liquidA state of matter in which the atomic-scale particles remain close together but are able to change their positions so that the matter takes the shape of its container to be electrolyzed. These electrodes are often made of an inertUnreactive. Used to describe coordination complexes that exchange ligands slowly or an electrode in an electrochemical cell that serves only as a surface where reaction can occur and is neither consumed nor added to during reaction. material such as stainless steel, platinum, or graphite. The liquid to be electrolyzed must be able to conduct electricity, and so it is usually an aqueousDescribing a solution in which the solvent is water. solutionA mixture of one or more substances dissolved in a solvent to give a homogeneous mixture. of an electrolyte or a molten ionic compoundA compound containing oppositely charged ions held together by electrostatic attraction. Usually the ions are in a crystal lattice with positive ions surrounded by negative ions and negative ions surrounded by positive ions.. The electrodes are connected by wires to a battery or other source of direct current. This current source may be thought of as an “electron pump” which takes in electrons from one electrode and forces them out into the other electrode. The electrode from which electrons are removed becomes positively charged, while the electrode to which they are supplied has an excess of electrons and a negative charge.
The negatively charged electrode will attract positive ions (cations) toward it from the solution. It can donate some of its excess electrons to such cations or to other species in the liquid being electrolyzed. Hence this electrode is in effect a reducing agent. In any electrochemical cellA system containing an oxidation-reduction reaction in which oxidation and reduction reactions are physically separated and the transferred electrons pass through an electrical circuit. Voltaic (galvanic) cells produce and electic current; electrolyic cells use electric current to force a reaction to occur. (electrolytic or galvanic) the electrode at which reduction occurs is called the cathode.
The positive electrode, on the other hand, will attract negative ions (anions) toward itself. This electrode can accept electrons from those negative ions or other species in the solution and hence behaves as an oxidizing agent. In any electrochemical cell the anode is the electrode at which oxidation occurs. An easy way to remember which electrode is which is that anode and oxidation begin with vowels while cathode and reduction begin with consonants.
The following video shows this process taking place in a neutral solution of water with some electrolytes present.
As an example of how electrolysis can cause a chemical reactionA process in which one or more substances, the reactant or reactants, change into one or more different substances, the products; chemical change involves rearrangement, combination, or separation of atoms. Also called chemical change. to occur, suppose we pass a direct electrical current through 1 M HCl. The H3O+ ions in this solution will be attracted to the cathode, and the Cl– ions will migrate toward the anode. At the cathode, H3O+ will be reduced to H2 gasA state of matter in which a substance occupies the full volume of its container and changes shape to match the shape of the container. In a gas the distance between particles is much greater than the diameters of the particles themselves; hence the distances between particles can change as necessary so that the matter uniformly occupies its container. according to the half-equation
2H+ + 2e– → H2 (1a)
(As seen in other sections, we shall write H+ instead of H3O+ in half-equations to save time.) At the anode, electrons will be accepted from Cl– ions, oxidizing them to Cl2:
2Cl– → Cl2 + 2e– (1b)
During electrolysis H2(g) and Cl2(g) bubble from the cathode and anode, respectively. The overall equation for the electrolysis is the sum of Eqs. (1a) and (1b):
2H+(aq) + 2Cl–(aq) → H2(g) + Cl2(g) (1)
or 2H3O+(aq) + 2Cl–(aq) → H2(g) + Cl2(g) + 2H2O(l)
The net reaction [Eq. (1)] is the reverse of the spontaneousCapable of proceeding without an outside source of energy; refers to a reaction in which the products are thermodynamically favored (product-favored reaction). combination of H2(g) with Cl2(g) to form HCl(aq). Such a result is true of electrolysis in general: electrical current supplied from outside the system causes a non-spontaneous chemical reaction to occur.
Although electrolysis always reverses a spontaneous redox reaction, the result of a given electrolysis may not always be the reaction we want. In an aqueous solution, for example, there are always a great many water molecules in the vicinity of both the anode and cathode. These water molecules can donate electrons to the anode or accept electrons from the cathode just as anions or cations can. Consequently the electrolysis may oxidize and/or reduce water instead of causing the dissolved electrolyte to react. An example of this problem is electrolysis of lithium fluoride, LiF. We might expect reduction of Li+ at the cathode and oxidation of F– at the anode, according to the half-equations
Li+(aq) + e– → Li(s) (2a)
2F–(aq) → F2(g) + 2e– (2b)
However, Li+ is a very poor electron acceptor, and so it is very difficult to force Eq. (2a) to occur. Consequently, excess electrons from the cathode are accepted by water molecules instead:
2H2O(l) + 2e– → 2OH–(aq) + H2(g) (3a)
A similar situation arises at the anode. F– ions are extremely weak reducing agents—much weaker than H2O molecules—so the half-equation is
2H2O(l) → O2(g) + 4H+(aq) + 4e– (3b)
The overall equation can be obtained by multiplying Eq. (3a) by 2, adding it to Eq. (3b) and combining H+ with OH– to form H2O:
2H2O(l) → 2H2(g) + O2(g) (3)
The following video shows the electrolysis of water taking place, using sulfuric acidIn Arrhenius theory, a substance that produces hydrogen ions (hydronium ions) in aqueous solution. In Bronsted-Lowry theory, a hydrogen-ion (proton) donor. In Lewis theory, a species that accepts a pair of electrons to form a covalent bond. as a bridge to allow for the transfer of charge. After the electrolysis is complete, the identities of the gases formed are verified using burning splint tests.
Thus this electrolysis reverses the spontaneous combination of H2 and O2 to form H2O. In discussing redox reactions we mention several oxidizing agents, such as which are strong enough to oxidize H2O. At the same time we describe reducing agents which are strong enough to reduce H2O such as the alkali metals and the heavier alkaline earths. As a general rule such substances cannot be produced by electrolysis of aqueous solutions because H2O is oxidized or reduced instead. Substances which undergo spontaneous redox reaction with H2O are usually produced by electrolysis of molten salts or in some other solventThe substance to which a solute is added to make a solution.. There are some exceptions to this rule, however, because some electrode reactions are slower than others. Using Table 11.5, for example, we would predict that H2O is a better reducing agent than Cl–.
Hence we would expect O2, not Cl2, to be produced by electrolysis of 1 M HCl, in contradiction of Eq. (1). It turns out that O2 is produced more slowly than Cl2, and the latter bubbles out of solution before the H2O can be oxidized. For this reason Table 1 found in the Redox Couples section cannot always be used to predict what will happen in an electrolysis.