Formaldehyde hydration constant

Pyrolysis of the ethylene acetal of bicyclo[4.2.0]octa-4,7-diene-2,3-dione yields a-(2-hydroxyphenyl)-y-butyrolactonc 11 a mechanism involving a phenyl ketene acetal is proposed. Tartrate reacts with methanediol (formaldehyde hydrate) in alkaline solution to give an acetal-type species (9) 12 the formation constant was measured as ca 0.15 by H-NMR. Hydroxyacetal (10a) exists mainly in a boat-chair conformation (boat cycloheptanol ling), whereas the methyl derivative (10b) is chair-boat,13 as shown by 1 H-NMR, supported by molecular mechanics calculations. 

The net effect of this cyclohexadiene/phenyl ring insertion at the carbonyl group is to cause an increase in the overall equilibrium constant for the addition of solvent water, from A dd = 2.3 x 103 for hydration of formaldehyde154 to A dd = 4.0 x 107 for hydration of the p-quinone methide l,3 so that Kj = AT. dd/Aiadd = 1.7 x 104 for transfer of the elements of water from formaldehyde hydrate to 1 (Scheme 43). We have proposed that the relatively small driving force of 6 kcal/mol for this transfer of water from CH2(OH)2 to 1 represents the balance between larger opposing effects3 ... 

The simplest of the aldehydes is formaldehyde, whose oxidation by Ce(IV) in 2.0 M perchlorate media has been studied by Husain (1977). It is presumed that formaldehyde exists as a ketohydrate in acid solution (the hydration constant is lO M ). Michaelis-Menten kinetics describe the results, indicating the formation of a precursor complex. A detailed mechanism permits a calculation of the equilibrium quotients for formation of the Ce(IV)-formohydrate complex and the ionization of a... 

FIGURE 16.29 The magnirnde of the equilibrium constant (A) depends on the relative energies of the carbonyl compounds and their hydrates. We know that for acetaldehyde the energies of hydrate and aldehyde are comparable, because A" 1. Acetone is more stable than acetaldehyde and formaldehyde is less stable. Steric factors make the acetone hydrate less stable than the acetaldehyde hydrate. By contrast, the formaldehyde hydrate is more stable than the acetaldehyde hydrate. The equilibrium constants reflect these changes. 

The exceptions are formaldehyde, which is nearly completely hydrated in aqueous solution, and aldehydes and ketones with highly electronegative substituents, such as trichloroacetaldehyde and hexafluoroacetone. The data given in Table 8.1 illustrate that the equilibrium constant for hydration decreases with increasing alkyl substitution. 

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. 

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. 

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. 

Formaldehyde has a large equilibrium constant for hydrate formation because it has no bulky, electron-donating alkyl 2 X I03 groups. It is more than 99.9% in the hydrated form in... 

These stability effects are apparent in the equilibrium constants for hydration of ketones and aldehydes. Ketones have values of Keq of about 10-4 to 10-2. For most aldehydes, the equilibrium constant for hydration is close to 1. Formaldehyde, with no alkyl groups bonded to the carbonyl carbon, has a hydration equilibrium constant of about 40. Strongly electron-withdrawing substituents on the alkyl group of a ketone or aldehyde also destabilize the carbonyl group and favor the hydrate. Chloral (trichloroacetaldehyde) has an electron-withdrawing trichloromethyl group that favors the hydrate. Chloral forms a stable, crystalline hydrate that became famous in the movies as knockout drops or a Mickey Finn. 

The equilibrium constant for hydration is especially large for formaldehyde, trichloroacetaldehyde, and cyclopropanone. Explain. 

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. 

A review in Russian on the synthesis of carbohydrates from formaldehyde has appeared. An investigation of the formose reaction by g.c.-n.m.r. has shown that intermediate glycoaldehyde, glyceraldehyde, and dihydroxyacetone are present as mixtures of monomers, e.g. hydroxy carbonyl compounds, epoxides and hydrates, and dimers such as half and full acetals. Further study of the barium chloride-catalysed formose reaction at pH 12 has shown that the main product forming 33% of the total sugars is the branched pentulose (1). The sugar yield reached a constant value at 70% completion of the reaction, i.e., within... 

As these equations indicate, hydrations of aldehydes and ketones are reversible. The equilibrium lies to the left for ketones and to the right for formaldehyde and aldehydes bearing inductively electron-withdrawing substituents. For ordinary aldehydes, the equilibrium constant approaches unity. 

Table 19.2 lists the values for some representative hydration reactions. These equilibrium constants show the same trends found for the addition of HCN to a carbonyl group. Aldehydes with low molecular weights readily form hydrates. Formaldehyde is over 99% hydrated. Its hydrate is called formalin, a 37% by weight solution of formaldehyde in water that was used in the past to preserve biological specimens. Other aldehydes are substantially less hydrated. Ketones are normally hydrated less than 1%. The hydrates of aldehydes and ketones usually cannot be isolated and exist only in solution. 

On the basis of these findings, Auerbach concluded that polymeric hy-di-ates other than the trimeric hydrate HO - (CH20)a-H ai C not present in appreciable proportions m solutions containing up to 30 per cent formaldehyde. Deviations encountered in the mass action constants for more methylene glycol form can be calculated from the apparent molecular weights. Values obtained in this way are shown in Table 3a. 

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