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Chapter 19: Problem 79

Describe an experiment that would enable you to determine which is the cathode and which is the anode in a galvanic cell using copper and zinc electrodes.

Short Answer

Expert verified
The experiment involves setting up a galvanic cell with zinc and copper electrodes and observing the direction of current on the ammeter and the behavior of the electrodes over time. The anode is the electrode from which the current is flowing (where oxidation is taking place), and in this case, it would be the zinc electrode. The cathode is the electrode where the current is flowing to (where reduction is taking place), which would be the copper electrode.

Step by step solution

01

Setup the Galvanic Cell

First, set up your galvanic cell. Place a copper electrode in one beaker filled with a copper (II) sulphate solution (CuSO4), and a zinc electrode in another beaker filled with a zinc sulphate solution (ZnSO4). Connect the two electrodes with a wire that has an ammeter in the middle, and connect the two solutions with a salt bridge (a U-tube that has a solution of a non-reactive electrolyte, such as potassium nitrate).
02

Observe the Ammeter

Next, observe the ammeter. The direction of the current will indicate which is the anode and which is the cathode, since electric current (by convention) flows from positive to negative. If the current is flowing from the Zinc electrode to the copper electrode, then the Zinc electrode is the anode as oxidation is taking place on it, and the copper electrode is the cathode, where reduction is taking place.
03

Observe the Electrodes

Another way to determine the anode and cathode is to observe the electrodes over time. As the anode undergoes oxidation, it will slowly dissolve into the solution, while the cathode, where reduction takes place, will have copper plating onto it due to the reduction of Cu ions from the solution.

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Most popular questions from this chapter

The oxidation of \(25.0 \mathrm{~mL}\) of a solution containing \(\mathrm{Fe}^{2+}\) requires \(26.0 \mathrm{~mL}\) of \(0.0250 \mathrm{M} \mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) in acidic solution. Balance the following equation and calculate the molar concentration of \(\mathrm{Fe}^{2+}\): $$ \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+\mathrm{Fe}^{2+}+\mathrm{H}^{+} \longrightarrow \mathrm{Cr}^{3+}+\mathrm{Fe}^{3+} $$

Consider a galvanic cell composed of the SHE and a half-cell using the reaction \(\mathrm{Ag}^{+}(a q)+e^{-} \rightarrow \operatorname{Ag}(s)\). (a) Calculate the standard cell potential. (b) What is the spontaneous cell reaction under standard-state conditions? (c) Calculate the cell potential when \(\left[\mathrm{H}^{+}\right]\) in the hydrogen electrode is changed to (i) \(1.0 \times\) \(10^{-2} M\) and (ii) \(1.0 \times 10^{-5} M,\) all other reagents being held at standard-state conditions. (d) Based on this cell arrangement, suggest a design for a pH meter.

An acidified solution was electrolyzed using copper electrodes. A constant current of 1.18 A caused the anode to lose \(0.584 \mathrm{~g}\) after \(1.52 \times 10^{3} \mathrm{~s}\). (a) What is the gas produced at the cathode and what is its volume at STP? (b) Given that the charge of an electron is \(1.6022 \times 10^{-19} \mathrm{C}\), calculate Avogadro's number. Assume that copper is oxidized to \(\mathrm{Cu}^{2+}\) ions.

For each of the following redox reactions, (i) write the half-reactions; (ii) write a balanced equation for the whole reaction, (iii) determine in which direction the reaction will proceed spontaneously under standard-state conditions: (a) \(\mathrm{H}_{2}(g)+\mathrm{Ni}^{2+}(a q) \longrightarrow \mathrm{H}^{+}(a q)+\mathrm{Ni}(s)\) (b) \(\mathrm{MnO}_{4}^{-}(a q)+\mathrm{Cl}^{-}(a q) \longrightarrow\) \(\mathrm{Mn}^{2+}(a q)+\mathrm{Cl}_{2}(g)\) (in acid solution) (c) \(\mathrm{Cr}(s)+\mathrm{Zn}^{2+}(a q) \longrightarrow \mathrm{Cr}^{3+}(a q)+\mathrm{Zn}(s)\)

Write the Nernst equation for the following processes at some temperature \(T\) : (a) \(\mathrm{Mg}(s)+\mathrm{Sn}^{2+}(a q) \longrightarrow \mathrm{Mg}^{2+}(a q)+\operatorname{Sn}(s)\) (b) \(2 \mathrm{Cr}(s)+3 \mathrm{~Pb}^{2+}(a q) \longrightarrow 2 \mathrm{Cr}^{3+}(a q)+3 \mathrm{~Pb}(s)\)
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