Oxidation in Organic Chemistry, Part B

253 Pages·1973·2.88 MB·English

by  WALTER S. TRAHANOVSKY (Eds.)

OXIDATION IN
ORGANIC CHEMISTRY
Edited by WALTER S. TRAHANOVSKY
IOWA STATE UNIVERSITY OF SCIENCE
AND TECHNOLOGY
AMES, IOWA
PART Β
1973
ACADEMIC PRESS New York and London
COPYRIGHT © 1973, BY ACADEMIC PRESS, INC.
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Contributors
Gordon A. Hamilton, Department of Chemistry, The Pennsylvania State
University, University Park, Pennsylvania
Donald G. Lee, Department of Chemistry, The University of Saskatchewan
Regina Campus, Regina, Saskatchewan, Canada
Patrick D. McDonald, Department of Chemistry, The Pennsylvania State
University, University Park, Pennsylvania
W. G. Nigh, Department of Chemistry, University of Puget Sound, Tacoma,
Washington,
Robert J. Ouellette, Department of Chemistry, The Ohio State University,
Columbus, Ohio
Matthijs van den Engh, Department of Chemistry, The University of Sas
katchewan Regina Campus, Regina, Saskatchewan, Canada
vii
Preface
Some of the most important and common reactions in organic chemistry
involve oxidation and reduction. Because of the importance of this class of
reactions, numerous reagents have been developed which will bring about
certain oxidations or reductions selectively. This treatise is devoted to detailed
discussions of specific oxidants or topics involving oxidation of organic
compounds.
In this volume, three chapters are devoted to specific oxidants, cupric ion,
thallium(III), and ruthenium tetroxide. In these chapters an attempt is made
to present the main results of all of the literature which refers to these
oxidants, organized according to oxidation of specific functional groups or
types of organic compounds. A fourth chapter is concerned with a type of
oxidation, the oxidative coupling of phenols. Only results pertinent to the
mechanism of this conversion are presented, and an attempt is made to unify
and correct some of the mechanistic thoughts about this very important
reaction.
The level of all chapters is such that experts in these areas of research and
students and researchers that wish a thorough and rigorous discussion of
these topics should find them useful. In general, emphasis is on the scope and
preparative use as well as the mechanistic aspects of the various oxidations.
WALTER S. TRAHANOVSKY
ix
Contents of Part A
Edited by Kenneth B. Wiberg
Ross STEWART, Oxidation by Permanganate
KENNETH B. WIBERG, Oxidation by Chromic Acid and Chromyl Compounds
W. A. WATERS AND J. S. LITTLER, Oxidation by Vanadium(V), Cobalt(III),
and Manganese(III)
WILLIAM H. RICHARDSON, Ceric Ion Oxidation of Organic Compounds
RUDOLF CRIEGEE, Oxidations with Lead Tetraacetate
C. A. BUNTON, Glycol Cleavage and Related Reactions
AUTHOR INDEX—SUBJECT INDEX
xi
C H A P T ER I
Oxidation by Cupric Ion
W. G. Nigh
I. Introduction 1
II. The Chemical Nature of Ionic Copper 2
III. Oxidation of Acetylenes 11
IV. Oxidation of Aldehydes 31
V. Oxidation of Alcohols 35
VI. Oxidation of Amines 51
VII. Oxidative Halogenation 67
VIII. Oxidation of Mercaptans 84
IX. Oxidation of Organometallic Compounds 85
X. Oxidation of Carboxylic Acids 91
XI. Miscellaneous Oxidations 95
I. Introduction
The use of copper(II) as an oxidizing agent for organic compounds dates
back to a medieval preparation called Egyptian ointment. This concoction
was prepared by heating a mixture of honey (fructose and glucose), vinegar
(acetic acid), and verdigris (cupric acetate). Alchemists dispensed this mixture
for both medicinal and cosmetic purposes. It was not until 1815 that the
reddish brown precipitate produced in this reaction was shown to be cuprous
oxide.1
The first indication of the potential value of this reaction occurred in 1841,
when it was observed that D-glucose precipitated cuprous oxide from an
alkaline solution of cupric sulfate, whereas sucrose was unreactive toward
this reagent.2 Further work with carbohydrates led Barreswil to suggest that
an alkaline solution of cupric tartrate might be used as a qualitative test for
1 Vogel, Schweigger'sJ. 13, 162 (1815).
2 Trommer, Ann. Chem. Pharm. 39, 360 (1841); Chem. Zentra. 12, 762, (1841).
1
2 W. G. NIGH
reducing sugars.3 A few years later, Fehling worked out a useful analytical
procedure based on BarreswiPs suggestion.4
Since these early beginnings, copper(II) has been found to be a useful
oxidizing agent for a wide range of organic substrates. It offers the advantages
of high selectivity as a result of its mild oxidizing power and its compatibility
with a variety of solvent systems.
Π. The Chemical Nature of Ionic Copper
Copper is known to exist in the 0, 1 + , 2 + , and 3 + oxidation states. Of
these, copper(III) is the least encountered because of its very large oxidation
potential. The standard oxidation potentials of a number of copper species
are presented in Table I5'8 for comparison. The few known compounds
TABLE I
STANDARD OXIDATION POTENTIALS OF COPPER
Electrode reactions0 Solvent E° Reference
Cu = Cu+ + e~ HO -0.521 5
a
Cu + 2 NH = Cu(NH)+ + e~ HO + 0.12 5
3 32 a
Cu + 2 C5H5N = Cu(CHN)+ + e~ C5H5N -0.175 6
5 5 2
Cu+ = Cu2+ + e~ HO -0.153 5
a
Cu(NH)+ = Cu(NH)2+ + e~ HO -0.308 7
32 32 a
Cu+ = Cu2+ + e~ C5H5N -0.688 6
Cu+ = Cu2+ + e~ CHCN -1.28 8
3
Cu+ = Cu2+ + e~ Dioxane -0.250 7
Cu(MO)+ = Cu(MO)2+ + e- H0 -0.250 7
2 2 2
Cu(Im)+ = Cu(Im)2+ + e~ HO -0.317 7
2 2 a
Cu(CHN)+ = Cu(CHN)2+ + e' HO -0.300 7
5 5 4 5 5 4 a
Cu(EDA)+ = Cu(EDA)2+ + e~ H0 + 0.360 7
2 2 2
Cu(Bip)+ = Cu(Bip)2+ + e~ HO -0.120 7
2 2 a
Cu(Bip)+ = Cu(Bip)2+ + e~ Dioxane -0.251 7
2 2
Cu(Biq)+ = Cu(Biq)2+ + e~ Dioxane -0.771 7
2 2
Cu(Phen)+ = Cu(Phen)2+ + e~ H0 -0.174 7
2 2 2
Cu(Phen)+ = Cu(Phen)2+ + e~ Dioxane -0.296 7
2 2
Cu(5-NOPhen)+ = Cu(5-N0Phen)2 + + e~ Dioxane -0.379 7
a 2 2 2
Cu(2,9-MePhen)+ = Cu(2,9-MePhen) 2+ + e~ HO -0.594 7
2 2 2 2 a
Cu(5-NHPhen)+ = Cu(5-NHPhen)2 + + e~ Dioxane -0.248 7
2 2 2 2
Cu2+ = Cu3+ + e~ HO <-1.8 5
a
0 MO, morpholine; Im, imidazole; EDA, ethylenediamine; Bip, 2,2/-bipyridine;
Biq, 2,2'-biquinoline; Phen, 1,10-phenanthroline.
3 C. Barreswil, /. Pharm. 6, 301 (1844).
4 H. Fehling, Ann. Chem. Pharm. 72, 106 (1849).
5 W. M. Latimer, "The Oxidation States of the Elements and their Potentials in Aqueous
Solutions," p. 183. Prentice-Hall, Englewood Cliffs, New Jersey, 1952.
/. Oxidation by Cupric Ion 3
containing copper(III) appear to exist as paramagnetic octahedral salts.
The steel-blue KCu0, however, is diamagnetic which suggests that it is a
2
square-planar complex. Copper(III) has been suggested as an intermediate
in certain reactions involving catalytic amounts of cupric ion and oxidizing
agents such as hydrogen peroxide, hexachloroiridate(IV), peroxydisulfate,
and hypochlorite9:
H H
2 2
H2C- \ H2CT \
Cu(II) + HO ,Cu(III) + OH + OH
2 a
HC. / H0
2 2
H H
2 2
H
2
-N
HCT
2 \
Puilll) HNCHCHNH + H + + Cu2 +
2 2 2
HC^ /
2
H
2
[O]
HNCHCHNH HNCHCH=NH + Η4
2 2 2 2 2
HNCHCH=NH + HO HNCHCHO + NH
2 2 a 2 2 3
It has also been suggested9 that the oxidation potential of copper(III) is
decreased by complex formation to a point which is low enough to allow its
formation with much milder oxidizing agents.
RCOCH3 + CuCl [RCOCH3—Cu(II)]
2
CuCl
[RCOCH3—Cu(II)] 2 [RCOCH3—Cu(III)]
[RCOCH3—Cu(III)] + ci- [RCOCH3—Cu(I)] + CV
HO
\
RCOCH3 C=CH
2
R
HO
C=CH + CI4 RCOCHCl + H +
2 2
R
There is, however, no direct evidence of the intermediacy of copper(III) in
the absence of strong oxidizing agents. To the contrary, cupric ion oxidations
appear to involve only the cupric-cuprous couple.
β A. K. Gupta, /. Chem. Soc, London p. 3473 (1952).
7 B. R. James and R. J. P. Williams, /. Chem. Soc, London p. 2007 (1961).
8 I. M. Kolthoff, /. Amer. Chem. Soc. 79, 1852 (1957).
9 M. Anbar, Advan. Chem. Ser. 49, 134 (1965).
4 W. G. NIGH
In contrast to copper(III), copper(II) is a relatively mild oxidizing agent.
Copper(II) is the most common valence state of the metal and generally
exhibits a coordination number of four or six. In the majority of cases it
possesses either a square-planar or an octahedral bond orientation. As a
result of its d9 electronic configuration, octahedral copper(II) complexes
usually exhibit a Jahn-Teller distortion. Therefore two trans metal-ligand
distances are greater than the other four, resulting in an elongated octahedral
structure. Less commonly, cupric ion may form distorted tetrahedral,
square-pyramidal, or trigonal-bipyramidal complexes. All mononuclear
copper(II) complexes are paramagnetic. However, in those cases where two
copper(II) ions are held close together, there is a considerable amount of
quenching of the spin moment. In the dimeric copper(II) salt of diazoamino-
benzene (CHNH—N=N—CH) the spins of the two cupric ions are so
e 5 e 5
strongly coupled that the salt is diamagnetic.10
L
As a result of its d10 electronic structure, copper(I) is always diamagnetic.
Cuprous ion is normally either two- or four-coordinated. With monodentate
ligands there is a preference for a linear configuration, while with bidentate
ligands a tetrahedral structure is common. The hydrate of copper(I) is
unstable and disproportionates into copper(II) and metallic copper. However,
in the presence of a suitable complexing agent, cuprous ion may be stabilized
in aqueous solutions and, in some cases, may even become more stable than
cupric ion. Because of its lower charge density, copper(I) should be stabilized
relative to copper(II) by decreasing the dielectric constant (c) of the solvent.
Thus the reduction potential of copper(II) is observed to be lower in water
( = 78.5) than in solvents such as acetonitrile (€ =38.8), pyridine (€ =
€25 20 25
12.3), or dioxane (€ = 2.2) (see Table I). This observed increase in stability
25
is also due in part to coordination with the solvent.
The sharing of the electron pair of a ligand (electron pair donor) by a
metal ion (electron pair acceptor) may, in the Lewis sense, be considered as
10 C. M. Harris and R. L. Martin, Proc. Chem. Soc. London p. 259 (1958).
t
/. Oxidation by Cupric Ion 5
an acid-base reaction. Therefore, it is not surprising that the stability of
coordination compounds generally increases in proportion to the basicity
of the ligand. The nature of the donor atom also affects the stability of the
metal complex. Copper(II) belongs to the group of metal ions which exhibit
the stability orders N>0>S and F"»C1~ >Br_ >I", while copper(I)
exhibits the orders Ν> S> Ο and I" > Br' >CI" »F~.
Providing that the steric requirements are not prohibitive, a ligand which
possesses two or more donor groups may chelate to a single metal ion.
Chelation generally results in a dramatic increase in the stability of the
complex. For example, the bidentate ligand ethylenediamine forms a chelate
with cupric ion which is nearly a billion times more stable than the corre­
sponding ammonia (monodentate) complex.11
ΉΝ^ NH
3 3
Cua++4NH Cu
3 / \
LHN NH
log Κ = 10.0 3 3
CHNH HNCH
2 2 2 2
Cua + +2 HNCHCHNH
a a a 3
CHaNHa HNCHJ
a a
log Κ = 18.7
If the chelating agent is tri- or tetradentate, two or three interlocking rings
may be formed, resulting in an even greater degree of stabilization.12
H=N \l=CH
I I
CH CH
3 3 E = -0.75
£ = +0.02 ll2
1/2
In contrast to the behavior of copper(II), the ethylenediamine complex
of copper(I) does not exhibit any appreciable stabilization relative to the
ammonia complex.7
Cu + +2 NH ^ [Cu(NH)] +
3 32
log Κ = 10.6
Cu+ + 2 HNCHCHNH ^ [Cu(HNCHCHNH)]+
2 2 2 2 2 2 2 22
log Κ = 11.4
This is a result of the formation of only a single coordinate bond between
the diamine and the cuprous ion. The preferential stabilization of copper(II)
11 L. G. Sillen and A. E. Martell, Chem. Soc, Spec. Publ. 17, (1964).
12 M. Calvin and R. H. Bailes, /. Amer. Chem. Soc. 68, 949 (1946).