Redox Reactions

FIELDS OF STUDY: Inorganic Chemistry; Organic Chemistry; Biochemistry

ABSTRACT

The concept of redox reactions and its importance in chemistry-related fields are elaborated. Oxidation and reduction are electron transfer processes that occur simultaneously in a chemical system; the amount of oxidation must exactly balance the amount of reduction.

The Give and Take of Redox Reactions

The term "redox" is derived from reduction and oxidation, two inseparable processes of electron transfer that occur simultaneously in a great many chemical reactions. The process is similar in concept to paying for a purchase: the buyer must give money to the seller, and at the same time, the seller must accept the money from the buyer for the transaction to take place. In redox reactions, the "buyer" is the reducing agent that gives up, or donates, a specific number of electrons, and the "seller" is the oxidizing agent that accepts that specific number of electrons.

In chemical terms, oxidation is the loss of electrons and reduction is the gain of electrons. This is perhaps easiest to keep in mind using a mnemonic such as "Let Every Orange Gorilla Eat Rice." There is no set formula for how to remember the relationship, and any means of recalling the acronym LEO-GER (Loss of Electrons is Oxidation, Gain of Electrons is Reduction) is equally valid.

The most fundamental aspect of redox reactions is that, in the same way that mass must balance throughout the process of a chemical reaction, the electrons must also balance. This can be thought of as the law of conservation of charge, in that electrical charge can be neither created nor destroyed in a chemical reaction.

The Meaning of Oxidation States

All chemical reactions take place at the level of the outermost electrons of the atoms involved in the interaction. Bonds are formed between atoms by the sharing of electrons, and a means of keeping track of the electrons is most helpful in understanding interactions at the atomic level. In many cases, it is the only way to determine the proper balance of a chemical reaction. The oxidation state concept is essentially a way for the chemist to quantify the relationship between two elements involved in a chemical reaction in which electron transfers take place.

The oxidation state of a single atom is essentially the number of electrons that particular atom has gained or given up control over in a chemical interaction. As a very simple example, a sodium atom that reacts with chlorine to produce sodium chloride (NaCl) has given up its one outermost electron to the chlorine atom. In so doing, the sodium atom has changed from the electrically neutral form of Na⁰ to the positively charged sodium ion Na⁺. At the same time, the chlorine atom has changed from the electrically neutral form of Cl⁰ to the negatively charged chloride ion Cl⁻. The oxidation state of the sodium atom has changed from 0 to +1, while the oxidation state of the chlorine atom has changed from 0 to −1. The sodium atom has been oxidized by the chlorine atom, just as the chlorine atom has been reduced by the sodium atom. Some confusion may arise from the terminology at this point, since the atom that was oxidized is labeled the reducing agent, and the atom that was reduced is labeled the oxidizing agent. This can be readily overcome, however, by recalling that Loss of Electrons is Oxidation, Gain of Electrons is Reduction.

The oxidation state of different atoms can seem overly large, ranging from −8 to +8, and one may easily mistake these values for the charge carried on the particular atom. It is important to remember that the oxidation state only represents the number of electrons that the atom has gained or lost control of in its interaction with another atom. It may be helpful to also remember that the oxidation-state concept also applies to covalent compounds in which ions are not involved and electrons are shared between atoms to form chemical bonds. In such compounds, the oxidation states are determined by the theoretical gain or loss of electrons based on how electronegative each atom is.

Assigning Oxidation States

To facilitate understanding and application of the oxidation-state concept, a set of rules is used to assign the oxidation state of any particular atom.

• Rule 1: A hydrogen atom is always assigned an oxidation state of +1.
• Rule 2: An oxygen atom is always assigned an oxidation state of −2.
• Rule 3: A fluorine atom is always assigned an oxidation state of −1.
• Rule 4: Single-atom ions are always assigned an oxidation state that corresponds to their formal charge.

There are exceptions to these rules, but they are covered by other rules for assigning oxidation states. In organic compounds, overall oxidation states are assigned to a particular carbon atom according to the bonds it has formed.

• Rule 5: A carbon atom is assigned an oxidation state of −1 for each bond to a less electronegative atom, 0 for each bond to another carbon atom, and +1 for each bond to a more electronegative atom.
• Rule 6: To complete mass and charge balance, add H₂O and H⁺ in acidic solution or H₂O and OH⁻ in basic solution as needed.

For example, perchloric acid has the chemical formula HClO₄. To balance the four oxygen atoms (each assigned an oxidation state of −2, for a total of −8) and the hydrogen atom (assigned an oxidation state of +1), the central chlorine atom must be assigned a state of +7. Since the chlorine atom is very electronegative and does not give up electrons to form a positively charged ion, it should be readily apparent that the oxidation state of +7 represents something entirely different from the atom’s ionic charge.

Half Reactions

To make the redox concept easier to comprehend, it is useful to break down a reaction to its component "half reactions." A half reaction isolates just the portion of the reaction corresponding to oxidation and reduction. For example, when magnesium metal, Mg, is dissolved in nitric acid, HNO₃, a redox reaction takes place in which aqueous magnesium nitrate, Mg(NO₃)₂, and hydrogen gas, H₂, are produced. The corresponding half reactions are as follows:

Mg⁰ → Mg²⁺ + 2e⁻

H⁺¹ + 1e⁻ → H⁰

These half reactions clearly indicate that the magnesium atom has given up two electrons and the hydrogen atoms have each accepted one electron. By balancing the number of electrons that are transferred and adding together the result, the balanced reaction can be determined. Here, there must be two molecules of HNO₃ to account for the two electrons accepted from the magnesium and the two atoms of hydrogen that are produced. Thus, the balanced equation is:

Mg + 2HNO₃ → Mg(NO₃)₂ + H₂

PRINCIPAL TERMS

• oxidation: the loss of one or more electrons by an atom, ion, or molecule.
• oxidation state: a number that indicates the degree to which an atom or ion in a chemical compound has been oxidized or reduced.
• oxidizing agent: any atom, ion, or molecule that accepts one or more electrons from another atom, ion, or molecule in a reduction-oxidation (redox) reaction and is thus reduced in the process.
• reducing agent: any atom, ion, or molecule that donates one or more electrons to another atom, ion, or molecule in a reduction-oxidation (redox) reaction and is thus oxidized in the process.
• reduction: the gain of one or more electrons by an atom, ion, or molecule.

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