Electro-Analytical Chemistry

Transcription

1 Electrochemistry involves redox systems. Therefore species amenable for analysis primarily are involved in red-ox systems. Electro-Analytical Chemistry Preliminaries Terminology: Red-ox reaction one species undergoes a loss of electrons another gains electrons. e - Fe(III) + V(II) Fe(II) + V(III) oxidant reductant Further the reaction is carried out so that oxidation and reduction occur at different locations - electrodes -in electrochemical set-ups. Fe(III) + V(II) Fe(II) + V(III) Two electrodes when coupled constitute an electrochemical cell. The electron movement; q = quantity of charge, i = rate of movement of charge. The electrical potentials of the electrodes, E el and the potential difference of the electrodes E cell that are involved and can be measured. Two types of cells are studied: a. Galvanic (Voltaic) cell: Galvanic cell uses spontaneous red-ox chemical reactions to produce electrical energy; that would result in a flow of electrons. G < 0. b. Electrolytic cell: An electrolytic cell decomposes chemical compounds by red-ox processes using electrical energy from an outside source - electrolysis. This is an energy demanding process non-spontaneous. G > 0. In any type of cell: Anode oxidation occurs Cathode reduction occurs Galvanic Cell Cd + 2 AgCl(s) Cd Ag + 2Cl - Analytical Chemistry Qualitative Analysis property characteristic to analyte E el i rate of reaction E cell G of reaction G = -nfe cell i Oxidation A reduction Quantitative analysis property related to concentration of analyte i, q q = quantity of current = i t = n F n = #moles electrons 1 mol e - = 96485C = 1F 1

2 Cd + 2Ag + Cd Ag V Half Cell Reactions: Cd (s) + 2 AgCl (s) Cd +2 (aq) + 2Ag (s) + 2Cl - (aq) Cd(s) Cd +2 (aq) + 2e anode; oxidation 2AgCl(s) + 2e 2Ag(s) + 2Cl - (aq) cathode; reduction The above two equations (half reactions) involve a physical transfer of electrons (Faradaic Process) By convention: anode (oxidation half reaction) - left Notation: Cd Cd +2 (aq) Ag +1 (aq) Ag phase boundary i, ampere; A t, sec F = coulombs/mole R resistance, C/mol E, volts, V i = q/t q = charge; C Construction of electrodes: The electrode is made out of species involved in the half reaction. If a metal is not involved, Pt provides electrical connectivity. 1 C/sec = 1 A (quite a large charge flow rate) Work (J) = Free energy from reaction = q E Coulombs volts Convention Negative terminal Left Black ANODE Convention Positive terminal Right Red CATHODE G = -nf E cell i = E/R Standard Hydrogen Electrode: SHE Standard Hydrogen Electrode: SHE E SHE = 25 0 C by definition. In any cell, when electrons move (current flows) between electrodes the potential difference drops to zero. 2

3 E cell = Difference in electrical potentials of the two electrodes. Measured E cell positive if anode connected to negative terminal, Positive terminal Red Negative terminal COM - Black ph meter is very close to an ideal voltmeter. Impedance Draws negligible current. Negative terminal - reference electrode slot. - BNC outer connector Electrodes (Equilibrium) Electrode Potentials: M(s) M + M + (aq) M(s) M + (aq) + e e if Metal acquires a negative potential w.r.t. solution Bayonet Neill Concelman connector The equilibrium set up at the electrodes is not the conventional chemical equilibrium, rather it is an electrochemical equilibrium. The potential difference developed across the interface also controls the equilibrium position of the half reaction. This type of equilibrium is also referred to as frustrated equilibrium. M(s) M + e Faradaic processes: an interfacial phenomenon. M + (aq) M(s) M + (aq) + e Depending on the position of equilibrium the metal acquires a negative/positive potential w.r.t. solution. C M+ (aq) is one determinant of the position of equilibrium. 3

4 How electrode potentials develops at an electrode: Electrode Interface : Ex: Cu/Cu +2 (aq) Electro-chemical potentials of the species Cu +2 ion is not the same in the two phases, that has not attained equilibrium. Natural tendency is to equalize the electrochemical potentials of the two species. In order to achieve such a state, ion concentration must change in the two phases at the interface. Cu(s) Cu +2 (aq) + 2e If the electrochemical equilibrium shifts to right; the excess e - would remain on the Cu metal making it more negative w. r. t. solution and vice versa. IHL OHL Bulk The higher the tendency for the oxidation process to occur, the higher would be the electron density on the metal. For cases where the oxidation is dominant it s electrode potential is more negative. The electrode potentials of red-ox systems are tabulated, relative to the standard hydrogen electrode (SHE). Standard Electrode Potentials, E o Electrode reaction Eº /V Electrode reaction Eº /V Li + + e Li AgI + e Ag + I K + + e K Sn e Sn Cs + + e Cs H + + 2e H exactly Ba e Ba AgBr + e Ag + Br Al e Al I 3 + 2e 3I Zn e Zn Fe 3+ + e Fe Ga e Ga Hg e 2Hg Fe e Fe Ag + + e Ag Cr 3+ + e Cr Hg 2+ +2e Hg Cd e Cd Pd e Pd V 3+ + e V e 2Cl Ni 2+ + e Ni Au e Au Equilibrium Electrode potential: The magnitude of the electrode potential depends on the excess charge that exists above that of the metal alone. If the excess of (negative charge) electrons is present in the metal of the electrode the potential of the electrode is negative, with the energy of the electrons in the electrode high and vice versa. An external power supply is capable of forcing an excess (or a depletion electrons) from the metal of the electrode, could lead to a non-equilibrium condition at the interface - electrolysis. In electrolysis, there is a net reaction forced by the applied power source later topic. 4

5 Absolute individual electrode potentials cannot be measured, only potential differences; i.e. only relative values can be measured. Electrode potentials are measured against a standard electrode and tabulated ; electrode potential of the standard hydrogen electrode (SHE) is defined as 0.00V at 25 0 C. Potentiometer Pt H 2 (g) (p=1atm) H + (aq) (a=1) - + E test or E el SHE Test electrode Potentiometer measures the potential difference between two electrodes. E cell = E RHS E LHS E cell = E test E LHS E cell = E test E SHE E cell = E test E cell = E test = E el = 0, definition for this set up. Measured E cell for this set up = Electrode potential of test electrode Hg(EDTA) -2 (aq, 0.005M)+2e Hg(l)+EDTA -4 (aq,0.015m) p H2 =1atm SHE a H =1 Potentiometer Standard Electrodes and Standard Electrode Potential E o el: When the activity (~concentration) of all species involved in the half reaction are unity standard electrode. Potential of such electrodes are defined as it s standard electrode potential. Potentiometer - + SHE (Test) Standard electrode High Impedance Voltmeter Electrode reaction Eº /V Electrode reaction Eº /V Li + + e Li AgI + e Ag + I K + + e K Sn e Sn Cs + + e Cs H + + 2e H exactly Ba e Ba AgBr + e Ag + Br Al e Al I 3 + 2e 3I Zn e Zn Fe 3+ + e Fe Ga e Ga Hg e 2Hg Fe e Fe Ag + + e Ag Cr 3+ + e Cr Hg 2+ +2e Hg Cd e Cd Pd e Pd V 3+ + e V e 2Cl Ni 2+ + e Ni Au e Au Note: Large and negative electrode potential means less tendency for reduction, tendency is to oxidize. Electrode potential, by convention is a measure of the ability to undergo reduction. 5

6 E - E V>0 + Reduction assumed To occur at test electrode If so E o el = positive a=1 Hg(EDTA) -2 (a=1), EDTA(a=1)/Hg Electrode potential is an interfacial phenomenon. Separation of charges across metal/solution interface brings about the potential difference between the solution and metal. The position of equilibrium governed by the activities (concentrations) of species (by way of reaction quotient Q) involved in the half reaction; determines the electrode potential and the inherent tendency to undergo reduction. Calculation of Electrode Potential: By convention electrode potentials are expressed for reduction reactions. If potential is positive, all what it means is that the reduction reaction of the test electrode has a higher propensity to happen compared to the reduction at the standard hydrogen electrode. Calculation of Electrode Potential (Nernst equation): 0 RT Eel E el ln 25 o C; volts el nf R = J/K mol T = temperature, K n= number of electrons involved in half reaction F = C Q el = reaction quotient in activity pure liquids, solvents, solids; a=1 6

7 Calculation of Electrode Potential: Electrode potentials are calculated for the reduction process as reduction potentials Write the half reaction as a reduction reaction, balanced in mass and charge. Write the expression for Q, determine the # electrons involved; Substitute in the Nernst Equation. E E.g. MnO 4- +8H + +5e = Mn H 2 O 0 RT a 2 Mn E 2 E 2 ln MnO4 Mn MnO4 Mn 5F a a E / / 8 MnO4 H RT ln Q nf 0 el el el Tabulation Calculation of Electrode Potential (Nernst equation): Eg. Hg(EDTA) 2- (aq)+ 2e = Hg(l) + EDTA 4- (aq) RT a a RT a E E E 2F a 2F a ln Hg EDTA ln EDTA el el el 2 2 HgEDTA HgEDTA Calculation of Cell Potential: Cell potentials are calculated for redox reaction. Write the reaction, balanced. Write the expression for Q of the reaction, recognize the # electrons involved; for the reaction. Substitute in the Nernst Equation. E RT nf 0 cell Ecell ln Qrx n Calculation of E 0 cell MnO 4- +8H + +5 Fe +2 = Mn Fe H 2 O Tabulated Electrode Elect Pot Reaction Pot EE MnO4 8H 5 emn 4 H2O 2 2 MnO4 Mn MnO4 / Mn E-E/ Fe Fe e E 0 cell = E 0 red E 0 oxd E 0 cell = E 0 + E 0 - Fe / Fe Fe / Fe E 0 cell = E 0 MnO4-/Mn+2 E 0 Fe+3/Fe+2 Calculation of Cell Potential: Eg. MnO 4- +8H + +5e = Mn H 2 O Fe +2 = Fe +3 + e Overall (all aqueous species) MnO 4- +8H + +5 Fe +2 = Mn Fe H 2 O E cell RT E E ln Q nf E 0 cell cell rxn 0 cell RT a a ln 5F a a a Mn Fe MnO4 H Fe E cell can be calculated by calculation each electrode potential (reduction) separately and subtracting the potential at oxidation half from the reduction half. E cell E cell = E red process E oxd process E cell = E + E - 0 RT Ecell Ecell ln Qrxn nf 5 0 RT a 2a 3 Mn Fe Ecell ln 5F a a a MnO4 H Fe 7

8 Alternate view of some electrodes: Electrons move from less positive electrode potential electrode to more positive electrode potential electrode, when connected by a conductor. Ag/AgCl(s), HCl(aq) AgCl(s) + e = Ag(s) + Cl - (aq) Ag + (aq) + e = Ag(s) E E log a 1 0 el AgCl / Ag Cl a E E log E log 1 1 K 0 0 el Ag/ Ag Ag/ Ag aag Cl sp, AgCl Alternate view of some electrodes: Pb/PbF 2 (s), HF(aq) PbF 2 (s) + 2e = Pb(s) + 2F - (aq) Pb +2 (aq) + 2e = Pb(s) E E log a el PbF 2/ Pb F K E E log E log 0 0 sp, PbF2 el Pb2/ Pb Pb2/ Pb 2 2 apb 2 af Use of electrodes in Analytical Chemistry The fact that the electrode potential (and the cell potential) is dependent on the concentration of species allows the use of electrodes as chemical probes. The electrode potential at an electrode measures the propensity of a reduction reaction to occur; at the concentrations (activities) of the species that has attained an electrochemical equilibrium at the interface. Not a viable cell configuration. Zn Cu Zn (aq) Cu (NO 3 ) 2 (aq) 8

9 Electrode Interface : Study of charge transfer reactions is the goal of most techniques electrochemistry. The interfacial system of is very complex and even in the absence of electron transfer and processes other than electron transfer do occur. Such processes can affect the electrical double layer and therefore the electrode behavior. The SHE is cumbersome to construct. Other half-cells are being used as secondary standards. Such processes include phenomena such as adsorption, desorption, and charging of the interface as a result of changing electrode potential. These are called non-faradaic processes. Silver/Silver chloride/kcl reference AgCl(s) +e- = Ag(s) + Cl - (aq) Silver/silver chloride: AgCl(s) +e - = Ag(s) + Cl - (aq) 25 vs. SHE vs. SCE Ag/AgCl, KCl (0.1M) Ag/AgCl, KCl (3M) Ag/AgCl, KCl (3.5M) Ag/AgCl, KCl (sat'd) Ag/AgCl, NaCl (3M) Ag/AgCl, NaCl (sat'd)

10 silver/silver-chloride electrode Calomel: Hg 2 (s) + 2e - = 2Hg (l) + 2Cl - (aq) Potential vs. SHE / V t / C 3.5 mol dm -3 sat. solution Hg Hg 2 (s) KCl, Hg 2 (aq,sat), KCl(s) frit Calomel: Hg 2 (s) + 2e - = 2Hg (l) + 2Cl - (aq) 25 vs. SHE vs. SCE, KCl (0.1M) , KCl (1M) NCE (Normal Calomel) , KCl (3.5M) , KCl (sat'd) SCE (Sat'd Calomel) , NaCl (sat'd) SSCE Mercury/mercurous sulfate: Hg 2 (s) + 2e - = 2Hg (l) + -2 (aq) 25 vs. SHE vs. SCE, H 2 (0.5M) , H 2 (1M) , K 2 (sat'd)