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Chapter 20: Problem 27

Indicate whether each statement is true or false: (a) The cathode is the electrode at which oxidation takes place. (b) A galvanic cell is another name for a voltaic cell. (c) Electrons flow spontaneously from anode to cathode in a voltaic cell.

Short Answer

Expert verified
(a) False, as cathode is the electrode where reduction occurs, not oxidation. (b) True, since galvanic cells and voltaic cells are the same and the terms can be used interchangeably. (c) True, as electrons flow spontaneously from the anode to the cathode in a voltaic cell due to the redox reaction.

Step by step solution

01

Understanding Cathode and Oxidation

In electrochemical cells, cathode is the electrode where reduction takes place, and anode is the electrode where oxidation occurs. Oxidation involves the loss of electrons, while reduction involves the gain of electrons. Now let's analyze the statement: (a) The cathode is the electrode at which oxidation takes place.
02

Verifying Statement (a)

Since we know that the cathode is where reduction occurs and not oxidation, the statement is false. #Statement b#
03

Understanding Galvanic Cells and Voltaic Cells

Galvanic cells and voltaic cells are names for the same type of electrochemical cell. These cells convert chemical energy into electrical energy through spontaneous redox reactions. The terms are used interchangeably. Now let's analyze the statement: (b) A galvanic cell is another name for a voltaic cell.
04

Verifying Statement (b)

Since galvanic cells and voltaic cells are the same and can be used interchangeably, the statement is true. #Statement c#
05

Understanding Electron Flow in Voltaic Cells

In a voltaic cell, the anode is where oxidation occurs, and the cathode is where reduction occurs. During the redox reaction, electrons are released at the anode (oxidation) and flow to the cathode (reduction) through an external circuit, generating an electric current. Now let's analyze the statement: (c) Electrons flow spontaneously from anode to cathode in a voltaic cell.
06

Verifying Statement (c)

Since electrons do flow spontaneously from the anode to the cathode in a voltaic cell due to the redox reaction, the statement is true.

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Most popular questions from this chapter

A voltaic cell that uses the reaction $$ \operatorname{PdCl}_{4}^{2-}(a q)+\mathrm{Cd}(s) \longrightarrow \mathrm{Pd}(s)+4 \mathrm{Cl}^{-}(a q)+\mathrm{Cd}^{2+}(a q) $$ has a measured standard cell potential of \(+1.03 \mathrm{V}\) . (a) Write the two half-cell reactions. (b) By using data from Appendix E, determine \(E_{\text { red }}^{\circ}\) for the reaction involving Pd. (c) Sketch the voltaic cell, label the anode and cathode, and indicate the direction of electron flow.

A voltaic cell similar to that shown in Figure 20.5 is constructed. One half- cell consists of an aluminum strip placed in a solution of \(\mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}\) , and the other has a nickel strip placed in a solution of \(\mathrm{NiSO}_{4}\) . The overall cell reaction is $$ 2 \mathrm{Al}(s)+3 \mathrm{Ni}^{2+}(a q) \longrightarrow 2 \mathrm{Al}^{3+}(a q)+3 \mathrm{Ni}(s) $$ (a) What is being oxidized, and what is being reduced? (b) Write the half-reactions that occur in the two half-cells. (c) Which electrode is the anode, and which is the cathode? (d) Indicate the signs of the electrodes. (e) Do electrons flow from the aluminum electrode to the nickel electrode or from the nickel to the aluminum? (f) In which directions do the cations and anions migrate through the solution? Assume the Al is not coated with its oxide.

Complete and balance the following equations, and identify the oxidizing and reducing agents. (Recall that the O atoms in hydrogen peroxide, \(\mathrm{H}_{2} \mathrm{O}_{2}\) , have an atypical oxidation state.) $$ \begin{array}{l}{\text { (a) } \mathrm{NO}_{2}^{-}(a q)+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q) \longrightarrow \mathrm{Cr}^{3+}(a q)+\mathrm{NO}_{3}^{-}(a q)} \\ \quad {\text { (acidic solution) }} \\\ {\text { (b) } \mathrm{S}(s)+\mathrm{HNO}_{3}(a q) \rightarrow \mathrm{H}_{2} \mathrm{SO}_{3}(a q)+\mathrm{N}_{2} \mathrm{O}(g)} \\ {\quad(\text { acidic solution })} \\ {\text { (c) } \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{CH}_{3} \mathrm{OH}(a q) \longrightarrow \mathrm{HCOOH}(a q)+} \\\ \quad {\mathrm{Cr}^{3+}(a q)(\text { acidic solution })} \\ {\text { (d) } \operatorname{BrO}_{3}^{-}(a q)+\mathrm{N}_{2} \mathrm{H}_{4}(g) \longrightarrow \mathrm{Br}^{-}(a q)+\mathrm{N}_{2}(g)} \\ \quad {\text { (acidic solution) }} \\ {\text { (e) } \mathrm{NO}_{2}^{-}(a q)+\mathrm{Al}(s) \longrightarrow \mathrm{NH}_{4}^{+}(a q)+\mathrm{AlO}_{2}^{-}(a q)} \\ \quad {\text { (basic solution) }} \\\ {\text { (f) } \mathrm{H}_{2} \mathrm{O}_{2}(a q)+\mathrm{ClO}_{2}(a q) \rightarrow \mathrm{ClO}_{2}^{-}(a q)+\mathrm{O}_{2}(g)} \\ \quad {\text { (basic solution) }}\end{array} $$

Gold exists in two common positive oxidation states, \(+1\) and \(+3 .\) The standard reduction potentials for these oxidation states are $$ \begin{array}{ll}{\mathrm{Au}^{+}(a q)+\mathrm{e}^{-}} \quad {\longrightarrow \mathrm{Au}(s) \quad E_{\mathrm{red}}^{\circ}=+1.69 \mathrm{V}} \\\ {\mathrm{Au}^{3+}(a q)+3 \mathrm{e}^{-} \longrightarrow \mathrm{Au}(s)} \quad {E_{\mathrm{red}}^{\circ}=+1.50 \mathrm{V}}\end{array} $$ (a) Can you use these data to explain why gold does not tarnish in the air? ( b) Suggest several substances that should be strong enough oxidizing agents to oxidize gold metal. (c) Miners obtain gold by soaking gold-containing ores in an aqueous solution of sodium cyanide. A very soluble complex ion of gold forms in the aqueous solution because of the redox reaction $$ \begin{array}{rl}{4 \mathrm{Au}(s)+8 \mathrm{NaCN}(a q)+2} & {\mathrm{H}_{2} \mathrm{O}(l)+\mathrm{O}_{2}(g)} \\ {\longrightarrow} & {4 \mathrm{Na}\left[\mathrm{Au}(\mathrm{CN})_{2}\right](a q)+4 \mathrm{NaOH}(a q)}\end{array} $$ What is being oxidized, and what is being reduced in this reaction? (d) Gold miners then react the basic aqueous product solution from part (c) with Zn dust to get gold metal. Write a balanced redox reaction for this process. What is being oxidized, and what is being reduced?

Hydrazine \(\left(\mathrm{N}_{2} \mathrm{H}_{4}\right)\) and dinitrogen tetroxide \(\left(\mathrm{N}_{2} \mathrm{O}_{4}\right)\) form a self-igniting mixture that has been used as a rocket propellant. The reaction products are \(\mathrm{N}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\) . (a) Write a balanced chemical equation for this reaction. (b) What is being oxidized, and what is being reduced? (c) Which substance serves as the reducing agent and which as the oxidizing agent?
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