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The ReactivityЦSelectivity Principle An Imperishable Myth in Organic Chemistry.

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Minireviews
H. Mayr and A. R. Ofial
DOI: 10.1002/anie.200503273
Reaction Kinetics
The Reactivity–Selectivity Principle: An Imperishable
Myth in Organic Chemistry**
Herbert Mayr* and Armin R. Ofial
Keywords:
frontier-orbital effect · Hammond effect · intrinsic
barriers · linear free energy relationships · reaction
kinetics
Dedicated to Professor Jrgen Sauer
on the occasion of his 75th birthday
The reactivity–selectivity principle (RSP), once a tenet of organic
chemistry, eroded during the 1970s and was more or less abandoned by
1980. Although it has been clear for more than 25 years that a decrease
in selectivity with increasing reactivity can only be expected with
certainty if diffusion control is approached, the RSP has survived as an
intuitively appealing rule. This Minireview shows why selectivity
cannot generally decrease with increasing reactivity and highlights the
weaknesses of the theoretical foundations of the RSP.
1. Introduction
Dozens of reviews on the reactivity–selectivity principle
(RSP) have been published during the past three decades,[1]
most of which agree in the conclusion that one cannot
generally expect a decrease in selectivity with increasing
reactivity.[2–14] Twenty-five years ago, Arnett summarized:
“The RSP seems to be such a reasonable proposition that
there have been a number of attempts to justify it on more or
less theoretical grounds and to relate it to the Hammond
postulate. However, as the RSP has been subjected to
increasing scrutiny, its credibility has eroded steadily. It now
seems plain that, although the concept of the principle is
attractive, it is virtually useless in practice as a general rule
since the number of times it does not work is approaching the
number of times that it does”.[7b]
Since we are in full consent with this statement, the
question arises, why another overview?
In spring 2005, the senior author presented an analysis of
reactivity–selectivity relationships as part of a lecture “Mythology in Organic Chemistry” at several universities. The
succeeding discussions revealed that a majority of chemists
still consider the decrease in selectivity with increasing
[*] Prof. Dr. H. Mayr, Dr. A. R. Ofial
Department Chemie und Biochemie
Ludwig-Maximilians-Universit3t M4nchen
Butenandtstrasse 5–13 (Haus F), 81 377 M4nchen (Germany)
Fax: (+ 49) 89-2180-77717
E-mail: herbert.mayr@cup.uni-muenchen.de
[**] We thank the Deutsche Forschungsgemeinschaft and the Fonds der
Chemischen Industrie for financial support of this work.
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reactivity as a rule, and violations of
the RSP as rare exceptions that are
only encountered in some special cases.[15–17] We believe that the kinetics of
electrophile–nucleophile
combinations that we have studied in recent years are ideally suited
to correct this view and to clarify why the RSP can only hold
in a narrow range. This Minireview highlights some milestones in the rise and the fall of the reactivity–selectivity
principle, presents some recent reactivity–selectivity relationships, and concludes with a simple geometrical argument as to
why it is impossible that selectivity generally decreases with
increasing reactivity.
2. Experimental Facts
Typically, the chemistry student first encounters the
inverse relationship between reactivity and selectivity when
studying the selectivities of radical-initiated halogenations of
alkanes. The different selectivities of chlorinations and
brominations (Figure 1) can be rationalized by the late,
productlike, transition states of hydrogen abstraction by the
bromine radical and the early, reactantlike, transition states of
hydrogen abstractions by the chlorine radical. The low
activation barriers for the hydrogen abstractions by ClC, which
are shown in Figure 1, have been determined experimentally.[20]
As fast kinetic methods were not readily available in the
1960s and 1970s, much information about the relative
stabilities of reactive intermediates was derived from studies
of selectivities under the assumption that low selectivity is
generally associated with high reactivity. Investigations by
Sneen and co-workers[21] on the selectivities of solvolytically
generated carbocations were systematically extended by
Raber, Harris, Hall, and Schleyer and seemed to confirm
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Figure 1. Product ratios[18] for the chlorination and bromination of 2methylbutane and calculated reaction enthalpies (DrH in [kJ mol1][19]).
the RSP:[22] When alkyl chlorides were solvolyzed in solutions
of sodium azide in aqueous acetone (80 %), highly stabilized
carbocations, such as the tritylium ion, which are formed
rapidly in SN1 reactions, show high selectivities, whereas lessstabilized, more-reactive carbocations, such as the adamantyl
cations, which are formed slowly in the solvolysis reactions,
show low selectivities (Figure 2).
Only one year later, Ritchie reported a fundamentally
different behavior in the reactions of stabilized carbocations
and diazonium ions with nucleophiles (Figure 3).[23] Morereactive electrophiles showed the same selectivities as lessreactive electrophiles, leading Ritchie to the formulation of
electrophile-independent nucleophilicity parameters N+ and
nucleophile-independent electrophilicity parameters log k0
[Eq. (1)].
log ðk=k0 Þ ¼ N þ
ð1Þ
As this finding was in sharp contrast to the reactivity–
selectivity principle, which had been considered as a fundaHerbert Mayr obtained his PhD in 1974 (R.
Huisgen; LMU Mnchen), and after postdoctoral studies (G. A. Olah; Cleveland), he
completed his habilitation in 1980
(P. von R. Schleyer; Erlangen). After professorships in Lbeck and Darmstadt he returned to the LMU Mnchen in 1996. In
2003 he became a member of the Bavarian
Academy of Sciences. He received the
Alexander von Humboldt Honorary Fellowship of the Foundation for Polish Science
(2004) and the Liebig Denkmnze (GDCh,
2006). His interests include quantitative
approaches to organic reactivity and the
theory of polar organic reactions.
Angew. Chem. Int. Ed. 2006, 45, 1844 – 1854
Figure 2. Stability–selectivity (N3/H2O) plot for carbocations derived
from solvolysis of alkyl chlorides (25 8C, 80 % aqueous acetone, data
from reference [22]).
mental law, a direct consequence of allegedly “basic principles”, that is, of the Bell–Evans–Polanyi principle and the
Leffler–Hammond relationship (see Section 3), numerous
attempts were made to explain why RitchieCs carbocations did
not follow the general rule.[15–17, 24] Better solvation of the lessstabilized carbocations was considered as one of the reasons
for the unexpected behavior shown in Figure 3.[25]
Armin R. Ofial studied chemistry at the TU
Darmstadt with Alarich Weiss and Herbert
Mayr (doctoral degree, 1996). Since 1997
he has been a Research Associate at the
LMU Mnchen. In 2005 he received a
postdoctoral fellowship of the Japan Society
for the Promotion of Science for a research
stay in the group of Shunichi Fukuzumi
(Osaka, Japan). His research interests include reactions of iminium ions, electrontransfer processes, and chemical kinetics.
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H. Mayr and A. R. Ofial
Figure 3. Constant selectivity relationships for the reactions of
carbenium and diazonium ions with nucleophiles.[23a]
Figure 4. Dependence of the second-order rate constants of SN2
reactions of 3- and 4-substituted pyridines with a variety of reagents
on the pKaH value of the pyridines:[26] A) MeSO3F in 2-nitropropane,
25 8C;[7b] B) PhSO2Cl in H2O, 25 8C;[27] C) EtSO3F in 2-nitropropane,
25 8C;[7b] D) allyl bromide in nitromethane, 40 8C;[28] E) methyl iodide in
nitrobenzene, 25 8C;[29] F) methyl iodide in 2-nitro-propane, 25 8C;[7b]
G) ethyl iodide in 2-nitropropane, 25 8C.[7b]
In retrospect, the astonishment about the Ritchie relationship is amazing, because at that time plenty of other
reactivity–selectivity relationships were known which showed
change in selectivity is caused by the variable rate constants
a similar behavior as described by Ritchie.[7b, 30] Figure 4 shows
kw for the reactions with water (Figure 7).[37d, 38, 39]
some examples.
The azide-clock method, later developed by Jencks and
Nowadays hundreds of correlations displaying constant
Richard, uses this phenomenon for deriving absolute rate
selectivity are known,[31] and an example from
our own laboratory is shown in Figure 5:[32–34]
The 1-nitroethyl anion, the most reactive carbanion of the series shown, differentiates between a series of quinone methides in the same
way as the much less nucleophilic carbanion
obtained from dimedone.
GieseCs observation that the selectivities of
radicals may be inverted by variation of the
temperature (Figure 6)[14] cast further doubt on
the generality of the reactivity–selectivity principle.
Although the importance of the diffusioncontrol limit for the interpretation of reactivity–
selectivity correlations had already been noted
by Kemp and Casey,[35] it was Rappoport and
Ta-Shma[13, 36] who demonstrated that the change
from activation to diffusion control is responsible for the different reactivity–selectivity relationships observed by Ritchie (Figure 3)[23] and
by Raber, Harris, Hall, and Schleyer (Figure 2)[22]: The decrease in selectivity with increasing reactivity of the carbocations shown in
Figure 2 is due to the fact that N3 undergoes
diffusion-controlled
reactions
(kaz = 5 F
109 m 1 s1) with all carbocations that react with Figure 5. Plot of the reactivities of various carbanions toward quinone methides versus
water at a rate faster than 105 s1,[37] and the the reactivities of the ethyl cyanoacetate anion toward these electrophiles.[32, 34]
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In the late 1980s and early 1990s, the senior author
reported a general method to determine rate constants for the
reactions of carbocations with p-systems (e.g., alkenes and
allyl silanes),[43–47] and it was found that RitchieCs constant
selectivity relationship [Eq. (1)] was not sufficient to describe
the whole body of data. Excellent correlations were obtained,
however, when an additional, nucleophile-specific slope
parameter s was introduced.[48, 49] In [Eq. (2)] k is the
log k20 o C ¼ sðE þ NÞ
Figure 6. Variation of the reaction temperature causes an inversion of
the selectivities of radical reactions (figure taken from reference [14a]).
Figure 7. Rate constants kaz ([m1 s1], 20 8C) for the azide trapping of
laser flash photolytically generated tritylium (triangles) and benzhydrylium ions (circles) as a function of kw ([s1], 20 8C), the rate constants
for the decay of carbenium ions in acetonitrile/water mixtures
(CH3CN/H2O 33:67 for tritylium ions;[37d] CH3CN/H2O 20:80 for
benzhydrylium ions[38]).
constants for the reactions of carbocations with solvents from
the ratio of azide-trapped products to solvolysis products
(R-N3/R-OSolv).[40–42]
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ð2Þ
second-order rate constant ([m 1 s1]), E is a nucleophileindependent electrophilicity parameter, and N is an electrophile-independent nucleophilicity parameter.
The resulting reactivity–selectivity correlations have been
discussed previously,[50, 51] and Figure 8 illustrates how one
might arrive at contradictory conclusions when the RSP is
employed to derive relative reactivities of carbocations. Let
us assume that we would try to determine the reactivity order
of carbenium ions solely from competition experiments
towards different allyl silanes:
* By choosing (2-methylallyl)trimethylsilane (A) and prenyltrimethylsilane (C) as a pair of competing nucleophiles,
we observe that selectivity (S = log kAlog kC) decreases
from left to right in Figure 8. By utilizing the RSP, we
would (correctly) conclude that the reactivities of the
carbenium ions increase from left to right.
* The choice of the pair prenyltrimethylsilane (C) and
allylchlorodimethylsilane (D) for an analogous series of
competition experiments would reveal an increase in
selectivity (S = log kClog kD) from left to right until kC
approaches the diffusion limit. This time, the RSP would
lead to the (incorrect) conclusion that the reactivities of
the carbenium ions in Figure 8 decrease from left to right.
* To complete the confusion, we might consider a series of
competition experiments with a third pair of nucleophiles,
for example, allyltrimethylsilane (B) and prenyltrimethylsilane (C). When moving from left to right in Figure 8 we
would first observe a decrease in selectivity (S =
log kBlog kC), then no selectivity for (4-MeOC6H4)2CH+, and finally an increase in selectivity (S =
log kClog kB). In this range, application of the reactivity–selectivity principle would again lead to an incorrect
order of carbocation reactivities.
* A comparison of A and B would finally reveal constant
selectivity in the range k < 108 m 1 s1.
These examples clearly show that changes in selectivity
alone do not allow one to derive relative reactivities of the
carbenium ions. Reactivity orders can only be derived from
competition experiments if one of the competition partners is
known to react with diffusion control.[50] Measurements of
absolute rate constants are, therefore, needed to determine
the correct reactivity order of electrophiles.
Since the publication of the relationships illustrated in
Figure 8, Equation (2) has become the basis for the most
comprehensive nucleophilicity scales presently available.[52–55]
As shown in Figure 9, reactions of carbanions,[33, 56, 57]
amines,[58] alkoxides,[58, 59] phosphanes,[60] enamines,[52, 61] diazo-
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Figure 8. Rate constants for the reactions of carbenium ions with allylsilanes.[50] Dots correspond to directly measured rate constants and shaded
bars to the results of competition experiments. Plots versus the electrophilicity parameter E, defined by Equation (2).
Figure 9. Direct comparison of the reactivities of different classes of nucleophiles (in CH2Cl2, 20 8C). Plots versus the electrophilicity parameter E,
defined by Equation (2).
alkanes,[62] ketene acetals,[52, 63] enol ethers,[52, 64] water,[65]
transition metal p-complexes,[52, 66] alkenes,[52] hydride donors,[49, 67] and arenes[52, 68] with benzhydrylium ions follow
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linear correlations with the electrophilicity parameter E. In
some cases it has been shown that these correlation lines bend
for k > 108 m 1 s1 and asymptotically approach the diffusion
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limit, which is typically at (2–4) F 109 m 1 s1 for cation–
molecule reactions in CH3CN (20 8C) and at (2–4) F
1010 m 1 s1 for cation–anion combinations in the same solvent.
The abscissa of Figure 9 covers 24 orders of magnitude
from very strong electrophiles, such as the unsubstituted
benzhydrylium ion on the right, to relatively weak electrophiles such as the amino-substituted quinone methides on the
left. Only few of the available data sets have been selected for
the construction of Figure 9 in order to avoid extensive
crossings. If selectivity were generally to decrease with
increasing reactivity, all correlation lines of Figure 9 would
converge when going from left to right, that is, from low
reactivity to high reactivity. For clear geometrical reasons, this
is not possible. One cannot even invent a scenario in which
selectivity generally decreases with increasing reactivity!
The reader can easily verify this statement by trying to
draw correlation lines for the reactions of about 100
substrates Bj with reagents Ai of variable reactivity in
Figure 10. Whereas it is easy to “invent” correlation lines
Figure 10. Correlations for the reactions of the substrates Bj with a
series of reagents Ai of variable reactivity: Try to formulate the RSP as
a general principle!
for about 10 substrates Bj, all of which converge when moving
towards the right according to the RSP, one will soon get into
trouble when constructing more correlation lines without
intersections, even if the search is not restricted to linear
correlations and takes into account correlation lines of any
curvature.
In view of the fact that inverse relationships between
reactivity and selectivity cannot even be formulated as a
general principle, it may appear curious that the RSP has so
long been considered a general rule. The reason is evident:
Most early applications were concerned with short-lived
intermediates whose reactivities were studied by competition
experiments and not by direct rate measurements. If one
extrapolates Figure 9 to higher reactivities, one arrives at a
picture related to Figure 8. If we only consider the high
reactivity part of this figure (log k > 6), one will often find a
decrease in selectivity when going from left to right, because
now the faster of two parallel reactions is close to diffusion
control, while the slower of the two reactions still shows
variable rates.
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If it is evident that one of two competing reactions is
diffusion-controlled, one can rely on a behavior as expected
by the RSP and, in addition, derive important information
from selectivities determined by competition experiments. In
such cases, the diffusion-controlled reaction can be considered as a clock, and product ratios can be utilized for
determining rate constants for the reactions of carbocations
with solvents[42] or propagation rate constants in ionic
polymerizations.[69, 70]
3. How Solid Are the “Theoretical Foundations” of
the Reactivity–Selectivity Principle?
The Bell–Evans–Polanyi principle rationalizes the increasing rates of many reactions with increasing exothermicities by describing the transition states C as a blend of reactant
and product configurations (Figure 11).[71] BrønstedCs catalysis law has been interpreted in the same way.[31]
Figure 11. The Bell–Evans–Polanyi principle.[71]
A quantitative relationship between the effects of structural variation on rate and equilibrium constants was given by
the Hammond–Leffler a value,[72, 73] which was supposed to
change from 0 for reactantlike transition states to 1 for
productlike transition states and was considered as a measure
of the position of the transition states (Figure 12).[5, 31, 72]
BordwellCs finding of a = 1.54 for the deprotonation of
nitrotoluenes in water[74] which corresponds to a = 0.54 for
Figure 12. Different positions of a) reactant- and b) productlike
transition states.
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H. Mayr and A. R. Ofial
the reverse reaction (Figure 13) cast doubts on the use of a as
an indicator of the position of the transition state, because it
was outside the defined range (0 < a < 1).
Figure 13. Deprotonation of nitrotoluenes.
The identity reactions depicted in Figure 14 have been
reported to be accelerated by electron-donating as well as by
electron-accepting substituents X.[75] As DG0 = 0 throughout
Figure 14. Identity reactions of benzyl chlorides.
these reaction series, a Hammond–Leffler a value of infinity
results,[5, 76, 77] which shows that BordwellCs observation of an
unusual a value cannot be a rare exception resulting from
special solvation effects.[78] Rather, it becomes clear that a
cannot be a measure of the position of the transition state.
As pointed out by Pross and Shaik, the transition state has
its own identity,[5, 6, 77, 79] and cannot simply be described as a
blend of reactant and product configurations, which is implied
in the Bell–Evans–Polanyi and the Leffler–Hammond treatment. The Marcus theory describes this phenomenon by
variable intrinsic barriers.[80]
In our work, we have found that the Leffler–Hammond
coefficients a in hydride-transfer reactions do not reflect the
position of the transition states. Kinetic investigations showed
that in hydride abstractions from CH groups by carbocations,
modification of DG0 caused by variation of the carbocation
has a much stronger effect on DG° than an equal change in
DG0 which is caused by variation of the hydride donor.[67d]
To rationalize this observation, we performed ab initio
quantum chemical calculations for the transfer of an allylic
hydrogen atom from substituted propenes to substituted allyl
cations (Figure 15).[81] As expected, electron-releasing substituents X and Y in the hydride acceptors diminish the
exothermicity (Figure 15 a), whereas electron-releasing substituents in the hydride donors increase the exothermicity of
the reactions (Figure 15 b).
The geometric parameters and the charge distribution in
the activated complexes are in line with the Hammond
postulate and show that increasing exothermicity shifts the
transition states on the reaction coordinate toward reactants.
By plotting DG° versus DG0, a value of 0.72 is found for the
Hammond–Leffler parameter a = dDG°/dDG0 when the
hydride acceptor is varied (Figure 15 a), and a = 0.28 when
the hydride donor is varied (Figure 15 b). If a were a measure
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Figure 15. Brønsted plots for the forward and reverse hydride-transfer
reactions. a) Variation of the hydride acceptor; b) variation of the
hydride donor; energies in [kcal mol1], MP2/6-31 + G(d,p)//RHF/631 + G(d,p).[81]
of the transition state, one would derive late transition states
for the reactions in Figure 15 a and early transition states for
the reactions in Figure 15 b. Because the identity reaction
(X = Y = CH3) is part of both reaction series, one might come
up with an early or a late transition state for the same
reaction, depending on the reaction series considered. Thus
the Hammond–Leffler a value cannot reflect the position of
the transition state.
A solution to this paradox came from a consideration of
the degenerate hydride transfers illustrated in Equation (3).[81]
MP2 calculations indicate that the activation free energies
of these reactions, that is, the intrinsic barriers DG0° as
defined by the Marcus theory,[80] increase with increasing
electron-releasing abilities of X and Y. Therefore, the intrinsic
barriers must also be variable in the reaction series of
Figure 15. Since substituent variation in the donor (Figure 15 b) influences reaction enthalpy and intrinsic barriers in
the opposite sense, while substituent variation in the acceptor
(Figure 15 a) affects both terms in the same sense, the slopes
of the two correlations are different, even though identical
transition states are involved.[81]
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It is well known that the Bell–Evans–Polanyi treatment
cannot be employed for explaining the variable reactivities of
cycloaddition reactions. Reactions of comparable DG0 (e.g.,
most Diels–Alder reactions) differ tremendously in rates,[82, 83]
which can be explained by perturbational MO theory (PMO
theory).[82] Since substituent variation has a larger effect on
DG°, the smaller the HOMO–LUMO gap is (Figure 16), in
frontier-orbital-controlled reactions selectivity generally increases with increasing reactivity,[14a] in contrast to the
postulate of the RSP.
the other hand, will increase the rate of the reaction of the
more-reactive nucleophile (higher HOMO) to a higher extent
and result in an increase in selectivity (frontier-orbital effect,
Figure 17 b). Though intellectually unsatisfactory, the linear
correlations depicted in Figure 9 are explained by a superposition of the curved lines of Figure 17.[84]
If the Hammond effect dominates (Figure 17 a) selectivities will decrease with increasing reactivities and if the
frontier-orbital term dominates (Figure 17 b) selectivities will
increase with increasing reactivities. Constant selectivity
relationships will result if both effects cancel each other.
Figure 18 illustrates this analysis for the reactions of
carbocations with allyltrimethylsilane and prenyltrimethylsilane. Both nucleophiles react at similar rates, even though the
Figure 16. Frontier-orbital interactions in Diels–Alder reactions.
Very often discussions of substituent effects consider only
one of these effects, that is, either the Leffler–Hammond (a
consequence of Bell–Evans–Polanyi) or the frontier-orbital
effect. Such treatments are sufficient if one of the effects
clearly dominates. If both terms operate in the same direction,
it is even difficult to identify the origin of the experimentally
observed effects. Thus, the increase in reactivity from left to
right in Figures 8 and 9 may either be explained by a decrease
in DG0 (Bell-Evans-Polanyi) or by a lowering of the LUMO
energies of the electrophiles.
Both effects will affect the relative reactivities of two
nucleophiles to a different degree, however (Figure 17). An
increase in electrophilicity shifts the transition states towards
reactants and, according to the Bell–Evans–Polanyi treatment, causes a decrease in selectivity (Hammond effect,
Figure 17 a). A lowering of the LUMO of the electrophile, on
Figure 18. Frontier-orbital interactions for the reactions of carbocations
R+ with allyltrimethylsilane and prenyltrimethylsilane.
HOMO of prenyltrimethylsilane is higher than that of
allyltrimethylsilane (Figure 8). Probably the more-favorable
electronic interaction in the case of prenylsilane is compensated by the stronger steric repulsion. As e1 > e2, the order of
the second-order perturbation energies is De2 > De1. A
variation in the carbocation, that is, of LU(R+), will then
affect De2 more than De1 because the HOMO–LUMO gap e2
is smaller than e1. The higher slopes of p-systems with methyl
groups at the position of electrophilic attack can thus be
explained.[52, 85]
Presently, we feel unable to evaluate the relative importance of the Hammond and the frontier-orbital effects, and for
that reason consider it problematic to predict generally slopes
of the correlation lines that imply changes in selectivities in
the activation-controlled range.[86]
4. Conclusion
Figure 17. Opposing effects on reactivity–selectivity relationships:
a) Hammond effect: effect of the variation of DG0 on the relative
reactivities of two nucleophiles; b) frontier-orbital effect: effect of the
variation of eLUMO on the relative reactivities of two nucleophiles.
Angew. Chem. Int. Ed. 2006, 45, 1844 – 1854
For geometrical reasons it is impossible to formulate the
RSP as a general principle (Figure 10). A general decrease in
selectivity with increasing reactivity can only be expected in
fast reactions that are close to diffusion control. In reactions
with rate constants k < 108 m 1 s1, selectivity can decrease,
increase, or remain constant when reactivity is increased, and
additional information is necessary to predict the effects of
structural variation on selectivity.
It is the relative importance of the Hammond effect and
the frontier-orbital effects that controls the slopes of the
correlations as shown in Figures 4, 5, 8, and 9 and thus
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accounts for the variable reactivity–selectivity relationships
commonly encountered.
Why is RSP, despite its frequently shown failures, still
considered as a rule from which only some rare exceptions
exist? Why doesnCt it follow the typical course of scientific
theories that are modified or abandoned when disproved by
experiments?[87] We feel that the reactivity–selectivity principle has adopted the status of a myth that is no longer subject
to common rules of scientific practice. Chemists like to
compare the behavior of molecules with that of living beings,
and our intuition is shaped by the experience that the hungry
lion shown in Figure 19 is highly (re)active and will unselec-
Figure 19. A catchy illustration of the reactivity–selectivity principle.[88]
tively chase the next prey, be it game, man, or rabbit. The
satisfied lion, on the other hand, will be selective and not hunt
a fast gazelle, but only grab a rabbit that comes within reach.
However, this parable is incomplete: Like the Bell–Evans–
Polanyi principle and the Leffler–Hammond treatment, this
consideration derives the behavior of the lion from the
condition of the animal before and after the meal. However,
chemical reactivity is not only controlled by the thermodynamic difference between reactants and products but also by
transition-state-specific properties (e.g., frontier-orbital interactions), which are included in the Marcus intrinsic
barriers. What is the counterpart of this term in Figure 19?
Hard to tell how the lion is feeling during the meal, but we
definitely know that we enjoy eating a delicious dessert, also if
we are not hungry.
Received: September 15, 2005
Published online: February 7, 2006
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One referee commented: “The statement, that constant selectivity relationships will result when Hammond effects and
frontier-orbital effects cancel, is not enough. It would be hard
to imagine that these effects accidentally cancel in such an
overwhelming number of cases. Is there something else that
makes constant selectivity relationships the rule?” We fully
agree with this statement. Unfortunately, we cannot give an
answer and hope that this Minireview will stimulate theoretical
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Reproduction of photos by courtesy of Ultimate Africa Safaris,
Seattle, USA.
2006 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Angew. Chem. Int. Ed. 2006, 45, 1844 – 1854
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