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Lesson 1: Redox Reactions
Part a: Oxidation and Reduction
Part a: Oxidation and Reduction
Part b:
Oxidation States
Part c:
Balancing Redox Reactions
Part d:
Corrosion
The Big Idea
Understanding oxidation and reduction reactions is essential in electrochemistry. Learn how electrons are transferred and why some substances act as oxidizing or reducing agents.
Electron Transfer Reactions
Chapter 15 of our Chemistry Tutorial involved a study of acid-base reactions. A central component of our model in that chapter was that acid-base reactions involved the transfer of a proton from one substance to another substance. One substance lost a proton while another substance gained the proton.
In Chapter 18, we will build a similar model for a type of reaction known as an oxidation-reduction reaction (abbreviated as redox). A redox reaction involves the transfer of an electron from one substance to another substance. All reactions in this chapter are electron-transfer reactions. One element will lose one or more electrons while another element will gain those that are lost.
Oxidation-reduction reactions are all around us. These include:
- the corrosion of iron and other metals,
- the burning of fossil fuels,
- the reaction of metals with acids,
- photosynthesis by plants, and
- cellular respiration in the human body.
Reactions which we previously classified as
combustion,
single replacement,
synthesis, and
decomposition usually involve electron transfer.
A Detailed Example of a Redox Reaction
When a strip of magnesium (Mg) metal is added to an aqueous solution hydrochloric acid (HCl), evidence of a reaction is quickly observed. Bubbles are formed and the size of the magnesium strip quickly decreases. The reaction that occurs is represented by the equation
Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g)
The above equation is the molecular equation representing the reaction. We know that both the HCl (a strong, completely dissociated acid) and the MgCl
2 (a soluble salt) exist as ions in the solution. We can represent the reaction with a
complete ionic equation.
Mg(s) + 2 H+(aq) + 2 Cl-(aq) → Mg2+(aq) + 2 Cl-(aq) + H2(g)
We observe that the Cl
- ion is a spectator ion that is not involved in the reaction. Writing the
net ionic equation will help us to see the electron transfer nature of the reaction.
Mg(s) + 2 H+(aq) → Mg2+(aq) + H2(g)
It is clear from the net ionic equation that a magnesium atom - Mg
(s) - is turning into a magnesium ion - Mg
2+(aq). The only way a neutral atom can turn into a 2+ ion is by losing two electrons. We can represent this by the equation:
Mg(s) → Mg2+(aq) + 2 e-
We can also make a claim that two hydrogen ions - 2 H
+(aq) - are changing into a diatomic hydrogen molecule - H
2(g). The only way two 1+ ions can turn into a neutral molecule is by gaining two electrons. We can represent this by the equation:
Oxidation and Reduction
We refer to these two changes as
half-reactions. Each represents one-half of the whole change that occurs. The Mg atom is said to have undergone oxidation.
Oxidation is defined as the loss of electrons. The H
+ ions are said to have undergone reduction.
Reduction is the gain of electrons.
Oxidation: Mg(s) → Mg2+(aq) + 2 e-
Reduction 2 H+(aq) + 2 e- → H2(g)
Two electrons are lost by the magnesium atom and transferred to the hydrogen ions. The oxidation and the reduction must occur together. One cannot happen without the other. The electrons that are lost by one element are the ones that are gained by the other element. Rather than viewing these as two separate events, they should be viewed as two dependent events.
A useful mnemonic for remembering the meaning of the terms oxidation and reduction is
OIL RIG. The OIL is short for oxidation is loss and the RIG is short for reduction is gain. The oxidation half-reaction involves an element losing electrons and the reduction half-reaction involves an element gaining electrons.
Finally, it is instructive to state that no electrons are created nor destroyed in this oxidation-reduction process. Electrons are simply being taken from one location (the element being oxidized) and transferred to a different location (the element being reduced).It is an electron transfer.
LEO the Lion says GER
Need a different mnemonic for remembering the meaning of oxidation and reduction? Try this one, named after our favorite lion Leo. You know what Leo says? "Ger".
LEO =
Loss of
Electrons is
Oxidation
GER =
Gain of
Electrons is
Reduction
Oxidizing and Reducing Agents
The reactants in a redox reaction are categorized as reducing agents and oxidizing agents. The
reducing agent is the reactant that undergoes oxidation. It loses electrons and transfers those electrons to the substance that is being reduced. The
oxidizing agent is the substance that is being reduced. It gains electrons during the process, receiving them from the substance that is being oxidized.
In the provided example, the Mg is being oxidized; it is the reducing agent. By losing electrons, Mg makes reduction possible. The H
+ ion is being reduced; it is the oxidizing agent. It’s willingness to gain or accept electrons makes the oxidation possible.
Another Redox Examples: Reaction of Na with Cl2
Let’s consider a second oxidation-reduction example - the reaction of sodium metal with chlorine gas. The balanced molecular equation is:
2 Na(s) + Cl2(g) → 2 NaCl(s)
The reactants are two neutral substances. Neither has a charge. The product is an ionic compound. We understand ionic compounds as consisting of two oppositely charged ions. In this case, NaCl consists of the Na
+ ion and the Cl
- ion. The two changes that occur in this reaction are
- The oxidation of Na to the Na+ ion
- The reduction of Cl2 to the Cl- ion.
The two half-reaction can be written as:
Oxidation: Na → Na+ + e-
Reduction Cl2 + 2 e- → 2 Cl-
The sodium metal is oxidized; it is the reducing agent. It helps the Cl
2 be reduced by donating two electrons to it. Upon receiving the two electrons, the Cl
2 is reduced, becoming Cl
- ion. Thus, Cl
2 is the oxidizing agent. The ionic bond between the oppositely charged ions results in the formation of NaCl
(s).
Next Up
The two examples on this page have been among the easier examples. In
the next part of Lesson 1, we will learn a strategy for identifying the half reactions, the oxidizing agents, and the reducing agents for the most difficult of situations.
Before You Leave - Practice and Reinforcement
Now that you've done the reading, take some time to strengthen your understanding and to put the ideas into practice. Here's some suggestions.
- The Check Your Understanding section below includes questions with answers and explanations. It provides a great chance to self-assess your understanding.
- Download our Study Card on Oxidation and Reduction. Save it to a safe location and use it as a review tool. (Coming Soon)
Check Your Understanding of Oxidation and Reduction
Use the following questions to assess your understanding of concepts related to oxidation and reduction. Tap the
Check Answer buttons when ready.
1. Complete the following sentence:
An atom of a metallic element is most likely to ________ (gain, lose, destroy, create) electrons. It will undergo ____________ (oxidation, reduction, hydrolysis, neutralization).
2. A double replacement reaction is represented by the following net ionic equation.
2 Al(s) + 3 Cu2+(aq) → 2 Al3+(aq) + 3 Cu(s)
a. Identify the element that is being oxidized.
b. Identify the element that is being reduced.
c. Identify the oxidizing agent.
d. Identify the reducing agent.
e. Write the equations for the two half reactions and label them as oxidation or reduction.
3. The equations for two half reactions are shown below. Identify each as being either an oxidation or a reduction half reaction:
a. Ni → Ni
2+ + 2 e
-
b. F
2 + 2 e
- → 2 F
-