1/26 Electrochemistry: Elektrolytic and galvanic cell co08 Galvanic series (Beketov, cca 1860): Li, Ca, Al, Mn, Cr Zn, Cd Fe, Pb, [H2], Cu, Ag, Au ≈ ≈ ⊕ Cell = system composed of two electrodes and an electrolyte. electrolytic cell: electric energy chemical reaction ! galvanic cell: chemical reaction electric energy ! reversible galvanic cell (zero current) Electrodes anode = electrode where oxidation occurs Cu Cu2+ + 2 e ! − 2 Cl Cl2 + 2 e − ! − cathode = electrode where reduction occurs 2 Cu + + 2 e Cu credit: Wikipedia (free) − ! Cl2 + 2 e 2 Cl − ! − Oxidation and reduction are separated in a cell. The charge flows through the circuit. 2/26 Anode and cathode co08 electrolytic cell galvanic cell '$ '$ - - ⊕&% &% ⊕ Cl2 Cu2+ Cu2+ Pt Cl ! ! 2 Cu Cl Cl − Cu − CuCl2(aq) CuCl2(aq) anode cathode anode cathode “anions go to the anode” 3/26 Galvanic cells: electrodes, convention co08 Electrodes(= half-cells) may be separated by a porous separator, polymeric mem- brane, salt bridge. Cathode is right (reduction) ⊕ Anode is left (oxidation) negative electrode (anode) positive electrode (cathode) ⊕ liquid junction phase boundary . (porous separator) j salt bridge .. semipermeable membrane k Examples: 3 Cu(s) CuCl2(c = 0.1 mol dm ) Cl2(p = 95 kPa) Pt j − j j ⊕ Ag s AgCl s NaCl m 4 mol kg 1 Na(Hg) ( ) ( ) ( = − ) j j j 1 NaCl(m = 0.1 mol kg ) AgCl(s) Ag(s) j − j j ⊕ 2 3 4 3 3 3 Pt Sn +(0.1 mol dm ) + Sn +(0.01 mol dm ) Fe +(0.2 mol dm ) Fe j − − jj − j ⊕ 4/26 Equilibrium cell potential co08 Also: electromotive potential/voltage, electromo- tive force (EMF). Should be measured at zero-current condition (balanced bridge, sensitive voltmeter) Cannot measure a potential of one electrode zero defined by the standard hydrogen electrode): 2 H+ (aq) 2 e H (g) + − 2 ! where H+ = 1 (pH=0) p pst and H2 = 1 ( H2 = ). Construction of a hydrogen electrode: Pt sheet covered by platinum black, saturated by hydrogen credit: wikipedie symbols: E, E, Δϕ; in physics: U 5/26 Cell potential II co08 Electrode potential of electrode X = voltage of cell H2 ( = 1) H+ ( = 1) X j j ⊕ NB: always the reduction potential Standard (reduction) potential of an electrode: all reacting compounds have unit activities Examples: E 0.337 V, E 1.360 V (at 25 C) e 2+ = Cle | Cl = Cu+ |Cu 2 2 − ◦ If the reaction are written in the way the cell produces energy: reaction = (reduction at cathode) + (oxidation at anode) E E red Eox = cathode + anode If all reactions are written as a reduction: reaction = (reduction at cathode) (reduction at anode) − E Ered Ered = cathode anode − 6/26 Termodynamics of a reversible cell co08 Reversibility = reactions can be reversed by a small voltage change. No irreversible processes (metal dissolution, diffusion, liquid junction. ) G W qE zFE p, T Δr m = el = = [ ] − − Nernst equation: ) RT Y E E ln ν = e zF − where Δ G zFE , K exp Δ G /RT exp zFE /RT r me = e = [ r me ] = [ e ] − − E E ,red E ,ox E ,red E ,red e = cathodee + anodee = cathodee anodee − ΔrG < 0 i.e. E > 0 the cell produces current ) E = 0 i.e. ΔrG = 0 = uncharged cell (equilibrium) distiguish from: equilibrium potential (at zero current) E E E E e 2 e 2 (oxidation) but e 1e Cu +|Cu = Cu|Cu + Cl2|2Cl = Cl |Cl − − 2 2 − hydrogen electrode right at 25 C: E = E0 pH 0.05916 V ◦ − · 7/26 Termodynamics of a reversible cell II + co08 only electric work Wel reversible [p, T] ∂ΔrGm ∂E Δ S zF r m = ∂T = ∂T − p p ∂(ΔrGm/T) ∂(E/T) Δ H T2 zFT2 r m = ∂T = ∂T − p p Qm = TΔrSm (II. Law for reversible processes) W Oops! p-V U Q W Q p V W Δr = + = Δ.r + el − H U pV [p] U p V Q W Δr = Δr + Δr( ) = Δr + Δr = + el Eq. Q = ΔrH holds true only if the only work is pressure-volume st And similarly for standard state (p = p , = 1), e.g.: ∂ G ∂E Δr e Δ S m zF e r me = ∂T = ∂T − p p 8/26 Reduction potentials and different valences + co08 Example. E 2+ . E 3+ . e(Cr Cr) = 0 913 V, e(Cr Cr) = 0 744 V. j E −3+ 2+ j − Calculate e(Cr Cr ). j Gibbs energies are additive (not potentials) Cr2+ 2 e Cr Δ G 2F 0.913 V + − r me = ( ) ! − · − Cr3+ 3 e Cr Δ G 3F 0.744 V + − r me = ( ) ! − · − Cr3+ 1 e Cr2+ Δ G 1F E Cr3+ Cr2+ + − r me = e( ) ! − · j F E 3+ 2+ F . F . 1 e(Cr Cr ) = 3 ( 0 744 V) + 2 ( 0 913 V) − · j − · − · − E 3+ 2+ . e(Cr Cr ) = 3 ( 0 744 V) 2 ( 0 913 V) = 0 406 V j · − − · − − 9/26 Electrodes co08 of the first kind (one reaction electrode—ion) – cationic, anionic – metal, amalgam (metal in Hg), nonmetallic, gas of the second kind (nonsoluble salt—two reactions) of the third kind redox (two exidation states) ion selective (membrane) 10/26 Electrodes of the first kind co08 cationic, metal ox: Zn Zn2+, red: Zn2+ Zn not for Fe,j Al + ions, whichj are⊕ covered by oxides amalgam ox: Na(Hg) Na+, red: Na+ Na(Hg) j j ⊕ RT Na(Hg) E E ln Na+|Na = Nae +|Na F − Na+ saturated amalgam: M(Hg)=1 cationic gas: hydrogen anionic gas: chlorine, oxygen −! Cl Cl Pt : Cl 2 e 2 Cl − 2 2 + j j ⊕ − ! − OH O Pt : 1 O 2 e H O 2 OH − 2 2 2 + + 2 j j ⊕ − ! − 1 + : 2 O2 + 2 e + 2H H2O − ! similarly: Br Br Pt − 2 j j ⊕ 11/26 Electrodes of the second kind co08 silver chloride Cl AgCl Ag − j j ⊕ AgCl(s) Ag+ Cl + − Ag+ e ! Ag(s) + − ! AgCl(s) e Ag Cl + − + − ! RT Ag Cl E E − e ln · AgCl | Ag | Cl = AgCl | Ag | Cl − F − − AgCl RT E ln = e Cl AgCl | Ag | Cl F − − − mercury chloride (calomel) Cl Hg Cl Hg − 2 2 j j ⊕ Hg Cl (s) 2 e 2 Hg 2 Cl 2 2 + − + − ! Q Q Q Q Usage: reference electrodes Q Q Pictures by: Analytical chemistry an introduction, 7th edition, Harcourt College, 2000 Q QQ 12/26 Curiosity: Electrodes of the third kind + co08 2 Ca +(aq) Ca(COO)2(s) Zn(COO)2(s) Zn(s) j j j ⊕ Three reactions: Ca2+ (COO) 2 Ca(COO) + 2 − 2 Zn(COO) ! (COO) 2 Zn2+ 2 2 − + 2 ! Zn + + 2 e Zn − ! 2 Ca + + Zn(COO)2 + 2 e Ca(COO)2 + Zn − ! RT 1 E E ln = e F − 2 Ca2+ c 2+ Can measure Ca2+ —there is no Ca Ca (aq) electrode, because Ca reacts with water j (Ion-selective electrodes are more advantageous) 13/26 Redox electrodes co08 metal: Sn4+ Sn2+ Pt j j ⊕ Sn4+ 2 e Sn2+ + − ! quinhydrone electrode (pH 1–8): quinone (p-benzoquinone) + hydroquinone 1:1, sat. in a buffer O O (s) 2 H+ 2 e HO OH (s) + + − quinone (Q)! hydroquinone (QH) Nernst equation: 1 RT . RT E E QH E Q|QH = e ln = e + ln H+ Q|QH 2F 2 Q QH F − Q H+ j . · = (0.699 0.059 pH) V − · Usage: measuring pH 14/26 Ion-selective electrodes co08 Semipermeable membrane (for some ions only) Glass electrode Membrane of special thin glass permeable for H+ (and other ions). Difference of chem. pot. of both solutions μ(H+, ) μ(H+, electrode) − 1 (H+, ) credit: www.ph-meter.info = ln RT (H+, electrode) ref. silver is compensated by the electric work FE. Nernst chloride − ) % equation glass ! credit: wikipedie RT RT E = const ln (H+, ) = const ln 10 pH − F − F · · Usage: measuring pH (2–12), other ions 15/26 Galvanic cells co08 By the source of ΔG: chemical concentration – electrolyte – electrode By ion transfer: one electrolyte with salt bridge with membrane credit: payitoweb.blogspot.com anode: Fe s Fe3+ aq 3 e 0.04 V ( ) ( ) + − ( ) ! 2 − Fe(s) Fe +(aq) + 2 e ( 0.44 V) ! − − cathode: 1 O 2 e 2H+ H O 1.23 V 2 2 + − + 2 ( ) or reduction of organic! compounds (vitamin C) [xcat ev/clanekagcl.ev]16/26 Simple chemical cell co08 Single electrolyte + electrodes Example. Consider a Pt electrode saturated by hydrogen and Ag wire covered by AgCl submerged in a solution of HCl (c 0.01 mol dm 3) on the top of the highest = − Czech mountain, Sneˇzkaˇ (1602 m above sea level). The standard reduction poten- tial of the Ag/AgCl/Cl electrode is 0.222 V (pst 101 325 Pa). Calculate the cell − = voltage. A sea-level-reduced pressure is 999 mbar, temperature 25 C. ◦ .66V(DH) V 0.4616 .61V(i.DH) (lim. V 0.4621 = ) 1 ( V 0.4561 γ = Pa 83128 p 17/26 Chemical cell—separated electrolytes co08 Porous barrier (liquid junction) (.).. Irreversible liquid junction (diffusion) potential. Reduced by the) salt bridge ( ). jj Example: Zn(s) ZnSO4 CuSO4 Cu(s) j jj j ⊕ Electrode concentration cell Examples: Pt—H2(p1) HCl(aq.) H2(p2)—Pt j j ⊕ (Given polarity for p1 > p2) Li(Hg)(1) LiCl(aq) Li(Hg)(2) j j ⊕ (Given polarity for 1 > 2) 18/26 Battery co08 (One or) more connected cells. Common disposable batteries: Alkaline battery (Zn, MnO2 + C) Zn KOH (gel) MnO2 j j ⊕ Zn 2 OH ZnO H O 2 e + − + 2 + − 2 MnO H O 2e ! Mn O 2 OH 2 + 2 + − 2 3 + − ! Lithium (Li metal is light, has high potential) Electrolyte = salt (e.g., LiBF4) in organic polar solvent Several possibilities, e.g.: Li Li+ e + − MnIVO Li+ e ! MnIIILiO 2 + + − 2 ! 19/26 Rechargeable batteries co08 Li-ion, Li-polymer: Li is in C (max.