Oxidizing Agents

Type of physical science: Chemistry

Field of study: Chemical compounds

Oxidizing agents remove electrons from other species, thus reducing themselves in redox reactions. Oxygen is the oxidizing agent in the processes of combustion, corrosion, and biochemical metabolism.

Overview

The idea of oxidation number was developed to describe aspects of the state of combination of elements. An element in a molecule is given an arbitrary number, called an oxidation number, to establish whether that element is combined with more or fewer electronegative elements. The concept is entirely artificial, and, unlike the charge of an ion, oxidation number cannot generally be determined experimentally. Nevertheless, its use is of paramount importance in understanding the distribution of electrons in covalent bonds.

An oxidation-reduction reaction, often referred to as redox reaction, involves an exchange of electrons. The species that loses electrons has its oxidation number increased and is said to be oxidized, while the species that gains electrons gets a lower oxidation number and is said to be reduced. As a result, the species that gains electrons creates the oxidation process by abstracting electrons and is called an oxidizing agent. Similarly, the species that loses the electrons is responsible for the reduction of the oxidizing agent and is called a reducing agent. It is worth noting that oxidation and reduction occur together in the same reaction, and that one process cannot occur without the other.

The electron-transfer reaction can be separated into two half-reaction equations, one representing the oxidation process and the other representing the reduction process.

A → A⁺ⁿ + ne^-(oxidation) B + ne^- → B^-n(reduction) where ne- is the number (n) of electrons (e-) involved in a half-reaction equation.

The half-reaction equation in which electrons appear on the right-hand side is called the oxidation half-reaction. The half-reaction equation in which electrons appear on the left-hand side is called the reduction half-reaction. The oxidation half- reaction supplies electrons to the reduction half-reaction. The reduction half-reaction potential (E^o_red) is a measure of the effectiveness of a reagent to act as an oxidizing agent and is measured in volts. It can be considered as equivalent to the free-energy change for the reduction process to occur. The more positive E^o_red is, the stronger the oxidizing agent.

The strongest oxidizing agent (and therefore the element with the highest reduction potential, 2.889 volts) is fluorine, F2. In other words, fluorine can oxidize any other element.

This is precisely the reason why it is seldom used as an oxidizing agent: It is too dangerous to work with. In contrast, a weak oxidizing agent does not gain electrons readily and is capable of reacting only with those species that are very easily oxidized. The weakest oxidizing agent is the lithium (Li+) ion, whose E^o_red is -3.040 volts. All standard potentials in water solution are tabulated in descending order (the F₂+2e^-→2F^- reaction first), and their E^o_red values are relative to that of the following reaction, whose reduction potential is arbitrarily given the value of 0.000 volts:

2H⁺(aqueous)+2e^-→H₂(gas) Since the half-reactions are listed in decreasing E^o, it follows that any oxidizing agent is stronger than any that stands below it in the table. On the other hand; the substances to the right of the half-equation are reducing agents; therefore, the order of relative strength as reductants is inverted. As a result, if F2 has a very high tendency to be reduced to F-, then F- must have a very low tendency to be oxidized to F2.

For a spontaneous redox reaction to take place, the voltage of the cell is always positive. Thus, if the calculated reaction is a positive quantity, the reaction will occur spontaneously in the laboratory. If the calculated voltage is negative, the reaction will not occur; instead, the reverse reaction will be spontaneous, since reversing the reaction leads to changing the sign of the E^o_red.

Applications

Historically, the term "oxidation" was used by scientists to denote the combination of an element with oxygen. It later took a broader meaning that includes reactions that do not involve oxygen. The first known examples of oxidation involved the combination of oxygen with other elements, whether metals or nonmetals. A number of such compounds, known also as oxides, exist naturally on earth or the atmosphere. Such examples include water, silicon dioxide in quartz, carbon dioxide as one of the constituents of atmospheric air, and several minerals (such as aluminum oxide in corundum and ferric oxide in hematite).

Oxidizing agents are of great importance in the synthesis of compounds and chemical analysis of specimens. The selection of a specific oxidizing agent will depend on its special properties. Chlorine, for example, is a powerful oxidant but is rarely used as such because it is a poisonous gas. Most common reagents are the permanganate (MnO₄^-), chromate (Cr₂O₄^-2), and dichromate (Cr₂O₇^-2) anions. They are all powerful oxygen-atom suppliers and very useful in organic synthesis.

Potassium permanganate (KMnO₄) is a popular reagent used in analyzing various metals in ores. For example, the amount of iron in an ore can be calculated by converting all iron into ferrous (Fe⁺²) ions, which can be titrated against a known solution of potassium permanganate. The purple color of permanganate will be converted to the practically colorless Mn⁺² ion at the equivalence point. Acidic solutions of potassium dichromate are used in the breath analyzer, a device used by the police to test drivers suspected of being drunk. The alcohol in the breath is blown by the driver into the orange dichromate solution and is converted into acetic acid (vinegar), while producing the green chromic ion. The degree of intoxication can be measured by the color change of the oxidizing agent (the dichromate ion) and is directly read from a calibrated meter on the instrument.

The amount of ozone in polluted air or the concentration of sodium hypochlorite (the active ingredient) in bleach can also be calculated via an oxidizing agent. The sample to be analyzed (containing ozone or hypochlorite) is first treated with a potassium iodide solution to produce iodine, which forms a dark-blue compound with starch. The solution is then titrated with a solution of sodium thiosulfate, whose exact concentration is known, and the equivalence point is determined when the blue iodine-starch solution abruptly changes to colorless.

Nitric acid (HNO₃) is a powerful oxidizing agent with relatively high E^o values (+0.96 volt for conversion to NO and +0.80 volt for conversion to NO₂). It is used extensively in the isolation and purification of metals from ores because it can oxidize almost all metals, except gold, platinum, rhodium, and iridium. The aqua regia (kingly water) mixture is a 3:1 relative ratio of hydrochloric and nitric acid that was discovered by the alchemists of the fourteenth century and is capable of dissolving even the above-mentioned metals. Another very useful oxidizing agent is hydrogen peroxide (H₂O₂), which is used in the bleaching of paper and textiles, purification of drinking water, sewage treatment, and sterilization of milk containers.

Oxychlorine trifluoride (OF₃Cl) has been synthesized for possible use as an oxidizer in rocket engines. Oxygen itself is a very important oxidizing agent. It is an oxidizing agent because its oxidation state changes from zero (in the molecular oxygen form) to minus two (after oxidation).

Photochemical excitation leads to an excited singlet electronic state. This is often accomplished through energy transfer from an organic compound, such as the dye fluorescein, which absorbs light and transfers it to oxygen to form an excited state of a relatively long life. Combustion involves the reaction of oxygen with an organic compound to produce carbon dioxide and water.

Incomplete combustion leads to the formation of carbon monoxide and is a result of insufficient presence of oxygen. The role of biochemical combustion is fundamental in oxidizing carbohydrates, fats, and proteins, thus providing the source of energy for sustaining the life process. The tragic explosion of the space shuttle Challenger in 1986 was a result of the uncontrolled exothermic reaction between hydrogen and oxygen. Hydrogen, stored as a liquid, is the fuel used to propel space shuttles outside the earth's gravitational pull, and oxygen was the oxidizing agent. The same reaction between oxygen and hydrogen to produce water and great amounts of heat was responsible also for the disastrous explosion of the German airship Hindenburg in 1937. Finally, the role of oxygen as the oxidizer in metal corrosion is very important. When exposed to air, metals are covered with an oxide coating, which in some cases (aluminum, chromium, nickel) is thin and protects the metal and its luster; however, in other cases, the effect is deteriorating. Thus, iron undergoes an oxidation-reduction reaction with oxygen to produce ferrous hydroxide, which leads to ferric hydroxide through another oxygen-initiating redox reaction, and eventually to rust. It is worth mentioning that iron does not rust in water unless oxygen is present. Rusting, however, also requires water, since oxygen dissolved in oil is not enough. The process is accelerated by other factors, such as the pH of the solution and the presence of salts, as well as stress on the iron.

The silver ion (Ag⁺) is the oxidizing agent in black-and-white photography. After exciting the grains of silver bromide on the thin gelatin coating of photographic paper, the silver ion oxidizes the developer, which is a solution containing the mild reducing agent hydroquinone, which is oxidized to quinone. The black metallic silver particles formed on the film are directly proportional to the amount of light that originally fell on the film. Finally, industry has produced numerous oxidizing agents that have found tremendous use in reactions that involve only one active site in synthetic processes. Such reagents include vanadium pentoxide V₂O₅, osmium tetroxide (OsO₄), sodium periodate (NaIO₄), manganese dioxide (MnO₂), and ruthenium tetroxide (RuO₄).

Electrochemical redox reactions are very useful in the preparation of pure elements. A variety of electrodes have been developed (mercury, carbon, gold, platinum) that act as the intermediary in the electron transfer. Synthetic organic electrochemistry is not very wide in reactions but encompasses several reactions that include anodic oxidations, such as the Kolbe reaction, in which carboxylic acid anions are oxidized to produce radicals that lead to the formation of dimer hydrocarbons and carbon dioxide.

Context

The role of oxidizing agents is of paramount importance to everyday life and is expected to play an even greater role in the future. Living beings will always use oxygen as the oxidizer to sustain life. The metabolism of carbohydrates, lipids, and proteins involves a series of oxidation processes that are catalyzed by the enzymes in the living organisms. Certain organisms are also found to be able to oxidize organic compounds and convert the C-H of alkanes into the C-OH of alcohols. These have been suggested for possible use in cleaning oil spills.

Oxygen as an oxidizer is always going to play an important role in the corrosion of metals (such as rusting of iron) and combustion of hydrocarbons. Corrosion is a major problem in modern life because of the frequent replacement of oxidized parts, such as the deterioration of the iron rods in the Statue of Liberty. Extensive research is in progress trying to understand the mechanism of corrosion and the role oxygen plays as the oxidizing agent.

Modern industry will always require oxidizers in the production of a greater variety of products. Processes such as bleaching, ozonolysis, and peroxide oxidations are essential in the synthesis and manufacture of polymers and their industrial products. Such types of oxidations will include vapor-phase (ammonia to nitric acid, hydrogen chloride to chlorine), liquid-phase (steel formation from air oxidation of molten iron), and solid-phase (potassium permanganate from manganese dioxide and potassium hydroxide).

Water acts as the oxidizing agent in photosynthesis, a process in which carbon dioxide is transformed into carbohydrates. This phenomenon, coupled with oxygen acting as the oxidizer, provides the assurance that life will be sustained as long as oxidation and oxidizing agents are present.

Principal terms

COMBUSTION: an exothermic oxidation reaction that often occurs with organic compounds (to produce carbon dioxide or monoxide and water) and metals (to yield oxides)

CORROSION: a destructive chemical process; often applied to the oxidation of a metal (such as the rusting of iron)

OXIDATION: a process that involves loss of electrons and leads to an increase in oxidation number

OXIDATION NUMBER: a number assigned to an atom in a molecule or an ion to reflect its state of oxidation

OXIDIZING AGENT: a species that accepts electrons in an oxidation-reduction reaction

REDUCING AGENT: a substance that gives up electrons in an oxidation-reduction reaction; the reducing agent is oxidized

REDUCTION: the gain of electrons from an element or compound, which leads to a decrease in oxidation number

Bibliography

Brady, James E., and Gerard E. Humiston. GENERAL CHEMISTRY: PRINCIPLES AND STRUCTURE. 4th ed. New York: Wiley, 1986. An undergraduate text for general chemistry students with a section (8.5) on oxidizing and reducing agents applications in the laboratory.

Chang, Raymond. CHEMISTRY. 4th ed. New York: McGraw-Hill, 1991. An excellent general chemistry undergraduate text. Section 4.7 discusses redox titrations and the significance of oxidizing agents such as the permanganate, dichromate, and thiosulfate anions in everyday-life applications of analytical chemistry. Section 3.6 includes a discussion on black-and-white photography.

Masterton, William L., and Cecile N. Hurley. CHEMISTRY: PRINCIPLES AND REACTIONS. Philadelphia: Saunders, 1989. A general chemistry book for undergraduates. Section 12.3, on oxidation-reduction reactions, covers the concept and rules of oxidation numbers. Chapter 21 discusses the strength of oxidizing agents and the corrosion of iron.

Onuchukwu, A. I. "Evaluation of Corrosion Susceptibility of a Metal: Student Corrosion Experiment II." JOURNAL OF CHEMICAL EDUCATION 65 (1988): 934. This is a simple article describing the corrosion process.

"Oxidizing Agent." In MCGRAW-HILL ENCYCLOPEDIA OF SCIENCE AND TECHNOLOGY. 20 vols. 6th ed. New York: McGraw-Hill, 1987. This article discusses the subject mainly from the electrochemical point of view, but also mentions briefly molecular oxygen and biochemical oxidations. It is followed by sections on oxides and oxidation-reduction.

Radel, Stanley, and Marjorie Navidi. CHEMISTRY. St. Paul, Minn.: West, 1990. A good text for undergraduate general chemistry students, with a good coverage of oxidation-reduction reactions in chapter 11.

The Chemistry of Air Pollution

Photochemistry, Plasma Chemistry, and Radiation Chemistry