Galvanic cells. Galvanic cells (2) Alessandro Giuseppe Antonio Anastasio Baron Volta. John Frederic Daniell

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A quote of the week (or camel of the week): I have no special talents.

permitting direct transformation of chemical energy (energy of chemical bonds) into electric energy (work).

whether a given device is a galvanic cell or not.

1 A quote of the week (or camel of the week): I have no special talents. I am only passionately curious Albert Einstein Physical Chemistry EPM/7 1 Galvanic cells Galvanic cells are devices permitting direct transformation of chemical energy (energy of chemical bonds) into electric energy (work). As you can see, the above definition does not say a word about the construction of the cells, while it leaves no doubt whether a given device is a galvanic cell or not. Physical Chemistry EPM/7 2 Galvanic cells (2) Alessandro Giuseppe Antonio Anastasio Baron Volta John Frederic Daniell First galvanic cell,, 18 University of Pavia, Italy Daniell cell,, 1836 King s College, London Physical Chemistry EPM/7 3

2 Galvanic cells (3) ZnSO 4 solution Zn plate anode (oxidation) Pt Zn(s) ZnSO 4 (aq) CuSO 4 (aq) Cu Pt CuSO 4 solution Cu plate cathode (reduction) Physical Chemistry EPM/7 4 Galvanic cells (4) Pt Zn(s) ZnSO 4 (aq) CuSO 4 (aq) Cu Pt Stockholm convention, cathode at the right side Shorter form may be permitted:zn(s) ZnSO 4 (aq) CuSO 4 (aq) Cu(s) Electromotive force (EMF) is voltage measured between a conductor connected to the right electrode and the same conductor connected to the left electrode at no current conditions. EMF = E cell = E cat E an Physical Chemistry EPM/7 5 Galvanic cells (5) Cathode is the half-cell, where reduction occurs. Anode is the half-cell, where oxidation occurs. From the structure of the cell, the reactions may be deduced (and vice versa). Zn(s) ZnSO 4 (aq) CuSO 4 (aq) Cu(s) Cathode (right): Cu +2 (aq) + 2e = Cu (s) Anode (left): Zn (s) = Zn +2 (aq) + 2e Overall (cell): Cu +2 (aq) + Zn (s) = Cu (s) + Zn +2 (aq) More examples will be given once you learn more types of half-cells. Physical Chemistry EPM/7 6

3 Galvanic cells (6) Left electrode in the cell shown at the right is SHE, whose potential is (by convention) equal to zero at any temperature. Therefore: cell = Ecat E E = E an cat Right electrode here is also in its standard state: Pt(black) H 2 (g,p=p ) H + (aq,a=1) Ce 4+ (aq,a=1),ce 3+ (aq,a=1) Pt Physical Chemistry EPM/7 7 Electrochemical series Why does zinc undergoes oxidation and copper ions are reduced in the Daniell cell? The answer is: these are the inherent properties of the two metals, zinc is less noble and it tends naturally to remain in ionic state rather than in metallic, while copper (a noble metal) just opposite. Noble metals cannot be dissolved in diluted solutions of acids, which means that they cannot be oxidized by hydrogen ions, while non-noble metals can be dissolved in diluted acids with evolution of gaseous hydrogen and formation of a salt. Physical Chemistry EPM/7 8 Electrochemical series (2) Zn(s) + 2H + (aq) = H 2 (g) + Zn 2+ (aq) H 2 (g) + Cu 2+ (aq) = Cu(s) + 2H + (aq) The above reactions are spontaneous, while the second one can be performed only with an active form of hydrogen (in statu nascendi). They can be, however, performed easily in electrochemical cells: Zn(s) Zn +2 (aq) H + (aq) H 2 (g) Pt(black) Pt(black) H 2 (g) H + (aq) Cu +2 (aq) Cu(s) Physical Chemistry EPM/7 9

4 Electrochemical series (3) If we measure potentials of the following cells: Pt(black) H 2 (g,p=p ) H + (aq, a=1) Zn +2 (aq,a=1) Zn(s) E cell1=,76v Pt(black) H 2 (g,p=p ) H + (aq,a=1) Cu +2 (aq,a=1) Cu(s) E cell2= +,34V where both half cells are in standard state, and E SEW=: E E cell1 = Ecat1 Ean 1 = Ecat1 = cell 2 = Ecat 2 Ean2 = Ecat 2 = +,76V,34V then we measure standard reduction potentials of the two halfcells (half-reactions) written as cathodes Physical Chemistry EPM/7 1 Electrochemical series (4) Positive standard reduction potential means that given reaction of reduction occurs spontaneously (at standard conditions), while negative one that a reverse reaction (hence, oxidation) is spontaneous. Redox couples characterized by positive standard reduction potentials are good oxidants (they tend spontaneously to their reduced state), while redox couples characterized by negative standard reduction potentials are good reductants (they tend spontaneously to their oxidized states). Physical Chemistry EPM/7 11 Electrochemical series (5) All redox couples or half reactions (formerly only metal metal ions systems were included), listed according to their standard reduction potentials at 298K, form what is known as the electrochemical series. These values are tabulated, see right, Full tables are available at the website of the Dep. of Physical Chemistry Half reaction strongest oxidant H 4 XeO 6 + 2H + + 2e XeO 3 + 3H 2 O F 2 + 2e 2F O 3 + 2H + + 2e O 2 + H 2 O S 2 O e 2SO 2 4 Ag 2+ + e Ag + Co 3+ + e Co 2+ E, V 3, 2,87 2,7 2,5 1,98 1,81 Physical Chemistry EPM/7 12

5 Galvanic cells (7) It may be demonstrated (we do not derive it, though) that G r = zfe Hence, at standard conditions (see former lectures): G r = zfe where: z is number of electrons exchanged in reaction F is Faraday constant F=N Av e =96487C (approx. 965C) These equations apply both to half rections and overall cell reactions. Physical Chemistry EPM/7 13 Nernst equation We remember, however, (lecture on chemical equilibria) that: Gr = G r + RT lnq Which means that Gibbs free energy of reaction depends on the inherent properties of the reactants and products expressed by its standard value and on actual composition of the reaction mixture expressed by the reaction quotient Q. So does the potential. If we divide both sides of the above equation by zf, we obtain E = E RT lnq RT Physical Chemistry EPM/7 14 Nernst equation (2) At 25 o C, RT/F=25,7 [mv], hence, one can write 25,7 59,2 E = E lnq[mv] E = E logq[mv] z z Nerstian equation may be applied to either overall cell reactions (and cell potentials, EMFs) or half-reactions (electrode potentials). Physical Chemistry EPM/7 15

6 Some conclusions Summarizing recent findings, one can conclude: E cell = E cat E an Calculating EMF, one should take potentials of both half-reactions as reduction potentials. If EMF is positive, then the reaction in the cell (as written) is spontaneous. If EMF is negative, then the reaction in the cell (as written) must be forced. However, if one closes the external circuit with a finite resistance, the current will flow its direction reversed to the predicted. It is rather obvious a reverse reaction (as compared with what is written) is spontaneous. When overall cell reaction is at equilibrium, the cell (battery) is exhausted G=, E=. When we measure EMF, reaction in the cell is not at equilibrium, (though it does not run), while both half-reactions are. In electrochemical process, value of Q may be maintained for long periods of time by opening the external circuit. Physical Chemistry EPM/7 16 Types of half-cells Type I half-cells are these, where metal is at equilibrium with its ions in solution (they are reversible vs cations). Basically we always write a single half-cell as a cathode. Cu 2+ (aq) + 2e = Cu (s) Ag + (aq) + e = Ag (s) Cu (s) Cu 2+ (aq) Ag (s) Ag + (aq) Hydrogen electrode is also treated as a type I electrode H + (aq) + e = ½H 2 (g) though, by construction it is a gaseous electrode E = E RT + zf z ln[me + Physical Chemistry EPM/7 17 ] Types of half-cells (2) In type II half-cells, metal is covered with a layer of solid poorly soluble salt of this metal and an anion, which is present in solution and perticipates in the equilibrium. They are reversible vs. these anions. AgI(s) + e = Ag (s) + I (aq) Hg 2 Cl 2 (s) + 2e = 2Hg (l) + 2Cl (aq) E = E RT zf Ag(s) AgI(s) I (aq) Hg(l) Hg 2 Cl 2 (s) Cl (aq) z ln[x Physical Chemistry EPM/7 18 ]

7 Types of half-cells (3) Reaction in type II half-cells may be represented as a sum of two reactions: Ag + + e = Ag (s) E 1 = +,8 V AgCl(s) = Ag + (aq) + Cl (aq) K SP = 1, AgCl(s) + e = Ag (s) + Cl (aq) E 3 = +,22 V G3 = G1 + G2 = E1 F RT ln KSP = E3 F E1 F + RT ln K IR RT E3 = = E1 + ln K SP F F Physical Chemistry EPM/7 19 Types of half-cells (4) Redox half-cells are made of an inert electronic conductor (Pt, C) immersed in a solution containing a redox couple. It means that both forms (oxidized and reduced) are present in solution and the conductor serves only either as a source or a drain of electrons. Ce 4+ (aq) + e = Ce 3+ (aq) Pt Ce 4+ (aq),ce 3+ (aq) MnO 4 (aq) + 8H + (aq) + 5e = Mn 2+ (aq) + 4H 2 O(l) Pt MnO 4 (aq),mn 2+ (aq),h + (aq) Please, note, that the potential of the second hafl-cell will depend also on ph. Physical Chemistry EPM/7 2 Types of half-cells (5) Oxide electrodes are made of a metal covered with a layer of solid oxide of this metal. HgO(s) + H 2 O(l) +2e = Hg(c) + 2OH (aq) Sb 2 O 3 (s) + 3H 2 O(l) + 6e = Sb(c) + 6OH (aq) E = E RT ln[oh ] F Hg(c) HgO(s) OH (aq) Sb(s) Sb 2 O 3 (s) OH (aq) H + or OH ions participate in the equilibria on their surfaces, hence, their potentials depend on ph and such electrodes may be used for ph measurements (and for years were actually used). Physical Chemistry EPM/7 21

8 Types of half-cells (6) Gas half-cells are made of an electronic conductor, whose surface must reveal adsorptive/electrocatalytical properties towards the gas in question. This conductor is in contact with the gas and immersed in a solution containing respective equilibrium ions (despite the hydrogen electrode these are anions). ½O 2 (g) + 2H + (aq) + 2e = H 2 O(l) Pt O 2 (g) H + (aq) ½Cl 2 (g) + e = Cl (aq) C Cl 2 (g) Cl (aq) Physical Chemistry EPM/7 22 Reference electrodes Usage of SHE as a reference electrode (left) is not practical (difficult and potentially dangerous). Therefore, other electrodes, characterized by stable and reproducible potentials are used. SCE Hg(l) Hg 2 Cl 2 (s) KCl(aq,sat) E=+,241 V Sat Ag AgCl Ag(s) AgCl(s) KCl(aq,sat) E=+,197 V,18V vs SHE? vs SCE SCE E If reference electrode other than SHE is used, it must be indicated. Physical Chemistry EPM/7 23 Potentiometry If we measure EMF of a cell: Pt(black) H 2 (g,p=p ) H + (aq,a=1) Cu +2 (aq,a=?) Cu(s) RT its potential is given as: E = E + ln a 2 +,where E = E 2 Cu / Cu 2F Cu Therefore, unknown activity (concentration) of copper(ii) ions may be calculated as: 2F( E E ) a = exp RT or at 25 o C a = 1 2( E E ),592 + Physical Chemistry EPM/7 24

9 Potentiometry (2) Last two relations are the basis of direct potentiometry. Instead of using them, however, dependence of EMF of the measuring cell (sensing electrode right, and reference electrode left) is usually determined experimentally, yielding a calibration curve, which is later used for determination of uknown concentrations. Potentiometric calibration curves - cations Potentiometric calibration curves - anions EMF, V Me Me pme EMF, V X X px Physical Chemistry EPM/7 25 Glass electrode Glass electrode (Klemensiewicz 197), is a basic modern electrode for ph measurements. This is a membrane electrode. Special glass used in the bulb can exchage protons with solutions (internal and external, measured). Internal equilibrium is always the same, while the external depends on ph of the solution in which the bulb is immersed. We measure difference of the potential jumps on both sides of the mebrane. wire Ag AgCl,1M HCl glass bulb Physical Chemistry EPM/7 26 Ion-selective electrodes Ion selective electrodes are membrane electrodes (see glass electrode) reversible versus different cations and anions and less influenced by presence of other ions in the sample then type I or type II electrodes. Many types of equilibria are utilized here: solubility of sparingly soluble salts ionic exchange complex formation reactions (e.g. calixerenes) Their potential is described by Nikolsky equation: RT E = const + ln ai + Kija nf j n / z j j Physical Chemistry EPM/7 27

10 Potentiometric titrations In titrations (see ealier lecture) before the equivalence point excess of the analyte exists, while after this point that of the titrant. Hence, concentrations of both substances changes by several orders of magnitude (titrant increases, analyte decreases). This may be detected potentiometrically using a suitable measuring cell. The shape of the titration curve obtained is shown at the right. Dashed curve indicates first derivative of the original curve. Physical Chemistry EPM/7 28

CHM 213 (INORGANIC CHEMISTRY): Applications of Standard Reduction Potentials. Compiled by. Dr. A.O. Oladebeye

CHM 213 (INORGANIC CHEMISTRY): Applications of Standard Reduction Potentials Compiled by Dr. A.O. Oladebeye Department of Chemistry University of Medical Sciences, Ondo, Nigeria Electrochemical Cell Electrochemical