Formaldehyde hydration rate constant

The type of catalyst influences the rate and reaction mechanism. Reactions catalyzed with both monovalent and divalent metal hydroxides, KOH, NaOH, LiOH and Ba(OH)2, Ca(OH)2, and Mg(OH)2, showed that both valence and ionic radius of hydrated cations affect the formation rate and final concentrations of various reaction intermediates and products.61 For the same valence, a linear relationship was observed between the formaldehyde disappearance rate and ionic radius of hydrated cations where larger cation radii gave rise to higher rate constants. In addition, irrespective of the ionic radii, divalent cations lead to faster formaldehyde disappearance rates titan monovalent cations. For the proposed mechanism where an intermediate chelate participates in the reaction (Fig. 7.30), an increase in positive charge density in smaller cations was suggested to improve the stability of the chelate complex and, therefore, decrease the rate of the reaction. The radii and valence also affect the formation and disappearance of various hydrox-ymethylated phenolic compounds which dictate the composition of final products. 

Calculations support a cooperative mechanism for the hydration of formaldehyde, acetaldehyde, acetone, and cyclohexanone in water. The results are supported by determination of the rate constant for the neutral hydration of acetone, using labelled acetone and water. Conclusions include ... 

For example, for formaldehyde (R = H) at neutral pH, the pseudo-first-order rate constant for the hydration reaction (forward reaction), k =k [H20], is about 10 s 1 and the first-order rate constant for dehydration, k2, is about 5x 10-3 s-1. In Chapter 20 we will use this example to show that the reactivity of compounds can influence the kinetics of air/water exchange if both processes (reaction and exchange) occur on similar time scales. 

Consider two aldehydes at neutral pH, formaldehyde and acetaldehyde. The hydration/ dehydration (pseudo-) first-order rate constants and the nondimensional Henry s law constants are summarized below. Since in the following discussion you are interested in orders of magnitude only, you assume that aqueous molecular diffusiv-ities of all involved species are the same as the value for C02, (DIW = 2 x 10 5 cm2s 1) and that the corresponding values in air are the same as the value for water vapor (Dwateri = 0.26 cm2s 1). This allows us (as a rough estimate) to calculate v,w and v,a directly from Eqs. 20-15 and 20-16, respectively. 

The usual means of finding general catalysis is to measure reaction rate with various concentrations of the general acids or bases but a constant concentration of H30 +. Since the pH depends only on the ratio of [HA] to [A-] and not on the absolute concentrations, this requirement may be satisfied by the use of buffers. Catalytic rate constants have been measured for a number of acids and bases in aldehyde hydration-dehydration, notably by Bell and co-workers.10 For formaldehyde, a = 0.24, /3 = 0.40 earlier work11 gave for acetaldehyde a = 0.54, /3 = 0.45 and for symmetrical dichloroacetone a = 0.27, /3 = 0.50. 

There are abundant kinetic data for addition of water to carbonyl compounds either uncatalyzed, acid catalyzed, or base catalyzed. For a considerable number of these reactions the equilibrium constant for hydration is also available and thus an extensive test of NBT is possible [73]. Over the entire range of reactivity for which data are available (from formaldehyde plus hydroxide to NA -dimethylacetamide plus water), the calculated AG values were in good agreement with experiment. At the time this paper was published [73], the rate of uncatalyzed hydrolysis of dimethylacetamide had not been reported. Since then Wolfenden and coworkers [74] have reported a rate constant at elevated temperatures extrapolating to a AG of 32kcal/mol in good agreement with our prediction of 31.11 kcal/mol. 

In principle the velocity of dehydration could be measured if a physical rather than a chemical method were available for removing the unhydrated carbonyl compound at a rate comparable to its hydration. It was claimed by Bieber and Triimpler (1947a) that this could be achieved by the removal of formaldehyde in a rapid gas stream, the rate of which appeared to be dependent on the pH of the solution. However, attempts to repeat their experiments have proved unsuccessful moreover, although they give no experimental details, calculation in terms of known kinetic and equilibrium constants shows that for a 1-ml liquid sample a gas flow of at least 30 litres/min would be required to produce an appreciable perturbation of equilibrium conditions (Bell and Evans, 1966). It is thus clear that this method has no practical application, at least to formaldehyde solutions. 

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