Structure–reactivity correlation for reactions of peroxide anion nucleophiles with substituted acyloxybenzenesulfonate bleach activators

Abstract

Second-order rate constants are reported for the reaction between a series of esteric bleach activators with different acyl substituents, namely, acetyloxybenzenesulfonate, n-butanoyloxybenzenesulfonate, n-nonanoyloxybenzenesulfonate and isononanoyloxybenzenesulfonate, and a set of peroxide nucleophiles whose basicity ranges from below that of the oxybenzenesulfonate leaving group to above it. The results conform to a Brønsted-type relationship with nuc , 0.42 ± 0.01 for acetyloxybenzenesulfonate, very similar to the value of 0.40 ± 0.01 for para-nitrophenyl acetate and a range of peroxide nucleophiles. A significantly larger value of nuc , 0.56 ± 0.05 is obtained with n-nonanoyloxybenzenesulfonate. This increase in nuc is interpreted in terms of steric and polar interactions between the acyl substituent and the attacking nucleophile. The reactivity of n-butanoyloxybenzenesulfonate is similar to that of n-nonanoyloxybenzenesulfonate whilst that of isononanoyloxybenzenesulfonate is less. These results are discussed in terms of the practical consequences for bleach activation and with regard to the mechanism of acyl transfer from esters to peroxides.

Borate‐Catalyzed Reactions of Hydrogen Peroxide: Kinetics and Mechanism of the Oxidation of Organic Sulfides by Peroxoborates

“…The pK a of HOOB(OH) 2 , the parent acid of diperoxoborate, is 5.7, compared with 8.98 for B(OH) 3 , the parent acid of monoperoxoborate (these values are consistent with HOO À being less electron donating than HO À ). [16] Figure 5 shows that the peroxoborates are less reactive than peroxycarboxylic acids, though this disadvantage lessens with increasing pH above their pK a values, at which the peroxycarboxylic acids become ionised.…”

“…The rate constants for the peroxoborates are very similar (dashed line on Figure 5) despite their different pK a values. This shows that boron-oxygen bond breaking is not significant in the transition state, consistent with the peroxoborates not acting as nucleophiles in which the neutral B(OH) 3 or HOOB(OH) 2 moieties are the respective leaving groups.…”

“…[15] It is, on the one hand, rather surprising, perhaps, that anionic peroxoborates should show a high reactivity towards nucleophiles, particularly anionic ones (and, as the present paper will report, especially since the uncharged, electron-deficient peroxoboric acid, HOOB(OH) 2 is less reactive). This may have led some authors to suggest that HOOB(OH) 3 À can behave as a nucleophile, with B(OH) 3 acting as an uncharged leaving group. [10,11] On the other hand the bisperoxotungstate monoanion is a well-known anionic electrophile, so the existence of effective negatively charged electrophiles cannot be contested.…”

“…[1,2] Although it is a powerful oxidant, the reactions of hydrogen peroxide are generally rather slow, and the challenge is to overcome this kinetic barrier in more cost-effective and "green chemical" ways. Hydrogen peroxide is activated through the formation of peroxycarboxylic acids, either from a reactive carboxylic acid derivative with a good leaving group (as in the peroxide bleach activators) [3] or from the parent carboxylic acid and a strong acid catalyst in preparative reactions. [1] Activation by the formation of peroxycarboxylic acids entails a sharp decrease in atom-efficiency that is avoided in true catalysis.…”

“…The present paper will report that borate is a better catalyst than hydrogen carbonate, with the Abstract: The kinetics of the oxidation of substituted phenyl methyl sulfides by hydrogen peroxide in borate/boric acid buffers were investigated as a function of pH, total peroxide concentration, and total boron concentration. Second-order rate constants at 25 8C for the reaction of methyl 4-nitrophenyl sulfide and H 2 O 2 , monoperoxoborate, HOOB(OH) 3 À , or diperoxoborate, (HOO) 2 [9] Peroxoborates are rapidly formed in solutions of hydrogen peroxide and borate and their salts are selectively crystallised from solution to form perborates, which are available commercially and have synthetic as well as bleaching and cleaning applications. [10,11] Over fifty years ago Edwards reported that the pH of borate buffer solution drops as hydrogen peroxide is added, due to the rapid equilibrium formation of peroxoborates.…”