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Chapter 14: Problem 3

What is an oxidation-reduction reaction?

Short Answer

Expert verified
An oxidation-reduction or redox reaction is a chemical reaction that involves change in oxidation states of atoms. It entails an oxidation process, where electrons are lost, and a reduction process, where electrons are gained. These processes take place simultaneously, with the substance that loses electrons (i.e., gets oxidized) causing reduction in other substances (thus is called the reducing agent), and the substance that gains electrons (i.e., gets reduced) causing oxidation in other substances (thus is called the oxidizing agent).

Step by step solution

01

Define Oxidation and Reduction

Oxidation is defined as the loss of electrons from an atom, ion or molecule. On the other hand, reduction is defined as the gain of electrons by an atom, ion or molecule. One example to remember this is through the acronym OIL RIG, where 'OIL' stands for Oxidation Is Loss (of electrons) and 'RIG' stands for Reduction Is Gain (of electrons).
02

Describe Redox Reaction

A redox reaction involves both oxidation and reduction processes occurring simultaneously. In a chemical reaction, the substance that loses electrons gets oxidized and is known as the reducing agent (because it causes other substance to be reduced). Likewise, the substance that gains electrons gets reduced and is known as the oxidizing agent (because it causes other substance to be oxidized). Therefore, it can be said that in a redox reaction, the reducing agent gets oxidized and the oxidizing agent gets reduced.
03

Illustrate with an example

Let's consider the chemical reaction between sodium (Na) and chlorine (Cl) to form sodium chloride (NaCl). The reaction can be written as: \(2Na + Cl2 \rightarrow 2NaCl\). This is a redox reaction. Here, sodium (Na) loses one electron to become \(Na^+\) (it is oxidized) and chlorine (Cl) gains one electron to become \(Cl^-\) (it is reduced). Sodium (Na) is the reducing agent because it causes chlorine to be reduced and chlorine (Cl2) is the oxidizing agent because it causes sodium to be oxidized.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Oxidation
Oxidation is a fundamental chemical process that plays a vital role in redox reactions. The term "oxidation" refers to the loss of electrons by an atom, ion, or molecule. When a substance undergoes oxidation, it becomes positively charged because it loses negatively charged electrons.
A helpful way to remember this is through the acronym OIL, which stands for "Oxidation Is Loss." This loss of electrons often results in the substance having a different oxidation state.
  • When iron rusts, it is an example of oxidation. Iron loses electrons to oxygen in the atmosphere to form iron oxide, or rust.
  • Burning, or combustion, is another common example of oxidation. In this case, a substance reacts with oxygen and loses electrons.
Remember, in any redox reaction, oxidation and reduction occur simultaneously.
Reduction
Reduction is the process where an atom, ion, or molecule gains electrons. By gaining electrons, the substance becomes more negatively charged, or less positive. Gaining electrons typically results in a change in the oxidation state of the substance.
Reduction can be neatly summarized by the acronym RIG, which stands for "Reduction Is Gain." This acronym succinctly captures the essence of the reduction process.
  • In the chemical equation where chlorine gas reacts with sodium to form sodium chloride, the chlorine is reduced because it gains electrons becoming chloride ions.
  • Photosynthesis is a biological process where carbon dioxide is reduced to glucose by gaining electrons, highlighting reduction in living organisms.
By understanding reduction, one appreciates the balance it maintains with oxidation in redox reactions.
Redox Reactions
Redox reactions are chemical reactions in which oxidation and reduction occur simultaneously. The term "redox" is a portmanteau of "reduction" and "oxidation," emphasizing their connected nature.
During a redox reaction, one substance loses electrons (is oxidized), while another gains those electrons (is reduced).
  • In a simple redox reaction like the one between hydrogen and oxygen to form water, hydrogen is oxidized, and oxygen is reduced.
  • Redox reactions are not limited to inorganic chemistry鈥攖hey occur in biological systems as well, such as in cellular respiration where glucose is oxidized.
The efficiency and balance of redox reactions enable essential energy transformations in both industrial and natural processes.
Oxidizing Agents
An oxidizing agent is a substance that makes the process of oxidation happen to another substance in a redox reaction. It achieves this by accepting electrons from the atom or molecule being oxidized.
Because it gains electrons, an oxidizing agent is reduced in the process.
  • Chlorine is a well-known oxidizing agent because it readily accepts electrons to form chloride ions.
  • In industrial processes, oxygen often acts as an oxidizing agent, facilitating combustion.
Understanding oxidizing agents is crucial because they dictate what gets oxidized in a redox reaction.
Reducing Agents
A reducing agent is a substance that donates electrons to another substance in a redox reaction, facilitating its reduction. When a reducing agent donates electrons, it undergoes oxidation.
This donation of electrons makes the reducing agent a key player in the completion of redox reactions.
  • Sodium in the sodium-chlorine reaction acts as a reducing agent by losing electrons, allowing chlorine to be reduced.
  • In organic chemistry, compounds like hydrides can serve as reducing agents, donating electrons to reduce carbonyl groups into alcohols.
Reducing agents are essential for countless chemical reactions, powering everything from basic industry to advanced biochemistry.

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Most popular questions from this chapter

In the following oxidation-reduction reactions, identify the oxidizing agent, the reducing agent, and the number of electrons transferred. (a) \(\mathrm{I}_{2}(a q)+2 \mathrm{OH}^{-}(a q) \longrightarrow \mathrm{I}^{-}(a q)+\mathrm{IO}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\) (b) \(\mathrm{Cr}(s)+2 \mathrm{H}^{+}(a q) \longrightarrow \mathrm{Cr}^{2+}(a q)+\mathrm{H}_{2}(g)\) (c) \(2 \mathrm{Cr}_{2} \mathrm{O}_{7}{ }^{2}(a q)+16 \mathrm{H}^{+}(a q) \longrightarrow\) \(4 \mathrm{Cr}^{3+}(a q)+3 \mathrm{O}_{2}(g)+8 \mathrm{H}_{2} \mathrm{O}(l)\) (d) \(3 \mathrm{Fe}^{3+}(a q)+\mathrm{Al}(s) \longrightarrow 3 \mathrm{Fe}^{2 *}(a q)+\mathrm{Al}^{34}(a q)\)

Balance the following halt-reactions, adding \(\mathrm{OH}^{-}(a q)\), \(\mathrm{H}_{2} \mathrm{O}(l)\), and electrons as appropriate. (a) \(\mathrm{CrO}_{4}{ }^{2}(a q) \longrightarrow \mathrm{Cr}(\mathrm{OH})_{3}(s)\) (b) \(\mathrm{ClO}_{2}(a q) \longrightarrow \mathrm{ClO}^{-}(a q)\) (c) \(\mathrm{MnO}_{4}^{-}(a q) \longrightarrow \mathrm{MnO}_{2}(s)\) (d) \(\mathrm{Br}^{-}(a q) \longrightarrow \mathrm{BrO}^{-}(a q)\)

Determine the oxidation number of chlorine in each of these chlorine oxides. (a) \(\mathrm{Cl}_{2} \mathrm{O}\) (d) \(\mathrm{ClO}_{2}\) (b) \(\mathrm{ClO}_{3}\) (e) \(\mathrm{Cl}_{2} \mathrm{O}_{5}\) (c) \(\mathrm{Cl}_{2} \mathrm{O}_{3}\)

In the following oxidation-reduction reactions, identify the oxidizing agent, the reducing agent, and the number of electrons transferred. (a) \(6 \mathrm{I}^{-}(a q)+\mathrm{BrO}_{3}^{-}(a q)+6 \mathrm{H}^{*}(a q) \longrightarrow\) \(3 \mathrm{I}_{2}(a q)+\mathrm{Br}^{-}(a q)+3 \mathrm{H}_{2} \mathrm{O}(l)\) (b) \(5 \mathrm{Br}^{-}(a q)+\mathrm{BrO}_{3}^{-}(a q)+6 \mathrm{H}^{*}(a q) \underset{3 \mathrm{Br}_{2}(a q)+3 \mathrm{H}_{2} \mathrm{O}(l)}{\longrightarrow}\) (c) \(\mathrm{XeF}_{4}(s)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Xe}(g)+4 \mathrm{HF}(a q)+\mathrm{O}_{2}(g)\) (d) \(\mathrm{Pb}(s)+2 \mathrm{H}^{+}(a q) \longrightarrow \mathrm{Pb}^{2+}(a q)+\mathrm{H}_{2}(g)\)

Consider the following reaction: $$ 6 \mathrm{~V}^{24}(a q)+\mathrm{Cr}_{2} \mathrm{O}_{7}{ }^{2}(a q)+14 \mathrm{H}^{4}(a q) \longrightarrow 6 \mathrm{~V}^{34}(a q)+2 \mathrm{Cr}^{34}(a q)+7 \mathrm{H}_{2} \mathrm{O}(l) $$ (a) Which species is oxidized? (b) Which species is reduced? (c) What is the oxidizing agent? (d) What is the reducing agent?
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