Presentation on theme: "Electrochemistry Electrochemical Cells Voltaic Cells"— Presentation transcript:
1 Electrochemistry Electrochemical Cells Voltaic Cells Standard Cell PotentialsEffect of Concentration on Cell PotentialsFree Energy and Cell PotentialBatteriesCorrosionElectrolytic CellsStoichiometry of Electrochemical ReactionsPractical Application: pH Electrode
2 Types of electrochemical cells Galvanic or VoltaicThe ‘spontaneous’ reaction.Produces electrical energy.ElectrolyticNon-spontaneous reaction.Requires electrical energy to occur.For reversible cells, the galvanic reaction can occur spontaneously and then be reversed electrolytically - rechargeable batteries.
3 Types of electrochemical cells Not all reactions are reversible.Examples of non-reversible reactionsIf a gas is produced which escapes.2H+ + 2 e H2 (g)If one or more of the species decomposes.
4 Voltaic cellsThere are two general ways to conduct an oxidation-reduction reactionMixing oxidant and reductant togetherCu2+ + Zn(s) Cu(s) + Zn2+This approach doesnot allow forcontrol of the reaction.
5 Voltaic cells Electrochemical cells Each half reaction is put in a separate ‘half cell.’ They can then be connected electrically.This permits better control over the system.
6 Voltaic cells Cu2+ + Zn(s) Cu(s) + Zn2+ Electrons are transferred from one half-cell tothe other usingan external metalconductor.CuZnZn2+Cu2+
7 Voltaic cells To complete the circuit, a salt bridge is used. e- e-
8 Voltaic cells KCl Salt bridge Allows ion migration in solution but prevents extensive mixing of electrolytes.It can be a simple porous disk or a gel saturated with a non-interfering, strong electrolyte like KCl.KClCl-K+Cl- is releasedto Zn side as Znis converted to Zn2+K+ is releasedas Cu2+ isconverted to Cu
9 Voltaic cells For our example, we have zinc ion being produced. This is an oxidation so:The electrode is- the anode- is positive (+).Zn Zn e-
10 Voltaic cells For our other half cell, we have Cu2+ + 2e CuFor our other half cell, we havecopper metal being produced.This is a reduction so:The electrode is- the cathode- is negative (-)
11 Zn | Zn2+ (1M) || Cu2+ (1M) | Cu Cell diagramsRather than drawing an entire cell, a type of shorthand can be used.For our copper - zinc cell, it would be:Zn | Zn2+ (1M) || Cu2+ (1M) | CuThe anode is always on the left.| = boundaries between phases|| = salt bridgeOther conditions like concentration are listed just after each species.
12 Cell diagrams Other examples Pt, H2 (1atm) | H+ (1M) || This is the SHE. Pt is used to maintain electrical contact so is listed. The pressure of H2 is given in atmospheres.Pt, H2 (1atm) | HCl (0.01M) || Ag+ (sat) | AgA saturated silver solution (1.8 x 10-8 M) based on the KSP AgCl and [Cl-]
13 Electrode potentialsA measure of how willing a species is to gain or lose electrons.Standard potentialsPotential of a cell acting as a cathode compared to a standard hydrogen electrode.Values also require other standard conditions.
14 Standard hydrogen electrode Hydrogen electrode (SHE)The ultimate reference electrode.H2 is constantly bubbledinto a 1 M HCl solutionPt | H2 (1atm), 1M H+ ||Eo = VAll other standard potentialsare then reported relative to SHE.H2Pt blackplate1 M HCl
15 Electrode potentialsStandard potentials are defined using specific concentrations.All soluble species are at 1 MSlightly soluble species must be at saturation.Any gas is constantly introduced at 1 atmAny metal must be in electrical contactOther solids must also be present and in contact.
16 Electrode potentials The standard potential for: Cu2+ + 2e- Cu (s) is V.This means that:If a sample of copper metal is placed in a M Cu2+ solution, we’ll measure a value of 0.337V if compared to:2H+ + 2e H2 (g)(1 M) (1 atm)
17 Half reactionsA common approach for listing species that undergo REDOX is as half-reactions.For 2Fe3+ + Zno(s) = 2Fe2+ + Zn2+Fe3+ + e Fe2+ (reduction)Zno(s) Zn2+ + 2e- (oxidation)You’ll find this approach useful for a number of reasons.
18 Half reactionsTables are available which list half reactions as either oxidations or reductions.Will provideStandard Eo values to help predict reactions and equilibria.Other species that participate in the reaction.Show the relative ability to gain or loss electrons.
19 Half reactions standard reduction potentials Half reaction Eo, VF2 (g) + 2H+ + e- 2HF (aq)Ce4+ + e Ce3+ (in 1M HCl) 1.28O2 (g) + 4H+ + 4e H2O (l)Ag+ + e Ag (s)2H+ + 2e H2 (g)Fe2+ + 2e Fe (s)Zn2+ + 2e Zn (s)Al3+ + 3e Al (s)Li+ + e Li (s)
20 Cell potentialsOne thing that we would like to know is the spontaneous direction for a reaction.This requires that we determine the Ecell.Since our standard potentials (E o) are commonly listed as reductions, we’ll base our definitions on that.Ecell = Ehalf-cell of reduction - Ehalf-cell of oxidationEocell = Eohalf-cell of reduction - Eohalf-cell of oxidation
21 Cell potentialsYou know that both an oxidation and a reduction must occur.One of your half reactions must be reversed.The spontaneous or galvanic direction for a reaction is the one where Ecell is a positive value.The half reaction with the largest E value will proceed as a reduction.The other will be reversed - oxidation.
22 Cell potentials For our copper - zinc cell at standard conditions: Eo redCu2+ + 2e Cu (s) VZn2+ + 2e Zn (s) VEcell VSpontaneous reaction at standard conditions.Cu2+ + Zn (s) Cu (s) + Zn2+
23 Concentration dependency of E Eo values are based on standard conditions.The E value will vary if any of the concentrations vary from standard conditions.This effect can be experimentally determined by measuring E versus a standard (indicator) electrode.Theoretically, the electrode potential can be determined by the Nernst equation.
24 Concentration dependency of E The Nernst equationFor Aa + ne BbE = Eo lnwhere: E o = standard electrode potentialR = gas constant, J/omolT = absolute temperatureF = Faraday’s constant, Cn = number of electrons involveda = activitya Aaa BbR Tn F
25 Concentration dependency of E If we assume that concentration is proportional to activity and limit our work to 25 oC, the equation becomes:E = E o logThis also includes a conversion from base e to base 10 logs.0.0592n[B]b[A]a
26 Concentration dependency of E ExampleDetermine the potential of a Pt indicator electrode if dipped in a solution containing 0.1M Sn4+ and 0.01M Sn2+.Sn4+ + 2e Sn2+ Eo = 0.15VE = 0.15V log= 0.18 V0.059220.01 M0.1M
27 Concentration dependency of E Another exampleDetermine the potential of a Pt indicating electrode if placed in a solution containing 0.05 M Cr2O72- and 1.5 M Cr3+, if pH = 0.00 (as 1 M HCl).Cr2O H+ + 6e Cr3+ + 7H2O (l)E o = 1.36 V
28 Concentration dependency of E E = E o log= 1.36 V log= 1.31 V0.05926[Cr3+]2[Cr2O72-][H+]140.05926(1.5)2(0.05)(1)14
29 Calculation of cell potentials To determine the galvanic Ecell at standard conditions using reduction potentials:Ecell = E ohalf-cell of reduction - E ohalf-cell of oxidationWhereEhalf-cell of reduction - half reaction with the larger or least negative E o value.Ehalf-cell of oxidation - half reaction with the smaller or more negative E o value.
30 Calculation of cell potentials At nonstandard conditions, we don’t know which will proceed as a reduction until we calculate each E value.Steps in determining the spontaneous direction and E of a cell.Calculate the E for each half reaction.The half reaction with the largest or least negative E value will proceed as a reduction.Calculate Ecell
31 Calculation of cell potentials ExampleDetermine the spontaneous direction and Ecell for the following system.Pb | Pb2+ (0.01M) || Sn2+ (2.5M) | SnHalf reaction EoPb2+ + 2e- Pb VSn2+ + 2e- Sn VNote: The above cell notation may or may not be correct.
32 Calculation of cell potentials Pb2+ + 2e Pb VSn2+ + 2e Sn VAt first glance, it would appear that Pb2+ would be reduced to Pb. However, we’re not at standard conditions.We need to determine the actual E for each half reaction before we know what will happen.
33 Calculation of cell potentials For lead:E = log= VFor tin:E = log= VUnder our conditions, tin will be reduced.0.0592210.010.0592212.5
34 Cell potential, equilibrium and DG We now know that changing concentrations will change Ecell. E is a measure of the equilibrium conditions of a REDOX reaction. It can be used to:Determine the direction of the reaction and Ecell at non-standard conditions.Calculate the equilibrium constant for a REDOX reaction.
35 Equilibrium constants At equilibrium EA = EB so0.0592nmEoA -log[ARED]n[AOX]n=EoB -[BRED]m[BOX]mE oB - E oA =0.0592nmlog[AOX]n[BRED]m[ARED]n[BOX]mK when atequilibrium,Q if not.log K =nm(E oB - E oA)0.0592A - species reducedB - species oxidized
36 Free energy and cell potential Earlier, we explained that DG and the equilibrium constant can be related. Since Ecell is also related to K, we know the following.Q DG EForward change, spontaneous < KAt equilibrium = K 0 0Reverse change, spontaneous > K
37 Batteries Portable voltaic cells These have become important to daily life.Dry cellsAll chemicals are in the form of a paste or solid. They are not really dry.Wet cellsA liquid solution is present.
38 Electrochemical reaction Zinc-carbon dry cellThe electrolyte, aqueous NH4Cl is made into a paste by adding an inert filler.Electrochemical reactionZn(s) + 2MnO2 (s) + 2 NH4- (aq)Zn2+ (aq) + Mn2O3 (s) + 2NH3 (aq) + H2O (l)This cell has a potential of 1.5 V when new.
40 Lead storage batteryThese are used when a large capacity and moderately high current is need.It has a potential of 2 V.Unlike the zinc-carbon dry cell, it can be recharged by applying a voltage.Car battery.This is the most common application.Most cars are designed to use a 12 V battery. As a result, six cells connected in a series are needed.
41 Lead storage battery Electrochemical reaction. 2PbSO4 (s) + 2H2O (l) Pb (s) + PbO2 (s) + 2H+ (aq) + 2HSO4- (aq)Note.Lead changes from a +2 to 0 and +4 oxidation state when a lead storage battery is discharged.Lead also remains in a solid form.
42 Lead storage battery A series of 6 cells in series are used to produce the12 volts thatmost carsrequire.
43 Corrosion Deterioration of metals by oxidation. Example. Rusting of iron and steel.EoAnode: Fe (s) Fe2+ + 2e VCathode: O2(g) + 2H2O(l) + 4e OH VRusting requires both oxygen and water.The presence of an acid enhances the rate of corrosion - more positive cathode.Cathode: O2(g) + 4H+(aq)+ 4e H2O(l) V
44 RustingO2 from airO2Fe2+Water drope-RustCathodeFeAnodeIron
45 Corrosion prevention Another example. Quite commonly a rod of magnesium is placed in a hot water tank.It will be oxidized to Mg2+ instead of the iron tank rusting.This greatly extends the life of the tank.Sacrificial anodePieces of reactive metal that are connected to an object to be protected by a conductor.
46 Electrolytic cells With voltaic cells, reactions occur spontaneously. With electrolytic cells, a potential is applied, forcing a reaction to go.- work is done on the system.- polarize the cell.- causes unexpected things to happen.- Ecell will change during the reaction.
47 Applying a voltageWhen we apply a voltage, it can be expressed as the following:Eapplied = Eback + iRWhereEback = voltage required to ‘cancel out’ the normal forward or galvanic reaction.iR = iR drop. The work applied to force the reaction to go. This is a function of cell resistance.
48 Applying a voltage Eback Increases as the reaction proceeds Actually consists of:Eback = Erev (galvanic) + overpotentialOverpotentialAn extra potential that must be applied beyond what we predict from the Nernst equation.
49 Overvoltage or overpotential A cell is polarized if its potential is made different than its normal reversible potential - as defined by the Nernst equation.The amount of polarization is called the overpotential or overvoltage. = E - Erev
50 Overvoltage or overpotential There are two types of .Concentration overpotential.This occurs when there is a difference in concentration at the electrode compared to the bulk of the solution.This can be observed when the rate of a reaction is fast compared to the diffusion rate for the species to reach the electrode.
51 Overvoltage or overpotential Concentration overpotential.Assume that we are electroplating copper.As the plating occurs, copper is leaving the solution at the electrode.This results in the [Cu2+] being lower near the electrode.[Cu2+]electrode[Cu2+]bulk
52 Overvoltage or overpotential Activation overpotentialResults from the shift in potential at the electrode simply to reverse the reaction.This effect is at its worst when a reaction becomes nonreversible.Effect is slight for deposition of metals.Can be over 0.5V if a gas is produced.Occurs at both electrodes making oxidations more ‘+’ and reductions more ‘-’.
53 Electrolytic cells In electrolytic cells The reaction requiring the smallest applied voltage will occur first.As the reaction proceeds, the applied E increases and other reactions may start.Lets look at an example to determine if a quantitative separation is possible.
54 Electrolytic exampleCan Pb2+ be quantitatively be separated from Cu2+ by electrodeposition?Assume that our solution starts with 0.1M of each metal ion.We’ll define quantitative as only 1 part in cross contamination (99.99%)Cu2+ + 2e = Cu Eo = VPb2+ + 2e = Pb Eo = V
55 Electrolytic example Copper We start with 0.1 M and begin our deposition. We don’t want any lead to deposit until at least 99.99% of the copper has been removed M Cu2+E = logE = V0.05922110-5
56 Electrolytic example Lead Pb would start depositing at: E = logE = VThe separation is possible but our calculations neglect any overpotential.0.0592210.1
57 Stoichiometry of electrochemical reactions Faraday determined that the the amount of product formed was proportional to the quantity of electricity transferred.A coulomb (C) is a quantity of electricity. Current is the rate of electrical flow.coulombs of electricity are are equivalent to one mole of electronscoulombs = 1 Faraday (F )Current = Amps = i = C / s
58 Stoichiometry of electrochemical reactions The number of equivalents deposited can be found by:gramsgram equivalent weightequivalents =g x e in transferformula weight=coulombs96 500=i t96 500=
59 Stoichiometry of electrochemical reactions The number of grams deposited then is:gdeposited =Where i = current in ampst = time in secondsFM = formula massn = number of electrons transferred per species( )i t FMnequivalentweight
60 ExampleDetermine the number of grams of Cu that could be converted to Cu2+, if a current of 6 A is applied for 5 minutes.Half reactionCu2+ (aq) + 2 e Cu (s)g == g(6 A) (5 min x ) ( )(96 500) ( 2e-)smingmol
61 ElectrogravimetryOne practical application of electrolysis is the method of electrodeposition.A quantitative analysis based on weight gain.It relies on the production of a metal or metal oxide on an electrode.The weight of the electrode is measured both before and after the material is deposited.The amount of material is determined by difference.
62 Electrogravimetry R - potentiometer A - ammeter V - Voltmeter Anode Pt cathodeStirbar
64 ElectrogravimetryOnly a limited number of species work well with electrodeposition.Cathode electrodepositions.Deposited from simple cations: Cu, Ni, ZnDeposited from cyanide complexes: Ag, Cd, AuAnode electrodepositionsDeposited as oxides.Pb PbO2Mn MnO2
65 pH electrode We can use one half of an electrochemical cell to measure properties of the other half.ReferenceelectrodeThe partof the cellthat is heldconstantIndicatorelectrodeThe part ofthe cell thatcontains thesolution weare interestedin measuring
66 pH electrode The earlier example would be too difficult for routine use.We can ‘repackage’a half cell in the formof an electrode.pH electrode- first discovered- still the most significant- relies on a glass wall or membrane.Ag wire0.1M HClAgClthinglasswall
67 pH electrode Combination pH electrode A reference electrode is inside the pHelectrode.
68 How a pH electrode works H3O+ partially populates both the inner and outer SiO2 surfaces of the glassmembrane.The concentration difference results in a potential across the glass membrane.A special glass is used:22% Na2O, 6% CaO, 72% SiO2