Indian Journal of Chemistry Vol. 42A, November 2003, pp. 2683-2697
Advances in Contemporary Research
Acid-base equilibria of hydroxamic acids: Spectroscopic investigation
Kallol K Ghosh School of Studies in Chcmistry, PI. Ravishankar Shukla University. Raipur 492010, CG., India E-mail: [email protected] Received 3 J March 2003: revised 22 August 2003
This work reviews the acid-base equilibria of hydroxamic acids. Hydroxamic acids and metal hydroxamates have recently enjoyed an enormous expansion boosted by new discoveries on their high applicability in the cardiovascular diseases, HIV, malaria, cancer, metal poisoning and Alzheimer's disease. Unfortunately the basic structural chemistry of hydroxamic acids still unresolved problems. Although there are various spectroscopic (NMR, IR, Ff-ICR-CID, Raman UV) and theoretical methods available for the study of acid-base behaviour, UV method is rapid, simple and reliable. In this review recent advances on acid-base behaviour has been discussed. Evaluation of the relative basicities and protonation parameters of polyfunctional acids and bases is an important task that has attracted much interest both as a source of information on the molecular electronic structure and as a tool to interpret the reactivity of acid-base catalysed reactions.
Dr. Kallol Kumar Ghosh is presently a Reader of Chemistry at Pt. Ravishankar Shukla University, Raipur. He received his M.Sc. degree (Physical Chemistry) in 1981 and Ph.D. degree in 1986 from Pt. (I) Ravishankar Shukla University under Prof. S.G. Tandon. He is fascinated by the unique properties of hydroxamic acids and hi s Despite the biological importance and rich research is focused on acid-base catalysis, micellar catalysis and chemistry of hydroxamic acids, their acid-base structure-reactivity correlations. He has done postdoctoral behaviours have received little attention so far. The work on artificial siderophores in the laboratory of Prof. Akira acid-base properties of organic molecules are among Katoh, Seikei University, Tokyo, on a INSA-JSPS visiting the most relevant factors affecting their reactivity, and fellowship. therefore studies leading to the definition of such Hydroxamic acids (I), both natural and synthetic, properties through equilibrium measurements yield linear as well as cyclic, play a variety of roles in extremely valuable data, particularly for the medicine, agriculture and engineering for example, as elucidation of reaction mechanisms. siderophores for iron (111) 1, as potent and selective An acid-base reaction (protonation and proton 2 3 inhibitors of enzymes such as peroxidases , ureases , abstraction) is the most fundamental elementary 4 5 and matrix metalloproteinases , and as hypotensive , chemical event (Eq. I), and leads to a large increase antioxidant6 and as corrosion inhibitors 7 and redox in the reactivity of neutral organic molecules when 8 switches . These acids have recently been widely used they are converted to their charged, conjugate acid Of as key functional groups of potential therapeutics base form. Hence, the long standing general interest targeting cardiovascular diseases, HIV, and in methods capable of probing the extent of the 9 l2 Alzheimer's disease , malarialo,ll, allergic disease , proton-transfer process. Moreover, hydroxamic acids cancer13 ' 14 an d d'la b etes.15 A mazmg . I y, suc h a b roa d are polyfunctional, i.e. , exhibit at least two spectrum of activity of hydroxamic acids is almost conceivable protonation or deprotonation sites, and exclu sively associated with structure, reactions and the nature of the ionized species is of interest in itself. complexation. The structural chemi stry of hydroxamic acids contains sti ll unresolved problems. ... (1 ) like the reason for their relatively hi gh acidity and the structure of the corresponding anions l6 . As th eir name An accurate knowledge of the protonation and implies, hydroxamic acids behave as weak acids (pK, deprotonation behaviour is required both for - 8-9) and, like most carbonyl compounds, they are structure-reactivity correlations and for the detailed I6 iS also weak bases - . analysis of acid-base catalysed reactions. 2684 INDIAN J CHEM, SEC A, NOVEMBER 2003
Several spectroscopic techniques have been employed to elucidate acid-base process of hydroxamic acids. The significant contributions of 19 22 23 24 8agno et al. - , Garcia et al. . and Crumbliss 25 28 et al. - deserve special mention in this context. Besides these spectroscopic studies several theoretical studies at the semi-empirical, INDO, AMi, and ab initio level have been reported on the protonation and ·H' 29 35 deprotonation of hydroxamic acids - . In recent years, our group has also been involved in 38 41 H-
... (2) Writing Eq. 3 in logarithmic form results in Eq. 4 .
(4) The mathematical definition of KBH+ is like that of Ka; writing 'a' for activities and 'I' for molar activity coefficients, Eq. (3) is obtained. where log! = log (CBH+/CB), called ionization ratio. GHOSH: ACID-BASE EQUILIBRIA OF HYDROXAMIC ACIDS
At any given acid concentration, I is calculated medium for their outstanding protonating ability. The from Eq. 5, which requires knowledge of the UV spectral curves change with increase in aciel observable 0 (absorbance or chemical shift) of the concentration. In order to prevent the substrate from free (013 ) and protonated bases (O BH+) any hydrolysis effect at hi gh acidity level s. th e solutions used were always prepared a fre sh. The 1= (0 - OB)/(013H+ - 0 ) ... (5) pioneering work on the protonation behaviour of 17 hydroxamic acid was reported by Tillett et (1 1 . The protonation equilibria of some para substitutcu The protonation constants pK13H+ have been benzohydroxamic acids (R.C H4.CO.NHOH; R = H. calculated from Eqs (6-11). 6 p-MeO, p-Me, p-OH, p-Cl) have been determineci in Hammett acidity function method (HAFM/ 9 sulphuric acids at 25 °C. The UV absorption spectra (1i" equimolar solutions at different acid concentrations log 1= - I1IHo + pKBH+ ... (6) may undergo changes not only because of th e protonation process, but also as a consequence of th e Corrected Hammett Oeyrup acidity function method medium change, so in many cases no isosbestic point (HAFM)50 is observed. The author and his group have been active in this area. The benzohydroxamic acid, BH A . .. (7) (II), and some 4-substituted derivatives ha ve been chosen for protonation studies in HCI , H 2 S0 ~ and Bunnett-Olsen method (BOM) 49 HCI04 (ref. 42) .
... (8)
52 Some authors represent Bunnett-Olsen equation in another form (Eq. 9). (X = H. OMe, Me, F, CI , Br, NO") ... (9) (II)
Cox-Yates excess acidity method (EAM)53 This important work expands knowledge of acid base behaviour of weak bases, carefully quantifyin g ... (10) acidities of the protonated forms. Medium and protonation effects have been adequately separated by Marziano-Cimino-Passerini method (MCP)54 characteristic vector analysis55.57 . This method decomposes the experimental absorbances readi ngs ... (11) into the mjnimum number of independent components capable of reproducing the spectral curves. frolll where CH+ is the molar concentration of solvated which reliable ionization ratios can be obtained. The protons, Ho and HA are acidity functions, Me is the absorbance A at r wavelengths was obtained in n acti vi ty coefficient function, X is the excess acidity different acidities, and the data were arrayed to form n fun ction and m, m*,
Table I-Protonation costant (P KB~ I +) of some hydroxamic acids
S. No. Hydroxmic acid Condition Method PKBH + Ref.
1. C.H.CONHOH Aqu.oUI H.SO, ' 2S"C UV -1 .g3 17. 24. 40 H-H-OH, 10% dioxin. H.SO, , 2S"C UV -2.10 • -E:)--<: • 0 10% dloxan. HCIO, , 2S"C UV -1.94 X • H, OM • • CI, M., OH, NO •. F
2. X-C.H,CON(OH)C.H. H.SO, 10% (v/v) 1 ,4-dloxan., 2S"C UV - 2.30 24, 17 ~t: x~ HCIO, 10% (v/v) l ,4-dloxan •• 2S"C UV - 2.g2 X • H, 4-CH., 4-CI
3. OH-C.H,CONHOH H.SO, H-N-OHt 10% (v/v) l ,4-dloxan., 2S"C UV -1.27 <9t -o 10% (v/v) 1 ,4-dloxane, 2S"C UV - 2.1 0 HCIO, 10% (v/v) l,4-dloxan., 25"C UV - 1.55
4. CH.CONHOH Aqu.oul H.SO, NMR - 1.15 10% (v/v) l,4-dloxane, 2S"C UV - 2.20 HCIO, 10% (v/v) l ,4-dloxIM, 2S"C UV - 2.21 10% (vlv) l,4-dloxane, 3S"C UV -2.47 10% (v/v) l ,4-dloxan., 45"C UV - 2.47
5. CH.~~-oH 10% dlell."., H.'O,2S"C UV -1." 24 (Q}-CH-CH -<:-0 10% dk»Iane, HCIO, 25"C UV - 2.1 1
6. uv - 2.37 42
40b observed . The shape of the absorption band is also been observed in hydrochloric and perchloric regular with \"ox 236 nm and absorbance AB =0.368 acids. The position of Amax of 4-substitutecl for P-CH3BHA. The absorbance is higher for BH+ benzohydroxamic acid remained constant in dil ut e
(An"" = 268 nm and ABIl+ = 0.689). This means that the solutions of mineral acids but in more concentrated
ba:'le under study undergoes protonation in highly acid acid there was a steady bathochromic shift in Il l11 ." . solutions. Such an explanation is rational and obvious. except for 4-NOr BHA. The wavelength of li ght The UV absorption spectra of benzohydroxamic absorption conesponds to the energy di rfercnce~ acid and some of its derivatives at different acidities between normal ground state and excited state.
(H}SO." 0.0 to 16.2 M; HCl, 0.0 to 9.3 M and HCI04 Incorporation of different substituents results in a 0.0 to 8.0 M) are illustrated in Figs 4-6. As the change in ground state and excited state energie'. substrate become protonated there was a definite shift which are reflected as shift in ultraviolet spectra (In
in spectrum. aq ueous solution Arnax of BHA = 228 nm. A, 11:1\ or In th e case of BHA in sulphuric acid at low electron donating 4-0CH3-BHA = 252 11111 and -+ acidities a peak is present at 228 nm, which shifts to CH1-BHA = 236 nm; Amo, of electron withdrawing -+ 236 nm at intermediate acidities and continues F-BHA = 228 nm, 4-CI-BHA = 236 nm, 4-Br-BHA = shifting to th e 240-250 nm region in strongly acidic 240 11m and 4-NOr BHA = 268 nm). The "hifts are soluti ons. Similar shifting and medium effects have generally m order of mesomenc effects of GHOSH: ACID-BASE EQUILIBRI A OF HYDROXAMIC ACIDS 26b7
2.2 HlSO, 2.0 ~H 1.0 I 1.11 CH~-O-C-O 1.11 0.' 1.4 •U •U C 1.2 C 0.11 ~ 1.0 ~ 0 0 .! 0.8 0.4 0( .a• 0.8 C 0.4 0.2 0.2
0.0 0.0 HIO 200 210 220 230 240 250 2110 270
180 200 220 240 2110 210 320 Wavelength (nm) Wavelength (nm) Fig. 4-UV absorption curves of BHA in aqueous perchloric ,Kid solutions: 1. (0.0 M), 2. (2.00 M), 3. (4.0 M ). 4. (5.0 M). 5. (o.n Fi g. I- UV spectra of molecular species in water (B) and M). 6. (7 .0 M) 7, (8.0 M). pro tonated species (DH+ 16.2 mol dm'3 H2S04) of 4-CHr BHA.
1.4 H,SO, 1.2 H-N-OH 1.2 I CH.O~C-O 1.0 1.0 u u 0.' " " ~ 0.11 0.11 I ~ 0.11 0 .a ! C C 0.4 0.4
0.2 0.2
0.0 0.0 u...--:":-:-...... ~I-...... _I-...:::;::!!!-.e:=...... I...- ...... -.l 240 2110 2110 300 320 Wavelength (nm) 180 200 240 NO 110 Wavelength (nm) Fi g, S- UV absorption curves of4-CH)O-BHA in aqucou" sulphuri c acid solutions: I. (0.0 M). 2. ( 1.80 M). 3. (3.00 At) . ... Fig. 2-UV spectra of molecular species in watcr (B) and (5.40 M), 5. (7.20 M), 6. (8. 10 M), 7. (9.00 M ). 8. ( 10.8 !Ii). () (12 .6 M). protonated species (BH+ 14.4 mol dm,3 H2S04) of 4-F-BHA.
0.' H,SO, 0:11 H,IO, 0.' K-N-OH 0.' H-N-OH I I 0.7 Br--C-O 0.7 8r-O-<:-o o.e
" 0.11 u 0.5 U "c ~ 0.5 .e 0.4 o 0.4 .a~ 0.3 C ! 0.3 0.2 0.2 0.1
0.1 0.0
0.0 8----- 1110 200 220 240 320 Wavelength (nm) 110 200 220 240 210 210 uo Wavelength (nm) Fi g. 6- UV absorpti on curves of 4-Br-BHA in aqueolls , ulphuri l' ac id solutions: I. (0.0 M). 2. (1.80 M). 3. (3.60 M). 4. (5 ... 0 Ml. ~ . Fi g. 3-UV spectra of molecular species in water (B) and (7 .20 M). 6. (8. 10 M), 7. (9.00 M). 8. ( 10.8 M ). 9. ( 12.0/11 ). Ill. prolonaled species (BH+ 14.4 mol dm,3 H2S04) of 4-Br-BHA . ( 12.6 M ), II. (14.4 M). 2688 INDIAN J CHEM, SEC A, NOVEMBER 2003
substituents. Substituents like f1uoro, chloro, bromo Tabl e 2-Protonation constants pKBH+ of 4-substillited and nitro have a relatively small effect on the benzohydroxamic ac id s determined by di fferent meth ods III benzenoid B band but polar substituents such as aqueous H2SO4
OCH, and CH3 groups alter the benzene spectrum 4-XC6H4CONHOH pKIlH+ considerably. HAFM BOM EAM Mer The UV absorption spectra of all the hydroxamic acids at different acid concentrations undergo changes OCH 3 - 1. 88 - 1.62 - 1.91 - 1.7::' not only because of the protonation process, but also CH3 - 1.95 - 1.90 - 1.93 - I.,) 1 H -2.06 -1.86 - 2.0 1 - 1. 71) as a consequence of the medium effect. Therefore, no F -2. 10 -1.89 - 2.07 - I .:oN good isosbestic points have been seen and the shifting CI -2.17 -1.92 -2. 16 - 1.9 1 with respect to acid concentration (aqueous and Bf -2.25 -2.01 - 2. 19 - ::' . 10 hi gher) is shown in Figs 4-6. In lower acid N02 -2.40 -2.1 9 - 2.23 - 2. 1:; concentration, the concentration of conjugate acid is small. The small bathochromic shift in this region can The influence of solvation may be assessed on a be ex plained as due to increasing interaction between semi-quantitative basis from the m* values. To a the carbonyl group and the solvent. In intermediate reasonable approximation m* -- ( 1-<1» and the two acidities, the concentration of unprotonated species is are inconvertible, Ho + 10gCH+ -- -X, these slope small. Again there is small bathochromic shift with parameters are of paramount importance in defining in crease in acid concentration as a result of changing the protonation behaviour of weak bases. The need for interaction between cation and solvent. The gradual two parameters to describe protonation equilibria Illay shift in Am" results in the convergence of the be considered as reflecting stabilizati on of th e absorbance curves into an apparent isosbestic point. protonated base both by internal delocalization of its Inspite of these problems, all the methods give rather cationic charge and by solvation through hydrogen similar results. By analogy with the larger spectral bonding. The m* values 0 f ami' d es 53 are In. t he range 58 'b h' changes found for benzamide , we ascn e t IS 0.5-0.6. These values, which are low if compared to s p ec~'a l change to the tautomeric change of the cation those of nitrogen bases such as primary anilines (m '" from N-protonated to O-protonated. At higher values of 1.00), are believed to be primarily evid ence concentration of acids, the activity of water falls of strong hydrogen-bonding of the O-protonated BH + which is essential for the stablization of N-protonated with water. cations. Similar changes have been reported in the The pKBH+ has been determined from spectroscopic . 'd59 UV spectra of N-methyl-p-tolylhydroxamlc aCI data by different methods (Eqs 6- 1 I). Some between 60 and 96% sulphuric acid, and have also representative pKBH + values are given in Table-2 and been ascribed to the tautomeric change in the their graphs are shown in Fig. 7. All the linear structure of the cation from N-protonated in aqueous regression plots are good but showed a scatteri ng of acid to O-protonated in concentrated acid. However, points, which could only be resolved into straight thi s controversial question of site of protonation and lines by omitting certain points. The condition of medium effect for hydroxamic acid have not been HAFM method is that the slope of 10gCst·tlCs against answered so far. Hx must be unity. All the seven compounds As solutions of hydroxamic acid cation are made investigated closely follow th e HA function. The still more acidic, the UV spectra continue to undergo slopes (m) of log I vs. HA plots bein g 0.91±1 , onl y an changes. Liler58 interpreted this type of changes as a appropriate acidity function will give a slope of unity. medi~m effect on the equilibrium with the N Unfortunately, no acidity function has been developed protonated forms, which will be strongly hydrogen for hydroxamic acid. Some reconstituted spectra of bonded to water, being displaced by the O-protonated hydroxamic acids after applying characteristic vector forms in the concentrated acid. The O-protonated ions analysis are shown in Figs (8-13). possess considerable charge delocalization, and it is More recently, Garcia et al. 23 investi gated th e argued that they wi II be less affected by the removal mechanism of protonation of acetohydroxamic acid of water from the medium than the N-protonated (CH CO.NHOH). Experimentally, the UV spectral forill. Accordingly, the equilibrium will be tipped 3 curves were recorded at di fferent temperatures. at from the N- to the O-protonated form as the constant dioxanelwater concentration. and at ve ry concentration of acid becomes high. high concentrations of strong mineral acids. The GHOSH: ACID-BASE EQUILIBRIA OF HYDROXAMIC ACIDS
1.2 Ti""0:::::111..... :O:::=:===;------, (HAFM) log I ;: -m HA+pK BH+ (BOM) 0.7 n ; 6 .0.3 log I • log C , = pK + ; ·2.25 H • BH (1 . ~)( H o +Iog Cw )+ pKew Q.2 m :~~~~ n ; 6 ; ·2.01 pK BH + ; 0.278 0.3 !c -C.S ~ ; 0.989 • () ~------~ 0> 0> .2 0.8 Sf , 1.3 • 1.3 0> .2 ,'.S • ' .8
2.3 -u~ ___+_----~--~--~----~--~ • 4 5 • 5
U Q.2 (EAM) (MCP) log 1·log C + ; m' X+pK + H eH log 1·log Cw ; .neMe + pK , n ; 6 BH n ;6 .0.3 pK + ; ·2.19 • ~ • BH pKBW ; ·2.10 m' ; O. 274 + r ; 0.994 I + ~ B : ~~~24 () -C.X I () -C.S 0> ,o 0> ,0 0> ' 1.) - ,1.l .2 0> .2 • " .X ',1.8
.2.1 _---+---t---_+_----I--_-~ .2.1 6 0 6 10 X Me
Fi g. 7-Different methods for the determination of pKBH+ of 4-Br-BHA in H2S04 before applying CV A. spectral curves o f Fig. 14 show no clear isosbestic studies on the acid-base behaviour of hydrox B !~ HCI04 10 10 U •U U• 1..5 c: c: tV tV D. D. L.. L.. 0 I.• 0 I.' D.• D.• c( c( t..I t..I ... o. A » 10 •U u c: tV D.... 0 1.0 D.• c( 0..5 "' -F~----~-----,------r---=--' Fig. 8-Experimental absorbances of BHA as a fUllcti on of mineral acids. The PBHA has one acidic proton, BHA has two The UV spectra of the investigated hycirox:llllic and SHA has three acidic protons and they are prone acids are di storted and exhibit shi :ts in w 3.0 -.------, 0.7 r--.------;:=====::::;-, H,SO, 0.' H--N-OH I 0.' sr-Q-C-O 0 .• 2.0 0.3 •u C l1li 0.2 Sl.... o 0.1 ~ 1. 0 0( o -0.1 ~--+---+__-_... __ -_--_--_--__I 1110 210 230 250 270 2110 310 330 Wavelength (nm) 00 --t-:1-- -...,----....----.,---.::=!!!II 21111 l Z=, 2511 27=, 31H) Fig. I O- Reconsti tuted spectra of 4-B r-benzohydroxa llli c ;lCid ill H ~ SO~ at differe nt acidities. 2.5 -.------, 1.a ,------;:=:0::=====9 1.3 2.0 1.1 u 0.8 • 1.5 "c u c 0.7 €o .0 "' .! o.a (; 1. 0 oC lit .0 0.3 0( 0.1 0.5 0.1 L---+--__ --2::::::::;=:::J 230 250 270 2110 310 0.0 --f.L..:..---.----,,..-----r-= Wavelength (nm) ~ oo 22 .5 2511 2i ~ 31H) Fi g. II - ReconstilLll ed spectra of 4-0CH -benzohyd ro x; lIlli c ;Il'id 3.0 -,------. 1 He) in H2S04 at different acid iti es. 1.2.______--;:::======::::;, H,SO. H-N-OH I 2.0 CH, -0 0.' U "C cg u " 0.8 .a.... c o .!.. WI o 0 .• .0 I/J 0( 1.0 .Q oC 0.2 o -0.2 "---+---+__---<__ -_--_--_---l 0.0 -+==-----.----~---__r-..: 180 210 230 250 270 2110 310 330 225 2511 275 JIH) Wavelength (nm) Fig. l}- Reconstitu ted UV spectra of BHA in different acidities of Fig. 12-Reconstituted spectra of 4-CHJ -benzohydroxa l11 ic acid mineral acids resulting from CVA. in H2S04 at different acid iti es. 2692 INDIAN J CHEM, SEC A, NOV EMBER 2003 0.5 -"T"------.------, 1.0 -,------ 0." 0 .3 ~ O . ~ 0.2 0 .1 0.0 200 22! 2SO 27!1 300 325 200 225 250 ). A. C.S 1.0 -r--- -- 0." 0 .3 0.5 ui ~ 0.2 0 .1 0.0 ~:..L-__ ----=;~~----__l 0.0 200 225 250 200 22S 2!10 275 300 316 A. ). Fig. 13- RcconstilU ted set of spcctral curves after vector analysis Fig. l4- ReconstilUted set of spectral curves aft er vector :l11 ;ti ) , i, trcatmen!. (rcatmen!. K12 The difference in the values of thermodynami c H,O + X + BH' ~ B + H' + X + H,O parameters can be explained on the basis of two different effects: the formation of ion pairs between K" II X (C I0 - , HS0 - and CI -) and the BW cations, and K,. II 4 4 thc stabili zati on of BH ~ due to the positive charge BH'/ H,o/ X ~ [BH', X l + H,O deloca li zation. The first effect explains the Scheme 4 dependence of the extent of protonation on the nature of the mineral acid; the relatively high concentrations th e [BH +, X] contact ion-pair (Scheme 4). Moreovcr. of X present in the solvent medium favour th e comparison of log aH20 values at the samc molar for mation of ion pairs with BH+. Water is also U4 concentration gives aH 20 (HCI ) > aH 20 (H :S0 ) > invo lved in th e ion-pairing process to an extent th at 4 aH20 (HCI0 ), and ac10 - i.e., ions ex hibit ;\ depends on th e medium permitivity, so that an 4 act-> 4 cr eq uilibrium is established between "solvent stronger tendency to form ion-pairs compared t\) C10 4 separated" and "contact" ion-pairstiJ. Therefore, the ions. The result s indicate K13H + (HCI ) < KBII+ (HcS04 ) protonated base is di stributed among the BH+ free < KBH + (HCI04 ) . Changing the nature of th e mine r;t\ forlll. th e BH +/H20/ X solvent-separated ion-pair, and acids brings about changes in th e p h ys lc o-c h L~lf1 i ca l GHOSH: ACID-BASE EQUILIBRIA OF HYDROX AMIC AC IDS 26<)~ pro perti es of the medium, which may affect the KI3 NMR method value as well. Concerning the second effect, i.e. , the None of the techniques employed so far can expl ain stabi li zati on of BH+ by charge delocalization, it the ioni zation (protonati on and deproto nati o n) si te should be noted that the structure of the o.-proto nated unambiguously. In fact, UV spectra suffer from th e hydroxami c acid involves a large delocali zation of the combined influence of strong medium effects and side positive charge, as a consequence, the number and reacti ons. With the NMR method, th e observed type of Nand C substituents are expected to play an quantity is a chemi cal shi ft, whi ch must he important role in the stability and solvation of determined relati ve to an internal standard. In addi tion BH+,w hi ch are refl ected in the p KBH+ and m* values, to the usual requirements, an appropriate chemi cal respecti veli s. The change of the p KBH+ values was in shift standard for this purpose should also mai nt ai n its reasonabl y good agreement with the sequence of the che mi cal identity within the entire acidity range. and catalyti c effi ciency of the mineral acids used, HCI > respond to solvent changes in the same mann er as thc free and protonated bases do. . H 2 So.~ > HC lo.4 S9 Simil arl y, Ghosh et al.44 studied protonation Lobo et ai. first measured the IH NMR spectrum behaviour of a cyclic hydroxamic acid (N . of RCo.N(o.H). CH3 (R = H, C H3, C H(CHck C (CH3h, C ti H , P-CH3-C6HS) in C OCl". The spect rulll hydroxyphthalimide VII) in H2SO.4 by UV method. s of parent compound in H2SO.4 « 2.5 M ) contained o two methyl and two formyl sig nals as in CDCI , II /c" attributed to VIII and IX. © /N-OH -"""'c oII (VII) The posItI on of Amax di d not shift significantly (300-308 nm) as the substrate became protonated. The pKJ3H+ (-2.30) was calcul ated at 308 nm where (VIII) (IX ) max imum vari ati ons are observed. The reactivity of At hi gher concentrati on of acids (3 M to 13 M) the cycli c compound changes with internal strain (I-strain). This I-strain is made up of various two methyl signals coll apsed to a singlet attri hut ahle components, in cl uding bond angle distorti on and bond to the presence of signi ficant contributio ns of the N oppositi on effects. proto nated form. No splitting of the N- meth yl signal was seen due to rapid exchange with the medium. At FT- /CR-CD stlldy hi gher concentrati on of H2SO.4 (> 13 M) th ere \\'a:- Gas-phase acidities of acetohydroxamic acid two sets of signals, for N-methyl and o ne formyl (AHA) and its o.-methyl and N-methyl derivati ves protons, an observati on that is rati ona li zed o n th e have been measured by Fourier transform io n basis of o.-protonati o n leading to species X and XI. tis cyclotron resonance . The first conclusion emerges that hydroxamic "acids" are also relati vely strong bases. I n the scale of organic bases (- LOO k cal) they fa ll in the vicini ty of strong bases (at 35% of the whole range). Compared with amides hydroxamic ac ids are stronger acids (by 15k cal) and weaker bases (by 7 k cal). In all the cases the reason is (X ) (X I) evidently the e lectron-attracting inductive effect of the oxygen atom. In the case of acidity of hydroxami c 1t is known that above 60% H 2 So.~ th e acti vi ty (11' acids th is effect is due to a change in the mesomeric water decreases very sharply and this mi ght render th L' stru cture of th e ani o n; the substi tution by an oxygen N-protonated cation more unstabl e th an the 0. induces an electron attraction from N and renders the proto nated o ne. T he sigmoid curve obtained by mesomerism within the o.=C-N unit better balanced. plott ing the N-meth yl shifts (re lati ve to The two effects seem to strengthen each other in an trimeth ylammo nium) against the ami de acidity efficient way. fu ncti on, HA , showed th e absence of any appreciable 2694 INDIAN J CHEM, SEC A, NOVEMBER 2003 1 protonation in thi s region of acidity. The pKBH+ values - °F======::-:-:-::-===::::::;:=:;;:;;;;:;;:;::= for all compounds studied by this NMR method fell I -12 within th e range of - to -3. o ~ The traditional ex perimental techniques for )( N quantitatively studying acid-base equilibria often fail ::. -14 to prov ide an answer concerning th e stru cture of the t;: - 21 :i: ion bein g form ed. Bagno et al. developed a new When H2S04 is added, a downfield shift of the water equivalent as to their solvating ability etc .. th ey are protons (which are now exchanging with the acid) the closest in practical terms, and it is noteworthy th at have been observed. The methyl singlet is observed at opposite types of behaviour were detected in so 1.568. Normally these singals undergo a downfield similar media. The different sites of ioni zati on sh ift (higher 8) with increasing acidity. But opposite (protonation and deprotonation) in th e two cases can trends have been observed here. Hence pKBH+ could also be rationalized on the basis of O-ioni za ti on of not be calculated. AHA giving rise to an oxyanion with a loca li zed charge (which is well stabilized in water). wh erea s for BHA the aromatic nng providing an ex tr ~ 1 Deprotonation Behaviour stabilization of the nitranion, which apparent" There has been considerable debate as to whether compensates for the loss of solvation. hyd roxamic acids are nitrogen or oxygen acids. The structure of the anion had been a matter of Conclusion 67 controversy until XPS studies, in the solid state , and Acidic and! or basic groups that are present in the 17 68 0 NMR studies, in methanol confirmed the N molecules usually dictate the chemical, physical and deprotonation. Correlations between Hand S of biological properties of hydroxami c acids. The 27 69 ion ization and some early studies led to the basicity and solvation are intrinsicall y import ant opposite conclusion (O-ionization). The procedure is parameters for understanding the behaviour 01' organi c complicated also because these studies differ with and biochemical systems. Whi Ie tremendous progress respect to the structure of the acids employed, the has been made over the past 60 years in understanding experimental conditions, and especially the solvent. and applying the chemistry of hydroxamic acids. it Bordwell el (/1. 70 performed a comparison of the occupies relati vely little space in textbooks. The ac iditi es of several types of oxygen, nitrogen, and present review gives the ex perimental spectruscopic carbon acids, including hydroxamic acids in OMSO investigation of acid-base equilibria of hydroxamic solution, They found that for both aceto- and benzo acids in solution. The methods described here ha ve hydroxamic acids, N-alkylation decreased the acidity been extensively treated theuretically and pract icall y. more than acid O-alkylation, indicating that the Both techniques (UV and NMR) normally employed parents are NH rather than OH acids in OMSO. They for the study of protonated eq uilibria are affected tl) observed that N-alkylhydroxamic acids exhibit strong some extent by the med ium effects that accompany homo-H-bonding, whereas the parent and 0- acid concentration changes. Characteristic vector alkylhydroxamic acids did not, supporting their analysis is used to separate medium effects from behaviour as NH acids. However, although the protonation equilibria (clear isosbestic point). U\ ' ox idation potentials of the hydroxamate anions in method is simple. economic and yields consistent ami DMSO are consistent with NH ionization, oxidation reliable results. potential measurements in MeOH admit OH ion ization. The change from NH to OH acidity is Acknowledgement probably caused by interactions with the protic The author is deeply indebted to Prof. M.F. Rua ssc. solvent. (University of Paris-7), Prof. Jose M, Leal (University 2696 INDIAN J CHEM, SEC A, NOVEMBER 2003 of Burgos), Dr Alessandro Bagno (University of 24 Garcia B. fbeas S, Hoy uelos F J & Leal J M. .I or,!.: Chelli. h() Padova) and Dr Robin A. Cox (University of Toronto) (200 I ) 7986. 25 Brink C P & Crumbli ss A L. } org Cilelli. 7 (19X2 ) I 17 1. for stimulating correspondence, guidance and kind 26 Brink C P. Fi sh L L & Crumbli ss A L. .I org Chell I . .'if) cooperation. The author is also thankful to Prof. G L (1985) 2277. Mundhara, Head, Chemistry, Pt. Ravishankar Shukla 27 Monzy k B & Crumbli ss A L,} org Che/li . 45 ( I L)X(), '+6 711. Uni versity. 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