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A combined microflow reactor and short-path-length spectroscopy cell along with the accompanying process controls are described to obtain real-time, in situ transmission IR spectra of reaction components of aqueous solutions up to 725 K and 335 bar. Quantitation of the spectra was required to obtain kinetics and equilibrium constants. The extinction coefficient of CO2 in H2O at 275 bar was found to increase monotonically from 1.52 × 106 at 298 K to 2.26 × 106 cm2 mol-1 at 573 K. Also, CO2 dissolved in H2O was rotationally quenched on the IR time scale below 375 K but progressed into rotational diffusion around 625 K and finally essentially free rotation above 700 K. The kinetics and pathway of hydrothermolysis of urea to CO2 and NH3 were determined directly from spectral data at 473−573 K. Good agreement was obtained between experimental and calculated concentration−time data by using a reaction model consisting of (NH2)2CO → NH4 + + OCN- and NH4 + + OCN- + H2O → CO2 + 2NH3. The Arrhenius parameters for the first-order reaction are E a = 84.2 kJ mol-1 and ln A (s-1) = 17.5, and for latter pseudo-second-order reaction are E a = 58.5 kJ mol-1 and ln A (L mol-1 s-1) = 17.1. The global rate of formation of CO2 without the kinetic model is first-order and has different Arrhenius parameters. As part of this study, the species of 0.1 m (NH4)2CO3 equilibrium were determined at 298−650 K and 275 bar. The equilibrium shifted from the hydrolyzed ionic components at lower temperature to the neutral CO2, NH3, and H2O components at higher temperature. Therefore, the (NH4)2CO3 equilibrium does not influence the kinetic model of urea above about 475 K.

Infrared spectroscopy measurements of the kinetics and decomposition pathways of aqueous urea ((NH2)2CO, 200−300 °C, 275 bar) and guanidinium nitrate ([(NH2)3C]NO3, 240−300 °C, 275 bar) are described. A Pt/Ir alloy flow cell with diamond wafer windows was used, and heat and fluid transport models show that isothermal and plug flow conditions exist. The hydrothermolysis of urea was modeled by the conversion of urea to NH4 + + OCN- followed by hydrolysis to CO2 + 2NH3. These reactions are a subset of those for hydrothermolysis of guanidinium nitrate. Decomposition of guanidinium nitrate is catalyzed by the formation of NH3. The reaction scheme involves deprotonation of the guanidinium ion by NH3 to produce neutral guanidine, which hydrolyzes to form urea as the rate-determining step. The subsequent hydrothermolysis chemistry follows that of urea. Thus, although the overall decomposition rate of guanidinium nitrate is slower than that of urea, the Arrhenius parameters for the step for formation of CO2 from guanidinium nitrate and urea are similar [E a ≅ 66 kJ/mol, ln A (s-1) ≅ 19]. Only a small difference in these values is incurred by using a 316 stainless steel−sapphire cell in place of the Pt/Ir−diamond cell. Hence, wall effects appear to be small for this reaction.

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The rates and pathways of decarboxylation of acetic acid derivatives, RCO2H, and their Na+ salts, RCO2Na, which possess electron-withdrawing groups (R = CCl3−, CF3−, HOC(O)CH2−, NH2C(O)CH2−, CF3CH2−, NCCH2−, CH3C(O)−) were determined in H2O at 100−260 °C and a pressure of 275 bar. Simple conversion to RH + CO2 occurs in most cases, except that H2O appears to be a required reactant for the anions. Real-time FTIR spectroscopy was used to determine the rate of formation of CO2 in flow reactors constructed of 316 stainless steel (SS) and of titanium. With a few exceptions, the rate of decarboxylation is similar within the 95% confidence interval in 316 SS and Ti and the difference is smaller than that caused by R. Therefore, while wall effects/catalysis may exist in some cases, it plays a lesser role in the relative rates than the substituent R. The acid form of the keto derivatives decarboxylates more rapidly than the anionic form, whereas the reverse is true for the nonketo derivatives. In keeping with the greater role of H2O as a reactant, the entropy of activation for the anions is smaller or more negative than for the acids. A Taft plot of the decarboxylation rates suggests that the mechanistic details can be interpreted in terms of the various roles of R. Where R = HOC(O)CH2− and NH2C(O)CH2−, decarboxylation occurs faster than expected, probably because a cyclic transition state can exist. The rate is slower than expected for R = CF3−, perhaps because of stabilization of the acid by hyperconjugation. The mechanism of decarboxylation of acids of the remaining R groups is similar and the steric effect of R is somewhat more influential than its electron withdrawing power.