Log out Go to App
Go to App

Learning Materials

Features

Discover

Chapter 18: Problem 38

Sketch the galvanic cells based on the following overall reactions. Show the direction of electron flow, the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation. Assume that all concentrations are 1.0 \(M\) and that all partial pressures are 1.0 atm. a. $\mathrm{IO}_{3}^{-}(a q)+\mathrm{Fe}^{2+}(a q) \Longrightarrow \mathrm{Fe}^{3+}(a q)+\mathrm{I}_{2}(a q)$ b. $\mathrm{Zn}(s)+\mathrm{Ag}^{+}(a q) \rightleftharpoons \mathrm{Zn}^{2+}(a q)+\mathrm{Ag}(s)$

Short Answer

Expert verified
a. For reaction a, the cathode is the IO3- electrode and the anode is the Fe2+ electrode. Electrons flow from the Fe2+ electrode to the IO3- electrode through the wire. The salt bridge allows positive ions to migrate towards the cathode and negative ions towards the anode. The overall balanced equation is \(6\mathrm{Fe}^{2+}(aq) + 2\mathrm{IO}_{3}^{-}(aq) \rightarrow 6\mathrm{Fe}^{3+}(aq) + I_{2}(aq)\). b. For reaction b, the cathode is the Ag+ electrode and the anode is the Zn electrode. Electrons flow from the Zn electrode to the Ag+ electrode through the wire. The salt bridge allows positive ions to migrate towards the cathode and negative ions towards the anode. The overall balanced equation is \(\mathrm{Zn}(s) + 2\mathrm{Ag}^{+}(aq) \rightleftharpoons \mathrm{Zn}^{2+}(aq) + 2\mathrm{Ag}(s)\).

Step by step solution

01

Identify reduction and oxidation half-reactions

First, we need to identify the reduction and oxidation half-reactions for each overall reaction. a. For the reaction \(\mathrm{IO}_{3}^{-}(a q)+\mathrm{Fe}^{2+}(a q) \Longrightarrow \mathrm{Fe}^{3+}(a q)+\mathrm{I}_{2}(a q)\) Oxidation half-reaction: \(\mathrm{Fe}^{2+}(aq) \rightarrow \mathrm{Fe}^{3+}(aq) + e^-\) Reduction half-reaction: \(\mathrm{IO}_{3}^{-}(aq) + 6 e^{-} \rightarrow \mathrm{I}_{2}(aq)\) b. For the reaction \(\mathrm{Zn}(s)+\mathrm{Ag}^{+}(a q) \rightleftharpoons \mathrm{Zn}^{2+}(a q)+\mathrm{Ag}(s)\) Oxidation half-reaction: \(\mathrm{Zn}(s) \rightarrow \mathrm{Zn}^{2+}(aq) + 2 e^-\) Reduction half-reaction: \(\mathrm{Ag}^{+}(aq) + e^- \rightarrow \mathrm{Ag}(s)\)
02

Determine electron flow, ion migration direction, cathode, and anode

Electrons flow from the anode (oxidation) to the cathode (reduction). Ions in the salt bridge migrate to maintain electrical neutrality. a. For reaction a, Anode: \(\mathrm{Fe}^{2+}(aq) \rightarrow \mathrm{Fe}^{3+}(aq) + e^-\) Cathode: \(\mathrm{IO}_{3}^{-}(aq) + 6 e^{-} \rightarrow \mathrm{I}_{2}(aq)\) Electron flow: from Fe to IO3- Ion migration: positive ions towards the cathode and negative ions towards the anode. b. For reaction b, Anode: \(\mathrm{Zn}(s) \rightarrow \mathrm{Zn}^{2+}(aq) + 2 e^-\) Cathode: \(\mathrm{Ag}^{+}(aq) + e^- \rightarrow \mathrm{Ag}(s)\) Electron flow: from Zn to Ag+ Ion migration: positive ions towards the cathode and negative ions towards the anode.
03

Sketch the galvanic cell

Now we will sketch the galvanic cell based on the defined electron flow, ion migration, cathode and anode. a. For reaction a: - Anode: Fe2+ electrode - Cathode: IO3- electrode - Salt bridge connecting the two half-cells with positive ions going towards the cathode and negative ions going towards the anode. - Electron flow: from Fe2+ electrode to IO3- electrode through the wire. b. For reaction b: - Anode: Zn electrode - Cathode: Ag+ electrode - Salt bridge connecting the two half-cells with positive ions going towards the cathode and negative ions going towards the anode. - Electron flow: from Zn electrode to Ag+ electrode through the wire.
04

Write the overall balanced equation

Finally, we write the overall balanced equation for both reaction a and b: a. For reaction a, multiply the oxidation half-reaction by 6 and add the two half-reactions: \(6(\mathrm{Fe}^{2+}(aq) \rightarrow \mathrm{Fe}^{3+}(aq) + e^-) + (\mathrm{IO}_{3}^{-}(aq) + 6 e^{-} \rightarrow \mathrm{I}_{2}(aq))\) The overall balanced equation is: \(6\mathrm{Fe}^{2+}(aq) + 2\mathrm{IO}_{3}^{-}(aq) \rightarrow 6\mathrm{Fe}^{3+}(aq) + I_{2}(aq)\) b. For reaction b, simply add the oxidation and reduction half-reactions: \(\mathrm{Zn}(s) \rightarrow \mathrm{Zn}^{2+}(aq) + 2 e^- + 2(\mathrm{Ag}^{+}(aq) + e^- \rightarrow \mathrm{Ag}(s))\) The overall balanced equation is: \(\mathrm{Zn}(s) + 2\mathrm{Ag}^{+}(aq) \rightleftharpoons \mathrm{Zn}^{2+}(aq) + 2\mathrm{Ag}(s)\)

Unlock Step-by-Step Solutions & Ace Your Exams!

  • Full Textbook Solutions

    Get detailed explanations and key concepts

  • Unlimited Al creation

    Al flashcards, explanations, exams and more...

  • Ads-free access

    To over 500 millions flashcards

  • Money-back guarantee

    We refund you if you fail your exam.

Start your free trial

Over 30 million students worldwide already upgrade their learning with Vaia!

One App. One Place for Learning.

All the tools & learning materials you need for study success - in one app.

Get started for free

Most popular questions from this chapter

What reactions take place at the cathode and the anode when each of the following is electrolyzed? a. molten \(\mathrm{NiBr}_{2} \quad\) b. molten \(\mathrm{AlF}_{3} \quad\) c. molten \(\mathrm{MgI}_{2}\)

An experimental fuel cell has been designed that uses carbon monoxide as fuel. The overall reaction is $$2 \mathrm{CO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)$$ The two half-cell reactions are $$\mathrm{CO}+\mathrm{O}^{2-} \longrightarrow \mathrm{CO}_{2}+2 \mathrm{e}^{-}$$ $$\mathrm{O}_{2}+4 \mathrm{e}^{-} \longrightarrow 2 \mathrm{O}^{2-}$$ The two half-reactions are carried out in separate compartments connected with a solid mixture of \(\mathrm{CeO}_{2}\) and \(\mathrm{Gd}_{2} \mathrm{O}_{3}\) . Oxide ions can move through this solid at high temperatures (about \(800^{\circ} \mathrm{C} ) . \Delta G\) for the overall reaction at $800^{\circ} \mathrm{C}\( under certain concentration conditions is \)-380 \mathrm{kJ}$ . Calculate the cell potential for this fuel cell at the same temperature and concentration conditions.

Under standard conditions, what reaction occurs, if any, when each of the following operations is performed? a. Crystals of \(\mathrm{I}_{2}\) are added to a solution of \(\mathrm{NaCl}\) . b. \(\mathrm{Cl}_{2}\) gas is bubbled into a solution of \(\mathrm{Nal}\). c. A silver wire is placed in a solution of \(\mathrm{CuCl}_{2}\) d. An acidic solution of \(\mathrm{FeSO}_{4}\) is exposed to air. For the reactions that occur, write a balanced equation and calculate \(\mathscr{E}^{\circ}, \Delta G^{\circ},\) and \(K\) at \(25^{\circ} \mathrm{C}\)

A galvanic cell is based on the following half-reactions: $$\mathrm{Cu}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{Cu}(s) \quad \mathscr{E}^{\circ}=0.34 \mathrm{V}$$ $$\mathrm{V}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{V}(s) \quad \mathscr{E}^{\circ}=-1.20 \mathrm{V}$$ In this cell, the copper compartment contains a copper electrode and \(\left[\mathrm{Cu}^{2+}\right]=1.00 M,\) and the vanadium compartment contains a vanadium electrode and \(\mathrm{V}^{2+}\) at an unknown concentration. The compartment containing the vanadium \((1.00 \mathrm{L}\) of solution) was titrated with 0.0800\(M \mathrm{H}_{2} \mathrm{EDTA}^{2-}\) , resulting in the reaction $$\mathrm{H}_{2} \mathrm{EDTA}^{2-}(a q)+\mathrm{V}^{2+}(a q) \rightleftharpoons \mathrm{VEDTA}^{2-}(a q)+2 \mathrm{H}^{+}(a q)$$ $$\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad K=?$$ The potential of the cell was monitored to determine the stoichiometric point for the process, which occurred at a volume of 500.0 $\mathrm{mL} \mathrm{H}_{2} \mathrm{EDTA}^{2-}$ solution added. At the stoichiometric point, \(\mathscr{E}_{\text {cell}}\) was observed to be 1.98 \(\mathrm{V}\) . The solution was buffered at a pH of 10.00 . a. Calculate\(\mathscr{E}_{\text {cell}}\) before the titration was carried out. b. Calculate the value of the equilibrium constant, \(K\), for the titration reaction. c. Calculate \(\mathscr{E}_{\text {cell}}\) at the halfway point in the titration.

When magnesium metal is added to a beaker of \(\mathrm{HCl}(\mathrm{aq})\), a gas is produced. Knowing that magnesium is oxidized and that hydrogen is reduced, write the balanced equation for the reaction. How many electrons are transferred in the balanced equation? What quantity of useful work can be obtained when \(\mathrm{Mg}\) is added directly to the beaker of \(\mathrm{HCl}\)? How can you harness this reaction to do useful work?
See all solutions

Recommended explanations on Chemistry Textbooks

Physical Chemistry

Read Explanation

Chemistry Branches

Read Explanation

Chemical Analysis

Read Explanation

Organic Chemistry

Read Explanation

Nuclear Chemistry

Read Explanation

The Earths Atmosphere

Read Explanation
View all explanations

What do you think about this solution?

We value your feedback to improve our textbook solutions.

Study anywhere. Anytime. Across all devices.

Sign-up for free