Gatewaylectrochemistry: Oxidation and Reduction


Table of Contents

Icon Legend
Ox & Red


Voltaic Cells

Cell Voltage

Calc. Potentials


Electrolytic Cells

Common Oxidation States

When iron is burned in pure oxygen, a brown soot is formed. This reaction can be described with the chemical equation shown below. Examine the oxidation states of the reactants with the mouse pointer and then answer the questions that follow.

roll over to select; 'click' here to reset.Roll over equation to see oxidation states of atoms

oxidation of iron

What is the oxidizing agent in the above redox reaction?

Fe O2 

What is the reducing agent in the above redox reaction?

Fe O2 


An oxidizing agent is itself reduced. Does iron gain electrons in this reaction?

A reducing agent is itself oxidized. Does oxygen lose electrons in this reaction?

Is there a relationship between whether the reactant is an oxidizing or reducing agent and the change, if any, in its oxidation state? As previously mentioned, oxidizing agents are reduced. In this reaction, oxygen
goes from {0 2–}; oxygen is therefore the oxidizing agent, not a surprise. The oxidation state of iron goes from {0 3+}; iron is therefore the reducing agent.

Note: knowing that oxygen, a common oxidizing agent, is typically reduced to {2–} allows us to deduce the oxidation state of iron in the product to be {3+} (from Fe2O3).

Figure 3

roll over to select; 'click' here to reset.Roll over groups to see oxidation states; click on mouse to reset.

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Being able to work backwards from known oxidation states is a powerful skill. We must sometimes "follow" the oxidation states of the reactants if we are to determine if they are agents of oxidation or reduction—or if a redox reaction has occurred at all! Knowing the common oxidation states of those substances listed in Figure 3 is necessary if we are to dissect more complex oxidation reactions.

Consider the following half-reaction:
 4 H+ +  MnO4   MnO2  +  2 H20

It may not be readily apparent whether MnO4 is acting as a reducing or an oxidizing agent. However, since we know the oxidation state of oxygen in both MnO4 and MnO2, we can determine the change in the oxidation state of managnese. If the oxidation state becomes more positive, then it must be transferring electrons to something else. It is a reducing agent. If, on the other hand, the oxidation state is becoming more negative, then it is accepting electrons from something else. It is an oxidizing agent.

Is manganese (in the permanganate ion) acting as an oxidizing or reducing agent in this reaction?

Oxidizing Agent Reducing Agent

The oxidation state of oxygen in both MnO42- and MnO2 is -2. What is are the oxidation states of manganese in these two substances?

Good! The oxidation state of manganese becomes more negative (from +7 to +4) so it is acting as an oxidizing agent.

The Periodic TableNote: there are periodic trends associated with preferred oxidation states. It would be useful to compare the states listed in Figure 3. to the periodic table in your textbook.There is also a link to a simple table on the right.

Before leaving this module, you should know how to identify oxidizing agents or reducing agents by their change in oxidation states. To that end, you should also commit to memory the preferred oxidation states of some common reactants. Please activate the debriefing before returning to the home page. After sucessfully completing the debriefing, you will be given the opportunity to move to the next module.

'click' here to engage the debriefing.

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