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Redox Main
Half-Reactions Balancing Redox Reactions in Neutral, Acidic, and Basic Solutions |
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Electrolysis
Unlike galvanic cells, electrolytic cells involve reactions that do not want to occur; a voltage source or galvanic cell is needed to push the electrons through the unfavorable reaction, forcing it to proceed. Therefore, we can define electrolytic cells as devices using electricity to create a chemical change (exactly opposite of galvanic cells). Electrolysis can occur in aqueous solutions or in molten salts.
A common electrolysis situation is extracting chlorine from seawater, according to the reaction below.
2 Cl-(aq)
Cl2(g) + 2 e- E° = -1.36 V
Since salt (NaCl) is dissolved in seawater, the reduction of the sodium ion could happen as well.
Na+(aq) + 2 e-
However, there's a problem with these reactions: both half-reactions have significantly negative potentials, and there are some other reactions that we need to consider before we hook up our voltage source and start pumping out chlorine and sodium. Water itself can be ripped apart by strong electric currents, as shown in the reactions below:
Reduction: 2 H2O(l) + 2 e-
Reduction: 2 H3O+(aq) + 2 e-
Oxidation: 6 H2O(l)Note that more factors than just the potentials of the reactions come into play; the speed with which the reaction can occur also helps decide. You must use some common sense in figuring out which reaction will take precedence. In this case, the voltage of the reduction of sodium is so low that water will be reduced to hydrogen gas and hydroxide ions at the cathode. The reference electrode (reduction of hydronium) has an even better voltage, but the concentration of hydronium in pure water is too low (10-7M) for this reaction to occur. Similarly, water could be oxidized to oxygen and hydronium at the anode, but this process is slow enough that the oxidization of chlorine is preferable. Therefore, we can add the two selected half-reactions as shown below:
Reduction: 2 H2O(l) + 2 e-
Oxidation: 2 Cl-(aq)Net reaction: 2 H2O(l) + 2 Cl-(aq)
Therefore, a voltage source or galvanic cell with a voltage of at least +2.19 V is necessary to cause this reaction to occur. In electrolytic cells, the polarities of the anode and cathode are reversed: the cathode is now labeled "-", because the external source is pumping electrons into the cathode. The anode is labeled "+", because the electrons return to the source from this pole. However, reduction still always occurs at the cathode, and oxidation always takes place at the anode. Be sure to remember the polarity differences between galvanic/voltaic and electrolytic cells.
The results of our experiment can be easily confirmed; bubbles will appear at both electrodes, with hydrogen appearing at the cathode and greenish chlorine being produced at the anode. In addition, an indicator or pH meter will show alkalinity at the cathode.
Electrolysis can also occur in molten salts, a process which is considerably less complicated than aqueous reactions, because the destruction of water is not a factor. Using the example above, molten NaCl can be electrolyzed according to the familiar half-reactions above. Because the temperature is very high, the reduction potentials are only estimates (due to the non-standard conditions). However, adding up the reduction potentials for the oxidation of chlorine and the reduction of sodium gives a voltage of about -4 V, illustrating the difficulty of reducing sodium. Note the differences caused by simply dissolving salt in water! Also keep in mind the reversal of polarity in electrolytic cells, the alternative reactions involving the destruction of water, and the dual factors of reduction potential and speed of reaction that determine which half-reaction will predominate in aqueous reactions.
Another important aspect of electrolysis reactions involves measuring the amount of electrons that pass through the cell. If we know the number of moles of electrons transferred, we can find the number of moles of products generated and reactants consumed. Therefore, a quantitative discussion of electron counting is essential.
Electric charge is usually measured in the coulombs (symbol: C), a unit equivalent in charge to 6.24 x 1018 electrons. Current on a circuit is measured in charge passed per unit of time, so the unit of current, the ampere (symbol: I, or amp for short), is defined as a current of one coulomb passed per second. Because one coulomb is 6.24 x 1018 electrons, and one mole contains 6.02 x 1023 units, one mole of electrons contains about 96,500 coulombs (more specifically, 96,485.21 coulombs are carried by one mole of electrons). This constant is named after Michael Faraday, a 19th-century scientist who pioneered the science of electrochemistry.
Knowing the definition of an ampere and the Faraday Constant, we can determine how many moles or grams of a substance can be electrolyzed by a given reaction, or how many amps of current are needed to electrolyze a given quantity of reactant in a given time.
For example, say we want to know how many grams of chlorine gas can be recovered from a molten sample of NaCl with a current of 5 amperes, run for 2.5 hours. First, let's rewrite the half-reaction from above:
2 Cl-(aq)Don't be confused if the electrons appear on the right side; since the reaction above is unfavorable, the electrons must be forced through the system. Now, since the 5-amp current ran for 2.5 hours, and since one amp equals one coulomb per second, we can determine the number of coulombs passed:
(5 C/sec)(3600 sec/hr)(2.5 hr) = 45000 coulombs.
Dividing by the Faraday Constant, about 96,500 C/mol, we find that about .466 moles of electrons passed through the circut. Since each mole of chlorine gas oxidized releases two electrons, we get a result of .233 moles of Cl2 gas generated. With a molar mass of about 71 grams, we know that 16.543 grams of chlorine gas were released. Using PV=nRT, we could alternately find the volume of chlorine produced, and use its density to find the weight.
Other problems might ask you to find a number of coulombs or the amperage of a current necessary to complete a given reaction. Using the information given above, you should be able to work out the calculations needed by defining what you have, what you want, and what steps are necessary to get there.
The final page in this chapter will explain some of the equations and mathematical principles behind redox reactions, including redox reactions in non-standard conditions and relations between E°, ΔG, and K, the equilibrium constant.
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