Analytical techniques confirm the porous and crystalline structure of ZnO‐MnO@C and the in situ growth of Mn deficient ZnMn2O4@C. The Zn/Mn‐d‐ZMO@C cell displays a promising capacity of 194 mAh g−1 at a current density of 100 mA g−1 with 84% of capacity retained after 2000 cycles (at 3000 mA g−1 rate). The improved performance of this cathode originates from in situ orientation, porosity, and carbon coating. Additionally, first‐principles calculations confirm the high electronic conductivity of Mn‐d‐ZMO@C cathode. Importantly, a good capacity retention (86%) is obtained with a year‐old cell (after 150 cycles) at 100 mA g−1 current density. This study, therefore, indicates that the in situ grown Mn‐d‐ZMO@C nanoarchitecture cathode is a promising material to prepare a durable AZIB.Manganese (Mn)‐based cathode materials have garnered huge research interest for rechargeable aqueous zinc‐ion batteries (AZIBs) due to the abundance and low cost of manganese and the plentiful advantages of manganese oxides including their different structures, wide range of phases, and various stoichiometries. A novel in situ generated Mn‐deficient ZnMn 2O4@C cathode has been developed by immersing ZnO‐MnO@C electrode in 2 m ZnSO4 + 0.2 m MnSO4. The in situ formation of a reversible ZnMn2‐xO4@C with Zn4(OH)6(SO4).xH2O surface layer exhibits extraordinary cycle performance.A novel battery consisting of zinc anode and Mn deficient ZnMn Among ion‐based batteries, lithium‐ion batteries (LIBs) are widely studied, which are made of highly explosive lithium metal containing with toxic flammable organic electrolyte. Toxic metal like lead is also used for fabrication of lead‐acid batteries. However, considering not only the risk factors for the environment and human health but also cost and lifetime, the future large‐scale applications of LIBs, lead‐acid batteries are hindered. Even the breakthrough research on battery field has been done, still their performances are lower than the satisfactory level for meeting up the ever‐growing energy demand globally. Therefore, finding highly efficient battery technology which should be made of readily available materials and environmentally friendly are still anticipated.Energy storage is a necessary part of the modern era of ubiquitous high technology. Ecofriendly and high voltage battery technology is required for the safe, long‐term, and convenient operation of advanced electronic devices, such as hand‐held devices, robots, electric vehicles, and medical equipment.4, 5 Also, they are comparatively safe and less toxic than alkali metals. Among the above‐mentioned batteries, ZIBs with aqueous electrolyte have been a target of particular investigation because of their considerable energy density (5851 mAh mL−1), substantial redox potential (‐0.76 V vs standard hydrogen electrode) and ecofriendly zinc anode. In addition, it is easy, convenient and does not require any extra‐expensive atmospheric control to manufacture which lead to the low cost.Multivalent ion batteries such as, aluminum, magnesium, nickel, calcium, and zinc‐ion batteries (ZIBs) have garnered significant attention due to the availability and the relatively low price of those metals used.17 Als−1). Also, vanadium‐based compounds have been reported as a potentially high capacity cathode material for rechargeable AZIBs. However, those cathode materials exhibit low voltage compared to that of manganese (Mn)‐based cathode materials. The rechargeable AZIB battery system was developed by Kang et al. where MnO2 was used as a cathode and exhibit a discharge capacity of 210 mAh g−1. In the meantime, Pan et al. obtained a better result by introducing an electrolyte containing ZnSO4 + MnSO4 in which MnSO4 compensates for the loss of Mn from MnO2 cathode. Long‐term stability (capacity retention of 92% over 5000 cycles) and capacity of 285 mAh g−1 in respect to the total MnO2 are achieved by using a mild aqueous electrolyte of 1 m ZnSO4 + 0.1 m MnSO4. They suggested that conversion reaction between MnOOH and MnO2 played a pivotal part in prolonging the lifetime of the AZIBs. In addition, capacity attenuation and battery failure due to the secondary products 2, Mn2O3) formation in the cathode side during cycling also remain great issues. Furthermore, Alfaruqi et al. also revealed the dissolution of manganese associated with the phase transformation of MnO2 in rechargeable AZIB.A suitable cathode material can add significant contribution on overall battery performances. Although, Prussian blue‐based cathode material can maintain a considerable cycle life of the battery, but exhibit very low capacity (nearly ≈ 50 mAh g−1).23 Als2O4 in the cathode due to Zn2+ ion insertion at the MnO2 cathode was observed for the first time in 2012. Afterwards, Chen et al. have also introduced ZnMn2O4 cathode with cation vacancy for an AZIB which exhibits high capacity retention (94%) over 500 cycles (at 500 mA g−1 current rate). Simultaneous Zn2+ and proton insertion into MnO2 in a mild acidic aqueous electrolyte was reported. Wei et al. developed a novel in situ formed Zn/MnO2 battery where zinc sulphate hydroxide (Zn4SO4(OH)6.5H2O) was used as a starting material that initiate an in situ Zn/MnO2 cell during the first charging. The development and investigation of Mn‐containing cathode materials for AZIBs is still ongoing. It is then worth mentioning that although the Mn‐based cathode are promising over many other transitional metal based cathode, they also possess several drawbacks; and thus they need to be improved.The formation of ZnMn in 2012.33 Aft2O4@C with Mn deficiency (denoted as Mn‐d‐ZMO@C) from the ZnO‐MnO@C nanocomposite synthesized by a simple solvent dry process for high‐performance rechargeable AZIB. The surface morphology of highly porous and crystalline ZnO‐MnO@C was investigated using SEM and HRTEM. The intrinsic structure and in situ changes of ZnO‐MnO@C were confirmed by X‐ray techniques. Time‐dependent in situ potentiostatic electrochemical impedance spectroscopy (PEIS) revealed the phase alteration of the cathode materials throughout the aging process. Both in situ and ex situ X‐ray diffraction analyses were performed to confirm the formation and reproducibility of Mn‐d‐ZMO@C after assembling the coin cell with ZnO‐MnO@C electrode. The Zn/Mn‐d‐ZMO@C cell exhibited excellent discharge capacity of 219 mAh g−1 at 100 mA g−1 and cycle stability with discharge capacity of 98 mAh g−1 after 2000 cycles with nearly 100% Coulombic efficiency at a high current density of 3000 mA g−1. A charged cell can be used after one year and it showed a high capacity retention of 86% over 150 cycles at 100 mA g−1 current density. First principle study was used to understand the improved conductivity of Mn deficient electrode. In addition, electrochemical properties of the ZnMn2O4 and ZnMn2O4‐MnO samples were also evaluated for comparison purpose.In this study, we have revealed the in situ growth of ZnMn2Figure 6 octahedral and hexagonal ZnO crystal built from ZnO4 tetrahedral units. Porous spherical‐shaped particles of the prepared sample were observed in the recorded FE‐SEM images  and that of 0.22 nm for MnO (200)  Figure . In orde2 adsorption–desorption study was performed to investigate the porosity features of the ZnO‐MnO@C powder in terms of their surface area, pore size, and pore volume. The resulting isotherm  reveals a type‐IV profile with a type‐H1 hysteresis loop, indicating mesoporous particles. The Brunaur‐Emmett‐Teller (BET) surface area was calculated to be 76.62 m2 g−1, indicating that such high surface area can enhance the electrochemical performance of ZnO‐MnO@C by offering a large interface between the electrolyte and the electrode. Further, the pore size distribution plot of the ZnO‐MnO@C powder  reveals a pore diameter of 9.55 nm and a pore volume of 0.183 cm3 g−1. This surface characterization result is therefore consistent with the porous properties observed using SEM analysis.The N2O4 (JCPDS # 010‐071‐2499) . And the powder treated at 500 °C produces ZnMn2O4 (JCPDS # 01‐077‐0470) with MnO (JCPDS # 01‐075‐0257 indicated with * sign) impurity . This indicates that MnO started growing with increasing temperature. As expected, the XRD peak corresponding to the graphitic carbon is observed in all samples. Furthermore, as shown in the SEM images , the shape of ZnMn2O4 and ZnMn2O4‐MnO particles, respectively, are similar to that of the porous spherical ZnO‐MnO@C particle.To study the evolution of the ZnO‐MnO phase at temperatures lower than the synthesis temperature, we studied the samples treated under Ar atmosphere at 400 °C and 500 °C. From the XRD pattern of the low‐temperature calcined powder, it is seen to be well‐indexed with ZnMn2O4, ZnMn2O4‐MnO, and ZnO‐MnO samples prepared at 400 °C, 500°C and 600 °C, respectively, in Ar. For the initial two samples containing ZnMn2O4, corresponding Raman bands were observed at 323, 373, and 669 cm−1 . The peak around 1146 cm−1 observed for the ZnO‐MnO sample corresponds to the ZnO. Because of photochemical irradiation, the MnO can be converted to Mn3O4, thereby reflecting the characteristic peak position around 634 cm−1. A couple of strong peaks at 1333 and 1594 cm−1 were observed for all three samples . Here, 1333 cm−1 corresponds to the disordered carbon (D‐bond) arising from A1g vibration mode. The Raman peak for graphitic carbon (G‐Bond) located near the 1594 cm−1 position is attributed to the E2g vibration mode of ordered carbon. The ratio of ID/IG intensity is close to 1, demonstrating a sufficient amount of graphitization which is beneficial for the smooth conduction of electrons. Further characterizations were focused mainly on ZnO‐MnO@C since it is the target sample of the present study.Raman analysis was performed for checking the structural information of carbon wrapped ZnMnrmation).48 The1/2 and Mn 2p3/2 respectively  which indicates the presence of manganese as Mn2+ in MnO. Further, the peaks located at 1021 and 1044 eV are characteristic of Zn 2p3/2 and Zn 2p1/2, respectively .In order to check elements and the oxidation state of the manganese in ZnO‐MnO@C, X‐ray photoelectron spectroscopy was performed. The survey scan for ZnO‐MnO@C confirms the position of notable elements  and is provided in the Figure S5a (Supporting Information). The binding energies of 641.0 and 653.37 eV can be ascribed to Mn 2pm ZnSO4 and 0.2 m MnSO4 was used as an electrolyte. The XRD of the electrode was checked after recovering the ZnO‐MnO@C electrode from 24 h aged coin cell. Figure 2O4 (JCPDS # 01‐077‐0470) and Zn4(OH)6(SO4).0.5 H2O (JCPDS # 00‐044‐0674) (ZBS). The peak near 12.5°, 25.0°, and 50.80° corresponds to the plan (001), (002), and (0‐14) of (Zn4(OH)6(SO4). H2O) and XRD peak at 29.03°, 33.11, 35.52°, 59.06°, 65.0° for the plane (112), (103), (211), (321), and (400) of ZnMn2O4. The ZBS can be formed depending on the pH of the electrolyte. The pH of the electrolyte of the aged coin cell was around 4.5 which is favorable for ZBS stability. Interestingly, highly reversible Mn‐d‐ZMO@C was formed simultaneously, which acts as a cathode material and initiates an in situ Zn/Mn‐d‐ZMO@C battery. Since the cathode surface was covered by ZBS precipitate, to obtain clear XRD pattern of the newly formed ZMO phase, we have then performed ex situ XRD analysis of the aged ZnO‐MnO@C electrode after washing with pH ≈ 3 water. As depicted in Figure S6 (Supporting Information), the XRD pattern is well indexed with the ZnMn2O4 (JCPDS # 01‐077‐0470). It is highly like that ZnO reacted with MnO produces Mn‐d‐ZMO@C in the presence of the electrolyte. Mn2+ may also dissolve from ZnO‐MnO@C into the electrolyte while the formation of ZnMn2O4 with Mn vacancy occurred. A similar kind of phenomenon was investigated by the Wang et al., where they performed electrodeposition to oxidize Mn2+ to Mn4+. Furthermore, a Mn2+/MnO2 battery was developed by the Chen et al. In their study, soluble Mn2+ and solid MnO2 was obtained while 4 m MnSO4 was used as an electrolyte. recently reported the participation of Mn2+ in a Zn/MnO2 aqueous batteries, where they calculated the Gibbs free energy of ZnMn2O4 in the Zn‐Mn‐H2O system was ‐293.439 kcal mol−1 indicating that the formation of ZnMn2O4 is more feasible in our Zn/ZnO‐MnO@C aqueous system. Moreover, ZBS reacts with the dissolved Mn2+ in the electrolyte produces more Mn‐d‐ZMO@C.The in situ cell of Zn/Mn‐d‐ZMO@C was obtained while aging the assembled coin cell for 24 h where zinc was used as an anode, the ZnO‐MnO@C electrode in the counterpart and an aqueous solution of 2 ctrolyte.50 The2SO4.H2O, we have treated the sample using mildly acidic water (pH 3) to remove the precipitate prior to the measurements. It was found that a mildly acidic water could effectively remove the precipitate. ICP and XPS analyses of the treated‐electrode suggest that the presence of Mn vacancy in the in situ growth Mn‐d‐ZMO@C, which will be beneficial for electrode applications. Precisely, the binding energy of around 642 and 653.5 eV observed for Mn 2p3/2 and Mn 2p1/2, which indicating peak shifting to higher energy region resulted after aging process . It also suggests the oxidation of manganese occurred during aging process. In addition, the atomic ratio of Zn to Mn was measured to be ≈1.72, depicting the chemical formula of the Mn deficient ZnMn2O4 is approximately ZnMn1.72O4; and thus the theoretical capacity of the sample can be calculated to be 239 mAh g−1. Moreover, the presence of Zn(OH)2SO4 H2O is easily distinguishable to the naked eye as a white layer covering on the electrode surface. Hence, during cleaning of the recovered electrode using mild acid, the white precipitative layer peeling off from the electrode is visible. Furthermore, manual scraping of the recovered electrode surface and the collection of the internal active material ensures that the contribution of Zn(OH)2SO4 H2O becomes insignificant during ex situ measurements. Finally, we compared the ICP results of the prepared powder and the aged electrode sample to clearly understand the manganese deficiency in the latter.We further attempted to measure the chemical composition of the cathode. To exclude additional contribution of Zn content from Zn(OH)cipitate.50 ICP2O4. As illustrated in Figure S8 (Supporting Information) , the lattice planes 111, 103, and 211 are for zinc inserted spinel ZnMn2O4. Elemental mapping  confirmed the elemental distribution of Zn, Mn, S, and O and the growth of Zn4(OH)6(SO4).xH2O layer. This analysis revealed the formation of Mn‐d‐ZMO@C after aging the electrode. As illustrated in Figure 2O4 is as follows:The morphological evolution of the cathode was also studied. It can be seen from the SEM image that the shape of the electrode remains unchanged after aging the assembled coin cell Figure . Further ZnMn2O4.54 Ele4 + MnSO4 electrolyte for 12 and 24 h without Zn anode. An ex situ XRD characterization for the electrode was performed after the immersion . It can be observed that the ZnO‐MnO@C electrode was converted after immersing in the electrolyte. Compared to the in situ aged electrode which shows prominent phase transformation, for the immersed electrode, the new phases of ZnSO4.3Zn (OH)2.4H2O and spinel ZMO were continuously formed; however, some portion of the ZnO.MnO phase remained unreacted. This can be associated with the potential difference arises in the assembled coin cell. Nevertheless, it can be inferred that ex situ XRD analysis after immersing the electrode without Zn anode indicates phase transformation of the cathode was thermodynamically favored.To obtain more insight into phase transformation of the electrode, the as‐prepared electrode was immersed in the ZnSO2O4, ZnMn2O4‐MnO, and ZnO‐MnO and in situ formed Mn‐d‐ZMO@C with respect to the standard materials MnO, Mn2O3, and MnO2. Figure S10a (Supporting Information) shows that the XANES curve for ZnO‐MnO and reference material MnO are close enough to overlap. Similarly, the curve for ZnMn2O4 and ZnMn2O4‐MnO are almost superimposed on that of the Mn2O3. Therefore, the oxidation state of Mn in ZnMn2O4, ZnMn2O4‐MnO is close to 3+ with respect to the reference materials Mn2O3. As expected, Mn2+ was observed for the ZnO‐MnO@C sample. It can be seen that the Mn3+ in ZnMn2O4 reduced with the increasing temperature. At 500 °C the oxidation state of Mn in ZnMn2O4 started to reduce, and it reached Mn2+ at 650 °C which is in agreement with the oxidation state of Mn in the molecular formula of ZnO‐MnO confirmed from the XRD pattern. Interestingly, the Mn in ZnO‐MnO oxidized back to Mn3+ after aging the ZnO‐MnO electrode in a zinc cell containing ZnSO4 and MnSO4 aqueous solution as an electrolyte. In agreement with the XANES analysis, the formation of ZnMn2O4 was confirmed from the XRD pattern of the recovered electrode. A clearer picture can be seen when comparing Mn valance versus energy in Figure S10b (Supporting Information). We also performed voltage change measurement of the Zn/ZnO‐MnO cell during aging process. Voltage change analysis suggests that the cell underwent an increased potential from 1.27 to 1.34 V and is stabilized after 24 h which further confirms the oxidation of Mn2+ to Mn3+ happened during the aging process . This observation also supports the in situ formation of ZnMn2O4.XANES analysis was performed to investigate the oxidation state of the Mn in carbon enfolded ZnMnFigure Rct, Rs, CPE, and Zw represent charge transfer resistance, the resistance of electrolyte, the constant phase element, and Warburg impedance, respectively. It can be observed from the Figure Rsurf). This is due to the formation of Zn4(OH)6SO4.xH2O (ZBS). Interestingly, a similar semicircle observed from 20th to 150th h and all impedance curves are overlapped on each other until 150 h. Indicating a stable phase formed after 24 h 6SO4.xH2O. In situ PEIS study is consistent with the ex situ XRD analysis.The aging process further studied using (PEIS). We have taken PEIS immediately after assembling the coin cell and recorded at different times. As depicted in h Figure . The fith Figure ,f. This 4 electrolyte with and without MnSO4 additives. As displayed in the Figure 4 in ZnSO4 electrolyte shows that impedance stabilizes after 24 h. Whereas more resistance has grown continuously in the absence of MnSO4 additives   that the initial charge capacity for the ZnO‐MnO@C electrode in an aqueous solution of ZnSO4 + MnSO4 is 151 mAh g−1 at 300 mA g−1 current density, whereas it is only 57 mAh g−1 in ZnSO4 electrolyte. Similarly, a high discharge capacity of 141 mAh g−1 was obtained in the presence of MnSO4 additives. In both cases, the initial activation process was observed while checking cycle ability .For comparison, we have investigated PEIS of the ZnO‐MnO@C nanocomposite in ZnSOs Figure . This re) Figure . From thy Figure . After 7−1. As shown in Figure S12b (Supporting Information), two distinct cathodic peaks were observed at 1.22 and 1.39 V. An increase of peak intensity in the subsequent cycle showed the activation process which agrees with the characteristics of charge/discharge. A similar phenomenon was observed in a previously reported article. Electrochemical charge–discharge was conducted within the potential window of 0.8 to 1.9 V. As shown in Figure m ZnSO4 + 0.2 m MnSO4 electrolyte registered 212 mAh g−1 capacity and a discharge capacity of 194 mAh g−1 at a current density of 100 mA g−1. At the fifth cycle, the cell registered a high capacity of 227 mAh g−1. When checking rate capability at different current rates, sacrificial ZnO‐MnO@C cathode materials display very promising performance. Figure −1 observed at 100, 200, 500, 1000, 2000, and 3000 mA g−1 current rate, respectively. After cycling at different current density, the electrode was able to register discharge capacity of 201 and 233 mAh g−1 when it was operated at 300 and 100 mA g−1 rate, respectively. The electrochemical charge–discharge curve at the various current rates is shown in Figure −1 and the remainder of the cycles at 3000 mA g−1 current density. For the first cycle, the Zn/Mn‐d‐ZMO@C cell registered a discharge capacity of 78 and 81 mAh g−1 charge capacity at a very high current density of 1000 mA g−1. A high capacity of 132 mAh g−1 was recorded after 10 cycles. The preliminary rise in capacity is due to the stimulation process, which agrees with previously reported articles. When the cell is cycled at 3000 mA g−1 current rate, 92.5 mAh g−1 discharge capacity was observed at the 11th cycle. It reached 98 mAh g−1 capacity at the 14th cycle and sustained a high number of cycles with good capacity retention. After 2000 cycles, it recovered 82.7 mAh g−1 discharge capacity, i.e., it retained 84% capacity (with respect to the highest capacity of 98 mAh g−1). The charge–discharge profile at different cycles as shown in Figure Cyclic voltammetry analysis was performed and is discussed for analysis reaction mechanism of Zn/ZnO‐MnO@C cell. From the CV profile, the oxidation peak at 1.58 V appeared at the very first cycle while the cell operated at a scan rate of 0.2 mV s article.54 Ele2O4@C, ZnMn2O4/MnO@C and ZnO‐MnO@C electrodes in MnSO4 additive containing ZnSO4 electrolyte, respectively, were measured and compared . Electrochemical cyclic voltammetry was performed within the potential range 0.8 to 1.9 V at a scan rate of 0.2 mV s−1. Figure S13a (Supporting Information) represents the anodic peak near 1.59 V and the cathodic peaks at 1.39 and 1.15 V were observed for ZnMn2O4@C. Similar results were observed for ZMO with MnO impurity . Both observations are consistent with the CV profile of in situ cell developed from the ZnO‐MnO@C nanostructure materials. Electrochemical charge–discharge properties of the ZnMn2O4@C and ZnMn2O4/MnO@C were also evaluated at a 500 mA g−1 rate. It can be observed that in situ generated cathode from ZnO‐MnO exhibited less polarization than those of ZnMn2O4 and ZnMn2O4‐MnO electrodes . When comparing their cycle ability performances, pure ZnMn2O4 registered only 20 mAh g−1 discharge capacity at 500 mA g−1 current density, whereas they were 56 and 118 mAh g−1 for ZnMn2O4‐MnO and ZnO‐MnO, respectively. Moreover, the Mn‐d‐ZMO@C cathode delivered stable long cycle life . It can be concluded that the incorporation of MnO enhanced the electrochemical performance of the ZMO electrode significantly. In addition, the electrochemical cycle performance of the Zn/ZnO‐MnO@C electrode was investigated after a long time (1 year). The assembled cell was kept after charging up to 1.9 V and discharged after 1 year. At the 150th cycle, this 1 year old cell can still deliver a discharge capacity of 60 mAh g−1 with no substantial capacity fading (nearly 86% capacity retention) at 100 mA g−1 current rate .To understand the superiority of the in situ generated Mn‐d‐ZMO@C electrode prepared in the present study, the electrochemical performances of ZnMn−1) and displayed in Figure The cyclic voltammetry (CV) analysis of the Zn/ZnO‐MnO@C cell was performed to study the further reaction mechanism and chemical kinetics. Investigation of chemical kinetic of the cell, cyclovoltammetry was employed at different scan rates, namely i is the current response at a static potential (V), v is the scan rate, k1 and k2 are constants use to determine the capacitive and diffusion‐limited behavior, respectively. The calculated capacitive behavior of the electrode was 48% at 0.2 mV s−1 rate and is shown by the shaded area in Figure −1 sweep rates, respectively.Here Figure 2O4 (indicated by the * sign) can be clearly seen, along with Zn4(OH)6SO4.5H2O (specified with the # sign) and an unavoidable peak for stainless steel (identified with the γ sign). More water content was identified for the in situ situation but is reduced while checking the XRD of the dried electrode. A total of 137 XRD measurements were taken during charge–discharge. When the electrode was subjected to charge, the peak for Zn4(OH)6SO4.5H2O (ZBS) started to weaken due to the dissolution of ZBS because of the increasing acidity of the electrolyte. While the electrode charged completely, the pH value was reduced to around 3 and ZBS precipitate can be easily dissolved. Previous report also suggested after charging, the XRD peak for ZBS was abolished completely. The plane 103 for ZnMn2O4 also almost disappeared at the charged state. Remarkable peak shifting of plane 211 is observed from the Figure θ angles from 11.5° to 13.5°. The plane 001 corresponding to ZBS disappeared after charging and reformed after discharging. Similar phenomena were observed for the 002 plane and the rest of the planes for ZBS. 112, 103 and 211 planes of ZMO were considered for further analysis. From Figure 1‐yMn2O4 (identified by an α sign), whereas the diffraction peak for 211 plane shifted substantially. New peaks near 22° and 35.76° are identified as ZnMn3O7.3H2O (PDF # 15‐0807). While some studies showed that during charging ZBS dissolution occurred due to the change of the electrolyte pH, we also observed the formation of ZnMn3O7.3H2O, which is in agreement to the reports by Zhao et al. and Huang et al. It was found that Mn2+ additive enhanced the deformation of ZBS by producing ZnMn3O7 which could also act as a Zn host and influenced the initial capacity significantly  is observed. This points to the feasibility of MnO2 getting deposited on to the electrode surface via the oxidation of Mn2+ additive ions in the electrolyte during the initial charge cycling. The overall XRD results, thus conclude that the deintercalation of Zn leading to Mn oxidation in the bulk occurs upon initial charging. Besides intercalation, a major role is suspected to be played mainly by the Mn2+ additive ions in the reaction mechanism. In specific, a portion of the Mn2+ ions, mostly from the additive, undergoes a conversion reaction with ZBS on the electrode surface to produce ZnMn3O7.3H2O. On further charging, the remnant Mn2+ ions electrodeposit as MnOx via oxidation. Moreover, Mn2+ also affects the Coulombic efficiency of the cell while cycling at a low current density due to the formation of ZnMn3O7 and MnO1.88 at the initial stage 6SO4.0.5H2O and ZnMn2O4, respectively, which became more prominent after cycling . It is worth noting that the overlapped diffraction peaks make it more difficult to analyze the phases; however, we can clearly observe ZnMn2O4 phase of the acid‐washed electrode . In addition, we have also performed ex situ XRD and SEM analyses of the in situ formed electrode after 20 cycles and the results are depicted in Figure S17 (Supporting Information). Strong ZnMn2O4 (JCPDS # 010‐077‐0470) along with ZBS (JCPDS # 039–0690) characteristics can be clearly observed after 20 cycles from XRD patterns. Notably, SEM image indicates the spherical shape of the particle preserved after cycling. These observations suggest the high reversibility and stability of the cathode during consecutive cycling.When the cell is discharged up to 0.8 V, the 112, 103, and 211 planes were recovered completely; this proved the structural reversibility of the in situ electrode. For further clarification, the XRD pattern of the electrode was also recorded after 10 cycles and compared with the pristine electrode . This peak shifts due to the existence of less water content in the ex situ samples.50 The−1) is worth noting. The present study clarifies that the reaction mechanism in the Mn‐d‐ZMO@C electrode can be explained by an intercalation and two conversion processes, respectively. The deintercalation, in general, is believed to be dominant at low current densities and, in that case, it is feasible that the deintercalated Zn2+ions are not fully inserted back into the electrode bulk during the discharge reaction either due to the presence of irreversible sites or because of some hindrance to the motion of the intercalating Zn2+ ions. The low current density can motivate the formation of a relatively thick ZBS layer on the electrode surface during discharge cycling. Thus, it can be possible that the difficulty in penetration of Zn2+ ions through this thick layer formed at the electrode/electrolyte interface may not allow the full intercalation of the deintercalated Zn2+ ions. Also, this irreversibility can be related to some unusual pH variation that can hinder the reversibility of either of the conversion reactions when cycled at low current densities.The relatively lower Coulombic efficiency during the initial few cycles at low current density . While charging, the energy of Mn K‐edge moves towards higher energy. At the fully charged state (1.90 V) the Mn K‐edge of the electrode (purple colored curve) is nearly superimposed on that of the MnO2 reference (black colored curve). This is highly likely due to the presence of ZnMn3O7 and Mn vacancy in the spinel phase electrode. During discharge, the reverse phenomena occurred . At the depth of discharge of 0.8 V, a very similar Mn K‐edge spectrum was obtained as that observed for the aged electrode. That means Mn3+ is regained after fully discharged. This result supported both the in situ and ex situ XRD analyses. It is also worth mentioning that in the discharge reaction Mn2+ exists in the electrolyte solution. As displayed in Figure S18 (Supporting Information), the Mn K‐edge for the first and second charges were comparable, showing the reversibility of the electrode reaction. Moreover, we have performed ex situ XPS analysis during cycling. As shown in Figure S20 (Supporting Information), Mn2p3/2 and Mn2p1/2 peaks move to higher energy after the aging process. Peak near the 641.5 eV corresponds to Mn3+. It can be seen that the binding energy increases further while the electrode charged to 1.9 V and majority of the Mn2p3/2 peak near the 642.2 eV indicating the oxidation of Mn and similar peak shifting observed for Mn2p1/2 peak. The ex situ XPS results were closely resemblance to the XANES analyses. Overall, the repetition of nearly Mn4+/Mn3+ redox couple was observed during charge/discharge.Ex situ synchrotron XAS analysis was performed at a different state of charge–discharge to study the change of manganese oxidation state of the ZnO‐MnO@C electrode. The Mn K‐edge spectrum of the aged electrode close to that of Mn1/2 peak.62 The4 + MnSO4 electrolyte. Further, it converted into spinel ZnMn2O4 phase and the of Zn4(OH)6.SO4.5H2O (ZBS) precipitate was formed on the surface of the electrode materials. Further, ZBS reacts with Mn2+ to produce more ZMO, which is consistent with the previous reports. When being charged, the oxidation of the Mn occurred. Specifically, Zn1‐yMn2‐xO4 and ZnMn3O7.3H2O were formed after charging. It is to be noted that at charged state, the pH of the electrolyte was reduced to 3.0 from 4.5, while in the fully discharged state, ZnMn2O4 reformed, whereas electrolyte pH was found to be 4.5. As also illustrated in the Figure In the light of above discussion, at first ZnO‐MnO@C undergoes a “self‐oxidation” in an aqueous solution of ZnSO reports.40 WheCharge:Discharge:2 system is hugely debated at present. Briefly, different type of reactions was proposed to describe the mechanism in manganese oxide cathodes for AZIBs. Initially, pure Zn‐intercalation was floated as the plausible reaction mechanism. Thereafter, the possibility of a dual Zn2+/H+ intercalation regulation was put forward. However, controversies still exist in assigning the specific ion related to the pathway of the main driving reaction. Meanwhile, conversion reactions based on a proton‐transfer process and a phase transfer process, respectively, were also deemed possible to govern the electrochemical regulations in manganese oxides of AZIBs.Also, the possibility of mixed intercalation‐cum‐chemical conversion reactions were proposed. Very recently, a manganese deposition/dissolution process that challenges the views of conventional intercalation was presented. In this regard, we propose that the reaction mechanism in the present Mn‐d‐ZMO@C cathode can be described by a combined intercalation‐conversion‐deposition reaction in an Mn2+‐containing ZnSO4 aqueous electrolyte medium.The reaction mechanism of manganese oxide cathodes in the Zn//MnO present.60 BriR) decreases from around 178 to 143 Ω at 1.35 V. This is maybe due to the gradual deformation of more ZBS which was well observed in ex situ XRD analysis. Interestingly, resistant further reduced in the subsequent charge state and only 18 Ω observed when it was charged to 1.90 V. This result suggest that the deformation of Mn‐d‐ZMO@C occurs during charging. When it is discharged, the R value increases dramatically, and this is due to the reversible formation of the ZnMn2O4 and ZBS layer . Similar phenomena were observed in the subsequent cycles . Based on electrochemical results, the in situ grown Mn‐d‐ZMO@C shows superior electrochemical performance. In addition, ICP analyses suggest the dissolution of Mn, which leads to the deficient amount of Mn in the in situ grown sample. With the Mn vacancy defect, it is expected that the material shows an improved electrochemical performance.Further in situ PEIS analysis was conducted for investigating the change of resistance and phase formation during cycling. From Figure S21a (Supporting Information) it can be elucidated that during first charge, over all resistant . Figure 2O4 and ZnMn2O4 with Mn vacancy defect along with their corresponding total density of states. It can be clearly seen that the defect‐free ZnMn2O4 exhibits semiconductor feature, while remarkably, the ZnMn2O4 with the Mn vacancy defect shows differently. Accordingly, it is worth noting that the Mn‐d‐ZMO@C sample showed a metallic behavior, which has a higher conductivity compared to the defect‐free ZnMn2O4. Thus, fast electron transport and an enhanced electrochemical property are expected. Therefore, the present in situ growth formation technique will open the new insight for fabricating advanced electrode materials with high performance.Therefore, to validate this, first‐principles calculations based on density functional theory66 wer2+ to Mn3+) reactions initially and leads to the in situ formed of Mn‐d‐ZMO@C cathode. In situ PEIS and XRD, ex situ XRD, and XANES study revealed the development of the reproducible Zn/Mn‐d‐ZMO@C system. Moreover, these characterizations revealed that the reaction mechanism of the present Mn‐d‐ZMO@C cathode can be described by a combined intercalation–conversion–deposition reaction. The in situ generated Zn/Mn‐d‐ZMO@C battery exhibits excellent electrochemical performance. A long cycle with high capacity retention was achievable with the in situ formed cell (84% capacity maintained after 2000 cycles). A comparative study of the in situ grown Zn/Mn‐d‐ZMO@C cell with Zn/ZnMn2O4@C and Zn/ZnMn2O4‐MnO@C, respectively, identified that the Zn//Mn‐d‐ZMO@C cell exhibited improved electrochemical performance than that of the conventional cells, respectively. However, there are some issues to be addressed further. For example, the relatively lower Coulombic efficiency during the initial few cycles at low current density (≈100 mA g−1) is worth noting in Figure 2+ ions are not fully inserted back into the electrode bulk during the discharge reaction either due to the presence of irreversible sites or because of some hindrance to the motion of the intercalating Zn2+ ions. The low current density can motivate the formation of a relatively thick ZBS layer (compared to that at higher current densities) on the electrode surface during discharge cycling. This layer can act as a barrier to Zn‐intercalation thus making it difficult for the intercalation of all the deintercalated Zn2+ ions during charge cycling. Also, this irreversibility can be related to some unusual pH variation that can hinder the reversibility of either of the conversion reactions when cycled at low current densities. More studies are required to establish the exact reason for the irreversibility at low current densities. The charged Zn//Mn‐d‐ZMO@C cell could be used even 12 months later and able to use for long cycle with fewer capacity decay (86% capacity retention). It exhibited a promising energy density of 296 Wh kg−1 at 135 W kg−1 power density, more superior than these of earlier reported works on Zn/MnO2, Zn/Na0.95MnO2, Zn/ZnMn2O4‐Mn2O3, Zn/Mn2O3, ZnMn2O4, and Zn/MnO2. Moreover, the in situ formed Zn/Mn‐d‐ZMO@C cell can deliver 178 Wh kg−1 specific energy at a high specific power of 4050 W kg−1. A Clearer picture has been depicted in the Ragone plot 6(SO4).xH2O phase was identified using in situ and ex situ XRD. The in situ formed Zn/Mn‐d‐ZMO@C cell delivered an initial discharge capacity of 194 mAh g−1 at 100 mA g−1 current rate. Moreover, it can retain 84% capacity over 2000 cycles. In situ and ex situ XRD proved the formation and reproducibility of the in situ formed Mn‐d‐ZMO@C electrode. In addition, in situ/ex situ X‐ray and electrochemical analyses revealed that a possible combined intercalation–conversion–deposition reaction governs the regulation mechanism in the present Mn‐d‐ZMO@C cathode. However, considering the existing debate in the regulation mechanism for manganese oxide cathodes of AZIBs, further studies are required to confirm the proposed reaction. Further, first‐principles calculations suggest higher conductivity of the Mn defected electrode compared to that of pure ZnMn2O4. The in situ formed Zn/Mn‐d‐ZMO@C cell can provide a high specific energy density of 296 Wh kg−1 at 135 W kg−1 specific power density. The charged in situ cell can also be cycled after being stored for a period of one year. This study opens new avenue for further development of cathode material for high performance and practical rechargeable AZIB system.In summary, ZnO‐MnO@C nanocomposite materials have been synthesized herein via the solvent dry process. An in situ Zn/Mn‐d‐ZMO@C cathode was formed in the Zn cell containing an aqueous 2 43, Zn(NO3)2.6H2O, Mn(NO3)2.4H2O from Sigma‐Aldrich and commercial acetylene black were used as a precursor. Initially, 2 m of Zn(NO3)2.6H2O and 4 m of Mn(NO3)2.4H2O were dissolved in 20 mL of concentrated HNO3 followed by the addition of 50 mg of acetylene black powder to the above solution under vigorous stirring. The solution was then dried at 80 °C in air. Finally, the dried powder was annealed at 400 °C, 500 °C, and 650 °C temperature separately for 4 h in an argon (Ar) atmosphere. Materials and electrochemical characterization are discussed in the Note S1 (Supporting Information).ZnO‐MnO@C composite materials were simply synthesis from a solvent dry process. 70% HNOThe authors declare no conflict of interest.Supporting InformationClick here for additional data file.


The oxide scale, formed by pre-oxidation, protects the substrate from the NaCl induced corrosion during the initial stage. However, the corrosion of the pre-oxidized sample is severely accelerated by solid NaCl after an incubation period. The chlorine, generated from the decomposition of solid NaCl, diffuses into the oxide/substrate interface as ions during the incubation period, which was observed by ToF–SIMS. The chlorine at the oxide/substrate interface induces the fast corrosion after the incubation period although the pre-oxidation scale is complete and compact.The effect of pre-oxidation on the corrosion behavior of pure Ti covered with a solid NaCl deposit in the humid O Compressor blades of airplanes and ships suffer severe corrosion in marine environment, which is mainly due to the fact that marine air contains abundant salts  and water vapor. At the temperature of 300–600 °C where compressor blades operate, NaCl appears in solid state and gets deposited on the metal surface. The corrosion of several metals and alloys of compressor blades gets accelerated under a synergistic effect of the solid NaCl deposit and humid oxygen.11 demonstrated that the corrosion behavior of pure Fe, pure Cr, Fe–Cr alloy, Ti alloy and Ni alloy at 600 ˚C under a deposit of solid NaCl deteriorated sharply. Similarly, the authors investigated the influence of KCl on the oxidation of the 304-type (Fe18Cr10Ni) austenitic stainless steel at 600 ˚C in 5% O2 and in 5% O2 + 40% H2O in paper17, and the results showed that small additions of potassium chloride strongly accelerated high temperature corrosion. The fast corrosion is mainly due to a chemical reaction of the solid NaCl/KCl deposit with oxides. This reaction destroys the protective oxide scale on the surface and forms a porous non-protective scale . It is also reported in these investigations that the key factor for the fast corrosion is the chlorine, HCl or Cl2, which cyclically reacts with the substrate at the oxide/substrate interface. The gaseous molecule of chlorine, which is formed during the chemical reaction of solid NaCl and oxide18, diffuses inward quickly through macro defects (like cracks and holes) in the scale of corrosion products. The chlorine then reacts cyclically with the substrate to generate volatile products (for example CrCl3) which break the protective scale 8.Extensive works have been done on the corrosion behavior of several metals and alloys of compressor blades. The authors in paper11, we observe that the scale of corrosion products of Ti60 alloys can be divided into an outer the corrosion layer and an inner corrosion layer after exposure at 600 ˚C in the environment of NaCl + H2O + O2. The chlorine exists in a Ti-Cl bond which is formed by replacing O with Cl within lattices of Ti oxides of the inner corrosion layer. We think the chlorine can diffuse inward as ion. In this case, Cl- comes from the decomposition of solid NaCl and the residual sodium reacts with oxides on the surface to form metallic acid salts (like Na2CrO41 and Na4Ti5O1210). In order to make clear the corrosion mechanism, we need to first figure out the state of the chlorine and study its diffusion in the scale of corrosion products, especially during the destruction process of the scale of protective oxide, like TiO2 and Cr2O3.In our previous work2 condition at 600 ˚C, a compact and even oxide scale forms on the surface10. This oxide scale can simulate the passive film to protect the substance from the corrosive environment. When the pre-oxidized pure Ti is exposed in the NaCl + H2O + O2 environment, this oxide scale can slow down the destruction process of the protective passive film such that the behavior of the solid NaCl can be investigated in detail. Thus, to clarify the micro-mechanism under the specific environment of NaCl + H2O + O2, we need to study the effects of pre-oxidation on the corrosion behavior of pure Ti.When pure Ti is oxidized in the pure O2 flow at 600 ˚C, using scanning electron microscope (SEM) equipped with an energy dispersive spectrometer (EDS), X-ray diffraction (XRD) and time of flight-secondary ion mass spectrometry (ToF–SIMS). After examining the detailed properties of the corrosion products and the diffusion of chlorine in the oxide scale, we discuss possible micro acceleration mechanisms of the chlorine based on experiment results.In this paper, we investigate the effects of pre-oxidation on the corrosion behavior of pure Ti underneath a solid NaCl deposit in humid O2O + O2 (about 0.17 mg/cm2) is slightly larger than that under the condition of O2 (about 0.09 mg/cm2). In fact, under both the above two conditions the corrosion is minor. In the initial stage (incubation period), the mass gain is slight when the pre-oxidized pure Ti samples under NaCl deposit. In contrast, the mass gain increases rapidly after this incubation period. After 20 h exposure in NaCl + H2O + O2 or NaCl + O2, the mass gains are about 1.8 mg/cm2 and 1.4 mg/cm2, respectively, which are about one order of magnitude larger than that in the absence of the solid NaCl deposit. Thus, we reach the conclusion that the solid NaCl accelerates the corrosion of the pre-oxidized pure Ti in O2 or humid O2.The mass gain curves of pure Ti under different conditions are shown in Fig. 2O + O2. After exposed for 20 h, the mass gain is about 6.0 mg/cm2, which is increased by 3.5 times compared with the pre-oxidized samples. Therefore, we can conclude that the pre-oxidation improves the corrosion resistance of pure Ti coated with a solid NaCl deposit in humid O2 at 600 ˚C, especially in the incubation period.For the bare samples, the mass gain increases dramatically during the whole exposure in NaCl + H2 flow for 20 h, the surface and the cross-sectional morphologies are shown in Fig. 2.For the pure Ti samples after pre-oxidation in O2O + O2, the surface and the cross-sectional morphologies are shown in Fig. For the bare samples after 20 h exposure in NaCl + H2O + O2 up to 20 h. The results show that the corrosion products mainly consist of TiO2 and some Na4Ti5O12, with the residual NaCl.In Fig. 2 flow, the surface and the cross-sectional morphologies are shown in Fig. 2 flow. From Fig. For the pre-oxidized pure Ti samples after 100 h exposure in dry O2 flow, the surface and the cross-sectional morphologies are presented in Fig. For the pre-oxidized samples after 100 h exposure in humid O2, the surface and the cross-sectional morphologies are displayed in Fig. For the pre-oxidized samples after 20 h exposure in NaCl + O2O + O2, the surface and the cross-sectional morphologies are depicted in Fig. For the pre-oxidized samples after 20 h exposure in NaCl + H2O + O2, the surface and the cross-sectional morphologies are shown in Fig. For the pre-oxidized samples after 1 h exposure in NaCl + H2O + O2, the surface and the cross-sectional morphologies are depicted in Fig. For the pre-oxidized samples after 5 h exposure in NaCl + H2 was formed on the pre-oxidized samples after 100 h exposure in dry O2 flow  on the surface. As shown in Fig. 2O + O2 (which is after the incubation period), the corrosion products also contain Na4Ti5O12. Since the diffracted intensity of Na4Ti5O12 in Fig. 4Ti5O12 is formed when the pre-oxidized samples are exposed in NaCl + H2O + O2 for a longer time period. As shown in Fig. 2O + O2 (which is during the incubation period), there are not Na4Ti5O12 among the corrosion products. We conclude the corrosion products under the above different conditions in Table The corrosion products of the pure Ti samples are identified by XRD. The XRD patterns in Fig. low Fig. a or in Hlow Fig. b. In Fig2O + O2. The Figure 2O + O2, Cl- ions first enrich on the surface of the pre-oxidized samples  and then decrease rapidly in the corrosion product scale. Afterwards, Cl- is enriched at the oxide/metal interface, which indicates that the chlorine penetrates the oxide and accumulates at the oxide/metal interface. In Fig. 2O + O2 for 5 h, a similar distribution of Cl- in the corrosion product scale was observed. Particularly, the Cl- ion is still enriched when the - around the oxide/metal interface is stronger in Fig. In Fig. - signal in Fig. - is similar with Cl-. First, the OH- ion enriches on the surface and then decrease rapidly in the corrosion product scale. This is the information of water vapor. Afterwards, OH- is enriched at the oxide/metal interface where Cl- is accumulated, which indicates that the H penetrates the oxide along with the chlorine.The OH2 flow at 600 ˚C for 20 h, a compact TiO2 scale with a thickness of about 1 μm was formed on the surface, which is caused by the following reaction between the metal substrate and O2:2O + O2) increased quickly and the mass gain is about 6.0 mg/cm2 after 20 h exposure . The resulting scale consists of plentiful and porous corrosion products , the mass gain of the pre-oxidized samples is smaller than that of the bare samples  and reduce the corrosion to a certain extent, especially in the incubation period.When pre-oxidized samples are exposed in the presence of solid NaCl  planes into free H atoms and OH- groups23 when the pre-oxidized samples is exposed in humid O2. The papers22 reported that along the crystal channels in c-axis direction, the diffusion rate of hydrogen is at least one order of magnitude larger than that in the perpendicular direction of c-axis. Thus, the generated hydrogen atoms are more prone to diffuse into TiO2 through the former one. This dissolution of hydrogen in TiO2 will form hydrogen defects, thus the hydrogen can increase the concentration of crystal defects and increase the outward diffusion of Ti ions23. This fast diffusion in the scale causes a larger corrosion rate and hence the hole forms in the oxide scale  than when exposed in the absence of the solid NaCl (O2 or H2O + O2); (2) after 20 h exposure, the pre-oxidized samples has a larger mass gain under the condition of NaCl + O2 or NaCl + H2O + O2 than under the condition of O2 or H2O + O2. Additionally, we can observe that the corrosion product scale of the pre-oxidized samples is thicker when exposed in the presence of the solid NaCl  at 600 ˚C. Hence, we think the chlorine permeates into the corrosion product scale as ion (Cl-).For the pre-oxidized samples, a thick TiOace Fig. a. Shu etace Fig. a while tete Fig. . This im2O + O2, the chlorine is also observed around the metal/oxide interface by SIMS. Additionally, the intensity of chlorine is still large when the intensity of oxygen decreases to a relatively small value. As aforementioned, the substrate locates at the places where the intensity of oxygen is small. Hence, we can infer that the chlorine diffuses into the substrate and directly binds with the metal of the substrate during the fast corrosion. The gathered compound diffuses outward, resulting in the loss of metal in the substrate. After 20 h exposure, a mass of the metal reacts with the chlorine and diffuses outward, and the inner corrosion layer is formed under the outer corrosion layer , which can protect the substrate from fast corrosion. However, the “protective oxide” is not completely inert to solid NaCl and the corrosion rate of the pre-oxidized sample is also greatly accelerated by solid NaCl after an incubation period. During the incubation period, the chlorine, generated in the decomposition of solid NaCl at 600 ˚C, diffuses into the oxide/substrate interface as ions to enhance the corrosion rate, and then the residual sodium reacts with the oxide on the surface. After the incubation period, the chlorine diffuses into the substrate and directly reacts with the substrate, resulting in the fast corrosion.During the pre-oxidation, a compact and thick TiO2 flow at 600 ˚C for 20 h for the pre-oxidization. The surfaces of preheated samples were covered with a layer of NaCl deposit by repeatedly brushing and drying with saturated NaCl solution24, until about 4 ± 0.2 mg/cm2 solid NaCl was deposited on.The material used in this study was pure Ti and the samples were cut into the size of 10 mm × 15 mm × 2.5 mm. Prior to experiments, all samples were mechanically grinded with 800# SiC paper, ultrasonically degreased in alcohol for about 20 min, and dried in the air. The bare samples were then exposed in the pure O24. The continuous mass gain during the corrosion experiment was obtained with a thermo-gravimetric analysis (TGA).The corrosion tests were carried out in a thermo-balance 2 was bubbled into the distilled water with a glass bubbler  to produce the test atmosphere (humid O2). According to the relationship between the vapor pressure of water and its temperature, we precisely set up the temperature of the distilled water in the glass bubbler to control the amount of the water vapor. In this study, the temperature of the distilled water was about 70 ˚C, producing about 30.8 vol.% water vapor. In order to avoid the water vapor condensing inside the thermo-balance, a counter flow of pure N2 was passed through the thermo-balance, whose flow rate was about 400 mL/min. After the furnace reached 600 ˚C and the gas flows of humid O2 as well as pure N2 stabilized, the samples were quickly lowered into the constant temperature zone of the furnace tube.The pure O2 flow at 600 ˚C (denoted as NaCl + H2O + O2); the second condition was carried out in a humid O2 flow at 600 ˚C (denoted as H2O + O2); the third condition was with a solid NaCl deposit layer in a dry O2 flow at 600 ˚C (denoted as NaCl + O2); and the fourth condition was carried out in a dry O2 flow at 600 ˚C (denoted as O2). We present the explicit environmental parameters in Table In this study, the corrosion experiments were respectively carried out under the following four conditions: the first condition was with a solid NaCl deposit layer in a humid OThe surface morphologies and the cross-sectional morphologies of corrosion products scale were collected by SEM–EDS. The chemical composition of corrosion products was identified by XRD.2O + O2 and O2, the mass gains of the pre-oxidized sample were very small , which allowed parallel mass registration with high sensitivity and high mass resolution. A cesium liquid–metal ion (LMI) gun at 20 keV beam energy was used for spatially resolved ToF–SIMS analysis.The pre-oxidized samples were first exposed in either NaCl + H


Emissions of CO2 from fossil fuel combustion and land use change reduce the ratio of 13C/12C in atmospheric CO2 (δ13CO2). This is because 12C is preferentially assimilated during photosynthesis and δ13C in plant‐derived carbon in terrestrial ecosystems and fossil fuels is lower than atmospheric δ13CO2. Emissions of CO2 from fossil fuel combustion also reduce the ratio of 14C/C in atmospheric CO2 (Δ14CO2) because 14C is absent in million‐year‐old fossil fuels, which have been stored for much longer than the radioactive decay time of 14C. Atmospheric Δ14CO2 rapidly increased in the 1950s to 1960s because of 14C produced during nuclear bomb testing. The resulting trends in δ13C and Δ14C in atmospheric CO2 are influenced not only by these human emissions but also by natural carbon exchanges that mix carbon between the atmosphere and ocean and terrestrial ecosystems. This mixing caused Δ14CO2 to return toward preindustrial levels in the first few decades after the spike from nuclear testing. More recently, as the bomb 14C excess is now mostly well mixed with the decadally overturning carbon reservoirs, fossil fuel emissions have become the main factor driving further decreases in atmospheric Δ14CO2. For δ13CO2, in addition to exchanges between reservoirs, the extent to which 12C is preferentially assimilated during photosynthesis appears to have increased, slowing down the recent δ13CO2 trend slightly. A new compilation of ice core and flask δ13CO2 observations indicates that the decline in δ13CO2 since the preindustrial period is less than some prior estimates, which may have incorporated artifacts owing to offsets from different laboratories' measurements. Atmospheric observations of δ13CO2 have been used to investigate carbon fluxes and the functioning of plants, and they are used for comparison with δ13C in other materials such as tree rings. Atmospheric observations of Δ14CO2 have been used to quantify the rate of air‐sea gas exchange and ocean circulation, and the rate of net primary production and the turnover time of carbon in plant material and soils. Atmospheric observations of Δ14CO2 are also used for comparison with Δ14C in other materials in many fields such as archaeology, forensics, and physiology. Another major application is the assessment of regional emissions of CO2 from fossil fuel combustion using Δ14CO2 observations and models. In the future, δ13CO2 and Δ14CO2 will continue to change. The sign and magnitude of the changes are mainly determined by global fossil fuel emissions. We present here simulations of future δ13CO2 and Δ14CO2 for six scenarios based on the shared socioeconomic pathways (SSPs) from the 6th Coupled Model Intercomparison Project (CMIP6). Applications using atmospheric δ13CO2 and Δ14CO2 observations in carbon cycle science and many other fields will be affected by these future changes. We recommend an increased effort toward making coordinated measurements of δ13C and Δ14C across the Earth System and for further development of isotopic modeling and model‐data analysis tools.In this “Grand Challenges” paper, we review how the carbon isotopic composition of atmospheric CO 14C and 13C, in atmospheric CO2 are changing in response to fossil fuel emissions and other human activitiesCarbon isotopes, 13C, BECCSFuture simulations using different SSPs show continued changes in isotopic ratios that depend on fossil fuel emissions and, for 14C and 13C in studies of the carbon cycle or other fields will be affected by future changesApplications using atmospheric  Radiocarbon is formed naturally in the upper atmosphere from cosmogenic radiation, which produces neutrons that react with atmospheric nitrogen. Because the isotopic composition of carbon is affected by physical, chemical, and biological processes, these ratios are not constant, and they vary across different carbon pools and over time and space. Precise measurements of small differences in these ratios, together with theoretical or empirical models of isotopic fractionation and mixing, enable the investigation of various aspects of the carbon cycle. Observing and analyzing the changes in carbon isotopic composition of atmospheric CO2 can help to understand the natural carbon cycle's response to human activities.Carbon isotopes are present in the atmosphere, ocean, and terrestrial biosphere in ratios of approximately 99% 13C refers to the deviation of the ratio 13C/12C in a sample from a standard ratio 13C/12C, referred to as Vienna Pee Dee Belemnite (VPDB). Typical measurement precision is ±0.01–0.03‰ for atmospheric CO2. The primary international reference material for δ13C is calcite . Calcite must be converted to CO2 to implement the VPDB scale at individual laboratories, which has been shown to result in significant laboratory offsets  over the Industrial Period and the processes driving these changes. We review key applications for atmospheric δ13CO2 and Δ14CO2 observations from the literature, with an emphasis on global or large‐scale processes. Then we present new simulations of future changes in atmospheric δ13CO2 and Δ14CO2 corresponding to future emission scenarios through 2100. We discuss the impacts of these future changes on applications for atmospheric δ13CO2 and Δ14CO2 observations and make recommendations for observational and modeling activities for δ13C and Δ14C.In this paper, we review the observed changes in the 214C because they have been stored in geological reservoirs for millions of years, much longer than the 14C half‐life of 5,700 years. This gives fossil fuels a Δ14C signature of −1,000‰. For 13C, the carbon in fossil fuels has an isotopic signature (δ13C) that ranges from −44‰ to −19‰  increases while coal (δ13C ~ −24‰) decreases.Cement manufacturing also involves “fossil” carbon in that the source material is geological and therefore free of any 13CO2 and Δ14CO2: net transfer of carbon from the biosphere to the atmosphere (or vice versa), changes to the average 13C discrimination and its spatial pattern through changes in plant type such as the conversion of forest to pasture, and changes in the residence time of carbon in the biosphere. Overall, land use appears to have had small effects on global mean δ13CO2 and Δ14CO2 over the Industrial Period, in part because of responses of natural biospheric and ocean fluxes that compensate for land use effects on δ13CO2 and Δ14CO2. However, land use effects could be important regionally and for some applications  at the surface, modified by fractionation during gas transfer. Fractionation during gas transfer includes both kinetic and equilibrium effects , while also incorporating recent flask data from other laboratories by quantifying laboratory offsets. Ongoing activities to distribute reference materials of pure CO2 or CO2 in whole air show promise for improving measurement compatibility  because, in this case, the bomb‐derived 14C is mixed into a lower concentration of atmospheric CO2  does not have a strong influence on the long‐term δ13CO2 trend.The potential for long‐term trends in plant discrimination had also been neglected in global studies using the double deconvolution. Using historical δg et al.  found th13CO2 measurements are commonly used to investigate terrestrial biosphere activity on local or regional scales by estimating isotopic signatures of photosynthesis or respiration using the “Keeling Plot” approach. The “Keeling Plot” . Evidence for a regional Suess Effect had already appeared in comparisons of tree ring data  . The process for deliberate CO2 removal included in the SSP scenarios is Bioenergy with Carbon Capture and Storage (BECCS). In this way, the CO2 removal is mediated by an initial uptake into the terrestrial biosphere, which has implications for atmospheric δ13CO2  are now based on a set of five narratives, called the shared socioeconomic pathways (SSPs)  show an inflection in δ13CO2 that is more pronounced than for Δ14CO2. In SSP1‐1.9, δ13CO2 is approximately −7‰ in 2100, about the same as it was in 1940. SSP1‐2.6 and SSP5‐3.4os both have δ13CO2 of approximately −7.5‰ in 2100, after a stronger decline and reversal in SSP5‐3.4os compared to SSP1‐2.6. SSP2‐4.5 reaches a minimum of −9.4‰ in 2070 and then returns to −9‰ by 2100. SSP3‐7.0 and SSP5‐8.5 decrease through the century and reach −11.4‰ and −12.8‰, respectively, in 2100.Simulated δ13CO2 and Δ14CO2 for SSP1‐2.6, SSP5‐3.4os, and SSP5‐8.5. The contributions for SSP1‐1.9, SSP2‐4.5, and SSP3‐7.0 are shown in Figure −1 to zero while the negative influence of fossil fuel emissions strengthens slightly. Positive influences from biospheric exchange and from 14C production by natural cosmogenic radiation and by nuclear power plants have relatively constant positive influences of 3–4‰ year−1 over 2000–2015. Trend contributions of similar magnitudes were found in the early 2000s by Levin et al.  or 2060 (SSP5‐8.5). Over this time, the overall trend in Δ14CO2 trend slows, the biospheric disequilibrium weakens because Δ14C of previously assimilated carbon is more similar to present Δ14CO2. Sign changes in the Δ14CO2 trend lead to sign changes in the biospheric disequilibrium and influence on the Δ14CO2 trend. The effect is modulated by the turnover time of carbon in the biosphere, between 36 and 56 years , Australia's Commonwealth Scientific and Industrial Research Organisation (CSIRO), and Scripps Institution of Oceanography (SIO). Only one global network for Δ14CO2 is currently being operated, by the University of Heidelberg, although other global networks have operated in the past  and in North America by NOAA and other laboratories. Urban‐scale networks have also been developed  or the ICOS portal (https://www.icos‐cp.eu/).Currently, observations of atmospheric δ14CO2 and δ13CO2 away from the surface. There have been some stratospheric observations of Δ14CO2 conducted since the late 1980s  and at a larger network of sites for δ13CO2  simulate carbon isotopes in the land and ocean modules of their Earth System Models using a specified atmospheric boundary condition (Figure 13CO2 and Δ14CO2 in the atmospheric models of Earth System Models has not yet been implemented, so it is currently not possible to do a fully coupled simulation of δ13C and Δ14C using the most state‐of‐the‐art models. Such fully coupled isotopic models would be useful not only for the modern period but also for paleoclimate modeling. Other shared tools enabling atmospheric modeling of δ13CO2 and Δ14CO2 would also help to exploit existing and future atmospheric measurements.In addition to efforts expanding the observations of Δn Figure . Only on14CO2 observations to monitor regional emissions from fossil fuel combustion, many more observations and better modeling capabilities on regional scales are needed. Studies of regional atmospheric measurement networks combined with high‐resolution atmospheric modeling have only recently been published (Basu et al., 2 (ffCO2, Equation 2 and then run an inversion for fossil fuel and biospheric fluxes. In contrast, Basu et al. (14CO2 and CO2 fluxes across North America, including all the processes that can influence Δ14CO2 in the inversion. Other best practices that are still under development include the location and sampling height of observation sites in the network. Sites that have lower sampling heights or that are located closer to emission sources have higher signals in ffCO2, whereas sites with higher sampling heights or located further from sources have lower signals, but they represent larger regions. Having more than one observation site within a particular region can be important to prevent biases from any unique site characteristics, for example, including both urban and ex‐urban sites that differ not only in their ffCO2 signals or representation scale but also in the atmospheric model's representation of the transport in different types of regions (Brophy et al., 14C emissions from nuclear power plants can cause enrichment of Δ14CO2, particularly near to high 14C‐emitting reactors in the United Kingdom and Canada (Bozhinova et al., 14C emissions from nuclear power plants and better 14C emissions data would enable their effect to be accurately accounted for and improve the utility of Δ14CO2 for ffCO2 quantification in regions with nuclear power plants. Further development of regional networks for Δ14CO2 and complementary measurements including satellite observations, as well as the model‐data analysis frameworks for interpreting the observations to constrain ffCO2 emissions, are needed (Ciais et al., To fully develop the use of Δu et al.  set up t112 has undergone dramatic changes as a result of human activities and the response of the natural carbon cycle to them. The relative amount of atmospheric 14C and 13C in CO2 has decreased because of the addition of 14C‐ and 13C‐depleted fossil carbon, while the nuclear bomb tests increased 14C in the atmosphere in the 1950s and 1960s. Measurements of Δ14CO2 and δ13CO2 have been used to make invaluable contributions to our knowledge of atmospheric mixing, air‐sea gas exchange, plant function, and fossil fuel emissions. As fossil fuel burning continues to grow, the Suess Effect on 14C and 13C in CO2 continues. However, lower‐emission scenarios would lead to stabilized Δ14CO2 and increases in δ13CO2 over this century. The different paths described by the SSP‐based scenarios show that there is a wide range of possible changes to 14C and 13C of CO2 in the future. Researchers should be aware of the possible changes and their effect on the continued utility of 14C and 13C measurements of CO2 for scientific applications across various fields. We recommend a concerted effort to increase the number of 14C and 13C measurements across the Earth System and more development of publicly available modeling tools that incorporate 14C and 13C, including Earth System Models.Since the Industrial Revolution, the carbon isotopic composition of atmospheric COSupporting Information S1Click here for additional data file.Table S1Click here for additional data file.Table S2Click here for additional data file.


KMT2D leading to Kabuki syndrome. Deletion of KMT2D has been thought to be lethal, but here we describe a patient with KMT2D deletion and classical Kabuki syndrome phenotype.We present the first patient described with haploinsufficency of  Dear Editor,KMT2D and KDM6A are found in 70% of patients with KS (type 1 MIM#:147920 and type 2 MIM#:300867 respectively). KMT2D and KDM6A are both highly constrained genes (pLI = 1.0) suggesting that both genes are intolerant to heterozygous loss‐of‐function variation and thus haploinsufficient. While true haploinsuffiency through deletion of the entire locus of KDM6A is a well‐established cause of KS2 . Most molecularly confirmed cases of KS are from heterozygous pathogenic variants in KMT2D (94% n = 621/660) but in published reports only six cases (~0.9% n = 6/621) involve any type of KMT2D deletion larger than 20 basepairs , also known as Niikawa‐Kuroki syndrome, is a Mendelian disorder of the epigenetic machinery which occurs in approximately 1:32,000 births characterized by intellectual disability, facial and limb dysmorphic features, and postnatal growth retardation. Pathogenic variants in genes KMT2D with classical KS. Our patient is the first child of healthy parents born at term with a 2‐vessel cord and intrauterine growth restriction. She required intervention for hypothermia and feeding intolerance and was subsequently admitted to the pediatric intensive care unit (PICU) for hyperinsulinemic hypoglycemia. At 10 months of age, she is G‐tube fed with persistent hyperinsulinemic hypoglycemia requiring diazoxide and continuous overnight feeds, a phenotype seen in other children with KS1. She is hypotonic (trunk > extremities), with left hip dysplasia, joint laxity, and premature thelarche, all features previously seen in KS1. She has moderate bilateral hearing loss without a history of ear infections. An echocardiogram revealed a small‐medium atrial septal defect and a renal ultrasound was normal. She has the characteristic KS facial gestalt including elongated palpebral fissures, long eyelashes, arched eyebrows with sparseness of the lateral third of the brow, ptosis, a broad nasal root, flat nasal tip, high arched palate, micrognathia, and large cup‐shaped ears with an abnormally‐formed right ear pinna with pits . Of these, only three diseases are inherited in an autosomal dominant pattern, Selective Tooth Agenesis‐8 (WNT10B MIM#:617073), susceptibility to early onset osteoporosis (WNT1 MIM#:615221) and KS (KMT2D MIM#:602113). At 10 months, our patient has seemingly normal dentition with eruption of three normally formed teeth, and there have been no issues suggesting abnormal bone density. Our patient is the first known case to date of a nonmosaic whole‐gene deletion of KMT2D and presents with the typical clinical phenotype of KS. Early lethality due to haploinsufficency of KMT2D may be supported by the fact that on our evaluation of DECIPHER deletions and duplications that include all of KMT2D, there is marked overrepresentation of duplications (11/12), with the only deletion being very large (101.3 Mb). Though we cannot rule out reduced viability of pregnancies with a KMT2D deletion, deletion of KMT2D does not appear to be embryonic lethal as our patient is alive and well. Additionally, one would expect our patient to have a more severe clinical presentation than the average individual with KS if heterozygous deletion of KMT2D were to cause embryonic or in utero lethality; however, our patient's clinical presentation is no more severe than the average Kabuki patient, and from a congenital anomaly standpoint, her malformations are on the milder end of the spectrum.Microarray analysis revealed a 434.6 kb deletion of 12q13.1 (Figure KMT2D in a patient with classical features of KS. The relatively small size of the deletion identified leads us to conclude that the loss of KMT2D is the principal cause of her clinical phenotype. This supports the notion that a primary cause of KS is decreased amounts of KMT2D protein rather than altered function which can be seen with gain‐of‐function or dominant‐negative mutations.Our patient is therefore the first described case of constitutional whole‐gene deletion of HTB is a consultant for Millennium Pharmaceuticals.


Heteropolyacids such as tungstosilicic acid (HSiW) are Brønsted acid catalysts with higher DME production rates than γ-Al2O3, especially at low temperatures (140–180 °C). In this work, we show that the performance of supported HSiW for the production of DME is strongly affected by the nature of the support. TiO2 and SiO2 supported HSiW display the highest DME production rates of ca. 50 mmolDME/h/gHSiW. Characterization of acid sites via 1H-NMR, NH3-isotherms and NH3-adsrobed DRIFT reveal that HSiW/X have Brønsted acid sites, HSiW/TiO2 showing more and stronger sites, being the most active catalyst. Methanol production increases with T until 200 °C where a rapid decay in methanol conversion is observed. This effect is not irreversible, and methanol conversion increases to ca. 90% by increasing reaction pressure to 10 bar, with DME being the only product detected at all reaction conditions studied in this work. The loss of catalytic activity with the increasing temperature and its increasing with reaction pressure accounts to the degree of contribution of the pseudo-liquid catalysis under the reaction conditions studied.Dimethyl ether (DME) is an advanced second-generation biofuel produced via methanol dehydration over acid catalysts such as γ-Al DME can be also produced from biomass-derived syngas thus being considered a second-generation biofuel that can be used as diesel substitute or blend or as substitute for liquefied petroleum gas (LPG)3. DME is a gas at ambient temperature and pressure, but liquefies at 5.9 bars and 25 °C4, which is advantageous in terms of storage and handling. It is a non-toxic, non-dangerous chemical. Moreover, it is rapidly photodegraded to H2O and CO2 if released to the atmosphere. DME is widely used as aerosol propellant and due to its physical properties such as flammability and stability are similar to that of LPG, so it is used as LPG blend or as fuel for power generation in gas turbines1. In recent times, however, the interest in DME has shifted and due to properties such as high cetane number between 55–60, low auto-ignition temperature and its low emissions of soot DME is currently viewed as an ideal diesel substitute or blend.The production of dimethyl ether (DME) from biomass is attracting a great deal of attention in industry and academia. DME is mainly produced from natural gas or coal-derived methanol via dehydration processes3OH → CH3OCH3 + H2O) in gas phase over acid catalysts such as γ-Al2O3, zeolites or silica-modified alumina at moderate pressures between 10 and 20 bar, and temperatures 250–300 °C14. However, under such conditions undesired products such as olefins are produced and react further to coke species, thus leading to a rapid deactivation of the catalyst. It has been recently demonstrated that TiO2 supported heteropolyacids (HPA) exhibit higher reaction rates in methanol dehydration reaction than γ-Al2O3 under milder reaction conditions (1 bar and 180 °C) with 100% selectivity for DME17, thus preventing catalyst deactivation by coke formation/deposition.DME is typically obtained via methanol dehydration (2CH6) stabilized by protons, that surround a central tetrahedral unit (XO4). They are usually formulated as H8-n[Xn+M12 O40], where X is the central atom, n is the oxidation state and M is the metal ion. This arrangement constitutes the primary structure of the HPA, while the way in which these units arrange three-dimensionally combined with crystallization water is known as the secondary structure, which may be flexible. The presence of double bonds M = O in these structures results in the delocalization of the negative charge, conferring a high mobility to the protons of the HPAs’ structures. As a consequence, HPAs present strong Brønsted acidity and as a matter of fact the strength and density of the acid sites depend on the content of water, which forms hydrogen bonds. Due to this singular structure, the catalytic behavior of these materials is complex, with active sites located both on the surface and within the solid. Therefore, in addition to the typical heterogeneous catalysis that takes place at surface sites, including pore walls, reactants are absorbed within the bulk, that is, into the interanion space, where they are able to react as well. This latter feature is referred to as bulk-type or pseudo-liquid catalysis20.The Keggin structure is one of the most common and best known types of HPA structures for catalytic applications, as well as the best-known one. It is formed by four units of three octahedrons  and silicotungstic acid  have been reported among the most active HPA structures for alcohol dehydration22, which have specifically demonstrated to achieve high conversion and selectivity in the DME synthesis reaction23 at temperatures as low as 140–160 °C and atmospheric pressure. For instance, when supported on inorganic carriers such as TiO2, methanol conversions as high as 80% and 100% DME selectivity at 200 °C have been reported23. In these works, the beneficial effect of supporting the HPAs on TiO2 was clearly demonstrated and TiO2 supported HSiW roughly doubles the methanol conversion and DME productivity achieved by non-supported HSiW at temperatures ranging from 140 to 160 °C24. Recent studies have also confirmed the beneficial effect of supporting HPA on carriers such as SiO2, BN, and carbon nanotubes26. However, comprehensive investigations about the effect of the support are still lacking in the literature.In particular, phosphotungstic acid (H28.The superior performance of supported HSiW and HPW for the synthesis of DME from methanol has been only demonstrated at low temperatures (below 200 °C) and pressures (1 bar). These reaction conditions may not be feasible and economically favorable for the production of DME in large-scale commercial plants, in which methanol production from syngas takes place at high temperatures (240 °C and higher) and pressures (20–70 bar). Moreover, this feature is critical for the direct synthesis of DME from syngas since the reactor must operate at sufficient temperatures and pressures (well above 180 °C and 1 bar) to efficiently transform syngas into methanol  previous to producing DME2O3, producing only DME at all reaction conditions studied in this work. The catalytic performance of the supported HSiW materials for DME production depends on the nature of the support, being the catalyst supported on TiO2, ZrO2 and SiO2 the ones displaying the highest methanol conversions. Whereas methanol conversion remains stable at temperatures below 190 °C, increasing reaction temperature at 200 °C or above leads to a rapid decline of methanol conversion with time on stream without affecting product selectivity. However, methanol conversion increases to its initial value by increasing reaction pressure. This behavior is ascribed to the contribution of pseudo-liquid catalysis under the reaction conditions studied in the manuscript.In this work, we show that when supported onto inorganic carriers, silicotungstic acid (HSiW) is a more active catalyst for the dehydration of methanol to DME than γ-Al4SiW12O40·xH2O, Sigma-Aldrich, ≥99.9%) in ethanol (Panreac) to the corresponding support. Different supports have been studied in this work namely TiO2 , SiO2 (BASF D11-11), CeO2 , BN , ZrO2  and Al2O3 . The quantities of HSiW and support were chosen by taking into account that the loading of HSiW in the catalyst should satisfy the relationship 4.5 Keggin units per square nanometer (KU/nm2) of the support. This loading has been previously demonstrated to lead to the highest conversion of methanol per gram of catalyst for heteropolyacids supported on TiO223. Maintaining constant this value means that the weight percentage of HSiW for each catalyst is different, depending on the specific surface area of the support. After impregnation, the catalysts were held for 24 hours at room temperature and then dried at 60 °C overnight. Catalysts were named as HSiW/X, where X identifies the support. Bulk HSiW and γ-Al2O3  were dried and used as reference catalysts.Supported HSiW catalysts  were synthesized following the impregnation at incipient wetness method, consisting in the dropwise addition of a solution of HSiW , using a powder X-ray X´Pert Pro PANalytical with a configuration θ − 2θ, using CuKα radiation, with an Anton Paar XRK900 for the x-ray diffraction data collected in a range of temperatures between 20 and 550 °C. The crystallite sizes of the HSiW phases on each catalyst were calculated by applying the Scherrer Eq. 4SiW12O40·6H2O) is stable in all catalysts.k is the Scherrer constant (k ≈ 0.94), λ is the incident wavelength, θ is the diffraction angle and βsize is the experimental full-width at half-maximum (FWHM). This equation is employed to obtain an estimated value of the size of the X-ray coherent diffracted crystalline domains (d).The specific surface areas of the catalysts were obtained from the NAs) using the size of the particles determined by XRD and approximating to a spherical geometry using Eq. 230:d is the diameter of the particles and ρ is the heteropolyacid bulk density calculated for H16O46SiW12·6H2O .We determined the mass specific surface areas  was used to record the DRIFT spectra. Before collecting the spectra, the catalysts were pretreated at 220 °C in a current of He (20 mL min−1) during 30 min. NH3 was used as probe molecule to identify acid sites on the catalysts. A flow of 20 mL min−1 of NH3/He (5/95 vol) was passed through the DRIFT cell during for 15 minutes. The cell was purged under a 20 mL min−1 flow of He; the spectra were recorded at 100 °C. For methanol adsorption experiments, the methanol was introduced into the DRIFT cell by bubbling He (20 mL min−1) through a glass saturator filled with methanol at room temperature during 30 min. After methanol adsorption/reaction, the samples were treated in He flow at increasing temperatures from 140 to 400 °C at 10 °C min−1.A Bruker AV400-WB spectrometer was used to collect the solid-state NMR spectra of the catalysts, at 2 (13% vol. methanol) were fed to the reactor with a HPLC pump and a TOHO TTM 005 mass flow controller, respectively. The methanol was evaporated and mixed with the N2 before being introduced into the reactor. The effluents of the reactor were measured on line with a Varian CP-3800 gas chromatograph equipped with a SPB-5, 60 m x 0.53 mm capillary column connected to a flame ionization detector (FID) where methanol and DME were separated and analyzed. Catalytic tests were carried out at different pressures , different temperatures (140–240 °C) and maintaining the methanol liquid hourly spatial velocity (LHSV) constant at 1.11 h−1. Catalysts were treated in situ in a N2 flow at 220 °C during one hour prior to each experiment. Methanol conversion and DME production rate were calculated from Eqs. 3OH flowin represent the number of moles of methanol fed to the reactor per unit of time.The catalytic activity of the bulk and supported HSiW samples in the methanol dehydration reaction was measured in a fixed bed reactor placed into a furnace to control its temperature, which was measured by a thermocouple in contact with the catalytic bed. The catalytic bed is formed by 0.2 g of catalyst diluted in SiC up to 1.8 mL to minimize the formation of significant exotherms in the bed, with a particle size between 250 and 300 µm. Methanol and N2 g−1 with no porosity.The textural properties of the catalysts and of the supports, along with the actual HSiW loading in each catalyst are shown in Table The specific surface area and the pore volume of the HSiW supported catalysts are lower than those of the bare supports, but higher than that of bulk HSiW. The pore diameters of the HSiW/X catalysts are characteristic of each material, but all lie within the mesoporous range, between 2 and 50 nm.2/N2 20/80 vol. As observed, concerning the HPA phase, the diffractograms of all HSiW/X samples recorded at room temperature show weak and broad reflections corresponding to H4SiW12O40·xH2O with a water content between 6 and 21 molecules32, probably H4SiW12O40·14H2O, with the most intense reflections around 8 and 29° 32. However, the high dispersion of the HPA species makes it complicated to determine accurately the exact structure and actual number of water molecules within the HSiW crystalline structure. When the temperature increases above room temperature, between 25 and 150 °C, the diffractograms display the expected diffraction lines  for hexahydrated HSiW (H4SiW12O40·6H2O)16, which has a crystal structure similar to H3PW12O40·6H2O33. The diffractograms collected at temperatures between ca. 200 and 450 °C show a displacement of the reflections into a dehydrated phase, namely H4SiW12O4034. It is difficult to determine the actual phase of HSiW in HSiW/Al2O3 from the XRD data, probably due to the high dispersion of the HPA on the support36. Eventually, with the increasing temperature, the structure of the HSiW collapses leading to the formation of WO3, which is the only W-containing crystalline species (together with the supports) observed in the diffractograms recorded at 550 °C.In order to investigate if the structure of HSiW has been modified during the impregnation process and to assess the stability of the HSiW phases with temperature, we rely on XRD and Raman analyses. Fig. 2, ZrO2, BN and CeO2) present the diffraction reflections of crystalline TiO2, ZrO2, BN, and CeO2 phases, respectively, in all the temperatures measured. However, HSiW/X (X = SiO2 and Al2O3) do not show the reflections of the supports . The absence of diffraction lines for silica and alumina-supported HSiW catalysts has been reported by other authors, and it is attributed to a loss of crystallinity of the support after the impregnation process36.Concerning the supports, HSiW/X  of HSiW/X calculated from Eq. 2 < HSiW/TiO2 < HSiW/ZrO2 < HSiW/BN < HSiW/CeO2 < HSiW/Al2O3.As shown in Table −1, assigned to νs(W = Ot) and νas(W = Ot), respectively37, which confirms that the heteropolyacid maintain the Keggin structure after the process of impregnation. The results obtained by Raman spectroscopy are in good agreement with the XRD analyses, even for the HSiW/Al2O3, and indicate that the HSiW structure is stable after deposition on the support.Fig. 1H-NMR and DRIFT of NH3-adsorbed. 1H-NMR spectra of the catalysts are shown in Fig. 4SiW12O4037. This peak can be observed in the spectra of all supported catalysts. However, both the actual position and the shape of the peaks are characteristic of the nature of the support, and therefore it can be used as a descriptor of the HPA-support interaction. Thus, the 1H-NMR peaks observed in the spectra for HSiW are almost identical to those of HSiW/SiO2 and HSiW/TiO2, indicating the lack of interactions between the protons of the HSiW and the TiO2 or SiO2 supports. On the other hand, the peaks in the spectra of HSiW/ZrO2, HSiW/BN, HSiW/CeO2, and HSiW/Al2O3 appear at lower chemical shifts than the peak of HSiW, suggesting the development of an interaction between the HSiW and the supports. The shifting of NMR peaks of supported HSiW vs. bulk HSiW has been observed previously and ascribed to the the weakening of the acidity of the protons in HSiW once they are supported38. As also observed in Fig. 2 and HSiW/CeO2 display two peaks at 8.1 and 7.5 ppm which reveal the presence of two types of protons with different acidity. Peaks at lower shifts, ascribed to the supports, are also observed in the spectra of HSiW/SiO2 39, HSiW/ZrO2 (at 5.3 ppm)40, HSiW/Al2O3 41 and HSiW/BN (at 1.6 ppm)42. 1H-NMR results show that the acidity of protons in HSiW/TiO2 and HSiW/SiO2 is the same as in the bulk HSiW.The acidity of the catalysts was evaluated by 3 as a probe molecule. The 3500–1000 cm−1 region of the spectra of NH3 adsorbed on bulk and supported HSiW catalysts at 100 °C is shown in Fig. 4+) bands at 3232, 3041 and 2809 cm−1, and δas(NH4+) and δs(NH4+) bands at ca. 1677 and 1425 cm−1, respectively46. The spectra of supported HSiW also show very weak bands due to ammonia adsorbed on Lewis acid sites at 1677–1620 and 1235–1331 cm−1 assigned to δas(NH3) and δs(NH3), respectively45. The identification of Lewis acid sites on HSiW/SiO2 is not possible because the bands overlap with the strong features of SiO2 between 1100 and 1350 cm−1. The spectrum for bulk HSiW shows no bands that can be attributed to Lewis acidity. It has been reported that the position of the δs(NH3) band is sensitive to the Lewis acid strength of the adsorbing sites, shifting at higher frequencies for stronger Lewis acid sites44. The bands assigned to δ(NH3) appear at higher wavenumbers values in the spectra of HSiW/Al2O3 than in the spectra of HSiW/TiO2 and HSiW/ZrO2 suggesting a stronger acidity of the Lewis acid sites on HSiW/Al2O3.The nature of the aid sites in bulk and supported HSiW has been further investigated by DRIFT spectroscopy by using NH2 and HSiW/TiO2 reach methanol conversions close to the equilibrium at temperature as low as 180 °C. Importantly, the selectivity for all HSiW catalysts towards DME is 100%, irrespectively of reaction temperature. Although all HSiW catalysts display high activity, the actual conversion of methanol is characteristic of each catalyst, ranging between 90 and 10%. HSiW/SiO2, HSiW/TiO2, and HSiW/ZrO2 display the highest methanol conversions in the series, of ca. 90%, 80% and 60%, respectively. These values are higher than that of the bulk HSiW (ca. 40%). HSiW/Al2O3 shows similar methanol conversion than HSiW at all temperatures studied. Finally, BN and CeO2 supported HSiW show the lowest conversions in the series of ca. 25% and 5%, respectively. Fig. i.e., the ones supported on ZrO2, TiO2 and SiO2 also record the highest DME production rates, with HSiW/TiO2 and HSiW/ZrO2 showing the highest DME production rate of 50 mmolDME h−1 gHSiW−1 at 180 °C.Bulk (HSiW) and supported (HSiW/X) samples were tested in the methanol dehydration reaction at 140, 160, and 180 °C and 1 bar; conversion values are reported after at least 4 hours on stream. Fig. 2O3 and zeolites (HZSM-5 and HY), which display high activity at temperatures above 260 °C, being inactive below 200 °C47. Mesoporous aluminosilicates and aluminium phosphates have been also reported active methanol dehydration catalysts, but only reach methanol conversions higher than 50% at temperatures above 300 °C or 250 °C, respectively47. Other Lewis acid catalysts such as WO3/TiO2 or Nb2O5/TiO2 are known to have limited methanol conversion activity, below 10%, with DME selectivity ranging between 80–95%13. Only few articles describing the role of supported HPA as catalyst for methanol dehydration have been reported48. Alharbi et al. reported a methanol conversion of 13% over HSiW/SiO2 at 120 °C48, in line with 30% methanol conversion at 140 °C reported in our work. Ciftci et al.15 studied a H3PW12O4-based catalyst for the DME synthesis and found that methanol conversion increased with temperature, but decreases above 200 °C. Similar results were observed by Schnee et al., who reported that methanol conversion over bulk H3PW12O4 at 200 °C declines with TOS49. These results are in line with the trends observed in our work (see below).These activity values are significantly higher than the values of state-of-the-art γ-Al2 and HSiW/Al2O3 show lower activity than bulk HSiW, with HSiW/CeO2 showing the lowest activity in the series. On the other hand, SiO2, TiO2 and ZrO2, which presented the highest catalytic activities, have smaller BET specific surface areas (except HSiW/BN) than HSiW/CeO2 and HSiW/Al2O3. This observation supports the idea that surface catalysis is not the only main route for the dehydration of methanol to DME on supported HSiW. Noticeably, the activity trend displays an inverse correlation with the As values, and the catalysts showing lower dispersions, i.e., SiO2 and TiO2 supported HSiW, show the highest methanol conversions in the series. This feature indicates that the presence of HSiW aggregates is beneficial for the catalytic process, which is in good agreement with the predominant role of the pseudo-liquid catalysis. Note also that SiO2 and TiO2 supported HSiW catalysts are the ones showing the more acid protons (see 1H-NMR results), probably because the lower degree of dispersion of HSiW onto SiO2 and TiO2 results in a lower interaction between HSiW and the support, favoring proton mobility50. By contrary, by being well dispersed on the supports, a stronger interaction between HSiW and Al2O3, BN and Ce2O3 is developed, decreasing proton mobility and consequently the acidity of the catalysts, thus explaining the shifting of the 1H-NMR peaks to lower values observed in Fig. A straightforward relationship between the catalytic activity and the BET specific surface areas of HSiW/X is not observed; the catalysts with the highest BET areas Table , namely 3 adsorption isotherms. The densities of acid sites normalized to the mass of the catalyst and to mass of HSiW are summarized in Table 3 adsorption isotherms do not discriminate between the acid sites of the support and the HSiW.The acidity of the catalysts that displayed the highest activity, that is, those with methanol conversions higher than that of bulk HSiW, was quantified from NHNH3/gcat for HSiW/ZrO2 to 0.476 mmolNH3/gcat for HSiW/Al2O3. These values are actually lower than that of bulk HSiW of 0.513 mmolNH3/gcat. The different number of acid sites of each catalyst appears to account for the different HSiW loading on each catalyst  of the catalysts, except for bulk HSiW and HSiW/Al2O3. The latter catalysts show high acidity but low activity. The lower activity of the alumina-supported HSiW in comparison to TiO2 or SiO2-supported HSiW has already been reported before51. However, correlating the DME production activity of HSiW/Al2O3 with its acidity should be taken cautiously since, as explained above, the NH3 adsorption technique cannot be used to discriminate between acid sites associated to HSiW or Al2O3, but the latter acid site are not active for DME production from methanol at the low temperatures studied in this work. These findings indicate that not only the number of acid sites, but also the nature of the support and its interactions with the HPA, contribute to the observed catalytic performance. In fact, considering again the 1H-NMR spectra it seems like the catalysts that showed the weakest HSiW-support interactions are the ones that record the highest methanol conversions (HSiW/TiO2 and HSiW/SiO2).As shown in Fig. 2, HSiW/TiO2, and HSiW/ZrO2 as compare to bulk HSiW despite the higher amount of acid sites in the latter. Complementary studies regarding the methanol adsorption mechanism were carried out by in situ DRIFTS with the aim to explain the behavior of the different supported catalysts. Fig. In addition, the accessibility of the methanol molecules to the heteropolyacid acid sites has been suggested to be critical for the catalytic performance. This feature can explain the highest methanol conversion activity observed with HSiW/SiO2O) mode of physisorbed water between 1616 and 1634 cm−1 52. As observed in Fig. 2O molecules were previously adsorbed. This pattern is observed in the spectra for HSiW/SiO2, HSiW/TiO2, and HSiW/ZrO2 but not in the case of HSiW/Al2O3. Indeed, the band ascribed to physisorbed H2O remains visible for the Al2O3 supported HSiW sample even after 30 minutes of methanol flow, suggesting a very strong retention of the H2O molecules that would prevent methanol from adsorbing and, consequently, reacting onto the HSiW/Al2O3. This explains the low conversion obtained with this catalyst, and suggests that the same situation might be happening with the HSiW/BN and HSiW/CeO2 catalysts.After thermal treatment, and prior to methanol adsorption, all samples showed a band ascribed to the δ and pressures (P ≥ 10 bar). This feature is of an utter relevance for the direct synthesis of DME from syngas whereby syngas transformation to methanol and subsequently to DME should be carried out in a single reactor with bifunctional catalysts.As discussed above, HSiW/X catalysts are significantly more active (and selective) for the production of DME from methanol than the benchmarking catalysts γ-Al2, ZrO2 and TiO2 supported HSiW catalysts, since they showed the highest methanol conversions in the series. Catalysts supported on BN, CeO2 and Al2O3 were excluded from these experiments since they recorded lower methanol conversions than non-supported HSiW. As observed in Fig. 2 catalysts from 76% to 53% after less than 3 hours on stream. The decrease in methanol conversion over HPA catalysts at 200 °C and 1 bar at increasing TOS has been also reported by other authors49. However, as observed in Fig. 2 and SiO2 supported HSiW catalysts, i.e., methanol conversion decreases at during time on stream at 200 °C and 1 bar, increasing back to its initial value by increasing reaction pressure to 10 bars. Increasing further reaction temperature to 220 °C (at 10 bar) results in a progressive  decline of the methanol conversion  and pressures (up to 24 bar) with SiOion Fig. . Again, 85% Fig. . Finally85% Fig. , that ca85% Fig. . This tr2O3 at 230 °C upon DME adsorption53. Another extensive work on coking of zeolites during methanol conversion was carried out by Schulz, in which the temperatures considered relevant to study this process start at 270 °C54. These temperatures are higher than the ones explored in our work and we observed that methanol conversion decreases swiftly at temperatures as low as 200 °C  can be ascribed to catalyst deactivation. The most typical cause of deactivation of acid catalysts during methanol  dehydration is the formation of coke deposits. Deactivation by coke formation is a progressive process that takes place at higher temperatures than those used in this work. For instance, it has been reported that no coke is formed over γ-Al °C Fig. . More im °C Fig. . MoreoveThe results presented in this work demonstrate that HSiW based catalysts can be used for the production of DME from methanol at high temperatures (above 200 °C) provided the reaction is conducted at higher pressures. The pressure should be determined as a function of the temperature to allow the pseudo-liquid catalysis to take place.2, TiO2, and SiO2, which are more active than bulk HSiW. This enhancement is ascribed to the better accessibility of the methanol to the active sites of the catalyst, enabling the interchange of crystallization water by the reacting molecule. The nature of the support is crucial in the performance of the supported HSiW; the development of strong support-HSiW interaction results in a lower mobility of the water located between the polyanions, preventing the contribution of the pseudo-liquid catalysis. Similarly, operating at temperatures above 180 °C prevent the access of methanol to the active sites between Keggin units and the loss of the pseudo-liquid character of the catalytic process. This effect can be avoided by increasing reaction pressure.A series of supported HSiW catalysts were synthesized through the incipient wetness impregnation method. XRD and Raman analyses reveal that the Keggin structure of the HSiW remained unaltered after deposition on the supports. All catalysts show high activity for the dehydration of methanol to DME at low temperatures, especially those supported on ZrOSupplementary information.