## Electrolytic Properties

#### Learning Objective

• Use a table of standard reduction potentials to determine which species in solution will be reduced or oxidized.

#### Key Points

• When an electrical current passes through a solution (often of electrolytes), a cation or neutral molecule gets reduced at the cathode, and an anion or neutral molecule gets oxidized at the anode.
• To determine which species in solution will be oxidized and which reduced, a table of standard reduction potentials can identify the most thermodynamically viable option.
• In practice, electrolysis of pure water can create hydrogen gas.

#### Terms

• electronthe subatomic particle that has a negative charge and orbits the nucleus; the flow of electrons in a conductor constitutes electricity.
• electrodethe terminal through which electric current passes between metallic and nonmetallic parts of an electric circuit; in electrolysis, the cathode and anode are placed in the solution separately.

## Electrolytic Properties

When electrodes are placed in an electrolyte solution and a voltage is applied, the electrolyte will conduct electricity. Lone electrons cannot usually pass through the electrolyte; instead, a chemical reaction occurs at the cathode that consumes electrons from the anode. Another reaction occurs at the anode, producing electrons that are eventually transferred to the cathode. As a result, a negative charge cloud develops in the electrolyte around the cathode, and a positive charge develops around the anode. The ions in the electrolyte neutralize these charges, enabling the electrons to keep flowing and the reactions to continue.

For example, in a solution of ordinary table salt (sodium chloride, NaCl) in water, the cathode reaction will be:

$2H_{2}O + 2e^{-} \rightarrow 2OH^{-} + H_{2}$

and hydrogen gas will bubble up. The anode reaction is:

$2NaCl \rightarrow 2 Na^{+} + Cl_2 + 2e^{-}$

and chlorine gas will be liberated. The positively-charged sodium ions Na+ will react toward the cathode, neutralizing the negative charge of OH there; the negatively-charged hydroxide ions OH will react toward the anode, neutralizing the positive charge of Na+ there. Without the ions from the electrolyte, the charges around the electrode slow continued electron flow; diffusion of H+ and OH through water to the other electrode takes longer than movement of the much more prevalent salt ions.

In other systems, the electrode reactions can involve electrode metal as well as electrolyte ions. In batteries for example, two materials with different electron affinities are used as electrodes: outside the battery, electrons flow from one electrode to the other; inside, the circuit is closed by the electrolyte’s ions. Here, the electrode reactions convert chemical energy to electrical energy.

## Oxidation and Reduction at the Electrodes

Oxidation of ions or neutral molecules occurs at the anode, and the reduction of ions or neutral molecules occurs at the cathode. Two mnemonics for remembering that reduction happens at the cathode and oxidation at the anode are: “Red Cat” (reduction – cathode) and “An Ox” (anode – oxidation). The mnemonic “LeO said GeR” is useful for remembering “lose an electron in oxidation” and “gain an electron in reduction.”

It is possible to oxidize ferrous ions to ferric ions at the anode. For example:

$Fe^{2+}(aq) \rightarrow Fe^{3+} (aq) + e^{-}$

Neutral molecules can also react at either electrode. For example, p-Benzoquinone can be reduced to hydroquinone at the cathode:

$+ 2 e^{-} + 2 H^{+} \rightarrow$

In the last example, H+ ions (hydrogen ions) also take part in the reaction, and are provided by an acid in the solution or by the solvent itself (water, methanol, etc.). Electrolysis reactions involving H+ ions are fairly common in acidic solutions, while reactions involving OH- (hydroxide ions) are common in alkaline water solutions.

The oxidized or reduced substances can also be the solvent (usually water) or electrodes. It is possible to have electrolysis involving gases.

In order to determine which species in solution will be oxidized and which will be reduced, the standard electrode potential of each species may be obtained from a table of standard reduction potentials, a small sampling of which is shown here:

Historically, oxidation potentials were tabulated and used in calculations, but the current standard is to only record the reduction potential in tables. If a problem demands use of oxidation potential, it may be interpreted as the negative of the recorded reduction potential. For example, referring to the data in the table above, the oxidation of elemental sodium (Na(s)) is a highly favorable process with a value of $E_{ox}^0 (V)$= + 2.71 V; this makes intuitive sense because the loss of one electron from a sodium atom produces a sodium cation, which has the same electron configuration as neon, a noble gas. The production of this low-energy and stable electron configuration is clearly a favorable process. Chlorine gas on the other hand is much more likely to be reduced under normal conditions, as can be inferred from the value of $E_{red}^0 (V)$= +1.36 V in the table. Recall that a more positive potential always means that that reaction will be favored; this will have consequences concerning redox reactions.