Solvent Effects in Chemistry E-Book

Erwin Buncel

• Herausgeber: John Wiley & Sons
• Kategorie: Wissenschaft und neue Technologien
• Sprache: Englisch

Seitenzahl: 408

Veröffentlichungsjahr: 2015

Beschreibung

This book introduces the concepts, theory and experimental knowledge concerning solvent effects on the rate and equilibrium of chemical reactions of all kinds. It begins with basic thermodynamics and kinetics, building on this foundation to demonstrate how a more detailed understanding of these effects may be used to aid in determination of reaction mechanisms, and to aid in planning syntheses. Consideration is given to theoretical calculations (quantum chemistry, molecular dynamics, etc.), to statistical methods (chemometrics), and to modern day concerns such as "green" chemistry, where utilization and disposal of chemical waste or by-products in an environmentally safe way is as important as achieving the desired end products by all chemists nowadays.  The treatment progresses from elementary to advanced material in straightforward fashion. The more advanced topics are not developed in an overly rigorous way so that upper-level undergraduates, graduates, and newcomers to the field can grasp the concepts easily.

Leseprobe

CONTENTS

COVER

TITLE PAGE

PREFACE TO THE SECOND EDITION

PREFACE TO THE FIRST EDITION

1 PHYSICOCHEMICAL FOUNDATIONS

1.1 GENERALITIES

1.2 CLASSIFICATION OF SOLVENTS

1.3 SOLVENTS IN THE WORKPLACE AND THE ENVIRONMENT

1.4 SOME ESSENTIAL THERMODYNAMICS AND KINETICS: TENDENCY AND RATE

1.5 EQUILIBRIUM CONSIDERATIONS

1.6 THERMODYNAMIC TRANSFER FUNCTIONS

1.7 KINETIC CONSIDERATIONS: COLLISION THEORY

1.8 TRANSITION-STATE THEORY

1.9 REACTIONS IN SOLUTION

1.10 DIFFUSION-CONTROLLED REACTIONS

1.11 REACTION IN SOLUTION AND THE TRANSITION-STATE THEORY

PROBLEMS

2 UNREACTIVE SOLVENTS

2.1 INTERMOLECULAR POTENTIALS

2.2 ACTIVITY AND EQUILIBRIUM IN NONELECTROLYTE SOLUTIONS

2.3 KINETIC SOLVENT EFFECTS

2.4 SOLVENT POLARITY

2.5 ELECTROSTATIC FORCES

2.6 ELECTROLYTES IN SOLUTION

2.7 SOLVATION

2.8 SINGLE ION SOLVATION

2.9 IONIC ASSOCIATION

2.10 SOLVENT MIXTURES

2.11 SALT EFFECTS

PROBLEMS

3 REACTIVE SOLVENTS

3.1 SPECIFIC SOLUTE/SOLVENT INTERACTIONS

3.2 HYDROGEN BONDING

3.3 ACIDS AND BASES IN SOLVENTS

3.4 BRØNSTED–LOWRY ACIDS AND BASES

3.5 ACIDITY FUNCTIONS

3.6 ACIDS AND BASES IN KINETICS

3.7 LEWIS ACIDS AND BASES

3.8 HARD AND SOFT ACIDS AND BASES (HSAB)

3.9 SCALES OF HARDNESS OR SOFTNESS

3.10 ACIDS AND BASES IN REACTIVE APROTIC SOLVENTS

3.11 EXTREMES OF ACIDITY AND BASICITY

3.12 OXIDATION AND REDUCTION

3.13 ACIDITY/REDOX DIAGRAMS

3.14 UNIFICATION OF ACID–BASE AND REDOX CONCEPTS

PROBLEMS

4 CHEMOMETRICS: SOLVENT EFFECTS AND STATISTICS

4.1 LINEAR FREE ENERGY RELATIONSHIPS

4.2 CORRELATIONS BETWEEN EMPIRICAL PARAMETERS AND OTHER MEASURABLE SOLVENT PROPERTIES

4.3 REPRESENTATION OF CORRELATION DATA ON THE HEMISPHERE

4.4 SOME PARTICULAR CASES

4.5 ACIDITY AND BASICITY PARAMETERS

4.6 BASE SOFTNESS PARAMETERS

4.7 CONCLUSION

5 THEORIES OF SOLVENT EFFECTS

5.1 INTRODUCTION: MODELING

5.2 QUANTUM-MECHANICAL METHODS

5.3 STATISTICAL-MECHANICAL METHODS

5.4 INTEGRAL EQUATION THEORIES

5.5 SOLVATION CALCULATIONS

5.6 SOME RESULTS

PROBLEMS

6 DIPOLAR APROTIC SOLVENTS

6.1 INTRODUCTION

6.2 ACIDITIES IN DMSO AND THE H-SCALE IN DMSO–H

2

O MIXTURES

6.3 USE OF THERMODYNAMIC TRANSFER FUNCTIONS

6.4 CLASSIFICATION OF RATE PROFILE-MEDIUM EFFECT REACTION TYPES

6.5 BIMOLECULAR NUCLEOPHILIC SUBSTITUTION

6.6 PROTON TRANSFER

6.7 D

2

–HO

−

EXCHANGE

PROBLEMS

7 EXAMPLES OF OTHER SOLVENT CLASSES

7.1 INTRODUCTION

7.2 ACIDIC SOLVENTS

7.3 BASIC SOLVENTS

7.4 CHIRAL SOLVENTS

8 NEW SOLVENTS AND GREEN CHEMISTRY

8.1 NEOTERIC SOLVENTS

8.2 SUPERCRITICAL FLUIDS

8.3 IONIC LIQUIDS

8.4 LOW-TRANSITION-TEMPERATURE MIXTURES

8.5 BIO-BASED SOLVENTS

8.6 FLUOROUS SOLVENTS

8.7 SWITCHABLE SOLVENTS

8.8 GREEN SOLVENT CHEMISTRY

9 CONCLUDING OBSERVATIONS

9.1 CHOOSING A SOLVENT

9.2

ENVOI

APPENDIX

ANSWERS

CHAPTER 1

CHAPTER 2

CHAPTER 3

CHAPTER 6

REFERENCES

INDEX

END USER LICENSE AGREEMENT

List of Tables

Chapter 01

Table 1.1 Limiting equivalent conductances of ions in amphiprotic solvents

Table 1.2 Molecular solvents

Table 1.3 Transfer free energies of reactants (

δ

tr

G

R

) and transition states (

δ

tr

G

T

) and solvent effects on reaction rates. Classification of reaction types

Chapter 03

Table 3.1 Autoionization constants of selected solvents

Table 3.2 Coefficients in the equation representing the excess acidity in aqueous solutions of sulfuric, perchloric, and hydrochloric acids as functions of the mole fraction of acid,

X

2

, with the standard error of estimate,

S

Table 3.3 Comparison of

α

/

β

with softness. Edwards’s acid parameters for 17 cations and the tentative values of the proposed softness parameter (see text)

Table 3.4 Trilinear correlation of the Marcus softness parameter

μ

with either the volume refraction

R

v

or molar refraction [

R

]

a

and with Swain’s parameters

A

j

and

B

j

Table 3.5 Kinetic data for the rearrangement of azoxybenzene to 4-hydroxyazobenzene at 25°C

Chapter 04

Table 4.1 Parameters correlated

Table 4.2 Pairwise correlation coefficients between selected solvent parameters

Table 4.3 Rate constants of three reactions of a square-planar Pt(II) complex in eight solvents, with values of seven parameters for each solvent

Table 4.4 Coefficients of correlation between (1) and A

j

(3) and B

j

(4) and with either β

μ

or (2)

Chapter 05

Table 5.1 Calculated (Hartree–Fock, second-order Møller–Plesset) and experimental Internal energy of

gauche–anti

rotation of 1,2-dichloroethane in vacuo and by qO-SCRF in solution

Table 5.2 Enthalpy changes for gas-phase reactions: M

+

(H

2

O)

n

−1

 + H

2

O = M

+

(H

2

O)

n

and X

−

(H

2

O)

n

−1

 + H

2

O = X

−

(H

2

O)

n

Table 5.3 Identity S

N

2 reactions: Comparison of central barrier heights calculated by various methods

Chapter 06

Table 6.1 Physical properties of some dipolar aprotic non-HBD solvents in Order of Increasing dipole moment

Table 6.2 Equilibrium acidities in dimethyl sulfoxide and in water

Table 6.3 Selected

H

−

data for aqueous binary mixtures with several dipolar aprotic solvents, each with 0.011M tetramethylammonium hydroxide

Table 6.4 Free energies of transfer of ions (

δ

Δ

G

tr

) from water to nonaqueous solvents at 25°C (molar scale, in kcal mol

−1

)

Table 6.5 Transfer free energies of reactants () and transition states () and solvent effects on reaction rates. Classification of reaction types

Table 6.6 Relative rates of the S

N

2 anion–molecule reactions A and B and of the S

N

Ar reactions C and D in protic and polar non-HBD solvents at 25°C

Table 6.7 Enthalpies of transfer (kcal mol

−1

) for reactants and for the transition state of the D

2

–HO

−

exchange process in the DMSO–H

2

O system

Chapter 07

Table 7.1 Substances soluble in HF with acid/base properties

Chapter 08

Table 8.1 Cumulative energy demand (CED) for three common solvents, in MJ kg

−1

)

Table 8.2 The Pfizer solvent selection guide for medicinal chemistry

Appendix

Table A.1 Properties of selected solvents (at 25°C except as noted)

Table A.2a Solvent property parameters: symmetric properties

Table A.2b Dual parameters

Table A.3 Values of selected parameters for selected solvents (in order of increasing )

List of Illustrations

Chapter 01

Figure 1.1 The Walden product, Λ

0

η

, for HCl in 1,4-dioxane/water mixtures versus percentage of dioxane at 25°C.

Figure 1.2 Ternary diagram for classification of liquids (schematic; location of points is conjectural); [bmim]PF

6

represents a room-temperature ionic liquid (see Section 8.3).

Figure 1.3 Vapor pressure over binary solutions. Dashed lines: ideal (Raoult’s Law). Solid curves: positive deviations from Raoult’s Law. Note that where

x

2

≪

 1,

P

1

is close to ideal, and vice versa.

Figure 1.4 Successful (a) and unsuccessful (b) transfer of a hydrogen atom from HI to Cl.

Figure 1.5 A portion of a potential-energy surface

E

(

x,y

), showing a saddle point.

Figure 1.6 (a) The reaction pathway of least energy and (b) the profile along the pathway, for the hydrogen atom–molecule exchange reaction (schematic).

Figure 1.7 A possible, more realistic reaction profile for a ligand-exchange reaction, showing reactants (a), precursor (b), and successor (d) complexes, the transition state (‡), the possibility of the formation of a reactive intermediate (c), and products (e).

Figure 1.8 Effect of solvation on activation energy. Potential energy

V

versus the reaction coordinate. Solid curves represent the energy profile in the absence of solvation. (a) Solvation of the reactant (increased activation energy). (b) Solvation of the activated complex (reduced activation energy).

Chapter 02

Figure 2.1 (a) Kinetic solvent effect in Etard’s reaction. Log

10

(

k

1

) versus Hildebrand’s solubility parameter. Solvents are (from Stairs, 1962): (1) 1,2-dichloroethane, (2) 1,1,2,2-tetrachloroethane, (3) trichloromethane, (4) pentachloroethane, (5) 1,1,2-trichloro-

f

-ethane (Freon 113), (6) tetrachloromethane, and (from Cook and Meyer, 1995) (7) cyclohexane. (b) Logarithms of rate constants as in (a) versus Kirkwood’s dielectric function. Numbers refer to solvents as in (a). The slope of the fitted line is 2.9 ± 0.2 (dimensionless).

Figure 2.2 Hildebrand’s solubility parameter

δ

versus Kirkwood’s dielectric function (

ε

r

 − 1)/(2

ε

r

 + 1) for a selection of solvents of low (squares) and high (circles) dipole moment.

Figure 2.3 Cycloaddition of diphenylketene to butyl vinyl ether. Common logarithm of the relative rate constant versus Kirkwood’s dielectric function.

Figure 2.4 The Debye–Hückel limiting law for a 1:1 electrolyte (Eq. 2.21, lower dashed curve), the corrected law (Eq. (2.22), upper dashed curve), Davies’s approximation (Eq. 2.23, black curve) and experimental data for HCl, all in water at 25°C. The activity coefficient is plotted against ionic strength,

μ

 = ∑

z

i

2

c

i

/2.

Figure 2.5 (a) Solvation in flatland. A solvated anion (black circle) and its cybotactic region. The solvent molecules within the innermost circle are virtually fixed in orientation toward the ion. Those within the next circle are less strongly oriented, but more closely packed, while those beyond are undisturbed. (b) Hydrophobic solvation. The innermost solvent molecules form a cage around the nonpolar solute, hydrogen-bonded to each other. As in the ionic case, these are surrounded by a region of disturbed structure, beyond which the solvent is normal.

Figure 2.6 Hydrophobic dimerization: 2 M(solv) → M2(solv). Twelve solvent molecules are shown as being liberated.

Figure 2.7 Schematic cross section of a micelle formed by an anionic detergent. The counterions shown would realistically be part of the ionic atmosphere about the micelle.

Figure 2.8 Acids in organic/aqueous solvents: (1) 2-propanol 5%, (2) methanol 10%, (3) ethanol 10%, (4) 2-propanol 10%, (5) methanol 20%, (6) ethanol 20%, (7) glycerol 50%, (8)1,4-dioxane 15%. Common logarithm of

K

a

(

S

)/

K

a

(

W

) versus the reciprocal of the relative permittivity of the mixed solvent. Acids are indicated by plotting symbols: formic, triangle; acetic, cross; propanoic, circle; butanoic, square; water, filled circle.

Figure 2.9 Bjerrum probability distribution. Probability that a species of given charge (0 or −1) will be found at a distance

r

from a central ion of charge +1. A negative ion inside the distance

q

is considered paired with the central cation. If the distance of closest approach is

a

, the probability that a pair exists is proportional to the shaded area under the upper curve,

a

 < 

r

 < 

q

. If

a

 > 

q

, pairing does not occur.

Figure 2.10 Visible absorption spectra of (triphenylmethyl)lithium: (a) at room temperature in diethyl ether (dash-dot), THF (solid) and 1,2-dimethoxyethane (dashes); (b) temperature dependence of the spectrum of the same substance in diethyl ether.

Figure 2.11 Ionic radius and ion-pair dissociation: p

K

diss

for Cs, Rb, K, Na nitrates in methanol (circles) and K, Na, Li picrates in acetonitrile (squares) versus the reciprocal of the cation crystal radius. The value of p

K

for NaNO

3

in methanol is reported as < 1.

Figure 2.12

E

T

(30) values for water and normal alcohols versus number of carbons. The curve is represented by

E

T

(30) = (

E

p

 + 

nE

0

)/(

n

 + 1), but the data were fitted using the linearized form, Equation 2.29. The polar group contribution

E

p

, = 63.37; the methylene group contribution

E

0

 = 46.27.

Figure 2.13 Kinetic salt effect; slope of the graph of log(

k

/

k

0

) vs.

I

1/2

/(1 + 

I

1/2

), where

I

1/2

is the ionic strength, for 11 reactions between ions of different charge type in aqueous solution at 25°C versus the product of the charges on the reacting ions.

Chapter 03

Figure 3.1 Excess acidity, X, in aqueous perchloric (light curve), sulfuric (black), and hydrochloric (dashed) acid solutions versus mole fraction of the acid,

x

2

.

Figure 3.2 Brønsted plot for the dehydration of acetone hydrate in acetone, catalyzed by acid. The logarithm of the catalytic rate constant

k

a

is plotted versus the negative of the dissociation constant of the acid. Points are shown for 32 carboxylic acids and 15 phenols. The point labeled 47 represents 2,4-dinitrophenol. Acids of other types (e.g., oximes) fell much further off the line (see Section 3.6.3).

Figure 3.3 Brønsted-type plot of log

k

Nu

versus p

K

a

for the reaction of nucleophiles with methyl

p

-nitrophenyl sulfate. The α-nucleophiles are shown as solid circles.

Figure 3.4 Plots showing the effect of solvent composition on the α-effect for the reactions of PNPA with Ox

−

(an alpha nucleophile) versus 4-ClPhO

−

(a normal nucleophile), open circles, left-hand scale, and with IBA

−

versus 4-ClPhO

−

, filled circles, right-hand scale, in DMSO–water mixture at 25.0 ± 0.1°C.

Figure 3.5 Effect of solvent composition on the α-effect for reactions of PNPA at 25.0 ± 0.1°C:

k

(Ox

−

)/

k

(4-ClPhO

−

) in MeCN–H

2

O (

●

);

k

(Ox

−

)/

k

(4-ClPhO

−

) in DMSO–H

2

O (

○

); and

k

(IBA

−

)/

k

(4-ClPhO

−

) in DMSO–H

2

O (

□

).

Figure 3.6 Effect of solvent composition on the α-effect for the reactions of PNPA, PNPDPP, and PNPBS with Ox

−

versus 4-ClPhO

−

in DMSO–water mixture at 25.0 ± 0.1°C.

Figure 3.7 Illustration of the relationship between the transfer free energies of reactant and the transition state and the free energies of activation for a reaction occurring in two solvent systems. The reference solvent, in this case water, is denoted by the zero subscript, and the solvent under investigation (such as DMSO–water mixture) by subscript S. Uppercase Δ refers to change during activation; lowercase

δ

refers to change of solvent.

Figure 3.8 Energy profile illustrating the effect of the zero-point energy of the initial state only, on the activation energy for carbon–deuterium versus carbon–hydrogen bond scission, leading to a primary isotope effect.

Figure 3.9 Linear activated complex. (a) Antisymmetric stretch: the H is being transferred from C to B. (b) Symmetric stretch: the H hardly moves.

Figure 3.10 A general energy profile that illustrates the origin of the kinetic isotope effect in terms of the zero-point energies in both the initial and transition states. The ZPE for the heavy isotope (subscript h) and for the light isotope (subscript l) arise from summation over the various vibrational modes (i).

Figure 3.11 Edwards’s parameters,

β

versus

α

, for 17 cations, with a proposed softness scale superposed. Circles represent hard ions, crosses soft, and squares borderline. Note the apparently anomalous position of Zn

2+

.

Figure 3.12 (a) Drago–Wayland parameters

E

A

and

C

A

for nonionic acids. Hard acids are shown as circles, soft as crosses, and borderline as squares. The bonding atom is indicated by color: black for hydrogen and red for others. Lines are superposed representing values of a proposed softness parameter, defined by . (b) Drago–Wayland parameters for nonionic bases:

E

B

versus

C

B

. Colors indicate bonding atoms: red O, green N, and black others (S, P, C, Cl). Hard bases are shown as circles, soft as crosses, and borderline as squares. The dashed lines are for indicated values of the proposed softness parameter for neutral bases, defined by .

Figure 3.13 Pourbaix diagram for iron species at various reduction potentials

E

and pH in aqueous environment at 25°C. Only the species indicated in the fields were considered. The dashed lines represent the thermodynamic thresholds for the emission of O

2

(upper) and H

2

(lower).

Figure 3.14 Pourbaix diagram for species derived from isobutane (2-methylpropane), represented as

i

C

4

H, in solution in HF. The neutral point in HF is at pH = 6.9, at which

H

0

 ≈ −15. Addition of KF or SbF

5

serves to adjust the pH to more basic or acidic values, within the practical limits, 1–13, approximately a range of

H

0

about −21 to −9 (double arrow).

Chapter 04

Figure 4.1 Scree plot: eigenvalues

λ

r

of the matrix of correlation coefficients of 23 parameters for 28 solvents, in descending order. Four eigenvalues are greater than unity, with a distinct break before the fifth, suggesting that four independent properties of the solvents are significant.

Figure 4.2 Representation of the 23 parameters on the hemisphere. Squares represent positive vector directions, crosses negative, that is, antipodes. The Cartesian axes corresponding to the three principal components are labeled N(North pole),

G

(Greenwich meridian on the Equator), and

E

(90°E on the Equator).

Figure 4.3 Hemisphere projection: three rate constants and seven solvent parameters. Circles represent points in their natural positions, crossed circles antipodes of points that fall in the hidden hemisphere, that is, the negatives of the indicated variables.

Figure 4.4 Rate constants (logarithms) for reactions, (a–c) in the seven solvents versus the corresponding elements of the parallel vectors

z

i

.

Figure 4.5 Kosower’s

Z

versus Reichardt’s

E

T

(30). Points marked with crosses are omitted from the statistics. They are for compounds flagged by the last authors as uncertain and a group of esters and ethers that seem to have values of

Z

clustered about 64–65 kcal mol

−1

, lying above the trend. No strong HBD solvents are included.

Figure 4.6 The dielectric parameter

β

μ

plotted against for 101 solvents. Subclasses 3a, 3b, and chloroform are combined as class 3.

Figure 4.7 versus the composite variable: ;

C

1

 = 0.003764,

C

2

 = 0.4974. In this and the following figure, the points marked by crosses are for acetic acid. The circled cross is for the dimer; the plain cross is for the monomer. Neither is included in the statistics. The standard error of estimate

s

 = 0.046.

Figure 4.8 versus the new composite variable: ;  = 0.04174,  = 0.4792,

s

′ = 0.037. See caption for Figure 4.5.

Figure 4.9 Similar to Figure 4.6, but plotting the square root of

β

μ

. Classes 1 and 2 are now combined.

Figure 4.10 Acidity parameters: Taft–Kamlet

α

versus Swain’s

A

j

. Circles are for O-acids and (filled) formamide; crosses are for C-acids. Separate least-squares lines are fitted for O-acids and for C-acids.

Figure 4.11 Swain’s acity parameter

A

j

versus Catalán’s (2001) SA. Circles are for O-acids (alcohols and acetic acid); crosses are for assorted C-acids. The straight line, fitted to the O-acids only, is represented by ; the curve, fitted to all, is represented by .

Figure 4.12 Nine basic parameters for 13 non-HBD solvents. Parameters represented are (latitude, longitude in degrees): 1.

β

(72, −107); 2.

B

j

(9, −3); 3. SB (62, −151); 4. DN (63, −177); 5.

C

b

(41, 75); 6.

E

b

(48, −15); 7. −Δ

H

BF3

(75, −168); 8.

D

s

(76, 36); 9. Δ

δ

(CHCl

3

) (78, 0) The error circles are somewhat arbitrary; they are drawn with radii proportional to the difference between the sum of squares of the three direction cosines and unity.

Figure 4.13 Six parameters for base softness. Three (

μ

,

D

s

, and ΔΔν(C–I)) are discussed by Chen

et al

. (2000).

R

v

is the volume polarization, is derived from the Drago parameters

C

B

and

E

B

, and

S

orb

is the reciprocal of the LUMO–HOMO energy difference.

Chapter 05

Figure 5.1 Illustration of the approximation of radial parts of hydrogen-like atomic orbitals (heavy curves) as sums of Gaussians: (a) 1

s

and (b) 3

p

. Not to scale and fitted by trial and error (not optimized).

Figure 5.2 Two Gaussian functions in one dimension,

G

1

and

G

2

, and their product (heavy curve).

Figure 5.3 Intermolecular potentials illustrated: hard sphere (dashed), Lennard-Jones (black). A Buckingham potential adjusted to have the same collision distance, depth of well and long-range energy as the latter was indistinguishable from it at the scale of this diagram.

Figure 5.4 A stereo view of the ST-2 model of water (Stillinger and Rahman, 1974). Charges are placed at the tetrahedral angles: +

q

at 100 pm from the center representing the proton positions and −

q

at 80 pm representing the nonbonding pairs. The electrostatic potential due to these charges is in addition to a Lennard-Jones potential acting between oxygen nuclei as centers. Other models used have included further properties such as polarizability. See, e.g., Kusalik and Svishchev (1994), Svishchev

et al.

(1996).

Figure 5.5 Periodic boundary conditions in two dimensions. Twelve molecules; neighbor list boundary shown for molecule

A

. The molecule entering the box just above

A

is leaving at the bottom.

Figure 5.6 Cavity in a continuous dielectric (schematic) containing a ketene molecule. Rotation about the molecular axis is not hindered.

Figure 5.7 Stepwise experimental, gas-phase enthalpies of hydration of H

3

O

+

, circles (Lau

et al.

, 1982), and corresponding theoretical energies, squares (Newton 1997), plotted versus the reciprocal of the number of water molecules. NB,

n

in this and the following figure corresponds to

n

 − 1 in Table 5.2.

Figure 5.8 Stepwise experimental, gas-phase enthalpies of hydration of alkali-metal ions (Kebarle, 1972) plotted versus the reciprocal of the number of water molecules. Circles, Li

+

; squares, Na

+

; oblique crosses, K

+

; plus signs, Rb

+

; triangles, Cs

+

.

Figure 5.9 Reaction profile (schematic) for an S

N

2 reaction in the gas phase, showing potential wells corresponding to ion–dipole complexes. The energy of the transition state may by higher (heavy curve) or lower (light curve) than the energy of the reactants. Entropy effects are important. See discussion in Shaik

et al.

(1992) or Olmstead and Brauman (1977).

Figure 5.10 Calculated internal energies in the gas phase (short dashes) and the potential of mean force in DMF (long dashes) and in aqueous solution (solid curve) for the reaction of Cl

−

with CH

3

Cl as a function of the reaction coordinate,

r

c

, in angstroms.

Figure 5.11 Taft–Kamlet

α

versus

α

1

, the H-bond acidity scale based on the acidity contribution to the solvatochromism of Reichardt’s dye 30 (Cerón-Carrasco

et al.

, 2014a, b). Acids of types O–H are shown as circles, C–H acids as crosses, and N–H as squares.

Figure 5.12 Gibbs energy reaction profile of the palladacycle-catalyzed methanolysis of methyl parathion (

4

) showing the Gibbs energy of activation calculated with and without inclusion of the rotational and translational contributions to the entropy change and the experimental value.

Figure 5.13 Calculated pair distribution functions for argon in water: argon–oxygen

g

Ar–O

and argon–hydrogen

g

Ar–H

(Guillot

et al.

, 1991). Note the near coincidence of the first maxima.

Figure 5.14 Calculated pair distribution functions for methane in water: carbon–oxygen

g

C–O

, carbon–hydrogen(water)

g

C–Hw

, hydrogen(methane)–hydrogen(water)

g

Hm–Hw

, and hydrogen(methane)–oxygen

g

Hm–O

. Note the near coincidence of the first maxima for C–Hw and C–O.

Figure 5.15 Enthalpy and entropy contributions to the Gibbs energy of solvation of methane in a number of solvents (Bagno, 1998).

Chapter 06

Scheme 1

Scheme 2

Figure 6.1 Reaction profile illustrations of some representative reactions for changes from protic (solid curves) to dipolar-aprotic (dashed) media, showing enthalpy as the ordinate. Using the classification in Table 6.5, cases 6.10 and 6.12 represent balancing situations, and 6.11 represents positively reinforced and 6.13 represents negative transition-state control.

Figure 6.2 Effect of added dimethyl sulfoxide on the rate constants,

k

2

, for the reaction of benzyl chloride with the 9-cyanofluorenyl anion in ethanolic media at 35.7°C.

Chapter 08

Scheme A

Figure 8.1 The process by which a switchable-hydrophilicity solvent can be used to extract soybean oil from soybean flakes without a distillation step. The dashed lines indicate the recycling of the solvent and the aqueous phase.

Scheme B

Guide

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SOLVENT EFFECTS IN CHEMISTRY

Second Edition

ERWIN BUNCELROBERT A. STAIRS

Copyright © 2016 by John Wiley & Sons, Inc. All rights reserved

Published by John Wiley & Sons, Inc., Hoboken, New JerseyPublished simultaneously in Canada

No part of this publication may be reproduced, stored in a retrieval system, or transmitted in any form or by any means, electronic, mechanical, photocopying, recording, scanning, or otherwise, except as permitted under Section 107 or 108 of the 1976 United States Copyright Act, without either the prior written permission of the Publisher, or authorization through payment of the appropriate per-copy fee to the Copyright Clearance Center, Inc., 222 Rosewood Drive, Danvers, MA 01923, [phone], fax [phone], or on the web at www.copyright.com. Requests to the Publisher for permission should be addressed to the Permissions Department, John Wiley & Sons, Inc., 111 River Street, Hoboken, NJ 07030, [phone], fax [phone], or online at http://www.wiley.com/go/permissions.

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Library of Congress Cataloging-in-Publication Data

Erwin Buncel     Solvent effects in chemistry / Erwin Buncel, Robert A. Stairs.          pages    cm    Includes bibliographical references and index.

    ISBN 978-1-119-03098-0 (cloth)1.  Solvation.    2.  Chemical reactions.    3.  Solvents.    I.  Buncel, E.    II.  Title.    QD543.S684 2015    541′.34–dc23                                    2015010522

Cover image courtesy of Professor Errol Lewars, Trent University.

PREFACE TO THE SECOND EDITION

The present work is in effect the second edition of Buncel, Stairs, and Wilson’s (2003) The Role of the Solvent in Chemical Reactions. In the years since the appearance of the first edition, the repertoire of solvents and their uses has changed considerably. Notable additions to the list of useful solvents include room-temperature ionic liquids, fluorous solvents, and solvents with properties “switchable” between different degrees of hydrophilicity or polarity. The use of substances at temperatures and pressures near or above their critical points as solvents of variable properties has increased. Theoretical advances toward understanding the role of the solvent in reactions continue. There is currently much activity in the field of kinetic solvent isotope effects. A search using this phrase in 2002 yielded 118 references to work on their use in elucidating a large variety of reaction mechanisms, nearly half in the preceding decade, ranging from the SN2 process (Fang et al., 1998) to electron transfer in DNA duplexes (Shafirovich et al., 2001). Nineteen countries were represented: see, for example, Blagoeva et al. (2001), Koo et al. (2001), Oh et al. (2002), Wood et al. (2002). A similar search in 2013 yielded over 25,000 “hits.”

The present edition follows the pattern of the first in that the introductory chapters review the basic thermodynamics and kinetics needed for describing and understanding solvent effects as phenomena. The next chapters have been revised mainly to improve the presentation. The most changed chapters are near the end, and attempt to describe recent advances.

Some of the chapters are followed by problems, some repeated or only slightly changed from the first edition, and a few new ones. Answers to most are provided.

We are grateful to two anonymous colleagues who reviewed the first edition when this one was first proposed, and who pointed out a number of errors and infelicities. One gently scolded us for using the term “transition state” when the physical entity, the activated complex, was meant. He or she is right, of course, but correcting it in a number of places required awkward circumlocutions, which we have shamelessly avoided (see also Atkins and de Paula, 2010, p. 844.). We hope that most of the remaining corrections have been made. We add further thanks to Christian Reichardt for steering us in new directions, and we also thank Nicholas Mosey for a contribution to the text and helpful discussions, and Chris Maxwell for Figure 5.11. We add David Poole, Keith Oldham, J. A. Arnot, and Jan Myland to the list of persons mentioned in the preface to the first edition who have helped in different ways. Finally, we thank the editorial staff at Wiley, in particular Anita Lekhwani and Cecilia Tsai, for patiently guiding us through the maze of modern publishing and Saravanan Purushothaman for careful copy-editing that saved us from many errors. Any errors that remain are, of course, our own.

EB, Kingston, OntarioRAS, Peterborough, OntarioApril 15, 2015