Oxidizing Agents

© Linda M. Sweeting 1998

Oxidation is defined in chemistry as gain of oxygen, loss of hydrogen or loss of electrons; the loss of electrons enables you to calculate an oxidation state.

Metal Oxyacids

Chromium Reagents (Chromate)

All forms of Cr(VI) are powerful oxidizing agents, and oxidize any CH bonds on a carbon with an oxygen as far as possible without breaking any carbon-carbon bonds; for example, secondary alcohols are converted to ketones, and aldehydes to carboxylic acids. The most common reagents are: H2CrO4; K2Cr2O7 + H2SO4; CrO3 + H2SO4; they are approximately equivalent. They oxidize the activated CH bonds next to an aromatic ring, the "benzyl" hydrogens, so completely that they usually convert any alkyl benzene to a benzoic acid. Cr(VI) reagents are so powerful that they can also oxidize alkenes and alkynes, breaking the carbon-carbon bond as ozone does, but this reaction is not synthetically useful.
In non-aqueous solutions, oxidation by Cr(VI) does not go to completion (the intermediate partially oxidized material containing Cr must be hydrolyzed for oxidation to continue); thus, under these conditions, primary alcohols may be oxidized to aldehydes without forming carboxylic acid. The most common reagents for this partial oxidation are: PCC, or pyridinium chlorochromate (formed by dissolving CrO3 and HCl in pyridine); Collins reagent (CrO3 in CH2Cl2); chromyl chloride (CrO2Cl2).
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 (see the essay on acids) 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.


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.