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Redox Main
Half-Reactions Balancing Redox Reactions in Neutral, Acidic, and Basic Solutions |
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Reduction Potentials
By comparing the electronegativities of elements in our "Reference" section, you can tell that some atoms want their electrons more than others. The halogens are electromaniacs, while the alkali metals almost couldn't care less. The noble gases don't want a new electron, but it's almost impossible to get rid of their old ones! We've already explained these trends using electron orbitals, but they play an important part in electrochemistry.
For a given oxidation or reduction half-reaction, scientists have come up with a number called a reduction potential (symbol: E°, where that superscript "o" once again means under standard thermodynamic conditions, when all gases are at 25 °C and 1 atm and all aqueous compounds are at 1M concentration) that measures how much that reaction wants to occur. Keeping up with scientists' ambivalence about whether negatives are "good" or "bad" (remember "Thermodynamics"?), a positive reduction potential indicates a product-favored reaction; the bigger the reduction potential, the more strongly the reaction will want to occur. A negative value represents a reactant-favored reaction; the more negative the reduction potential, the more unwilling the reaction is to occur.
Since reduction potentials are a measure of how strongly electrons are pulled or pushed from one species to another, they are not dependent on the amount of reactants present. In other words, multiplying the entire reaction by a number, say 2, will not change the reduction potential. However, reversing the reaction will reverse the sign of the reduction potential. These values are measured in volts--the same volts that your 9-volt Duracell, 12-volt car battery, and 120-volt residential electric socket (if you live in the USA) are measured in.
For example, let's look at two half-reactions and their reduction potentials: the oxidation of lithium and the reduction of fluorine.
Li(s)
Li+ + e- Eº = 3.05
F2 + 2 e-As you can see, both of these reactions have positive E° values--they really want to occur. In fact, these two reactions have the highest voltages for an oxidation and reduction reaction, respectively. Now, according to the definition of a half-reaction and Hess' Law, which we discussed last chapter, you should be able to add these reactions and the E° values; indeed, you can do this. Adding the two half-reactions together (multiplying the oxidation reaction by 2, since fluorine gas needs 2 electrons to reduce) yields the following redox reaction and E°:
2 Li(s) + F2(g)
We can also use E° values to predict which atom or molecule will get the electrons in a redox reaction. For instance, the potentials for the oxidation of manganese and zinc are both positive; both atoms' electrons are easily removed, according to the following half-reactions:
Zn(s)
Mn(s)By comparing the reduction potentials, we can see that manganese wants to lose its electrons more than zinc. So, if you have a solution with both solid and aqueous manganese and zinc, the following half-reactions will occur:
Zn2+(aq) + 2 e-
Mn(s)Even though zinc does not want to reduce, the stronger potential of the manganese reaction forces it to absorb the extra electrons:
Zn2+(aq) + Mn(s)
The same principle can be applied to competing reduction reactions; for example, fluorine can force iodine to oxidize, stealing its electrons. The reduction reaction with the stronger potential will proceed, while the other reduction reaction will be reversed to provide electrons.
Since some reduction potentials are positive and some negative, you probably wonder how they are measured. As in the enthalpy concept of last chapter, there is no "absolute standard" against which reduction potentials can be measured. Instead, scientists decided that the reduction of hydronium ions into hydrogen gas would have a reduction potential of 0.00. This system is called the reference electrode or standard hydrogen electrode, because all other redox reactions are measured against it. Stronger oxidizing agents would force hydrogen to oxidize, symbolized by a positive reduction potential; weaker ones would be oxidized by the H+, giving a negative potential. However, keep in mind that these values are for standard conditions only. Later in this chapter, we will discuss reduction potentials in nonstandard conditions.
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