《有机化学》课程教学资源（参考书籍）Solvents and Solvent Effects in Organic Chemistry, Third, Updated and Enlarged Edition, 2003, Christian Reichardt.pdf_P31-P35

chair-configured, non-dipolar solvent 1,4-dioxane, which often behaves like a polar solvent even though its relative permittivity is low (er ¼ 2:2), is caused by its large nonideal quadrupolar charge distribution [411]. 2.2.3 Dipole-Induced Dipole Forces [32] The electric dipole of a molecule possessing a permanent dipole moment m can induce a dipole moment in a neighbouring molecule. This induced moment always lies in the direction of the inducing dipole. Thus, attraction always exists between the two partners, which is independent of temperature. The induced dipole moment*) will be bigger the larger the polarizability a of the apolar molecule experiencing the induction of the permanent dipole. The net dipole/induced dipole energy of interaction for two di¤erent molecules, each possessing a permanent dipole moment m1 and m2 and polarizabilities a1 and a2, often referred to as the induction or Debye interaction [32], is given by Eq. (2-4). Udipole-induced dipole ¼ 1 ð4p e0Þ 2 a1 m2 2 þ a2 m2 1 r6 ð2-4Þ For a dipolar molecule of m ¼ 3:3 1030 Cm (1 D; e.g. HaaCl) separated from a molecule of polarization volume a ¼ 10 1030 m3 (e.g. C6H6) by a distance of 300 pm, the temperature-independent interaction energy is about 0.8 kJ/mol [26d]. Similarly, a charged particle such as an ion introduced into the neighbourhood of an uncharged, apolar molecule will distort the electron cloud of this molecule in the same way. The polarization of the neutral molecule will depend upon its inherent polarizability a, and on the polarizing field a¤orded by the ion with charge z e. The energy of such an interaction is given by Eq. (2-5). Uion-induced dipole ¼ 1 ð4p e0Þ 2 z2 e2 a 2 r4 ð2-5Þ The importance of both of these interactions is limited to situations such as solutions of dipolar or ionic compounds in nonpolar solvents. 2.2.4 Instantaneous Dipole-Induced Dipole Forces [33, 34, 186] Even in atoms and molecules possessing no permanent dipole moment, the continuous electronic movement results, at any instant, in a small dipole moment m, which can fluctuatingly polarize the electron system of the neighbouring atoms or molecules. This coupling causes the electronic movements to be synchronized in such a way that a mutual attraction results. The energy of such so-called dispersion or London [33] inter- * The induced dipole moment is defined as mind ¼ 4p e0 a E (e0 permittivity of vacuum; a electric polarizability of the molecule; E electric field strength). 2.2 Intermolecular Forces 13

14 2 Solute-Solvent Interactions actions can be expressed as ) 1 (2-6a) where a and are the polarizabilities and I and 2 are the ionization potentials of the two different interacting species 33)When applied to two molecules of the same sub- stance,Eq.(2-6a)reduces to Eq.(2-6b). Udispersion = 3x2.1 (2-6b) (4r·c0)24r6 Dispersion forces are extremely short-range inaction(dep ending on 1/61). sion fo ar all nd mole alone mol which dipol greate p cia espe th high po li part is due to n d ergy of liquid 2b e at con 14 5 100 10g interac 11.3 kJ/mol (-2.7kcal/mol)[35a].T d the average nce r corre pond to thos liquid HCI.It is instructive to compare the these einteractions or two dipoles oth momen m and oriented as in Fig.2-3a,the interaction energy is only -5.3 kJ/mo (-1.1 kcal/mol)[35a].Thus,for HCI and most other compounds the dispersion forces are considerably stronger than the dipole-dipole forces of nearest neighbour tance in the liquid state.However,at larger distances the dispersion energy falls off rapidly As a result of the term in Eq.(2-6b),dispersion forces increase rapidly with the molecular volume and the number of polarizable electrons.The polarizabilityx is con nected with the molar refraction and the index of refraction, according to the equation of Lorenz-Lorentz.Therefore.solvents with a large index of refraction.and hence large optical polarizability,should be capable of enjoying particularly strong dispersion forces.As indicated in Table A-1(Appendix),all aromatic compounds possess relatively high indices of refraction,e.g.quinoline (n=1.6273),iodobenzene(n=1.6200),aniline (n=1.5863),and diphenyl ether (n=1.5763);of all organic solvents,carbon disulfide (n=1.6275)and diiodomethane(n=1.738)have the highest indices of refraction Solvents with high polarizability are often good solvators for anions which also bossess high polarizability.This is due to the fact that the dispersional interactions between the solvents and the large,polarizable anions like II,SCN or the picrate anion are significantly larger than for the smaller anions like Fe,HOe,or R2Ne [36] Perfluorohydrocarbons have unusually low boiling points because tightly held electrons in fluorine have only a small polarizability

actions can be expressed as Udispersion ¼ 1 ð4p e0Þ 2 3a1 a2 2r6 I1 I2 I1 þ I2 ð2-6aÞ where a1 and a2 are the polarizabilities and I1 and I2 are the ionization potentials of the two di¤erent interacting species [33]. When applied to two molecules of the same substance, Eq. (2-6a) reduces to Eq. (2-6b). Udispersion ¼ 1 ð4p e0Þ 2 3a2 I 4r6 ð2-6bÞ Dispersion forces are extremely short-range in action (depending on 1=r6!). Dispersion forces are universal for all atoms and molecules; they alone are responsible for the aggregation of molecules which possess neither free charges nor electric dipole moments. Due to the greater polarizability of p-electrons, especially strong dispersion forces exist between molecules with conjugated p-electron systems (e.g. aromatic hydrocarbons). For many other dipole molecules with high polarizability as well, the major part of the cohesion is due to dispersion forces. For example, the calculated cohesion energy of liquid 2-butanone at 40 C consists of 8% orientational energy, 14% inductional energy, and 78% dispersion energy [35]. Two molecules with a ¼ 3 1030 m3, I ¼ 20 1019 J, and r ¼ 3 1010 m have an interaction potential of 11.3 kJ/mol (2.7 kcal/mol) [35a]. These values of a, I, and the average intermolecular distance r correspond to those for liquid HCl. It is instructive to compare the magnitude of these dispersion forces with that of the dipole-dipole interactions. For two dipoles, both with dipole moments of 3:3 1030 Cm (1.0 D), separated by a distance of r ¼ 3 1010 m and oriented as in Fig. 2-3a, the interaction energy is only 5.3 kJ/mol (1.1 kcal/mol) [35a]. Thus, for HCl and most other compounds, the dispersion forces are considerably stronger than the dipole-dipole forces of nearest neighbour distance in the liquid state. However, at larger distances the dispersion energy falls o¤ rapidly. As a result of the a2 term in Eq. (2-6b), dispersion forces increase rapidly with the molecular volume and the number of polarizable electrons. The polarizability a is connected with the molar refraction and the index of refraction, according to the equation of Lorenz-Lorentz. Therefore, solvents with a large index of refraction, and hence large optical polarizability, should be capable of enjoying particularly strong dispersion forces. As indicated in Table A-1 (Appendix), all aromatic compounds possess relatively high indices of refraction, e.g. quinoline (n ¼ 1:6273), iodobenzene (n ¼ 1:6200), aniline (n ¼ 1:5863), and diphenyl ether (n ¼ 1:5763); of all organic solvents, carbon disulfide (n ¼ 1:6275) and diiodomethane (n ¼ 1:738) have the highest indices of refraction. Solvents with high polarizability are often good solvators for anions which also possess high polarizability. This is due to the fact that the dispersional interactions between the solvents and the large, polarizable anions like Im 3 , Im, SCNm or the picrate anion are significantly larger than for the smaller anions like Fm, HOm, or R2Nm [36]. Perfluorohydrocarbons have unusually low boiling points because tightly held electrons in fluorine have only a small polarizability. 14 2 Solute-Solvent Interactions

2.2.5 Hydrogen Bonding [37–46, 187–190, 306] Liquids possessing hydroxy groups or other groups with a hydrogen atom bound to an electronegative atom X are strongly associated and have abnormal boiling points. This observation led to the contention that particular intermolecular forces apply here. These are designated as hydrogen bridges, or hydrogen bonds, characterized by a coordinative divalency of the hydrogen atom involved. A general definition of the hydrogen bond is: when a covalently bound hydrogen atom forms a second bond to another atom, the second bond is referred to as a hydrogen bond [38]. The concept of hydrogen bonding was introduced in 1919 by Huggins [37]. The first definitive paper on hydrogen bonding – applied to the association of water molecules – was published in 1920 by Latimer and Rodebush [191]. All three were working in the Laboratory of G. N. Lewis, University of California, Berkeley/USA. A hydrogen bond is formed by the interaction between the partners RaaXaaH and :YaaR0 according to Eq. (2-7). ð2-7Þ RaaXaaH is the proton donor and :YaaR0 makes available an electron pair for the bridging bond. Thus, hydrogen bonding can be regarded as a preliminary step in a Brønsted acid-base reaction which would lead to a dipolar reaction product RaaXm HaaYlaaR0 . X and Y are atoms of higher electronegativity than hydrogen (e.g. C, N, P, O, S, F, Cl, Br, I). Both inter- and intramolecular hydrogen bonding are possible, the latter when X and Y belong to the same molecule. The most important electron pair donors (i.e. hydrogen bond acceptors) are the oxygen atoms in alcohols, ethers, and carbonyl compounds, as well as nitrogen atoms in amines and N-heterocycles. Hydroxy-, amino-, carboxyl-, and amide groups are the most important proton donor groups. Strong hydrogen bonds are formed by the pairs OaaH O, OaaH N, and NaaH O, weaker ones by NaaH N, and the weakest by Cl2CaaH O and Cl2CaaH N. The p-electron systems of aromatic compounds, alkenes, and alkynes can also act as weak hydrogen bond acceptors [189]. When two or more molecules of the same type associate, so-called homointermolecular hydrogen bonds are formed (Fig. 2-4). The association of di¤erent molecules (e.g. RaaOaaH NR3) results in hetero-intermolecular hydrogen bonds. The designations homo- and heteromolecular [192] as well as homo- and heteroconjugated hydrogen bond are also in use. A remarkable example of a competitive solventdependent equilibrium between homo- and hetero-intermolecular hydrogen-bond associated species has been found in solutions of 4-hydroxyacetophenone and 2-(2- hexyloxyethoxy)ethanol [319]. Hydrogen bonds can be either intermolecular or intramolecular. Both types of hydrogen bonds are found in solutions of 2-nitrophenol, depending on the Lewis basicity of the solvent [298]. The intramolecularly hydrogen-bonded form exists in nonhydrogen-bonding solvents (e.g. cyclohexane, tetrachloromethane). 2-Nitrophenol breaks its intramolecular hydrogen bond to form an intermolecular one in electron-pair donor (EPD) solvents (e.g. anisole, HMPT ). 2.2 Intermolecular Forces 15

Circular hydrogen bonds have been found in the hexahydrate of a-cyclodextrin (cyclohexaamylose) [193]. Hydration water molecules and hydroxy groups of the macromolecule cooperate to form a network-like pattern with circular OaaH O hydrogen bonds. If the OaaH O hydrogen bonds run in the same direction, the circle is called homodromic. Circles with the two counter-running chains are called antidromic, and circles with more randomly oriented chains are designated heterodromic [193]; cf. Fig. 2-4a. Such circular hydrogen bonds can be of importance with respect to the inner molecular structure of water and alcohols (cf. also Fig. 2-1). The question of the exact geometry of hydrogen bonds (distances, angles, lonepair directionality) has been reviewed [194]. The bond dissociation enthalpy for normal hydrogen bonds is ca. 13 ... 42 kJ/mol (3 ... 10 kcal/mol)*). For comparison, covalent single bonds have dissociation enthalpies of 210 ... 420 kJ/mol (50 ... 100 kcal/mol). Thus, hydrogen bonds are approx. ten times weaker than covalent single bonds, but also approx. ten times stronger than the nonFig. 2-4. Homo-intermolecular hydrogen bonds in alcohols, carboxylic acids, and amides (the hydrogen bonds are denoted by dotted lines). Fig. 2-4a. Three types of circular hydrogen bonds: (a) homodromic, (b) antidromic, and (c) heterodromic hydrogen bonds [193]. * Bond dissociation enthalpies outside these limits are, however, known. Examples of weak, normal, and strong hydrogen bonds are found in the following pairs: phenol/benzene (DH ¼ 5 kJ/ mol) [47], phenol/triethylamine (DH ¼ 37 kJ/mol) [47], and trichloroacetic acid/triphenylphosphane oxide (DH ¼ 67 kJ/mol) [48]. An extremely strong hydrogen bond is found in Me4NþHF 2 (DH ¼ 155 kJ/mol) [38]. The strength of a hydrogen bond correlates with the basicity of the proton-acceptor and the acidity of the proton-donor molecule. Compounds with very strong hydrogen bonds have been reviewed [320]. 16 2 Solute-Solvent Interactions

2.2 Intermolecular Forces 17 specific intermolecular interaction forces.The question as to whether or not a hydrogen nd is stronger than the equivalent deuterum bond is addressed in reference 321:the D-bond seems to be somewhat stronger than the H-bond in the case of neutral hydro- gen-bonded complexes,but the reverse is true for charged complexes. Hydrogen bonds are characterized by the following structural and spectroscopic features [39]:(a)the distances between the neighbouring atoms involved in the hydrogen bond [X and Y in Eq.(2-7)]are considerably smaller than the sum of their van der Waals radii;(b)the X-H bond length is increased and hydrogen bond formation causes its IR stretching mode to be shifted towards lower frequencies(for exceptions see reference 1901):(c)the dipolarity of the X-H bond increases on hydrogen bond for mation,leading to a larger dipole moment of the complex than expected from vectorial addition of its dipolar components R-X-H and Y-R';(d)due to the reduced elec- tron density at H-atoms involved in hydrogen bonds,they are deshielded,resulting in substantial downfield shifts of theirH NMR signals;(e)in hetero-molecular hydrogen bonds,a shift of the Bronsted acid/base equilibrium R-X-H...Y-R'-R- x...H-Y-R'to the right-hand side with increasing solvent polarity is found (cf. Section 4.4.1 and references [195,322]for impressive examples). Up until now there has been no general agreement as to the best description of the nature of the forces in the hydrogen bond [42-46.The hydrogen bond can be described as a dipole-dipole or resonance interaction.Since hydrogen bonding occurs only when the hydrogen is bound to an electronegative atom,the first assumption concerning the ature nd was that it consists of a dipole-dipole interaction such as orted d by the fact that the strongest hydrogen bonds are formed in pairs in which the dre gen is bonded to the most elec tronegative elem nts (e.g.F-H...F,AH =-155 kJ/mol).The ate str gth of he hydr bole inte ons is due to the uch smaller size of the hydr n atom relative to any which allows it to oach another dipole closely.Thiss ole model nts for the try of the hyd d a linea arranger nt ximizes the ons to believ that more is involved in hvdro ogen bonding than simply ted dipole-dipole tio The f hy dicate siderable erlap o A h should le o rep of sym hydrogen rt,a of th of the x_ to a According t tion can d by two contributing"protomeric" structures,which only in the position of the proton* R-X-H--Y-R' R-X...H-YO-R (2-8) *The pounds such as benzene in terms of a resonance hybrid 13231

specific intermolecular interaction forces. The question as to whether or not a hydrogen bond is stronger than the equivalent deuterium bond is addressed in reference [321]: the D-bond seems to be somewhat stronger than the H-bond in the case of neutral hydrogen-bonded complexes, but the reverse is true for charged complexes. Hydrogen bonds are characterized by the following structural and spectroscopic features [39]: (a) the distances between the neighbouring atoms involved in the hydrogen bond [X and Y in Eq. (2-7)] are considerably smaller than the sum of their van der Waals radii; (b) the XaaH bond length is increased and hydrogen bond formation causes its IR stretching mode to be shifted towards lower frequencies (for exceptions see reference [190]); (c) the dipolarity of the XaaH bond increases on hydrogen bond formation, leading to a larger dipole moment of the complex than expected from vectorial addition of its dipolar components RaaXaaH and YaaR0 ; (d) due to the reduced electron density at H-atoms involved in hydrogen bonds, they are deshielded, resulting in substantial downfield shifts of their 1H NMR signals; (e) in hetero-molecular hydrogen bonds, a shift of the Brønsted acid/base equilibrium RaaXaaH YaaR0 S Raa Xm HaaYlaaR0 to the right-hand side with increasing solvent polarity is found (cf. Section 4.4.1 and references [195, 322] for impressive examples). Up until now there has been no general agreement as to the best description of the nature of the forces in the hydrogen bond [42–46]. The hydrogen bond can be described as a dipole-dipole or resonance interaction. Since hydrogen bonding occurs only when the hydrogen is bound to an electronegative atom, the first assumption concerning the nature of the hydrogen bond was that it consists of a dipole-dipole interaction such as RaaXdmaaHdl YdmaaR0 . This viewpoint is supported by the fact that the strongest hydrogen bonds are formed in pairs in which the hydrogen is bonded to the most electronegative elements (e.g. FaaH Fm, DH ¼ 155 kJ/mol). The greater strength of the hydrogen bond compared with non-specific dipole-dipole interactions is due to the much smaller size of the hydrogen atom relative to any other atom, which allows it to approach another dipole more closely. This simple dipole model accounts for the usual linear geometry of the hydrogen bond, because a linear arrangement maximizes the attractive forces and minimizes the repulsion. However, there are reasons to believe that more is involved in hydrogen bonding than simply an exaggerated dipole-dipole interaction. The shortness of hydrogen bonds indicates considerable overlap of van der Waals radii and this should lead to repulsive forces unless otherwise compensated. Also, the existence of symmetrical hydrogen bonds of the type Fdm H Fdm cannot be explained in terms of the electrostatic model. When the XaaY distance is su‰ciently short, an overlap of the orbitals of the XaaH bond and the electron pair of :Y can lead to a covalent interaction. According to Eq. (2-8), this situation can be described by two contributing ‘‘protomeric’’ structures, which di¤er only in the position of the proton*). ð2-8Þ * The term ‘‘protomeric structure’’ was obviously introduced in analogy to the well-known ‘‘mesomeric structures’’, which are used to describe the electronic ground state of aromatic compounds such as benzene in terms of a resonance hybrid [323]. 2.2 Intermolecular Forces 17