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2. Electrochemistry EP&M. Chemistry. Physical chemistry. Electrochemistry. 2. Electrochemistry. A phenomenon of electric current in solid conductors (class I conductors) is a flow of electrons caused by the electric field applied. For this reason, such conductors are also known as electronic conductors (metals and some forms of conducting non-metals, e.g., graphite). The same phenomenon in liquid conductors (class II conductors) is a flow of ions caused by the electric field applied. For this reason, such conductors are also known as ionic conductors (solutions of dissociated species and molten salts, known also as electrolytes). The question arises, what is the phenomenon causing the current flow across a boundary between the two types of conductors (known as an interface), i.e., when the circuit looks as in Fig.1. DC source Fig. 1. A schematic diagram of an electric circuit consis- flow of electrons ting of both electronic and ionic conductors. The phenomenon occuring at the interface must involve both the electrons and ions to ensure the continuity of the circuit. metal metal flow of ions ions containing solution When a conventional redox reaction occurs in a solution, the charge transfer (electron transfer) happens when the two reacting species get in touch with each other. For example: Ce4+ + Fe 2+→ Ce 3+ + Fe 3+ (2.1) In this reaction cerium (IV) draws an electron from iron (II) that leads to formation of cerium (III) and iron (III) ions. Cerium (IV), having a strong affinity for electrons, and therefore tending to extract them from other species, is called an oxidizing agent or an oxidant. Iron (II), that readily donates electrons, is called a reducing agent or a reductant. Iron (II) reduces cerium (IV) and at the same time is being oxidized. Cerium (IV) oxidizes iron (II) at the same time being reduced in the process. It is possible to separate the processes of oxidation and reduction in space. The two species do not come into contact with each other and the electron transfer is achieved via an external electrical circuit. Such an arrangement is called an electrochemical cell. An electrochemical cell in which reaction (2.1) occurs is shown in Fig. 2. 2.1. Electrochemical cells. electrons Fig. 2. In the cell shown at left, the overall reac- µA tion is the same as reaction (2.1). How- ever, oxidation of iron (II) occurs in the electrolytic bridge left beaker (half-cell) and reduction of cerium (IV) in the right one. Electrons are drawn by cerium from the platinum sheet and delivered by iron to a twin one (Pt terminals are inert in the process). 2+ 3+ 4+ 3+ - - One can write: Fe Fe +e Ce +e Ce 4+ - 3+ Pt Pt Ce + e→ Ce (2.1a) Fe2+→ Fe 3+ + e - (2.1b) 2+ 4+ Fe solution Ce solution Reactions written in the caption to Fig. 2 are called half-reactions. They can be handled in the usual way as any other chemical reactions (normal rules apply). If we add both half-reactions, we will get back the overall reaction (2.1). The electrode reaction is the phenomenon we asked at the beginning, connecting the 15 EP&M. Chemistry. Physical chemistry. Electrochemistry. flow of electrons in the external circuit with flow of ions in the electrolyte. To permit ion flow and close the circuit, solutions in both half-cells are connected by a reversed U-tube (with porous tips) filled with solution of salt, e.g., KCl or KNO3, known as an electrolytic bridge. Some cells are built of electrodes that share a common electrolyte. They are known as cells withou liquid junction. 2.1.1. Electrodes (half-cells). Anode and cathode. Half cells are also called electrodes, though, the term is frequently confused with the terminals of the external circuit (in the example shown - Pt plates). Sometimes, the material of the terminal is actually involved in the half reaction, for example: Cu2+ + Zn 0→ Zn 2+ + Cu 0 (2.2) Cu2+ + 2e -→ Cu 0 (2.2a) Zn02+-→ Zn + 2e (2.2b) when the two terminals are made of copper and zinc. The name of the half cell or the electrode should rather be associated with the half-reaction than with the material of the terminal. One can also say, with a redox couple involved. It is quite obvious, that when both constituents of the couple are ions, or one is in gaseous state, then the terminal is not involved and is made of a neutral material, like platinum or carbon (graphite, glassy carbon). By definition, the electrode at which an oxidation reaction occurs is called anode and that one at which a reduction takes place is cathode. An important electrode involving gas in the redox couple is shown in Fig.3. We will frequently comment on this example in future considerations. R ∞ Fig. 3. The half-reaction occurring in the left mV half-cell is: +- electrolytic bridge H2 (g)→ 2H (aq) + 2e (2.3a) H Note, that the reaction is written in a way 2 p=1 Atm indicating an equilibrium. Pt 4+ 3+ - Ce +e Ce Pt Standard a=1Ce4+ Hydrogen a=1Ce3+ Electrode 4+ 3+ a=1H+ Ce /Ce solution In general there are two types of electrochemical cells. Galvanic or voltaic cells are these which are capable to deliver electricity at the expense of a spontaneous reaction occurring in the cell. On the contrary, electrolytic cells are those in which we perform desired chemical changes (reactions) at the expense of an external source of electric energy (non-spontaneous, externally driven redox reactions). Both kinds have numerous practical applications. Voltaic cells are batteries (storage of chemical energy to be retrieved in a form of electric energy) and electrolytic cells are used in manufacturing processes (e.g., those of metals). 2.1.2. Conventional cell notation. A half-cell, or a redox couple is usually written as a following scheme: Oxidized form/Reduced form (2.4) + 2+ 4+ 3+ or Ox/Red for short, e.g., H /H2, Cu /Cu, Ce /Ce . A cell notation, by convention (Stockholm convention) is written in such a way, that anode is written at the left side and cathode - at the right. For example, the cell shown in Fig. 2 and in reaction (2.1) may be written as follows: 2+ 3+ 4+ 3+ Pt Fe (aq,a1 ),Fe (aq,a2 ) Ce (aq,a3 ),Ce (aq,a4 ) Pt (2.5) 16 EP&M. Chemistry. Physical chemistry. Electrochemistry. where symbol | indicates a phase boundary (interface) and symbol || indicates the liquid junction. Activities are indicated (we will simplify it using molar concentrations). Some other examples: cell corresponding to reaction (2.2) 2+ 2+ Zn Zn (aq,a1 ) Cu (aq,a2 ) Cu (2.6) cell shown in Fig.3 +4+3+ Pt H2 (g,p = 1 Atm)H (aq,a = 1) Ce (aq,a = 1),Ce (aq,a = 1)Pt (2.7) 2.1.3. Electrode reaction. Overall cell reaction. The basics of this subject have been discussed above. One of common problems, that will become important in future considerations is how to design a cell to perform a desired overall reaction. Example 2.1: Write a suitable diagram of a cell in which the following overall reaction occurs (in aqueous solution): -2++ 2+4+ 2MnO4 + 5Sn + 16H→ 2Mn + 5Sn + 8H2 O (2.8) Solution: The reaction (as written) suggests the permanganate is the oxidant (undergoes reduction) and tin (II) is the reductant (undergoes oxidation). Hence, the Sn4+/Sn2+ redox couple should be placed at the left side (oxidation → anode). 4+ 2+ - 2+ + Pt Sn (aq,a1 ),Sn (aq,a24 ) MnO (aq,a3 ),Mn (aq,a4 ),H (aq,a5 ) Pt (2.9) Example 2.2: Design a cell diagram in which the following overall reaction occurs in aqueous solution: K AgCl(s)←→s Ag+- (aq) + Cl (aq) (2.10) Solution: The reaction can be "decomposed" into two reactions (each one temporarily written as reduction): Ag+- (aq) + e→ Ag(s) (2.11) AgCl(s) + e--→ Ag(s) + Cl (aq) (2.12) One can handle the electrochemical reactions in a way analogous to that utilized when one treats thermochemical reactions using the Hess's law. However, one doesn't multiply the potentials when the reaction (stoichiometric coefficients) is multiplied. In the above case, one can reverse the first reaction and add both of them. The outcome is exactly reaction (2.10), as desired. The scheme suggests then, reaction (2.11) to be oxidation (anodic) and reaction (2.12) to be reduction. Therefore, the following scheme should be the solution to the problem: Ag|AgCl(s)|Cl-+ (aq)||Ag (aq)|Ag (2.13) We will see later, whether the above scheme represents an actual galvanic cell. In addition to the galvanic/electrolytic classification of the cells, there is another, important one. Namely, one can distinguish the reversible and irreversible cells. If, after reversing the direction of the electron flow, the overall reaction doesn't change and simply runs backward, the cell is reversible. In the irreversible cell, changing the direction of current causes an entirely different half-reaction to undergo at one or even both electrodes. Fuel cells. Let's consider the cell shown schematically below: + - Pt H21 (g,p ) H (aq,a1 ) OH (aq,a22 ) O (g,p 2 ) Pt (2.14) We see, that at the anode gaseous hydrogen is oxidized and at the cathode oxygen is reduced. This is so called fuel cell. The hydrogen/oxygen fuel cell is the source of the most clean, environment friendly energy. If hydrogen is produced and stored, then fuel cells can deliver its energy and the sole 17 EP&M.
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