Question

(a) Assuming standard conditions, arrange the following in order of increasing strength as oxidizing agents in acidic solution: \(\mathrm{Cr}_{2} \mathrm{O}_{7}{ }^{2-}, \mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{Cu}^{2+}, \mathrm{Cl}_{2}, \mathrm{O}_{2} .\) (b) Arrange the following in order of increasing strength as reducing agents in acidic solution: \(\mathrm{Zn}, \mathrm{I}^{-}, \mathrm{Sn}^{2+}, \mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{Al}\).

Short answer

The order of increasing strength as oxidizing agents is: H2O2 < Cu^2+ < O2 < Cr2O7^2- < Cl2. The order of increasing strength as reducing agents is: 2H2O < 2I^- < Sn^2+ < Zn < Al.

Step-by-step solution

Obtain standard reduction potentials from a table

Look up the standard reduction potentials of Cr2O7^2-, H2O2, Cu^2+, Cl2, and O2 in a table of standard reduction potentials. Using a table of standard reduction potentials, the following values are found: Cr2O7^2-(aq) + 4H2O(l) + 6e^- ⟶ 2Cr^3+(aq) + 7H2O(l) E° = \(1.33\:V\) O2(g) + 4H^+(aq) + 4e^- ⟶ 2H2O(l) E° = \(1.23\:V\) Cu^2+(aq) + 2e^- ⟶ Cu(s) E° = \(0.34\:V\) 2H2O2(aq) ⟶ 2H2O(l) + O2(g) + 2e^- E° = \(-0.68\:V\) Cl2(g) + 2e^- ⟶ 2Cl^-(aq) E° = \(1.36\:V\)

Rank the species by their standard reduction potentials

Sort the standard reduction potentials from lowest to highest to rank the species in order of increasing strength as oxidizing agents. The sorted list in order of increasing strength as oxidizing agents is: H2O2 ⟶ H2O + O2 + 2e^-(E° = -0.68V) < Cu^2+ + 2e^- ⟶ Cu (E° = 0.34V) < O2 + 4H^+ + 4e^- ⟶ 2H2O (E° = 1.23V) < Cr2O7^2- + 14H^+ + 6e^- ⟶ 2Cr^3+ + 7H2O (E° = 1.33V) < Cl2 + 2e^- ⟶ 2Cl^- (E° = 1.36V) (b) To rank the given species in order of increasing strength as reducing agents, we will:

Obtain standard reduction potentials from a table

Look up the standard reduction potentials of Zn, I^-, Sn^2+, H2O2, and Al in a table of standard reduction potentials. Using a table of standard reduction potentials, the following values are found: Al^3+(aq) + 3e^- ⟶ Al(s) E° = \(-1.66\:V\) Zn^2+(aq) + 2e^- ⟶ Zn(s) E° = \(-0.76\:V\) H2O2(aq) + 2H^+(aq) + 2e^- ⟶ 2H2O(l) E° = \(1.78\:V\) Sn^4+(aq) + 2e^- ⟶ Sn^2+(aq) E° = \(0.15\:V\) I2(s) + 2e^- ⟶ 2I^-(aq) E° = \(0.54\:V\)

Invert the standard reduction potentials for reducing agents

The inverted standard reduction potentials will now represent the reducing agents' tendency to lose electrons. Al(s) ⟶ Al^3+(aq) + 3e^- E° = \(1.66\:V\) Zn(s)⟶ Zn^2+(aq) + 2e^- E° = \(0.76\:V\) 2H2O(l) ⟶ H2O2(aq) + 2H^+(aq) + 2e^- E° = \(-1.78\:V\) Sn^2+(aq) ⟶ Sn^4+(aq) + 2e^- E° = \(-0.15\:V\) 2I^-(aq) ⟶ I2(s) + 2e^- E° = \(-0.54\:V\)

Rank the species by their inverted standard reduction potentials

Sort the inverted standard reduction potentials from lowest to highest to rank the species in order of increasing strength as reducing agents. The sorted list in order of increasing strength as reducing agents is: 2H2O ⟶ H2O2 + 2H^+ + 2e^-(E° = -1.78V) < 2I^- ⟶ I2 + 2e^-(E° = -0.54V) < Sn^2+ ⟶ Sn^4+ + 2e^-(E° = -0.15V) < Zn ⟶ Zn^2+ + 2e^-(E° = 0.76V) < Al ⟶ Al^3+ + 3e^-(E° = 1.66V)

Key concepts

Standard Reduction Potentials

Understanding standard reduction potentials is crucial when exploring reactivity in electrochemical processes. These potentials, often denoted as E° values, provide insights into the tendency of species to gain electrons and thus reduce. In essence, they are a measure of the driving force behind a reduction reaction.

Higher reduction potential values indicate a greater tendency to undergo reduction, making such species strong oxidizing agents. Conversely, substances with lower or negative reduction potentials are less inclined to be reduced, instead, they are more likely to act as reducing agents, thereby losing electrons. The values are typically measured under 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 bar for any gases involved.

Listed in increasing order of their standard reduction potentials, we can predict which substances are stronger or weaker oxidizing agents. For example, in our exercise, the standard reduction potential for Cl₂ (1.36 V) is higher than that for H₂O₂ (-0.68 V), indicating that chlorine gas is a stronger oxidizing agent than hydrogen peroxide in an acidic solution.

When solving homework problems or studying electrochemical series, always refer to standardized tables to look up these values accurately. They form the foundation for comparisons and predictions in redox chemistry.

Electrochemistry

Electrochemistry bridges chemical reactivity with electrical energy, focusing on the interplay between electrons and chemical species in redox reactions. The heart of electrochemistry lies in the understanding of how electrons transfer between reactants, defining their oxidation states. It encompasses everything from batteries to corrosion and even biological energy systems.

In our textbook exercise, one facet of electrochemistry under inspection is the behavior of agents in their oxidized and reduced forms. It involves applying the previously mentioned standard reduction potential values to determine the strength of these agents as either oxidizers or reducers.

Important to note is that the redox behavior under standard conditions can be vastly different from non-standard conditions. Factors such as concentration, temperature, and pressure can influence the potential and thus change the ranking order of oxidizing and reducing agents. When engaging with electrochemical concepts, always be clear about the experimental conditions to accurately interpret the results.

Chemical Reactivity

Chemical reactivity is an essential concept in understanding how and why chemical species engage in reactions. It encompasses the willingness of a substance to undergo a chemical change. This reactivity is influenced by a myriad of factors, including the nature of the chemical bonds, the presence of catalysts, and the electron configuration of the reactants.

In the context of redox reactions, as shown in our exercise, chemical reactivity expresses itself in the form of oxidizing and reducing agents. Oxidizing agents are substances keen on receiving electrons, thus causing other substances to lose electrons or be oxidized. Reducing agents, on the other hand, readily donate electrons, resulting in other substances gaining electrons, or being reduced.

The strength of an oxidizing or reducing agent is directly related to standard reduction potentials. An agent with a high potential is a strong oxidizer because it 'prefers' to gain electrons, whereas a negative potential suggests a strong reducer due to its tendency to lose electrons. Therefore, considering standard reduction potentials can provide a deeper understanding of a substance's reactivity and predict possible reaction outcomes in a given environment.