Cr(VI) reagents have been shown to be carcinogenic, upon ingestion either
through the stomach or the lungs. Not many years ago chromic acid solutions
were the way to clean glassware (they remove organic compounds very well
and leave the glass sparkling). They have been shown to leave traces of Cr(VI)
on the glass, which can be death to a Grignard reagent, for example. The
cancer-causing properties have resulted in strict regulations for disposal (you
can't!), and thus these reagents are no longer used routinely. Interestingly,
the chromate (CrO4-) looks like sulfate to cells, and is
readily incorporated. Once in the cell, it oxidizes something and is converted
to Cr(III); the Cr(III) looks a lot like Zn(II) and other biologically important
ions. It is the Cr(III) that actually causes the damage that leads to cancer but
the Cr(III) itself cannot get into the cells - it has to enter as Cr(VI).
Manganese reagents (Permanganate)
Permanganate ion, KMnO4, will accomplish many of the reactions
that chromate will. It is most commonly used in basic solution where it is
reduced to the brownish-black solid, MnO2; in acid solution, it is
reduced to the pink, soluble ion Mn++. With heat in base, it oxidizes
alcohols without breaking carbon-carbon bonds and benzyl CH bonds completely to
benzoic acids (their salts since the solution is basic).
There are no deactivated forms of permanganate that permit oxidation of
primary alcohols to aldehydes. However, careful reaction of cold basic
permanganate with alkenes will give 1,2-diols stereospecifically as osmium
tetroxide does. Heating that mixture will cause cleavage of the CC bond.
Osmium Tetroxide
Another extremely toxic compound, OsO4, is used to selectively
and stereospecifically oxidize alkenes to 1,2-diols. The reaction is usually
done with another oxidizing agent in the solution to regenerate the
OsO4 so that only a catalytic amount of thne reagent need be used.
Its toxicity is greater than lead or mercury, and similar in effect, attacking
the nervous system and the liver.
Sodium metaperiodate and lead tetracetate also oxidize 1,2-diols, but
they cleave the carbon-carbon bond between them like hot basic permanganate.
Nitric Acid and Nitrous Acid
Concentrated nitric acid is a 69% aqueous solution, but is extremely
dangerous because it is a rapid and powerful oxidizing agent. For example (do
not try this except in a fume hood and with supervision), it reacts with copper
to form greenish cupric nitrate in solution and gaseous oxides of nitrogen which
include NO2, the brown gas which makes photochemical air pollution so
corrosive to tissues. In the presence of sulfuric acid it forms NO2+ which adds to aromatic rings;
since they subsequently lose a hydrogen, the overall reaction is oxidation. Some
of the resulting products, such as TNT and TNB, are explosive: the high
concentration of oxygen in the molecule provides an internal source of oxygen
for very rapid oxidation. (An explosion is a redox reaction that propagates so
fast that the heat cannot be dissipated.) Concentrated nitric acid will convert
alcohols to nitrate esters. Although the esterification is not an oxidation, the
products of that reaction can also be explosives - such as nitroglycerin and
nitrocellulose.
Nitrous acid, HNO2, can be used (with sulfuric acid as a
catalyst) to put an NO (nitroso) substituent on an aromatic ring; this
nitrosation reaction is exactly analogous to the nitration reaction. But nitrous
acid is far more famous for its ability to oxidize amines and amides to
N-nitroso compounds (R2-N-N=O). With anilines, the end product is a
diazonium salt which can be used to make azo dyes or, with N2 as a
leaving group, to attach nucleophiles to the aromatic ring. With other amines
and amides, the resulting N-nitroso compounds are quite carcinogenic and can be
formed in meats that have been cured with sodium nitrite, the salt of nitrous
acid. There is no doubt that cultures using a lot of nitrate- and nitrite-cured
meats have a higher incidence of cancer of the digestive tract.
Halogens
In oxidizing ability, the halogens follow the expected order: F2
> Cl2 > Br2 > I2. All should be
treated with great respect when working with them, as they can do a great deal
of tissue damage very quickly. Fluorine reacts so explosively with so many
organic compounds that working with it requires special facilities; most people
send things out for fluorine treatment to the experts. The most commonly used
halogens are Cl2 and Br2; Cl2, a gas, can be
generated in dilute solution from bleach, NaOCl (more about this later) and
Br2 can be purchased - it is a volatile reddish-brown corrosive
liquid.
Their power as oxidizing agents is revealed by the ease with which they
oxidize CH bonds: benzyl > alkyl, and tertiary > secondary > primary.
For less reactive CH bonds, Cl2, Br2 and I2
need heat and/or light to react in a reasonable time; a hydrogen on an
sp2 or sp carbon can be assumed not to react (see below for other
reactions of alkenes and alkynes). The mechanism of this oxidation is
well-understood and included in nearly every organic chemistry course as an
example of a reaction with free radicals as intermediates. Chlorine and bromine
atoms (radicals) formed from organohalogen compounds that make it to the
stratosphere are responsible for the partial destruction of the ozone layer;
they catalyze the conversion of O3 back to O2.
In addition, halogens undergo polar reactions. For example, they add to the
carbon-carbon pi bond of alkenes and alkynes and the halogenate very rapidly the
alpha-CH bond of aldehydes, ketones and carboxylic acid esters (via the enol,
which has a carbon-carbon double bond). To understand these oxidations, it is
helpful to think of a Br-Br molecule as having a small contribution to its
resonance hybrid from Br+Br-; the Br+ is the
oxidizing agent (it gains to electrons to become effectively Br-),
and the Br- is left over. The addition of Br2 to double
bonds illustrates this model well: first the Br2 adds Br+
stereospecifically syn, then the Br- remaining adds from the
other side. One reagent that often substituted for chlorination is NaOCl (sodium
hypochlorite, common household bleach) in water; part of the NaOCl is converted
to HOCl, which reacts as if it were HO-Cl+. These
reactions are similar to oxidations by peroxides.
Selective oxidations require a less reactive reagent. For example,
N-bromosuccinimide will react with allyl or benzyl compounds to
substitute for one of the allyl hydrogens without adding to any double
bonds. Resistant, very stable substrates, like aromatic compounds, need an extra
boost to be able to react with halogens. To oxidize aromatic rings with
halogens, use a strong Lewis acid
catalyst, FeCl3
Forms of Oxygen and Peroxides
Ozone, O3, is bent and polar. It is made in the upper atmosphere
by collision of an oxygen atom with an oxygen molecule and in the lower
atmosphere by reaction of oxygen with oxides of nitrogen formed in photochemical
smog. Its polarity makes it fairly reactive. One common reaction is an addition
to alkenes to form a five-membered ring, with the three oxygens still attached
to one another called a molozonide, which rearranges and ultimately converts the
carbon-carbon double bond into two carbon-oxygen double bonds. If the reaction
is done in the presence of oxidizing agents (like hydrogen peroxide, below),
carboxylic acids and ketones are formed, but it is possible to collect aldehydes
(and ketones) if the reaction is done in the presence of zinc. This reaction is
at least in part responsible for the destructive effect of ozone on rubber
(tires, for example) and lungs.
Dioxygen, O2, is rarely used as a reagent in a laboratory
setting, but is commonly used in the chemical industry, where cost it THE most
important thing. Usually a catalyst is needed and/or heat, or even better
light. It is UV light that makes it possible for dioxygen to dissociate
to form oxygen atoms, which react with dioxygen to form ozone; lower energy
light will produce an excited state of oxygen which reacts with alkenes. Even
without light, oxygen will slowly oxidize aldehydes when they are exposed to air
(to the carboxylic acid), but most reactions with oxygen at the concentration of
the atmosphere are too slow at room temperature to be noticed. The small amount
of reaction of oxygen with ethers produces peroxides which can detonate with the
slightest touch; anhydrous ethers should be danted and used up quickly to
minimize the chance of this happening. Never touch a bottle of an ether that is
supposed to be liquid if there are crystals in the bottom! And avoid diisopropyl
ether, the most dangerous of all. Mammals use some pretty clever tricks to get
oxygen to react only where it's needed, such as complex transporter molecules
like hemoglobin and myoglobin to get it to the site, where it is delivered with
high local concentration.
Hydrogen peroxide, H2O2, is a moderately strong
oxidizing agent; it even slowly oxidizes water. Its reactions tend to occur via
radical paths, but it is easy to understand the reactivity if you imagine the
HOOH to split into OH- and OH+; the OH+ would
of course be the oxidizing agent, as it desperately needs electrons. Its
reactivity increases if the peroxy (OOH) part of the molecule is linked to an
electron withdrawing group, as in peroxyacids.
Peroxyacids (also known as peracids), RCO3H, have an OOH group
attached to the carbonyl group. It is even more helpful here to think of the
actual oxidizing agent as OH+, since the RCO2-
is such a good leaving group. Peroxyacids are stronger oxidizing agents than
hydrogen peroxide itself, converting alkenes into epoxides, for example. This
reaction with alkenes is mechanistically similar to the addition of bromine to
alkenes; the bromonium ion intermediate looks a lot like epoxide (except for the
charge). See halogens
above.
Weak Oxidizing Agents
Very mild oxidizing agents such Ag+, Cu++ are usually
used to do easy oxidations, such as oxidation of an aldehyde to a carboxylic
acid in the presence of other oxidizable groups such as alcohols. Both are used
to test for reducing sugars (i.e. oxidizable sugars which can reduce the
reagent), which all have aldehyde groups. Tollens' reagent consists of a
solution of Ag(NH3)2+ which forms a silver
mirror on the glass when reduced. Fehling's and Benedict's tests consist of blue
Cu++ complexed with tartrate or citrate respectively; both produce a
brick-red precipate of Cu2O upon reduction.
March 20, 1998, last revised February 1999.