Why does oxidation and reduction occur simultaneously

In H 2 O 2 , oxygen has an oxidation state of In H 2 O, its oxidation state is -2, and it has been reduced. In O 2 however, its oxidation state is 0, and it has been oxidized.

Oxygen has been both oxidized and reduced in the reaction, making this a disproportionation reaction. The general form for this reaction is as follows:. Every balanced redox reaction is composed of two half-reactions: the oxidation half-reaction, and the reduction half-reaction. For example, look at the following redox reaction between iron and copper:. In this reaction, iron is oxidized, and copper is reduced or, iron is the reducing agent , and copper is the oxidizing agent.

We can split this reaction into two half-reactions. The oxidation half-reaction looks as follows:. This shows the oxidation of iron and the loss of two electrons. Notice that this equation is balanced in both mass and charge: we have one atom of iron on each side of the equation mass is balanced , and the net charge on each side of the equation is equal to zero charge is balanced.

This half-reaction explicitly shows the copper II ion gaining two electrons. Note again that the equation is balanced in mass and charge. Now that we have our two balanced half-reactions, we can combine them to get the full redox reaction:.

Adding the two halves of a redox reaction : These two halves of the reaction can be added like any other chemical equation.

Once the equations are added, the electrons on each side cancel out. Note that the two electrons on each side of the equation cancel out. This is very important, because the final balanced equation for any redox reaction should never contain any electrons. First, we need to split this reaction into its two half-reactions.

We need to balance this equation by mass. The equation is already balanced in nitrogens, but not oxygens. You can balance oxygen by adding the appropriate number of water molecules:. Now the equation is balanced in oxygens, but not hydrogens. The equation is now balanced in mass, but not charge. To balance the charge, we will add two electrons to the right side of the equation:. The equation is now balanced for mass, and we need only balance for charge. We therefore need to add 6 electrons to the right side of the equation to balance the charges:.

Lastly, in order to get our full balanced redox equation, we need to add our half-reactions so that all the electrons cancel out. For this reaction, we can multiply the first half-reaction by Although this example seems intimidating, balancing redox reactions in acidic solution becomes much easier with careful practice.

For instance:. The half-reaction above is balanced for mass in acidic solution. If we are in basic solution, however, we would need to add 2 hydroxides to both sides of the equation:. These species will neutralize each other to form water, so we can rewrite this as follows:. Lastly, because we have water molecules on both sides of the equation, we cancel out like terms to give us:. This half-reaction is now balanced for mass in basic solution.

From here, we proceed just as we did above in acidic solution: balance the charge by adding the appropriate number of electrons. Balancing redox equations can certainly be complicated and time-consuming, so it is wise to practice them extensively. Redox titration determines the concentration of an analyte containing either an oxidizing or a reducing agent. As with acid-base titrations, a redox titration also called an oxidation- reduction titration can accurately determine the concentration of an unknown analyte by measuring it against a standardized titrant.

In reality, electrons are lost by some atoms and gained by other atoms simultaneously. However, mentally we can separate the two processes. Oxidation is defined as the loss of one or more electrons by an atom. Reduction is defined as the gain of one or more electrons by an atom. In reality, oxidation and reduction always occur together; it is only mentally that we can separate them. Chemical reactions that involve the transfer of electrons are called oxidation-reduction or redox reactions.

Redox reactions require that we keep track of the electrons assigned to each atom in a chemical reaction. How do we do that? We use an artificial count called the oxidation number to keep track of electrons in atoms. Oxidation numbers are assigned to atoms based on a series of rules. Oxidation numbers are not necessarily equal to the charge on the atom; we must keep the concepts of charge and oxidation numbers separate. Let us work through a few examples for practice. In H 2 , both hydrogen atoms have an oxidation number of 0, by rule 1.

We can use rule 4 to determine oxidation numbers for the atoms in SO 2. All redox reactions occur with a simultaneous change in the oxidation numbers of some atoms. At least two elements must change their oxidation numbers.

When an oxidation number of an atom is increased in the course of a redox reaction, that atom is being oxidized. When an oxidation number of an atom is decreased in the course of a redox reaction, that atom is being reduced.

Oxidation and reduction are thus also defined in terms of increasing or decreasing oxidation numbers, respectively. Consider the reactants. Because both reactants are the elemental forms of their atoms, the Na and Cl atoms as reactants have oxidation numbers of 0.

Both reactants are the elemental forms of their atoms, so the Na and Br atoms have oxidation numbers of 0. Because oxidation numbers are changing, this is a redox reaction. The total number of electrons being lost by sodium two, one lost from each Na atom is gained by bromine two, one gained for each Br atom.

Oxidation reactions can become quite complex, as attested by the following redox reaction:. To demonstrate that this is a redox reaction, the oxidation numbers of the species being oxidized and reduced are listed; can you determine what is being oxidized and what is being reduced?

This is also an example of a net ionic reaction; spectator ions that do not change oxidation numbers are not displayed in the equation. Eventually, we will need to learn techniques for writing correct i. Iron is an essential mineral in our diet; iron-containing compounds like the heme protein in hemoglobin could not function without it. To ensure that we ingest enough iron, many foods are enriched with iron. Reduced iron is simply iron metal; iron is added as a fine metallic powder.

Although it is difficult to establish conclusive reasons, a search of scientific and medical literature suggests a few reasons. One reason is that fine iron filings do not affect the taste of the product. Finally, of the common iron substances that might be used, metallic iron is the least expensive. These factors appear to be among the reasons why metallic iron is the supplement of choice in some foods. Learning Objectives Define oxidation and reduction.

Assign oxidation numbers to atoms in simple compounds. Recognize a reaction as an oxidation-reduction reaction.