Nucleophilicity, Basicity, and the BrÃ¸nsted Equation - American

9 Nucleophilicity, Basicity, and the Brønsted Equation

Downloaded by UNIV OF PITTSBURGH on May 4, 2015 | http://pubs.acs.org Publication Date: July 1, 1987 | doi: 10.1021/ba-1987-0215.ch009

Frederick G. Bordwell, Thomas A. Cripe, and David L. Hughes Department of Chemistry, Northwestern University, Evanston, IL 60201

Application of the Brønsted relationship reveals that, when measurements of rate and equilibrium constants are made in dimethyl sulfoxide solution using families of bases wherein donor atom and steric effects are kept constant, nucleophilicities depend on only two factors, (a) the relative basicities of Nu- and (b) the sensitivities of the rates to changes in these basicities (the Brønsted β). All combinations of nucleophiles and electrophiles appear to fit this pattern. Points for para π-acceptor substituents deviate from the Brønsted lines because of enhanced solvation effects that introduce a kinetic barrier that is not modeled properly by the equilibrium acidities. The carbon basicities of carbanion, nitranion, oxanion, and thianion families were calculated in the gas phase and shown to correlate linearly with their experimental gas-phase hydrogen basicities. Evidence is presented for a rough, general linear correlation between log k and the oxidation potential of anions in nucleophile-electrophile combinations.

TTHE

T E R M N U C L E O P H I L I C I T Y refers to the relative rate of reaction of an electron donor with a given electrophile, as distinct from basicity, which refers to its relative affinity for a proton in an acid-base equilibrium. A quantitative relationship between rate and equilibrium constants was dis covered by Br0nsted and Pedersen (i) in 1924. These authors found that the rate constants for the catalytic decomposition of nitramide by a family of bases, such as carboxylate ions ( G C H C 0 ~ ) , could be linearly correlated with the acidities of their conjugate acids, p K . This observation led to the discovery of general base catalysis and the first linear free-energy relation ship, which later became known as the Br0nsted equation: 2

2

H B

0065-2393/87/0215-0137$06.00/0 © 1987 American Chemical Society In Nucleophilicity; Harris, J., et al.; Advances in Chemistry; American Chemical Society: Washington, DC, 1987.

138

NUCLEOPHILICITY

log V =

βΡ*ΗΒ + C

(1)

Hammett (2) recognized the general nature of the Br0nsted relationship in 1935 and showed that this relationship could be applied to several other kinds of reactions, including a methyl-transfer ( C H ) reaction between A r N ( C H ) and C H I , as well as to proton-transfer ( H ) reactions. (He pointed out that the C H equilibrium constant would no doubt provide a better model than the H equilibrium for the rate constants for C H , but that such equilibrium constants were not available.) Until recently, however, the Br0nsted equation has been used primarily to correlate rate-equilibrium data for H reactions in aqueous mediae Now that the mechanism of the decomposition of nitramide is understood to involve a base-promoted deprotonation of a tautomer of nitramide, accomplished by elimination of hydroxide ion (equation 2) (3, 4), it is clear that Br0nsted and Pedersen were measuring the relative nucleophilicities of various families of bases toward hydrogen with reference to the relative acidities of their conjugate acids. The Br0nsted relationship can be cast in the form of the Hammett equation to bring out this feature (equation 3). Therefore, the Hammett equation is really a special case of the Br0nsted relationship +

3 T

+

3

2

3

T

+

3 T

+

+

T

3 T

+


T

+ /

o-

B- + H N = N (

• B-H + N = N - 0 - + HO"

(2)

log (k/k ) = β log (K/Ko)

(3)

0

In 1953 Swain and Scott (5) assumed that nucleophilicity (n) in S 2 and related reactions was an inherent property, which could be defined by equation 4, where s is the sensitivity of the rate constants to variations in (or of) the electrophile. S 2 reactions of nucleophiles with C H B r in water were used as a standard (s = 1.0). N

N

3

log (k/k ) = sn

(4)

0

It was soon realized, however, that this simple definition of nu cleophilicity would not suffice. Shortly afterward, Edwards (6) attempted to define relative nucleophilicities in terms of two parameters, H (basicity) and E (oxidation potential), using the (variable) coefficients α and β to relate these properties to changes in the electrophile (equation 5). Later, Edwards and Pearson substituted a polarizibility parameter, P, for E . In essence, equation 5 is a Br0nsted equation with a second parameter added. o x

o x

1

For other applications of the Br0nsted equation to S 2 reactions, see Smith, G . F. /. Chem. Soc. 1943, 521-523. H u d s o n , R. F ; Klopman, G . / . Chem. Soc. 1962, 1062-1067. Hudson, R. F Chemical Reactivity and Reaction Paths; Klopman, G . , E d . ; Wiley-Interscience: New York, 1974; Chapter 5. For applications of the Br0nsted equation to acyl transfer and other reactions, see H a m m e t t (2) and Jencks, W. P. Catalysis in Chemistry and Enzvmolo&u; M c G r a w - H i l l : New York, 1969. N

In Nucleophilicity; Harris, J., et al.; Advances in Chemistry; American Chemical Society: Washington, DC, 1987.

9. BORDWELL ET AL.

Nucleophilicity, Basicity, and the Br0nsted Equation 139 log (k/ko) = aH + β £

(5)

ο χ

The demonstration by Parker (8) that nucleophilicities in S 2 reactions could be enhanced by as much as 10 by changing from a hydroxylic solvent [ H 0 , C H O H , H O C ( 0 ) C H , etc.] to a non-H-bond donor "aprotic" solvent [(CH ) SO, C H C N , H C O N ( C H ) , etc.] called attention once again to the dominant role that solvation plays in rates of nucleophile-electrophile com binations in solution (9, 10). Because stabilization by H-bond-donor solvents varies greatly with nucleophile size, charge, extent of electron pair delocalization, and the nature of the donor atom, the solvation parameter alone makes it difficult, if not impossible, to design an equation capable of quan titative correlation of rate-equilibrium data in such media. In contrast, solvation, as well as other factors controlling nucleophilicity, can be held relatively constant in dimethyl sulfoxide and like non-H-bond donor sol vents. This result can lead to a simple relationship between nucleophilicity and basicity. N

8

2

3

3

2


3

2

3

3

2

Precision and Scope of the Br0nsted Equation. The key to eliminating factors, other than basicity, that dictate nucleophilicity lies in the use of families of nucleophiles in dipolar non-H-bond donor solvents. By a family of nucleophiles, we generally mean a family of anionic bases, A ~ , wherein the basicity can be changed by remote substitution. For example, the basicity of fluorenide carbanions can be changed by 10 or more p K units in (CH ) SO solution by introducing remote substitutents (11). [By contrast, with H bond-donor solvents, such as water, the solvent leveling effects usually restrict basicity changes in a family to a practical limit (1-2 p K units) .] At the same time, the nucleophile donor atom (carbon) is kept constant, and steric and solvation effects are kept nearly constant. Br0nsted plots of the rate constants (log k) for fluorenide ions ( A ) reacting with electrophiles plotted against equilibrium acidities of their conjugate acids (pK ), both in ( C H ) S O solution, exhibit excellent linearity for all electrophiles studied to date. Table I summarizes the results of Br0nsted correlations with families of fluorenide ions and related carbanions (C~), as well as other families of anions (nitranions, N ~ , oxanions, O " , and thianions, S"). The reaction types include S 2 (12-14), S 2 ' (A. H . Clemens and J.-P. Cheng, unpublished results), E 2 (15), H + (16), S A r (12), and e ~ (17, 18). For all 20 of the combinations of nucleophiles and electrophiles shown in Table I, and others that we have studied, the relative rate constants depend on only two factors: (a) the basicity of the anion as defined by the acidity of its conjugate acid, p K , and (b) the sensitivity of the rate constant 3

2

3

-

HA

3

2

N

N

T

N

T

H A

2

These solvents all contain protons that react with strong bases. The term "aprotic" is a misnomer that should be abandoned. 3

For reasons elaborated in the section on solvation effects on nucleophilicity, use of substituents of the type p - N 0 , p - C N , p - S 0 R , and p - C O R to extend the lines in most Br0nsted plots is impractical. 2

2


140

NUCLEOPHILICITY

Table I. Br0nsted Correlations for Reactions of Anion Families with Electrophiles in (CH ) SO Solution 3 2

Reaction

Electrophile

S2 S 2' E2 H+ S Ar

RX 0 ~ > N ~ > C ~ when attack is on hydrogen in an E 2 reaction (15), and a similar order, S", O " > N ~ > C ~ , was observed for attack on hydrogen in the base-catalyzed isomerization of butenenitrile (16). Why does the Br0nsted relationship hold for these diverse reactions and give promise of holding for all combinations of nucleophiles and electrophiles (Table I)? In other words, why does the thermodynamics of proton transfer serve as a model, at least as a first approximation, not only for rates of reactions involving proton transfers ( H ) but also for rates of reactions involving alkyl transfers (R ), bromine atom transfers (Br ), e ~, and other transfers? A partial answer to this question comes from recent results of Arnett and co-workers (28-30), who found in carbanion-carbocation combi nation equilibria that the hydrogen basicities for the carbanions correlated linearly with their carbon basicities. Also, in our laboratory, Cripe (24) observed that the intrinsic, gas-phase, hydrogen basicities for several anion families correlate linearly with their carbon basicities. Furthermore, there is good reason to believe that these relative intrinsic basicities for anion families observed in the gas phase carry over to the solution phase (24). These conclusions were arrived at by using a modification of a method developed by Hine and Weimar (31) to calculate gas-phase (intrinsic) carbon basicities.


3

2

3

+

T

+

+

T

T

T

4

Basicities of Anions toward Hydrogen and toward Carbon. Hine and Weimar used the position of the equilibrium in equation 7 to define carbon basicity, relative to hydrogen basicity. β CH OH + HA
S" >, Ν " , I", and B r " > C l " > O " and F " with the total range being almost 20 orders of magnitude. This represents a difference of 27 kcal/mol in the free energies of reactions of these anions with a common electrophile ( C H O H ) , even after the differences associated with p K are factored out. The difference in oxanion and carbanion basicities toward carbon at the same hydrogen basicity is about 16 kcal/mol (for example, compare C H 0 and C N " in Table II). We might wonder whether or not solvation effects are likely to mask this intrinsically greater carbon basicity of carbanions than oxanions. Comparison of equations 8 and 10 shows that the effect on the equilibria of going from the gas phase into the solution phase will depend primarily on the relative solvation energies of the two different anions, because (a) H O is common to the two equations and (b) solvation effects on the neutrals are expected to be small and to differ but little. Fortuitously, the free energies of aqueous solvation of C H 0 and C N ions are almost identical ( — 74 and — 73 kcal/mol, respectively) (33), which means that the ~16 kcal/mol greater basicity of C N ion toward carbon will be retained in solution. Other solvation differences will be larger, but the relative intrinsic basicities of anions toward hydrogen and carbon revealed by the gas-phase data will be retained in solution because solvation stabilization of the anions will affect equations 8 and 10 to exactly the same degree. A

+

3

H

T

A

-

+

T

-

_

6

+

5

14

+

T

3 T

_

_

3

+

+

3

5

T

-8

3 T

+

+

T

3 T

C H a A

H A

H


A

+

+

T

3 T

+

+

3T

T

3

H A

-

6

-

_

6

-

5

-


5

150

NUCLEOPHILICITY +

Table II. Experimental Free-Energies of Reactions AG°(H ) and Equilibrium Constants (K%_) for the Reaction in Equation 8, and Calculated Equilibrium Constants, K™ , for Reaction in Equation 10 T


e

No.

A"

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

MePhCH CH CHCH

a

b

AG°(//+)

2

2

2

CH =CHHC=CMeC^C0 NCH 0 NCH(Me)CNCNCH MeC(0)CH CHOCH F CHOMeOEtOPhO" HC0 H NMeNH" Me N" HSMeSPhSHFciBrI2

2

2

2

2

2

2

3

2

2

2

b

a

-25.0 11.7 0.2 -12.0 -20.0 16.8 12.1 32.0 32.3 38.2 19.6 22.4 24.4 16.1 (0) 11.4 14.5 39.5 43.4 -12.1 -11.7 -5.2 36.9 31.3 56.0 ~-10 18.3 56.9 68.5 77.8

1.77 2.88 7.15 5.74 3.97 5.46 1.47 4.39 2.65 1.31 4.94 4.46 1.55 1.77

X X X X X X X X X X X X X X

AG°(Mef)

1018 10-9 10-1 10» 1014 1013 10-9

4.43 3.20 8.37 4.35 1.21 7.21 1.91 6.72 2.22 9.75 7.71 9.14 1.72 1.48

10-24 10-24

10-28 10-15 10-17 10-1» 10-12 (1) 4.77 X 10-9 2.61 X l O - i i 1.47 X 10-29 2.09 X 10-32 6.79 X 108 3.47 X 108 6.25 X 103 1.16 X 10-27 1.42 X 10-23 1.33 X 10-41 ~2 X 107 4.39 X 10-14 2.93·X 10-42 1.00 X 10-50 1.63 X 10-57

x 1026 x 101 x 109 x 1021 x 1027 x 101 x 104 x 10-13 x lO-ii x 10-14 x 10-2 x 10-7 x 10-7 x 103 (1) x 10-6 x 10-26 x 10-26 x 10-30 x ion x 1013 x 109 x 10-20 x 10-15 x 10-33 x 1027 x 10-13 x 10-37

3.86 1.56 2.73 2.99 6.93 7.08 8.81 3.72 6.72 4.50 -5.2 6.50 3.78 1.36 X 10-43 2.17 x 10-48

-36.5 -2.1 -13.6 -29.6 -37.1 -2.5 -5.9 16.7 14.6 17.8 1.5 8.3 9.3 -4.3 (0) 7.4 10.7 35.0 40.4 -16.2 -19.0 -13.6 26.6 19.4 44.3 -38.0 16.7 49.9 58.7 65.3

Values taken from the following two references: Bartmess, J. E . ; Mclver, R. T., Jr. "Gas Phase Ion Chemistry Volume 2"; Bowers, M . T . , E d . ; Academic Press: New York, 1979, Chapter 11. Taft, R. W. Prog. Phys. Org. Chem. 1983, 14, 305-346.

K Me = K H X £ MeA A

A

HA

In the previous section, anions of the same basicity in ( C H ) S O solution showed a different order of reactivities when the anions attacked hydrogen than when they attacked carbon. The major difference was an enhanced nucleophilicity for carbanions when forming a bond to carbon. This differ ence may have a thermodynamic origin because, in the gas phase, carbanions have an enhanced carbon basicity, relative to nitranions and oxanions, at the same hydrogen basicity. Although the rates (nucleophilicities) of the S 2 reactions for different donor-atom anions in ( C H ) S O were compared at the same hydrogen basicity (PK A)> order is not intrinsic, that is, a nu cleophilic order where the rates have been adjusted for A G ° differences. Indeed, when the n - C H ( C H ) C l S 2 reactivities are adjusted for estimated 3

2

N

3

t

n

2

e

H

3

2

3

N



9. BORDWELL ET AL.

Nucleophilicity, Basicity, and the Br0nsted Equation 151

+

+

Figure 7. Plot of AG°(CH ) values against &G°(H ) values for the 30 anions in Table II (2Λ). 3T

T

differences in gas-phase basicities toward carbon at the same hydrogen basicity, by using the families of lines in Figure 7 as a model, the intrinsic nucleophilicity order toward carbon in ( C H ) S O at constant hydrogen basicity and with AG° = 0 (22, 23), will probably be S" > O " > N " > C " . This order is the same as the intrinsic nucleophilicity order of these donoratom anions toward hydrogen in ( C H ) S O (16). Other manifestations of the high carbon basicity of carbanions, relative to oxygen, include the wellknown tendency for enolate ions to alkylate on carbon and the tendency of carbanions to effect S 2 substitution on cyclohexyl substrates under condi tions where oxanions of the same hydrogen basicity effect E 2 elimination. 3

2

5

3

2

N

5

Because S" and N ~ family lines appear to be nearly colinear and about midway between the O " and C family lines in Figure 7, their relative reactivities in Br0nsted plots will not change on converting the χ axis from hydrogen to carbon basicity. But the O " family is less basic toward carbon and the C ~ family is more basic toward carbon, at the same hydrogen basicity, than the S" and N " families. T h e values obtained by taking the difference in carbon basicities of the O " and C " lines, relative to the N ~ and S~ lines, and multiplying by the β value indicate that the apparent C " rates are about 10 too fast and the apparent O " rates are about 10 too slow when corrected for the thermodynamic driving force. -

2


2

152

NUCLEOPHILICITY

Nucleophilicity, Basicity, and Redox Potentials. Several authors have shown that nucleophilicities in various reactions can be correlated with redox potentials of nucleophiles. Edwards (6) used E values in water in his twoparameter equation (equation 5). Dessy et al. (34) observed a linear correla tion between oxidation potentials of various transition metal nucleophiles, M L " , and the rate constants for reactions with C H I . Recently, Ritchie (33) found a correlation between oxidation potentials of nucleophiles and their rates of combination with pyronin cation. Linear correlations between oxida tion potentials of anions and the p K values of their conjugate acids were also observed by a number of investigators (35-37). In our laboratory, a good linear correlation between oxidation potentials of 2-fluorenide ions and the pK values of their conjugate acids was found and a similar, but poorer, correlation for meta-substituted phenylcyanomethide ions was observed. Points for para donor substituents deviate from the lines in these plots because the p K values fail to take into account the radical-stabilizing abilities of these substituents (18). The radical-stabilizing effects of remote substituents on radicals are relatively small, however (~0.5-3.0 kcal/mol) (18). As a consequence, in view of the general correlations observed between log k and p K for nucleophile-electrophile combinations (Table I), a gen eral, but not precise, relationship between log k and E is expected. o x

3

H A


H A

H A

H A

ox

Acknowledgment We are grateful to the National Science Foundation for support of this research. Literature Cited 1. Brønsted, J. N.; Pedersen, K. J. Z . Phys. Chem., Stoechiom. Verwandschaftsl. 1924, 108, 185-235. 2. Hammett, L. P. Chem. Rev. 1935, 17, 125-136. 3. Pedersen, K. J. J. Phys. Chem. 1934, 38, 581-600. 4. Bell, R. P. The Proton in Chemistry, 2nd ed.; Cornell University: Ithaca, NY, 1973; p 161. 5. Swain, C. G.; Scott, C. B. J. Am. Chem. Soc. 1953, 75, 141-147. 6. Edwards, J. O. J. Am. Chem. Soc. 1954, 76, 1540-1547. 7. Edwards, J. O.; Pearson, R. G. J. Am. Chem. Soc. 1962, 84, 16-24. 8. Parker, A. J. Chem. Rev. 1969, 69, 1-32. 9. Ogg, R. Α.; Polanyi, M . Trans. Faraday Soc. 1935, 31, 604-620. 10. Glew, D. N . ; Moelwyn-Hughes, E. A. Proc. R. Soc. London, Ser. A 1952, 211, 254-265. 11. Bordwell, F. G.; Hughes, D . L . J. Org. Chem. 1980, 45, 3314-3320, 3320-3325. 12. Hughes, D. L. Ph.D. Dissertation, Northwestern University, 1981. 13. Bordwell, F. G.; Hughes, D. L. J. Org. Chem. 1983, 48, 2206-2215. 14. Bordwell, F. G.; Hughes, D. L. J. Org. Chem. 1982, 47, 3224-3232. 15. Bordwell, F. G.; Mrozack, S. B. J. Org. Chem. 1982, 47, 4813-4815.


9. B O R D W E L L ET A L .

16. 17. 18. 19. 20. 21. 22. 23.


24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37.

Nucleophilicity, Basicity, and the Br0nsted Equation 153

Bordwell, F. G.; Hughes, D. L. J. Am. Chem. Soc. 1985, 107, 4737-4744. Bordwell, F. G.; Clemens, A. H. J. Org. Chem. 1985, 50, 1151-1156. Bordwell, F. G.; Bausch, M . J. J. Am. Chem. Soc., 1986, 108, 1979-1988. Bordwell, F. G.; Bartmess, J. E . ; Drucker, G. E.; Margolin, Z.; Matthews, W. S. J. Am. Chem. Soc. 1975, 97, 3226-3227. Taft, R. W. Prog. Phys. Org. Chem. 1983, 14, 246-350. Bordwell, F. G.; Branca, J. C.; Cripe, T. A. Isr. J. Chem. 1985, 26, 357-366. Marcus, R. A. J. Phys. Chem. 1968, 72, 891-899. Hassid, A. I.; Kreevoy, M . M . ; Liang, T.-M. Symp. Faraday Soc. 1975, 10, 69-77. Cripe, T. A. Ph.D. Dissertation, Northwestern University, Sept 1985. Mishima, M.; McIver, R. T., Jr.; Taft, R. W.; Bordwell, F. G.; Olmstead, W. N . J. Am. Chem. Soc. 1984, 106, 2717-2718. Bordwell, F. G.; Hughes, D. L. J. Org. Chem. 1982, 47, 169-170. Bordwell, F. G.; Hughes, D. L. J. Am. Chem. Soc. 1984, 107, 3234-3239. Arnett, E. M . ; Troughton, Ε. B. Tetrahedron Lett. 1983, 24, 3299-3302. Arnett, Ε. M . ; Troughton, E . B.; McPhail, A. T.; Molter, Κ. E. J. Am. Chem. Soc. 1983, 105, 6172-6173. Troughton, Ε. B.; Molter, Κ. E . ; Arnett, Ε. M . J. Am. Chem. Soc. 1984, 106, 6726-6735. Hine, J.; Weimar, R. D., Jr. J. Am. Chem. Soc. 1965, 87, 3387-3396. Bartmess, J. E . ; McIver, R. F., Jr. Gas Phase Ion Chemistry; Bowers, M . T., Ed.; Academic: New York, 1979; Vol. 2, Chapter 11. Ritchie, C. D. J. Am. Chem. Soc. 1983, 105, 7313-7318. Dessy, R. E . ; Pohl, R. L.; King, R. B. J. Am. Chem. Soc. 1966, 88, 5121-5129. Breyer, B. Ber. Dtsch. Chem. Ges. 1938, 71, 163-171. Kern, J. M . ; Sauer, J. D.; Federlin, P. Tetrahedron 1982, 38, 3032-3033. Bank, S.; Schepartz, Α.; Giammateo, P.; Zubieta, J. A. J. Org. Chem. 1983, 48, 3458-3464.

RECEIVED

for

review October 2 1 , 1985.

ACCEPTED

February 10, 1986.


Nucleophilicity, Basicity, and the BrÃ¸nsted Equation - American

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