From each of the following pairs of substances, use data in Appendix E to
choose the one that is the stronger reducing agent:
(a) Fe(s) or \(\mathrm{Mg}(s)\)
(b) \(\mathrm{Ca}(s)\) or \(\mathrm{Al}(s)\)
(c) \(\mathrm{H}_{2}\) (g, acidic solution) or \(\mathrm{H}_{2}
\mathrm{~S}(\mathrm{~g})\)
(d) \(\mathrm{BrO}_{3}^{-}(a q)\) or \(\mathrm{lO}_{3}^{-}(a q)\)
20.44 From each of the following pairs of substances, use data in Appendix E
to choose the one that is the stronger oxidizing agent:
(a) \(\mathrm{Cl}_{2}(g)\) or \(\mathrm{Br}_{2}(l)\)
(b) \(\mathrm{Zn}^{2+}(a q)\) or \(\operatorname{Cd}^{2+}(a q)\)
(c) \(\mathrm{Cl}^{-}(a q)\) or \(\mathrm{ClO}_{3}^{-}(a q)\)
(d) \(\mathrm{H}_{2} \mathrm{O}_{2}(a q)\) or \(\mathrm{O}_{2}(\mathrm{~g})\)
20.45 By using the data in Appendix E, determine whether each of the following
substances is likely to serve as an exidant or a reductant: (a)
\(\mathrm{Cl}_{2}\) (g), (b) \(\mathrm{MnO}_{4}^{-}\)(aq, acidic solution), (c)
\(\mathrm{Ba}\) (s), (d) \(\mathrm{Zn}(\) s).
20.46 Is each of the following substances likely to serve as an oxidant or a
reductant: (a) \(\mathrm{Ce}^{3+}(\mathrm{aq})\), (b) \(\mathrm{Ca}(\mathrm{s})\),
(c) \(\mathrm{CO}_{3}^{-}(\mathrm{aq})\), (d) \(\mathrm{N}_{2} \mathrm{O}_{5}(g)\)
?
\(20.47\) (a) Assuming standard conditions, arrange the following in order of
increasing strength as oxidizing agents in acidic solution: \(\mathrm{Cr}_{2}
\mathrm{O}_{7}^{2-}, \mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{Cu}^{2+},
\mathrm{C}_{2}, \mathrm{O}_{2}\) (b) Arrange the following in order of
increasing strength as reducing agents in acidic solution: \(\mathrm{Zn},
\mathrm{I}^{-}, \mathrm{Sn}^{2+}, \mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{AL}\).
20.48 Based on the data in Appendix E, (a) which of the following is the
strongest oxidizing agent and which is the weakest in acidic solution:
\(\mathrm{Br}_{2}, \mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{Zn}, \mathrm{Cr}_{2}
\mathrm{O}_{7}{ }^{2-} ?\) (b) Which of the following is the strongest reducing
agent, and which is the weaket in acidic solution: \(\mathrm{F}^{-},
\mathrm{Zn}, \mathrm{N}_{1}{ }^{+}\), \(\mathrm{I}_{\mathrm{n}} \mathrm{NO}\) ?
The stronger reducing agents from the given pairs are as follows:
(a) Mg(s)
(b) Ca(s)
(c) H₂(g, acidic solution)
The stronger oxidizing agent in the last pair is:
(d) BrO₃⁻(aq)
Step by step solution
01
Identify reduction potentials
From Appendix E, we have the standard reduction potentials for these substances:
Fe: \(Fe^{2+}(aq) + 2e^- → Fe(s)\); E° = -0.44 V
Mg: \(Mg^{2+}(aq) + 2e^- → Mg(s)\); E° = -2.37 V
02
Compare reduction potentials
Mg(s) has a more negative reduction potential, which means it is a stronger reducing agent compared to Fe(s).
Stronger reducing agent: Mg(s)
(b) Ca(s) or Al(s)
03
Identify reduction potentials
From Appendix E, we have the standard reduction potentials for these substances:
Ca: \(Ca^{2+}(aq) + 2e^- → Ca(s)\); E° = -2.87 V
Al: \(Al^{3+}(aq) + 3e^- → Al(s)\); E° = -1.66 V
04
Compare reduction potentials
Ca(s) has a more negative reduction potential, which means it is a stronger reducing agent compared to Al(s).
Stronger reducing agent: Ca(s)
(c) H₂ (g, acidic solution) or H₂S(g)
05
Identify reduction potentials
From Appendix E, we have the standard reduction potentials for these substances:
H₂: \(2H^{+}(aq) + 2e^- → H_2(g)\); E° = 0.00 V
H₂S: \(2H^{+}(aq) + 2e^- + H_2S(g) → 2H_2O(l) + S(s)\); E° = 0.14 V
06
Compare reduction potentials
In this case H₂ has a more negative reduction potential compared to H₂S, meaning H₂ is a better reducing agent.
Stronger reducing agent: H₂(g, acidic solution)
(d) BrO₃⁻(aq) or IO₃⁻(aq)
07
Identify reduction potentials
From Appendix E, we have the standard reduction potentials for these substances:
BrO₃⁻: \(2BrO_3^-(aq) + 12H^{+}(aq) + 10 e^- → Br_2(aq) + 6 H_2O(l)\); E° = 1.52 V
IO₃⁻: \(2IO_3^-(aq) + 12H^{+}(aq) + 10 e^- → I_2(aq) + 6 H_2O(l)\); E° = 0.77 V
08
Compare reduction potentials
Due to the more positive reduction potential, BrO₃⁻ is a better oxidizing agent compared to IO₃⁻.
Stronger oxidizing agent: BrO₃⁻(aq)
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Standard Reduction Potential
Understanding the standard reduction potential is crucial in identifying stronger reducing agents in chemical reactions. The standard reduction potential (E°) is a measure of the tendency of a chemical species to acquire electrons and be reduced. These values are measured at standard conditions, which include a temperature of 298 K, a 1 M concentration for each ion participating in the reaction, and a pressure of 1 atm for gases.
For the selection of a stronger reducing agent, one needs to consider the substance with the more negative E° value. This is because a more negative value indicates a greater tendency to donate electrons. Thus, when comparing substances like Fe(s) and Mg(s), Mg(s) emerges as a stronger reducing agent because of its lower E° of -2.37 V compared to -0.44 V for Fe(s).
It is crucial for students to understand that the more negative the standard reduction potential, the stronger the reducing power of the element. This principle is applied consistently across various elements and compounds to determine chemical reactivity in redox reactions.
Oxidation-Reduction Reactions
Oxidation-reduction reactions, also known as redox reactions, are processes that involve the transfer of electrons between two species. Oxidation is the loss of electrons, while reduction is the gain of electrons. In these reactions, the substance that loses electrons is called the reducing agent, and the substance that gains electrons is called the oxidizing agent.
Oxidation and reduction always occur simultaneously, as one species cannot lose electrons without another gaining them. The strength of oxidizing and reducing agents can be evaluated by looking at standard reduction potentials. For instance, in a pair such as H₂ (g, acidic solution) and H₂S(g), H₂ is the better reducing agent as it has a more negative reduction potential (E° = 0.00 V) compared to H₂S (E° = 0.14 V).
When you grasp the concept of electron transfer in redox reactions, you can analyze and predict the outcomes of chemical reactions, which is an essential skill in chemistry.
Chemical Reactivity
Chemical reactivity refers to the propensity of chemical substances to undergo chemical change. The reactivity of a substance can be influenced by various factors including its standard reduction potential. Compounds with high reactivity are more likely to undergo chemical transformations that release or acquire electrons.
In the context of redox reactions, a substance that readily donates electrons, having a low (negative) standard reduction potential, is considered chemically reactive as a reducing agent. Conversely, a substance with a high (positive) standard reduction potential tends to accept electrons, making it a reactive oxidizing agent.
For example, Ca(s) with an E° of -2.87 V is more reactive as a reducing agent than Al(s) with an E° of -1.66 V. This indicates that calcium metal will more readily lose electrons in a redox reaction, emphasizing the connection between chemical reactivity and reduction potentials.
Galvanic Cells
Galvanic cells, also known as voltaic cells, are devices that convert chemical energy into electrical energy through spontaneous redox reactions. They consist of two half-cells, each containing an electrode and an electrolyte. When these half-cells are connected, the flow of electrons from the stronger reducing agent to the stronger oxidizing agent through the external circuit generates electric current.
Standard reduction potentials are integral to understanding how galvanic cells work, as the difference in these potentials between the two half-cells determines the voltage of the cell. The greater the difference in reduction potential, the higher the cell's voltage. In the classroom setting, students can consider a galvanic cell made of Mg(s) and Cu²⁺ as an example. Magnesium's lower reduction potential makes it an anode where oxidation occurs, while copper's higher reduction potential makes it a cathode where reduction happens.
The principles of galvanic cells not only have educational value but also practical applications, such as in batteries and corrosion prevention, making them an essential aspect of electrochemistry.
