MTRL 358Electrowinning and Electrorefining 2014 Metal recovery is the final step in most hydrometallurgical processes. This is commonly practiced for aluminum, copper, zinc, nickel, cobalt and gold. (Aluminum cannot be obtained by electrolysis from water; it is too strongly reducing. Molten salt electrolysis is used instead.) In all hydrometallurgical processes metals are present in solution as complexes of the metals in positive oxidation states, e.g. [Zn(H2O)6]+2. All metal recovery processes then necessarily involve reduction. Hence all these processes require a reducing agent. The process may be thermodynamically favourable (like hydrogen gas reduction of [Ni(NH 3)n]+2 complexes (n = 2, for instance) or unfavourable (like electrowinning of copper in which water is forced to be the reducing agent). When the process is thermodynamically unfavourable (E < 0, by definition) it is categorized as an electrolysis. Electrolysis for metal production is called electrowinning. Briefly, a copper EW plant may have many cells (hundreds). Each cell is a little over 1 m wide, ~1.5-2 m deep and several meters long. They contain several dozen cathodes and the same number + 1 anodes. Metal is plated onto both sides of the cathode sheets, while water is oxidized to form O 2 and H+ at the anodes. A schematic illustration of a cell is shown in the diagram below. Enriched electrolyte supplied from solvent extraction stripping is fed into the cells. It passes through a cell once and then is returned to SX stripping as the lean electrolyte. Once the copper has been plated to a thickness of about 0.5 cm, the cathodes are removed from the cell and the copper is prepared for sale. Figure 1. Schematic illustration of a copper electrowinning cell. 1 Pyrometallurgically produced metals are often not sufficiently pure to be sold as high purity products. They are usually further refined and often using electrolysis. Molten, as-produced metal is cast into electrodes (~1 m x ~1 m) and these are interleaved with metal sheets in cells. The cast, impure metal electrodes are anodically polarized to electrochemically corrode them (a form of leaching), while the interleaved sheets are cathodically polarized to plate out the dissolved metal ions. A very pure metal product is formed. This is called electrorefining. Background Electrochemistry The relevant electrochemistry was developed in the Eh-pH diagram course notes and should be consulted. Reminder on calculating E° or E: The potential difference or voltage generated by an electrochemical cell at a certain temperature is strictly a function of the composition of the cell, i.e. activities of the reactants and products. It does not depend on the charge that passes through that potential (i.e. nF). The energy associated with passage of the charge through the potential difference does depend on the amount that is passed: Energy = voltage x charge. But, the potential itself generated by the cell has nothing to do with the charge that is passed. Therefore DO NOT multiply E°’s by n numbers to calculate E° or E for a cell! Electrowinning Equations (Faraday's Law Relationships) (1) Farady’s Law Faraday’s law states that the number of moles of metal produced in an electrolysis is directly proportional to the charge passed. The constant of proportionality is nF, where n = moles of electrons per mole of metal produced (an integer) and F is the Faraday which is 96485 C/mole e -; a mole of electrons has a charge of 96485 C. (i.e. 6.02205 x 10 23 e-/mol e- x 1.60218 x 10-19 C/e-.) Taking into account the fact that charge, q = current x time for fixed current (or the integral of I vs. t for a varying current) and that the moles of metal produced = mass/atomic weight leads to the formula provided. Moles of metal plated q (charge passed in the electrolysis) {1} moles metal = nM = q/nF {2} The units of q/nF are C/(mole e-/mol metal x C/mole e-) = mol metal. Charge passed at constant current is q = It. Moles of metal = M/AW where M is the mass of metal plated and AW is the atomic weight in g/mol. Then moles metal plated is, nM = It/nF = M/AW {3} 2 Rearranging gives, M = It AW nF {4} From Faraday’s law it is obvious that the lower n is, the less electricity that will be required per unit mass of metal plated. Some metal ions have more than one oxidation state. For typical copper electrowinning the cathodic half reaction is, Cu+2 + 2e- = Cu (1) Alternative leaching processes have been developed that form cuprous complexes: Copper sulfides air, Cl- CuCl2-aq (2) In this case the cathodic half reaction is, CuCl2- + e- = Cu + 2Cl- (3) which would use half the electricity. (No such process is currently commercially applied.) (2) Current Efficiency The simple formula above determines mass of metal plated for a given current and time, or vice versa. If the only cathodic (reduction) process operative is metal ion reduction to metal, then the formula gives an accurate indication of the mass of metal for a given current and time. However, other reduction half reactions may also occur simultaneously. These unwanted side reactions also consume electricity (current) and result in a lowered efficiency of use of current for metal plating. This leads to the idea of current efficiency. Current efficiency (CE) for metal plating then (or any electrolytic process), is the ratio of actual mass of metal plated to the theoretical mass based on Faraday’s law. It is usually given in %. CE = actual mass metal plated x 100 theoretical mass expected {5} The theoretical mass is given by Faraday’s law. CE = 100M It AW/nF = 100nFM It AW {6} Now M is actual mass of metal plated. For instance, calculate the current efficiency for the following conditions: 100 g of copper was plated in a copper electrolysis experiment using a constant current of 4 A for 23 hours. What was the current efficiency? 3 4 . reversible. Energy efficiency is the ratio of the theoretical energy required to the actual energy required. (an electrolysis) is applied. For practical rates we require Eappl > E . where We’rev and we'rev are the electrical work under reversible conditions.7% (23 hr x 3600 sec/hr x 4 C/sec x 63. The current is forced to go in the opposite direction to the natural tendency of the cell. this is the difference between we'rev and We'rev. E < 0). and due to less than 100% current efficiency the charge passed may be greater than the theoretical minimum.CE = 100 g Cu x 100 x 2 mol e-/mol Cu x 96485 C/mole e. The applied voltage opposes the thermodynamic voltage and hence it is taken to be positive. ∫Vdq if the voltage varies with charge passed. The actual applied voltage is designated Eappl. in percent. The theoretical charge required is given by Faraday’s law. i. For now suffice it to say that if an opposing voltage equal to E .= 91. By definition 1 VC = 1 J (1 volt·coulomb = 1 joule). The symbol W e'rev represents the work in joules. i. -G = nFE = we'rev {10} units in J/mol.* In electrowinning E is negative (G > 0. the thermodynamic cell voltage. q = nMnF {8} Hence the minimum energy requirement is. the reaction as written is not favourable so we'rev < 0. which equals Vq at constant voltage). then the thermodynamic tendency is just matched or just overcome and the reaction is exceedingly slow.546 g Cu/mol Cu) {7} (3) Energy Efficiency There is a theoretical minimum energy required to electroplate a metal.e. The theoretical voltage is E. On both counts the energy required will be greater than the theoretical limit. The units of w e'rev are J/mol. Electrical work (energy) is voltage times charge (more generally.) CE = 100 nFM = 100 nFnM It AW q {11} * Reversible and irreversible processes are reviewed in the next section. In practice the actual voltage will be greater than the thermodynamic minimum for a number of reasons. We’rev = E nMnF units in J {9} (Work and cell thermodynamic voltage are related by. work done on a system is negative. where E < 0. This is the thermodynamic minimum voltage times the charge passed at 100% current efficiency.e. As per the engineering convention. 6 x 10 6 J.) The energy efficiency is: 100 E nMnF 100Eappl nMnF/CE EE = 100We’rev = We’ = E CE Eappl {14} Take an example again of 100 g of copper plated as above at a voltage of 2. J/mol).) EE = 0.e. as will be the current.89 V = the necessary applied voltage to just overcome the thermodynamic negative cell voltage. (If the cell is large and relatively little copper is plated. i. A kWh is 1000 watts for 1 hour = 1000 J/sec x 3600 sec = 3. To convert the energy consumption number to kWh/t metal involves only unit conversions: -w’e J x 1 kWh x 1 mol metal x 10-6 g {18} 5 .7% current efficiency.0 = 40. The derivation above employed the energy consumption. q = 100nMnF CE {12} Then the actual energy input is. the applied voltage will be about constant. (In reality the Nernst equation E would be required for a real cell with non-standard activities.Since M/AW = nM and It = q. Eappl q.7/ 2. -We’ = 100Eappl nM nF CE {16} Divide both sides by the moles of metal (n M here) to get the specific energy consumption in J/mol (designated -we’): -w’e = 100 nFEappl CE {17} A more conventional unit is kilowatt-hours per tonne of metal.8% {15} This is not very high. 0. (4) Specific Energy Consumption This is the actual energy requirement in units of energy per unit amount of metal plated (e.89 x 91. Reasons for this will be explained later. -We’ = Eappl q = 100EapplnMnF CE {13} (-We' is a positive number.89 V (= E° assuming standard conditions.0 V with 91.g.) E = -0. J/mol For copper electrowon as above at 2 V and 91. Only CE% (e.7%) of the total charge passed goes to plate metal. where Ac is the surface area per cathode sheet. Next. the current being passed is j times the total plating area.mol 3.e.g. I = jA cN.e.7 x 0.0043713 kWh mol / J t = 1840 kWh/t Cu {19} {20} In practice a typical energy requirement for copper EW is about 1900-2000 kWh/t. j (C/sec m2) x surface area (m2) x time (sec) = charge (C). q = 100nMnF CE {21} nM = {22} q CE 100nF where nM is the moles of metal produced. The charge at constant current is q = It.6 x 106 J AW g t For copper this works out to -we’ J/mol x 0.209 x 105 J/mol x 0. rearrange to obtain: nM = dnM = jAcN CE t dt 100nF mol/sec {25} AcN is the total plating surface area.0043713 kWh/t. nM = It CE 100nF {23} Metal is plated onto both sides of a cathode starter sheet (e.01 = 4.7% current efficiency. or with N being the number of plating surfaces and A c being the surface 6 . nM = jAcNt CE 100nF {24} i.209 x 105 J/mol Cu = 4. (5) Metal Production Rate Starting with Faraday’s law and current efficiency again.x 96485 C x 2. The total plating surface area for a number N cathodes is AcN. i. -w’e = 2 mol e. a steel sheet in copper electrowinning). 91. Taking j as the current density in A/m2.g. This may be obtained in two ways. either using N to be the number of cathode starter sheets with A c being the area of both sides combined.0 V mol Cu mol e91. 49037 dnCu t Cu/day h d dt {27} A typical EW cell would contain 60 cathodes 1 m wide x 1-1.546 g x 10-6 t x 3600 dt dt s mol g s x 24 h = 5.x 96485 C mol mol e- {28} = 0. dMCu = dnCu mol x 63. edge strips may be used to prevent plating there. In practice. multiply dn M/dt by appropriate conversion factors. The area of the narrow sides is negligible and little copper plates there since the electric field is rather diffuse at the sides anyway.917 2 dnM = sec m sheet dt 2 mol e. The calculation above allows estimation of copper production based on current and current efficiency. the fact that metal is plated on two sides is factored in. each with length x width = 1 m x 1m. at 200 A/m 2 current density and 91. To get plating rate in mass per unit time. Either way is equivalent.6262 t Cu/day/cell x 365 days/y x S cells = 50. This aids in design of an actual EW tankhouse. The plating area on one side is 1 m2. Copper is plated on both sides of the cathodes. This makes removal of the plated metal sheets much easier.000 t/yr of copper production under the conditions we have been using in the calculations above? 0.6262 t Cu/day Note: the cathode sheet has dimensions of 1 m x 1 m. For tonnes per day: dM = dnM mol x AW g x 10-6 t x 3600 sec x 24 h dt dt sec mol g h d {26} For copper.7% CE is: 200 C x 2 m2 x 60 sheets x 0. This depends on the metal being plated since it involves the atomic weight. For copper the cathode production rate for a cell with 60 cathodes per cell.49037 = 0.area of just one side. Typical current densities range from 200-350 A/m2. How many cells of 60 cathodes each would be needed to achieve 50.000 t/y {29} S = 218.8 7 . we plate on both sides.11405 x 5.2 m deep and 61 anodes. The number of cells needed to achieve a desired production per year can be readily determined. The plating area of the whole sheet is 2 m 2.11405 mol/sec dMCu/dt = 0. Regardless. 7 x 63. the average weight of a cathode would be: 100 cm x 100 cm x 0.We would need 219 cells.385 x 105 sec 2 200 C x 1 m x 91.5 cm x 8. (We are considering the time to grow a single cathode copper sheet.7% CE: t= 100 x 2mol e.55 days Obviously the higher the current density.x 96485 C x 44. i.546 g/sec t 100 nF {31} t= {32} 100 nF MCu sec jAcN CE x 63. Now Ac is the area for a single sheet (on one side).6 kg (0. Summary of the Faraday's Law relationships.600 g mol mol e = 7.) For MCu = 44.92 g/cm3 = 44.600 g as above. The five relationships above are summarized in the table below. N now is 1.546 g sec m2 mol {33} = 8. dM/dt is constant as well. The total number of cathode sheets that must be employed then is 219 x 60 = 13. so the mass M plated over a specified time t equals dM/dt: MCu = jAcN CE x 63. If the current density was increased the number of cells required could be lower.92 g/cm3.5 cm? The metal production rate equation can be rearranged to obtain time. Table 1. 1 m 2 in this case. and given that the density of copper is 8.5 cm thick on average and 1 m long on each side.140. M = I t AW in g Faraday’s law nF CE = 100nFM Current efficiency It AW EE = E CE Energy efficiency Eappl -w’e = 100 nFEappl in J/mol Specific energy consumption CE Metal production rate dnM = jAcN CE in mol/s 8 .0446 t) {30} This means that 1.6 kg = 44. the shorter the plating time.546 where MCu in g. Since the current is constant. If a fully plated copper cathode is 0. plated at 200 A/m 2 with 91. How long would it take to plate a copper cathode to a thickness of 0.e.1211 x 106 sheets of copper have to be handled per year. or 3071 per day. Often cathode quality issues limit the current density. Real processes sometimes can approach. an infinitesimal change. the heat loss is larger than would be the case under reversible conditions. In the limit of infinite time a finite extent of reaction will have occurred. but never truly attain reversibility. Increase the concentration of A by d[A]. Continue to increase the concentration of A in infinitesimal steps. Stopping or reversing the process requires a finite change in a variable. if a process is operating reversibly. The reaction proceeds to the right to an infinitesimal degree. And suppose the system is at equilibrium.g. A reversible process always has associated with it the minimum possible heat flow. Note that at any stage during the process the reaction can be reversed by adding an infinitesimal concentration of B. (Irreversible does not mean that it cannot be reversed. i. Heat loss from the system is negative and q irrev < qrev. while the work that can be done is the maximum. If a process is exothermic. they occur infinitely slowly. If you very slowly and gently set down a large rock on a surface there will be little or no perceptible change in temperature of the rock and the surface.e. If you drop the rock it will hit the floor with substantial force and generate a substantial rise in temperature. A = B. "What is the minimum possible heat we can put into a process to make it go?" Or. they are a condition or case that thermodynamics can use to tell us something about theoretical limiting possibilities. "What is the maximum possible work we can get out of a process?" Naturally then. e. The latter was the most irreversible case. This is a reversible process. real processes are always irreversible.dt 100nF Thermodynamics of Electrochemical Cells Reversible and Irreversible Processes In thermodynamics a reversible process is one for which the direction of a process (such as a reaction) can be reversed by an infinitesimal change. rather the opposite of thermodynamically reversible. They move spontaneously towards equilibrium. But. They are also at equilibrium in other respects. How then can there be any actual change of state? Suppose there is a chemical reaction occurring. but. With respect to galvanic electrochemical cells (favourable reaction) under hypothetical reversible conditions. an infinitesimal change in pressure or temperature or concentration can reverse the direction of the process. the heat flow is the minimum. Truly reversible processes are of no practical use. Both cases involved the same change in gravitational potential energy. For example. qirrev is a bigger negative number than qrev (qirrev > qrev). Recall that the change in internal energy for a 9 . The reaction proceeds to produce additional concentration of B by d[B] increments. It helps us to answer questions like. for instance.) They have a finite (not infinitesimal) driving force to proceed in one direction. chemically or mechanically. Reversible processes are in thermal equilibrium with their surroundings.) Real processes involve conditions that are far from equilibrium. (This can be qualitatively understood from an example. 5. a refresher on some aspects of the thermodynamics related to electrochemical cells is needed.w. the more work that was extracted from that change.change of state (e. where P = pressure and V = volume. and in example 4 the battery goes dead. or potential for the process to occur. so the work obtainable is less.) Internal energy is a state function. and when this is done the real process is also irreversible. Extending the idea that reversible processes run infinitely slowly. the more irreversible the process.) The reversible case defines the limiting possibility. (U = q . Then. Water flowing downhill 3. (Recall that for a spontaneous process. The reverse (opposite) process is actually favoured and naturally does tend to occur. the faster the process is run. and the more of the energy that is lost as heat. There is a driving force. Flow of heat from a hot body to a cold one 2. In example 1. not how you get from one to the other. the more rapidly the battery is discharged. For a real cell. the process is necessarily irreversible and the heat flow is greater than in the reversible case.e. The Gibbs free energy function expresses the requirement that spontaneous processes must increase the net entropy of the system plus its surroundings. G < 0.e. Once the final equilibrium state is reached the capacity of the system to do further work is exhausted. the more the heat and the less the obtainable work. i. i.) If the process is not favourable it is not spontaneous and does not naturally tend to occur. it depends only on the final and initial states.g. Enthalpy is the sum of internal energy + PV. operated under real conditions. The farther from reversibility (or the more irreversible the process). Hydrometallurgical leaching reactions 4. work done by the system is positive by definition and heat exiting the system is negative. the less heat evolved for a given change. the two bodies reach the same temperature. The conversion of chemical energy into electrical energy in galvanic cells. If the change is thermodynamically favourable then the process is spontaneous and irreversible. H = U + PV {34} 10 . The non-spontaneous process can be forced to occur by input of sufficient energy. A Review of Some Relevant Thermodynamics Next. 1 mol A aq 1 mol B aq. Real processes involve finite changes of state with finite energy changes. Some examples of irreversible processes include: 1. An electrolysis. as above) is the sum of heat flow minus work flow. TdS {39} dG = dU + d(PV) – TdS {40} dG = dqrev – dwrev + d(PV) . i. dqrev = TdS {36} where T is the absolute temperature and dS is the entropy* change of the system (this from the definition of entropy).e. T).U = q – w. Then. dH = dq – dw + d(PV) {35} (H is very similar to U. At constant temperature.TdS dG = TdS – dwrev + d(PV) – TdS {41} {42} Generally. staying with the engineering convention that work done by the system is positive and heat flow out of the system is negative. dG = -PdV – dw’rev + PdV + VdP = -dw’rev {44} G = -w'rev {45} since dP = 0 (constant pressure). gravitational etc. but more convenient at constant pressure. The work term is. This is a limiting case. and this is obtainable only under reversible conditions. -dwrev = -PdV – dw’rev {43} where dw’rev is the non-PV work (electrical work here) under reversible conditions. reversible conditions: dG = dH . such as electrical. the work is comprised of pressure-volume work and non-PV work. However. (The cancelling of the PdV terms is what makes the enthalpy function convenient. real cells cannot operate under truly reversible conditions. At constant pressure.) Hence the maximum non-PV work (w’ rev) is equal to -G (at fixed P. Sometimes real cells may come moderately close. and the heat flow is the minimum possible.) For a reversible process the heat flow is denoted dqrev. Electrochemical cells commonly do operate under conditions of constant temperature and pressure. Under reversible conditions. and. 11 . the system can do its maximum possible work. (the former is of interest here). -dw = -dwrev = maximum work possible {37} G = H –TS by definition {38} where G is the Gibbs free energy. PV + PV = q = heat flow at constant pressure. q < 0 means heat flows out of the system into the surroundings (exothermic) q > 0 means heat flows into system from surroundings (endothermic) w < 0 means work flows into the system from surroundings w > 0 means work flows out of system into the surroundings Heat flow is considered from the perspective of the system. Thus the "concentration" of energy always tends to drop. pressure-volume work. Thus if G < 0 there is a net increase in entropy within the system + surroundings. If G > 0 the reaction is unfavourable (not spontaneous). This is expressed mathematically as. and the process is spontaneous.S. Then H = q . entropy naturally tends to increase. either means of energy flow out 12 . Either means of energy flow into the system is positive.TS accounts for the fact that heat flow out of the system into the surroundings is the opposite of heat flow in the surroundings to the system. U = q . unequal concentrations tend to equalize. energy wants to become more diffuse.* Entropy can be thought of as the inverse of the "concentration" or "quality" of energy. while work flow is considered from the perspective of the surroundings. and so on. If a process is spontaneous (or favourable) it means that overall there is a net increase in entropy. Energy relations for electrochemical cells Recall the first law of thermodynamics. which states “the energy of the universe is constant.e. It just happens.e.TS at constant temperature. if it were to occur there would be a net decrease in entropy. the dispersal of energy increases. Energy naturally tends to disperse: heat flows to cooler bodies.” In any system energy can be transferred to or from the surroundings as heat or work. it takes care of the sign convention issues. light moves away from its source. To provide a brief and less than rigorous rationale for this. Then H/T is the entropy change in the surroundings. where S is the entropy change for the system. it is the inviolable requirement for any process to be spontaneous. though it can be forced with energy input. THIS IS WHAT DRIVES ALL SPONTANEOUS PROCESSES.(TS) = H . consider that G = H . (In the SI convention both work and heat flows are considered from the perspective of the system. Regardless. i.w {46} The equation follows the engineering sign convention where. a degradation of the "concentration" of energy. All this occurs naturally without having to be forced.w and at constant pressure w = P V. These are forms of energy in transit.” or “energy is neither created nor destroyed. It may occur within the system of interest (a reaction in a cell. This is the entropy effect. Then G/T = H/T .) Under reversible conditions H/T = qrev/T = the heat flow into the surroundings over T. H/T is q/T when the pressure is constant (U = q . i. The minus sign in G = H . And. often denoted q P. both are flows of energy. Thermodynamically speaking. for instance) or it may occur in the environment surrounding the system (the "surroundings") or both. such as the expansion of a gas. This cannot naturally occur. What is so marvelous about the Gibbs free energy function is that it accounts for both the change in entropy in the system and in its surroundings. including work other than pressure-volume work (PV work. G2 – G1 = U2 – U1 + (P2V2 – P1V1) – (T2S2 – T1S1) {48} Since U2 – U1 = q . G = H . e.w {49} G2 – G1 = q . w = w’ + P(V2 – V1) {52} where w’ represents the electrochemical work.w + P(V2 – V1) – T(S2 – S1) {51} w is the total work. For a change of state achieved reversibly. H = q – w’ {56} which is equal to the heat flow at constant pressure. Since.g. at constant temperature and pressure. Substituting this into the preceding equation yields. Most chemistry texts follow the latter convention. we would have the equation.w + (P2V2 – P1V1) – (T2S2 – T1S1) {50} and at fixed P and T. G = U + PV – TS = H – TS {47} If a system undergoes a change from state 1 to state 2.TS {54} Again.TS {57} An electrochemical cell for which G < 0 (favourable or spontaneous) operated reversibly will generate the maximum possible amount of electrical work. G = qrev – w’rev .TS {55} It is then apparent that. expansion of a gas against some external pressure P). G2 – G1 = q . this applies to a process at fixed P and T. or G2 – G1 = q – w’ – T(S2 – S1) {53} G = q – w’ . For an electrochemical cell where electrical work is also possible (by means of electrons flowing through an external circuit). -w’rev = -w’max = G {58} 13 .) By definition.of the system is negative. When E > 0 the cell reaction is spontaneous. It represents the driving force for electron transfer. or.= Zn E° = -0. the current is infinitely small). It is the voltage of a cell under reversible conditions (no current is flowing. The H+/H2 half reaction is rapid on platinum. for which E > 0. (Rates of electron transfer depend strongly on the surface at which they occur. Work (in joules) = voltage (V) x charge (C). The other electrode is a piece of platinum. the minimum possible work that we can input to force a reaction to go in the unfavourable direction is again -w'rev (< 0). the actual work input will be > -w'rev. i. G = qrev + G .as per equation {45}. The four common types of cells are illustrated in the Figure 2 below. Hydrogen gas is bubbled over the platinum surface. as we would expect. qrev = TS (at constant temperature) {60} This is the flow of energy as heat in a cell operated reversibly and represents the minimum possible heat loss (again for a cell where G < 0 which does electrical work).76 V (5) There are two possibilities. Real cells. not favourable).TS {59} or. The favourable reaction may occur. Definitions The cell potential is also called the EMF (electromotive force). Alternatively the reaction can be forced to go in the opposite direction by applying a 14 . A piece of zinc is suspended in a solution of ZnSO 4 and H2SO4. in order to distinguish charge from heat flow. not the SI convention) and are related as follows. For an electrolytic cell (G > 0. To make the reaction go at a practical rate. the charge will be symbolized as q c. is not favoured.= H2 E° = 0 V (4) Zn+2 + 2e. The cell potential E and the reversible electrical work w’rev have the same sign (for the engineering convention. -G = nFE = w’rev = qcE {61} Types of Electrochemical Cells Now different types of electrochemical cells can be compared along with their energy relations.) The half reactions are: 2H+ + 2e. operated irreversibly will generate less work and more heat. making the process irreversible. When E < 0 the reaction as written is not spontaneous. (Recall that G = -nFE. Substituting this into the equation above.e.) In this discussion. Schematic depiction of four different types of electrochemical cells [1]. as illustrated in Figure 3 below. Figure 2. Because the reduction of H + on pure Zn is so slow. The favourable reaction involves oxidation of Zn to Zn+2 and reduction of H+ to H2. the cell can be set up with Zn metal in direct contact with H +. In fact there is no thermodynamic reason why the reaction should not spontaneously occur directly on the zinc surface. (a) (b) 15 . whereas it is quite rapid on Pt. the reduction of H + on very pure Zn is very slow. However. Otherwise two half cells with provision for ionic conduction would be used. This is electrolysis.suitably high opposing voltage. 1. Schematic illustration of hydrogen evolution on (a) a zinc surface (slow) and (b) catalyzed by Pt metal in contact with the Zn (fast).47 x 105 J/mol = -147 kJ/mol {62} H = q .= H2 E° = 0 V (7) The overall reaction is: Zn + 2H+ = Zn+2 + H2 E° = 0 . the potential difference is zero (w’ = 0 = Eqc. The reaction proceeds spontaneously since E° >0 (or E > 0 for non-standard conditions) and the process is favourable. So 16 . which is the role of Pt. (If you have ever short-circuited a battery by connecting a wire across both ends.76) = 0.76 V = -1. Since the reduction of H + on pure Zn is slow.) Since there is no work being extracted. Any redox reaction occurring directly between reagents without running the electron transfer through an external circuit is a short-circuited cell. you know this from experience.76 V (6) 2H+ + 2e. Therefore E = 0).76 V (8) This is really equivalent to the situation in Figure 3 (b). Zn = Zn+2 + 2e. and so by convention is negative. First consider the short-circuited cell.x 0. G° = -nFE° = -2 mol e-/mol x 96485 C/mol e.(-0. there is no load).47 x 105 VC/mol = -1. H = q {64} All the energy is dissipated as heat. the wire can get red hot and the rate of discharge of the cell can get dangerously fast.E° = -0.Figure 3. Cementation reactions and corrosion processes are examples of shortcircuited cells. but when short-circuited this is not realized. Note that the cell in principle is capable of manifesting a voltage E. The process is exothermic. all that is required for the reaction to proceed at a substantial rate is a catalyst. Heat flows from the system (the cell) to the surroundings.w’ and w’ = 0 {63} Since. (no electrical work is being done. There is no load (no electrical work is extracted) in the system. If the conditions were non-standard the potential difference would differ from E°. it is often not so easy for many reasons. Then the current flow is so small that the rate of reaction also is extremely slow. Hence the heat loss must be somewhat greater than the minimum reversible process heat loss. however. The current passes at a fairly low. Then the heat loss is greater and the work that can be extracted is lower. PH2 = 1 atm. but. (Alternatively. This indicates the thermodynamic potential difference. as per the Nernst equation. This very nearly approaches the reversible case (a process carried out infinitely slowly). no current flows. Other examples include oxidative leaching processes of sulfides and combustion reactions. the faster the process and the greater the extent of departure from reversibility. For all practical purposes.) Under standard conditions (298 K. the rate of the reactions is infinitely slow by virtue of the open circuit. Open-circuit cell. For practical work we use the cell galvanically or electrolytically. The half reactions are the same as in (1).76 V. An open circuit cell is essentially a cell working reversibly. The reaction is favourable. 3. In this case the electrodes are connected to a moderately high resistance device that uses electrical energy as work. The cell voltage equals the thermodynamic potential. open circuit cells can be used to measure thermodynamic potentials. unit activities of the ions) the measured potential difference is 0. At this point the opposing voltage is equal to the cell voltage. In practice. finite rate.is a cell where a wire is connected across the poles. the process is getting closer to operating like a short-circuited cell. The thermodynamic potential for the reaction is given by the Nernst equation: 17 . called potentiometers. However. The greater the current. the maximum possible voltage). A voltmeter can be used to measure the potential difference. we can't extract work from such a cell in a finite time. It is of use for measuring cell potentials. Galvanic cell (battery). It's of no practical use as far as obtaining electrical work. and this would indicate the driving force under those conditions. or driving force for the reaction. an opposing voltage can be applied until the current is zero. though it is irreversible. In principle. where all the energy is dissipated as heat. The half reactions are the same as in (1). These devices. 2. are not much in use anymore. The opposing voltage slows the reaction until a point where the reaction stops. Since the same charge is being passed as it would be under reversible conditions (2e . if the electrical connection between the two electrodes is broken. and the work obtained must be somewhat less than it would be under truly reversible conditions. One way to rationalize this is that as the current gets high.per Zn+2) and the available work is less. the cell voltage (denoted V) must be lower: (w’ = V qc) < (w’rev = E qc) {65} Therefore V < E (E is the cell voltage under reversible conditions. The cell can do its maximum possible work. such as a radio. This may employ a very large resistance inside the meter. This might be close to reversible conditions of operation. This is why fuel cells are naturally quite efficient.E = E° . Recall that H is a state function. A fuel cell is simply a galvanic cell in which the reactants are continuously replenished. no further change occurs. the left side of reaction 8 at a given temperature. The sign of q is negative to account for heat flow from a system being taken to be < 0. at which point G and E for the cell both go to zero. galvanically or as a short-circuited cell. the efficiency approaches that of a reversibly operated cell. As the reaction proceeds [Zn +2] increases.RTln aZn+2 PH2 nF aH+2 {66} Say PH2 is kept constant at 1 atm. At this point the battery is dead. Eventually equilibrium is reached and no further reaction occurs. if the cell is discharged slowly. the right side of reaction 8 with specified temperature. 18 . but the sum. pressure and concentrations) to a specified final state (e. Then E goes to zero. is always the same. In principle qrev could be positive as well. The process will continue until chemical equilibrium is reached.w'. pressure and concentrations) the change in enthalpy is the same no matter how the change is effected.g. What does depend on how we run the cell is q and w'. At equilibrium there is no more driving force for the reaction to proceed. and the products are continuously removed. The ln term thus increases as the reaction proceeds and E drops. E then is a measure of how far away from equilibrium the system is. while aH+ decreases. Figure 4.g. q . while work done by a system is > 0. how great the driving force is for chemical reaction to occur. a galvanic cell is one for which w'rev is positive. The thermodynamics can be conveniently represented on a diagram as shown below. meaning that for going from a specified initial state (e. However. energy both as heat and work leave the system. be it reversibly. The limiting case is the reversible cell. All real galvanic cells operate at less than the thermodynamic limit of efficiency. By definition. work is done by the system. Note that in this case heat flows from the system. Summary of thermodynamic effects for galvanic and short-circuited cells. In this case H < 0. relative to the limiting. Note that the net heat flow depends on how much work is extracted from the cell. Thus the additional heat loss arises from inefficiency in operating the cell. while work done by a system is > 0. qrev is the minimum possible heat flow and w rev is the maximum possible work. Another possibility for a galvanic cell is when qrev > 0. The net heat flow is q = qrev + q. q > qrev and w' < w'rev. The quantity q is the difference bewteen w'rev and w': q = qrev + q {67} G = -w'rev = -nFE {68} Based on equation {54}. reversible case. are irreversible in the thermodynamic sense and exhibit larger heat losses and lesser capabilities for work. Galvanic cells operate at finite rates. In addition. For a galvanic cell q < 0 always. The sign of qrev is positive to account for heat into a system being taken to be > 0. the faster the cell is operated (the greater the current) the more irreversible it is and the greater the heat flow and the less the work. 19 . q = G + w' + TS {69} q = G + w' + qrev {70} q = -w'rev + w' + qrev {71} q = -w'rev + w' {72} -q = w'rev . If the cell is short-circuited no work can be extracted and all the energy output is lost as heat. The heat loss is the sum of q rev + q (which is < 0).w' {73} Then. This is the other limiting case. This is depicted in the alternative diagram below. Alternative summary of thermodynamic effects for galvanic and shortcircuited cells. Figure 5. some potential work is lost as heat. In practice. Zn+2 + 2e. The barren electrolyte after EW is recycled to increase the metal ion tenor.= Zn (9) H2 = 2H+ + 2e- (10) The overall reaction is how the reverse of reaction (8): Zn+2 + H2 = Zn + 2H+ E° = -0. As reactants are depleted. To make the reaction proceed further still. 20 .02958log(PH2 aZn+2) 2F aH+2 aH+2 {75} even a 50% change in concentrations has only a small effect on the cell voltage. Eappl > 0. electrowinning is carried out under conditions of controlled current.76 V. The flow of electrons is reversed and so are the electrode reactions relative to the galvanic or short-circuited cases. electrowinning usually takes less than 50% of the desired metal ion from the solution. To overcome this.76 V (under standard conditions) is applied externally to force the reaction to go as written above. 2.e. The reaction will reach equilibrium when the thermodynamic cell voltage reaches Eappl. one would have to increase Eappl. Considering that E changes by. Hence the log term increases and consequently E decreases (becomes a larger negative number) as the reaction proceeds. a voltage of >0. In electrolysis the reaction is being forced in the opposite direction of its natural or spontaneous direction. Under conditions of fixed external potential. the applied voltage must increase to maintain the constant current. as was explained in the section on Faraday's Law relationships. This is what is employed in electrowinning. The thermodynamic potential for the reaction is given by the Nernst equation: E = E° . the reaction rate would decrease as the thermodynamic cell voltage decreases (because the driving force. In practice. The electrodes are connected to an external power supply such that the voltage exceeds and opposes the thermodynamic cell voltage. Then the thermodynamic cell potential is just balanced by the applied potential and there is no net potential difference between the electrodes. As the reaction proceeds aH+ increases and aZn+2 decreases. Electrolytic cell. which is the difference between Eappl and the thermodynamic E . Then it will stop. decreases). The reaction stops. i.303RTlog(PH2 aZn+2) = 0. The half reactions now are.RTln aH+2 nF aZn+2PH2 {74} Say PH2 is fixed at 1 atm.76 V (11) This is NOT favourable and will not occur naturally.4. The H 2/H+ reaction now becomes the anode and the Zn +2/Zn process becomes the cathode. This will continue until E and Eappl are equal. rather than controlled potential. an external voltage is applied. w’ = -nFEappl {76} where Eappl is a positive number. and the work done on the cell is given by.w’ {77} Note that w' < 0. w’ < w’ rev and -w' > -w'rev. H = q . Then w’ < 0. The energy changes are summarized in the diagrams below.Now work is being done on the cell by the surroundings. The work done on the cell is directly proportional to the applied voltage. The minimum work required to drive (a) (b) 21 . (a) q rev > 0. Summary of thermodynamic effects for an electrolytic cell. Referring again to the equation. Additional heat over and above qrev must be dissipated to the surroundings. Once Eappl gets large enough there is a net heat flow from the cell into the surroundings. In practice more electrical work is required to obtain reasonable rates. q = -nFE . (b) q rev < 0. which in turn depends on the magnitude of Eappl (Figure 6a). excess applied energy is lost as heat. the sign of q depends on the magnitude of w'. Some of that supplied energy does work on the cell. w’ = -nFEappl and w'rev = -nFE {80} then. Note that q is always < 0. This discussion assumed an isothermal system. Some of the supplied work energy results in an increase 22 . some of it ends up as heat. loss to surroundings and the heat capacity of the solution. Electrical work is supplied to the cell in order to overcome the cell's natural thermodynamic tendency. Practical electrowinning usually uses Eappl >> -E (the thermodynamic cell voltage) so that heat will be evolved. Depending on the metal being electrowon. as we would expect. -w' > -w'rev. If qrev < 0 the sign of q is always negative. but also heating of the electrolyte. The temperature of the solution will depend on the heat generated. There will be heat loss to the surroundings. Once -w' becomes large enough excess heat is dissipated to the surroundings. The net heat is the sum: q = qrev + q = qrev -w'rev + w' {78} If qrev > 0. excess heat may need to be deliberately withdrawn by heat exchangers.nFEappl + qrev {81} q = -nF(E + Eappl) + qrev {82} <0 >0 When E = Eappl q = qrev. the applied voltage then just matches E and the cell operates reversibly. the reaction against its favourable direction (backwards) to effect electrolysis is -w'rev. As Eappl exceeds E then excess heat begins to be evolved. q = G + w’ + TS {79} Substituting in. For an actual industrial cell the electrolyte temperature will rise and reach some steady state (although it may fluctuate with environmental conditions).Figure 6. as indicated in Figure 6b. Hence the cathode is positively polarized and the anode is negative. H-H bonds and Zn-O bonds in [Zn(H2O)6]+2. solvent extraction is the common method for solution purification. This is the net effect of breaking bonds (e. so that now the cathode is negatively polarized and the anode is positive.in chemical potential energy. I = C/sec moles e-/sec moles Cu+2/sec reacted. H3O+. Rates of Electron Transfer and the Effects of an Applied Voltage The rate of metal plating is directly proportional to the current through the cell. When Eappl > -E the reaction proceeds at a finite rate. concentrated electrolyte may be available for electrowinning. {83} In electrowinning we set the current and allow the voltage to adjust accordingly (as governed by V = IR). Eappl = -E. the higher the applied voltage must be.) Most commonly electrowinning is performed using sulfate solutions with oxygen evolution as the anodic half reaction: 23 .g. When the applied voltage precisely matches the thermodynamic cell voltage. This accords with the fact that the thermodynamic cell voltage is positive. Thus power supply (+) goes to the cell (+) and likewise the (-) of the power supply goes to (-) of the cell. Electrode polarity For a spontaneous reaction electrons flow from (-) to (+). finally. [Zn+2] decreases). SO42-) and dipoles (such as partial charge separation in H-O-H. but the flow of electrons is opposite. and the greater the difference the greater the rate. E = Ecathode .g. In other words. This forces the cell to run in the opposite direction. the oxygen being more electronegative and developing a negative charge. (In the case of copper. the electrodes retain the same polarity in either the galvanic or the electrolytic cases.Eanode > 0 {84} In an electrolysis the applied voltage opposes the thermodynamic voltage (E < 0) and is greater than -E. the cell operates reversibly and the reaction is infinitely slow. since. Electrostatic interactions involve charged ions (Zn +2. Zn-Zn metalmetal bonds and H-O bonds in H 3O+) and. Electrowinning Fundamentals After suitable solution purification a quite pure. etc. as is the direction of the chemical reaction. Thus the higher the current. the hydrogens having a positive charge). repelled from the negative electrode and attracted to the positive one. and forming new ones (e. changes in electrostatic interactions in the solution due to changes in composition ([H +] increases. the standard H +/H2 24 .g. i.= M (cathode) (12) H2O = 0. the higher the current. Then the measured voltage. The minimum energy required corresponds to the thermodynamic potential difference. The nature of the relationships between the applied voltage and logI is illustrated in the figure below.(anode) MSO4 aq + H2O l = M s + 0. corresponds to E. Then Cl .M+2 + 2e.89 V E° = 1. However. The rate of metal plating is directly related to the current.34 .23 V = -0. Electrowinning may also be carried out from chloride solutions in some instances.= Cu E° = 0. As the current approaches zero (logI -) the cell runs increasingly slowly.5O2 g (15) The reduction reaction is: Cu+2 + 2e.5O2 g + H2SO4 aq (13) (overall reaction) (14) In such cases acid is generated during electrowinning. we cannot run a half reaction in isolation. the half reaction potential.. the graph focuses on the log(current)-voltage graph for a single half reaction of interest *. and the greater the metal plating rate. so energy must be supplied to make it go.5O2 + 2e. in principle. 0.23 V (17) {85} The reaction is not thermodynamically favourable. Of course another half reaction has to be at play as well.is oxidized to Cl2.+ 2H+ = H2O Then E° = 0. CuSO4 aq + H2O Cu s + H2SO4 aq + 0. The electrowinning reaction for copper is. In practice higher voltages are used to attain practical rates.e. This depicts the current-voltage relationship for a single half reaction. approaching reversible behaviour. relative to some other half reaction (e.34 V (16) The oxidation half reaction is the reverse of the oxygen reduction half reaction.1.5O2 + 2H+ + 2e. the more e-/sec transferred. It is given the symbol . the number of electrons/mol of metal plated. Consider the cathodic branch. (The form of the curves is well understood from electrochemical theory. The half reaction of interest occurs at a "working" electrode. This is called the Tafel region. half cell.g. Schematic illustration of a polarization curve for a half reaction plotted as voltage vs. are needed. Ni+2. As indicated earlier. Eventually a nearly linear region occurs. As the potential is decreased M n+ is reduced. Mn+ + ne. but that is beyond the scope of this introduction. the shapes of the curves are often quite similar (there are subtle differences). For the upper branch the oxidation occurs. Co+2 or Zn+2).) At this point a minute change in potential can reverse the direction of the half reaction (again.= M s (19) It all depends on how the electrode is polarized (how the external potential is applied). at a given current Ni+2/Ni is greater than Cu+2/Cu. At 25 . M s = Mn+ + ne- (18) For the lower branch the reduction half reaction occurs.g. Its slope is directly proportional to n. This is illustrated in the figure below. logI. the value of n will have a direct effect on how big the overvoltage is. higher voltages. The current is measured between the working electrode and a "counter" electrode. what really matters is where the curves lie horizontally. The reduction of Cu+2 on Cu is considerably faster than the reduction of Ni +2 onto Ni. This appears as a shift of the Ni +2/Ni curve further to the left (to lower currents) relative to Cu+2/Cu.) The curve indicates that to obtain higher currents. consistent with reversible behaviour). Overpotential (or overvoltage) is the additional potential needed beyond the thermodynamic potential E required to make the half reaction go at the desired rate. For a given n value (2 for EW involving Cu+2.Figure 7. beyond the reversible value E. As a result. where logI is virtually linear with applied potential. The overpotential This leads to the idea of the overpotential. e. The potential is measured relative to a reference electrode. * A "three-electrode" cell can be used for this. However. e. This has implications for the energy consumption in electrowinning of the two metals. At first a substantial increase in reduction current results from relatively small decreases in applied voltage. to CNO-. Electron transfer to the metal cation. or energy hurdle that must be overcome. they all have an activation barrier. Migration of the metal atom on the surface to a suitable crystallographic site. Chloride oxidation to Cl 2 is also employed in Ni EW. and common for Ni and Co. This is ubiquitous in Cu. Schematic illustration of reduction branches of polarization curves for Cu+2/Cu and Ni+2/Ni couples. H2O is the reactant. Each of these processes contributes to the overpotential. Loss of water molecules may occur in concert with electron transfer. The anode half reaction also has an associated overpotential. 26 .a given current a greater overvoltage necessarily implies greater energy consumption. Figure 8. For electroplating simple metal aquoions three broad classes of processes must occur: Deformation of the aquocation complex as it approaches the metal surface and loss of coordinated water molecules. the anode half reaction is oxidation of CN. and Zn EW. Oxygen evolution has some advantages: No additional costly reagents are needed. For a given current the Ni +2/Ni couple requires greater overpotential than does the Cu+2/Cu couple. Oxygen evolution is the most common anode half reaction in electrowinning. In gold EW from cyanide solution. Sometimes one of these steps can be a major contributor to the overpotential. 89 V Ni+2 + H2O = Ni + 2H+ + 1/2O2 (20) E° = -1.5% Sn. This contributes to highly negative thermodynamic cell voltages: Cu+2 + H2O = Cu + 2H+ + 1/2O2 E° = -0. Relatively less corrosive sulfate medium is suitable. DSA anodes are of interest because the platinum group metal 27 . 0.) This results in a high anodic overpotential. the main disadvantage is that it contributes to high energy consumption. The intent is to shift the polarization curve further to the right. (This is an area of considerable economic and technical importance for fuel cell development too. There may be other effects as well. which are inexpensive. First the E° for the O2/H2O half reaction is 1. Lead has one of the highest overpotentials for oxygen evolution! However. Calcium is added to increase strength. Tin lowers the oxygen evolution overpotential and improves corrosion resistance. lead is commonly used due to its low cost. Much work has gone into trying to find ways to lower the oxygen evolution overpotential as it is a significant factor in the operating costs of an EW plant. The anodes are cold rolled for the same reason.1% Ca). Lead anodes are commonly alloyed with minor amounts of other elements to try to lower the overvoltage.23 V. water oxidation on most electrode surfaces is very slow.46 V Zn+2 + H2O = Zn + 2H+ + 1/2O2 E° = -1.99 V (21) (22) In addition. This is illustrated in the figure below. However. Lead anodes may be used. The most common anode material in use today in Cu EW is a PbSn-Ca alloy (~1. and sulfate medium is the least expensive. At 300 A/m 2 current density the approximate O2 evolution overpotentials under different conditions are provided in the Table below. There are numerous effects of solutes and alloy elements on anode performance and longevity. oxides used to coat the anodes (usually a titanium substrate) are very good at promoting O2 evolution. without which there would be no current. This is necessary since anodes and cathodes must be removed (anodes in order to be replaced. 6% Sb) ~0.5%) 1 ~0. cobalt is expensive. However. The other substantial resistance in the circuit is the solution resistance.5O2 (anode) (23) (24) Co+3 is a powerful oxidant and rapidly oxidizes water. (This was one of the key observations that lead scientists to conclude that some compounds were comprised of discrete cations and anions.9 V . This is lost as heat. cathodes for Cu metal harvesting). beyond which it has little beneficial effect. Likewise.7 +2 Pb/Sb (6% Sb) 100-150 mg/L Co ~0. However.very high! 2Co+3 + H2O = 2Co+2 + 2H+ + 0.6 2 DSA ~0. Oxygen evolution overpotentials from sulfuric acid solutions under several conditions at 300 A/m2 current density [2]. Various amendments may be applied to shift the polarization curve to lower . It also lessens anode corrosion and improves stability of the PbO2 layer.6 Pb/Ca(0. Cations move towards the cathode (negatively polarized) and anions move toward the anode (positively polarized). these extend over the length of the tankhouse and overall the resistance is enough to cause a moderate voltage drop. Thus electrolyte conductivity is an important technical consideration in EW. Anode material Other conditions Overpotential (V) Pb/Sb (e. Schematic illustration of oxygen evolution polarization curves. Table 2. This allows for a complete circuit.Figure 9.35 1 Preferred due to better mechanical and corrosion properties compared with Pb/Sb. contact between anodes and the current distribution conductors (busbars) and between cathode sheets and the busbars also has a resistance loss (contact resistance). Resistance losses Conductors that carry electricity to the electrodes have an intrinsically low resistance. Cobalt addition is commonly practiced in Cu EW. 2 DSA = dimensionally stable anodes. It is added at 100-200 mg/L.) A strong electrolyte is a salt that is soluble in water 28 .1%)/Sn(1. rigid titanium sheet or mesh coated with platinum-group metal oxides. It is believed that Co +2 oxidizes at the anode and catalyzes H2O oxidation [3]: Co+2 = Co+3 + e- E° = 1. Ions in solution are capable of conducting electricity. In solution the current is carried by migration of ions.g. H + is present as H3O+. Copper EW electrolytes contain on the * The reason is believed to be due to the "proton jump" mechanism. The greater the cross sectional area of the resistor. the more current it can support (more pathways for electrons/ions to move through). Much beyond this and corrosion of the steel starter sheets and other parts of the plant becomes problematic. and where feasible. H +.as well.and which fully dissociates into ions. so based on Ohm's law the voltage drops due to resistance losses can be summed up: VIR = IRi {86} i The resistance of a resistor is a function of its geometry and its inherent resistivity (the resistance of a standard geometry). e. Thus the presence of H 2SO4 in the electrolyte is quite beneficial. Thus a more concentrated solution of ions.g. recall that SO42. then other cations. The longer the distance through which the electrons (or ions) must travel. is by far the best conductor. The hydrogen ion. Hence R 1/A where A is the cross sectional area. Thus in EW large surface area electrodes are used with a minimum practical spacing between cathodes and anodes. R L/A {87} R = L/A {88} and where = resistivity. This allows it to move through solution much more rapidly than the diffusion or migration of other ions. Clearly the lower the resistivity the better. Ions vary in their ability to conduct electricity (per unit concentration). This is illustrated in the figure below. The same mechanism may be in effect for OH -. high concentrations of H+ are desirable. This essentially the formation of a complex. though it is somewhat less effective. Some example 29 . Limitations on these features are discussed later. In electrowinning of divalent metal ions from aqueous solution there is a complicating factor and that is ion pairing.is also a weak base and may form HSO4. followed by OH-. the greater the total resistance. EW is operated at constant current. H+ is at least 3 times more conducting (per unit concentration) than other cations*. temperature and pH. M+2aq + SO42-aq = MSO4 aq (25) The extent to which this happens depends on the metal ion. A neutral ion pair does not contribute significantly to the conductivity of the solution. An H+ in an H3O+ ion can "jump" to a neighbouring water molecule as shown below. Then. as indicated by its lower conductance. order of 180 g/L H 2SO4. R L where L = length. Table 3. data for simple CuSO4-H2SO4 electrolytes are shown in Table 3. A typical electrolyte might contain 40 g/L Cu+2.6 1. The inverse of resistance is conductance.50 40 182.Figure 10.3 2. The resistance then is: 30 .2 1. This also will affect the resistivity of the electrolyte. 180 g/L sulfuric acid and have an operating temperature of 50°C (although there may be a significant range of conditions in practice).6 1.4 2.26 35 143.4 1.42 50 163.76 55 182.e. Now higher conductance (lower resistance) is desirable.42 cm.57 55 163.71 50 182.15 As an example. while higher acid concentrations lower the resistivity. i.4 1. The specific conductance is 1/ in -1cm-1.29 2. Illustrative example data for resistivity of simple CuSO 4-H2SO4 solutions at different temperatures [4]. Note that for a given acid concentration (50°C column) the resistivity increases with increasing [Cu+2]. The corresponding resistivity then is about 1. Resistivity cm Cu+2 g/L H2SO4 g/L 30°C 40°C 50°C 25 71.92 1. Illustration of the dependence of resistance on conductor length and cross sectional area.7 3.37 2.6 1. [CuSO4] (as well as decreased [H+] due to formation of HSO 4-). take the cathode-anode gap to be 4.74 40 163. Iron is also commonly present in copper EW electrolytes at concentrations of up to 5 g/L.2 m 2 (100 cm x 120 cm). and the immersed plating area of 1.4 1.09 35 191. It is also common to consider the inverse of resistance or resistivity instead.63 60 167.52 3.75 3. It will be present in its two common valence states: Fe+3 and Fe+2.8 cm. The last term is the sum of the IR voltage drops. This is what determines the energy consumption and associated cost.2 m2 plating area x 300 A/m2 = 0. solution resistance ~0. a total voltage drop dues to resistance losses of up to 0. depending on whether the hardware and rectifier resistances drops are included or not.8 cm / 12000 cm2 = 0. For a 300 A/m 2 current density. Depending on the source cited.1 V (depends on current density) IRi ~ 0. at higher currents the 31 . Limiting and practical current densities From Faraday's law and related relationships. Ballpark estimates for the four terms in copper EW are given below: -E ~ 0. for a given plating surface area production increases with increasing current density. -E is constant for a given electrolyte composition and changes only a little through the cell as metal is plated.2 V (The rectifier converts AC power to DC power.42 cm x 4.000568 x 1. In general industrial plants employ cell voltages of 1.) Total applied voltage ~2V in the example above. However.6 V (depends on current). In the end. what matters is the total applied voltage and the total current. Note too that increasing the current increases this term. it is clear that plating at higher current densities is desirable. there are two important limitations.5 V (depends on current density) Cathodic overpotential (C) ~ 0.5-0. Likewise the overpotentials are a function of current (or current density). there is some variability in values quoted for solution resistance and cell hardware resistances.000568 {89} The voltage drop across the resistance is V = IR.3 V and i rectifier + cell hardware resistances ~0.204 V {90} The four contributions to applied cell voltage Putting the pieces together we have the following relationship for the applied voltage: Eapplied = -E + C + A + IRi {91} i The first term is the thermodynamic potential (E is negative and must be opposed).9 V Anodic overpotential (A) ~ 0. Next are the two overpotentials.1. 0. However. The first is a fundamental limit. Note that this significantly exceeds -E.6 V or so is typical.2 V.8-2. Referring back to Figure 7. The actual cell voltage (measured between anode and cathode) and the total applied voltage (including the power source and current distribution network) may differ somewhat. which induces mass transport of Mn+ towards the cathode. Distance is measured from the cathode surface. This creates a concentration gradient. Schematic illustration of concentration profiles for a metal ion being plated at an electrode surface. impurities content in the cathodes and make them offspec. 32 . The greater the plating rate. making cathode harvesting difficult and costly. the lower the concentration at the surface. In practice significantly lower current densities are employed for conventional EW. At lower. the deposit may be more adherent. Figure 11. the bigger the concentration drop. This is related to a practical limitation.). etc. As metal ion is plated. relative to the bulk concentration. At some critical current density the concentration of metal ion at the surface goes to zero. oxygen (from water. This makes for very crumbly deposits that adhere very weakly to the cathode surface. This is the diffusion limited current and represents the maximum current that can be sustained under the given conditions. iron. This is illustrated in the figure below. but still too high plating rates too many crystals are still forming. Thus in conventional copper EW maximum current densities are about 350 A/m2 [6]. This wastes electricity and would never be attempted in a normal electrowinning situation. Reduction of H+ to H2 may then occur.23 V.polarization curves start to approach infinite slope (tend toward vertical lines). This will increase the sulfur (from electrolyte sulfate). etc. for which E° = -1. filtered and washed. the concentration drops near the surface. i. rather than growing existing crystal faces. For copper EW the diffusion limited current density is about 500 A/m 2 [5]. An increase in the applied voltage then would also necessarily occur (at constant current). The dashed line represents the boundary layer thickness. The net result is that large numbers of new crystals form on the surface. Fine grained copper particles would spall off the cathode and have to be collected. At very high plating rates the copper atoms deposited on the cathode surface do not have time to migrate to a suitable crystallographic site. Then the next available half reaction will begin to occur. but it will be porous as crystals grow together rapidly and trap solution between their faces. Further increases in current cannot be sustained by Mn+ reduction. At this point the polarization curve approaches infinite slope.e. However this requires energy. copper depleted electrolyte near the cathode and rising oxygen gas bubbles near the anode.01 mm. At the anode surface oxygen gas is evolved. This in turn promotes uniform plating rate over the surface. This promotes a uniform distribution of the electric field over the surfaces (other than at the edges).01 mm. If the boundary layer is made thinner by agitation then the concentration gradient steepens and higher limiting (and practical) current densities are possible. which is important for growing a smooth. However. though not to the limit of ~0. This is called the boundary layer. Directing electrolyte up between electrodes using a header further thins the boundary layer. There are two features in an EW cell that result in a measure of natural agitation (actually convection) [5]. lowering the [Cu +2] and decreasing the solution density. compact deposit. Figure 12. if electrolyte is directed up between cathodes and anodes using a header. which also helps to thin the boundary layer near the cathode surface. This causes solution to naturally rise near the surface of the cathode.The 500 A/m2 limiting current density is for a solution without intentional agitation. This lowers porosity and helps prevent occlusion of electrolyte with the attendant increase in impurities in the 33 . as shown in the figure below. Agitation then becomes difficult.2 mm [7]. closely spaced together. At the electrode surface Cu +2 is depleted. which requires intensive agitation [7] and is not achieved in EW. This gas pushes up solution as it rises. The net result is two counter-rotating loops. Overall this displaces solution in an upward motion in the vicinity of the anode. The bubbles accumulate at the surface before rupture. The boundary layer thickness achieved by this natural convection is about 0. A typical Cu EW tankhouse may have several hundred cells. Illustration of natural convection in an EW cell based on the rising of lower density. each with as many as 60 cathodes and 61 anodes. The vertical line in Figure 11 demarcates the point closest to the electrode surface where the solution metal ion concentration is equal to the bulk solution concentration. Current distribution and protrusions Anodes and cathodes are large plane sheets that are kept parallel. The minimum thickness of the boundary layer is about 0. this can impart some additional agitation allowing current densities to be at the higher end of the practical range.1-0. Schematic illustration of the effects of protrusions and depressions on electric field distribution. This in part acts to limit how close cathodes and anodes can be to each other. An operator will then go to the cell and use a bar to dislodge the dendrites to resume plating. This can be done by using "leveling" agents. However. Then there are high and low spots on the surface. The copper electrode now becomes the anode for a short time. making electroplating easier. Finally. further increasing surface roughness and promoting growth of protrusions. deposits are polycrystalline and surface roughness will eventually develop. As these extend further out from the surface their growth rate increases. the high points are closer to the anode and the depressions are further away. The figure at right is a schematic illustration of a short-circuit caused by a dendrite. The current follows the path of least resistance through the short circuit. At this point all the electrical energy in the cell is lost as heat. Due to differing distances IR" > IR > IR'. eventually making contact and causing a shortcircuit. This increases the resistance of the protrusion and helps slow down their growth [8]. Plant operators take pains to minimize this problem. These are often complex chemical mixtures that act by selectively adsorbing on the fastest growing sites. Figure 13. This makes the tips grow faster and the depressions grow slower. Electric field lines tend to converge at points and diverge at depressions. rather than through the solution to plate copper. Further. Then a polycrystalline agglomeration of copper called a dendrite will extend out towards the anode. The 34 . Another approach to controlling surface roughness is periodic current reversal (PCR). decreasing the IR voltage drop due to solution resistance for the high points and increasing it for the depressions.deposit. This leads to the situation shown in the diagram below. The polarity of the cathodes and anodes is switch briefly at frequent intervals. high points extend out into the boundary layer where the [M n+] is higher. All totaled these effects can result in more rapidly growing protrusions. It is common practice to use infrared scanners to detect hot spots in cells where a short circuit has occurred. to which copper adheres very weakly) and placed into the main EW cells. Sometimes copper starter sheets are used instead. with some variation.993% A major source of sulfur impurity is sulfate from occluded electrolyte. briefly. usually ~3 mm thick and 1 m X 1-1. however. This may obviate the need for addition of leveling agents and associated costs. A typical Cu EW cell will have up to 60 cathodes and 61 anodes. Long conductive bars called busbars conduct electricity to the electrodes. Cell dimensions are on the order of 6 m long x 1. Previously the electrolytic cells were made of concrete and had to be lined to prevent corrosion. Cells and hardware A diagram depicting a simplified cell is shown below in Figure 14. making the most of the available plating surface area. A cell will have n cathodes and n +1 anodes. The starter sheets are removed from a cathode substrate (such as titanium. dissolving them. Copper metal will plate on the lead electrode. as indicated above. are most rapidly oxidized. Then metal may be plated onto both sides of every cathode. A schematic illustration of a cell is shown in Figure 15 below. Stainless steel is used to minimize corrosion. Cathode copper is often plated onto stainless steel blanks or starter sheets.protrusions being the most active sites. This effectively prevents growth of protrusions on the cathode surface.4 m deep by 1. Copper Electrowinning Practice Specifications LME (London specifications: Metals Exchange) grade A copper has the following Pb < 5 ppm S < 15 ppm Other impurities <49 ppm Cu >99. it does cost electricity. but this will again be completely oxidized off.2 m surface dimensions [6]. PVC was a common liner. Cathodes and anodes are connected in a parallel arrangement. Cathodes are interleaved between pairs of anodes. These are plated separately and are 0.25 m wide [9].5-1 mm thick. Modern cells tend to be made 35 . The duration of the polarity switch is brief. 36 .Figure 14. Schematic illustration of a copper electrowinning cell electrode arrangement. A typical tankhouse will have hundreds of cells. Figure 15. Cells are electrically insulated from ground to minimize stray currents that lower current efficiency. Excessive build-up can eventually compromise cathode quality [10]. Schematic illustration of a copper EW cell. of polymer concrete (aggregate in a plastic binder). Regular cleaning is required to remove lead sludge due to slow anode corrosion (mainly as PbO2). In order to keep the cathode quality similar throughout the cell the total drop in electrolyte copper concentration within each cell is typically quite small (2-5 g/L). Anodes are typically made of a lead alloy as mentioned earlier. Addition rates corresponding to 150-300 g/tonne of copper produced are typical.3 cm at the start of cathode growth and 3. another beneficial effect is that it seems to promote the adherence of PbO2 to the anode surface [8. A cold-rolled Pb alloy anode can last 5-10 years [13]. The lower the rich electrolyte copper concentration.6 cm thick anodes.8 cm at the end. Anode consumption is 0. Guar gum is a polysaccharide. 37 . Steel starter sheets. Low rich electrolyte Cu+2 concentrations and low concentration changes between electrolyte entering and exiting the cells may necessitate smaller numbers of cathodes per cell (<60). However. Again. this coats on the fast growing protrusions and slows their growth. up to 150 mg/L Co +2 and 30-40 g/L Cl-. although the DSA's are starting to appear.013-0. Rich electrolytes can have as little as 25 g/L Cu +2 [12].5-11. PbSO4 is also quite insoluble. a PbO 2 particle lodges on the cathode surface. a polymer of mannose and glucose. being more rigid than thin copper starter sheets. Due to the formation of the PbO 2 layer the corrosion of lead is very slow.4 cm [11].3 cm thick steel cathode sheets and 0. Cathodes are grown to a thickness of about 0. Then the cathodes at the front of the cell experience nearly the same composition as those near the exit. Lead contamination of cathode copper can be kept to <1 ppm [5]. Most plants have rich electrolytes with 30-38 g/L Cu +2 [6] and concentrations of up to 50 g/L are possible.6 cm thick [6] and about 3 cm shorter on each side to promote uniform current distribution on the cathodes [12]. A PbO2 layer forms on the anode surface.The cathode-anode gap needs to be as small as possible to minimize the IR loss due to solution resistance. have allowed for the narrowing of the cathode-anode gap. With tin alloy electrodes a SnO2 layer also forms.8-4. Hence lower rich electrolyte copper tenors may limit a plant to lower current densities [12]. Gradually small particles of PbO2 spall off the anodes. As mentioned previously Co+2 is added to lower the O2 evolution overpotential. Typically anodes are ~0. if say.13]. Their disadvantage is a very high overvoltage for O2 evolution. which lowers resistance losses. which is also conductive. A common leveling agent in Cu EW is a derivative of guar gum [8]. Lead anodes are chosen for their longevity and insolubility. the anode-cathode gap is about 4. and cathodes that grow to 0. up to 3 g/L Fe+2/+3. With 0. the less tolerance there is for change in copper concentration.17 anode/tones copper plated (~1-12 kg/tonne Cu) [11]. The cathode-cathode centre-line spacing is commonly 9.5 cm and due to misshapen electrodes and cathode protrusions there is a limit on how small the gap can be. This can lead to some lead contamination of the cathodes.35. Electrolyte Keeping the Cu+2 concentration in the solution high helps with sustaining a high current density. Other compounds in the electrolyte include 140-150 g/L H 2SO4.5 cm thickness. The largest plants produce about 100. Thus Fe +3/+2 cycles back and forth and consumes current. Figure 16 shows photographs of a modern EW tankhouse.) Copper plated onto copper starter sheets is plated to about 100 kg of mass. Typical energy consumptions are 1900-2000 kWh/t Cu [6.) The rest remain in the cell so that it can keep operating [9] (called live stripping).+ 8H+ + 5e- (26) Permanganate is a powerful oxidant that can oxidize organic extractant and solvent. Manganese may enter the electrolyte from solvent extraction by entrainment of aqueous solution in the loaded organic if the PLS is high in manganese. This can cause various problems in EW. In either process solid MnO 2 may form: MnO4.+ 3Fe+2 + 4H+ = MnO2 + 3Fe+3 + 2H2O (27) Manganese dioxide is itself a strong oxidant. In part this arises from chemical absorption of Fe+3 by the hydroxyoxime extractants. otherwise it would be too difficult to remove. (It is important that the copper not adhere too strongly to the cathode substrate. It helps to promote a smooth copper deposit [14].9]. iron also lowers current efficiency. This is a major source of current losses in copper EW. Excess chloride is harmful as it promotes corrosion of the steel cathode blanks. and in part it comes from entrainment of aqueous phase in the loaded organic. Current efficiency is typically 85% . Iron up to a point has some beneficial effects.95% . Ferric is reduced to ferrous at the cathode and ferrous is oxidized back to ferric at the anode. (Hence the number of cathodes per cell must be evenly divisible by three or two. lowering copper plating current efficiency. Chloride at 30-40 mg/L promotes smooth copper deposits [15]. An average annual production rate is 50. In addition the fine particles may foul cathodes and cause crud formation in the stripping operation of SX [16]. Up to about 3 g/L iron appears to provide a good balance between promoting smooth Cu deposits and unduly lowering current efficiency [8]. Much beyond this thickness and the mass becomes too great and the cathode may dislodge from the steel sheet under its own weight.Iron enters the electrolyte from solvent extraction. It can also oxidize ferrous to ferric.000 tonnes per year. Mn+2 can be oxidized at the anode to form MnO4-: Mn+2 + 4H2O = MnO4.000 t/y. Either 1/3 or 1/2 of the cathodes are harvested from a cell at any given time. 38 . Plant operations Copper is usually plated to a thickness of about 5 mm (40-60 kg each) over the course of 5-10 days. Cathodes are mechanically stripped from the steel sheets. In a good plant the ratio of Cu:Fe entering the electrolyte is 1000:1 on a mass basis. However. This arises in part from the electrical resistance heating directly from the EW process. Oxygen evolution at the anodes forms an acid mist when the gas breaks the surface. It may be necessary to cool the lean electrolyte returning to stripping to minimize organic degradation reactions [12]. Control measures include polyethylene balls on top of the surface to promote coalescence.10]. It was found that introducing the electrolyte via a distribution manifold led to consistently better cathode quality with the added 39 .Cells are operated at 45-50°C solution temperature [8]. The electrolyte heated by the waste heat from electrolysis (solution resistance) is used to preheat incoming electrolyte [10]. surfactants that produce a foam layer and prevent mist escaping and very good cross ventilation. In many older plants electrolyte was introduced at one end of a cell and exited from the other without any deliberate agitation. It also is a small loss of copper values. Sometimes poor quality (rough) cathode deposits were formed. dense deposits (fine-grained) [6. Good control of temperature is important for obtaining smooth. This is corrosive and toxic and much effort goes into controlling it. (a) (b) Figure 16 [17]. Notice that the tankhouse accommodates two parallel rows of cells. Openings along the right wall are for fans for ventilation. 40 . Each parallel bank of electrodes is a cell. (b) An illustration of “live stripping” where a fraction of the cathodes are removed and electrowinning continues in the cell. This example uses steel starter sheets. (a) A copper electrowinning tankhouse. Bleeding for impurity control Inevitably impurities will build up in the EW electrolyte [5]. NO3-. Uneven current distribution at the edges produces dendritic growths and this copper is not readily recovered.8]. making cathode stripping difficult. Roughly half of the world’s EW plants in 1999 had electrolyte inlet manifolds [13]. Several of the deleterious effects of impurity elements have been mentioned previously. A small fraction of the lean electrolyte is removed from the circuit (commonly 1-3% [8]). This leads to accumulation of iron especially. Electrodes are not shown. This occurs because the flow of electrolyte between EW and SX stripping is a closed loop. The maximum is about 0. Plating on the edges would bind the cathodes on each face together.8]. whether by entrainment of aqueous solution in the loaded organic or chemical extraction. Removing impurities is commonly done by bleeding. P. Illustration of electrolyte circulation manifold developed at San Manuel copper plant [5. If the electrolyte flow into the cells from SX is too low to sustain this then additional recirculation pumps may be needed. The manifold directs electrolyte flow upwards between the electrodes from all across the bottom of the cell. Strips or moldings or wax are placed on the edges of the cathodes to prevent copper plating on the edges [5]. This is illustrated in Figure 17 below. means plant electrolyte. A typical good flow rate is 0. and depending on the ore being leached. also Cl-.E. Spacers are also used to keep anodes and cathodes from coming into physical contact. Spacers on various sorts keep a sufficient gap between anodes and cathodes so that exfoliating PbO2 does not as readily come in contact with the cathodes [9]. These may be made of PVC and attached to anodes at the corners and in the middle.14 m 3/h/m2 [6] (too high and PbO2 particles can stay suspended). Mn+2 and others.advantage of allowing higher current densities (around 320 A/m 2 at the plant where this was initiated) [5].12 m 3/h/m2 of plating area (number of cathode faces x area per face) [5. 41 . Figure 17. Another means of preventing contact is anode buttons [12]. it is important to keep the bleed stream flow to a minimum.15 t Fe/day x 1000 kg/t) / 2 kg Fe / m3 = 75 m3/day = 3.55% of the electrolyte. Obviously.e. The flows to and from this tank are shown in the figure below. High copper and somewhat higher acid than in the PLS can displace iron and with it Cl . Plant configuration and electrolyte storage Rich electrolyte from SX goes into an electrolyte storage tank. Some of the bleed (diluted) can be used for this purpose.000 t/yr copper plant with Cu from SX = 10 g/L (difference between rich electrolyte and lean electrolyte [Cu +2]). The great advantage is that no cobalt is lost from the electrolyte. If the iron concentration that can be tolerated in the electrolyte is 2 kg/m 3 the bleed rate can be calculated: (0. This is best achieved by having a very well operated SX plant to keep the Cu:Fe ratio high. then the resin can be reused. the concentration of iron in the electrolyte is 1/1000 that of the copper.750 t/year) the iron entering the electrolyte is 0. which is very good selectivity. Complete removal from the electrolyte by bleeding is not feasible. Once the resin capacity is saturated the resin is stripped using a Cu +2/Cu+/H2SO4 solution. in a good plant the Cu:Fe ratio entering the electrolyte from SX is 1000:1. The acid concentration is quite high. The resin selectively removes Fe+3. 42 .or NO3-. The higher the iron content that enters EW. Then another stage is added in SX prior to stripping. consider a 50. This is called scrubbing [18]. An ion exchange process for iron removal (ferric only) from EW electrolytes is also available [19]. the loaded organic may have to be treated to lower the impurities concentrations [8]. Where iron. The bleed solution may be treated in a number of ways. Obviously this should be kept to a minimum. If we take as an example 150 t/day copper production (54. It is common practice to split the lean electrolyte flow between storage and return to SX [6]. Then too a little iron has beneficial effects on cathode quality.As mentioned. again. nitrate or chloride are at high levels in the PLS. The electrolyte flow rate then through EW is 571 m 3/hr. Since cobalt is quite expensive this adds a significant cost.125 m3/hr {92} For instance. then the whole electrolyte would have to be treated. the higher the bleed rate must be. This is necessary because the change in copper concentration between rich and lean electrolytes may be greater than the change in copper concentration between the inlet and outlet of the cells. The bleeding process necessarily removes cobalt from the electrolyte. The bleed rate then represents only about 0.15 t/day. It can be returned to leaching or to solvent extraction loading. Note that electrolyte makes one pass through the cells and then is returned to stripping as lean electrolyte. i. so it may have a mildly negative effect on extraction. 76 kg/h. It depends on the total copper plating rate in the cell and the electrolyte flow rate through the cell. Also as noted previously. Note that the total mass flow through EW must equal the total mass flow through stripping due to mass balance. but in most cases it is about 20-40% [6]. . The strip solution flow rate is then given by the copper production rate and the copper : As = dMCu/dt = Cu+2] strip 5707. and this must equal the mass flow through EW. i.546 g x 10-3 kg x 3600 sec dt 100nF sec mol g h {94} 43 .! Thus the rich electrolyte cannot be run directly into the EW cells! A flow configuration that accommodates these requirements is shown in Figure 18. The storage tank provides for a consistent composition electrolyte in each cell in the EW plant. ensuring that cathode quality is uniform across the cell.The fraction returning to SX ranges from 3-100% (100% means no splitting of the flow). The rich electrolyte copper concentration is [Cu+2]R. For a plant producing 50.) However. The lean electrolyte exiting EW is used to dilute the rich electrolyte from stripping so that after passage through EW. This requires that the flow rate through EW (designated E) must differ from that through stripping.[Cu+2]lean elec. The mass flow through stripping is As([Cu+2]rich elec.518 m3/h 15 kg/m3 {93} (mass per unit time / mass per unit volume).).76 kg/h = 380. Take as an example a rich electrolyte with 45 g/L Cu +2 and a strip copper of 15 g/L. the concentration changes through stripping and EW may differ significantly. The difference in concentrations [Cu+2]R – [Cu+2]L = copper . the specific flow rate of electrolyte to each cell should be around 0. and for the lean electrolyte it is [Cu +2]L. The electrolyte flow to EW is E and its concentration is [Cu +2]t.e. This in turn means that [Cu+2]into EW [Cu+2]rich elec. The arrangement results in small change in copper concentration across the cell (about 2-3 g/L). its concentration equals that of the lean electrolyte. all of which is transferred to the electrolyte during stripping.12 m3/h/m2 plating area.[Cu+2]lean elec.000 t/y copper the hourly production rate is 5707. The change in copper concentration across the cell can be represented as [Cu+2]tL. The copper delta for stripping (difference in copper concentrations for the rich and lean electrolytes) is often quite a bit higher (up to 15 g/L) [6]. (We are omitting the lean electrolyte bleed at this point. The strip solution flow from/to SX is designated A s m3/h. The copper production rate is set by the net amount of copper taken up by the organic in extraction. The copper concentration drops by 2-5 g/L as it passes through each cell.). Then the lean electrolyte [Cu +2] is 30 g/L. The copper production rate from equation {25} and equation {26} is: dMCu = (j CE)(Ac N S) mol x 63. E([Cu+2]into EW . at least when higher end current densities are used. Designate the specific flow rate as fs. When that is considered then mass flow through stripping = mass flow through EW + bleed. and values for j. The total plating area is Ac N S.185499 x 10-5 (j CE)(Ac N S) [Cu+2]tL {96} Ideally. N and S are to be chosen. where Ac is the total plating area per sheet (both sides). CE. Similarly.12Ac N S = 1. Some of this returns to stripping.12-0. Then.12 (or a similar value) Ac N S {97} With these two equations we can solve for [Cu+2]tL: E = 0.14 m3/h/m2 plating area. is given by the copper production rate divided by the change in copper concentration across each cell: E = 1.= 1. fs = E = 0. The [Cu+2]t is such that after EW the electrolyte [Cu +2] = lean electrolyte [Cu+2].185499 x 10-5 (j CE)(Ac N S) kg Cu/h {95} Figure 18.12 44 . the flow rate to the cells.185499 x 10-5 (j CE) {99} 0. S is the number of cells. Ac. Schematic of flows between SX and EW and how electrolyte storage is used. This allows for [Cu+2]stripping to be > [Cu+2]EW. and [Cu+2]R to be > [Cu+2]t into EW. the flow rate E should be such that the specific flow rate is 0. E. and some to the tank.185499 x 10-5 (j CE)(Ac N S) [Cu+2]tL {98} [Cu+2]tL = 1. 12 m3/m2/h = 140.816 kg/m3 {100} 0.5 m3/h E = 2027 m3/h E – As = 1647 m3/h The fraction of spent electrolyte going to SX stripping is: Fraction to SX = As = 18. Then for each cell.12 If fs is smaller.8 {104} There will be 141 cells in use. [Cu+2]tL = 1.2 kg/h = 50.8% E {106} 45 . maintenance and cleaning. The mass balance for copper in Figure 18 is: dMCu = As([Cu+2]R – [Cu+2]L) = E([Cu+2]t – [Cu+2]L) {101} dt Hence.) The nominal copper production rate will be: 1.078 t/y {105} The flow rates through EW will be: As = 380. with a central cathode harvesting area.816 kg/m3 = 2027.185499 x 10-5 x 300 x 95 x 2 x 60 x 141 kg/h = 11433. In a real tankhouse. Then a tankhouse will have an even number of cells. the number of cells then is given by using either equation {96} or equation {97}: S= E Ac N fs = 2027.22 m3/h {103} Assuming a conventional cell with 60 steel cathode starter sheets. E = As [Cu+2]strip {102} [Cu ]tL +2 E = 350. (Typically the tankhouse is split in two.518 m3/h x 15 kg/m3 2.185499 x 10-5 x 300 x 95 = 2.Take as an example j = 300 A/m2 and CE = 95%. longer residence time for Cu+2 to be depleted).22 m3/h 2 m2 x 60 x 0. then the concentration change will be larger (lower flow. there will be more cells in order to allow for downtime. with 1 m x 1 m plating area per side. As mentioned previously.g. the coatings are expensive.The strip [Cu+2] is 15 g/L (as specified) and the [Cu+2] from the tank to the cells is 32. column flotation to remove entrained organic) [16]. nodular deposits with consequent increase in short circuits and loss of current and energy efficiency. However. In copper EW hoods are starting to be used to prevent acid mist 46 . too high and spalling PbO 2 particles from the anodes may remain suspended in solution longer and contaminate the cathodes. there has been a move towards adopting DSA's into EW plants [20. Removing lead metal and compounds from the plant also improves worker health and safety.) Thus air sparging of the electrolyte in the cells is also practiced to improve mass transport of Cu+2 to the cathode surfaces. and results in cost savings in energy utilization. higher current density and higher cathode surface area increase the copper plating rate in the cells. Thus. A high concentration is needed to allow for good quality copper plating at high current density. What is required is a decrease in the boundary layer thickness to steepen the Cu +2 concentration gradient and increase the limiting current density. The gas is much too toxic to allow it to escape into the tankhouse and the atmosphere. The lower the copper concentration drop across the cell. synthetic additives. and this too requires a higher rate of flow of electrolyte to the cells. which in turn lowers the applied voltage.21]. Higher current efficiency. it is also captured for use in leaching operations. Electrolyte clean-up processes may be conducted on solution flowing from the first tank (e. Maintaining evenly spaced and parallel cathodes and anodes is also increasingly important as current density increases. these are composed of a titanium substrate coated with platinum-group metal oxides. both of which are cost savings. there is a limit to how high the specific flow rate can be. (Recall that practical current densities are roughly 2/3 of the limiting current density. These have been in use for many years in nickel EW plants where chlorine gas evolution occurs at the anodes. this will tend to produce rough. However. Sometimes there are two electrolyte storage tanks. Of course. There has been a move to add hoods or covers to copper EW cells [21]. These anodes do not required CoSO 4 addition either and labour and down time to clean lead corrosion sludge out of cells is no longer necessary. This has necessitated development of new. Current densities of up to 450 A/m2 are being employed in some modern plants [21]. This necessitates a higher circulation rate through the cells (E). The lean electrolyte should probably not have <30 g/L copper. the higher the specific flow rate must be. Higher flow rates also incur greater cost for pumping. the cost savings in energy consumption appear to more than compensate for the cost. Another problem is that traditional additives used to promote smooth plating (guar derivatives) begin to break down more rapidly as current density begins to exceed 300 A/m 2. They dramatically lower the oxygen evolution overpotential. Nevertheless.8 g/L. Recent developments Dimensionally stable anodes (DSA) have been in development for many years. This becomes all the more important at higher current densities and with air sparging. it also prevents manual inspection of the cells to look for problems.05916/n)log(1/aMn+) {107} Note again that orders of magnitude change are needed in ionic activities to effect large changes in E. M (28) The slopes are derived from the Nernst equation at 25°C: E = E° . From equation {107} as the metal ion concentration drops. In electrowinning it is desirable to have high electrolyte concentrations to maintain high energy efficiency (lower applied voltages). It obviates the need for foaming agents to reduce acid mist. and E 47 . E Mn+/M also decreases. However.accumulation in the tankhouse.(0. The values at the right hand axis are E° values for the various metal ion reductions (where a = 1 m): Mn+ + ne. Mists are directed to scrubbers. These E values are essentially plots of E for the half reactions relative to the standard hydrogen electrode (as per the Nernst equation) versus ion activity (on the molal scale). Electrowinning other Metals Figure 19 shows potential/molal ion activity relationships. New types of sensors (such as online monitoring of cell temperatures and voltages) are needed to detect short circuits. If one were to set up a cell with the H +/H2 half reaction occurring on Pt electrodes. The exchange current is the current when both the anodic (H2 = 2H+ + 2e. the more noble the metal is said to be. Plots of rate of hydrogen evolution on the surfaces of various pure metals are shown in Figure 20. Potential-ionic activity relationships for several metal ions and H +/H2 on Pt and O2/H2O on Pt at 25°C. The higher this E° value. and aluminum corrodes. It is this current that is estimated by the * But. why then can aluminum be put in contact with water without noticeable effect? Because of a very strong. or as per the Nernst equation under other conditions).EO2/H2O. making the operation inefficient. This is illustrated in Figure 21. the thermodynamic potential for the couple (0 V when aH+ = 1. The potentials are so negative that only H2 evolution could occur. making the thermodynamically favourable process very slow. The rate at which electrons are being exchanged between H + and H2 is the exchange current density. No voltage would exist between the two electrodes and there would be no change in composition in either half cell. becomes a bigger negative number (Eappl has to increase): E = EMn+/M . usually at a potential close to E. thermodynamically speaking. both at identical rates. Au + (E°Au+/Au = 1. PH2 = 1 atm. Referring back to Figure 7 this is the point where the linear regions of the polarization curves intersect. H 2 would be undergoing oxidation and H+ would be undergoing reduction to H 2. For these metals hydrogen evolution is not a concern. The same considerations apply to titanium 48 . For metals with E°Mn+/M < 0 V the rate of hydrogen evolution relative to metal plating rate is a crucially important question. but. Most metal sulfates are soluble to a maximum of around 1 M. EO2/H2O is near 1. metal ions with E° > E°O2/H2O cannot electrowon from aqueous solution.= H2 here) half reactions for a given couple have the same rate. that oxide coating gets broken down. so that hydrogen evolution is the favoured process. adherent oxide coating that forms on the Al surface and protects it from corrosion. They are too strongly oxidizing and the simple cations do not persist in water.* Relatively few metals have E° Mn+/M > 0 V.) Most metal ions have E° Mn+/M < 0 V. Generally. there would be no net effect.here) and cathodic (2H+ + 2e.69 V) cannot be formed in water at any practical concentration. for instance. The values plotted are the exchange current densities (symbol Jo). lower concentrations will necessarily result in lower limiting current densities.Figure 19. So. water would be oxidized by them. and with both half cells under identical conditions. (Some complexes of gold have lower E° values and can persist in aqueous solutions. However. In a sufficiently acidic or basic solution. so electrolysis on concentrated solutions tends to be feasible. Aluminum and magnesium cannot be electrowon from aqueous solution. Note the positions of the H+/H2 and O2/H2O lines. This is another case of kinetic stabilization.23 V {108} In addition. the exchange current densities can be at least partially correlated with the strength of the metal-hydrogen bond. atomic number. This spans almost 12 orders of magnitude! Platinum has the highest exchange current density. This forms an adsorbed H atom. consider the nature of the H+ reduction process on a metal surface. They range from about 10 A/m 2 to 5 x 10-9 A/m2. unfortunately it makes the technology very expensive. The system just described is at equilibrium. but degrades under reducing conditions.6 V). For instance.) * There are correlations that provide something of an explanation for the variation in exchange current densities. Data from [22]. but it is even more expensive. This layer is stable in an acidic and oxidizing environment. but the forward and backward rates are identical. Two of these H atoms 49 . To make sense of this. Figure 20. but equilibrium is a dynamic situation. In essence.(E°Ti+2/Ti = -1. then accept an electron. It is a very good catalyst for promoting hydrogen evolution. H+ may adsorb onto the surface. Much research has gone into searching for alternative electrode materials that are cheaper. reaction is still occurring. See text for explanation of exchange current. which also forms a stable TiO 2 layer on its surface. Plots of log (exchange current density) for H + reduction vs.* Note the enormous range of current densities for the H +/H2 exchange current density. intersection of the Tafel regions in Figure 21. (This is why fuel cells have used platinum electrodes. a higher J o means that the overvoltage for the hydrogen evolution half reaction is smaller. Rhenium has a very similar Jo to Pt. a high-energy intermediate state (adsorbed H) involves a highly endothermic reaction that is not very favourable. Log Jo. Figure 21. But. Mn(OH)2. the thermodynamic half reaction potential.76 V. Indeed. At all pH where Zn+2 can predominate H2 evolution is the thermodynamically favoured process. though not so high as to precipitate an Mn(II) hydrolysis product. but much slower H + reduction. which is well below H+/H2. This is illustrated from the Eh-pH diagram in Figure 22. The solution must have a relatively high pH. and again this would slow the overall reaction. Thus it is practically feasible to electrowin zinc metal from ZnSO 4 solution. If the metal-hydrogen bond is very weak.) This same phenomenon allows Mn EW as well.2 Zn+2 is no longer dominant. But if the metal-hydrogen bond is very strong. will readily deposit the parent metals on the zinc surface. And in this case high acidity cannot be tolerated. then the electron transfer is favourable. At this point the anodic and cathodic currents are identical. E°Zn+2/Zn is -0.must then combine to form H2 gas. Small amounts of Co +2. Zn EW is practiced from quite acidic solutions containing on the order of 100 g/L H2SO4. Beyond pH ~6. The lines usually meet near E. but rather solid ZnO is. which uses acid generated in EW. The rate of Zn plating is much greater and outruns the much more thermodynamically favourable. E° Ni+2/Ni and +2 50 . the rate of hydrogen evolution on pure Zn is very slow (J o = 3. this inhibits the electron transfer. Purity of the electrolyte is a critical issue in Zn EW.2 x 10-7 A/m2) which indicates a high overvoltage for H2 evolution. albeit with at best about 60% current efficiency. Ni . Extrapolation of linear Tafel region lines to obtain the exchange current density. the rest going to make H2. then the formation of H2 is inhibited. etc. Consider the example of the Zn +2/Zn couple. Both half reactions proceed at this same rate so that there is no net change in composition. If the metal-hydrogen bond is fairly strong. (This helps with conductivity of the electrolyte and also is important for the leaching process. The H +/H2 and Ni+2/Ni lines intersect at pH 3.01 atm). the cathodes and anodes are physically separated using porous cloth bags that surround either the cathodes or the anodes. Below this pH H+ is easier to reduce to H2 than reduction of Ni+2 to Ni. Metals in the Group VIII block of the periodic table exhibit especially high rates of hydrogen evolution. Here bags surround the cathodes. This is illustrated in Figure 23. thus opening up a kinetically facile pathway for the thermodynamically favoured process. so now hydrogen evolution has a site on the Zn cathode surface where it can readily occur. EhNi+2/Ni > EhH+/H2. introduce hydrogen and oxygen impurities into the solid and passivate the electrode surface. Figure 22. above pH 6. E°Ni+2/Ni = -0. Hence one of the biggest concerns in zinc hydrometallurgy is electrolyte purification. However. relative to the level in the anode 51 . Above this pH Ni+2 is the stronger oxidant.24 V and the rate of hydrogen evolution on Ni is relatively high as well. Pd and Pt column have the highest rates of any group.4 Ni(OH) 2 s precipitates and this would foul the cathodes. So how can we electrowin Ni metal and those like it? Consider again the Eh-pH diagram in Figure 22. The rate of H2 evolution on Co and Ni is quite high. Partial Eh-pH diagrams for three M-H 2O systems. The Ni. and in practice. This maintains a small hydrostatic head in the bags. Then current efficiency for zinc EW plummets and hydrogen gas is produced. Hence the pH must be kept at 3 < pH < 6. at least). In practice. This means that Ni+2 can be electrowon above this pH. Then regions of say Co metal form on the zinc surface. Nickel sulfate solution is pumped into the cathode bags (called the catholyte). it would be kept low to avoid the danger of explosion.4 (at these ionic activities.E°Co+2/Co > E°Zn+2/Zn. The H2 pressure is low (0. For instance. Numerous other configurations are also possible. All this adds considerable expense. This prevents anolyte solution from coming into contact with the cathodes. In the anode chamber acid is produced via oxygen evolution (in sulfate medium): SO42. Nickel may also be electrowon from a chloride electrolyte. Then the anode reaction is 2Cl . Figure 23. Another variation on Ni EW is to use chloride medium. Toxic Cl2 gas must be collected and is used in leaching operations. This adds to the energy costs. or EW will cease. The conductivity of the electrolyte is also lowered by having to avoid H+ (which is the best ionic conductor) and use Na 2SO4 instead. Hence again separate cathode and anode chambers are required. though it is more corrosive. The anolyte is pumped out at the same rate that the catholyte is pumped in. neutralized catholyte is pumped back into the cathode bags. Again. Then Cl 2 gas is formed at the anode. Hence solution flows from the cathode chamber to the anode chamber. it is imperative that powerfully oxidizing Cl2 not come into contact with the cathodes.compartment (the anolyte solution).+ H2O = H2SO4 + 1/2O2 + 2e- (29) The increasing acid concentration is prevented from coming into contact with the cathodes using the cloth bags.= Cl2 g + 2e-. 52 . it is highly toxic and it is very useful in leaching. The anolyte is continuously neutralized with NaOH. However. Now the anode chambers must be enclosed and the Cl2 gas must be collected. Schematic illustration of a nickel electrowinning cell. nickel is valuable enough that this can be tolerated. This has advantages in leaching. Nickelenriched. This is generally not pure enough for most applications. The reduction of H+ to H2 at the electrode surface necessarily raises the local pH there.e. To overcome this B(OH)3 (boric acid) is commonly added. copper metal is oxidized at the anode. Staying with the example of copper electrorefining. The most viable method for purification of copper is electrorefining (ER). The process produces high purity copper (>99.8% pure. Cu = Cu+2 + 2e- E° = 0. In this process. despite being about 99. This can lead to precipitation of Ni(OH) 2 with attendant passivation. Electrochemistry figures very prominently in metal production. Although this represents a loss of current efficiency. The Cu+2/Cu reduction potential lies well above the H+/H2 line at all pH where Cu+2 is predominant. This weak acid reacts with OH . However. the impure copper is cast into anodes and placed into electrochemical cells.34 V (31) and reduced at the cathode. neglecting scrap) involves either electrorefining or electrowinning. there are other voltage drops that result in a cell voltage that is somewhat less than 0 V. Cu+2 is simply a much stronger oxidant and so copper can be electrowon from acid solution.34 V (30) E° = 0. Clearly E° for this process is 0 V.= Cu+2 The cathodes are negatively polarized and the anodes are positively polarized. The impure copper is anodically dissolved by means of an electric current. Cu+2 + 2e. In the process the impurities present in the anode copper either stay with the anode or are precipitated as insoluble salts or are removed from the electrolyte by taking a bleed stream and purifying it. <0.99% Cu. The activity of copper in the anodes is only slightly less than one. Nickel and cobalt also may be purified by electrorefining. Small bubbles growing and dislodging generate significant mass transport. and the copper concentration in the electrolyte is typically on the order of 1 m. (The cell voltage is E cathode . Electrorefining [23] Most of the world’s copper is produced pyrometallurgically. The Eh-pH diagram in Figure 21 also illustrates why copper can be so easily electrowon. The dissolved Cu +2 is plated onto high-purity copper cathode starter sheets.Eanode < 0. one side benefit is that it produces a great deal of mixing in the vicinity of the cathode. Hence the great majority of the world’s primary copper production (i. Hence the thermodynamic cell voltage in practice is actually quite close to 0 V.In practice some degree of hydrogen evolution occurs in nickel electrowinning.) The following also contribute to the required cell voltage: 53 .to help buffer the local pH in an acceptable range.004% S + traces of other metals). Electrical resistance (0.01-0.02 V. Leads used to connect the electrodes to the power supply - called busbars.) Electrolyte resistance (0.11-0.13 V. Resistance to electrical conduction by means of ionic migration; lower than for EW electrolytes.) Cathodic overvoltage (0.04-0.08 V. This is due largely to organic coatings on the cathode copper surfaces. Organic additives are added to promote a smooth, continuous deposit. Alternatively periodic current reversal may also be employed to promote smooth deposit growth.) Anode and cathode connections (0.03-0.06 V. Electrical resistance of the connections.) Anode polarization (up to 0.01 V. Voltage required to corrode the anode copper.) The total cell voltage is typically 0.25-0.3 V. The greatest contributor to the voltage is the electrical resistance of the electrolyte. Sulfuric acid concentrations are kept high to keep the resistance of the electrolyte low. Since the solution resistance of the electrolyte increases directly with the separation between anode and cathode, a narrow anode-cathode spacing is desirable. (ER tankhouses also involve very large numbers of cathodes and anodes, so minimizing the gap also is good for lowering capital costs of the associated infrastructure.) However, cathodes need to grow to a specified thickness (about 1.4-1.9 cm in electrorefining), starter sheets (the substrates onto which the metal is plated) and anodes may not be perfectly planar nor vertical and polycrystalline protrusions called dendrites can grow out from the cathodes towards the anodes, causing short circuits and loss of metal plating efficiency. For all these reasons the gap between cathodes and anodes in ER is roughly at least 2.5 cm. Note that in ER as the cathode grows, the anode gets thinner. The conductivity of typical ER electrolytes is 0.5-0.7 -1 cm-1. Typical current densities are 190-260 A/m2 and cathodes typically are 1 m x 1 m. Given a cathode-anode spacing of 2.7 cm, a current of 250 A and a conductivity of 0.7 -1 cm-1, the resistance of the electrolyte would be, R = (1/0.7 -1 cm-1) x 2.7 cm x (1/104 cm2) = 0.000386 {109} From V = IR the voltage drop would be, V = 250 A x 0.000386 = 0.097 Volts {110} which is close to the lower end of the typical range of voltage drops due to solution resistance listed previously. Impurities 54 The table below lists the common impurities in anodes and the purer cathodes. The main impurities are As, Bi, Ni, Pb, Sb, Se and Te. If an element has an E° that is less than E°Cu+2/Cu (0.34 V), then it can dissolve from the anode. (A lower standard reduction potential means that the element is more easily oxidized.) For an impurity metal in the anode, Manode = Mn+aq + ne(32) The reduction potential (for the reduction half reaction) is given by, E = E° - 2.303RT log aManode nF aMn+ {111} The activity of metal impurity in the anode will be low. If the activity of the ion in solution is also low, the reduction potential will be close to E°. Table 4. Impurities in copper before and after ER. Impurities Anodes, % Cathodes, % As 0-0.3 <0.0002 Bi 0-0.001 <0.0001 Fe 0.002-0.03 <0.002 Ni 0-0.5 <0.001 Pb 0-0.1 <0.0005 Sb 0-0.3 <0.0002 S 0.001-0.003 <0.001 Se 0-0.02 <0.0002 Te 0-0.001 <0.0001 Ag trace-0.1 <0.001 Au 0-0.005 <0.00001 PGM* trace *PGM = platinum group metals (Rh, Ir, Ru, Os, Pd, Pt) Iron, for example, has E°Fe+2/Fe = -0.41 V. It is much easier to oxidize than copper. Now, however, since the reduction potential for Fe +2 + 2e- = Fe is substantially lower than that for Cu+2 to Cu, copper will preferentially plate over iron. (Fe+2 having a lower E° than Cu+2 means that Cu+2 is a stronger oxidant; Fe+2 is harder to reduce.) On the other hand, elements with an E° greater than that of Cu+2 will not dissolve from the anode. Thus electrorefining is very selective. Only species with reduction potentials similar to that of E° Cu+2/Cu will dissolve from the anode (at least by anodic oxidation) and plate at the cathode. Even then, kinetic factors also would come into play. If a species has a substantial cathodic overpotential for reduction, its tendency to reduce and co-plate with copper at the cathode would be diminished. A partial table of standard reduction potentials is reproduced below. The first column indicates the principal species that would form in solution. Inspection of the data indicates that mainly Fe, Ni, Pb, Sb, As and Bi should dissolve. 55 The other way that impurities can enter the cathodes is by occlusion (or entrainment). Cathode surfaces are polycrystalline and not perfectly smooth. Microscopic cavities form, which contain the electrolyte. These may grow over with Table 5. Some standard reduction potentials relevant to Cu ER. Oxidized Half reaction E° V form Au+3 Au+3 + 3e- = Au 1.42 + Ag Ag+ + e- = Ag 0.80 +2 +2 Cu Cu + 2e = Cu 0.34 BiO+ BiO+ + 2H+ + 3e- = Bi + H2O 0.32 HAsO2* HAsO2 + 3H+ + 3e = As + 2H2O 0.25 SbO+ SbO+ + 2H+ + 3e- = Sb + H2O 0.21 + + H 2H + 2e = H2 (for reference) 0 Pb+2 Pb+2 + 2e- = Pb -0.13 Ni+2 Ni+2 + 2e- = Ni -0.23 Fe+2 Fe+2 + 2e- = Fe -0.41 * Or perhaps H2AsO3. copper metal and trap small amounts of electrolyte within the cathodes. Impurities in the electrolyte, including copper sulfate and sulfuric acid are then occluded. This is the principal origin of impurities in the cathodes. Notes on specific impurities follow. (a) Gold and platinum group metals do not dissolve from the anode. Their reduction potentials are too high. These metals stay in the residues of the anodes. Anodes are not completely dissolved. They are removed and replaced after about 85% reaction. Precious metals may be a significant source of revenue. (b) Silver does partially dissolve, Aganode = Ag+ + e- (33) but a low concentration of sodium chloride (~0.05 M NaCl) in the electrolyte precipitates most of this as insoluble AgCl: Ag+ + Cl- = AgCl s 1/Ksp = 5.6 x 109 (34) The net reduction reaction is, Aganode + Cl- = AgCl + e- (35) In fact, E° for Ag+/Ag is 0.80 V. This is considerably higher than that for Cu +2/Cu (0.34 V). This should preclude silver dissolving in the first place. But, 56 g. but form quite insoluble precipitates of PbSO 4 and a basic tin(II) sulfate. they are more easily oxidized) and can enter the electrolyte as solution species (e. Most of the AgCl reports with the anode residues. cobalt. and chloride is present in the electrolyte.g. The differences in the two potentials are quite small and this might allow some Bi to be electrochemically deposited.). The potential at the cathode is around E° Cu+2/Cu = 0. see table of reduction potentials on previous page).e. The E° for BiO +/Bi is 0.= Ag + Cl- E° = 0.5 m3 of electrolyte per tonne of copper metal plated. Cu2S etc. (d) Lead and tin (tin was not mentioned in the tables above) are readily corroded. the strong acid of the electrolyte would redissolve them in some instances. in principle. A bleed stream is taken from the electrolyte and subjected to purification to remove the impurities from the circuit.22 V (36) Since this is less than E°Cu +2/Cu.) Electrolyte purification 57 .) This illustrates the beauty of electrorefining.34 V) can’t electrochemically plate at the cathode. while those that have E° >0. The case of bismuth is different. (e) Arsenic. In principle the impurities could plate at the cathode. which are not electrochemically oxidized.g. This is the main mechanism by which impurities enter the cathodes. respectively. even if they could plate. As a result these elements do not dissolve to any significant degree in the electrolyte. depending on relative concentrations. Fe+2 and Ni+2.g.34 V. i. Some cathode contamination by occlusion occurs. selenium and tellurium. which is quite similar to E° for Cu+2/Cu. Cu2Se. These compounds are too stable to corrode. the potential for copper reduction is too oxidizing to allow less noble ions to reduce. selenium and tellurium form stable copper compounds with many metals (e.e. Nevertheless. E°Co+2/Co = -0. Whatever does end up in the cathode results from occlusion of suspended particles in the electrolyte. all the corroded species build up in the electrolyte over time and eventually would contaminate the cathodes by occlusion. Ag2Se. These elements mainly report with the anode residues. possibly as a result of forming stable compounds with copper (e. (c) Sulfur.34 V can’t be corroded from the anodes.10. HAsO2. Much of the antimony (~60%) and some of the arsenic (~25%) end up staying in the anode residues. conversion of anode silver to solid AgCl is favourable.AgCl + e.28 V). (Besides. It is evident then that mechanisms of dissolution may be complex. Silver may be present as a copper-silver solid solution and as compounds with sulfur. The purification method is thermodynamically tuned to be very selective for copper. arsenides) during anode casting. Because of this they have to be removed from the electrolyte. iron and nickel are all less noble than copper (i. BiO+. the bleed rates are also low (0. Since the impurities concentrations in the anodes are low. bismuth. Ag2Te4.32 V. Co+2. but many have E° < E° H+/H2 (e. The impurity metals that are more easily corroded than copper (E° < 0. The first stage produces quite high purity copper cathodes. The adsorption of glue to varying degrees on the cathode surfaces results in a slightly higher overall resistance and a consequent voltage drop. Other additives are also added for various reasons. antimony and bismuth are removed as well. What is left is a high concentration H2SO4 solution (1000 g/L or ~10 M). In the second stage an impure copper product is formed that can be recycled to the furnace for generation of copper anodes. The result is that protrusions are prevented from growing. high voltages are required. In order to maintain high plating rates. This is done in three stages. The cathodes may be used for production of arsenic or recycled to the smelter. Plating conditions Organic additives are introduced to the electrolyte for a number of reasons. Water is evaporated and the sulfate salts of the metals precipitate. repeated regularly. The copper product is highly contaminated with these elements. This is mostly recycled to electrorefining to make up for lost acid. This is done by adding copper metal (“shot”) and reacting it with oxygen: Cu + 1/2O2 + H2SO4 = CuSO4 + H2O (37) Evaporation and cooling of the solution precipitates much of the solution copper as CuSO4·5H2O. One method is to first plate copper electrochemically from the electrolyte. Arsine is extremely toxic and must be collected. This results in an insulating layer that inhibits further deposition of metal. This is essentially an electrowinning operation. These are called “liberator” cells. as was noted previously.) The third stage in particular is quite energy inefficient due to the low concentrations of ions in solution. One typical additive is “bone glue. the cell polarity is switched so that plating occurs briefly on the anodes and corrosion of copper from the cathodes. An alternative to the first EW stage is to remove copper as CuSO 4·5H2O. It may be sold. flocculants 58 . cobalt and iron sulfates. The same effect can be attained by a technique called periodic current reversal (PCR). The protrusions are the most rapidly dissolved.There are a number of methods available for purifying the bled electrolyte. This results in strongly reducing conditions at the cathode and some arsine gas (AsH3) is also evolved. nickel. A side benefit is that higher current densities can also be employed. For a brief interval.” a mixture of natural proteins. resulting in greater overall plating rates. The precipitation is driven in part by the resulting high H 2SO4 concentration (source of high sulfate concentration). This avoids build-up of these impurities in the process. After the third stage the solution contains mainly sulfuric acid. (In the smelter these elements are partially rejected as slags and dusts. In the third stage most of the arsenic. In arsenic-containing EW electrolytes there is also a danger of arsine generation. Glue acts by adsorbing particularly on the tips of protrusions. The product contains a little iron and nickel as well. It is added to minimize the growth of protrusions on the cathode surface. For instance. It is added at 1-10 mg/L. while the indentations are the slowest to dissolve. The resistance of the anode greatly increases and copper corrosions ceases. very high current densities result in rough cathodes with more occlusion of electrolyte. maintaining good circulation and keeping the current density low enough. the faster the copper cathode production rate. which increases the solubility of copper sulfate). In addition. The electrolyte is circulated at a slow rate (~0. Stray currents arise from fortuitous conduction pathways. This ensures that Cu+2 in solution is transported from anode to cathode. copper metal corrosion is much more favourable than water oxidation. The main sources of loss are stray currents to ground (1-3%). Titanium is expensive. much above 250 A/m 2 (or 300 A/m2 with PCR) anode passivation occurs. This has been attributed to rapid buildup of CuSO 4 in the electrolyte at the anode surface and resultant precipitation of CuSO 4·xH2O. the copper sheets are easily separated from the substrate. These tend to be thin sheets of copper metal (0. so the number of titanium sheets needed is a small fraction of the total number of cathodes being plated. However. These may be detected by infrared scanners. Efficiency and energy consumption Current efficiency for ER is on the order of 90-96%. conductive film on titanium. The cathode is only weakly negatively polarized (the thermodynamic cell voltage is almost zero). This would occur if the solubility of copper sulfate was exceeded. the higher the current density. The electrolyte is heated by steam coils to 60-65°C entering the circuit and leaves at 55-60°C.may be added to promote flocculation and settling of fine precipitates and solids derived from the anodes. after which they are mechanically removed from the titanium substrates. oxygen is slightly soluble in the aqueous 59 . Some heating of the electrolyte occurs as a result of the solution resistance. A cathode substrate is needed for plating copper. whereas the starter sheets are plated for one day. Due to an adherent. This occurs to a small extent. However. They are plated onto titanium blanks. Hence oxygen in solution can attack the copper in acidic solution and oxidize it.5-1 mm). The precipitate is nonconductive and forms an insulating coating on the anodes. This helps prevent impurity occlusion in the cathodes. operating the cells at higher temperature (~60°C. Passivation is prevented by keeping the copper concentration in the electrolyte well below saturation (typically around 40-45 g Cu+2/L). Electrorefining cathodes are plated for about 10-14 days. but the process is easier than previously practiced alternatives. Copper dendrites growing from cathode to anode cause short circuits. Cu + 1/2O2 + H2SO4 = + H2O + CuSO4 (38) Oxygen is not generated at the anode. The nodules are broken off to resume plating. They are plated for about 24 hours. such as spilled electrolyte. anode-cathode short circuits (1-3%) and reoxidation of copper cathode (1%). It also carries away impurities and aids in distribution of additives to the solution.02 m 3/min). In principle. slightly lower potential. The electrolysis is carried out in a parallel circuit. washed to displace the residues (“slimes”) and recycled back to anode casting. Oxygen may also oxidize ferrous to ferric to a small extent. Blister copper is the term for primary copper from pyrometallurgical processing. depending on the plant. Care is taken to ensure that cathode starter sheets and anodes are planar and aligned properly in the cells. In EW the energy requirement is 1900-2000 kWh/t Cu. Approximate mass and volumetric flows are included. which can be a substantial source of revenue. Other operational aspects Anodes and cathodes are set in cells in a parallel arrangement with each anode between two cathodes. and ferric can corrode copper: 2Fe+3aq + Cu = 2Fe+2 + Cu+2 (39) Quantities such as energy requirement and copper production rate may be calculated from the Faraday’s Law relationships. This ensures good uniform corrosion of anodes and plating of cathodes. This is about 1/10 the energy for copper cathode production (including the starter sheets). The residues may be collected and processed for precious metals recovery. The difference is due mainly to the low cell voltage in ER compared to the higher voltage in EW (~2 V). Note the substantial recycle of copper anode material and that a small amount of high-purity copper metal is produced from the bleed stream purification process. the purification process for the bleed stream requires some energy (roughly 25 kWh/t Cu produced). Hence ER for copper is not all that energetically demanding. At the start of the process a set of anodes and a set of cathode starter sheets are placed into the 60 . etc. spent anodes and residues. The energy requirement for copper ER with 96% current efficiency and 0. The process takes on the order of 10-14 days per set of cathodes. In addition to the ER process itself.25 V applied potential is 220 kWh/t Cu. The electrical contacts are profiled to keep the cathodes and anodes at the right spacing. Flowsheet Once the anodes have been about 85% consumed they are removed from the cells. which is about twice as long as electrowinning of copper. as well as avoiding direct short circuits between cathodes and anodes. The number of starter sheet cells is relatively small. Electrical contact is made through a system of copper metal contacts. This keeps all anodes at the same potential and all cathodes at the same. Cells are made of concrete with polymer liners or aggregate in plastic. Cells are operated in series and can be isolated for purposes of installing or recovering cathodes. Some statistics for typical copper electrorefining operations are provide in Table 6 that follows below. Starter sheets are quite thin.solution and enters from the air. not in series. Each cell contains 30-50 anodes. A flowsheet is presented in Figure 24 that schematically depicts an ER process with the bleed stream being purified by electrowinning as described previously. Figure 24.cells. The spent anodes are then recycled to anode casting. Spent anodes are washed to remove residues and the cells are washed out to recover the residues. Simplified flowsheet for a copper electrorefining process. Cathodes are melted and cast into now very pure form. then a new set of cathode starter sheets are installed and the process continues. The bleed rate to purification is intermediate (about 30 m3/100 t of copper produced). 61 . which is done by melting the copper metal. These may be treated for precious metals recovery. After 10-14 days the cathodes are removed. Each set of cathodes is removed and washed. For n anodes there are n+1 cathodes (opposite of electrowinning). Approximate distribution of copper mass is indicated. After two sets of cathodes are recovered from one set of anodes the electrolyte is drained away. Centre for Metallurgical and Process Engineering. V. If power is converted from AC to DC there is a small loss here as well. 62 . Fundamentals of Electrochemistry..04 Temperature (°C) 60-65 Circulation rate (m3/min per cell) 0. Short Course Notes.25-0.3 Energy consumption (kWh/t Cu) 220-270 * 150.15 Cl 0. and not be occluded in the cathodes.01-10 Fe 0. and Conrad.003 3 Electrolyte composition (kg/m ): Cu 43-50 H2SO4 180-200 Ni 2-10 As <0. X-7.04-0.35 Bi 0. V. Energy consumption refers to DC voltage requirements.7 L x W x thickness (m) ~1 x ~1 x ~0.15-0. in Hydrometallurgy: Theory and Practice Course Notes.02 2 Power: Cathode current density (A/m ) 190-280 Current efficiency (%) 90-96 Cell voltage (V) 0. Table 6.The circulation rate is high enough to ensure good migration of copper.000-400. B.000 tonnes/year! References [1] Adapted from Ettel.05 Weight (kg) 310-380 Center-line spacing (cm) 10 Lifetime (days) 20-28 Anode consumption (%) 80-87 Cathodes: Lifetime (days) 10-14 Starter sheet weight (kg) 5-6 Weight (kg) 130-165 Total metallic impurities (%) 0.A. UBC.1 x 1.5-99.001-0.03-0. However. p. [2] Ettel.2-1 Sb 0. circulation rates need to be low enough to allow dislodging particles of anode residues to fall to the bottom of the cell. Some typical copper electrorefining plant statistics. Copper production (t/d) 400-1100* Number of cells 800-2400 Cell inner dimensions L x W x D (m) 4-5 x 1.2 Anodes: %Cu 99.015-0. Schlesinger and A.V. Phoenix AZ. vol. 493-567.G.com/NR/rdonlyres/97088921-AD31-461E-A1209C2B791ACD26/0/lixsolve. 41-56.R. “Principles and practical considerations of copper electrorefining and electrowinning. 11B. Dixon. 217-221. solvent extraction and electrowinning world operating data. “Do’s and don’ts of tankhouse design and operation. 38.A. University of Capetown. Kenyen and D. 8. pp.” Proceedings of the Copper 99 – Cobre 99 International Conference. 2002. 21-24. 159-163. Nicol. ed. Journal of Applied Electrochemistry. 24-26. A. S. W. 19.. Feb. 10-13. “Plant practices and innovations at Magma Copper Company’s San Manuel SX EW plant. T. King. “Electrolytic copper -leach.” Hydrometallurgy Conference. Montreal. Retrieved Feb.. pp. eds. D. Oct.. 1999... Society for Mining. Hackl and D.G.. Prasad.N.L.. Influence of cobalt ions on the anodic oxidation of a lead alloy under conditions typical of copper electrowinning. 1980. 4th edn. [13] J..S. Extractive Metallurgy of Copper. Vol.” in Copper Leaching. 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P.C.K. and Electrowinning Technology. 2nd edn. electrical conductivities and viscosities of CuSO4/H2SO4 solutions in the range of modern electrorefining and electrowinning electrolytes. pp.[3] Nguyen. 169-186.” in Copper Leaching. 4. 20th Annual Hydrometallurgical Meeting. Hiskey.G.. Harris.T. Lix Reagent Solvent Extraction Plant Operating Manual. D. Oct. W. p. 63 . Beukes and J. 213-240.G. [10] C. Biswas.B. Eamon. Pergamon. Badenhorst.B. and Davenport. R. Pfalzgraff. Kennedy and T. [4] Price. G. Cognis Corp. M. Solvent Extraction. Metallurgical Transactions B. Davenport. G. 2009 at: http://www. [11] M. Solvent Extraction.cognis. “Copper electrowinning: theoretical and practical design.V. 18.K. .K. 22 pp. Glossary of Some Common Terms 64 . Illescas. 757-769.cognis.. Biswas. 1-4. C.” Randol Conference. Miller. 467-481.. Pergamon. Readett and P. 1635-1647. 1999. and electrochemical behaviour of metals. 10. Hutchinson. 2011. 6-10. B. 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S. “Solvent extraction . Soderstrom. 2000. 2009 at: http://www. and Electrowinning Technology. M. Armstrong. Scotland.R.J. Trasatti." Proceedings of Copper 2010 International Conference.[14] G.. (b) A very slightly different version of this article can be accessed at (Feb.. “The ISA process and its contribution to electrolytic copper. Kordosky. Dominguez. QC. Vol. Clayton. S. W. Extractive Metallurgy of Copper.J.. Electroanalytical Chemistry and Interfacial Electrochemistry. 1972. J. No. “Iron control strategies in copper solvent extraction plants. 1997.pdf [18] (a) J. 25/09 at: http://www.D.. ed. G. Olafson and C. [17] W. Hamburg. D. S. 4. Elsevier. Retrieved Feb.V. 2010.E. King..” Cognis Corp. 18. K. 3rd. Vance. Oct. R. electronegativity. “Ion exchange for iron impurity control in the base metal industry. Hahn. [20] Sandoval. M.. [21] Schlesinger. D. Jun. T. Davenport. Electrolytic hydrogen evolution in acid solutions. 349-372. Spence and M.. Retrieved Feb. Unger. pp. C. Mackenzie.G. S. M. Vol. Montreal.C and Davenport. pp. 17. 5. “Practical aspects of copper solvent extraction from acidic leach liquors. 5th edn.com/NR/rdonlyres/B159035F-B210-47AD-81ED96771D437E82/0/Iron_Control_Strategies_Randol_Conference. and Wassink.. and Robinson. Aqueous solutions are highly non-ideal. i. The thickness of the boundary layer is determined in large part by the extent of turbulence (agitation). 1 C of charged passed through a potential difference of 1 V involves 1 J of energy. Electrorefining An electrolytic refining process where impure metal is anodically corroded and replated onto cathodes. 1 VC 1 J. Activity Effective concentration. Bleed A solution taken from a process stream. If all the current goes to only one half reaction. Upon contact with the anode a short-circuit occurs. conductivity See resistance. The thinner it is. Bleeding A term for the action of removing a fraction of a process solution stream. 1 C/sec = 1 A (1 A = 1 amp). CE = 100%. Current density The current in A per unit area. 65 . See also bleeding. Side reactions that consume some of the electricity lower CE to <100%. AW Atomic weight in g/mol.Acid mist Gas bubbles breaking the surface of the electrolyte in the cells forms a fine mist in the tankhouse. May refer to nominal area (as in 1 m2 for a 1 m x 1 m sheet) or true area (which requires knowledge of actual surface area). Dendrite Same as a protrusion. Anode The electrode where oxidation occurs (loss of electrons). the higher the limiting current density can be. a polycrystalline mass or relatively fast-growing metal crystals arising from a metal cathode surface and growing towards an anode. This makes a solute behave as if its concentration is either higher or lower than its analytical concentration. meaning that there are strong interactions between solutes and water. Conductance. Coulomb The unit of electric charge. In EW it is taken from the lean electrolyte to minimize build-up of harmful impurities in electrowinning. often to prevent excessive build-up of impurities. Boundary layer The stagnant film in the vicinity of a solid in a solution (here an electrode). Current efficiency (CE) The fraction of the current (in %) that goes to the specified electrochemical half reaction. Cathode The electrode where reduction occurs (gain of electrons).e. Current The flow of electricity in C/sec. By definition. its Electrolyte A solution that contains ions.TS < 0. E = E° . E° must also be known at T. Electrolysis An electrochemical process forced to go opposite to thermodynamically favourable direction by application of an external voltage. Galvanic Describes the operation of an electrochemical cell when current flows spontaneously. Limiting current density The maximum possible current density for an electrolysis. which is expressed mathematically as G = H . A mole of reaction means that the molar quantities of reactants form the molar quantities of products for a reaction as written. kWh Kilowatthour a unit of energy (not power): 1 kWh = 1000 J/sec for 3600 seconds = 3. such as diffusion coefficient and ionic mobility of the cation. Lean electrolyte (or LE) The copper-depleted (somewhat depleted. acid enriched electrolyte solution that comes from electrowinning and returns to SX stripping. Favourable Same as spontaneous. F = the Faraday constant (96485 C/mole e-) and Q is the reaction quotient (has the same form as the equilibrium constant. represents a disequilibrium condition). Thermodynamically speaking. but does no equal K. not completely).314 nF J/mol K). ER Short for electrorefining. The Nernst equation allows calculation of the voltage for a cell under any conditions. T = absolute temperature. It is a function of physical properties of the metal cation (in EW). Faraday The charge of a mole of electrons = 96484. n = moles e-/mol of reaction.6 x 106 J. Live stripping See Stripping. In EW. At the limiting current density the concentration of the electroactive species at the cathode surface goes to zero.RTlnQ where R is the ideal gas constant (8. At constant temperature this becomes G = H . Nernst equation n The number of moles of electrons passed for a half reaction or a cell reaction per mole of reaction. the solution to be electrolyzed.(TS) < 0. EW Short for electrowinning. as well as mass transport factors such as the bulk concentration and the thickness of the boundary layer. the naturally favoured direction of a chemical reaction.Electrowinning Electrolysis for primary metal production. If the half 66 .6 C/mol e -. A reaction is favourable when there is a net increase in the entropy of the system plus its surroundings. the harder it is to oxidize the metal. symbol . E°Au+/Au = 1. Potential (electric) See voltage. there are ways to study one half reaction while not considering the other.69 V.g. Protrusion See dendrite. relating current to voltage. A measure of the tendency for a chemical species to accept electrons and be reduced relative to the standard H+/H2 half reaction. J/s or kW (kilowatts). resistance per unit length per unit area: /cm/cm2 = cm. Reduction potential A voltage. Short circuit Direct electric contact between anode and cathode. Overpotential Same as overvoltage.g. the additional voltage required to make a half reaction proceed at a measurable rate (As always half reactions never run in isolation. PLS Pregnant leach solution (from leaching. Specific conductivity is the inverse of resistivity in W-1cm-1.) Overvoltage See overpotential. The standard reduction potential is the special case where all species are present at unit activity. Polarization A potential difference. Resistance The constant of proportionality in Ohm’s law. e. Spent electrolyte Same as lean electrolyte.5 mole H2. Power Energy per unit time. then n = 24 and 24 mole H + is reduced to form 12 mole H 2. Resistance is the inverse of conductance or conductivity. Rich electrolyte (or RE) The copper-enriched. Also the extent to which current flow is impeded. 67 . also called Siemens. Noble metals Metals whose cations have relatively high reduction potentials.reaction is written H+ + e. Unit = ohms. Resistivity The intrinsic resistance of a material or solution. cathodic and anodic half reactions must be coupled. Current flows directly between cathode and anode and is lost entirely as heat.= 1/2H2. acid-depleted (somewhat depleted. The higher E° is.= 12H2. If the half reaction is written 24H + + 24e. V = IR. n = 24 mole e-/mol of reaction = 24 mole e -/24 mol H+ = 24 mol e-/12 mole H2. n = 1 mole e-/mol of reaction = 1 mole e-/mol H+ = 1 mol e-/0. e. then n = 1 and 1 mole H + is reduced to form ½ mole H2. However. not completely). electrolyte that returns from stripping to EW. causing one electrode to be positive relative to the other (rendered negative). by growth of a protrusion from a cathode. Au is the most noble of all metals. fed to solvent extraction). the cell can continue to be operated. By definition 1 VC = 1J. Voltage in itself is not a potential energy. Whether energy is released or input depends on whether the charge is attracted or repelled. A voltage is the electrical potential energy per unit charge. When 1/2 or 1/3 of the cathodes are harvested at a time from a cell.Stripping In EW the process of mechanically dislodging cathode sheets from starter sheets. If the potential is high. 68 . Tankhouse The building where EW or ER is conducted. the energy involved in passing the charge is high. The "per unit charge" term is crucial. This improves productivity and is called live stripping. Voltage Voltage. The units of voltage are volts (V). One can think of voltage as a potential for passing a charge. energy and charge cannot be defined other than in terms of each other.