figure S29).

S25

Figure S23. 1H NMR spectra of equilibrium between 2, 3, and DBU under 1 atm of H2 in CD3CN at temperatures increasing from 20 oC to 60 oC (400 MHz, CD3CN).

Figure S24. 1H NMR spectra of equilibrium between 2, 3, and Vk’s under 1 atm of H2 in toluene-d8 at temperatures increasing from 60 oC to 100 oC (400 MHz, toluene-d8).

S26

Figure S25. 1H NMR spectra of equilibrium between 2, 3, and TBD under 1 atm of H2 in THF-d8 at temperatures increasing from 0 oC to 50 oC (400 MHz, THF-d8). Construction of van ’t Hoff plots. Concentrations of 2, 3, and base were determined by

integrating against the mesitylene internal standard and used to determine Keq for the H2 heterolysis

reaction of Eq. S5 at each temperature. van ’t Hoff plots were constructed by plotting the inverse

of temperature against ln(Keq) at each point. ΔΗ˚1 and ΔS˚1 values for H2 heterolysis were then be

determined by the slope and intercept, respectively. Figures S26-28 below show the van ‘t Hoff

plot from a single VT NMR experiment in each solvent.

H2 + M + B⇄ MH– + HB+ –RTln(Keq) = ∆G˚1 ; 𝐾%& =[)*+][*-.]/01[)][-]

Eq. S5

S27

Figure S26. van ’t Hoff plot of H2 heterolysis based on the equilibrium between 2, 3, and DBU in CD3CN.

Figure S27. van ’t Hoff plot of H2 heterolysis based on the equilibrium between 2, 3, and TBD in THF-d8.

S28

Figure S28. van ’t Hoff plot of H2 heterolysis based on the equilibrium between 2, 3, and Vk’s in toluene-d8.

Figure S29. Comparison of H2 filling methods on H2 splitting Keq values of Eq S5. Teflon-sealed NMR tubes were filled with 1 atm H2 at 0 ˚C (blue), 25 ˚C (black), 50 ˚C (red) and then Keq was determined at three NMR probe temperatures (0, 25, and 45 C, 400 MHz, THF-d8). The Keq values are within error of each other regardless of the temperature of H2 backfilling.

S29

VII. Hydricity as a Function of Temperature This section derives the thermochemical parameters needed to estimate the thermodynamic

hydricity of [(POCOP)Re(CO)2H]– as a function of temperature. The approach utilizes the “H2

heterolysis” method for hydricity determination, but with each free energy determined at the

desired temperature (rather than the standard state temperature of 298 K), as outlined in equations

S6-S9. Each parameter is discussed independently in the following sections. Note that equations

S6-S8 are solving for free energies away from standard states, and therefore we have included RT

in the equations instead of multiplying by a factor of 1.364 (which is specific to 298 K).

MH– + HB+ ⇄ H2 + M + B –∆G1 = –RTln(Keq) Eq. S6 H+ + B ⇄ HB+ ∆G2 = –2.303·RT·pKa Eq. S7 H2 ⇄ H+ + H– ∆GH2 = ∆HºH2 – T∆SºH2 Eq. S8

MH– ⇄ M + H– ∆GH– Eq. S9 Free energy of H2 splitting by the Re complex, ∆G1. The value of ∆G1 at various temperatures

was obtained experimentally, as described in the main text.

Free energy of proton dissociation, ∆G2. The pKa was assumed to be constant. This assumption

is supported by prior experimental studies of acidity temperature dependence of cationic N-

heterocycles in acetonitrile, which exhibited nearly negligible changes in pKa (ca. 0.1 pKa unit per

10 K).10 Even when the pKa is constant, however, free energy ∆G2 does change with temperature

according to eq S7.

Free energy of H2 heterolysis, ∆GH2. The ∆GH2 value was determined from the enthalpy and

entropy associated with the transfer of ions from water to acetonitrile, following the general

approach used for ∆GºH2 at 298 K.4,11 Here, the free energy for the reactions of Eqs S10-S12 was

determined at the desired temperature in acetonitrile and the summed to give ∆GH2 at that

temperature.

S30

H2(g) ⇄ H+(aq) + H–(aq) ∆GH2 = –46.12·[(–1.055) + ((T – 298)·(–0.00148))] Eq. S10 H+(aq) ⇄ H+(MeCN) ∆Gtr(H+) = ∆Hºtr(H+) – (T·∆Sºtr(H+)·0.001) Eq. S11 H–(aq) ⇄ H–(MeCN) ∆Gtr(H–) = ∆Hºtr(H–) – (T·∆Sºtr(H–)·0.001) Eq. S12 H2(g) ⇄ H+(MeCN) + H–(MeCN) Eq. S13 Bratsch reported the temperature dependence of the potential for the reduction of H+(aq) to H–(aq),

–1.48 mV/K;12 eq S10 experiences the same shift because the reduction of H+(aq) to H2(g) is

temperature independent (defined as 0 V vs NHE at all temperatures). The reduction potentials for

H+(aq) to H•(aq) (–2.29 V vs NHE) and H•(aq) to H–(aq) (+0.18 V vs NHE) were chosen to align

with the original hydricity scale in acetonitrile developed by DuBois.6 Eq S11, the free energy to

transfer H+ from water to acetonitrile, can be determined from the enthalpy and entropy

components. The H+ transfer enthalpy from water to MeCN was reported by Hefter and Marcus,13

and can be used with the H+ transfer free energy at 298 K to calculate the H+ transfer entropy from

water to MeCN (Table S2). The transfer free energy at any temperature can then be calculated

according to eq S11.

Table S2. Thermodynamic and physical properties of proton, halide, and hydride ions.a

Ion 1/r ∆Gsolv,aq in kJ/mol

∆Hsolv,aq in kJ/mol (kcal/mol)

∆Gtr in kJ/mol

(kcal/mol)

∆Htr in kJ/mol

(kcal/mol)

∆Str in J/mol·K

(cal/mol·K)

∆Gsolv,MeCN in kJ/mol (kcal/mol)

∆Hsolv,MeCN in kJ/mol (kcal/mol)

H+ 1104 59.0 41.1 –60.1

I– 0.4545 243 247 19.6 –9.3 –88 262.6 237.7 Br– 0.5102 278 289 33.0 7.6 –89 311 296.6

Cl– 0.5525 304 320 42.1 18.4 –87 346.1 338.4

F– 0.7519 429 464 47.2 476.2

H– 0.7246 402 444.5 (106.2)

57.9 (13.8)

71.8 (17.2)

46.7 (11.2)

459.9 (109.9)

516.3 (123.4)

aValues in bold were determined by interpolation see additional SI documents. Values in italics determined from enthalpy of reaction and free energy of reaction at 298 K. Because the free H– ion is unstable for prolonged periods in solution, its properties are estimated.

The currently adopted method for estimating solvation properties is based on extrapolations based

on the ionic radii and hydration energies of the halides.4,14 Prior derivations utilized the free energy

S31

of hydration and the ion transfer free energies at 298 K. The halide ion transfer enthalpies from

Hefter and Marcus,13 and the halide ion solvation enthalpies from Jenkins,15 were used to estimate

the hydride ion transfer enthalpy. With the hydride ion transfer enthalpy and the transfer free

energy at 298 K, the hydride ion transfer entropy can be calculated, eqs S14–S16.

H–(g) ⇄ H–(MeCN) ∆Hºsolv,MeCN Eq. S14 H–(aq) ⇄ H–(g) –∆Hºsolv,H2O Eq. S15 H–(aq) ⇄ H–(MeCN) ∆Hºtr Eq. S16 Plotting the enthalpy of solvation of the halides in water against the inverse of the ionic radius

enabled interpolation of the enthalpy of solvation of H– in water, ∆Hºsolv,H2O = 106.2 kcal·mol–1.

From the known ion transfer enthalpies from water to acetonitrile of Cl–, Br–, and I–, the enthalpy

of solvation of the halides in acetonitrile can be obtained (Table S2). Plotting enthalpy of solvation

of halides in MeCN vs the inverse of the ionic radius provides the enthalpy of solvation of H– in

MeCN, ∆Hºsolv,MeCN = 123.4 kcal·mol–1, which in turn can be related to the ion transfer enthalpy

from water to MeCN (∆Hºtr = 17.2 kcal·mol–1). With the hydride ion transfer enthalpy and transfer

free energy at 298 K, the hydride ion transfer entropy can be calculated according to ∆G = ∆Hº –

T∆Sº: ∆Sºtr = 11.15 cal·mol–1·K–1. The free energy of H2 heterolysis in acetonitrile (Eq S13) can

be calculated at a given temperature by summing the free energies of the Eqs S10, S11, and S12

at the given temperature. An Excel spreadsheet calculator is supplied as Supporting Information.

Hydricity calculation. The hydricity at a given temperature can be calculated according to Eqs

S6-S9, above. The value of ∆G1 (Eq S6) comes from variable temperature H2 splitting equilibrium

experiments. The pKa (Eq S7) is assumed to be constant, so the standard 298 K value in acetonitrile

is used. Finally, the value of ∆GH2 (Eq S8) at the given temperature is calculated as described

above, and the Eqs S6-S8 are summed to give the hydricity. As discussed and tabulated in the main

text, ∆GH– of [(tBuPOCOP)Re(CO)2H]– in MeCN decreases as temperature increases (~ – 0.1

S32

kcal·mol–1·K–1). This is shown in Table S3, wherein values for DG1 (in MeCN, using DBU as the

base, pKa = 24.316), DG2, and ∆GH2 are shown alongside ∆GH–.

Table S3. Impact of temperature on DG1, DGH2, and DGH─ in MeCN assuming a temperature-independent pKa.

Temp (K) ∆G1 (kcal/mol)

∆G2 (kcal/mol)

∆GH2 (kcal/mol)

∆GH─ (kcal/mol)

293.13 1.84 ± 0.03 –32.6 76.2 41.8 297.22 2.17 ± 0.21 –33.1 76.5 41.3 301.29 2.30 ± 0.18 –33.5 76.8 41.0 306.16 2.76 ± 0.26 –34.0 77.2 40.4 310.36 2.80 ± 0.23 –34.5 77.5 40.2 315.19 3.31 ± 0.36 –35.0 77.8 39.4 319.71 3.55 ± 0.22 –35.6 78.1 39.0

Figure S30. Plot of ∆∆G1 (∆∆G1 = ∆G1 – ∆G°1, made using experimental data in MeCN using DBU as base; red squares), ∆∆G2 (∆∆G2 = ∆G2 – ∆G°2; green squares), ∆∆GH2 (∆∆G2 = ∆GH2 – ∆G°H2; blue squares), and ∆∆GH– (∆∆GH– = ∆GH– – ∆G°H–; black squares) as a function of temperature. Free energy of formic acid and formate formation in water. This section derives the

thermochemical parameters to determine the free energy of forming formic acid or formate in

water.

S33

CO2(g) + H2(g) ⇄ HCO2H(g) ∆G17 = ∆Hºrxn – T•∆Sºrxn Eq. S17

HCO2H(g) ⇄ HCO2H(aq) ∆Ghyd(HCO2H) = ∆Hºhyd(HCO2H) – T•∆Sºhyd(HCO2H) Eq. S18

CO2(g) + H2(g) ⇄ HCO2H(aq) ∆G19(aq) of (HCO2H) Eq. S19

The free energy of equation S17 was determined using the gas-phase enthalpy of formation and

gas-phase molar entropy of formic acid (for CO2(g), ∆Hfº = –94.1 kcal/mol, Sº = 51.1 cal·mol–1·K–

1; for H2(g), ∆Hfº = 0 kcal/mol, S˚ = 31.2 cal·mol–1·K–1; for HCO2H(g) ∆Hfº = –90.5 kcal/mol, Sº

= 59.4 cal·mol–1·K–1).17 For equation S18, the free energy of hydration of gas-phase formic acid

was determined based on the enthalpy and entropy parameters. The free energy at 298 K

(∆Gºhyd(HCO2H) = –4.94 kcal/mol)18 and the enthalpy of hydration (∆H˚hyd(HCO2H) = –11.3

kcal/mol)19 were used to determine the entropy of hydration (∆S˚hyd(HCO2H) = –21.3 cal·mol–1·K–

1). The sum of ∆G17 and ∆Ghyd gives ∆G19(aq) of (HCO2H) as a function of temperature (Eq. S19). A

common strategy to make the net hydrogenation of CO2 to formic acid be thermodynamically

favorably is to deprotonate formic acid and form the formate anion. The free energy of this

subsequent step is dictated by the following:

HCO2H(aq) ⇄ H+(aq) + HCO2–(aq) ∆Gion = 2.303•RT•pKa@T of HCO2H Eq. S20

B(aq) + H+(aq) ⇄ HB+(aq) ∆Gion = –2.303•RT•pKa@T of B Eq. S21

CO2(g) + H2(g) ⇄ H+(aq) + HCO2–(aq) ∆Gdeprot = –2.303•RT•DpKa@T Eq. S22

𝛥𝑝𝐾A@𝑇 = p𝐾A@TofB– p𝐾A@TofHCO2H

For equation S20, the free energy of proton dissociation from formic acid requires that the pKa at

temperature, pKa@T, be known. In water it has been shown that the pKa of formic acid, 3.75, is

essentially invariant with temperatures up to at least 100 ˚C.20,21 The second equation, S21;

involves the temperature-dependent pKa of the added base. Combining equations S20 and S21

gives the free energy of formic acid deprotonation, ∆Gdeprot, as a function of DpKa and as a function

of temperature.

S34

Free energy of formate hydricity, ∆GH– of HCO2–. This section derives the thermochemical

parameters needed to estimate the thermodynamic hydricity of formate as a function of

temperature in both water and acetonitrile. The following equations were used to determine the

temperature dependent hydricity in water:

CO2(g) + H2(g) ⇄ HCO2H(g) ∆G17 = ∆Hºrxn – T•∆Sºrxn Eq. S17

HCO2H(g) ⇄ HCO2H(aq) ∆Ghyd(HCO2H) = ∆Hºhyd(HCO2H) – T•∆Sºhyd(HCO2H) Eq. S18

HCO2H(aq) ⇄ H+(aq) + HCO2–(aq) ∆Gion = 2.303•RT•pKa@T Eq. S20

CO2(g) + H2(g) ⇄ H+(aq) + HCO2–(aq) ∆GH–(aq) of HCO2

– + H+ Eq. S23

H+(aq) + HCO2–(aq) ⇄ CO2(g) + H2(g) –∆GH–(aq) of HCO2

– + H+ Eq. S24

H2(g) ⇄ H+(aq) + H–(aq) ∆GH2 = –46.12·[(–0.74) + ((T – 298)·(–0.00148))] Eq. S25

HCO2–(aq) ⇄CO2(g) + H–(aq) ∆GH–(aq) of HCO2

– Eq S26

The free energy of equation S17, S18, and S20 were determined above. ∆G17, ∆Ghyd, and ∆Gion

could then be determined at a given temperature T and summed to give ∆GH–(aq) of HCO2– + H+ at

T (equation S23). The hydricity of formate can then be calculated from Equation S24 (the reverse

of equation S23) and Equation S25 (from the temperature dependence for the reduction of H+(aq)

to H–(aq),12 using the reduction potentials of H+(aq) to H•(aq) (Eº = –2.29 V vs NHE) and H•(aq)

to H–(aq) (Eº = 0.81 V vs NHE) proposed by Rosseinsky and adopted by Appel in developing the

aqueous hydricity scale.22,23 Summation of equations S24 and S25 gives ∆GH–(aq) of HCO2–. The

hydricity of aqueous formate at 298 K is calculated to be 23.6 kcal/mol, which is very similar to

the accepted literature value of 24.1 kcal/mol.4

Having determined the temperature-dependent hydricity of formate in water, we estimated the

temperature dependency of hydricity in acetonitrile using the following equations:

S35

H+(aq) + HCO2–(aq) ⇄ CO2(g) + H2(g) –∆GH–(aq) of HCO2

– Eq. S27

H2(g) ⇄ H+(aq) + H–(aq) ∆GH2 = –46.12·[(–1.055) + ((T – 298)·(–0.00148))] Eq. S28

HCO2–(MeCN) ⇄ HCO2

–(aq) –∆Gtr(HCO2–) = –∆Hºtr(HCO2

–) + T•∆Sºtr(HCO2–) Eq. S29

H–(aq) ⇄ H–(MeCN) ∆Gtr(H–) = ∆Hºtr(H–) – (T·∆Sºtr(H–)·0.001) Eq. S12

CO2(g) + H2(g) ⇄ H+(aq) + HCO2–

(aq) ∆GH–(MeCN) of HCO2– Eq. S30

It is important to note that different values for the reduction potential of H•/H–(aq) underpin the

aqueous and MeCN hydricity scales. Equation S25 uses the reduction potential of +0.81 V vs NHE

for the aqueous scale, while equation S28 uses the reduction potential of +0.18 V vs NHE for the

acetonitrile scale.4 The transfer free energy for formate anion from water to acetonitrile at 298 K

was recently reported (∆Gºtr(HCO2–) = 10.8 kcal/mol).24 Because neither the transfer enthalpy

(∆Hºtr(HCO2–)) nor the transfer entropy (∆Sºtr(HCO2–)) of formate anion from water to acetonitrile

could be located in the literature, we assumed formate and trifluoroacetate to have similar transfer

enthalpies (∆Hºtr(HCO2–) ≈ ∆Hºtr(CF3CO2–) = 3.2 kcal/mol).13 This enthalpy was combined with

∆Gºtr(HCO2–) at 298 K to obtain ∆Sºtr(HCO2–) = 25.4 cal·mol–1·K–1. As a check, this entropy value

was compared to that obtained from the transfer enthalpy of formate anion from water into 60/40

acetonitrile/water mixtures ((∆Htr,60%(HCO2–) = 1.9 kcal/mol),13 which gives ∆Sºtr(HCO2–) = 29.6

cal·mol–1·K–1. The entropies generated by these different approximations differ by only 4 cal·mol–

1·K–1. Equation S12 was derived above, and the values of ∆Hºtr(H–) and ∆Sºtr(H–) can be found in

Table S2. To determine the hydricity of formate in acetonitrile at a given temperature, the free

energies of equations S27, S28, S29, and S12 at that temperature are summed to afford ∆GH–

(MeCN) of HCO2–. At 298 K, the estimated of ∆GH–(MeCN) of HCO2–, was estimated to be 41.2

kcal/mol. This value is in reasonable agreement with the conventionally accepted value of formate

hydricity in acetonitrile (44 kcal/mol).4 To maintain consistency with the literature, we added

S36

correction factor of +2.84 kcal/mol to the value of ∆GH–(MeCN) of HCO2– determined at any

temperature.

S37

VIII. Impact of H2 Splitting Thermodynamics on cReH As DG1 changes as a function of temperature or solvent, the equilibrium concentration of 3 can

vary dramatically. 3 is the active hydride species in the hydrogenation of CO2 and its equilibrium

concentration may be related to observed trends in catalytic activity. To explore this potential

relationship, we sought to express the temperature-dependence on the equilibrium concentration

of 3 by determining the mole fraction of 3 (cReH) as a function of temperature. We did so by

utilizing the enthalpies and entropies derived in SI Section VI to predict Keq at a given temperature

and then converting Keq to cReH (see below for derivation). We also demonstrate in the sections

below how tuning reaction conditions (amount of base, pressure of H2, and base pKa) influence

cReH, so as to guide selection of conditions for screening catalyst activity

Derivation of cReH as a function of temperature.

A well-defined equilibrium expression for H2 splitting will have several restrictions relating

concentrations of various species, noted below.

D𝐺T =–𝑅𝑇 ∗ ln=𝐾%&> = D𝐻T– 𝑇D𝑆T

𝐾%& =[𝑀𝐻Z] ∗ [𝐻𝐵\]𝑃*1 ∗ [𝑀] ∗ [𝐵]

[𝑴𝑯Z] + [𝑴] = [𝑴]𝒕𝒐𝒕𝒂𝒍

[𝑴𝑯Z] = [𝑯𝑩\]; which is true for every point along the equilibrium continuum

[𝑯𝑩\] + [𝑩] = [𝑩]𝒕𝒐𝒕𝒂𝒍 and therefore [𝑩] = [𝑩]𝒕𝒐𝒕𝒂𝒍 −[𝑯𝑩\] = [𝑩]𝒕𝒐𝒕𝒂𝒍 − [𝑴𝑯Z]

In a given experiment, we know [M]total and [B]total, and can set these as the initial concentrations

of [M] and [B]. Given Keq for a particular temperature derived from the thermodynamic parameters,

we can solve for the [MH_] at a given equilibrium as shown below.

S38

[M] [B] [HB+] [MH–] Initial Concentration [M]total [B]total 0 0

Change in Concentration -x -x +x +x Equilibrium Concentration [M]total - x [B]total - x x x

Plugging this into Keq:

𝐾%& =[𝑀𝐻Z] ∗ [𝐻𝐵\]𝑃*1 ∗ [𝑀] ∗ [𝐵]

𝑃*1 ∗ 𝐾%& =(𝑥) ∗ (𝑥)

([M]ghgAi − x) ∗ ([B]total − x)

𝑃*1 ∗ 𝐾%& =m1

m1Z([-]nonpq\[r]nonpq)m\([-]nonpq∗[r]nonpq)

𝑃*1 ∗ 𝐾%& ∗ 𝑥s − (𝑃*1 ∗ 𝐾%&) ∗ =[𝐵]ghgAi + [M]𝑡𝑜𝑡𝑎𝑙>𝑥 + (𝑃*1 ∗ 𝐾%&) ∗ =[𝐵]ghgAi ∗ [M]𝑡𝑜𝑡𝑎𝑙> = 𝑥s

(𝑃*1 ∗ 𝐾%& − 1) ∗ 𝑥s − (𝑃*1 ∗ 𝐾%&) ∗ =[𝐵]ghgAi + [M]𝑡𝑜𝑡𝑎𝑙>𝑥 + (𝑃*1 ∗ 𝐾%&) ∗ =[𝐵]ghgAi ∗ [M]𝑡𝑜𝑡𝑎𝑙> = 0

Then solve the quadratic:

𝑥 = −𝑏 +√𝑏s − 4𝑎𝑐

2𝑎

• 𝒂 = 𝑷𝑯𝟐 ∗ 𝑲𝒆𝒒 − 𝟏

• 𝒃 = −(𝑷𝑯𝟐 ∗ 𝑲𝒆𝒒) ∗ =[𝑩]𝒕𝒐𝒕𝒂𝒍 + [𝐌]𝒕𝒐𝒕𝒂𝒍>

• 𝒄 = (𝑷𝑯𝟐 ∗ 𝑲𝒆𝒒) ∗ =[𝑩]𝒕𝒐𝒕𝒂𝒍 ∗ [𝐌]𝒕𝒐𝒕𝒂𝒍>

𝑥 =Z(/01∗���)∗=[-]nonpq\[M]𝑡𝑜𝑡𝑎𝑙>Z�(Z(/01∗���)∗=[-]nonpq\[M]𝑡𝑜𝑡𝑎𝑙>)1Z�=/01∗���ZT>∗(/01∗���)∗=[-]nonpq∗[M]𝑡𝑜𝑡𝑎𝑙>

s(/01∗���ZT)

Can then substitute for Keq to relate back to DG1

𝑥 =

Z�/01∗𝑒

−𝛥𝐺1𝑅𝑇 �∗=[-]nonpq\[M]𝑡𝑜𝑡𝑎𝑙>Z�(Z(/01∗𝑒

−𝛥𝐺1𝑅𝑇 )∗=[-]nonpq\[M]𝑡𝑜𝑡𝑎𝑙>)1Z��/01∗𝑒

−𝛥𝐺1𝑅𝑇 ZT�∗(/01∗𝑒

−𝛥𝐺1𝑅𝑇 )∗=[-]nonpq∗[M]𝑡𝑜𝑡𝑎𝑙>

s(/01∗𝑒−𝛥𝐺1𝑅𝑇 ZT)

At equilibrium, 𝑥 = [𝑀𝐻Z] and therefore:

S39

c𝑴𝑯+ =[𝑴𝑯Z][𝑴]𝒕𝒐𝒕𝒂𝒍

= 𝒙

[𝑴]𝒕𝒐𝒕𝒂𝒍

This then allows us to plot c)*+ as a function of temperature. Unless otherwise specified, the

following conditions were used:

• [𝑴𝑯Z] +[𝑴] = [𝑴]𝒕𝒐𝒕𝒂𝒍 = 𝟏

• [𝑯𝑩\] + [𝑩] = [𝑩]𝒕𝒐𝒕𝒂𝒍 = 𝟏

• 𝑷𝑯𝟐 = 𝟏

This simplifies the equation to:

c𝑴𝑯+ =[𝑴𝑯Z][𝑴]𝒕𝒐𝒕𝒂𝒍

= 𝒙𝟏 = 𝒙

𝑥 =Z�𝑒

−𝛥𝐺1𝑅𝑇 �∗(s)Z�(Z(𝑒

−𝛥𝐺1𝑅𝑇 )∗(s))1Z��𝑒

−𝛥𝐺1𝑅𝑇 ZT�∗(𝑒

−𝛥𝐺1𝑅𝑇 )∗(T)

s(𝑒−𝛥𝐺1𝑅𝑇 ZT)

Plotting cReH versus T. Having derived cReH as a function of temperature using the conditions

outlined above, plotting the data shows a sigmoidal curve. At low temperatures, cReH approaches

1 and is essentially independent of temperature. Conversely, at high temperatures cReH approaches

0 and again becomes independent of temperature.

S40

Figure S31. Plot of cReH as a function of temperature (T) in toluene (ΔH˚ = –26.7 kcal·mol–1; ΔS˚1 of –76.0 cal·mol–1·K–1; using Vk’s)

Figure S32. Plot of cReH as a function of temperature (T) in THF (ΔH˚ = –18.9 kcal·mol–1; ΔS˚1 of –72.0 cal·mol–1·K–1; using TBD).

S41

Figure S33. Plot of cReH as a function of temperature (T) in MeCN (ΔH˚ = –16.7 kcal·mol–1; ΔS˚1 of –63.3 cal·mol–1·K–1; using DBU).

Figure S34. Plots of cReH as a function of temperature (T) in all solvents.

Theoretical Impact of Thermodynamic Parameters on cMH.

Effect of enthalpy, ΔΗ˚1. The effect of changing the enthalpy of H2 splitting is shown in Figure

S32. As the enthalpy changes from –22 to –32 kcal·mol–1 and the entropy is kept constant at –76.0

cal·mol–1·K–1, the hydride remains the dominant species at higher temperatures, and the entire

sigmoidal curve simply shifts along the temperature axis.

S42

Figure S35. Plot of cMH as a function of T at varying ΔH˚1 and a constant ΔS˚1 of –76.0 cal·mol–1·K–1. Effect of entropy, ΔS˚1. The effect of changing the entropy of H2 splitting is shown in Figure S33.

As the entropy was varied from –50 to –70 cal·mol–1·K–1 at a constant enthalpy of –20 kcal·mol,

the temperature at which the hydride mole fraction begins to drop changes. Furthermore, as the

entropy becomes more negative, the transition temperature range narrows.

Figure S36. Plot of cMH as a function of T at varying ΔS˚1 and a constant ΔH˚1 of –20 kcal·mol–1.

S43

Effect of changing the pKa of the base. The effect of temperature-independent changes in pKa was

explored by correcting DG1 of H2 splitting with the energy of DDG2 (related to the DpKa). This is

based on the fact that the hydricity, DGH–, is a thermodynamic parameter. Any changes in the free

energy derived from the pKa of the base must be equally offset by changes in the free energy of H2

splitting. This was exemplified in Figure S37 using the thermodynamic data determined in THF

(ΔH˚1 = –18.9 kcal·mol–1, ΔS˚1 = –72.0 cal·mol–1·K–1, pKa = 21.0) and varying the pKa.

–D𝐺T +D𝐺s +D𝐺*s = D𝐺*– =–D𝐺T′ +D𝐺s′ +D𝐺*s

–D𝐺T +D𝐺s =–D𝐺T′ +D𝐺s′

D𝐺s −D𝐺s� =–D𝐺T� + D𝐺T

DD𝐺s = DD𝐺T = −2.303 ∗ 𝑅𝑇 ∗ D𝑝𝐾A

c𝑴𝑯+ =[𝑴𝑯Z][𝑴]𝒕𝒐𝒕𝒂𝒍

= 𝒙

[𝑴]𝒕𝒐𝒕𝒂𝒍

𝑥 =

Z�/01∗𝑒−𝛥𝐺1+DD𝐺2

𝑅𝑇 �∗([-]nonpq\[r]�����)Z�(Z(/01∗𝑒−𝛥𝐺1+DD𝐺2

𝑅𝑇 )∗([-]nonpq\[r]�����))1Z��/01∗𝑒−𝛥𝐺1+DD𝐺2

𝑅𝑇 ZT�∗(/01∗𝑒−𝛥𝐺1+DD𝐺2

𝑅𝑇 )∗([-]nonpq∗[r]�����)

s(/01∗𝑒−𝛥𝐺1+DD𝐺2

𝑅𝑇 ZT)

Under our standard conditions of:

[𝑀𝐻Z] +[𝑀] = [𝑀]ghgAi = 1 [𝐻𝐵\] + [𝐵] = [𝐵]ghgAi = 1

𝑃*s = 1 This simplifies the equation to:

c𝑴𝑯+ =[𝑴𝑯+][𝑴]𝒕𝒐𝒕𝒂𝒍

= 𝒙𝟏= 𝒙

𝑥 =Z�𝑒

−𝛥𝐺1+DD𝐺2𝑅𝑇 �∗(s)Z�(Z(𝑒

−𝛥𝐺1+DD𝐺2𝑅𝑇 )∗(s))1Z��𝑒

−𝛥𝐺1+DD𝐺2𝑅𝑇 ZT�∗(𝑒

−𝛥𝐺1+DD𝐺2𝑅𝑇 )∗(T)

s(𝑒−𝛥𝐺1+DD𝐺2

𝑅𝑇 ZT)

Varying the pKa at constant ΔH˚1 and ΔS˚1 results in a dramatic shift in the temperature-

independent region favoring formation of MH. This clearly shows why changing the acid identity

S44

can strongly influence catalyst hydrogenation activity. If the pKa of the conjugate acid is too low,

the mole fraction of hydride at even ambient temperatures becomes exceptionally low.

Figure S37. Plot of cMH as a function of temperature (T) while varying pKa. Table S4. Impact of temperature and pKa on the formation of MH determined using ΔH˚1 = –18.9 kcal·mol–1, ΔS˚1 = –72.0 cal·mol–1·K–1.

T (˚C) cMH; pKa = 18.0

cMH; pKa = 20.3

cMH; pKa = 21.0

cMH; pKa = 24.0

–50 0.44 0.92 0.96 1.00 0 0.02 0.18 0.33 0.94

50 0.001 0.01 0.03 0.51 100 0.0001 0.002 0.005 0.13 150 0.00003 0.0005 0.001 0.03

Effect of changing the metal:base ratio. Catalytically hydrogenations often utilize a large excess

of base. This directly influences Keq of H2 heterolysis while leaving the free energy of H2

heterolysis unchanged. This is because the speciation of MH in the presence of H2 with larger

equivalencies of base

S45

favoring formation of MH. As shown in Figure S38 (calculated using ΔH˚1 = –26.7 kcal·mol–1

and ΔS˚1 = –76.0 cal·mol–1·K–1), we find that increasing the relative concentration of base allows

for higher concentrations of MH even as temperatures increase. Additionally, Figure S38 shows

that the consumption of base during a catalytic reaction can have a marked impact on the formation

of MH (Table S5). Changing the M:B ratio causes a deviation from our standard conditions, and

so we used the following conditions:

• [𝑴𝑯Z] +[𝑴] = [𝑴]𝒕𝒐𝒕𝒂𝒍 = 𝟏

• 𝑷𝑯𝟐 = 𝟏

This simplifies the equation to:

c𝑴𝑯+ =[𝑴𝑯+][𝑴]𝒕𝒐𝒕𝒂𝒍

= 𝒙𝟏= 𝒙

𝑥 =Z�𝑒

−𝛥𝐺1𝑅𝑇 �∗([-]nonpq\T)Z�(Z(𝑒

−𝛥𝐺1𝑅𝑇 )∗([-]nonpq\T))1Z��𝑒

−𝛥𝐺1𝑅𝑇 ZT�∗(𝑒

−𝛥𝐺1𝑅𝑇 )∗([-]nonpq)

s(𝑒−𝛥𝐺1𝑅𝑇 ZT)

Figure S38. Plot of cMH as a function of T while varying equivalencies of base. Determined using ΔH˚1 = –26.7 kcal·mol–1 and ΔS˚1 = –76.0 cal·mol–1·K–1.

S46

Table S5. Impact of temperature and ratio of M:B on the formation of MH determined using ΔH˚1 = –26.7 kcal·mol–1 and ΔS˚1 = –76.0 cal·mol–1·K–1.

T (˚C) cMH; 1:1, M:B

cMH; 1:10, M:B

cMH; 1:100, M:B

cMH; 1:1000, M:B

0 1.00 1.00 1.00 1.00 50 1.00 1.00 1.00 1.00 100 0.25 0.62 0.92 0.99 150 0.04 0.12 0.32 0.69 200 0.007 0.02 0.07 0.20

Effect of changing the H2 pressure. High pressures of H2 are often used to increase the rate of

hydrogenation. This has a similar effect to varying the concentration of base due its direct influence

on Keq and is shown in Figure S39. In practice, the impact of varying H2 pressure will be more

muted than varying equivalencies of base due to the smaller variance in conditions (testing 1 to 50

atm is common in H2 pressure but sometimes thousands of excess equivalents of base are used).

Changing the H2 pressure causes a deviation from our standard conditions, and so we used the

following conditions:

• [𝑴𝑯Z] +[𝑴] = [𝑴]𝒕𝒐𝒕𝒂𝒍 = 𝟏

• [𝑯𝑩\] + [𝑩] = [𝑩]𝒕𝒐𝒕𝒂𝒍 = 𝟏

This simplifies the equation to:

c𝑴𝑯+ =[𝑴𝑯+][𝑴]𝒕𝒐𝒕𝒂𝒍

= 𝒙𝟏= 𝒙

𝑥 =Z�/01∗𝑒

−𝛥𝐺1𝑅𝑇 �∗(s)Z�(Z(/01∗𝑒

−𝛥𝐺1𝑅𝑇 )∗s)1Z��/01∗𝑒

−𝛥𝐺1𝑅𝑇 ZT�∗(/01∗𝑒

−𝛥𝐺1𝑅𝑇 )

s(/01∗𝑒−𝛥𝐺1𝑅𝑇 ZT)

S47

Figure S39. Plot of cReH as a function of T while varying H2 pressure. Determined using ΔH˚1 = –26.7 kcal·mol–1 and ΔS˚1 = –76.0 cal·mol–1·K–1. Table S6. Impact of temperature and H2 pressure on the formation of MH determined using ΔH˚1 = –26.7 kcal·mol–1 and ΔS˚1 = –76.0 cal·mol–1·K–1.

T (˚C) cMH; 1 atm

cMH; 5 atm

cMH; 25 atm

cMH; 50 atm

0 1.00 1.00 1.00 1.00 50 0.86 0.93 0.97 0.98 100 0.27 0.45 0.65 0.72 150 0.04 0.09 0.18 0.24 200 0.01 0.02 0.04 0.06

S48

IX. Catalytic Hydrogenation of CO2

Analysis of initial rates. In a N2-filled glovebox, a Teflon-sealed flask was charged with 400 µL

of a 1 mg/mL solution of 2 in THF, MeCN, or toluene (0.4 mg, 0.63 µmol) and 100 equiv (62.5

µmol) base. In THF, the base was tBuP1(pyrr)3 (pyrr = pyrrolidinyl); in MeCN, the base was DBU;

in toluene, Vk’s base was used. The solution was diluted with additional solvent to a total volume

of approximately 5 mL before sealing. The reaction flask was then subjected to three freeze-pump-

thaw cycles before coming to thermal equilibrium in an ice bath or temperature-controlled oil bath

for at least 5 minutes. The flask was then filled with 1:1 mixture of H2:CO2 at 1 atm and manually

agitated for 5 seconds. The flask was then left to stir open to a bubbler for a given reaction time

(t). The reaction time was chosen to ensure conversion remained under 10%. After the chosen

reaction time, the flask was sealed and frozen in liquid nitrogen. The headspace was removed

under vacuum and the solvent was removed in vacuo as the solution slowly thawed, yielding either

colorless oils or colorless solids depending on the base/solvent combination. To the residue was

added a 0.05 M or 0.005 M solution of NaOTs in D2O and quantitative NMR analysis was used to

determine the final concentration of the formate product.

S49

Figure S40. 1H NMR spectrum after CO2 hydrogenation in toluene at 100 ˚C for 2 minutes (400 MHz, D2O).

Figure S41. Plot of initial rates TOF in MeCN using DBU as the base as a function of ∆∆GH– (blue squares, ∆∆GH– = ∆GH–(MeCN) of HCO2– – ∆∆GH–(MeCN) of 3) or as a function of ∆Grxn (red triangles, ∆Grxn = ∆∆GH– + ∆G1) or using Vk’s as the base as a function of ∆∆GH–(blue star) or as a function of ∆Grxn (red circle).

TsO–

D2O

OCHO–

Vk’s

Vk’s

Vk’s

Vk’s

S50

Analysis of the effect of pressure on TOF. In a N2-filled glovebox, a Parr multi-reactor pressure

vessel (see details in General Considerations above) with a glass liner was charged with 400 µL

of a 0.1 mg/mL solution of 2 in toluene (0.04 mg, 0.063 µmol) and 10,000 equiv (625 µmol) Vk’s.

The solution was diluted with additional solvent to a total volume of approximately 5 mL and a

stir bar was added before sealing under N2 inside the glovebox. The pressure vessels were

connected to the reactor manifold and stirred at a rate of 250 rpm. The temperature in each vessel

was 25.0 ± 1.0 ˚C. The reactor manifold was purged with 1:1 H2:CO2 and then each reactor vessel

was pressurized to either 20 or 34 atm to initiate the reaction. After 10 minutes, the pressure was

vented and the resulting colorless solution was transferred to a Teflon-sealed flask, sealed, and

frozen in liquid nitrogen to cease the reaction. The flask was attached to a Schlenk line and the

headspace was removed under vacuum. The solution was allowed to thaw and the solvent was

removed in vacuo yielding colorless oils. To the residue was added a 0.018 M NaOTs in D2O and

quantitative NMR analysis was used to determine the final concentration of the formate product.

Maximum TON Determination. In a N2-filled glovebox, 400 µL of a 1 mg/mL solution of 2 in

toluene (0.4 mg, 0.63 µmol) was added to a Teflon-sealed flask followed by 1000 equiv of Vk's

base (added as 484 µL of a 442 mg/mL solution in toluene). The solution was diluted with

additional solvent to a total volume of approximately 5 mL before sealing. The reaction flask was

then subjected to three freeze-pump-thaw cycles on a Schlenk line before being allowed to

thermally equilibrate in a temperature-controlled oil bath for at least 5 minutes at 25 ˚C. The flask

was then filled with 1:1 mixture of H2:CO2 at 1 atm and manually agitated for 5 seconds. The flask

was then left to stir open to a bubbler for 270 minutes (based on the initial rates estimate). Upon

completion of the reaction, the flask was sealed and frozen in liquid nitrogen. The headspace was

S51

removed under vacuum and the solvent was removed in vacuo as the solution slowly thawed,

yielding an off-white solid. To the residue was added a 0.1 M solution of NaOTs in D2O and

quantitative NMR analysis was used to determine the final concentration of the formate product.

The average yield over two runs was 102% ± 3%, giving an average TON of 1023 ± 27 with an

average TOF of 227 ± 6 h–1.

Figure S42. 1H NMR spectrum after CO2 hydrogenation in toluene at 25 ˚C for 270 minutes (400 MHz, D2O). Catalyst Stability Tests. To test the stability of the Re complexes during catalysis, CO2

hydrogenation was undertaken in a Teflon-sealed NMR tube. A Teflon-sealed NMR tube was

charged with a toluene-d8 solution of 4.8 mg of 2 (7.5 µmol), 48.0 mg of Vk’s base (140 µmol, 19

equiv), and trimethoxybenzene as an internal standard. Initial NMR spectra were collected in the

absence of H2 and CO2. Afterwards, the brown solution was subjected to three freeze-pump-thaw

cycles before backfilling with 1:1 mixture of H2:CO2 at 298 K. Upon mixing, the brown solution

turned colorless. The solution was then left to react at room temperature for 24 h, during which

time 2 quantitatively converted to 3 with concomitant formation of several equivalents of formate.

TsO–

D2O OCHO–

Vk’s

Vk’s

Vk’s

Vk’s

Vk’s Vk’s

S52

The Teflon-sealed NMR tube was then heated at 50 °C to investigate the stability at higher

temperatures. Over the course of 11 h, partial conversion of 3 to tricarbonyl complex 1 occurred

(87% 3, 13% 1, Figure S41).

Figure S43. 1H NMR spectra monitoring catalyst during CO2 hydrogenation by 2 in toluene-d8 with Vk’s base and 1:1 mix of H2:CO2 at 1 atm of pressure (162 MHz, toluene-d8).

Figure S44. Concentration of 1-3 overtime during CO2 hydrogenation by 2 in toluene-d8 with Vk’s base and 1:1 mix of H2:CO2 at 1 atm of pressure. Dashed line represents change from monitoring reaction at 25 ˚C to heating reaction at 50 ˚C. Mass balance is based on total [Re]. Construction of Quantitative Reaction Coordinate Diagrams in MeCN. Reaction coordinate

diagrams (RCDs) in MeCN were constructed using quantitative values for the free energy of H2

0

0.05

0.1

0.15

0.2

0.25

0.3

0.35

0 5 10 15 20 25 30 35 40

[Re]

Time (h)

Mass Balance

2

3

1

2

3

1

Vk’s

[HVk’s]+

t = 0 h, 25 ˚C

t = 3 h, 25 ˚C

t = 24 h, 25 ˚C t = 3 h, 50 ˚C

t = 11 h, 50 ˚C

S53

splitting (DG1), free energy of hydride transfer (DDGH–), and the free energy of the overall reaction

(DGrxn), Figure S45. Through the RCD analysis, we found that over the temperature range studied

during catalysis, DGrxn was always downhill. The RCD also demonstrates the effect of temperature

during catalysis (Figure S45A), during which we observe the rate of catalysis to increase as

temperatures increase. Additionally, we found that changing the base from DBU to Vk’s had no

effect on the rate (and therefore the barrier), even though DG1 and DGrxn became considerably

more favorable (Figure S45B).

Figure S45. Idealized reaction coordinate diagrams for outer-sphere CO2 hydrogenation by catalyst 2 in MeCN using quantitative values for DG1, DDGH–, and DGrxn. Temperature effects are shown in (A) and the effect of changing the base pKa is shown in (B).

S54

Table S7. Comparison of catalyst activity for CO2 hydrogenation to formate at ambient temperatures and pressure (20 to 25 ˚C, 1 atm of 1:1 H2:CO2 blend) in non-aqueous solvents. TON is calculated as moles of formate / moles of catalyst. TOF is calculated as TON/time. iPrVk’s = 2,8,9-triisopropyl-2,5,8,9-tetraaza-1-phosphabicyclo[3.3.3]undecane.

Precatalyst Solvent Base (equiv.)

Time (h)

TON (% Yield)

TOF initial (h–1)

TOF final (h–1)

Reference

THF iPrVk’s (2035) <1 2000

(98%) 3400 25

THF iPrVk’s (667) 1.5 160

(24%) 64 8

THF iPrVk’s (275) 6 250

(91%) 67 41 26

DMF DBU (50000) 12 5652

(11%) 471 27

Toluene Vk’s (1000) 4.5 1023

(102%) 245 227 This work

S55

Table S8. Comparison of catalyst activity for CO2 hydrogenation to formate at ambient temperatures and pressure (25 ˚C, 1 atm of 1:1 H2:CO2) in 1 M NaHCO3(aq). TON is calculated as moles of formate / moles of catalyst. TOF is calculated as TON/time.

Precatalyst [Cat] (µM)

Time (h) TON

TOF initial (h–1)

TOF final (h–1)

Reference

50 24 92 7 3.8 28

50 33 330 27 10 29

20 18 30

20 14700 168 30

20 144 30

20 72 655 198 9 30

S56

Table S9. Comparison of catalyst activity for CO2 hydrogenation to formate at ambient temperatures and varying pressures (20 to 25 ˚C, 1:1 H2:CO2 blend) in non-aqueous solvents. TON is calculated as moles of formate / moles of catalyst. TOF is calculated as TON/time. iPrVk’s = 2,8,9-triisopropyl-2,5,8,9-tetraaza-1-phosphabicyclo[3.3.3]undecane.

Precatalyst Solvent Pressure (atm)

Base (equiv.)

Time (h)

TON (%

Yield)

TOF initial (h–1)

TOF final (h–1)

Reference

THF 20 iPrVk’s (12750) <1 9400

(74%) -- 74000 25

THF 34 iPrVk’s (26667) 1.5 19,000

(72%) 27,000 12666 8

THF 34 iPrVk’s (3200) 6 3150

(99%) 9700 6900 26

DMF 2 DBU (50000) 12 5434

(50%) -- 453 27

Toluene 20 Vk’s (10000) 0.16 556

(5.6%) 3334 -- This work

S57

X. Crystallographic Details

(tBuPOCOP)Re(CO)3 (1) – x1806002

Crystallographic Considerations. A colorless crystal (approximate dimension 0.200 x 0.100 x

0.100 mm3) was placed onto the tip of MiTeGen and mounted on a Bruker SMART Apex II

diffractometer and measured at 100 K. A preliminary set of cell constants was calculated from

reflections harvested from three sets of 12 frames. These initial sets of frames were oriented such

that orthogonal wedges of reciprocal space were surveyed and were used to produce initial

orientation matrices. The data collection was carried out using Cu Ka radiation (graphite

monochromator) with a frame time of 10 seconds and a detector distance of 4.0 cm. Sections of

frames were collected with 0.50º steps in w and f scans. Data to a resolution of 0.83 Å were

considered in the reduction. Final cell constants were calculated from the xyz centroids of 4839

strong reflections from the actual data collection after integration (SAINT).31 The intensity data

were corrected for absorption (SADABS).32 Please refer to Table S11 for additional crystal and

refinement information.

Structure solution and refinement. The space group P–1 was determined based on intensity

statistics and systematic absences. The structure was solved using XT33 and refined using XL

refinement program34 using least squares minimization in the Olex2 software.35 All non-hydrogen

atoms were refined with anisotropic displacement parameters. The hydrogen atoms were placed in

ideal positions and refined as riding atoms with individual relative isotropic displacement

parameters. The final full matrix least squares refinement converged to R1 = 0.0225 and wR2 =

0.0566 (F2, all data). The remaining electron density is located around the Re metal center.

S58

Table S11 Crystal data and structure refinement for x1806002. Identification code x1806002 Empirical formula C50H78O10P4Re2 Formula weight 1335.40 Temperature/K 100.15 Crystal system triclinic Space group P-1 a/Å 8.677(7) b/Å 12.469(8) c/Å 13.453(10) α/° 98.51(2) β/° 98.60(4) γ/° 105.218(15) Volume/Å3 1361.8(17) Z 1 ρcalcg/cm3 1.628 µ/mm-1 10.098 F(000) 668.0 Crystal size/mm3 0.200 × 0.100 × 0.100 Radiation CuKα (λ = 1.54184) 2Θ range for data collection/° 6.776 to 136.31 Index ranges -10 ≤ h ≤ 10, -14 ≤ k ≤ 15, -16 ≤ l ≤ 16 Reflections collected 26212 Independent reflections 4839 [Rint = 0.0373, Rsigma = 0.0239] Data/restraints/parameters 4839/0/310 Goodness-of-fit on F2 1.164 Final R indexes [I>=2σ (I)] R1 = 0.0225, wR2 = 0.0563 Final R indexes [all data] R1 = 0.0232, wR2 = 0.0566 Largest diff. peak/hole / e Å-3 1.37/-0.60

Table S12 Bond Lengths for x1806002. Atom Atom Length/Å Atom Atom Length/Å Re1 P2 2.4395(15) C1 C2 1.393(4) Re1 P1 2.4344(15) C6 C5 1.388(4) Re1 C1 2.175(3) C5 C4 1.392(4) Re1 C24 1.950(3) C4 C3 1.388(4) Re1 C25 1.995(4) C11 C13 1.533(5) Re1 C23 1.961(4) C11 C14 1.539(4) P2 O2 1.657(2) C11 C12 1.537(4) P2 C11 1.885(3) C3 C2 1.387(4) P2 C7 1.896(3) C8 C7 1.555(4)

S59

P1 O1 1.652(2) C9 C7 1.544(4) P1 C15 1.892(4) C7 C10 1.532(5) P1 C19 1.885(3) C15 C16 1.543(5) O1 C2 1.403(4) C15 C17 1.546(5) O2 C6 1.399(4) C15 C18 1.541(5) O3 C23 1.157(4) C19 C21 1.540(5) O4 C24 1.155(4) C19 C20 1.539(5) O5 C25 1.157(4) C19 C22 1.531(5) C1 C6 1.394(4)

Table S13 Bond Angles for x1806002. Atom Atom Atom Angle/˚ Atom Atom Atom Angle/˚ P1 Re1 P2 150.39(3) C5 C6 O2 117.5(3) C1 Re1 P2 75.61(9) C5 C6 C1 123.7(3) C1 Re1 P1 75.55(9) C6 C5 C4 117.7(3) C24 Re1 P2 105.45(10) C3 C4 C5 121.5(3) C24 Re1 P1 103.97(11) O5 C25 Re1 178.9(3) C24 Re1 C1 173.43(12) C13 C11 P2 109.5(2) C24 Re1 C25 91.91(15) C13 C11 C14 108.2(3) C24 Re1 C23 82.30(14) C13 C11 C12 110.7(3) C25 Re1 P2 91.78(10) C14 C11 P2 109.0(2) C25 Re1 P1 90.53(11) C12 C11 P2 113.1(2) C25 Re1 C1 81.56(13) C12 C11 C14 106.2(3) C23 Re1 P2 90.37(10) C2 C3 C4 117.7(3) C23 Re1 P1 90.26(10) C8 C7 P2 105.4(2) C23 Re1 C1 104.23(13) C9 C7 P2 110.5(2) C23 Re1 C25 174.17(13) C9 C7 C8 111.5(3) O2 P2 Re1 105.17(8) C10 C7 P2 115.2(2) O2 P2 C11 99.88(14) C10 C7 C8 105.7(3) O2 P2 C7 99.16(13) C10 C7 C9 108.4(3) C11 P2 Re1 122.87(10) C16 C15 P1 111.4(2) C11 P2 C7 108.73(14) C16 C15 C17 111.1(3) C7 P2 Re1 116.55(11) C17 C15 P1 105.8(2) O1 P1 Re1 105.03(9) C18 C15 P1 114.9(2) O1 P1 C15 98.67(14) C18 C15 C16 108.2(3) O1 P1 C19 100.24(14) C18 C15 C17 105.4(3) C15 P1 Re1 116.71(11) C1 C2 O1 118.7(3) C19 P1 Re1 122.34(12) C3 C2 O1 117.4(3) C19 P1 C15 109.29(15) C3 C2 C1 123.8(3) C2 O1 P1 116.42(18) O3 C23 Re1 172.5(3) C6 O2 P2 116.81(19) C21 C19 P1 112.5(2)

S60

C6 C1 Re1 122.4(2) C20 C19 P1 109.1(2) C2 C1 Re1 122.0(2) C20 C19 C21 110.0(3) C2 C1 C6 115.2(3) C22 C19 P1 109.5(2) O4 C24 Re1 177.6(3) C22 C19 C21 107.0(3) C1 C6 O2 118.8(3) C22 C19 C20 108.6(3)

S61

(tBuPOCOP)Re(CO)2 (2) – x1806007

Crystallographic Considerations. An orange crystal (0.05 x 0.184 x 0.247 mm3) was placed onto

the tip of MiTeGen and mounted on a Bruker SMART Apex II diffractometer and measured at

100 K. A preliminary set of cell constants was calculated from reflections harvested from three

sets of 12 frames. These initial sets of frames were oriented such that orthogonal wedges of

reciprocal space were surveyed and were used to produce initial orientation matrices. The data

collection was carried out using Cu Ka radiation (graphite monochromator) with a frame time of

10 seconds and a detector distance of 4.0 cm. Sections of frames were collected with 0.50º steps

in w and f scans. Data to a resolution of 0.84 Å were considered in the reduction. Final cell

constants were calculated from the xyz centroids of 8986 strong reflections from the actual data

collection after integration (SAINT).31 The intensity data were corrected for absorption

(SADABS).32 Please refer to Table S14 for additional crystal and refinement information.

Structure solution and refinement. The space group P–1 was determined based on intensity

statistics and systematic absences. The structure was solved using XT33 and refined using XL

refinement program34 using least squares minimization in the Olex2 software.35 The asymmetric

unit contained two molecules of (tBuPOCOP)Re(CO)3. All non-hydrogen atoms were refined with

anisotropic displacement parameters. The hydrogen atoms were placed in ideal positions and

refined as riding atoms with individual relative isotropic displacement parameters. The final full

matrix least squares refinement converged to R1 = 0.0215 and wR2 = 0.0503 (F2, all data). The

remaining electron density is located around the Re metal center.

S62

Table S14 Crystal data and structure refinement for 1806007. Identification code 1806007 Empirical formula C24H39O4P2Re Formula weight 639.69 Temperature/K 100.15 Crystal system triclinic Space group P-1 a/Å 8.1984(5) b/Å 13.3760(7) c/Å 24.3513(17) α/° 79.938(3) β/° 86.858(3) γ/° 84.443(2) Volume/Å3 2615.0(3) Z 4 ρcalcg/cm3 1.625 µ/mm-1 10.456 F(000) 1280.0 Crystal size/mm3 0.247 × 0.184 × 0.05 Radiation CuKα (λ = 1.54178) 2Θ range for data collection/° 3.688 to 133.328 Index ranges -9 ≤ h ≤ 9, -15 ≤ k ≤ 15, -28 ≤ l ≤ 28 Reflections collected 53996 Independent reflections 8986 [Rint = 0.0316, Rsigma = 0.0209] Data/restraints/parameters 8986/0/583 Goodness-of-fit on F2 1.164 Final R indexes [I>=2σ (I)] R1 = 0.0215, wR2 = 0.0499 Final R indexes [all data] R1 = 0.0223, wR2 = 0.0503 Largest diff. peak/hole / e Å-3 1.18/-1.17 Table S15 Bond Lengths for 1806007. Atom Atom Length/Å Atom Atom Length/Å Re1 P1 2.3876(7) C15 C18 1.533(4) Re1 P2 2.3750(6) C15 C16 1.543(4) Re1 C1 2.130(3) C6 C5 1.382(4) Re1 C24 1.840(3) C25 C26 1.398(4) Re1 C23 1.930(3) C25 C30 1.403(4) Re2 P3 2.3827(7) C29 C30 1.391(4) Re2 P4 2.3752(7) C29 C28 1.392(4) Re2 C25 2.127(3) C19 C20 1.535(4) Re2 C47 1.944(3) C19 C22 1.529(4)

S63

Re2 C48 1.837(3) C19 C21 1.541(4) P3 O5 1.666(2) C26 C27 1.387(4) P3 C31 1.865(3) C5 C4 1.395(4) P3 C35 1.864(3) C27 C28 1.396(4) P1 O1 1.6608(19) C11 C12 1.532(4) P1 C15 1.867(3) C11 C13 1.536(4) P1 C19 1.868(3) C11 C14 1.538(4) P2 O2 1.6636(19) C37 C35 1.538(4) P2 C11 1.860(3) C7 C10 1.536(4) P2 C7 1.856(3) C7 C9 1.549(4) P4 O6 1.664(2) C7 C8 1.529(4) P4 C43 1.866(3) C31 C32 1.538(4) P4 C39 1.860(3) C31 C34 1.545(5) O2 C6 1.394(3) C31 C33 1.537(4) O5 C26 1.390(3) C2 C3 1.387(4) O1 C2 1.394(3) C45 C43 1.539(4) O4 C24 1.182(4) C35 C38 1.533(5) O6 C30 1.389(3) C35 C36 1.537(4) O8 C48 1.175(4) C43 C46 1.532(4) O7 C47 1.150(4) C43 C44 1.536(5) O3 C23 1.154(4) C3 C4 1.396(4) C1 C6 1.404(4) C39 C42 1.533(4) C1 C2 1.399(4) C39 C41 1.534(4) C15 C17 1.533(4) C39 C40 1.541(5)

Table S16 Bond Angles for 1806007. Atom Atom Atom Angle/˚ Atom Atom Atom Angle/˚ P2 Re1 P1 153.00(2) C30 C25 Re2 122.1(2) C1 Re1 P1 76.71(8) C30 C29 C28 117.8(3) C1 Re1 P2 76.80(8) C20 C19 P1 109.29(19) C24 Re1 P1 96.57(9) C20 C19 C21 109.3(2) C24 Re1 P2 95.82(8) C22 C19 P1 114.8(2) C24 Re1 C1 106.49(12) C22 C19 C20 109.1(2) C24 Re1 C23 88.22(14) C22 C19 C21 108.0(2) C23 Re1 P1 103.03(9) C21 C19 P1 106.28(19) C23 Re1 P2 101.24(9) O5 C26 C25 118.8(2) C23 Re1 C1 165.26(13) C27 C26 O5 117.5(2) P4 Re2 P3 153.23(2) C27 C26 C25 123.7(2) C25 Re2 P3 76.84(8) C6 C5 C4 117.4(2) C25 Re2 P4 76.91(8) O6 C30 C25 118.8(2)

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C47 Re2 P3 103.84(8) O6 C30 C29 117.8(2) C47 Re2 P4 100.85(8) C29 C30 C25 123.4(3) C47 Re2 C25 168.63(12) C26 C27 C28 117.7(3) C48 Re2 P3 95.26(9) C12 C11 P2 110.1(2) C48 Re2 P4 96.77(9) C12 C11 C13 109.2(2) C48 Re2 C25 105.65(12) C12 C11 C14 109.3(3) C48 Re2 C47 85.64(13) C13 C11 P2 114.5(2) O5 P3 Re2 105.85(7) C13 C11 C14 107.9(2) O5 P3 C31 99.16(12) C14 C11 P2 105.67(19) O5 P3 C35 100.16(12) C10 C7 P2 115.4(2) C31 P3 Re2 114.61(11) C10 C7 C9 107.8(3) C35 P3 Re2 119.91(10) C9 C7 P2 103.60(19) C35 P3 C31 113.21(14) C8 C7 P2 110.0(2) O1 P1 Re1 105.83(7) C8 C7 C10 110.8(2) O1 P1 C15 99.62(12) C8 C7 C9 108.9(2) O1 P1 C19 100.72(12) C32 C31 P3 110.4(2) C15 P1 Re1 114.47(9) C32 C31 C34 108.3(3) C15 P1 C19 112.34(13) C34 C31 P3 103.8(2) C19 P1 Re1 120.23(9) C33 C31 P3 115.1(2) O2 P2 Re1 106.41(7) C33 C31 C32 110.7(3) O2 P2 C11 99.85(11) C33 C31 C34 108.1(3) O2 P2 C7 100.02(12) O1 C2 C1 118.7(2) C11 P2 Re1 121.27(9) C3 C2 O1 117.7(2) C7 P2 Re1 112.99(9) C3 C2 C1 123.6(2) C7 P2 C11 112.70(13) C37 C35 P3 108.8(2) O6 P4 Re2 106.00(7) C38 C35 P3 106.6(2) O6 P4 C43 100.20(12) C38 C35 C37 110.7(3) O6 P4 C39 100.31(13) C38 C35 C36 107.9(3) C43 P4 Re2 122.00(10) C36 C35 P3 113.8(2) C39 P4 Re2 111.98(10) C36 C35 C37 109.1(2) C39 P4 C43 112.82(14) C45 C43 P4 108.4(2) C6 O2 P2 115.76(16) C46 C43 P4 114.0(2) C26 O5 P3 116.08(17) C46 C43 C45 109.7(2) C2 O1 P1 116.17(17) C46 C43 C44 108.3(3) C30 O6 P4 116.02(17) C44 C43 P4 106.9(2) C6 C1 Re1 122.2(2) C44 C43 C45 109.4(3) C2 C1 Re1 122.20(19) O4 C24 Re1 178.3(3) C2 C1 C6 115.5(2) C2 C3 C4 117.5(3) C17 C15 P1 109.8(2) C42 C39 P4 110.7(2) C17 C15 C16 108.8(3) C42 C39 C41 110.8(3) C18 C15 P1 115.1(2) C42 C39 C40 107.9(3) C18 C15 C17 110.6(2) C41 C39 P4 115.2(2)

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C18 C15 C16 108.0(3) C41 C39 C40 108.1(3) C16 C15 P1 104.13(19) C40 C39 P4 103.8(2) O2 C6 C1 118.7(2) C29 C28 C27 121.8(3) C5 C6 O2 117.5(2) C5 C4 C3 122.1(3) C5 C6 C1 123.8(2) O7 C47 Re2 178.8(3) C26 C25 Re2 122.3(2) O8 C48 Re2 179.2(3) C26 C25 C30 115.6(3) O3 C23 Re1 176.8(3)

Structural Comparison of 1 and 2.

Decarbonylation of 1 (pseudo-octahedral geometry) to generate 2 (pseudo-square pyramidal

geometry) results in notable changes in the solid state structure. The carbonyl ligand trans to the

phenyl backbone of the pincer in 1 and 2 are fairly similar, with Re–C bond lengths of 1.950(3)

and 1.937 Å and C–O bond lengths of 1.155(4) and 1.152 Å respectively. Conversely, the carbonyl

ligand cis (or axial) to the phenyl backbone in 1 and 2 are strikingly different with Re–C bond

lengths of 1.978 and 1.839 Å, which is concurrent with the drastic difference in their respective

C–O bond lengths lengthening from 1.157 Å in 1 to 1.178 Å in 2. Additionally, the Re–P bond

distances contract from 2.437 Å in 1 to 2.375 Å in 2 concomitant with an increase in the pincer P–

Re–P angle from 150.39(3)˚ in 1 to 153.12˚ in 2.

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