Oxidation: Involves the loss of electrons and an increase in the oxidation number.
Reduction: Involves the gain of electrons and a decrease in the oxidation number.
Identifying oxidizing and reducing agents requires determining how oxidation numbers change in a reaction.
Rules for Oxidation States
The oxidation state of any pure element in its natural state is zero.
In compounds, group 2 metals like magnesium typically have a +2 oxidation state.
Halogens like bromine typically have a -1 charge in compounds.
Example Analyses
Example 1: Magnesium and Bromine
Mg (Magnesium): Oxidation state goes from 0 to +2, indicating oxidation.
Br (Bromine): Oxidation state goes from 0 to -1, indicating reduction.
Agents:
Magnesium is the reducing agent (causes reduction of bromine).
Bromine is the oxidizing agent (causes oxidation of magnesium).
Example 2: Zinc and Hydrochloric Acid
Zn (Zinc): Oxidation state goes from 0 to +2, indicating it was oxidized.
H (Hydrogen): Oxidation state goes from +1 to 0, indicating it was reduced.
Agents:
Zinc is the reducing agent (causes reduction of hydrochloric acid).
Hydrochloric acid is the oxidizing agent (causes oxidation of zinc).
Example 3: Sodium Bromide and Chlorine Gas
NaBr (Sodium Bromide):
Sodium is a spectator ion.
Bromine oxidation state increases from -1 to 0, so it is oxidized.
Cl2 (Chlorine Gas): Oxidation state decreases from 0 to -1, indicating reduction.
Agents:
Sodium bromide is the reducing agent.
Chlorine is the oxidizing agent.
Additional Notes
Metals as Reducing Agents: Many metals like zinc and magnesium are good reducing agents.
Nonmetals as Oxidizing Agents: Nonmetals such as chlorine, bromine, and oxygen often act as oxidizing agents.
Special Cases for Oxygen:
Typically has a -2 oxidation state, except when bonded to fluorine or in peroxides/superoxides.
Final Example: Chlorite and Perchlorate
ClO2- (Chlorite): Chlorine has an oxidation state of +3.
ClO4- (Perchlorate): Chlorine has an oxidation state of +7.
ClO3- (Chlorate): Chlorine has an oxidation state of +5.
Changes:
Chlorite (ClO2-) is oxidized to chlorate (from +3 to +5).
Perchlorate (ClO4-) is reduced to chlorate (from +7 to +5).
Agents:
Chlorite is the reducing agent.
Perchlorate is the oxidizing agent.
Conclusion
Understanding the roles of substances as oxidizing and reducing agents depends on their ability to gain or lose electrons, thus affecting the oxidation states of the elements involved in a chemical reaction.