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Chapter 17: Problem 76

Which process (oxidation/reduction) takes place at the anode of a galvanic cell?

Short Answer

Expert verified
In a galvanic cell, the anode is the electrode where oxidation occurs. Thus, the process that takes place at the anode of a galvanic cell is oxidation.

Step by step solution

01

Understand Galvanic Cells

A galvanic cell, also known as a voltaic cell, converts chemical energy into electrical energy through redox reactions. It consists of two half-cells, each containing an electrode (anode or cathode) and an electrolyte. The two half-cells are connected through a salt bridge which allows the flow of ions between the half-cells.
02

Remember Redox Reactions

Redox reactions involve the transfer of electrons between species. In a redox reaction, one species undergoes oxidation (loses electrons) while the other undergoes reduction (gains electrons).
03

Identify Anode and Cathode

In a galvanic cell, the anode is the electrode where oxidation occurs, and the cathode is the electrode where reduction takes place. Electrons flow from the anode to the cathode through an external circuit, generating an electric current.
04

Determine the Process at the Anode

As we now know that oxidation occurs at the anode of a galvanic cell, we can conclude that the process taking place at the anode is oxidation.

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Most popular questions from this chapter

Balance each of the following half-reactions, which take place in acidic solution. a. \(\mathrm{SiO}_{2}(s) \rightarrow \mathrm{Si}(s)\) b. \(\mathrm{S}(s) \rightarrow \mathrm{H}_{2} \mathrm{S}(a q)\) c. \(\mathrm{NO}_{3}^{-}(a q) \rightarrow \mathrm{HNO}_{2}(a q)\) d. \(\mathrm{NO}_{3}(a q) \rightarrow \mathrm{NO}(g)\)

Balance each of the following half-reactions. a. \(\mathrm{I}^{-}(a q) \rightarrow \mathrm{I}_{2}(s)\) b. \(\mathrm{O}_{2}(g) \rightarrow \mathrm{O}^{2-}(s)\) c. \(\mathrm{P}_{4}(s) \rightarrow \mathrm{P}^{3-}(s)\) d. \(\mathrm{Cl}_{2}(g) \rightarrow \mathrm{Cl}^{-}(a q)\)

At which electrode (anode/cathode) do species gain electrons in a galvanic cell?

In each of the following reactions, identify which element is oxidized and which is reduced by assigning oxidation states. a. \(2 \mathrm{Cu}(s)+\mathrm{S}(s) \rightarrow \mathrm{Cu}_{2} \mathrm{S}(s)\) b. \(2 \mathrm{Cu}_{2} \mathrm{O}(s)+\mathrm{O}_{2}(g) \rightarrow 4 \mathrm{CuO}(s)\) c. \(4 \mathrm{B}(s)+3 \mathrm{O}_{2}(g) \rightarrow 2 \mathrm{B}_{2} \mathrm{O}_{3}(s)\) d. \(6 \mathrm{Na}(s)+\mathrm{N}_{2}(g) \rightarrow 2 \mathrm{Na}_{3} \mathrm{N}(s)\)

Pure iron ordinarily rusts quickly, but steel does not corrode nearly as fast. How does steel resist corrosion?
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