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P R E P R I N T – ICPWS XV Berlin, September 8–11, 2008 Kinetics and Equilibrium of the Reversible Formic Acid Decomposition in Hot Water Yoshiro Yasaka, Ken Yoshida, Chihiro Wakai, Nobuyuki Matubayasi, and Masaru Nakahara Institute for Chemical Research, Kyoto University Email :[email protected] During the past decades, a variety of technology has been proposed for the application of hydrogen energy. The fundamental problems are its storage and the transportation, and they have not yet been solved. This article explains how to use formic acid as a fuel; the advantages of carbonaceous fuels and hydrogen are fulfilled by formic acid. The strategy is shown here how to produce and utilize the formic-acid fuel on the basis of the detailed kinetic and equilibrium studies on the reversible formic acid decomposition in hot water, Introduction To overcome the environmental issues caused by extensive and rapid fuel consumption and hazardous chemical processes in industry is one of the greatest subjects for the chemistry of the 21st century. Hot water is a promising candidate that would play a key role in recovering the sustainable energy and material (carbon) cycles on earth. More and more attention is paid for the hydrogen cycle; see Figure 1. The most attractive potential of hot water for this application is that hot water can dissolve various chemical species including C1 (carbon one) molecules, plastics and biomass, and can induce unique hidden reactions among them in the absence of metal catalysts, thus allowing us to convert less useful materials into more useful ones [1]. The atmospheric carbon dioxide issue, which is now the most serious and urgent environmental problem all over the world, has its root in human activities interfering with the natural carbon cycle. The accumulation of carbon dioxide in the atmosphere can be interpreted by considering a large gap between the oxidation rate of fossil fuels by human beings (fast step) and the reduction rate of carbon dioxide by nature (slow step); see Figure 2. We are going to completely exhaust the millionsof-years-old fossil fuels in just a few centuries. The carbon dioxide problem is strongly related with the energy supply problem because the oxidation process is accompanied by useful energy of up to 800 kJ per one mole of carbon (44 g CO2). Fortunately, we receive at any moment plenty of solar energy in comparison to the energy we need for our daily activities. The solar energy we receive on earth amounts to 5730 TW, while the global energy consumption by human beings amounts to the one-thousandth [2]. A key point of the energy problem is the fact that so far we make use of only a small portion of solar energy. We have just started a large-scale fixation of solar energy in the form of a controversial man-made fuel called bioethanol. Can we take advantage of the unique potential of Figure 1. Carbon and hydrogen cycles for energy through formic acid. the gaseous species. This is a kind of chemical evolution. The carbonaceous energetic compounds produced from solar energy are very convenient to store because they are in the liquid or solid state. Typical examples are such fossil fuels as oil and coal. When they are burned, however, they are not so clean as hydrogen energy because of the accompanying production of nitrogen oxides, sulphur oxides, and carbon micro particles in exhaust fumes. Hydrogen is ultimately a clean energy resource but is extremely difficult to store because of the explosivity and lowest condensation temperature as seen in Figure 3. This is a huge obstacle we are faced with in developing hydrogen energy technology. The WGS reaction via formic acid is a key reaction that allows us to take advantage of the benefits of both types of fuels; we can store and transport formic acid just as we store petroleum and can use it as if it were hydrogen by the application of the second step of reaction (2). Formic acid is neither explosive, flammable nor very volatile; as a drawback it is weakly acidic. For these reasons, formic acid can be a potential manmade fuel alternative to bioethanol. Any weakly reducing carbonaceous materials can be used for the input energy for the generation of this new fuel from water. When we succeed, we need not rely upon corn or sugarcane for bioethanol. The strategy for the synthesis is shown in Figure 4. As can be seen, the partial combustion of waste carbonaceous materials yields carbon monoxide, hot water to go beyond bioethanol? We must reduce CO2 emission by 60-80 % by the year of 2050 (EU summit meeting on May 2007). Formic Acid as a Clean Fuel The water-gas-shift (WGS) reaction [3] CO + H2O CO2 + H2 (1) is one of the oldest hydrothermal reactions studied. It was industrially developed in the early 20th century for the steam-reforming process. However, the reaction mechanism under non-catalytic conditions has not been established for a long time. The first attempt to detect formic acid as an intermediate of the WGS reaction was done by Rice et al. using a flow system made of metal in supercritical water [4]. After their unsuccessful results, we have carried out series of studies in much milder reaction conditions [5], and after all succeeded in detecting formic acid produced from carbon monoxide in sub-critical water, which then decomposes into carbon dioxide and hydrogen [5c]. The WGS reaction is then rediscovered as a reversible transformation process between carboncontained chemical energy (HCOOH) and hydrogen-contained chemical energy (H2): CO + H2O HCOOH CO2 + H2 (2) The equilibrium can be shifted to either side (CO+H2O or H2+CO2) by tuning the thermodynamic conditions. Furthermore, we have shown that the reaction can be kinetically controlled in order to synthesize formic acid from CO2 Sun breathing Artificial Fuels in Future photosynthesis fossilization Life Fossile Fuels Figure 2. Materials and energy flow on earth. 2 (a) cell technology (electricity H2 ) and the reversible reaction (2) make this strategy very hopeful. The grand design of the formic acid fuel drafted above requires fine control of the WGS reaction via formic acid. The following sections introduce the basic hydrothermal chemistry of formic acid [5]. CO + H2O HCOOH Formic Acid Fuel CO2 + H2 Formic Acid Decomposition in Relation to the WGS Reaction 100 °C (b) H2 N2 In recent years, [5-10] formic acid is found to decompose without catalysts in the following ways: HCOOH CO +H2O (decarbonylation) (3) HCOOH CO2 +H2 (decarboxylation) (4) Using NMR, we have found that the decarbonylation is reversible and that carbon monoxide can be directly converted into formic acid in hot water [5c]. We also made a theoretical analysis on the Gibbs energies of the species (HCOOH, CO, CO2, H2, and H2O) involved in reactions (3) and (4) [5d]. The reversibility and the coupling of the reactions given by eqs 3 and 4 clearly indicate that formic acid exists as an intermediate in the WGS reaction as in eq 2. To comprehensively understand and control the WGS reaction, it is essential to determine the rate and equilibrium constants for the elemental steps shown in eqs 3 and 4. The kinetics and equilibrium study was hampered by the following: (i) the reaction order of the reactions (3) and (4) has not yet been established, (ii) these reactions can be catalyzed by metals used for reaction cells [5b,7,10], (iii) although the non-catalytic decarbonylation can be dominant in hot water, the reaction rate is very slow (the time scale of days to weeks below ~200 °C), and (iv) it is rather difficult to quantitatively analyze such main gaseous products as CO, CO2, and H2 in the heterogeneous system because of their distribution between the gas and liquid phases. Here we show that the most of these difficulties can be overcome by studying the decomposition of the formic acid in the presence of acid but in the absence of metal surface, because under such condions the two competitive decomposition pathways in eqs. 3 and 4 can be separated to obtain each reaction rate constant. The difficulty in the quantification of the products can be overcome by applying 1H and 13C NMR spectroscopy to the liquid and gas phases of the sealed reactor, with the carbon mass balance confirmed. 101 °C CH4 H2O HCOOH -160 °C -196 °C -253 °C Figure 3. The comparison of the condensation temperature of fuels and related compounds (b) involved in the formic-acid formation reaction (a). which is then hydrothermally converted into formic acid. The strategy as a whole is equivalent to the fixation of the solar energy, since the biomass used is grown by sunlight. The growth of plantlife is the most efficient and cheapest way as the first step of the solar energy. At the present stage, in fact, the global energy fixation rate by plants is estimated to be ~10 times faster than the global energy consumption rate by human activities [2]. Another energy source for the production of formic acid fuel is the electricity from power plant working with natural energy. Although wind power and solar power plants show a high performance in extracting natural energy, the output is unstable due to the temporal fluctuation and discontinuity of the given natural energy. This is a drawback and significantly reduces the overall efficiency of the power plants. It is thus essential to develop a way for the large-scale temporary energy storage to balance the energy supply and demand. The use of a formic-acid chemical tank is one of possibilities for this energy storage. The combination of the fuel 3 Partial combustion biomass waste plastic H2O CO Formic Acid Fuel WGS Reaction Fuel Cell Natural energy power plant H2 CO2 Figure 4. Two ways of producing formic acid fuel. that the main product was CO2 [6]. Other authors also reported at higher concentrations (0.13 M and 1 M) that the decarboxylation is the main decomposition path [7,8]. McCollom et al. examined the decarboxylation of formic acid using a gold bag reactor with and without minerals and clarified that the catalytic effect is not given by minerals but by the transition metals except for gold [10]. They also observed the formation of formate ion from CO2 and H2 in basic conditions, which indicates the reversibility of the decarboxylation. In the following part, we review our most recent studies [5] and discuss the results in relation to the formic acid fuel technology. Hydrothermal chemistry receives further attention because a number of organic reactions in hot water involve formic acid as a key product or intermediate in hot water [5,11]. We need to establish the kinetics and mechanisms of the hydrothermal formation and decomposition reactions of formic acid, in view of the following three processes [11]. (1) Formic acid, as an aldehyde (C1, single-carbon contained), is involved in the non-catalytic cross-disproportionations of aldehydes; an aldehyde is reduced to the corresponding alcohol and formic acid is oxidized to carbonic acid or carbon dioxide. (2) Formic acid reacts with formaldehyde (C1 aldehyde) in the presence of acid to form a C2 compound, glycolic acid, through C-C bond formation as a chemical evolution process. (3) Carbon monoxide, which is in a hydrothermal equilibrium with formic acid, is released through the decarbonylation of aldehydes that can be generated by the heterolytic fragmentation of ethers. The decomposition of formic acid has been studied for a long time, but much attention was paid to the catalytic effect of the metal surface. Our group and Bröll et al. reported the dominance of the decarbonylation path in the decomposition at low temperatures without kinetic analysis [5a,5b,9]. The decarboxylation pathway was much more quantitatively studied [6-10], but the reaction rate was obtained mainly under the surface catalytic effect of metals. Savage and co-workers studied the decomposition of formic acid in a flow reactor made of Hastelloy C-276 at 0.02 M (mol dm-3) in the temperature range of 320-500 °C and showed Experimental A quartz tube reactor was used in this study as shown in Figure 5. The reaction solution was introduced to occupy a specified portion (controlled with an accuracy of 1%) of the tube reactor. The volume (length, lTliq) ratio of the liquid phase to the total (L) is hereafter denoted as the filling factor: ρT = liq the volume of the liquid phase lT = the total volume of the reactor L (5) where the suffix T represents the temperature. Note that the filling factor ρT at a reaction temperature T is larger than that at room temperature 30 °C, ρr, due to the expansion of water when ρr > 0.32 (at densities higher than the critical density of water 0.32 g cm-3). The gaseous reaction product is in a dissolution equilibrium at each temperature. The 4 value of ρT can be calculated from ρr by using the tabulated density of water along the saturation curve [12]. The quartz tube reactor containing the reaction mixture was heated for a specified reaction time in a programmable electric furnace kept at high reaction temperature (170-330 °C) with an accuracy of 1 °C. The reaction proceeds in an equilibrium with the gases. After the sample was quenched by air, 1H and 1H-decoupled 13C NMR spectra were observed at 30 °C with a NMR spectrometer (JEOL ECA400, magnetic field strength of 9.4 T). Such gaseous products as CO, CO2, and H2 coexist in the liquid and gas phases and the NMR spectra were measured for each phase by the method described elsewhere [11a]. The time evolution of the concentrations was measured by repeating the above procedure. The concentrations of all species (HCOOH, CO, CO2, and H2) were quantified on the basis of their integrated peak intensity. For example, the concentration of formic acid [HCOOH] was estimated by the 1H peak intensity relative to that of solvent H2O (55.5 M) in the liquid phase. Then the fraction of residual formic acid, denoted as xHCOOH, is expressed by x HCOOH = of the gaseous products such as CO was summed over the liquid and gas phases. For the computation we used the following equation: [CO]liq + [CO] gas × xCO = 1 − ρr ρr [HCOOH]0 (7) Here, [CO] and [CO] are the concentrations of CO in the liquid and gas phases, respectively, obtained at 30 °C. The same procedure was used for CO2. No other side-reactions took place because the carbon mass balance was maintained during the reaction within 10%. The source of error is mainly the absence of an internal reference in the estimate of the gaseous products. The maintained mass balance also indicates that CO and CO2 do not leak through the reactor walls. However, H2 leaked through the quartz walls very slowly as pointed out in the original paper [5f]. Because the reaction is carried out in heterogeneous (sub-critical) conditions, the gaseous reaction products are present in both phases. For this reason, we need to take into account the dissolution equilibrium of the gases at the reaction temperature when determining the equilibrium constant for the decarbonylation defined as: liq [ HCOOH ] [ HCOOH ]0 K CO = (6) where [HCOOH]0 is the initial concentration of formic acid. The yields of gaseous products were determined in a similar way but by using an exteral reference for integrated intensity. The concentration gas [CO ] liq [HCOOH ] liq (8) eq where the symbol |eq represents the value evaluated at equilibrium. See the following sections in detail. Figure 5. The experimental procedure used in this study and the accompanying change of the filling factor at each step. The illustration shows the reactor tube (a) at room temperature before the initiation of the reaction, (b) during the reaction at a high temperature, and (c) after being rapidly cooled to room temperature. The filling factor at the reaction temperature is larger than that at room temperature due to the decrease of water density at elevated temperature for ρr > 0.32. For (b) and (c), there are also shown the reversible decarbonylation of formic acid and the dissolution equilibrium of carbon monoxide in the reactor. Note that the decarbonylation reaction is quenched in (c). 5 Decarbonylation Kinetics Decarbonylation Equilibrium Under the catalytic effect of metals, the decarbonylation pathway can be controlled to be a minor one. However, we have found that in the absence of the metal catalysts, formic acid decarbonylates much faster than the decarboxylation. When the HCl concentration is larger than 2.0×10-2 mol kg-1, the decarbonylation is exceedingly faster (more than 50 times) than the decarboxylation and become practically the only decomposition pathway. In this situation, the initial rate of carbon-monoxide production is equal to that of the formic-acid decomposition. We can assume the rate law of the decarbonylation as We have determined the equilibrium constant of the decarbonylation of formic acid in hot water in the temperature range of 170-280 °C. We have separated the time scale of the decarbonylation (fast) from that of the decarboxylation (slow). What is directly obtainable from the yield analysis is not KCO but a quantity denoted as the decarbonylation quotient QCO: v+1 = − QCO = d[ HCOOH ] b = k +1[H + ]a [ HCOOH]0 dt 0 log (k+1/ mol-1 kg s-1) -4.82 -4.22 -3.71 -3.29 -2.98 -2.75 -2.43 -1.91 -1.20 x HCOOH eq (11) where xHCOOH and xCO are provided by eqs 6 and 7, respectively. As the reaction takes place in subcritical water along the saturation curve, not all of the CO produced by the decarbonylation remains in the liquid phase, but a considerable amount of CO is accommodated in the gas phase of the sample at high pressures. Hence QCO is larger than KCO. To determine KCO we need to calculate the fraction, β, of CO distributed in the liquid phase. (9) Here v+1 is the reaction velocity, k+1 is the rate constant, [HCOOH]0 is the initial concentration of formic acid. The parameters a and b are the reaction orders with respect to the proton (HCl) and formic acid, respectively. As is shown in ref [5f], they were determined as unity. Thus we have v+1=k+1[H+][HCOOH] (10) This rate law for the decarbonylation is considered to be valid also in the absence of acid catalysts or in basic conditions; the decomposition of sodium formate produced no detectable carbon monoxide. On the basis of the reaction rate law (10), we determined the second-order rate constant k+1 in the temperature range of 170-280 °C on the liquid branch of the saturation curve. The results are summarized in Table 1. T/C 170 185 200 210 220 230 240 250 260 280 300 x CO β= moles of CO in the liquid phase total moles of CO (12) Then the relation of QCO and KCO is written as: K CO = βxCO xHCOOH = βQCO (13) where formic acid is considered to be present only in the liquid phase. The value of β is dependent on the solubility of CO at the reaction temperature and the filling factor defined by eq 5. The value of β has been determined on the basis of the following assumptions: (1) the dissolution of the gases is sufficiently faster than the decarbonylation and the decarboxylation. (2) The Henry law holds for the dissolution of CO into hot water. (3) The equation of state for the ideal gas holds for CO in the gas phase. The Henry’s constant kH is the mole ratio of liquid H2O to CO dissolved in hot water which is in a dissolution equilibrium with hypothetical ideal gas of CO at 1 MPa. When the partial pressure of CO is given by pCO/MPa, we have the molar (volume) concentration of dissolved CO as KCO 0.15 0.33 0.44 0.80 1.2 4.2 Table 1. The rate constant k+1 and the equilibrium constant KCO for the decarbonylation, and the acid dissociation constant Ka of formic acid in the temperature range of 170-330 °C on the liquid branch of the saturation curve. [CO]liq = 6 pCO dT kH (14) where dT is the molar concentration of water at the reaction temperature T on the saturation curve. The gas-phase molar concentration of CO is expressed by [CO]gas = pCO RT Decarboxylation Kinetics The kinetics of decarboxylation were found to be more complicated than those of the decarbonylation. The proton concentration does not linearly increase the decarobxyration rate. We measured the decarboxylation rate of formic acid in the presence of HCl as well as the decarboxylation rate of sodium formate. The results in Table 2 are summarized on the following rate law: (15) according to the ideal gas law; R is the gas constant. Using eqs 8 and 11-15, we can evaluate the value of β as d[CO 2 ] = v+ 2 = k [ HCOO − ] dt = k acid [ HCOOH ] + k base [ HCOO − ] [CO ]liq ρ T β= [CO ]liq ρ T + [CO ]gas (1 − ρ T ) = ρT (19) K D (CO) (1 − ρ T ) + ρ T (16) This rate law was previously postulated by Maiella et al. [8] and McCollom et al. [10]. The precision of the rate constants in Table 2 is relatively low (~10 %) because the surface of the quartz tube is corroded in basic conditions and the reaction analysis can be carried out only at an early stage. Two rate constants kacid and kbase are of the same order. This indicates that the decarboxylation rate changes only weakly with pH in contrast to the decarbonylation. The magnitude relation between kacid and kbase changes at ~280 °C, and at the higher temperature of ~ 330 °C, kbase overwhelms kacid. Taking advantage of the different pH dependence of the decarbonylation and decarboxylation rates, the reaction path control of the hydrothermal decomposition of formic acid can be realized. The decarbonylation is the major path in acidic conditions due to the catalytic effect of the proton, while the decarboxylation becomes dominant when pH > 4. where the liquid-gas distribution constant is introduced by K D (CO ) = kH [CO ]gas = RTd T [CO ]liq (17) and ρT is the filling factor at the reaction temperature. The Henry’s constant kH taken from ref [13] and the molar concentration dT of pure water tabulated in ref [12] are used to evaluate KD(CO). As the temperature increases along the saturation curve, liquid water is expanded and CO becomes more soluble in water in the temperature range studied here; KD(CO) = 29.6, 13.8, and 8.4 at 170, 240, and 280 °C, respectively. The value of ln KD(CO) and its temperature dependence are in good agreement with a previous computational study [5d]. Now the equilibrium constant KCO is explicitly written in terms of the directly measurable quantities QCO and ρT as K CO = ρT K D (CO)(1 − ρT ) + ρT T / °C 200 230 240 260 280 300 330 QCO (18) On the basis of this equation, we can determine KCO by measuring the equilibrium mole fraction xHCOOH(eq) of formic acid under a given filling factor ρT. We can evaluate QCO in eq 18 from eq 11; xCO(eq) = 1-xHCOOH(eq). The equilibrium constant KCO as the average of two measurements (13 measurements for 240 °C) of different filling factors are listed in Table 1 at 170-300 °C. log (kacid / s-1) -6.6 -5.7 -6.1 -5.9 -5.4 -4.6 -4.1 log (kbase / s-1) -6.4 -5.5 -4.9 -4.3 -3.4 Table 2. The rate constants kacid and kbase in the temperature range of 200-330 °C on the liquid branch of the saturation curve. 7 Decarboxylation Equilibrium Control of WGS Reaction for Formic Acid Fuel. Here we discuss the decarboxylation equilibrium. We followed the decarboxylation for a sufficiently long time and the time evolution is shown in Table 3 at 330 °C. As can be seen, the fast decarbonylation attains equilibrium within an hour, and this pre-equilibrium is maintained (xCO/xHCOOH ~ 5) during the course of the slow decarboxylation. The parallel decrease of formic acid and carbon monoxide almost stops at ~20 h and this indicates that the decarboxylation equilibrium is attained. The observed yield of hydrogen is slightly smaller than the 1 equiv. of CO2, and this probably indicates the leakage of hydrogen through the reactor walls (the time scale of the leakage is estimated to be ~100 h from long-time experiments). The equilibrium constant KCO2 of the decarboxylation is defined by In the previous sections, we have established the kinetics and equilibrium of the decarbonylation and the decarboxylation of formic acid which constitute the hydrothermal water-gas-shift (WGS) reaction. Based on the above analysis, here we propose a strategy how we can store energy in the form of formic acid fuel. As we have mentioned in the introduction the energy source for the production of formic acid fuel is carbon monoxide obtained from any carbonaceous materials. To convert carbon monoxide into formic acid, it is necessary to suppress the WGS reaction at the stage of the formic acid intermediate; see Figure 6a. The above kinetic study on the decomposition of formic acid has proven that in the presence of HCl, the decarbonylation has a much shorter time scale than does the decarboxylation. This is an ideal situation for the formic acid fuel production. The conversion ratio of CO into formic acid is, in turn, determined by the equilibrium constant KCO of the decarbonylation. The expected yield of formic acid xHCOOH was calculated using the data in Table 1 and is shown in Figure 7 as a function of temperature and the filling factor ρT at the reaction temperature. Better formic acid yield is obtained at a lower temperature and higher filling factor due to the favorable KCO and the increased partition of carbon monoxide into water. K CO2 = [CO 2 ]liq [ H 2 ]liq [HCOOH ]liq eq (20) By taking account of the liquid-gas distribution of H2 and CO2, the value of KCO2 at 330 °C is estimated to be 1×102 mol kg-1. Our result shows that the decarboxylation equilibrium strongly favors the CO2 side. This is consistent with the thermodynamic calculations by McCollom et al. [10], and our statistical-mechanical calculations [5d]. xHCOOH 1 0.27 0.04 0.014 0.0067 xCO 0 0.49 0.30 0.10 0.056 0.036 0.034 0.039 xCO2 0 0.25 0.66 0.89 0.94 0.96 0.96 0.96 (a) xH2 0 0.09 0.66 0.79 0.78 0.77 0.75 0.77 CO+H2O (b) CO2+H2 very slow fast HCOOH CO2+H2 very slow fast HCOOH CO+H2O Figure 6. Control of the WGS reaction to produce formic acid. xHCOOH t/h 0 0.33 3.2 6.0 9.0 12.0 15.0 19.7 Table 3. The time evolution of the reactant and product yields for the decomposition of 1.0 mol kg-1 formic acid at 330 °C and ρT ≈ 0.99 (the volume of the gas phase is negligible). The temperature is too high to follow the decarbonylation kinetics. 1.0 ρ T = 0.30 0.8 0.50 0.80 0.6 0.4 0.90 0.2 1.00 (totally filled) 0.0 180 200 220 240 o T/ C 260 280 Figure 7. The temperature dependence of formic acid yield xHOOH. Each of the contours corresponds to a fixed filling factor ρT indicated in the figure. 8 Another possible source of the formic acid fuel is the electric energy from wind and solar power plants. The electricity is first used to electrolysis water into hydrogen and oxygen. The hydrogen is then converted into formic acid which can be conveniently stored. The conversion of hydrogen into formic acid accompanied by carbon dioxide fixation is the reverse reaction of the decarboxylation of formic acid. To convert hydrogen into formic acid, the reaction conditions should be chosen so that the decarbonyaltion rate of formic acid is much slower than that of the reverse decarboxylation; see Fugure 6b. Such conditions would be satisfied in the presence of such a weak base as sodium hydrogencarbonate (note: H2 + NaHCO3 HCOONa + H2O) or metal catalyst such as SUS316 or more powerful but expensive rhodium complexes. The equilibrium formic-acid yield from hydrogen and carbon dioxide is determined by the value of the equilibrium constant KCO2. In basic conditions, the neutralization Gibbs energy change for formic acid and carbon dioxide should be also taken into account. Although the KCO2 value determined in acidic conditions implies an unfavorable formation of formic acid, a higher formic acid yield can be expected in basic conditions [10] and at a lower temperature [5d] in the presence of metal catalysts. References [1] (a) R. Noyori: Supercritical fluids: introduction. Chem. 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