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Chapter 20: Problem 1

In the Brønsted-Lowry concept of acids and bases, acidbase reactions are viewed as proton-transfer reactions. The stronger the acid, the weaker is its conjugate base. If we were to think of redox reactions in a similar way, what particle would be analogous to the proton? Would strong oxidizing agents be analogous to strong acids or strong bases?

Short Answer

Expert verified
In the Brønsted-Lowry concept, acids are proton (H+) donors and bases are proton (H+) acceptors. The analogous particle to the proton in redox reactions would be the electron (e-), as redox reactions involve the transfer of electrons between chemical species. Strong oxidizing agents, which are electron acceptors and have higher affinity for accepting electrons, would be analogous to strong acids, as both are donating a particle (protons or electrons) to other species in their respective reactions.

Step by step solution

01

Understand the Brønsted-Lowry concept

The Brønsted-Lowry concept defines acids as proton (H+) donors and bases as proton (H+) acceptors. In this concept, acid-base reactions are viewed as proton-transfer reactions.
02

Identify the analogous particle in redox reactions

Redox reactions involve transfer of electrons between chemical species. So, the analogous particle to the proton in redox reactions would be electron (e-).
03

Identify the analogous species for strong oxidizing agents

In redox reactions, oxidizing agents are electron acceptors, and reducing agents are electron donors. Strong oxidizing agents are those that have higher affinity for accepting electrons and will get reduced easily. Acids are proton donors in Brønsted-Lowry concept, while oxidizing agents are electron acceptors in redox reactions. So, strong oxidizing agents would be analogous to strong acids in this context.

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Most popular questions from this chapter

If the equilibrium constant for a two-electron redox reaction at $298 \mathrm{~K}\( is \)2.2 \times 10^{5},\( calculate the corresponding \)\Delta G^{\circ}\( and \)E^{\circ}$.

Is each of the following substances likely to serve as an oxidant or a reductant: $(\mathbf{a}) \mathrm{Ce}^{3+}(a q),(\mathbf{b}) \mathrm{Ca}(s),(\mathbf{c}) \mathrm{ClO}_{3}^{-}(a q),\( (d) \)\mathrm{N}_{2} \mathrm{O}_{5}(g) ?$

A cell has a standard cell potential of \(+0.257 \mathrm{~V}\) at $298 \mathrm{~K}$. What is the value of the equilibrium constant for the reaction \((\mathbf{a})\) if \(n=1 ?(\mathbf{b})\) if \(n=2 ?(\mathbf{c})\) if \(n=3 ?\)

During a period of discharge of a lead-acid battery, \(300 \mathrm{~g}\) of \(\mathrm{PbO}_{2}(s)\) from the cathode is converted into \(\mathrm{PbSO}_{4}(s)\). (a) What mass of \(\mathrm{Pb}(s)\) is oxidized at the anode during this same period? (b) How many coulombs of electrical charge are transferred from \(\mathrm{Pb}\) to \(\mathrm{PbO}_{2}\) ?

The standard reduction potentials of the following half-reactions are given in Appendix E: $$ \begin{array}{l} \mathrm{Fe}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{Fe}(s) \\ \mathrm{Cd}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \operatorname{Cd}(s) \\\ \mathrm{Sn}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \operatorname{Sn}(s) \\\ \mathrm{Ag}^{+}(a q)+\mathrm{e}^{-} \longrightarrow \operatorname{Ag}(s) \end{array} $$ (a) Determine which combination of these half-cell reactions leads to the cell reaction with the largest positive cell potential and calculate the value. (b) Determine which combination of these half-cell reactions leads to the cell reaction with the smallest positive cell potential and calculate the value.
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