By measuring the dependence of the RWGS reaction rates on the intensity and spectral power distribution of filtered light incident onto the nanostructured Pd@Nb2O5 catalyst, it is determined that the RWGS reaction is activated photothermally. That is the RWGS reaction is initiated by heat generated from thermalization of charge carriers in the Pd nanocrystals that are excited by interband and intraband absorption of visible and NIR light. Taking advantage of this photothermal effect, a visible and NIR responsive Pd@Nb2O5 hybrid catalyst that efficiently hydrogenates CO2 to CO at an impressive rate as high as 1.8 mmol gcat−1 h−1 is developed. The mechanism of this photothermal reaction involves H2 dissociation on Pd nanocrystals and subsequent spillover of H to the Nb2O5 nanorods whereupon adsorbed CO2 is hydrogenated to CO. This work represents a significant enhancement in our understanding of the underlying mechanism of photothermally driven CO2 reduction and will help guide the way toward the development of highly efficient catalysts that exploit the full solar spectrum to convert gas‐phase CO2 to valuable chemicals and fuels.The reverse water gas shift (RWGS) reaction driven by Nb One appealing route to the production of solar fuels is to use renewable sources of hydrogen to reduce the greenhouse gas CO2 to useful fuels has been studied.2 into useful products, gas‐phase light‐assisted reduction of CO2 is a practical option, since gas‐phase processes can be easily scaled and integrated with existing chemical and petrochemical industry infrastructure.2 to useful fuels over supported metallic catalysts, which are traditionally known to be activated thermally at elevated temperatures, has been investigated. In this context, the photothermal approach to the catalytic hydrogenation of CO2 has emerged with reported conversion rates as high as mmol gcat−1 h−1 to mol gcat−1 h−1.2 reduction reactions. Today however, the mechanism of photothermal catalysis remains unclear. By identifying the key factors which underpin the photothermal effect that promotes catalytic CO2 reduction reactions, it will prove possible to design and make nanoscale catalysts with structural, chemical, and physical properties that serve to maximize the rate of reduction of gaseous CO2 to chemicals and fuels, using both the heat and light from sunlight.The activity of various photocatalyts toward the light‐assisted reduction of COd band and s‐p conduction band) and intraband transitions (between filled and empty states in the d and s–p bands).For the case of metal nanocrystals, the photothermal effect can arise in different ways that include surface plasmon resonance (SPR) and nonradiative relaxation, and intraband or interband nonradiative relaxation of photoexcited charge carriers. The SPR effect occurs when absorbed radiant energy is stored in the collective resonant oscillations of conduction electrons.2 reduction have focused on metal nanocrystals dispersed on insulating metal oxide supports.2 to the surface of the catalysts, prolong photogenerated charge‐carrier lifetimes and enhance the reactivity of CO2.The catalyst support may also influence the photothermal effect. Recent efforts toward the development of photothermal catalysts for gas‐phase CO2 + H2→ CO + H2O) driven with visible and near‐infrared photons over hybrid catalysts composed of palladium nanocrystals dispersed on Nb2O5 nanorods denoted Pd@Nb2O5. Relevant prior work on Pd nanocrystals includes electrocatalytic,2,2 photothermal methanation,2 hydrogenation at high rates but require high operating temperatures (523 K–900 K) and high pressures.Herein we report the reverse water gas shift (RWGS) reaction  was prepared via hydrothermal synthesis as described in the experimental section and illustrated in Figure3O7(OH) reflect the orthorhombic crystal structure and underlying symmetry of the unit cell.3O7(OH) nanorods. Figure S1a (Supporting Information) is a high angle annular dark field (HAADF) image of the Nb3O7(OH) nanorods at low magnification. Figure S1b (Supporting Information) shows an atomic resolution HAADF image from a nanorod with its [001] direction parallel to the electron beam. The FFT of the image is shown in Figure S1c (Supporting Information), the (2 0 0), (1 1 0), and (0 2 0) spots can be clearly resolved and the viewing direction is confirmed to be [001]. From both the image and the FFT, [010] is found to be the preferred growth direction of the Nb3O7(OH) nanorods. Figure S1d (Supporting Information) is an atomic model of Nb3O7(OH) marked by the red rectangle in Figure S1b (Supporting Information). The dotted blue rectangle in Figure S1d (Supporting Information) indicates one unit cell.Nb3O7(OH) nanorods with Pd nanoparticles  nanorods is believed to provide binding sites for Pd nanocrystals. Therefore, the hybrid structures were obtained devoid of free Pd nanocrystals, as seen in the TEM image shown in Figure In the literature, Pd nanocrystals are typically loaded onto catalyst supports using an impregnation method wherein the supports are immersed in a palladium precursor solution prior to being thermally treated. However, in this work we used a straightforward microwave‐assisted hydrothermal reduction to decorate the Nbs Figure a. With t3O7(OH) nanorods to Nb2O5 nanorods, we pre‐treated Pd@Nb3O7(OH) at 120 °C for 24 h. The PXRD patterns in Figure S2 (Supporting Information) show that before the pre‐treatment the sample consists of orthorhombic Nb3O7(OH), while the treated sample underwent a structural transformation from orthorhombic Nb3O7(OH) to monoclinic Nb2O5. Due to the low Pd loading (0.5 wt%) and the small size of the nanocrystals, diffraction peaks associated with palladium species were not observed. As shown in the TEM images of pure Nb3O7(OH) nanorods before the pre‐treatment  and Nb2O5 nanorods after the pre‐treatment , the nanorod morphology is retained. The TEM image in Figure 2O5 nanorods in Pd@Nb2O5. The Pd nanocrystals were found to be well dispersed on the surface of the Nb2O5 nanorods with diameter in the range of 2–10 nm.To transform Nb3O7(OH), Nb2O5 and Pd@Nb2O5 samples, determined using the Brunauer–Emmett–Teller (BET) method, were similar, namely 84.9, 82.1, and 82.4 m2 g−1, respectively.It is well documented in the heterogeneous catalyst literature that specific surface area is important, where catalytic reaction sites and rates scale with surface area.2O5 nanorod is shown in Figure2O5 nanorod. The cyan arrow indicates the growth direction of the nanorod, and the lattice fringes corresponding to the (1 1 4¯) surface are visible in the nanorod. An enlarged image of the Pd nanocrystal, viewed down the [112] direction, is shown in Figure (1 1 1¯) and (2 2¯ 0)0 planes are 2.37 Å and 1.40 Å, respectively. Since the contrast of a HAADF image is not only determined by the Z number, but can also depend on thickness and channeling effects, we used the fast Fourier transform (FFT) to further verify the distribution of the Pd nanocrystals. The FFT of Figure 2O5 nanostructures are visible, and the Bragg peaks from (1 1 1¯) and (2 2¯ 0) planes are labeled. Figure In order to determine the chemical state of the Pd nanocrystals, the samples were further analyzed by STEM and XPS. A HAADF Z‐contrast image of a Pd nanocrystal decorated Nb2O5 nanorods. Individual EEL spectra extracted from the nanocrystal, nanocrystal–nanorod interface and nanorod are plotted in Figure 4,5‐edge and the Pd M3‐edge, while the spectra collected from the nanorod exhibits the Nb M3, Nb M2, and O K‐edges. As expected, the spectra collected at the interface between the Pd nanocrystal and the Nb2O5 nanorod show a superposition of these two individual spectra. Since the Pd M4,5‐edge overlaps with the Nb M2‐edge, and the Pd M3‐edge overlaps with the O K‐edge, we performed a multiple linear least square (MLLS) fitting analysis on the EELS mapping to distinguish those signals from each other. The spectra taken from Pd and Nb2O5 were used as references to obtain a fitting coefficient of Pd and Nb2O5. Figure 2O5 are mapped in Figure 2O5 nanorod. Therefore, both atomic structure and elemental mapping validate the architecture of the Pd nanocrystal decoration on the Nb2O5 nanorod.We used electron energy‐loss spectroscopy (EELS) to further study the chemical composition of the Pd nanocrystal decorated Nb2O5 sample with double the Pd loading (1 wt%), the Pd XPS signal could be discerned where binding energies of Pd 3d at 335.6 eV and 340.8 eV are assigned to the 3d5/2 and 3d3/2 spin‐orbit peaks of metallic Pd, further validating the accumulating evidence that Pd nanocrystals are in the zero‐valent metallic state in the Pd@Nb2O5 sample .At the lowest Pd loading (0.5 wt%), we were not able to determine the chemical state of Pd by performing XPS measurements. However, by using a Pd@NbFigure2O5 sample is much less than that of the Nb3O7 (OH) and Nb2O5 samples throughout the visible and NIR spectral regions because a significant amount of light is absorbed by the Pd nanocrystals. In the literature, the difference between insulating wide bandgap catalyst supports with and without Pd loaded onto their surface has been attributed to light absorption by the Pd nanoparticles.2O5 was determined to be 3.44 eV by fitting the diffuse reflectance spectra with a modified Kubelka–Munk function, as shown in Figure S5 (Supporting Information). From our XPS results , the valence band maximum (VB) and Fermi energy (Ef) of the Nb2O5 nanorods were determined to be 7.31 eV and 4.01 eV, respectively. The work function of Pd is reported in the literature.2O5 can be constructed  in the dark at room temperature (RT), (2) under irradiation from a 300 W Xe lamp at 25 kW m−2, and (3) in the dark at a temperature of 160 °C. Moreover, in this work all reported rates are normalized to the weight of Pd since the active sites reside at the Pd@Nb2O5 interface, as will be described further in Section In order to determine the photothermal catalytic activity of nanostructured Pd@NbFigure13CO products originated from 13CO2 rather than adventitious carbon sources . The results in Figure −2, we observe a CO2—to—CO hydrogenation rate as high as 4.9 mmol gcat−1 h−1 over the hybrid Pd@Nb2O5. For comparison, the reaction proceeded at a rate of 1 mmol gcat−1 h−1in the dark at 160 °C. The catalytic rate in the light is roughly five times the rate in the dark under thermal heating, demonstrating that the RWGS reaction can be efficiently driven solely by light over the nanostructured Pd@Nb2O5 catalyst. It can also be noted that no 13C labeled CH4 product was present for either the photo‐ or thermal catalytic reaction tests, with the exception of some tests carried out under the highest light intensity. That is, for some tests wherein the intensity of the incident light was 25 kW m−2, small amounts of 13CH4 product were visible. However, it should be noted that even for these cases the CH4 production rate was typically less than two orders of magnitude of the 13CO production rate and never exceeded 5% of the CO production rate for any given test.The duration of each test was 3 h and the results are plotted in 2.32O5 catalyst sample was irradiated with a 300 W Xe lamp, fitted with different high‐pass cut‐off filters, and the sample was exposed to air for 24 h between each test. The relative spectral distribution for this series of tests, labeled A through F, is shown in Figure −1 h−1. The CO production rates for batch reactions A through F are plotted in Figure cat−1 h−1, 1.1 mmol gcat−1 h−1, 0.78 mmol gcat−1 h−1 and 0.41 mmol gcat−1 h−1, respectively. For test C, the CO2 hydrogenation reaction proceeded at a rate of 1.8 mmol gcat‐1 h−1 using photons in the visible and NIR region. This suggests the RWGS reaction can be driven by visible and NIR light over nanostructured Pd@Nb2O5 catalysts.In order to quantitatively investigate the effects of the spectral power distribution on the CO production rate, a series of tests was carried out wherein the Pd@NbA = 1 − T − R, where R and T are the diffuse reflectance and transmittance, respectively.We also investigated the light intensity dependence of the CO production rates. Specifically, a series of tests labeled I through V was carried out wherein the intensity of the incident light from the 300 W Xe Lamp was sequentially reduced, and the sample was exposed to air for 24 h between different tests. No filters were used and the relative spectral irradiance for tests I through V are plotted in Figure −2 and 10.2 kW m−2, respectively. The RWGS rates for these two tests are also nearly identical despite the fact that the nanostructured Pd@Nb2O5 catalyst is primarily irradiated with NIR photons (λ > 600 nm) for test E whereas for test III more than 50% of the incident radiant energy is from the spectral region with λ < 600 nm. Here, it should also be noted that photons with energy greater than the bandgap of Nb2O5 (λ < 360 nm) account for more than 5% of the incident radiant energy in test III whereas no photons in this spectral region are incident on the Pd@Nb2O5 catalyst in test E. These results imply that electron–hole pairs photogenerated in the Nb2O5 nanorods (Eg = 3.44 eV) upon absorption of UV photons or ultraviolet plasmonic excitation of the Pd nanocrystals are not the primary factors driving the RWGS reaction. Thus, we can make the following two statements based on the results shown in Figure 1)2O5 and less than the plasmon resonance energy of Pd nanocrystals, one can deduce that photons absorbed in the nanostructured Pd@Nb2O5 catalyst provide thermal energy caused by intraband and/or interband electronic transitions and nonradiative relaxation. This activates the RWGS reaction photothermally.Considering the visible and NIR spectral range, where the energy of incident photons is less than the bandgap of Nb2)2O5 support or by exciting local plasmon resonance in Pd nanocrystals. While these photochemical reaction pathways contribute to the overal reaction rates, they are not the main driving force for the RWGS reaction over the nanostructured Pd@Nb2O5 catalysts. If the photochemical contribution was the main driving force then the dashed black line corresponding to tests I through V in Figure Considering the UV spectral range, photons have sufficient energy to activate the RWGS photochemically by generating electron–hole pairs in the NbThe RWGS reaction rates for the series of tests A through F and I through V are plotted as the red and black dashed lines in Figure  Eg = 3.4 eV upon2O5 catalyst, which photothermally accelerates the RWGS reaction.Therefore, we propose that, under illumination, absorbed photons ultimately thermalize and generate heat, resulting in high local temperatures in the Pd@NbA TEM image of the sample taken after these photothermal catalytic tests is shown in Figure S8 (Supporting Information), notably illustrating that there is no observable change in nanostructure morphology.2.42O5 catalyst as a function of the intensity and spectral distribution of the incident radiant power.Since it is clear that the RWGS is driven photothermally, we were interested in determining the temperature of the nanostructured Pd@NbTe, defined herein as the catalyst temperature associated with a given reaction rate attained when the catalyst was tested under similar reaction conditions but in the dark. Therefore, in order to estimate Te for the nanostructured Pd@Nb2O5 catalyst during tests at RT when subjected to incident radiant energy, the film sample was tested at temperatures of 60 °C, 85 °C, 110 °C, 135 °C, and 160 °C in the dark.One approach was to determine an “effective” reaction temperature of the catalyst, Figure−2 to 18 kW m−2.We fit the reaction rates measured for these dark tests, which are plotted in 2.52‐to‐CO over M@M'Ox catalysts.Figure2 on M, which then migrate (spillover) to the M‐M'Ox interface and diffuse onto the M'Ox, acting as a reducing species. Formate is formed in the CO2 chemisorption and hydrogenation process on the surface of the M'Ox support, followed by a rearrangement that generates CO and surface OH.Our justification for using an Arrhenius rate law in the previous subsection is based on previous studies that report a kinetic model for the thermally driven catalytic conversion of CO2O5, we first conducted controlled catalytic measurements over Pd@Nb2O5 under illumination in the presence and absence of H2, as shown in Figure S9 (Supporting Information). The results confirm that the reduction of CO2 can happen only in the presence of H2.In order to investigate if the RWGS proceeds by means of the formate decomposition pathway shown in Figure 2 adsorption capacity of the Pd@Nb2O5 catalysts using thermogravimetric analysis. The CO2 adsorption capacities are normalized to the surface area of each sample, which was measured using the BET method. The CO2 adsorption capacities are 1.11 μmol CO2 m−2 and 1.13 μmol CO2 m−2 for Nb2O5 and Pd@Nb2O5, respectively. These values are very similar. The similarity in CO2 adsorption for the Nb2O5 and Pd@Nb2O5 samples suggests that Pd does not significantly affect the availability of surface‐bound CO2 species.We investigated the COWe also considered the reaction rate activation energy. In this context it has been shown that the most likely rate‐limiting step for CO formation involves the conversion of formate intermediates to CO at the metal/oxide interface according to the reaction scheme: −1.2O5 an activation energy of 76 kJ mol−1. Indeed, the fact that these activation energies are comparable provides further evidence that the RWGS likely proceeds over the Pd@Nb2O5 catalyst via the formate decomposition pathway shown in Figure The activation energy of this rate‐limiting step is reported to be 66 kJ mol2O5 catalyst during the reaction, we performed both hydrogenation and thermogravimetric analysis (TGA) measurements on fresh and used catalysts. To conduct the hydrogenation experiment over used Pd@Nb2O5 catalysts that had been tested under the high light intensity testing conditions for over 40 h, we infiltrated the photoreactor with pure hydrogen gas at a pressure of 2 atmospheres for a duration of 3 h under an intensity of 25 kW m−2. No carbonaceous species were detected in the hydrogen gas after the 3 h test. We then repeated this hydrogenation test with the reactor heated to a temperature of 150 °C, and were still unable to detect any carbonaceous species. The TGA experiments were performed on Pd@Nb2O5 catalysts that had been used under high light intensity testing conditions reported in the paper for over 40 h.2 and the other flowing air. The TGA results for the used catalyst are identical to those for a fresh catalyst. Thus, the results from both our hydrogenation and TGA experiments imply that the surface of the Pd@Nb2O5 catalyst is devoid of surface carbon after catalyzing the photothermal reduction of CO2.To clarify whether or not any carbon was deposited on the surface of the Pd@Nb2.62O5 and Nb2O5 in order to more directly measure the local temperature that the samples attain in the light. In 2001, Bell conducted a pivotal investigation of the effects of temperature on the intensities and positions of Raman bands of metal oxides.2O5, a linear shift in the position of the band for niobium‐oxygen vibrations was observed as the acquisition temperature increased.2O5 and Nb2O5 at 6 power levels, as shown in Figure S11 (Supporting Information). The νNb=O stretching modes for Pd@Nb2O5 and Nb2O5 at around 990 cm−1 under different power levels are shown in Figure2O5 exhibited a shift to low frequencies. Under the lowest power level where useful Raman spectra are attainable, which is found to be 12 μW (laser power density is 2.4 kW cm−2), we assume that the temperatures of Pd@Nb2O5 and Nb2O5 are both approximately 300 K. This assumption is based on recent work by Kim et al. wherein Raman measurements performed with laser excitation at a wavelength of 532 nm and power density of 0.670 kW cm−2 on Bi2Se3 and Sb2Te3 crystals did not increase their temperature above 300 K.2Se3 and Sb2Te3 (≈0.3 eV), we do not expect the temperature of our Pd@Nb2O5 and Nb2O5 to be significantly elevated above room temperature when subjected to laser excitation at the lowest laser power density of 2.4 kW cm−2.It is important to examine the local temperature of the nanostructure surface under irradiation in order to understand the underlying mechanisms associated with the photothermal effect and determine the specific roles played by the metal nanocrystals and supports. To further understand the impact that the photothermal effect has on the temperature of the sample, Raman measurements were conducted on nanostructured Pd@Nb2O5 and Nb2O5 at the other 5 power levels using the linear dependence of the band shift with temperature based on Bell's findings, as shown in Figure 2O5 at the highest laser power level of 12 mW (laser power density is 2388 kW cm−2) is estimated to be 777 K while the temperature of Nb2O5 at the same power level is estimated to be 368 K.We estimated the temperatures of Pd@Nb2O5 and Nb2O5 sample using the correlation between the intensities of the Stokes and anti‐Stokes components of the Raman spectra where the local temperature of the sample can be determined from the expression:In addition to the aforementioned approach involving temperature‐dependent Raman mode frequency shifts, one can directly and quantitatively determine the local temperature of the Pd@NbISt and IASt, are the integral of the collected Raman signals from the Stokes (S) and anti‐Stokes (AS) Raman modes, respectively; T is the temperature in degrees Kelvin; ER = hc/λexcitation − hc/λemission = hc ωV, where ωV is the temperature‐dependent wave number of the Raman mode (cm−1); ωL is the wave number of the exciting laser light (12739 cm−1), giving ((ωL + ωV)/(ωL – ωV))4 = 1.492.Here the intensities of the signals, 2O5 and Nb2O5 were measured under a similar power density of 3050 kW cm−2 (24 mW) using a pump wavelength of 785 nm. The Raman signals of the Stokes and anti‐Stokes bands for the νNb—O—Nb stretching mode around 640 cm−1 are plotted in Figure S12 (Supporting Information).2O5 and Nb2O5 are estimated to be 776 K and 385 K, respectively.In order to determine the local temperature of the sample, both the Stokes and weaker anti‐Stokes Raman spectra of Pd@Nb2O5 and Nb2O5 determined by the two methods are consistent and clearly demonstrate that there is a photothermal effect of the Pd nanocrystals decorated on the Nb2O5 nanorods. The increase in the temperature of the Nb2O5 nanorods alone at different laser power levels is minimal, while that of Nb2O5 nanorods decorated with the Pd “nano‐heaters” can be increased from 300 K to around 777 K by increasing the pump laser power densities. This can be attributed to the photothermal effect of the Pd nanocrystals. Since we were using 633 nm and 785 nm laser light, it is also confirmed that the photothermal effect of the Pd nanocrystals can be induced by visible irradiation. It is important to note that the Raman measurements presented in this work demonstrate the photothermal effect on the Pd@Nb2O5 catalyst and also provide valuable insight directly demonstrating how the Pd nanocrystals heat up to a much greater extent than the Nb2O5 nanorod supports. However, it is equally important to point out the distinction between the laser excitation used in the Raman measurement when compared to broadband solar irradiation. Both have different brightness and associated spectral width, and as such, the measurements reported here are not used to determine exact estimates of the temperature of the Pd@Nb2O5 catalyst during the reaction in the light because the laser light source is monochromatic compared to the polychromatic output of the Xe lamp used to measure the RWGS reaction rates. Nevertheless, the Raman experiments performed herein allow us to infer that the comparative photothermal heating and temperature increase of the Pd@Nb2O5 nanocrystal–nanorod can be attributed to the interband and intraband absorption contributed by the Pd nanocrystals. The measurements provide direct evidence that this optical absorption is the underlying physicochemical phenomenon driving the RWGS reaction over the nanostructured Pd@Nb2O5.The local temperatures of Pd@Nb2.7Figure2O5 sample was initially irradiated, a CO production rate of 4.5 mmol gcat−1 h−1 was observed for a 3 h test. The second run produced CO at a rate of 6.7 mmol gcat−1 h−1 and the third run produced CO at a rate of 10.2 mmol gcat−1 h−1. For the fourth run, the RWGS proceeded at a rate of 11.7 mmol gcat−1 h−1, which is 2.6 times that of the first run. Finally, the film sample was exposed to air for 24 h and a repeat measurement was conducted, more or less reproducing the original rate of 4.7 mmol gcat−1 h−1. These results demonstrate that a certain kind of chemical/surface species were generated in situ during the catalyst testing process and these species, which exist under vacuum or H2 and CO2 gas atmosphere, are influencing the photothermal catalytic reaction in a positive way. Typically, on metal oxide supports such as TiO2, CeO2, and Nb2O5, oxygen vacancies (Vo) and reduced surface states  can be formed by annealing in an inert environment at high temperatures or under illumination.2 to the surface, lifetime of photogenerated charge carriers and visible light absorption.Often, the activity of photocatalysts decrease gradually with sequential tests due to numerous possible reasons including sintering, photodegradation, catalyst poisoning, and adventitious contamination. However, in contrast, our photothermal catalysts display a gradual rate enhancement with the time of testing. To gain a deeper insight into this unusual behavior, a set of photothermal catalytic measurements was carried out under similar reaction conditions as the tests reported in Figure 2O5 film sample before and after conducting the photothermal catalytic tests. When oxygen vacancies are generated in metal oxides, they can create mid‐gap color centers, enhancing the harvesting of visible light.4+) are generated under irradiation conditions, low temperature EPR spectra were recorded on Nb2O5 nanorod samples . Without any illumination, no EPR signal associated with Nb4+ could be observed.2O5 nanorod sample had been illuminated under a 300 W Xe lamp, an EPR signal with g = 1.9 was detected, which is diagnostic of Nb4+. This confirms that Nb4+ species can be formed through the reduction of Nb5+ under irradiation.Figure 2 on the surface of the metal oxide support, thereby promoting the formation of reaction intermediates and products.2O5 under illumination. Firstly, for 3h tests with just CO2 at a reaction pressure of 27 psi, in the absence of H2, no CO was detected. Furthermore, 3 h tests with CO2 and H2 at 27 psi and a 1:1 gas ratio carried out in the dark did not yield any CO product. Based on the reaction rate measurements presented in Figure 2 is dissociated on the photothermally heated Pd nanocrystals and so‐formed H migrates to the surface of Nb2O5. Concomitantly, Pd nanocrystals behave as local nanoscale heat sources enabling the photothermal effect on Nb2O5 nanorods to occur, where adsorbed CO2 and H2 based surface intermediates transform to CO and H2O. When oxygen vacancies and/or surface reduced niobium states are formed, the affinity of the surface of Nb2O5 for CO2 is increased. Moreover, the existence of oxygen vacancies and reduced niobium oxide surface states favour the decomposition of intermediates and are found to increase the CO production rate by a factor of 2.6.It has been demonstrated in the literature that oxygen vacancies and reduced oxidation state surface metal sites can increase the proclivity towards the adsorption and dissociation of CO32O5 nanorod‐supported Pd nanocrystals towards the reduction of CO2 to CO in the presence of H2 has been demonstrated. By powering this photothermal reaction with visible and near‐infrared light, conversion rates as high as 1.8 mmol gcat−1 h−1 have been achieved.In this study, the efficient photothermal catalytic activity of Nb2O5 support.By comparing the intensity and spectral power dependence of the RWGS conversion rate, it was found that the photothermal catalytic effect originates from intraband and/or interband optical excitation and nonradiative relaxation of Pd nanocrystals rather than being driven by ultraviolet plasmon excitation of Pd nanocrystals or electron–hole pair generation upon absorption of UV photons in the Nb2O5 nanorod supports and also provide a valuable probe of the local temperature. The results of this study also suggest that the observed RWGS reaction rate enhancement by a factor of 2.6 on nanostructured Pd@Nb2O5 that occurs during catalytic measurements is associated with in situ generation of oxygen vacancies and/or reduced niobium oxidation states on the surface of the Nb2O5 nanorods, which provides an additional driving force to promote the adsorption and reaction of CO2 with H2 to form CO and H2O.The Raman measurements presented in this study demonstrate the photothermal effect of the Pd “nanoheaters” on the Nb2 to valuable chemicals and fuels.It is hoped that the knowledge gleaned from this investigation will contribute to the design, synthesis and optimization of highly efficient photothermal catalysts that use the full solar spectrum to convert gas‐phase CO3O7(OH) NanorodsSynthesis of Nb: Niobium powder  was dissolved in a hydrochloric acid solution  in a pyrex beaker. The aqueous solution was ultrasonicated for half an hour and then stirred for 15 min. The solution was subsequently placed in a Teflon‐lined stainless steel autoclave with 100 mL capacity. The hydrothermal reaction was performed at T = 200 °C for 24 h.T = 70 °C for 12 h.2O5Synthesis of Nanostructured Pd@Nb: Nanostructured Pd@Nb2O5 was synthesized via a microwave‐assisted reaction. Typically, 50 mg of Nb3O7(OH) nanorods were suspended in anhydrous ethanol (20 mL) in a pyrex vessel with 40 mL capacity. A stock solution of Na2PdCl4  was prepared in anhydrous ethanol (20 mL), the concentration of which is 1 mg mL−1. 0.25 mL of Pd precursor solution, equivalent to 0.5 wt%, was added to the dispersion of 50 mg Nb3O7(OH) under sonication. After 30 min of sonication, the vessel was capped and transferred to the microwave reactor . After centrifugation and being washed with deionized water, the sample was placed into a vacuum oven at T = 70 °C for 24 h. The dried Pd@Nb3O7(OH) sample was then treated at T = 120 °C in air for 24 h to obtain the final Pd@Nb2O5 nanostructured samples ready for characterization and catalytic testing.Characterization: Sample morphology was determined using a Hitachi H‐7000 transmission electron microscopy (TEM) at 100 kV. The scanning transmission electron microscopy (STEM), high angle annular dark field (HAADF) imaging, and electron energy‐loss spectroscopy (EELS) analysis were performed using an aberration‐corrected Nion UltraSTEM‐200, equipped with a cold‐field emission gun. The Nion UltraSTEM‐200 was operated at 200 kV when performing HAADF imaging and EELS mapping. EEL spectra were collected using a Gatan Enfinium Dual EELS spectrometer. Principal component analysis (PCA) was performed on the EELS data to remove random noise.Powder X‐ray diffraction (PXRD) was performed on a Bruker D2‐Phaser X‐ray diffractometer, using Cu Kα radiation at 30 kV.The surface area of each sample was determined through volumetric nitrogen adsorption at 77 K using a Quantachrome Autosorb‐1‐C and calculated using Brunauer‐Emmet‐Teller (BET) theory. Samples were outgassed overnight at 80 °C prior to being analyzed.2O5, Nb3O7(OH), and Nb2O5 samples onto borosilicate glass microfiber filters.Diffuse reflectance spectra of the samples were measured using a Lambda 1050 UV/VIS/NIR spectrometer from Perkin Elmer equipped with an integrating sphere with a diameter of 150 mm. The film samples were prepared by drop‐casting aqueous dispersions of Pd@Nb2O5 and Nb2O5 samples onto p‐doped Si(100) wafers. All measurements were conducted in an ultrahigh vacuum chamber with a base pressure of 1 × 10−9 Torr. Data analyses were performed using the Multipak program and all binding energies were referenced to the NIST‐XPS database and the Handbook of X‐ray Photoelectron Spectroscopy.35,36X‐ray photoelectron spectroscopy (XPS) was performed using a Perkin Elmer Phi 5500 ESCA spectrometer. The spectrometer uses an Al Kα X‐ray source to generate X‐rays with an energy of 1486.7 eV. The samples used in XPS analyses were prepared by drop‐casting aqueous dispersions of Pd@Nb2 adsorption capacity was determined using thermogravimetric analysis  with the following procedure. Each sample (≈5–10 mg) was first heated at a rate of 10 °C min−1 to a final temperature of 120 °C, under a nitrogen flow of 100 mL min−1. This temperature and nitrogen flow was maintained for 1 h to remove any adsorbed moisture. The gas flow was then switched to carbon dioxide, maintaining the same flow rate and temperature for 2 h. The observed weight gain between the N2 and CO2 gas streams was directly used to calculate CO2 capacity, and normalized against the surface area of each sample.CO2, at a rate of 100 mL min−1. The samples were heated at a rate of 1 °C min−1 to 700 °C. The same TGA experiments were also performed using air since it is known to efficiently remove carbon deposits.TGA was performed to determine whether or not carbon deposits reside on the used catalyst.λ = 633 nm, with r = 400 nm) and a CW diode laser . The maximum laser power available using the HeNe laser was 12 mW. The samples were illuminated using a 100× objective. For the Raman spectra, which were obtained as a function of laser power, the exposure time varied from 1 to 40 s, with averaging over 4 spectra depending on the power used. The exposure time was increased as the incident power was reduced to maintain a signal to background ratio conducive to accurate peak fitting. Prior to fitting the Raman peaks, the spectra were identified and isolated from the measured data by subtracting a polynomial curve fit to the background signal. The stokes/anti‐stokes spectra were obtained using the 785 nm laser with a power of 24 mW. The samples for Raman measurements were prepared by drop‐casting aqueous dispersions of Pd@Nb2O5 and Nb2O5 samples onto borosilicate glass microfiber filters.The Raman spectra were obtained using the Horiba Jobin Yvon confocal system; namely a LabRAM HR 800 spectrometer equipped with a continuous wave HeNe laser  measurements were performed at 20 K using a Bruker ECS‐EMX X‐band EPR spectrometer equipped with an ER4119HS cavity. An Oxford ESR 900 helium cryostat controlled by an Oxford ITC503 temperature controller was utilized. Typical operating parameters were as follows: microwave frequency 9.381799030 Ghz, microwave power 21.63 mW, modulation amplitude 4 G, sweep width 512 G centered at 3455.45 G, time constant 0.01 ms, total sweep time 512 s, 4096 points were acquired. The EPR analysis was applied to nanostructured NbGas‐Phase Catalytic Measurements: For gas‐phase photothermal catalytic testing, samples were prepared by drop casting nanostuctured Nb2O5 and Pd@Nb2O5 from an aqueous dispersion onto binder free borosilicate glass microfiber filters having an area of ≈2.2 cm2 .2 (99.9995%) at a flow rate of 20 mL min−1. After purging, the reactor was infiltrated with H2 and CO2 gas at a 1:1 pressure ratio to a total pressure of 27 psi prior to being sealed. The pressure inside the reactor was monitored during the reaction using an Omega PX309 pressure transducer. Reactors were irradiated with a 300 W Xe lamp for a duration of 3 h. The spectral output from the 300 W Xe lamp was measured using a StellarNet Inc spectrophotometer and the power of the incident irradiation was measured using a Spectra‐Physics Power meter (model 407A). The radiant power absorbed by the sample for each run was determined using the equation P=∫Pλ⋅tλ⋅aλ dλ; where λP and tλ, are the power emitted by the Xe lamp and the fraction of light transmitted through the quartz window as well as any filters used in the measurement, respectively. The light absorbed by the sample was calculated using the equation λ = 1 – Rλa, where λR is the diffuse reflectance measured at wavelength λ. The reactors were not connected to a heating source for the photothermal tests, although the reactor temperature increased to 50 °C, over the 3 h duration of the test due to the heat generated from the incident irradiation . For thermal tests in the dark, the reactor temperature was controlled by an OMEGA temperature controller combined with a thermocouple placed in the reactor. Product gases were analyzed with a flame ionization detector (FID) installed in a SRI‐8610 gas chromatograph (GC) with a 6′ Haysep D column. Isotope tracing experiments were performed using 13CO2 . The reactor was evacuated prior to being injected with 13CO2 followed by H2. Isotopically labeled product gases were measured using an Agilent 7890A gas chromatographic mass spectrometer (GC‐MS) with a 60 m GS‐Carbonplot column fed to the mass spectrometer.The gas‐phase photothermal catalytic measurements were conducted in a custom‐built 1.8 mL stainless steel batch reactor with a fused silica view port sealed with a Viton O‐ring. The reactor was evacuated using an Alcatel dry pump prior to being purged with HAs a service to our authors and readers, this journal provides supporting information supplied by the authors. Such materials are peer reviewed and may be re‐organized for online delivery, but are not copy‐edited or typeset. Technical support issues arising from supporting information (other than missing files) should be addressed to the authors.SupplementaryClick here for additional data file.


Rechargeable Li–O2 batteries have amongst the highest formal energy and could store significantly more energy than other rechargeable batteries in practice if at least a large part of their promise could be realized. 2 batteries have amongst the highest formal energy and could store significantly more energy than other rechargeable batteries in practice if at least a large part of their promise could be realized. Realization, however, still faces many challenges than can only be overcome by fundamental understanding of the processes taking place. Here, we review recent advances in understanding the chemistry of the Li–O2 cathode and provide a perspective on dominant research needs. We put particular emphasis on issues that are often grossly misunderstood: realistic performance metrics and their reporting as well as identifying reversibility and quantitative measures to do so. Parasitic reactions are the prime obstacle for reversible cell operation and have recently been identified to be predominantly caused by singlet oxygen and not by reduced oxygen species as thought before. We discuss the far reaching implications of this finding on electrolyte and cathode stability, electrocatalysis, and future research needs.Rechargeable Li–O These include replacing the graphite anode by Li metal and intercalation cathodes by the O2 cathode to form the Li–O2 (air) battery, which is considered the battery with the highest specific energy. The O2 cathode comprises a porous, electrolyte filled electron conducting matrix, wherein O2 from the ambiance is reduced on discharge to form Li2O2 and the reverse process on charge.Raising energy storage beyond the limits of current battery technology has become a societal demand and focus of much Frontier research. The achievable limits of Li-ion batteries with respect to energy, material sustainability and cost will likely not satisfy the needs. This motivates ambitious approaches with ‘beyond-intercalation chemistries’.2 cell could store, despite it being the motivation for the research. The problem arises from confusing formal capacity  with theoretical capacity  and achieved true capacity (Li2O2 including the used electron and ion conductor).4 We thus shall first discuss realistic performance metrics.There is lots of ambiguity of what energy a Li–O2 battery combines two challenging electrodes. In most cases a Li metal anode is used, which is, despite decades of research, still associated with poor coulombic efficiencies.3 Other high capacity anodes such as Si may also be considered but likely require in either case protection against O2 ingress from the cathode. Until a couple of years ago there was hardly any knowledge on the O2/Li2O2 redox couple in aprotic media. Reactive species involved in the cycling mechanism, which challenge the stability of electrolyte and electrode material, turned out to be another critical research direction. Next to these scientific and materials challenges practical realization further faces engineering challenges with cell construction and air handling. Only understanding the science may thus form the foundation for tackling the engineering.The Li–O2 cathode. After dealing with performance we discuss the current understanding of Li2O2 formation and decomposition on cycling, followed by measures of reversibility, mechanisms that degrade electrolyte and electrode components and porous cathode design. Potentially transformative ideas start with much enthusiasm and hyped expectations. Thereafter, illusions of low hanging fruit fall and only going the extra mile for true understanding can allow progress to continue. The field of Li–O2 batteries is now in the latter stage. Real progress has been achieved with mechanistic understanding in the last years, which now puts us in a better position than ever to state that the disillusioned view of Li–O2 never leaving the state of a cell with low capacity, rate, energy efficiency and cycle life is too pessimistic. Yet it is unclear whether it can eventually lead to a technology.In this perspective we focus on the science underpinning the O2.5When performance is the argument for research work then data need to stand up to it. An alleged 5 to 10 fold theoretical higher specific energy in comparison to current LIBs is often found as justification when papers on the topic are introduced. However, these numbers stem from very simplistic views and are not realistic even in theory.2 cathode this means that the full volume of Li2O2 forms/disappears during discharge/charge. The basic charge storage process at the cathode is linking the redox moiety O2 to electron and ion transport according to O2 + 2e– + 2Li+ ↔ Li2O2. However, Li2O2 can, akin to most storage materials, not be cycled anywhere near to the bulk substance. Ion and electron transport are too poor to allow for practical bulk material electrodes. To provide simultaneous contact with ionic pathways to the electrolyte and electronic pathways to the current collector an electrolyte filled porous cathode is used. The capacity at a given initial porosity is determined by the degree of pore filling. Thus, beyond the scale of the single Li2O2 particle, the electrode including electrolyte becomes the actual Li+ host structure, which is required to fulfil the charge storage processes of linking formal ion host particles (Li2O2) to electron and ion transport. This introduces a ‘super-host structure’ that becomes an indispensable and integral part of the cell chemistry in a given electrode architecture and needs to be accounted for when reporting performance.Departing from the intercalation concept of LIBs does generally not allow for a stable framework in the active material. In the Li–O2 cathode with a reasonable super-host structure at best achieve and how does it compare to intercalation chemistries? One can assume that the Li2O2 particles could at best be packed with 74% volume occupation into a face centred cubic (fcc) structure, the theoretically limiting case. When charged the porous electrode is filled by electrolyte, which is displaced upon Li2O2 growth. total–1, which is higher than intercalation electrodes.What capacity could the Li–O2 batteries are special in that the positive electrode does not contain the redox material in the charged state that could be taken as a reference for capacity. Thus, it is convenient to report capacities per weight of porous electron conductor. Often, up to several 10 000 mA h gcarbon–1 are reached as first discharge capacities, which compare superficially favourably with some 100 mA h g–1 for intercalation materials.7 As a result of difficulties to cycle at full capacity it has become a habit to report cycling at, e.g., 1000 mA h gcarbon–1, which may still seem a lot in comparison to LIBs. Formal capacities are, however, easily misjudged since true capacities strongly depend on initial porosity and thus the substrate/electrolyte ratio as indicated by the vertical dotted line in carbon–1 grows with decreasing initial porosity due to the growing electrolyte/Li2O2 ratio at shallow discharge, Metal–O4Limited-capacity cycling often allows simulating a large possible cycle number even if the same cell at full discharge would not reach more than a few cycles and cumulative capacity equating to only a few full cycles. Clearly, overly capacity-limited cycling is not suitable to demonstrate large reversible capacities for many cycles. Yet, it is a common feature of beyond-intercalation chemistries that reasonably capacity-limited cycling can be enabling for cyclability and at the same time yield significant improvement over intercalation if the capacity on a total weight basis is kept in mind as shown in 2O2 and to minimize inactive mass and volume. Low packing density and overly restricted depth-of-cycling likely result in no advantage over intercalation electrodes as demonstrated in For higher true capacities than intercalation electrodes it is crucial to achieve an as high as possible packing density of Li(1) The thickness of the porous electrode.(2) The mass fractions of all components  as obtained from the mass fractions of all solids, their loading and the loading of electrolyte per unit electrode area.(3) The volume fractions of all electrode components are then obtained from the mass fractions and the densities.(4) With these measures it is straightforward to convert the capacity with respect to substrate into true capacity per mass and volume of total electrode including electrolyte.2 cathode to achieve higher true capacity than intercalation cathodes also in practice; key is high active material packing density and a small inactive/active material ratio.These parameters are easily obtained and journal editors and referees are urged to insist on them being provided for papers. There is no theoretical barrier for the Li–O3.2O2 is forming/decomposing attains paramount importance. It further directly impacts the stability of the cell components and rechargeability via the reactivity of the intermediates.With true energy depending crucially on filling of the available pore space, the mechanism by which Li3.1.2 reduction in aprotic Li+-electrolytes results in superoxide (O2–), which associates with Li+ and in the second step either undergoes a second 1e– reduction or disproportionates to form Li2O2. Two mechanisms have been proposed for how these steps proceed. The first involves a solution process, where O2– is solubilized to precipitate Li2O2 from the electrolyte solution,8 and the second considers the intermediate as surface bound throughout the process.10 Recently, a unified mechanism was described, where the solution and surface mechanism, respectively, are limiting cases.11 It describes the partition between these cases by the solubilisation of LiO2 in the equilibrium14 O2– solvation, correlated with the solvent acceptor number (AN), is usually weaker.13 The typical classes of electrolyte solvents span a wide range of DN from nitriles and sulfones (DN = 14–16), via glymes (DN = 20–24), amides (DN ∼ 26), sulfoxide (DN ∼ 30).14The first step of O+-coordinating salt anions can shift above equilibrium to the right to a similar extent as the solvent DN by competing association with the Li+ ion between salt anion and O2–.18 Similarly, Lewis acidic additives enhance O2– solvation as shown with water, alcohols, and onium cations.22 Unfortunately, both high DN solvents and protic additives that favour the solution mechanism can enhance parasitic reactions.24Lewis basicity (DN) or acidity (AN) of the electrolyte solution can be influenced by additives. Strongly Livia a second 1e– reduction or disproportionation  accelerate in the same electrolytes the second reduction at the expense of disproportionation and change the mechanism to surface growth. Recently, it was suggested that not only effective solvation strength controls surface/solution growth based on a correlation between effective solvent polarity (ENτ), tuned by additives, and the onset of the second reduction in DMSO.27 This was explained by an increasing solvent rearrangement energy with polarity and thus, according to Marcus–Hush theory, an increasing activation barrier for the second reduction.Irrespective of whether the surface or solution pathway prevails, the second electron transfer may either proceed ionation . With a 2 species and leads to a conformal coating of the electrode with discharge ceasing after only ∼5 to 10 nm, corresponding to low capacity.25 Beyond this thickness the charge transport resistance increases greatly as determined by impedance spectroscopy and does not permit sustaining the current any longer.28 The solution mechanism, in contrast, keeps electrode area open for longer and allows for larger capacity by the growth of large (micrometer sized), toroidal particles composed of lamellae, which can fill larger pores to a bigger extent  slightly above the OCV at ∼3 V and ever increasing voltage as charging progresses. Three underlying phenomena appear to be consolidated although details are still under debate: (1) electrochemical oxidation of Li2O2 is possible with low kinetic barriers at high rates; (2) increasingly difficult electron transfer along recharge contributes a minor fraction of the voltage rise; (3) rising voltage is mostly caused by accumulating parasitic products, which cause a mixed potential.Galvanostatic charging of Li–O+, e–, and O2 can be removed from Li2O2.33 They suggested that Li+ and e– can be removed starting below 0.2 V overpotential, leading to either surface  or bulk Lix2–O2via topotactic delithiation.32 Bulk Lix2–O2 appears to comprise Li2O2 and LiO2 domains.32 O2 evolution was initially suggested to have the highest barrier along the whole path.31 Later, a pathway was shown where O2 evolves facile after progressive delithiation at ∼0.3 V overpotential via Lix2–O2 to LiO2, which then decomposes rapidly at ∼2.6 V vs. Li/Li+  One mole of O2 and two moles e– produce exactly one mole of Li2O2 on discharge. On charge two moles e– consume one mole of Li2O2 and release one mole of O2. Thus:(B) One mole of O–/O2 = 2 is not a strict requirement for a rechargeable Li–O2 battery if Li2O2 is not the discharge product as occasionally claimed.50 For example, if Li2O2 with a certain fraction of ‘LiO2-like’ species or even pure LiO2 is the product e–/O2 may be lower than 2. However, in any case e–/O2 = e–/LixO2 must be identical on discharge and charge. These conditions further imply:The ratio e2 evolves during discharge and charge and no soluble or solid product other than Li2O2 (or H2O2) is produced.(C) All electrons involved contribute to the oxygen reduction reaction (ORR) or oxygen evolution reaction (OER). Thus, no other gas than OQORR = QOER) the O2 released on charge matches the amount consumed, thus n2,ORRO = n2,OERO.(D) For cycling with equal capacity on discharge and charge  to (D).Importantly, none of these measures can be taken for granted to be mutually met even if, for example, ee.g., 1000 mA h gsubstrate–1, thus forcing Qcharge = Qdischarge. Truncating discharge reasonably to avoid full electrode blockage may enable cyclability and appears justified as long as the true capacity based on the total weight is obeyed ; and (b) that as the capacity approaches full recharge the depletion of the Li2O2 must cause the voltage to rise ever steeper before it transits into a plateau at the electrolyte oxidation potential, blue curve in Q/dU must approach zero at full recharge irrespective of whether the preceding discharge was limited by voltage or capacity, As discussed in Section 2 on performance metrics, it has become a habit to cycle cells at fixed discharge/charge capacities of, eyed see . CapacitQ/dU remaining high, 2O2 oxidation, and would continue with a continuing flat plateau if such a cell were over charged (extrapolation by the red dotted curve). An exception is cells with oxidation mediators with large electrolyte-to-Li2O2 ratio. Notably, remaining within the oxidative stability window (without Li2O2) does not ensure absence of oxidative electrolyte decomposition nor does a lower recharge plateau per se  indicate less parasitic chemistry.51Frequently observed load curves of the type as shown in the red curve in 2 consumption/evolution has been described by two methods: (A) quantitative operando online mass spectrometry (MS), where the cell head space is continuously or intermittently purged to a MS.52 Using an O2/Ar mix allows also quantification of O2 consumption and of any other gases evolved on discharge.53 (B) Measuring the pressure in a closed cell head space.54 Peroxide or superoxide content of electrodes has been measured ex situ using either iodometric titration49 or spectrophotometry using the coloured [Ti(O22–)]2+ complex.55 This method was combined with equally MS based quantification of Li2CO3 and organic products by treatment with acid and Fenton's reagent to separately evolve CO2 from inorganic and organic compounds.44 The latter may also be quantified by 1H-NMR after immersing the electrode in D2O which further allows for speciation of the compounds.49 Importantly, all these methods capture the integral electrode. Qualitative spectroscopic or microscopic methods such as Raman, FTIR, XRD, XPS, or SEM cannot replace the mentioned or similar quantitative integral methods and cannot support claims of reversibility.Concluding about reversibility by the measures (A) to (D) requires multiple quantitative analyses. Measuring O4.2.57 Primary species are O2– and O22–, which in protic environments form species including HOO˙, HOO–, and HO˙. Their nucleophilicity, basicity, and/or radical nature cause reactivity via three major routes: nucleophilic substitutions, H+ and H-atom abstraction. The latter may also proceed with dioxygen in auto-oxidation reactions. Additionally, O2– can transfer electrons.Reduced reactive oxygen species (RROS) are well known for their reactivity with a wide range of organic substrates, which has both been used as a reactant and recognized as a source of unwanted reactions.via heteroatoms to dissolve a Li-salt. The polarity in turn makes adjacent C and H atoms reactive. After the complete failure of carbonate electrolytes for Li–O2 chemistry was rightly associated with the nucleophilicity of O2–, this and the other reactivities of RROS and O2 were taken to explain lesser but significant parasitic chemistry of all so far investigated alternative solvents and cell components.58 In the era of computational chemistry the likelihood for this assignment can be judged based on activation and reaction free energies. Bryantsev et al. have pioneered this field for O2– and O2 reacting with organic electrolytes via nucleophilic substitution, H-atom, and H+ abstraction.61 Their data for activation energies together with those of other researchers are summarized in Polarity is introduced to aprotic electrolyte solvents 2 exposure or the reversibility of the O2/O2– couple, reactions with activation energies beyond 100 kJ mol–1 can be considered not to contribute noticeably.58 Hence, only esters and lactones are expected to react via nucleophilic substitution with O2– and possibly ethers via H-abstraction with Li2O2.62 With ethers, for example, all pathways with O2–, Li2O2 and O2 require high activation energy and are strongly endothermic. The only exception is one study that found H-abstraction by Li2O2 clusters slightly exothermic with Eact < 100 kJ mol–1.62 Solvent coordination with Li+ was reported to further stabilize against H-abstraction by O2– and O2.63Strikingly, activation energies for all considered reactions involving the major classes of solvents – except for the very unstable solvents like carbonates – are too high to expect these reactions to strongly contribute to decomposition. Based on solvent stability screening experiments with KO2– forms viaeqn (10)–(13) HOO˙, HOO–, and HO˙, which are more reactive than the primary RROS.64In presence of proton sources such as water or weak acids O– is a stronger base than O2– and more readily abstracts protons to form R–. HO˙ could serve as the initiator to form R˙, which undergoes fast and thermodynamically favourable onwards chain reactions in the presence of O2.61 Overall, direct reactivity of O2–, Li2O2 and O2 with the most important classes of non-aqueous solvents for the Li–O2 cathode is unfavourable. Increasing parasitic chemistry with increasing water content is consistent with the protonated species being more reactive.12 Yet, much higher side reactions on charge than on discharge, which opposes superoxide occurrence, points at RROS not to be the prime cause for parasitic chemistry.HOO4.3.2O2 was hypothesised by Hassoun et al. to be able to generate singlet oxygen (1Δg or 1O2), the highly reactive first excited state of triplet ground state dioxygen (3Σg– or 3O2).69 This view was motivated by the known formation of 1O2 by chemical oxidation of H2O2 or alkaline peroxides.70 Based on the reversible potential of Li2O2 formation and the energy difference between triplet and singlet oxygen of ∼1 eV, 1O2 formation in the Li–O2 cell has been considered possible at charging potentials exceeding 3.5 to 3.9 V vs. Li/Li+.71 The idea was picked up by several reports but could not be verified due to the difficulties with detecting 1O2 except for one work, which identified small quantities of 1O2 between 3.55 and 3.75 V and explained it on thermodynamic grounds for the process Li2O2 → O2 + 2Li+ + 2e–.71 It could thus contribute to explaining parasitic chemistry above 3.55 V. Yet it has been found that from the start of charging below 3.5 V both a substantial amount of parasitic products is generated49 and that less than 1 mol O2 evolves per 1 mol Li2O2 consumed.49 Both could not be explained by reactivity of reduced O2 species and formation of 1O2 above 3.55 V.Electrochemically oxidizing Liet al. have shown that 1O2 forms also during discharge and from the onset of charge and that it accounts for the majority of parasitic reaction products.24 The amount of 1O2 increases during discharge, early stages of charge, and charging at higher voltages, and is enhanced by the presence of trace water. They used the 1O2 specific conversion of 9,10-dimethylanthracene (DMA) into its endoperoxide (DMA–O2) to probe 1O2 in the cell. Operando fluorescence detection on discharge and charge has shown rather small 1O2 abundance on discharge and significant 1O2 formation immediately after switching to charging, starting from ∼3 V  as quencher they have shown even more substantial reduction of side products  or (15). This pathway for 1O2 formation can be active from the first onset of charge as soon as Li+ and e– are extracted. Second, a further 1e– oxidation of the surface LiO2 species could give 1O2 above  + E(1Δg ← 3Σg–). With the thermodynamic equilibrium potential  estimated to be between 2.29 and 2.46 V (E(1Δg ← 3Σg–) ∼ 1 eV, a thermodynamic voltage for 1O2 evolution of 3.26 to 3.43 V can be estimated. Finally, above ∼3.55 V the pathway can sets in as suggested by Hassoun et al. and shown by Wandt et al. with 1O2 evolving by 2e––oxidation of Li2O2 (Li2O2 → O2 + 2Li+ + 2e–).71Overall, the disproportionation of superoxide in the presence of either Lid 2.46 V  and E and not solely on the superoxide concentration. These sources and sinks are both electrochemical and chemical and change with discharge/charge, electrolyte, current, and potential. Current density and electrolyte properties will influence the 1O2 formation in much the same way as they govern the occurrence of superoxide on discharge and charge below 3.5 V.77 Charge current will drive 1O2 production if it causes charging voltages above ∼3.5 V.Net formation of 4.4.2O2 with remaining stable LiO2 species, resembling the Lix2–O2 intermediate on charge.50 Based on these findings an Ir–graphene based cathode was reported to cycle in ether electrolyte via crystalline LiO2.7 These assignments were mostly based on Raman spectra that can distinguish O–O stretch vibrations in Li2O2 and LiO2. However, it was recently shown that PVDF binder decomposition can lead to vibrations mimicking those of LiO2, thus concluding that cycling was not based on LiO2.78In occasional reports the discharge product was reported to be Li2O2 as the main discharge product rather than LiOH as one could intuitively assume.19 Instead, water promotes discharge via the solution mechanism. Whether Li2O2 or LiOH forms was suggested to be governed by the effective pKa value of water in the electrolyte.20 A value of 35 in MeCN compared to 47 in DME was used to explain LiOH to form in the former and Li2O2 in latter. At water concentrations beyond 1% LiOOH was shown to form together with LiOH.79 LiOH was also found to form in a 4e–/O2 reduction in presence of LiI.80 Unfortunately, O2 evolution from LiOH could so far not be shown and apparent cyclability must be accounted for I– electrochemistry and parasitic chemistry.80Surprisingly large water contamination up to 1000s of ppm in ether electrolyte has been shown to still lead to Li2CO3 not only forms as a result of parasitic reactions, but also when the O2 supply is CO2 contaminated.82 CO2 has a high barrier for direct reduction but reacts readily with O2– along with the formation of Li2O2. If Li2CO3 could be decomposed on charge to the educts it would make the cell insensitive to CO2 in the O2 supply. While Li2CO3 can be decomposed from ∼3.8 V, it does not evolve O2 together with the CO2, which suggests that reactive intermediate form that decompose the electrolyte.82 Apparent cyclability of O2/CO2 cells that was shown in some cases was so far not compellingly associated with reversible chemistry. Making Li–O2 chemistry insensitive to H2O and CO2 contamination should thus remain a high priority, which foremost calls for rigorously investigating the associated parasitic chemistry.Li5.5.1.2–cathode due to low cost, high conductivity, and easily tuneable surface area and pore sizes. Carbon was, however, found to decompose itself on cycling and to promote electrolyte decomposition.83 Using 13C carbon black and MS analysis of the gaseous and solid products at various stages of cycling, carbon was observed to be relatively stable on discharge despite thermodynamic instability in contact with Li2O2; the majority of side products stems from the electrolyte. From the onset of charge, however, carbon decomposes to form Li2CO3 with increasing rate as the potential grows.49 Defect rich hydrophilic carbon is both much more vulnerable itself and promotes more strongly electrolyte decomposition during discharge and charge than hydrophobic carbon.83Carbon is the most common porous electron conducting matrix for the O2– attack. The perfectly opposing trend of O2– abundance – highest on discharge and ever decreasing as charge voltage grows – to decomposition rates makes this interpretation unsatisfactory. Therefore, reactive intermediates on oxidizing Li2O2 have been suggested.45 Carbon and electrolyte decomposition rates both follow the trend of 1O2 abundance as shown in 1O2 being the dominant driver of parasitic chemistry; possibly the nearly exclusive one on charge.As with electrolyte decomposition, the carbon instability was related to O53 For TiC a thin passivating layer of TiOx2– and TiOC has been identified to be critical for stability and conductivity.84 The metallic Magnéli phase Ti4O7 was equally shown to form surface TiOx2– and to allow for cyclability similar to Au and TiC.85 Limited binder stability adds another dimension to cathode design. Standard PVDF binder as used in LIB was found to react with Li2O2 and KO2,87 but equally 1O2 can be expected to contribute. More stable alternatives include PTFE and Nafion.87Given the instability of carbon, alternative corrosion resistant materials have been sought that at the same time do not promote electrolyte decomposition. They include Ti ceramics and nanoporous Au that allow for more stable cycling.2O2. High porosity becomes even more important when going from C to much denser metals or ceramics. So far the surface chemistry and electrochemistry of these alternative materials have been investigated with bound nanoparticles forming low porosity electrodes. Achieving high capacity based on the total electrode weight requires, however, filling highly porous electrodes to a large extent with Li5.2. = 2.96 V are ∼0.3 V on discharge and ever rising values on charge from nearly zero to, in some cases, up to 2 V. Inspired by aqueous O2 electrochemistry, these overpotentials were accounted to sluggish kinetics and evoked substantial efforts in finding efficient electrocatalysts including noble metals, transition metal oxides, and doped carbons.51 Considering the mechanism for Li2O2 formation/decomposition, there are, however, three major barriers questioning true effectiveness of solid electrocatalysts. First, ideal Li–O2 chemistry does not involve O–O bond breaking, the step necessitating catalysts in aqueous media. O2 reduction to superoxide and peroxide, in contrast, has been shown to be facile even on bare glassy carbon.33 Hence, perceived overpotentials are not due to ORR or OER kinetics. The discharge potential of ∼2.7 V is not due to overpotential of the O2/Li2O2 couple but is pinned there by the reversible potential of the O2/O2– couple, the first step on O2 reduction to Li2O2 and can thus not be raised by electrocatalysts.88 Li2O2 oxidation starts with nearly zero overpotential at ∼3 V. Rising charge potentials are, as discussed in Section 3.2, caused partly by increasingly difficult electron transfer in the receding Li2O2 and mostly by parasitic chemistry. Second, conventional electrocatalysis would only act on dissolved redox species. On discharge there would thus be no effect beyond monolayers of Li2O2 forming with the surface mechanism. Deep discharge with both surface and solution mechanism blocks eventually all electrochemically active surface. Third, charge transport limitations through forming Li2O2  may alter the delithiation kinetics of Li2O2 by forming Li-transition metal-oxides.77 Acting beyond the initial stages of charge requires maintaining the electrode/Li2O2 contact, for which the driving force is not clear as the Li2O2 closest to the electrode will necessarily be oxidized first. In the case of large Li2O2 deposits forming by the solution mechanism, this contact may never be fulfilled for a large fraction of the Li2O2. Not least did the habit of extended cycling at a small fraction of a possible single discharge capacity (see Section 2) arise from catalyst studies; with deep discharge the same cells fail typically very quickly, which hints at very limited effectiveness of the catalysts to oxidize large amounts of detached Li2O2.Whether solid electrocatalysts can have any effect on charge depends on the processes at the buried cathode/Li91 The exact pathways are not fully clarified but are at least in part associated with the catalyst's ability to dissociate the O–O bond. Concluding about a catalyst's ability to enhance efficiency and cyclability requires quantitative measures of reactant turnover and parasitic products without which any claim is inadequate (see Section 4.1).Given the paramount importance of parasitic chemistry, electrocatalysts must not catalyse parasitic reactions with the electrolyte or electrode. Unfortunately, materials identified as electrocatalysts do enhance parasitic reactions.6.2 to electron and ion transport according to O2 + 2e– + 2Li+ ↔ Li2O2. However, different to other Li+ storage materials, Li+ and e– transport into/out of the bulk Li2O2 particle is not even required since growth/dissolution occurs in any case at its surface. This unique feature of Li2O2 can be turned into a major advantage in terms of rate capability. It is the several orders of magnitude slower ion diffusivity compared to liquids that makes batteries slow and the very fast charge transport in liquids that makes supercapacitors high-power devices.92 Bypassing Li2O2 for ion and electron transport through a phase, where both are facile may thus enable high-power Li–O2 cells. Li2O2 would then only serve as the charge storage medium. While the liquid electrolyte (where the reaction takes place) provides facile ions transport, moving electrons through the liquid is more difficult. Possibilities involve: (a) giving solubility to LiO2 for it to act as an electron mediator during discharge ) were shown to form Li2O2 in an inner sphere process without the involvement of free superoxide .101 Early work has found that some oxidation mediators with suitable redox potentials oxidize Li2O2 with the expected amount of O2 evolution whereas others with similar potential evolve considerably less O2.98 As mechanistic descriptor for this behaviour, the position of the HOMO level of M+ was put forward, which, when close to the HOMO of the electrolyte, is prone to oxidize the electrolyte.99Oxidation mediators allow, in principle, charging at nearly zero overpotential and numerous oxidation mediators have been explored for redox potential and O+ may, instead of oxidizing Li2O2, diffuse out of the cathode to the counter electrode to cause leak-current by shuttling. This may be avoided by using Li+ conducting diffusion barriers such as ceramics or polymers as the separator.101 Second, mediators, which are mostly organic molecules, are themselves susceptible to decomposition. Both issues make it imperative to quantitatively measure O2 consumption/evolution, Li2O2 formation/oxidation, and parasitic products as discussed in Section 4.1. Any claim about performance improvements is inadequate without these measures.Two pitfalls have to be considered with mediators: first, oxidation mediators M7.2 cathode. The two central issues are: (A) discharge/charge mechanisms of Li2O2, and (B) mechanisms of parasitic chemistry.The past few years have brought substantial progress with the mechanisms underpinning the operation of the Li–Ovia a surface passivating mechanism (giving low capacities) or a solution based process to form large Li2O2 particles (required for high capacities). The deciding factor is the solvation of the superoxide intermediate by tuning the electrolyte interaction with Li+ or O2–via solvent, salt, and additive Lewis basicity/acidity. Oxidation of Li2O2 proceeds at low kinetic overpotentials and can thus, in principle, take place at high rates close to the thermodynamic potential. Rising voltage is predominantly associated with parasitic chemistry. Nevertheless, when discharge proceeds via the solution mechanism charge transport from large Li2O2 particles will contribute to overpotentials.A central issue was to identify conditions leading to discharge 2 chemistry appears to be the way forward for high capacity and rate capability and low overpotentials, here, Li2O2 only serves as the storage medium and is bypassed for charge transport through the electrolyte by means of redox mediators. With reduction mediators the pathway to form Li2O2 may be altered such that there is not free superoxide, which is a source of singlet oxygen and thus parasitic chemistry on discharge. Oxidation mediators allow, in principle, for charging at nearly zero overpotential. The biggest open question with mediators is their own susceptibility to decomposition and their impact on singlet oxygen formation.Solution based Li–O2O2 being the major cause was only recently overturned by finding that singlet oxygen (1O2) is formed on discharge and charge; the extent matches the pattern of parasitic reactions with relatively little on discharge and much more on charge. Practical realization of Li–O2 batteries will, in our opinion, stand or fall with mastering 1O2 formation. Open questions centre around: (1) factors influencing 1O2 formation including catalysts, electrolytes, mediators, and protic additives; (2) more detailed insight into formation mechanisms; (3) finding efficient quenchers; (4) finding mechanism to prevent 1O2 formation. Given that so far significant parasitic chemistry is to be expected during both discharge and charge, concluding about the efficacy of any measure to improve capacity, efficiency and cyclability requires quantitative analysis of reactant turnover and parasitic products without which any claim of improvement is inadequate.The major barrier for reversible cell operation is parasitic chemistry with electrolyte and cell components. The previous view of superoxide and Li2 cathode to achieve much higher capacity than intercalation cathodes. However, to do so it is crucial to achieve an as high as possible packing density of Li2O2 in the cathode and to minimize the inactive mass and volume including the electrolyte. Two habits make tracing progress in the field difficult: first, reporting capacity with respect to porous substrate mass, which represents a minor and widely varying fraction of the total electrode mass; second, reporting cycling at, e.g., 1000 mA h gcarbon–1, which may still seem a lot in comparison to intercalation electrodes. The problem is that in most cases true performance is below intercalation electrodes and that it masks irreversible reactions. Therefore, it is important to report performance with respect to the full electrode to allow for a fair assessment of energy, power, and cycle life.There is no theoretical barrier for the Li–O


The Epstein Barr Virus (EBV) has been associated with the autoimmune disease, Systemic Lupus Erythematosus (SLE). EBV nuclear antigen‐I (EBNA‐1) is the major nuclear protein of EBV. We previously generated an IgG monoclonal antibody (MAb) to EBNA‐1, 3D4, and demonstrated that it cross‐reacts with double stranded DNA (dsDNA) and binds the 148 amino acid viral binding site (VBS) in the carboxyl region of EBNA‐1. The aim of the present study was to characterize another antibody to EBNA‐1 that cross‐reacts with dsDNA, compare its immunoglobulin genes to 3D4, and finely map the epitope in EBNA‐1 that is recognized by these cross‐reactive antibodies.We generated an IgM MAb to EBNA‐1, 16D2, from EBNA‐1 injected mice and demonstrated by ELISA that it cross‐reacts with dsDNA and binds the 148 amino acid VBS. We sequenced the variable heavy and light chain genes of 3D4 and 16D2 and compared V gene usage. To more finely map the epitope in EBNA‐1 recognized by these MAbs, we examined their binding by ELISA to 15 overlapping peptides spanning the 148 amino acid domain.H and VL genes but identical JH and Jk regions with minimal junctional diversity. This accounts for similarities in their CDR3 regions and may explain their similar dual binding specificity. Epitope mapping revealed 3D4 and 16D2 bind the same peptide in the VBS. Based on the crystal structure of EBNA‐1, we observed that this peptide resides at the base of an exposed proline rich loop in EBNA‐1.Sequence analysis revealed that 3D4 and 16D2 utilize different VWe have demonstrated that two MAbs that bind EBNA‐1 and cross‐react with dsDNA, recognize the same peptide in the VBS. This peptide may serve as a mimetope for dsDNA and may be of diagnostic and therapeutic value in SLE. Autoimmune diseases such as Systemic Lupus Erythematosus (SLE) are caused by a constellation of genetic, environmental, and hormonal factors. Environmental triggers include pathogens that may elicit an autoimmune response as a consequence of molecular mimicry. Some viruses that have been associated with SLE are herpes zoster, cytomegalovirus (CMV), human endogenous retroviruses and the Epstein Barr Virus (EBV) et al. The association of EBV with SLE originally stems from observations that the prevalence of EBV infection is higher in SLE adolescents than age matched controls, that EBV infection sometimes precedes SLE, and that EBV DNA and EBV viral load are higher in SLE patients than healthy individuals et al. James and Harley Antibodies to double stranded DNA (dsDNA) are the hallmark of SLE and are found in 50–80% of SLE patients. Although dsDNA is a target of anti‐dsDNA antibodies, it is not known whether dsDNA is the eliciting antigen. Screening of phage peptide libraries has identified peptides derived from autoantigens that serve as molecular mimics for dsDNA We previously demonstrated that mice injected with an EBNA‐1 expression vector or the EBNA‐1 protein, developed antibodies to dsDNA We now describe a recently generated IgM MAb to EBNA‐1 (16D2) that also cross‐reacts with dsDNA. Although this MAb utilizes different heavy and light chain genes than 3D4, it recognizes the same carboxyl domain in EBNA‐1. The present study was undertaken to more finely map the peptide epitope in EBNA‐1 that can serve as a peptide mimic for dsDNA. By screening an overlapping peptide library derived from this 148 amino acid domain, we have identified a peptide target of both 3D4 and 16D2. Understanding how a peptide derived from a pathogen can serve as a molecular mimic for dsDNA may have important implications for diagnostic and therapeutic strategies in SLE.2+‐NTA beads.Recombinant EBNA‐1 (rEBNA‐1) lacking most of the G‐A repeat was purified from a baculovirus expression vector kindly provided by Dr. Lori Frappier , as previously described et al. BALB/c mice were injected intraperitoneally with 50 μg of rEBNA‐1 in complete Freund's adjuvant (CFA) 1:1(v/v) and boosted at weeks 3, 7, and 12 with 25 μg of rEBNA‐1 in incomplete Freund's adjuvant (IFA) 1:1(v/v). Mice were serially bled and sera examined by ELISA for antibodies to EBNA‐1 and dsDNA. Three days after the last boost, mouse splenocytes were harvested and fused to NSO cells according to Iliev Escherichia coli expression vector carrying an N‐terminal His tag, as previously described E. coli lysates and purified on Ni2+‐NTA beads according to Yadav et al. PCR fragments encoding the amino and carboxyl regions of rEBNA‐1 were previously cloned into a pET28a 452‐641, EBNA459‐619, EBNA459‐607, were a generous gift from Dr. Lori Frappier E. coli strain BL21 (DE3) and isolated from cell‐lysates on Ni2+‐NTA beads Plasmids (pET15b) carrying an N‐terminal His tag and expressing the following truncated EBNA‐1 proteins; EBNAReady to use Crithidia slides  were immunostained with MAb, 16D2 at 10 μg/ml. Slides were incubated in a moist chamber for 30 min at room temperature, washed extensively in PBS and then immunostained with a 1:1000 dilution of goat anti‐mouse IgM FITC  for 30 min at room temperature. Slides were washed again in PBS and then were examined by fluorescence microscopy using a Nikon Eclipse microscope, TE 2000‐S at 40× magnification.Total RNA was isolated from hybridomas producing MAbs 3D4 or 16D2, using Trizol reagent  according to the manufacturers protocol. First strand cDNA was prepared by RT‐PCR using 3.5 μg of RNA, 50 ng of random hexamers and 200U of superscript III RT  according to the manufacturers protocol.et al. Two rounds of PCR amplification were performed for the Ig heavy and light chain genes using 2.0 μl cDNA (for the first PCR) or 3.0 μl of first PCR product (for the second PCR), 250 μM of each dNTP, 0.5 μM of each primer, and 1.5 U of Hot Star Taq polymerase  in a total reaction mixture of 50 μl. The first round of PCR was performed at 94°C for 15 min followed by 50 cycles of 94°C for 30 sec, 56°C (IgH) or 50°C (Igκ) for 30 sec, 72°C for 55 sec, and a final extension at 72°C for 10 min according to Tiller ELISAs to examine direct antibody binding to EBNA‐1, dsDNA, truncated fragments of EBNA‐1, LPS, Sm, and Proteinase 3 were as previously described Fifteen peptides containing a five amino acid overlap, derived from a 148 amino acid region in the viral dimerization/binding domain in the carboxyl region of EBNA, were synthesized and biotinylated at the amino termini by LifeTein LLC . Purity of all peptides was greater than 80%. Costar ELISA plates  were coated overnight at 4°C with a 1:1000 dilution of streptavidin in PBS (Southern Biotech). The next day, plates were blocked in PBS, 1% BSA, for 2 h at room temperature. Plates were then incubated with 100 μg/ml of biotinylated‐peptide (50 μl per well) for 1 h at 37°C followed by six washes in PBS, 0.05% tween 20. Plates were next incubated for 1 h at 37°C with 5 μg/ml of 3D4 or 10 μg/ml of 16D2 diluted in PBS, 0.1% BSA. Plates were washed and then incubated with a 1:1000 dilution of goat anti‐mouse IgG‐AP or goat anti‐mouse IgM‐AP to detect 3D4 or 16D2 binding, respectively. Plates were developed with 4‐nitrophenyl phosphate disodium hexahydrate as substrate and read at 405 nm on a Titertek Multiscan ELISA reader. To compare the binding of 16D2 or 3D4 to biotinylated EBNA‐1 versus biotinylated PFM‐15 peptide, ELISAs were performed as described above except that following the coating with streptavidin, plates were incubated with mole equivalents of biotinylated PFM‐15 or biotinylated EBNA‐1. EBNA‐1 was biotinylated using EZ‐Link Sulfo‐NHS‐Biotinylation Kit . Following biotinylation, biotinylated EBNA‐1 (B‐EBNA‐1) was applied to a Zeba Desalt Spin Column (Thermo Scientific) and then dialyzed against PBS to remove excess biotin.For peptide inhibition of MAb binding to dsDNA, 3D4 (2.5 μg/ml) and 16D2 (5 μg/ml) were pre‐incubated for 1 h at 37°C with increasing concentrations of peptide PFM‐15 or GC‐15 (concentrations of peptide ranging from 0.1 μg/ml to 100 μg/ml) and then transferred to pre‐blocked dsDNA coated Immulon‐2 plates . Plates were then incubated for 2 h at 37°C, washed, and then incubated with 1:1000 dilution of goat anti‐mouse IgG‐AP or goat anti‐mouse IgM‐AP to detect 3D4 or 16D2 binding, respectively. Plates were developed with substrate as described above. MAbs were used at a concentration resulting in 50% maximal binding to dsDNA.For peptide inhibition of MAb binding to EBNA‐1, 3D4 (0.05 μg/ml) and 16D2 (5 μg/ml) were incubated with increasing concentrations of peptide PFM‐15 or GC‐15, ranging from 0.1 μg/ml to 1 mg/ml and then transferred to pre‐blocked EBNA‐1 coated Costar plates. Plates were then incubated for 1 h at 37°C, and the remainder of the ELISA was performed as described above. MAbs were used at a concentration resulting in 50% optimal binding to EBNA‐1.et al. We previously demonstrated that some mice immunized with EBNA‐1 developed antibodies to EBNA‐1 that cross‐reacted with dsDNA. We generated anti‐EBNA‐1 MAbs from one of these mice and observed that 71% of the anti‐EBNA‐1 MAbs cross‐reacted with dsDNA In the present study, we characterized an IgM MAb to EBNA‐1, designated 16D2, that was generated from somatic cell fusion following injection of a mouse with recombinant EBNA‐1 protein lacking the glycine alanine repeat region, as previously described E. coli plasmid expression vectors as described by Yadav et al. 452‐641, EBNA459‐619, and EBNA459‐607 . 16D2 was shown to bind strongly to all 3 fragments , corresponds to the viral dimerization/DNA binding site (VBS) of EBNA‐1. 3D4 was previously shown to bind strongly to this fragment as well Since 16D2, was observed to bind EBNA‐1 and cross‐react with dsDNA, we wondered whether it would recognize the same epitope in EBNA‐1 as 3D4. We therefore examined 16D2 for reactivity with the amino and/or carboxyl regions of EBNA‐1. Recombinant EBNA‐1 proteins representing the amino and carboxyl regions of EBNA‐1 were purified in our laboratory from Comparison of the relative binding of 16D2 and 3D4 to EBNA‐1 was demonstrated by ELISA. Both antibodies were serially diluted and incubated on the same EBNA‐1 coated ELISA plate and goat anti‐mouse kappa‐AP was added as secondary antibody  and light chain (VL) genes of 3D4 and 16D2 were sequenced to determine their identity. Sequences were analyzed using Ig BLAST (NCBI database) to determine the closest germline genes of origin. Both antibodies were found to utilize different VH and VL genes but the same JH and Jk gene segments  to the germline gene, VH J558.16 and 100% homology to the germline VL gene, 8–30. 3D4 has three mutations in its VH region  and five mutations in its VL region . 16D2 has 12 mutations in its VH region, six of which are silent . The 16D2VL is identical to the germline gene 8–30 . Interestingly, 16D2 and 3D4 utilize the same JH2 segment in their VH regions and the same JK1 in their VL regions. Furthermore, the D‐J joints of both MAbs lack junctional diversity as the carboxy terminal YFDYW amino acids of both CDR3s are JH2 germline encoded  are slightly longer than normal a characteristic of many autoreactive and polyreactive antibodies L CDR3 regions of 3D4 and 16D2 (30 amino acids) are average. The length of the CDR3 regions were determined by counting the amino acid residues after the conserved cysteine at the end of FR3 to the conserved tryptophan‐glycine in all mouse JH segments or the conserved phenylalanine‐glycine in all Jκ segments The variable heavy (V459‐607) were synthesized and biotinylated at the amino terminal end . All peptides were either 15 or 12 amino acids long with a 4–6 amino acid overlap  suggests that amino acids upstream of the VBS, and 5′ of aa 452, are critical for recognition by these antibodies. It is interesting that only the cross‐reactive antibodies target an epitope that resides within the VBS.In the present study, we also generated four MAbs  to EBNA‐1 that do not cross‐react with dsDNA. We were interested in determining the region in EBNA‐1 that these antibodies bind relative to the cross‐reactive antibodies. The isotype used by all four of these antibodies was shown to be IgG1; the same isotype used by 3D4 Fig. A. BindinIn an earlier study, we identified an IgG antibody to EBNA‐1, 3D4, which cross‐reacts with dsDNA. In the present study we identified another antibody to EBNA‐1, 16D2, of the IgM isotype, which also cross‐reacts with dsDNA. We have compared the binding properties, sequences, and epitope specificity of both antibodies to determine if there are similarities that could help explain the structural properties that confer dual specificity.Fine mapping of the target epitopes in EBNA‐1 for 3D4 and 16D2, revealed that they recognize an epitope localized to a 148 amino acid domain in the carboxyl region of EBNA‐1 that coincides with the VBS (aa 459–607). Epitope mapping using overlapping peptides that span the entire148 amino acid domain revealed that both 3D4 and 16D2 bind more strongly to peptide, PFM‐15 (PQPGPLRESIVCYFM) than any other linear peptide in this region. The binding of these MAbs to PFM‐15 is not due to charge interactions as PFM‐15 is hydrophobic and has a net charge of zero. Inhibition assays revealed that 10‐fold more PFM‐15 was required to inhibit 3D4 compared to 16D2 binding EBNA‐1. Furthermore, 3D4 bound equally well to PFM‐15 and the overlapping peptide, GL‐15, while 16D2 displayed significantly less binding to GL‐15 than PFM‐15  which mediates this dimerization The observation that PFM‐15 could inhibit binding of both 3D4 and 16D2 to dsDNA suggests that this peptide can serve as a peptide mimetope for dsDNA at least at high molar concentrations. However, it is not yet clear from our amino acid sequence analyses and known structural data why 3D4 and 16D2 also recognize DNA and what aspects of PFM‐15 serve as the dsDNA mimic. The proline loop of EBNA‐1 has some structural flexibility and it is thought to undergo subtle structural changes when EBNA‐1 binds dsDNA H and VL genes indicating that more than one anti‐EBNA‐1 antibody can cross‐react with dsDNA. 3D4 utilizes a gene from the VH7183 family and 16D2 utilizes a gene from the VHJ558 family. Members of these two gene families are frequently observed among polyreactive antibodies that arise in the early immune response and among anti‐dsDNA autoantibodies H and JK regions and lack junctional diversity at the D‐J joint. They share a stretch of germline encoded amino acids at the carboxyl ends of their CDR3 regions. These may be important contact residues with the antigen. Moreover, both 3D4 and 16D2 contain arginine residues and phenylalanine and tyrosine residues in their VH CDR3 regions and these basic and aromatic amino acids appear to be conserved in proteins that bind dsDNA H of 3D4 and 16D2 are slightly longer than normal , a characteristic of many autoantibodies The sequence analysis of the variable regions of 3D4 and 16D2 reveal differences and similarities that may help explain their shared dual binding properties. 3D4 and 16D2 utilize different VWhile EBNA‐1 has long been associated with SLE, most people who get infected with EBV do not develop lupus. This may be because healthy individuals who are exposed to EBV may preferentially develop antibodies to parts of EBNA‐1 that do not elicit cross‐reactivity, as observed with the non cross‐reactive MAbs that were generated in this study. Studies have shown that SLE patients have an altered humoral immune response to epitopes in EBNA‐1 relative to healthy individuals It is interesting to speculate that other pathogens besides EBV may have proteins containing peptides homologous to PFM‐15 that can elicit anti‐dsDNA antibodies and therefore serve as a peptide mimetope of dsDNA. Protein BLAST was used to search for pathogen proteins containing peptides with homology to a stretch of amino acids in PFM‐15. Of particular interest were proteins that contain the YFM amino acid motif present at the 5′ terminus of PFM‐15. The presence of this motif appears to be critical for recognition by 3D4 and 16D2 since both MAbs failed to bind GC‐15, which shares a common peptide core with PFM‐15 but lacks YFM. The protein BLAST search identified a transporter protein derived from the bacteria Lachnospiraceae bacterium A2, which has 100% homology with seven contiguous amino acids in PFM‐15 (SIVCYFM)  nucleic acid and amino acid sequences of 3D4 and comparison to most homologousVH germline gene as determined by IgBLAST (NCBI database).Figure S2. Variable light chain (VL) nucleic acid and amino acid sequences of 3D4 and comparison to most homologousVL germline gene as determined by IgBLAST (NCBI database).Figure S3. Variable heavy chain (VH) nucleic acid and amino acid sequences of 16D2 and comparison to most homologousVH germline gene as determined by IgBLAST (NCBI database).Figure S4. Variable heavy chain (VL) nucleic acid and amino acid sequences of 16D2 and comparison to most homologousVL germline gene as determined by IgBLAST (NCBI database).Click here for additional data file.