Kaplan MCAT General Chemistry Review

Electrolytic cells, in almost all of their characteristics and behavior, are the opposite of galvanic cells. Whereas galvanic cells house spontaneous redox reactions, which when allowed to proceed, generate current and deliver electrical energy, electrolytic cells house nonspontaneous reactions, which require the input of energy to proceed. Galvanic cells supply energy; electrolytic cells consume it. The change in Gibbs function for the redox reaction of an electrolytic cell is positive. The particular type of redox reaction that is driven by the external voltage source is called electrolysis, in which chemical compounds are decomposed. For example, electrolytic cells can be used to drive the nonspontaneous decomposition of water into oxygen and hydrogen gas. Another example, the electrolysis of molten NaCl, is illustrated in Figure 11.2.

Figure 11.2

Mnemonic

Anions are negatively charged and travel to the anode. Cations are positively charged and travel to the cathode. The same principles apply to electrophoresis, a laboratory technique used to purify, separate, or identify compounds.

In this electrolytic cell, molten NaCl is decomposed into Cl₂ (g) and Na (l). The external voltage source (“battery” in Figure 11.2) supplies energy sufficient to drive a redox reaction in the direction that is thermodynamically unfavorable (i.e., nonspontaneous). In this example, Na⁺ ions migrate toward the cathode, where they are reduced to Na (l). At the same time, Cl^- ions migrate toward the anode, where they are oxidized to Cl₂ (g). You’ll notice that the half-reactions do not need to be separated into different compartments because the desired reaction is nonspontaneous. Furthermore, irrespective of the fact that a nonspontaneous reaction is being driven by an external voltage, it is still the case—as it always is for every type of electrochemical cell—that oxidation occurs at the anode and reduction occurs at the cathode. (Note that sodium is a liquid at the temperature of molten NaCl; it is also less dense than the molten salt and, thus, is easily removed as it floats to the top of the reaction vessel.) This cell is used in industry as the major means of sodium and chlorine production. You may be wondering why industry goes to all this trouble to produce these compounds and why these compounds just can’t be dug up or pumped up from somewhere, rather than manufactured. Well, think about the thermodynamics of electrolysis. Energy is supplied to drive a nonspontaneous process. This means that the products of the reaction are not naturally favored. In fact—and we’ve mentioned this before (see Chapter 2, The Periodic Table)—sodium and chlorine are never found naturally in their elemental form because they are so reactive. So before we can use elemental sodium or chlorine gas, we have to make it ourselves first.

Michael Faraday was the first to define certain quantitative principles governing the behavior of electrolytic cells. He theorized that the amount of chemical change induced in an electrolytic cell is directly proportional to the number of moles of electrons that are exchanged during a redox reaction. The number of moles exchanged can be determined from the balanced half-reaction. In general, for a reaction that involves the transfer of n electrons per atom, M,

Mⁿ⁺ + ne^- M (s)

According to this equation, one mole of M(s) will be produced if n moles of electrons are supplied. Additionally, the number of moles of electrons needed to produce a certain amount of M(s) can now be related to the measurable electrical property of charge. One electron carries a charge of 1.6 × 10^-19 coulombs (C). The charge carried by one mole of electrons can be calculated by multiplying this number by Avogadro’s number, as follows:

(1.6 × 10^-19 C/e^-)(6.022 × 10²³ e^-/mol e^-) = 96,487 C/mol e^-

Key Concept

(1.6 × 10^–19)(6.022 × 10²³) = 96,487 C/mol e^– (~10⁵C/mol e^–)

This number is called Faraday’s constant, and one faraday (F) is equivalent to the amount of charge contained in one mole of electrons (1 F = 96,487 coulombs, or J/V).

This number is called Faraday’s constant, and one faraday (F) is equivalent to the amount of charge contained in one mole of electrons (1 F = 96,487 coulombs, or J/V) or one equivalent. On the MCAT, you should round up this number to 100,000 C/mol e^- to make calculations more manageable.

CONCENTRATION CELLS