After Chapter 11.1, you will be able to:
Reactions that involve the transfer of electrons from one chemical species to another can be classified as oxidation–reduction (redox) reactions.
The law of conservation of charge states that electrical charge can be neither created nor destroyed. Thus, an isolated loss or gain of electrons cannot occur; oxidation (loss of electrons) and reduction (gain of electrons) must occur simultaneously, resulting in an electron transfer called a redox reaction. An oxidizing agent causes another atom in a redox reaction to undergo oxidation and is itself reduced. A reducing agent causes the other atom to be reduced and is itself oxidized. There are various mnemonics to remember these terms, as highlighted in the sidebar.
Redox reactions: choose one of the mnemonics and stick with it!
Being familiar with some common oxidizing and reducing agents can save significant time on Test Day, especially in organic chemistry reactions. Some of the commonly used agents on the MCAT are listed in Table 11.1. Note that almost all oxidizing agents contain oxygen or another strongly electronegative element (such as a halogen). Reducing agents often contain metal ions or hydrides (H–).
| Oxidizing Agents | Reducing Agents |
|---|---|
| O2 | CO |
| H2O2 | C |
| F2, Cl2, Br2, I2 (halogens) | B2H6 |
| H2SO4 | Sn2+ and other pure metals |
| HNO3 | Hydrazine* |
| NaClO | Zn(Hg)* |
| KMnO4* | Lindlar’s catalyst* |
| CrO3, Na2Cr2O7* | NaBH4* |
| Pyridinium chlorochromate (PCC)* | LiAlH4* |
| NAD+, FADH** | NADH, FADH2** |
| * These oxidizing agents and reducing agents are commonly seen in organic chemistry reactions. ** These and other biochemical redox reagents often act as energy carriers in biochemistry reactions. | |
Note that biochemical redox reagents such as NAD+ tend to act as both oxidizing and reducing agents at different times during metabolic pathways. As such, they act as mediators of energy transfer during many metabolic processes, as shown in Figure 11.1.
On a technical level, the term oxidizing agent or reducing agent is applied specifically to the atom that gains or loses electrons, respectively. However, many science texts will describe the compound as a whole (CrO3, rather than Cr6+) as the oxidizing or reducing agent.
It is important, of course, to know which atom is oxidized and which is reduced. Oxidation numbers are assigned to atoms in order to keep track of the redistribution of electrons during chemical reactions. Based on the oxidation numbers of the reactants and products, it is possible to determine how many electrons are gained or lost by each atom.
In Chapter 3 of MCAT General Chemistry Review, we illustrated that metals form cations and nonmetals form anions. To form a cation, a metal must lose electrons. Therefore, metals like to get oxidized (lose electrons) and act as good reducing agents. Nonmetals, on the other hand, like to get reduced (gain electrons) and act as good oxidizing agents.
The oxidation number of an atom in a compound is assigned according to the following rules:
Think of the oxidation number as the typical charge of an element based on its group number, metallicity, and general location in the periodic table.
The conventions of formula writing put cation first and anion second. Thus HCl implies H+, and NaH implies H–. Use the way the compound is written on the MCAT along with the periodic table to determine oxidation states.
Don’t forget that you can click on “Periodic table” to pull it up on Test Day. Note the trends for assigning oxidation numbers; these general rules will help reduce the need to memorize individual oxidation numbers. Be aware that the transition metals can take on multiple oxidation states and therefore multiple oxidation numbers.
Oxidation number is often confused with formal charge, discussed in Chapter 3 of MCAT General Chemistry Review. Both account for the perceived charge on an element, but do so in different ways. Oxidation number assumes unequal division of electrons in bonds, “awarding” the electrons to the more electronegative element. Formal charge, on the other hand, assumes equal division of electrons in bonds, “awarding” one electron to each atom in the bond. In reality, the distribution of electron density lies somewhere between these two extremes. The assigning of oxidation number can be seen in Figure 11.2.
When assigning oxidation numbers, start with the known atoms (usually Group I and II, halides, and oxygen) and use this information to determine the oxidation states of the other atoms. Keep in mind that most transition metals can take on multiple oxidation states. When transition metals are oxidized or reduced, the absorption and emission of light from a metal is altered such that different frequencies are absorbed. For this reason, changes of oxidation state in transition metals usually correspond to a color change.
By assigning oxidation numbers to the reactants and products, one can determine how many moles of each species are required for conservation of charge and mass, which is necessary to balance the equation. To balance a redox reaction, both the net charge and the number of atoms must be equal on both sides of the equation. The most common method for balancing redox equations is the half-reaction method, also known as the ion–electron method, in which the equation is separated into two half-reactions—the oxidation part and the reduction part. Each half-reaction is balanced separately, and they are then added to give a balanced overall reaction.
Oxidizing agents oxidize other molecules, but are themselves reduced. Reducing agents reduce other molecules, but are themselves oxidized. If you determine one ion to be an oxidizing agent then the other must be a reducing agent.
Methodical, step-by-step approaches like the half-reaction method are great for the MCAT. Usually, you will not have to go through all of these steps before you can narrow down your answer choices and may be able to find the correct answer partway through the problem with a little critical thinking.