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Consider the oxidation/reduction reaction below: Fe(CN)6 3- + Ag(s) + Br- ⇌ Fe(CN)6 4- + AgBr(s) (a) Write the reaction in terms of two balanced half-reactions. (b) Express each half-reaction as a reduction reaction with its corresponding sign and value of standard electrode potential. (c) Arrange the oxidizing agents in (b) in order of decreasing effectiveness as electron acceptors.

Consider the oxidation/reduction reaction below: Fe(CN)6 3- + Ag(s) + Br- ⇌ Fe(CN)6 4- + AgBr(s) (a) Write the reaction in terms of two balanced half-reactions. (b) Express each half-reaction as a reduction reaction with its corresponding sign and value of standard electrode potential. (c) Arrange the oxidizing agents in (b) in order of decreasing effectiveness as electron acceptors.

Chemistry
Chemistry
10th Edition
ISBN:9781305957404
Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Chapter1: Chemical Foundations
Section: Chapter Questions
Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...
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Consider the oxidation/reduction reaction below:

Fe(CN)6

3- + Ag(s) + Br- ⇌ Fe(CN)6

4- + AgBr(s)

(a) Write the reaction in terms of two balanced half-reactions.

(b) Express each half-reaction as a reduction reaction with its

corresponding sign and value of standard electrode

potential.

(c) Arrange the oxidizing agents in (b) in order of decreasing

effectiveness as electron acceptors.

Exercise 4: Consider the oxidation/reduction reaction below: Fe(CN)63 + Ag(s) + Br" = Fe(CN)6 + AgBr(s) (a) Write the reaction in terms of two balanced half-reactions. (b) Express each half-reaction as a reduction reaction with its corresponding sign and value of standard electrode potential. (c) Arrange the oxidizing agents in (b) in order of decreasing effectiveness as electron acceptors.
Transcribed Image Text:Exercise 4: Consider the oxidation/reduction reaction below: Fe(CN)63 + Ag(s) + Br" = Fe(CN)6 + AgBr(s) (a) Write the reaction in terms of two balanced half-reactions. (b) Express each half-reaction as a reduction reaction with its corresponding sign and value of standard electrode potential. (c) Arrange the oxidizing agents in (b) in order of decreasing effectiveness as electron acceptors.
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