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Chapter 20: Problem 86

(a) What is an electrolytic cell? (b) The negative terminal of a voltage source is connected to an electrode of an electrolytic cell. Is the electrode the anode or the cathode of the cell? Explain. (c) The electrolysis of water is often done with a small amount of sulfuric acid added to the water. What is the role of the sulfuric acid?

Short Answer

Expert verified
(a) An electrolytic cell is a device that uses an electric current to drive chemical reactions, typically reduction-oxidation, at the electrodes through a process known as electrolysis. (b) The electrode connected to the negative terminal of a voltage source is the cathode of the electrolytic cell, where reduction occurs due to the gain of electrons. (c) Sulfuric acid, when added to water during electrolysis, serves as an electrolyte, enhancing the electrical conductivity of the solution and facilitating the flow of electric current for a more efficient process.

Step by step solution

01

(a) Definition of Electrolytic Cell

An electrolytic cell is a device in which an electric current is passed through a solution containing ions, causing a chemical reaction (typically reduction-oxidation) to occur at the electrodes. This process, known as electrolysis, can result in the deposition of a substance, the production of gases, or the dissociation of a solution.
02

(b) Identification of the Anode or Cathode

Since the negative terminal of a voltage source is connected to the electrode, this electrode is the cathode of the electrolytic cell. In an electrolytic cell, the cathode is where reduction occurs, meaning a gain of electrons. Electrons flow from the negative terminal of the voltage source to the cathode, allowing for the reduction reaction to take place at this electrode.
03

(c) Role of Sulfuric Acid in Water Electrolysis

The role of sulfuric acid in the electrolysis of water is to act as an electrolyte, or a substance that increases the electrical conductivity of the solution. This helps to facilitate the flow of electric current through the solution, making the electrolysis process more efficient. Since pure water has very low electrical conductivity, the addition of a small amount of sulfuric acid (H₂SO₄) significantly improves the rate of electrolysis by providing a greater number of ions to carry the electric current.

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Most popular questions from this chapter

(a) \(\mathrm{A} \mathrm{Cr}^{3+}(a q)\) solution is electrolyzed, using a current of \(7.60 \mathrm{~A}\). What mass of \(\mathrm{Cr}(s)\) is plated out after \(2.00\) days? (b) What amperage is required to plate out \(0.250 \mathrm{~mol} \mathrm{Cr}\) from a \(\mathrm{Cr}^{3+}\) solution in a period of \(8.00 \mathrm{~h}\) ?

Using data from Appendix \(\mathrm{E}\), calculate the equilibrium constant for the disproportionation of the copper(I) ion at room temperature: \(2 \mathrm{Cu}^{+}(a q) \longrightarrow \mathrm{Cu}^{2+}(a q)+\mathrm{Cu}(s)\)

A voltaic cell similar to that shown in Figure \(20.5\) is constructed. One electrode compartment 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 electrode compartments. (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 \(\mathrm{Al}\) is not coated with its oxide.

Cytochrome, a complicated molecule that we will represent as \(\mathrm{CyFe}^{2+}\), reacts with the air we breathe to supply energy required to synthesize adenosine triphosphate (ATP). The body uses ATP as an energy source to drive other reactions. (Section 19.7) At \(\mathrm{pH} 7.0\) the following reduction potentials pertain to this oxidation of \(\mathrm{CyFe}^{2+}\) : $$ \begin{aligned} \mathrm{O}_{2}(\mathrm{~g})+4 \mathrm{H}^{+}(a q)+4 \mathrm{e}^{-}--\rightarrow 2 \mathrm{H}_{2} \mathrm{O}(l) & E_{\mathrm{red}}^{\mathrm{o}}=+0.82 \mathrm{~V} \\ \mathrm{CyFe}^{3+}(a q)+\mathrm{e}^{-}--\rightarrow \mathrm{CyFe}^{2+}(a q) & E_{\mathrm{red}}^{\circ}=+0.22 \mathrm{~V} \end{aligned} $$ (a) What is \(\Delta G\) for the oxidation of \(\mathrm{CyFe}^{2+}\) by air? (b) If the synthesis of \(1.00\) mol of ATP from adenosine diphosphate (ADP) requires a \(\Delta G\) of \(37.7 \mathrm{~kJ}\), how many moles of ATP are synthesized per mole of \(\mathrm{O}_{2}\) ?

A voltaic cell is constructed that uses the following reaction and operates at \(298 \mathrm{~K}\) : $$ \mathrm{Zn}(s)+\mathrm{Ni}^{2+}(a q) \longrightarrow \mathrm{Zn}^{2+}(a q)+\mathrm{Ni}(s) $$ (a) What is the emf of this cell under standard conditions? (b) What is the emf of this cell when \(\left[\mathrm{Ni}^{2+}\right]=3.00 \mathrm{M}\) and \(\left[\mathrm{Zn}^{2+}\right]=0.100 \mathrm{M} ?\) (c) What is the emf of the cell when \(\left[\mathrm{Ni}^{2+}\right]=0.200 M\), and \(\left[\mathrm{Zn}^{2+}\right]=0.900 \mathrm{M} ?\)
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