ELECTROWINNING AND ELECTROREFINING OF

COPPER

A COURSE PRESENTED TO AMIRA P705A SPONSORS

April 2008

M J Nicol

Murdoch University

ELECTROWINNING AND ELECTROREFINING OF COPPER

Introduction 4 1 Redox Equilibria 5 2

2.1 Electrochemical Potentials 5 2.2 Single Electrode Potentials 5 2.3 Non Standard Electrode Potentials 7

Electrochemical Kinetics10 3 3.1 Faradays Laws 10 3.2 Reactions at electrodes10 3.3 Potential Dependence of Electrode Kinetics 11 3.4 Characteristics of the Butler-Volmer (BV) Equation 13

Mass Transport to Electrodes 16 4 4.1 Mass Transport Processes 16 4.2 Diffusion Layer Model 16 4.3 Mass Transport Correlations 19 4.4 Influence of Mass Transport on Electrochemical Kinetics 21 4.5 Mass Transport of Ions 21

Mass Transfer at Vertical Electrodes 24 5 5.1 Natural Convection 24 5.2 Application to Copper Deposition 26 5.3 Effect of Gas Evolution 26

Electrocrystallization 29 6 6.1 Influence of Kinetics on Deposit Morphology 30 6.2 Application to Metal Deposits 32

Current distribution in a cell 35 7 7.1 Types of Current Distribution 35 7.2 Primary Current Distribution 36 7.3 Secondary Current Distribution 37 7.4 Tertiary Current Distribution 40 7.5 Current Distribution in 3-Dimensional Electrodes 42

Energy consumption 43 8 Materials for cells and electrodes 44 9

9.1 Anodes and cathodes 45 Cell design 47 10 Tankhouse current distribution 48 11 Copper electrorefining 54 12

12.1 Cathodes and starter sheets 54

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12.2 Electrolytes 55 12.3 Current density 55 12.4 Anode passivation 56 12.5 Deportment of anode impurities 56

Electrowinning of copper57 13 13.1 Competing Reactions 58 13.2 Quality of Copper Cathodes 59 13.3 Additives in copper deposition 63 13.4 Anodes for Electrowinning 65

Electrowinning in Novel Cells 66 14 14.1 Packed or Fluidized Bed Cells 66 14.2 Forced Flow Cells 68

Appendix 70 15 15.1 Class Problems - EW Course 70 15.2 Workshop Problems - EW Course 75 Workshop Problem 1 75 Workshop Problem 2 76

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ELECTROWINNING AND ELECTROREFINING OF COPPER

1 Introduction his course examines the use of electrical power to reduce copper ions in solution to metal on a cathode from which the metal is subsequently removed. The power can also be used to oxidize copper metal to copper ions. The following sections

will deal with both the general theoretical aspects of the electrodeposition of metals in general with a focus on copper and also on the more practical conditions under which copper is recovered.

T Electrowinning is the recovery onto a cathode of a metal which will contain a range of impurities which are codeposited from the impure feed solution. The feed solution may be directly from a leach operation or may be partly or wholly purified prior to electrowinning such as in the SX/EW process. Electrorefining involves the electrodissolution of an impure anode (often from pyrometallurgical processes or preliminary electrowinning) into a solution from which it is electrodeposited as a highly purified metal onto a cathode which is often made from the pure metal. The crude anodes contain material which either dissolves and accumulates in the solution, or is insoluble and falls to the bottom of the cell as anode slime. The build-up of impurities in solution is controlled by removing part of the electrolyte (a bleed) and removing the unwanted metals by precipitation, or other means. The anode slimes are generally treated to recover valuable metals such as Ag, Au, Te, PGMs).

For many of the base metals, minimization of the energy cost in electrowinning or refining processes is an important consideration.

For every metal, the specifications are dictated by the major commodity metal exchanges, such as that in London (LME). For LME Grade A and Comex Grade 1 copper cathodes the specifications are given in Table 1.1

Table 1.1 Chemical specifications for copper cathode

Element LME Grade A Comex Grade 1 2 Se 2 2 Te 2 2 Bi 1 4 Sb 4 5 As 5 5 Pb 5

15 S 15 5 Sn 5

10 Ni 10 10 Fe 10 25 Ag 25

4

Clearly, the high purity of the copper required for commodity sales is difficult to produce except under tightly controlled conditions and many other methods of metal production (notably pyrometallurgy) cannot produce a suitably pure product. The copper produced by smelters is typically around 99.8% Cu, the remainder being comprised of a range of elements, including iron, arsenic, silver, gold and PGM.

2 Redox Equilibria This section is largely revision of the most important concepts and applications of electrochemical potentials as they apply to the reactions taking part during the electrowinning of copper, with the emphasis on reactions at electrodes.

2.1 Electrochemical Potentials

The reaction

2H2O + 2CuSO4 (aq) = Cu + O2(g) + SO42-(aq) +4H+(aq)

occurs in the cell during the electrowinning of copper .

It can (theoretically and practically) be separated into two half reactions

2H2O = O2(g) + 4H+(aq) + 4e

Cu2+ + 2e = Cu

Consider a cell composed of these two half-reactions as shown in Fig. 2.1.

Reduction Cathode

V

Cu Pb

eI

Oxidation Anode

Zn2+

Cu2+

Figure 2.1 An electrochemical cell

2.2 Single Electrode Potentials

In the case of the above cell, a voltage (or potential difference) of 0.89V would be measured by the voltmeter if the solutions contained copper and hydrogen ions of concentrations equivalent

5

to unit activity and is saturated with oxygen at one bar pressure. The copper electrode would be negative with respect to the lead electrode.

The cell voltage can be considered to be composed of a difference between the potential of the copper electrode and that of the lead/oxygen electrode,

i.e. E = V = EO2 - ECu It is only possible to measure the potential difference (or voltage) between two electrodes.

A scale of relative potentials can be devised if all potentials are measured relative to one electrode. For various reasons this standard reference electrode has been chosen as the Standard Hydrogen Electrode (SHE)

2H+ + 2e = H2

for which , at unit activity H+ and unit fugacity H2, E0 = 0.0V.

Thus for any redox couple,

Ox + ne = Red or

Mn+ + ne = M

the electrode potential(E) is the potential difference of a cell composed of this couple and a SHE.

i.e. E = E - E0HFor the above reaction the cell voltage(under standard conditions) given by

E = EO2o - ECuo = 1.229 0.340 = 0.889V The following table gives a selection of the Eo values for some couples of importance in the electrowinning and refining of copper.

Table 2.1 Some relevant standard reduction potentials.

Half reaction E0 , volts

Mn2+ + 2e = Mn -1.18 Fe2+ + 2e = Fe -0.44 PbSO4 + 2e = Pb + SO42- -0.35 Co2+ + 2e = Co -0.28 Ni2+ + 2e = Ni -0.26 Pb2+ + 2e = Pb -0.13 2H+ + 2e = H2 0 AgCl + e = Ag + Cl- 0.222 Cu2+ + 2e = Cu 0.34 Fe3+ + e = Fe2+ 0.77 Ag+ + e = Ag 0.80 O2 + 4H

+ + 4e = 2H2O 1.23 MnO2 + 4H+ + 2e = Mn2+ + 2H2O 1.33 Cl2 + 2e = 2Cl- 1.36 Mn3+ + e = Mn2+ 1.49 MnO4- + 8H+ + 5e = Mn2+ + 4H2O 1.51 PbO2 + SO42- + 4H+ + 2e = PbSO4 + 2H2O 1.70

Preferred Reduction or Cathodic Reaction Preferred Oxidation or Anodic Reaction (Reverse)

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The equilibrium potentials for the reduction of the more important metal ions are compared in Figure 2.2. Thus, the only metal which will be co-deposited with copper is silver because its potential is more positive than that for the reduction of copper ions. From a thermodynamic point of view, we would not expect lead, iron or nickel to codeposit with copper.

Order of Reducibility

-1.2

-1

-0.8

-0.6

-0.4

-0.2

0

0.2

0.4

0.6

0.8

Ag Cu H Pb Ni Co Fe Cd Cr Zn Mn

Eo ,

V

Figure 2.2 Standard reduction potentials for metal ions.

2.3 Non Standard Electrode Potentials

Most reactions are carried out under conditions in which the activities of the species taking part are not unity and one must be able to quantitatively account for deviations from the standard states. The well-known Nernst equation can be applied to calculate the equilibrium potential under any conditions. For a general half-reaction

Ox + ne = Red

E = E0 - RT/nF. ln{ ared / aox }

The formation of complexes between a metal ion and a ligand can have a significant effect on the value of the standard reduction potential for couples involving the metal ion. Fortunately in the case of acidic metal sulfate solutions, this is not a major effect.

For example, for the reduction of copper ions

Cu2+ + 2e = Cu E0 = 0.34V

and the presence of complexing ligands such as SO42- can change this potential. The effect of complexation can be viewed as a reduction in the activity of the free Cu2+ ion which will make it more difficult for it to be reduced to metallic copper i.e. one will require a more negative potential.

This effect can be quantitatively accounted for if the stability constants for the formation of the complexes are known.

Thus , for the relatively weak outer-sphere sulfate complex,

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2+Cu + SO42- = CuSO4(a) K1 = 5

from which

[Cu2+] = [CuSO4]/ [SO42-].K1

= [Cu]T/(1+K1.[SO42-]) where [Cu]T is the total copper concentration.

and one can calculate [Cu2+] after substitution of [SO42-]. However, this calculation is more complex because of the simultaneous equilibria involving the sulfate and bisulfate ions in acid solutions. One can make use of thermodynamic software packages to derive species distribution diagrams such as that shown in Figure 2.3 from which one can estimate the free cupric ion concentration under various conditions of acidity.

Note that the free cupric ion concentration increases as the acidity increases as a result of the fact that the free sulfate ion concentration decreases due to protonation to HSO4- as the acid concentration increases. Thus, although the acivity of copper ions in solution decreases with increasing concentration, addition of protons can increase the activity of the cupric ion.

0.0 0.5 1.0 1.5 2.0 2.50.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

M

M

H(+a)

Cu(+2a)

CuSO4(a)

SO4(-2a)

Figure 2.3 Effect of acid concentration on the species distribution in a sulfate solution containing 40 g/l (0.63M) copper at 40oC.

Thus, at 40oC, [Cu2+] = 0.61M and substitution in the Nernst equation gives

ECu = E0 - RT/2F. ln {1/[Cu2+]} = 0.340 0.006 = 0.334V. Cu

i.e. the effect is very small in this case and can effectively be ignored.

The thermodynamics of such systems can be conveniently summarized in EH/pH diagrams, with the one for the copper system in uncomplexing media shown in Figure 2.4.

Lower dotted line: 2H+ + 2e = H2+Upper dotted line: O2 + 4H + 4e = 2H2O

8

-1

-0.5

0

0.5

1

1.5

2

0 2 4 6 8 10 12

pH

EH

14

Cu

Cu2O

CuOCu2+

Figure 2.4 EH/pH diagram for the copper water system at 298K (0.1M activity of aqueous ions) The areas of stability of the various species can now be identified on the diagram. Thus, the Cu2+ ion is only stable at a concentration equivalent to unit activity in the grey area in the left-hand side of the diagram.

Several general features of these diagrams should be pointed out:

The oxidation state of the metal (or other central species) increases as one moves vertically in the diagram.

Thus, at pH 8 a vertical movement starts with Cu(0), moves through Cu(1) and into the Cu(II) area of stability.

The extent of hydrolysis increases from left to right. In the above case, the extent of hydrolysis increases from Cu2+ to Cu(OH)2 (or CuO) as we move from left to right at potentials above about 0.5V.

Soluble species are generally present on the left and sometimes (for amphoteric metals) on the right of the diagram.

The area of stability of water is between the two dotted lines. The consequence of this is that species with an area of stability below the lower line will be unstable in aqueous solutions reducing water to hydrogen gas.

Thus, because the line separating Cu2+ from Cu(s) is above the lower line, the reaction

Cu(s) + 2H+ 2+ = Cu + H2(g)

is not thermodynamically favourable at pH values below about 4 and, in fact, the reverse reaction is possible, i.e. we can use hydrogen to reduce copper ions.

Similarly, species with an area of stability which lies above the upper line will oxidize water to oxygen .

It should be pointed out that although these diagrams are invaluable in assessing and predicting what processes are possible from a thermodynamic point of view, any practical application of

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a possible reaction will necessarily require that the rate of the reaction is such that it can be made to occur in the time available for processing.

Thus, reference to the copper-water diagram above shows that it should be possible to electrodeposit copper from aqueous solutions. This follows from the fact that the line for the reduction of copper ions lies above the lower line i.e. the preferred reaction at the cathode in an electrowinning cell will be

2+Cu + 2e = Cu and not the reaction

2H+ + 2e = H2 which requires a more negative potential. You should be able to define the conditions required to produce Cu2O by reduction of cupric ions.

3 Electrochemical Kinetics These reactions involve the transfer of electrical charge (generally electrons) across interfaces between phases. Thus, in addition to the influence of reactant concentrations on the rate of such reactions, the electrochemical potential difference between the phases (or the concentration of electrons, if you like) also has a profound effect on the rate.

3.1 Faradays Laws

The rate of production of copper at a cathode is determined by Faradays law

m = Mit/nF (= 0.448 Mi g for t = 1 day and n=2) where m - mass of metal deposited (g) i - applied current (A) t - time (s) n -no. of electrons/mole of metal (e.g. 2 for Cu since the solution comprises Cu2+ ions ) M - atomic mass of the metal g/mole (63.5 for Cu) F is the Faraday (96487 coulomb/equivalent) is the current efficiency i.e. fraction of the current resulting in metal production. 3.2 Reactions at electrodes

Consider the simple cell shown below which consists of two copper electrodes in a solution or electrolyte of aqueous copper sulphate . A reference electrode has been added to the cell close to one of the copper electrodes. This enables us to measure the potential of the copper electrode as a function of the current passed (which can be varied by changing the voltage,V, applied by the power supply).

Note that the reference electrode could be another copper electrode or any suitable reference such as a Ag/AgCl electrode. Assume that it is the former.

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-+ PowerSupply

V

e(I) Ec

Cu2+

Cu Cu

Figure 3.1 An electrochemical cell

For the current I=0, Ec= Ea = Ee (equilibrium potential) = 0 .

As current is allowed to flow in the direction indicated, Ec will decrease and Ea will increase. The overpotential () as defined by = E - Eewill increase at the anode and decrease at the cathode.

Sign Conventions:

By above definition and that of E, cathodic overpotentials are negative and anodic overpotentials positive. Similarly, anodic currents are positive and cathodic currents negative

Rate of reaction at the cathode v = No of mol of Cu deposited/ unit time.

= I/2F (by Faradays law)

As the process is heterogenous, the current density (i = I/A, where I is the current and A the active electrode area) is normally used.

3.3 Potential Dependence of Electrode Kinetics

For a chemical reaction, the rate v is related to the free energy of activation (G#) by v = kT/h. a. exp(-G#/RT) where k and h are constants and a is the product of the activities or concentrations of the reactants taking part in the rate-determining step.

For an electrode process, we have seen that G =-nFE and it is reasonable, therefore, to expect the rate to be some exponential function of the potential.

Thus, for the above cell, we can draw a schematic current-potential curve for the reactions which occur as we change the potential. As we increase the voltage applied to the cell with the flow in the direction indicated, the right hand electrode will become an anode and copper will be dissolved at a rate proportional to the magnitude of the current. If we reverse the polarity and increase the current in the opposite direction, the right hand electrode will become a

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cathode on which copper will deposit. If we measure the potential of the electrode (relative to a copper zinc reference electrode) at each of the currents applied we can plot the current as a function of the potential as shown in Fig 3.2. (Remembering the sign conventions).

Figure 3.2. Schematic current-potential curve for the Cu2+/Cu couple

The above current-potential relationship is, in fact, a composite of two exponential curves as shown in Fig. 3.3. Note that the net current is zero at the equilibrium potential, Ee and that we have changed our reference electrode to, for example, a standard hydrogen electrode so that, for unit activity of copper ions, Ee is 0.34V.

Figure 3.3 Schematic current potential curves for electrochemical reactions.

Cu = Cu2+ + 2e

Cu2+ + 2e = Cu

ia

ia

Cu = Cu2+ + 2e

Cu2+ + 2e = Cuic

Ee

io

E, V vs NHE

12

Thus, the overall curve (black) is the algebraic sum of the currents due to the anodic (red) and cathodic (blue) reactions. At equilibrium (i=0), the rates of both reactions are equal but opposite and the current due to each reaction at equilibrium is known as the exchange current density (io).

These curves can be described quantitatively as follows

For the cathodic reaction, 2+Cu + 2e = Cu

ic = -nFkc [Cu2+] exp( -c FE /RT)

where kc is a heterogeneous rate constant and c (see below) a parameter known as the cathodic transfer coefficient. E is the potential with respect to any reference electrode. Note the negative exponent because the cathodic current will decrease with increasing potential and also the negative sign for the current.

Similarly, for the anodic reaction,

ia = nFka exp( a FE /RT) Note that in this case, [Cu] = 1 (solid )

At any potential, the net observed current i = ia + ic At the equilibrium potential, Ee , i = 0 and ia = -ic i.e. nFka exp(a FEe/RT) = nFkc[Cu2+] exp(-c FEe/RT) = i0 (Exchange current density)

Thus, remembering that = E - Ee one can write the current/potential relationship as follows, i = i0 [ exp{ aF(E-Ee)/RT} - exp{ -cF(E-Ee)/RT}] = i0 [ exp( a F/RT) - exp(- c F/RT)] This is the well known Butler-Volmer equation.

Note that the exchange current density is a function of the concentrations of the reactants and in this case varies with [Cu2+]1/ .

3.4 Characteristics of the Butler-Volmer (BV) Equation

For >0, the first exponential term in the BV equation is greater than unity while the second is less than one - the net current density is positive(anodic). For

Under these conditions,

Butler- Volmer Equationio = 1 A/m2, beta(anodic)= 0.4

-12

-10

-8

-6

-4

-2

0

2

4

6

-100 -50 0 50 100Overpotential, mV

Current Density,

A/m2

i = i0 [1+ (1-)F/RT -1 - (-F/RT)] = i0 F/RT This enables one to obtain i0 using the so-called linear polarization technique which is very common in corrosion science.

Figure 3.4 Calculated current- potential relationship

High-field Region

At large overpotentials, the cathodic contribution to the anodic reaction can be neglected and

ia = i0 exp[ (1-)F/RT] or

ln ia = ln i0 + (1-)F/RT which is the form of the empirical Tafel equation for an anodic process,

a = a + b.log i10 a with a = -b.log i10 oi.e. a = b.log10 (ia /io) The so-called Tafel Slope b = 2.303RT/(1-)F and is normally given in mV/decade current. Note that

i) the Tafel equation is written in terms of logs to the base 10.

ii) the exchange current density can be obtained by extrapolation of the Tafel line to the equilibrium potential ( = 0) iii) for a cathodic reaction, the corresponding Tafel equation is

14

c = b.log10 (io /ic) Typical schematic Tafel plots are shown in Figure 3.5. Note the common intercept on the ordinate of the anodic and cathodic extrapolations. From these plots, the parameters , c a and i0 can be obtained.

Tafel Plot of Butler-Volmer Equation

-0.6

-0.4

-0.2

0

0.2

0.4

0.6

0.8

1

1.2

-150 -100 -50 0 50 100 150 200Overpotential, mV

log10(i)

Figure 3.5 Calculated Tafel plots using the data from Fig. 3.4

Most electrochemical processes of importance in electrowinning occur in the high-field region.

Table 3.1 gives the electrochemical kinetic parameters for some common reactions encountered in the electrowinning of copper. Note that the exchange current density (the electrochemical equivalent of rate constant) for a particular reaction is dependent on the nature of the electrode and also the nature of and concentrations of the electroactive species.

Table 3.1 Electrochemical Parameters for Some Reactions at 25oC

Reaction Electrode E0V

Solution Tafel i0 slope, mV A m-2

2H+ + 2e = H2 Cu 0 1M H+ -5120 1.102+ -5Ni + 2e = Ni Ni -0.26 1M NiSO 120 2.1042+Cu + 2e = Cu Cu 0.34 1M CuSO 40 0.2 4+Ag + e = Ag Ag 0.80 0.1M Ag+ 460 1.10

2H2O = O2 + 4H+ -4 + 4e Pt 1.23 1M H2SO 120 1.104 -PbO2 1.23 120 1.10-6

Fe3+ + e = Fe2+ Pt 0.77 0.1M Fe2+,3+ 3120 4.10

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Calculate the overpotentials at the cathode and anode for a Cu electrowinning cell operating at 300 A/m2 with unit activity copper ions and protons. Use the data in Table 3.1.

4 Mass Transport to Electrodes The overall rate determining step in the deposition or dissolution of a metal may not be the electrochemical reaction but the rate at which material can be transported to or from the electrode surface. Under these conditions, we say that the rate of the reaction is mass-transport controlled.

4.1 Mass Transport Processes

Transport of species to and from electrodes can be accomplished by 1) Convection

Hydrodynamic transport can be induced by

Stirring Pumping Air Sparging Gas Evolution Density Gradients Thermal Gradients Convection is generally the most effective means of enhancing mass transport and will be dealt with in more detail at a later stage. 2) Diffusion

A concentration gradient acts as the driving force for diffusion from regions of high to low concentration. 3) Migration

The contribution of migration(movement of ions under the influence of the electrical field) to the overall transport process is of importance only in processes such as mass transport to electrodes or in ion exchange materials or membranes. In theory, this can be calculated as we shall see later. In many processes, this is a rather small contribution (particularly in acidic solutions) and is therefore often neglected. 4.2 Diffusion Layer Model In this section, we will introduce a simple but very effective method for describing mass transport to or from an electrode surface. It uses the concept of a stationary liquid film adjacent

What will the actual potentials (relative to SHE) of the cathode and anode and the overall cell voltage be under these conditions?

16

to the surface. Mass transport through this film is considered to occur solely by diffusion (and possibly by migration in an electrical field in the case of electrochemical reactions). Beyond the film (or diffusion layer), mass transport is enhanced by means of convection involving relative movement of the surface/liquid interface by stirring or other convective processes.

Thus, Fig. 4.1 schematically depicts this model in which the concentration of a reacting species such as Cu2+ ions is plotted as a function of distance from the electrode surface. The diffusion layer of thickness is shown by the vertical line. It is assumed that the concentration is maintained by convection at its bulk value beyond the diffusion layer.

Figure 4.1 Diffusion layer model for mass transport.

There are two limiting cases for the rate of the reaction at the surface,

i) The rate is controlled by the rate of the electrochemical chemical reaction. In this case, mass transport of the reacting species to the surface is very much faster than the rate of the reaction at the surface. Under these conditions, the concentration of the reactant at the surface will be essentially the same as its concentration in the bulk of the solution i.e. its concentration at the surface is not perturbed by the fact that there is an electrochemical reaction. Under these conditions, Co Cb and the concentration profile is as shown by the dotted blue line.

ii) The rate of the reaction is controlled by mass transport of the reactant to the surface. In this case, the electrochemical reaction is very much faster than that of mass transport. Under these conditions, the reactant is consumed by electrode reaction as soon as it arrives at the surface. Under these conditions, Co 0 and the concentration profile is shown by the solid black line.

For a process involving diffusion, the flux (j) at the surface (or rate at which the reacting species arrives per unit area of surface) is given by Ficks First Law of Diffusion,

j = D (C/x)x=0 i.e. the rate of diffusion is proportional to the concentration gradient with D being the diffusion coefficient.

CChemical C b

Co

Diffusion

Distance, x

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In the above case, the concentration gradient at the surface is approximated by that over the diffusion layer as shown by the dotted red line. Thus,

j = D(Cb - C0 )/ The maximum flux that can be supported under these conditions will be achieved for C0 = 0,

jL = Cb D/ = Cb . kLwhere kL = D/ is often used and is known as the mass transfer coefficient.. The table below lists some typical values of the diffusion coefficients of common species in aqueous solution.

Table 4.1 Some Diffusion Coefficients Infinite Dilution at 25oC for ions

Species

Diffusion Coefficient m2 s-1 x 109

H+ 9.31 -OH 5.26

-Cl 2.03 SO42- 1.06 Na+ 1.33

2+Cu , Ni2+, Zn2+ 0.72 Fe2+ 0.71 Fe3+ 0.60

O2(aq) 2.10 O2(g) 23000

Note that most metal ions have diffusivities close to 5 x 10-6 2 cm s-1 while the proton and the hydroxide ion have unusually high values for the same reason that their conductivities are high. The effect of concentration on the diffusion coefficients is not as pronounced as in the case of the conductivity and is therefore often ignored.

The thickness of the diffusion layer (and, of course the mass transfer coefficient) is strongly dependent on the effectiveness of the convective processes in the bulk of the aqueous phase. Some typical values are shown in Table 4.2.

Table 4.2 Some Approximate Diffusion Layer Thicknesses

Surface (mm) kL ( m s-1 ) Freely suspended 100m particle of density 4g/cm3 0.01 50 x10-6Suspended gas bubble 0.02 25 x 10-6

Plane surface with longitudinal flow (25cm s-1) 0.1 5 x 10-6

Rotating cylinder (100 cm s-1 ) 0.04 12.5 x 10-6

Natural convection on vertical electrode 0.1 0.2 5-10 x 10-6

*Assumes D = 5 x 10-10 m2 s-1

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4.3 Mass Transport Correlations

We have seen that it is possible to estimate the rate of mass transport of a species to a surface if we know the relevant mass transfer coefficient. Unfortunately, as we could expect, the coefficient will depend strongly on the geometry and flow conditions close to the surface. Thus, unlike diffusion coefficients, tables of mass transfer coefficients do not exist. A number of useful semi-empirical correlations have been established which enable one to calculate the Sherwood Number, Sh (the dimensionless surface concentration gradient averaged over the mass transfer surface) for various surface geometries and operating conditions. These correlations use several so-called dimensionless numbers which are given in Table 4.3.

Table 4.3 Some Dimensionless Groups

Number(or Group) Physical significance Form

Sherwood number (Sh) Mass transfer/Diffusion kL.l/D

Schmidt number (Sc) Momentum diffusion/Mass diff. /D or /.D Reynolds number (Re) Inertial forces/Viscous forces l.U/ Grashof number (Gr) Buoyancy forces/Viscous forces l3g/(.2) Peclet number (Pe) Flow/Diffusion velocity U.l/D

The symbols and their dimensions are : D diffusion coefficient (L2/t) g acceleration due to gravity (L/t2) kL mass transfer coefficient (L/t) l characteristic length (L) U fluid velocity (L/t) kinematic viscosity (L2/t) / fractional density change viscosity (M/L.t) It should be remembered that the exact definition of each of these dimensionless groups implies a specific physical system. Thus, for example, the characteristic length l in the Sherwood number will be the height from the bottom of a planar electrode for mass transfer by natural convection but will be the diameter of a spherical particle which is dissolving.

Many of these correlations have the same general form. They involve the Sherwood number which contains the mass transfer coefficient. In the case of diffusion, the Sherwood number varies with the Schmidt number. The variation of the Sherwood number with flow is more complex because of the different physical origins of flow i.e. forced convection due to pumping or other means and natural convection due to density or thermal gradients. The former involves the Reynolds number while the latter is described by the Grashof number.

i) Natural Convection

Sh = 0,67 ( Sc. Gr)0,25 for 104 < Sc. Gr < 1013 (Laminar) Sh = 0,31 ( Sc. Gr)0,28 for 1013 < Sc. Gr < 1015 (Turbulent) Sh = kL.B/D Gr = g(o - s) B3 /( o 2) where o and s are bulk and surface solution densities. Sc = /( D) where = solution viscosity and = kinematic viscosity (/).

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ii) Gas Evolving Surfaces Sh = 0,93 Re0,5 0,49 Sc

where Re = Vg .d./A. and Vg is the volumetric gas evolution rate. d is the bubble break-off diameter ( 40m for H2 in alkaline solutions and O2 in acid solutions and 60m for H2 in acid and O2 in alkaline solutions.) In many cases, this can be simplified to

kL = m ign

where m=5-10. 10-6 A-1 m3.s-1 -2 for H2 evolution and iH in Am

iii) Forced Convection

For laminar flow as shown between two large planar surfaces (height>>gap)

S

Flow

B

L

1/3Sh = 1,85 { Re. Sc. de/L}

Sh = kL . de /D where de =2S

Re = U. de / and, for surfaces of finite height,

1/3 {2/(1+)}1/3Sh = 1,47{ Re. Sc. de/L}where = S/B For fully-developed turbulent flow,

0,8Sh = 0,023 Re Sc1/3

iv) Particulate Beds

For a packed bed of spheres with Re < 0,1,

Sh = 1,09/ . ( Re. Sc )1/3For a packed bed of spheres with Re > 0,1,

Sh = 1,44 Re0,58 Sc1/3

For a fluidized bed (< 25% expansion),

Sh = (1-)1/2. Re1/2 Sc1/3 / where, Re = U d / and Sc = /(.D) and U = superficial fluid velocity, = bed voidage and d = particle diameter

vi) Spinning Disk Sh = 0.62 Re1/2. Sc1/3

where Sh = kd/D, Re = d2/ , Sc = /D, d is the diameter of the disk and is the angular velocity of the disk.

This relationship is the well known Levich equation for mass transport to a rotating disk, an important characteristic of which is the local mass transfer coefficient is independent of the position on the disk.

20

4.4 Influence of Mass Transport on Electrochemical Kinetics

The current- potential curves discussed above were derived and depicted on the assumption that the rates of the reactions at the electrodes are controlled by the rates of the electrochemical processes. Thus, for example, the cathodic curve for the reduction of Cu2+ ions shown in Fig. 3.3, shows that the current increases exponentially with decreasing potential. However, the concentration of Cu2+ ions at the electrode surface will decrease due to reduction and at a sufficiently high overpotential, this concentration will decrease to effectively zero. Further increases in the overpotential cannot increase the rate of reduction which has now become controlled by mass transport and the curve will take on the shape as shown below in Fig. 4.3. The maximum current that can be passed is known as the limiting current density.

ic

E

Cu2+ + 2e = Cu

il

Figure 4.2 Current-potential curve for a cathodic reaction involving mass transport.

The maximum current that can be passed is known as the limiting current density. The equation for the current-potential curve including mass transport in terms of the limiting current density iL, can be shown to be of the following form for a cathodic reaction with a high exchange current density,

= -(RT/nF). ln{(iL i)/i} where = E E1/2 and E1/2 is the half-wave potential

4.5 Mass Transport of Ions

In addition to convection and diffusion, the mass transport of ions in aqueous solutions can be influenced by the presence of an electrical field and it is necessary to briefly review these additional modes of transport.

a) Conductivity

Consider two electrodes of cross-sectional area A placed a distance l apart in a solution of an electrolyte as shown,

21

l

A

l

As in the case of any electrical conductor, the resistance and its reciprocal, the conductance or conductivity, is given by

Resistance : R = l /A where = resistivity in ohm.m Conductivity : = 1 / (ohm-1 m-1 or S m-1 where S is the Siemen) Unlike metallic conductors, the conductivity of electrolytes is concentration dependent, and therefore we define,

Molar Conductivity = / c ( S m2 -1 mol ) Generally decreases with increasing concentration and, approximately,

= 0 - Ac1/2where 0 is Limiting Molar Conductivity at infinite dilution and A is a constant for a particular metal ion. These values can be found in tables.

Transport Number tj = Ij / Ii where Ij is the current carried by species i and Ii is the total current carried by all ions i. This enables one to calculate contributions of single ions to the overall conductivity and

tj = cj j / (ci i) where j is the molar conductivity of species i. The summation is carried out over all ions. The conductivity of any salt can be calculated from a combination of cation and anion values

= (cj 0j) for dilute solutions.

As shown in Table 4.4 there are published values of the conductivity of typical electrolytes as a function of composition and temperature. There are also empirical equations such as

Resistivity(ohm m) ={3200 + 7.3[Cu] + 4.5[Fe] + 1.3[As] + 9.6[Ni] - 5.6[H2SO4] 14.6T}/105

22

Table 4.4 Specific Conductivity ( ) of Copper Electrolytes [Cu] [Acid] Temp Resistivity Conductivityg/L g/L deg C ohm m 1/(ohm m) 30 150 45 0.0192 52.0 30 160 45 0.0187 53.6 30 170 45 0.0181 55.2 30 180 45 0.0175 57.0 35 150 45 0.0196 51.1 35 160 45 0.0190 52.6 35 170 45 0.0185 54.2 35 180 45 0.0179 55.9 40 150 45 0.0200 50.1 40 160 45 0.0194 51.6 40 170 45 0.0188 53.1 40 180 45 0.0183 54.7 35 150 40 0.0203 49.2 35 160 40 0.0198 50.6 35 170 40 0.0192 52.1 35 180 40 0.0186 53.7 35 150 50 0.0189 53.0 35 160 50 0.0183 54.7 35 170 50 0.0177 56.4 35 180 50 0.0172 58.2

In the electrowinning of copper, anodes and cathodes of dimensions about 1m x 1m are used and the anode-cathode

b) Combined Diffusion and Migration

The total flux/unit area of an ionic species to the electrode surface is given by the sum of that due to diffusion and migration,

J = D(C/x)x=0 + t.i/nF where t is the transference number of the ion in question.

The limiting current density is therefore given by

iL = nFDCb / (1-t) For example, in a 1 mol/dm3 solution of CuSO4

tCu = 53.6/(53.6 + 80) = 0.40

Thus, the limiting current will be some 1.67 times that for diffusion alone.

separation is 2cm. Calculate the resistance between the electrodes and the voltage drop due to the electrolyte at a current density of 300 A m-2 for an electrolyte containing 40gpl Cu as sulfate and 170 gpl H2SO4 at 45oC .

23

On the other hand, a typical copper electrorefining solution also contains about 1.7 mol/ dm3 H

Velocity Profiles at Half-Height

-10-505

1015202530

0 5 10 15 20 25 30Distance from Cathode, mm

Vel

ocity

, mm

/s

RefiningWinning

2SO4 , and in this case tCu = 0.19 and the limiting current will be only some 1.23 times that for diffusion alone. Why is acid added if this is detrimental to mass transport of Cu2+ ions to the cathode?

5 Mass Transfer at Vertical Electrodes In industrial cells with vertical electrodes of length up to 2 m and narrow cathode-anode gaps, mass transport to the cathode is an important quantity as it limits the current density which can be applied. Thus, most processes operate at current densities of 1/3 to 1/2 of the limiting current density. The flow of electrolyte due to circulation of the electrolyte in the gap between the electrodes is generally relatively low with values of 1 to 10 cm/min being used. Mass transport to the cathode is also important as it can determine the rate at which impurities can be co-deposited with the copper.

5.1 Natural Convection

In the absence of gas evolution at the cathode (such as in the electrowinning of copper), natural or free-convective flow is generally the most effective means whereby the electroactive ions are transported from the bulk of the electrolyte to the surface of the cathode. Natural convection is caused by density or temperature differences in the electrolyte which results in a concentration gradient from the cathode. For a vertical cathode, the electrolyte is partially depleted of relevant metal ion at the cathode surface. This results in a lower density of the electrolyte in the vicinity of the cathode which, in turn, induces a vertical flow of electrolyte close to the cathode as shown in Figure 5.1.

Obviously, in a refining cell in which a metal is being dissolved at the anode, the electrolyte becomes denser close to the anode surface and the flow generated by natural convection is in the opposite direction, i.e. downwards. In combination with the upward flow at the cathode, this sets up a circulating flow between the electrodes which can be quite effective in enhancing mass transport as shown in Figure 5.1.

Figure 5.1 Variation of flow as a function of distance from cathode in copper refining and winning cells.

24

For constant concentration in the bulk and at the electrode surface, i.e. constant current density, the following correlation has been found to describe mass transport due to natural convection induced by density differences,

Sh = C.(Sc.Gr)1/4

where C is a constant and the Sherwood number Sh = kL.L/D, the Schmidt number

Sc = /D and the Grashof number Gr = {g(

H+

Cu2

C

Distance from Cathode

o - e)L3} / (o2) = g{(Co Ce)L3} / 2

and kL is the mass transfer coefficient, L the length of the electrode, D the diffusion coefficient of the electroactive species, the kinematic viscosity of the electrolyte and o,e the density of the electrolyte in the bulk and at the cathode surface respectively.

is the so-called densification factor which relates the density of an electrolyte to the concentration.

For most electrolytes, Sc>103 and C = 0.67

i) For a simple binary electrolyte such as copper sulfate, the above can be rearranged to give an expression for the average mass transfer coefficient over the electrode,

kL = 1.19 {(Co Ce)}1/4 1/4 -1/4 -1/4 D L m/s Note that mass transport is a relatively weak function of all the relevant parameters and that the average value decreases as the length of the electrode increases. The difference Co Ce is obviously related to the current density at the electrode.

In the case of a mixture of electrolytes, such as CuSO4 and H2SO4, the situation is somewhat more complicated because of the presence of the acid. Thus, as we have previously seen, in solutions typical of copper refining or winning, most of the electrical current is carried by the protons in the electrolyte, i.e. the transference number t H ~ 1 and tCu ~ 0. Thus, protons migrate to the cathode but are not reduced (in the case of copper deposition) and they therefore accumulate in the vicinity of the cathode. At steady-state, the rate of migration to the electrode is balanced by a diffusion flux in the opposite direction. This increase in acidity at the cathode is shown is shown in Figure 5.2.

Figure 5.2 Concentration profiles near the cathode in a refining cell

25

This increase in acidity affects the density of the electrolyte at the cathode surface and therefore the density gradient. It tends to reduce the density gradient and thereby also the mass transport of copper ions to the cathode.

Under these conditions, the calculation of the mass transfer coefficient is more complex and only the final result will be shown, expressed as the limiting current density, i.e for Ce = 0,

il = 114835 z (1-tCu)-1 DCu

3/4 Co5/4 (/L)1/4 2 A/m

where z is the no of electrons involved and = - Cu H.{tH/(1 tCu)}. (DCu/DH)3/4where Cu is the value of the densification factor for CuSO4 and H the value for H2SO4. 5.2 Application to Copper Deposition

The above can be used to estimate the limiting current density during electrowinning or refining of copper from an electrolyte containing 40 g/l of copper and 150 g/l sulfuric acid in cells containing vertical electrodes of height 1m at 60oC . The electrolyte is normally circulated through the cells at such a slow flow rate that the influence of the forced flow on mass transport to the cathode can be neglected with the major component in mass transport being natural convection.

The relevant values of the required parameters are

DCu = 1.2 x 10-9 m2/s, DH = 4.3 x 10-9 m2/s at 60oC

= 6.4 x 10-7 m2/s tH = 0.6, t = 0.07 CuCu = 1.5 x 10-4 m3/mol for 630 mol/m3 of CuSO4H = 5.2 x 10-5 m3/mol for 1530 mol/m3 of H2SO4from which, = 1.37 x 10-4 m3/mol

2and il = 830 A/m

Operating plants generally use current densities between 250 and 350 A/m2 which is in the range of about 1/3 to 1/2 of the limiting current density.

In the absence of the supporting electrolyte, i.e. the acid in this case, one can calculate the limiting current density using the above equation with = and t = 0.4 to give Cu Cu

2il = 1000 A/m

Thus, although the use of acid in the electrolyte to increase the conductivity is desirable, it has a negative impact on the mass transport of copper ions to the cathode surface.

5.3 Effect of Gas Evolution

Mass transfer to the cathode can be considerably enhanced as a result of the flow generated by the rising oxygen bubbles from the anode. This effect is even more pronounced in the case of hydrogen evolution at the cathode which is applicable in zinc and some extent nickel but not copper electrowinning. This effect is shown schematically in Figure 5.3 and the consequences are that

26

i) mass transfer at the anode is greater than that at the cathode. This has consequences for the speciation of, for example, iron in the electrolyte as we shall see later.

ii) Mass transport is greater towards the top of both anode and cathode. iii) On the other hand, the increased volume of gas in the electrolyte towards the

top of the inter-electrode gap results in an increased resistivity of the electrolyte and a lower current density on the upper surfaces of the electrodes. This hold up as a percentage of the inter-electrode volume is shown in Figure 5.3 for increasing current densities and its effect on the current distribution in Figure 5.4.

+ -

k

1.0E-04

1.5E-04

2.0E-04

2.5E-04

3.0E-04

200 250 300 350 400Current density, A/m2

Mas

s Tra

nsfe

r C

oeff,

cm

/s

4567891011

Mass Transfer

Gas Holdup

Figure 5.3 Effect of gas evolution at the anode on mass transport

0.8

0.9

1

1.1

1.2

1.3

0 0.2 0.4 0.6 0.8 1Fraction of Height from Bottom

i/im

300A/m2

500A/m2

Figure 5.4 Effect of gas hold-up on the current distribution in a typical Cu electrowinning cell. The ratio i/im is that of the local current density to the mean value for the overall cathode.

The results of recent CFD (Computational Fluid Dynamic) simulation of a typical copper EW cell are shown in Figure 5.5. The strong vertical movement of electrolyte close to the anode

27

and the very much weaker flow at the cathode is apparent. It is not clear from the diagram, but the flow is upward at the cathode in the lower section and downward in the upper section. Also apparent is greater extent of oxygen gas holdup in the upper section of the cell although the absolute maximum value of about 3-4% is considerably less than those shown in Figure 5.3. This difference could be attributed to the assumptions made for bubble size in the simulations.

Fig. 5.5. CFD simulation of flow and oxygen gas holdup in a cross section of a copper EW cell anode is on right

One of the most effective methods of increasing mass transport in cells is to sparge air between the anode and cathode and the results of testwork aimed at quantifying the effect on the mass transport coefficient at the cathode using a tracer of silver ions that are deposited at the limiting current during the deposition of copper are shown in Figure 5.6.

28

Diffusion Layer

0102030405060708090

100

0 0.1 0.2 0.3Diff Layer Thickness, mm

Dis

tanc

e fr

om b

otto

m, c

m

Air, 5 l/minNo air

Agitation by Air Sparging

0

0.5

1

1.5

2

2.5

3

3.5

0.1 1 10Air Flow, l min-1 dm-2

103 k

L, cm

/s

Natural convection

Figure 5.6 Effect of air sparging on mass transport at the cathode during the electrowinning of copper.

Calculate the current density due to the mass-transport controlled reduction of ferric ions (5.6 g/L) on the top and bottom half of a copper cathode assuming that the mean

6 Electrocrystallization

The purity and properties such as electrical conductivity of metal deposits depend on their crystallographic structure, texture and surface morphology. These properties will depend on a number of operating parameters during the electrocrystallization process such as the current density (or, more importantly, the current as a fraction of the limiting current), metal ion concentration, pH, temperature, agitation, nature of the anions and cations present and the presence of additives such as smoothing agents.

The cathodic deposition of a hydrated metal ion M(H2O)n2+ onto a metal electrode M occurs in several steps

i) Charge transfer to the hydrated metal ion in the double layer with the formation of a neutral or partially charged ad-atom or ad-ion which may still be partially hydrated.

ii) Migration of the ad-atom by surface diffusion to a growth site (such as an edge or kink) at which further cordination to metal atoms in the lattice occurs.

iii) Growth of the crystal lattice by successive transfer and diffusion steps.

This sequence of steps, any one of which may be rate-determining, is shown schematically below in Figure 6.1.

thickness of the diffusion layer is 0.17mm on the top half and 0.25mm on the bottom half. Assume that the diffusion coefficient for ferric ions in the electrolyte is 1 x 10-9 m/s.

29

The essential difference between an aqueous metal ion and a metal ion in a lattice is that the co-ordination sphere of the former is composed of water molecules while that of the latter is composed of other metal ions i.e. metal - H2O bonds have to be broken and replaced by metal - metal bonds. As could be expected, the rate of this process is influenced by a number of factors such as the nature and concentration of the metal ion and the anions present, the presence of inhibitors and the nature of the substrate. In the case of electrocrystallization, an important additional parameter is the potential difference between the solution and the substrate.

30

A

Aqeo

Aqueous ion

i)

Ad-ion/atom

(ii)

(iii)

Figure 6.1 Steps in the electrocrystallisation process

6.1 Influence of Kinetics on Deposit Morphology

i) Charge Transfer

The exchange current density of metal-metal ion electrodes varies between wide limits depending on the nature of the metal ion and its state in solution. For simple ions, the following table gives some idea of the range of reversibility. Replacement of water molecules by other molecules or anions in the metal co-ordination sphere can either accelerate or inhibit charge transfer e.g. chloride tends to increase the i0 for the deposition of the iron-group

metals while cyanide can prevent deposition completely. These effects are complex combinations of thermodynamic and kinetic factors.

Table 6.1 Classification of Metal Deposition Reactions

i0 ( A cm2 ) 10 - 10-3 -310 - 10-8 -810 - 10-15

Lead(II) Copper(II) Iron(II) Tin(II) Zinc(II) Cobalt(II) Mercury(II) Manganese(II) Nickel(II) Cadmium(II) Chromium(II) Silver(I) Platinum(II) Gold(I)

ii) Crystallization

The formation and morphology of electro-deposits is dependent on several factors the most important of which is the relative rates of nucleation and growth. In crystallization in general, the rate of nucleation is a strong function of the degree of supersaturation of the solution. In the case of electrocrystallization, the degree of supersaturation is determined in addition by the overpotential which therefore determines the nature of the deposit as follows:

Unlike other crystallization processes, the additional variable of the overpotential determines the nature of the deposit as follows:

Rate of Nucleation Conc of Ad-atoms Overpotential SEM micrographs of copper nuclei formed on a platinum surface are shown on Figure 6.2. In this case, the overpotential for nucleation on a foreign surface exceeds that for growth of the nuclei and therefore the density of nuclei is not great. This has obvious implications in regard to the adhesion of metal deposits to foreign substrates.

A high nucleation rate will create a large number of small crystallites i.e. a fine-grained smooth deposit. This is generally observed and, in the extreme, powders are formed as the current density approaches its limiting value.

Figure 6.2 Nucleation of copper on stainless steel

31

The use of complexed metal ions, such as Ag(CN)2- also tends to promote the formation of dense fine-grained deposits. Furthermore, growth can be inhibited by the addition of various levelling or brightening agents which absorb preferentially on the crystal defects which are the growth sites. Higher periodic current densities produced in superimposed pulsed or A.C current also act to promote nucleation and produce finer deposits at the same mean current density.

iii) Lateral versus Outward Growth

Outward growth of a perfect crystal requires nucleation of new 2-dimensional layer. On the other hand, lateral growth can occur at edge and kink sites and is generally preferred. Few crystals are perfect and the presence of defects such as screw dislocations enables lateral growth to be translated into an outward growth.

As we have seen, the electrical field is concentrated at a tip of a growing crystal and results in increased deposition in this region. Metals with high exchange current densities, i.e. large changes in rate for small changes in overpotential, will be most susceptible to this phenomenon.

The metals in the Periodic Table can be roughly sub-divided into various groups depending on the thermodynamic and kinetic characteristics of metal deposition as shown in Table 6.1. The most important properties of the two outermost groups can be summarized as shown in Table 6.2. The normal (or soft) metals generally have low melting points and are not highly stressed when deposited. The opposite is true of the inert (or hard) metals as shown by the data in Table 6.3. Note the high stress associated with the iron-group metals, particularly Co and Cr.

Table 6.2 Characteristics of Metal Deposition Reactions

Characteristic Normal Inert

Reversibility High Low

Overpotential Low High

Grain Size Large Small

Current Efficiency High(~100%) Low(30-80%)

H2 evolution rate Slow Fast

Stress Low High

Hardness Soft Hard

6.2 Application to Metal Deposits

The physical and , to a lesser extent, the chemical properties of cathodes are determined by a number of factors, on of the most important of which is the current density. The effect of current density on the morphology is shown in a general case in Figure 6.3. As the current density (or more accurately the ratio i/iL) increases, the mean size of the crystallites making up the deposit decreases with well formed large crystals giving way eventually to very fine crystals or powdery deposits. This is expected in terms of the effect of overpotential on the relative rates of nucleation and growth.

32

-E

i/iL

M2+ + 2e = M

2H+ + 2e = H2

Nodules, dendrites Polycrystalline Bunched layers, ridges Spirals, blocks

Course powders

Fine powders with hyroxides

1.0

Increasing rate of

nucleation

Figure 6.3 Effect of current density on the nature of deposits.

Other factors also contribute to the size of the crystallites and the morphology as shown qualitatively in Figure 6.4.

Increasing

Metal ion Conc. Current Density Temperature Mass Transfer Smoothing Agents

Figure 6.4 Effects of various parameters on the size of crystallites

The maximum current density that can be applied in industrial operations is normally about 30-40% of the limiting current density. In the case of base metal electrowinning, a rule-of-thumb is that that the ratio

Current density (A/m2)/Metal concentration (g/L) in spent < 10

Thus, for a cell concentration of 40 g/L Cu, the maximum current density should be about 400 A/m2.

The simplified Winand diagram (Figure 6.5) shows "stability domains" for the main types of electrodeposits according to a system proposed by Winand. The main types identified are

i) the field-oriented isolated type (FI) that consists of dendrites, whiskers and powders,

33

ii) the basis-oriented reproduction (BR) type typified by a dense deposit with large diameter crystals oriented parallel to the substrate and a high surface roughness,

iii) the field-oriented texture (FT) type that consists of columnar crystals with increased density and a low surface roughness,

iv) the unoriented dispersion (UD) type that consists of randomly oriented small crystals with a low surface roughness.

This diagram follows from observations that the inhibition intensity and/or increased deposition rate lead to changes from one type to another, with FI and BR at lower inhibition and FT and UD at high inhibition.

Figure 6.5 The Simplified Winand Diagram

In the context of this diagram, an inhibitor is any species that can adsorb on the metal surface thereby reducing the rate of electron transfer and thus nucleation and/or the rate of growth of the crystallites. In some cases, adsorption is specific to particular crystal faces leading to preferential growth in one crystal orientation. The vertical axis can also be viewed as decreasing exchange current density. Thus, for example, deposition of Ag from uncomplexed solutions such as nitrate has a very high io and could be described by a horizontal line near the top of the diagram(whiskers) whereas deposition from cyanide solutions involves a low io (due to inhibition by strongly adsorbed cyanide) and could be described by a line near the bottom of the diagram (smooth microcrystalline).

Thus, in the case of copper, sulfate is a weak inhibitor and io is reasonably high so that deposits obtained from pure sulfate solutions are of BR or FT types and eventually become the undesirable FI type. Organic additives are added to increase the inhibition intensity and, in electrowinning/refining plants, the desired FT/UD type can be produced. Winand has pointed out limitations in the diagram, particularly in relation to organic inhibitors.

34

7 Current distribution in a cell he distribution of current in electrochemical cells affects the distribution of the extent of reaction ( or product)

T

To individual cells and electrodes in a multi- cell, multi-electrode operation.

To various macroscopic areas on individual electrodes.

To various microscopic areas on the surface of individual electrodes.

In general, the current density distribution over electrode surfaces is not uniform and can depend on a number of factors such as size and geometry of the cell, current density, conductivity of the solution, kinetics of the electrode reactions and mass transport to the electrode.

7.1 Types of Current Distribution

There are three different types of current distribution which can be distinguished as summarized in Table 7.1.

Table 7.1 Types of Current Distribution

Type of current Assumptions Parameters distribution

Primary absence of overvoltages geometry of cell and electrodes

activation overvoltage without concentration variation near electrode

geometry Secondary activation overvoltage conductivity of electrolyte/ electrodesgeometry activation and concentration

overvoltages Tertiary activation conductivity

Consider the simple so-called Haring and Blum cell below consisting of two cathodes at distances Lf and Ln from a central anode.

Figure 7.1 Haring Blum Cell If the same voltage is applied across each electrode pair, then the currents flowing through the cells will be in = V/Rn and if = V/Rf

where Rf and Rn are the total resistances of the electrode pairs.

35

The total resistance (Rf ) to current flow at e.g. the far electrode pair,

R

- +

f = Rc + Rs + Rp where Rc is due to contacts, bus-bars etc

Rs is the resistance of the solution (anode to cathode)

Rp is the sum of the polarization resistances at each electrode

NB Rp is the sum of the reciprocals of the slopes of the current/potential curves for cathode and anode of the far electrode pair.

It should also be noted that in the above simple case, the current distribution is uniform over the surface of each cathode. This follows from the obvious observation that the potential gradient (E/x) is zero in all directions except that normal to the plane of the cathodes. The potential (Ex) at any point in the electrolyte at a distance x from one of the cathodes is simply given by

E = E + x(E - Ex c a c )/d

where E and Ea c are the anode and cathode potentials and d is the distance between them. This relationship is the solution (for a 1-dimensional case) of the more general Laplace equation

E / x + E / y + E / z2 2 2 2 2 2 = 0 with boundary conditions E = E at x = 0 and E = E at x = d. c a Another reactor which has a uniform current distribution is the so-called concentric cylinder cell which is typical of, for example the EMEW cell. Even without mathematical examination of this system, it is evident that the current distribution is uniform and that the equipotential surface will be cylindrical as shown. Unfortunately, very few other types of geometry give uniform current distribution and the full Laplace equation has to be solved with the relevant boundary conditions. 7.2 Primary Current Distribution

Where the electrode reactions have no overpotential, the only contribution to the total resistance, R, of an electrode pair is the resistance of the solution, therefore

and the current density distribution in the system is controlled by the geometry of the electrodes with those areas less readily accessible having a lower current density. This is the

distribution and is a simple first approximation. primary current

Quite often the labour involved in establishing a primary current distribution is not worthwhile because, as we shall see, it is seldom valid in real systems. Thus, for the relatively simple parallel plate reactor in which the electrodes do not fully occupy the cross- sectional area of the reactor, as shown, it can be shown that the ratio of the local

36

current density(i) at a point x from the edge of one electrode to the average current density (iav) is given by the rather complicated expression

( )21

22

2

lLx2

mk/

)/).(sinh(sinh

).(tanh/cosh.

i/i

0.60.8

11.21.41.61.8

22.2

-0.5 0 0.5x/L

i/iA

v

2.42.62.8

L/h=1, Wa=0

L/h=50, Wa=0

L/h=50, Wa=10

av =

where = IIL/2h and k(m) is the complete elliptical integral. The primary current distribution for various values of L/h and the Wagner number(Wa) (see below) are shown in Figure 7.2. Note that for Wa 0, the interelectrode spacing is very large.

Figure 7.2 Effect of edges on primary current distribution Apart from the infinite current densities predicted at x=0 and x = L, the current density is reasonably uniform over a substantial part of the electrode. The large current densities at the edges can at least be partly eliminated in real parallel plate cells by the use of edge strips of a non-conducting material. Another interesting case which will be discussed in more detail at a later stage is that of a three-dimensional particulate electrode such as a packed - or fluidized-bed. In terms of the primary current distribution only, the bed can be treated as a homogeneous combination of a zero- resistance conducting solid phase in an electrolyte whose resistivity is considerably higher (by a factor 1/ where is the voidage of the bed) than that of bulk electrolyte. In this case, the primary current distribution is simply a linear function of the depth into the bed. 7.3 Secondary Current Distribution

It has been shown that the primary current distribution predicts a ratio of 5 for the metal distribution in the simple Haring and Blum cell for Lf=5Ln. Actual experiments with an electrolyte typical of copper electrowinning practice reveal that the current density ratio is

Lh

x

37

less than 5 it is about 3.5. The question arises: what has made the current distribution more uniform than dictated by geometry alone? The answer is not difficult to see. It will be recalled that an electrochemical cell consists of three components viz.: the two electrode/electrolyte interfaces and the electrolyte and that each of these regions offer resistance to the flow of current. The primary current distribution takes only the electrolyte resistance into account. To take into account the resistance associated with the electrode/electrolyte interfaces, one has to derive it from the current/potential relationships, i.e. the Butler-Volmer equation. In theory, this can then be combined into an overall model by way of the Laplace equation. However, as indicated previously, the solution to these equations is difficult even for relatively simple geometries. It is instructive, however, to derive the relevant relationships for the Haring and Blum cell under two limiting conditions. (i) For electrode processes which operate close to equilibrium where the so-called low-field approximation to the Butler-Volmer equation is valid, i.e.

i = nFi / RT 0i - current density (A/m2) n - number of electrons F - Faraday's constant (96500 C mol-1) i0 - exchange current density (A/m2) - overpotential (V) This can be rearranged to

/i = RT/nFio = Rp giving the polarisation or charge-transfer resistance, R 2p (ohm m )

Thus, the above equation for the HB cell becomes

n

f

np

fp

np

fp

f

n

W/11W/11

LR

LR

RRRR

ii

++=

++=+

+=

where is the specific conductivity of the solution (S m-1) and

+=

cocaoap ininF

RTR,,

11

is the polarisation or charge-transfer resistance, Rp (ohm m2 ) of the cell. It is the sum of the resistances of the electrode reactions at the anode and the cathode. Note that it is the reciprocal of the slope of the current/potential curve (/i = RT/nFio = Rp) The second form is obtained by making the substitution Lf/ = Rf which is the resistance between two parallel electrodes of 1m2 area and Lf m apart and is the specific conductivity of the solution (S m-1)

Note that the ratio Rp/Rf =Wf is a dimensionless quantity called the Wagner Number.

Since all quantities are positive, in / if < Lf / Ln and therefore the current distribution is more uniform than in the previous case. Also, for Rp >> Lf , in / if 1 i.e. electrode reactions with low exchange current densities (io) and/or solutions with high conductivity favour more uniform current distribution. It should also be noted that, in this case of low overpotential, the current

38

distribution is independent of the current density. The effect of the exchange current density is shown by the curves in Figure 7.3. Thus, by complexing the copper ions with, for example, cyanide, the exchange current density is reduced by several orders of magnitude resulting in a very even current distribution in a cell designed to exaggerate the difference between point A and point B on the cathode. The same effect albeit at a reduced effect is observed on addition of a smoothing agent such as gelatine which adsorbs on the surface and lowers the exchange current density for deposition. Thus, the excvhange current density for reduction of copper ions is about 5x lower in the presence of 3ppm gelatine. This property is called throwing power by the electroplating industry.

Figure 7.3 Effect of exchange current density on current distribution. (ii) For electrode processes which operate far from equilibrium where the high-field approximation to the Butler-Volmer equation is valid, it will be recalled that, in this case, the current density increases exponentially with increasing overpotential and the value of the change-transfer resistance, i.e. the slope of the potential/current curve is not constant but decreases with increasing current density. Thus, although the addition of the charge-transfer resistance will improve the current distribution over that based on the primary effect, the inverse dependence of the change-transfer resistance on the current density results in decreasing uniformity of deposition with increasing current density. Thus for a Haring and Blum cell operating within the region where Tafel's law applies, it can be shown that

f

nf

n

f

fn

f

n

n

f

f

n

iilnW

LL

iLiilnb

LL

ii =

=

where b = ba + bc is the sum of the Tafel slopes for the anode and cathode reactions (based on natural logs!), and W = b/(Lnif), is a modified Wagner number. The effects of unequal cathode-anode spacing can be seen in the so-called Hull cell arrangement shown in Figure 7.4. This would be typically expected with a non-vertical cathode

39

with a higher (than average) current density on the one side of the blank and a lower current on the other at the same depth. The equalizing effect of operating in the high field region is apparent.

Primary3 i/i Secondaryav x=0

Figure 7.4 Current distribution in a Hull Cell It is worth emphasizing that the primary current distribution (dotted line) predicts a less uniform distribution than that which takes into account the secondary current distribution (calculated assuming the high-field approximation).

7.4 Tertiary Current Distribution

Previously it has been shown that, as the overpotential increases at an electrode, the current density increases until, in the limit, it reaches a plateau value determined by the rate of mass-transport of the electroactive species to the electrode. In principle, therefore, such a situation in a Haring and Blum cell would result in equal current densities at both cathodes, i.e. under mass-transport controlled conditions, the current distribution is uniform. Another way of looking at this is to note that the slope of the current/potential curve tends to zero under these conditions, i.e. Rp . The various regions of such a curve are shown in Figure 7.5

Figure 7.5 Overall current/potential curve for cathodic reduction Note that the value of the polarization or charge-transfer resistance(Rp) is the reciprocal of the slopes of the lines as drawn.

Mass transport

High Field

Low Field

i

E

1/Rp

x/d 1

2

1

d

x=d

40

A tertiary level of current distribution is also associated with the microdistribution of current due to imperfections on the surface of the electrode. Consider a microscopic cross-section of an electrode surface....

iR

Figure 7.6 Schematic of tertiary current distribution Three factors tend to influence the current distribution in the vicinity of the surface protrusion.

a) The current density is higher at a point or edge than on a plane - as depicted by the field lines.

b) The increased iR drop between the plane and the protrusion will enhance reaction on the latter.

c) The rate of diffusion to the tip is greater than that to the plane due to the shorter diffusion path and, even under mass-transport conditions, reaction will be faster at the tip.

These observations explain why metal deposits tend to become rougher with time of deposition and, in extreme cases such as silver deposition from nitrate solutions (very high i0) dendrites can be seen to grow visibly. The opposite effect is put to good use in electropolishing which preferentially removes the protrusions by anodic dissolution.

Combining all of these models and observations leads to the conclusions that the current distribution is more uniform...

the greater the slope of the polarisation curve (i.e. a high polarisation resistance) the larger the conductivity of the solution the smaller the distance between anode and cathode Applying this to industrial situations...

the current distribution depends on the composition of the electrolyte and adding supporting electrolyte (increasing the solution conductivity) or of additives which inhibit the desired reaction (increasing the polarisation resistance) lead to a more uniform current density

many industrial processes operate under activation control and the Tafel approximation applies, under this condition the higher the mean current density the less uniform the current density

since all three current distributions are dependent upon the geometry of the system, scale-up from laboratory to industrial (or even pilot) scale needs to account for the changes in current distribution due to larger electrodes and different cell designs

41

7.5 Current Distribution in 3-Dimensional Electrodes Within the general category of three dimensional electrodes we can include those which are genuinely porous such as a sintered or pasted electrode (e.g. a lead-acid battery electrode) and those which are composed of discrete particles with no mechanical adhesions. These may be in the form of a packed bed or the particles may be fluidized by flow of the electrolyte. There are various configurations of current and electrolyte flow possible and we shall consider only the case depicted below, viz. parallel flow of solution and current with the current feeder to the three-dimensional cathode and the anode arranged as shown in Figure 7.5.

Figure 7.5 Potential distribution in a 3-dimensional cell. As already mentioned, the ohmic drop in the electrolyte within the interstices of the bed will result in a current distribution which decreases with penetration of the bed (i.e. as x increases). Another way of looking at this is to note that the potential of the solution s will decrease with x as shown. If the bed material is highly conductive (e.g. metallic or graphite), the potential of the metal phase ( m ) is constant. The difference s - m is the overpotential and it can be seen that, at some value of x = heff , it will be zero. Thus the bed will be active for a depth of heff only. If s - m is great enough to produce mass-transport controlled reaction over part of the bed, then the current (and metal) distribution will take the form shown below Theoretical models for such a system have been developed which, as will by now be appreciated, are exceedingly complex. Experimental investigations of the scale-up of the height of the bed have confirmed the existence of the so-called effective bed height and a semi-empirical but useful relationship is heff = 0.4 ( /i)

xheff

i or metal

masstrans

s

m

x heff

packed

fluidized

+

x _

flow

42

where is the specific conductivity of the solution phase (ohm-1 cm-1 ) and (/i) is the slope of the polarisation curve measured at a bed height h < heff . Notice that this relationship can be written as Weff = 1/0.4 = 2.5 where Weff is the effective Wagner number. This restriction on the effective height of the bed can be relaxed somewhat if the bed is fluidized. In this case, m is not constant but also decreases as shown in the figure, i.e. the resistance of the solid phase is no longer very much smaller than that of the solution phase. Under these conditions the difference - s m decreases less rapidly with distance x and the effective bed height therefore increases. To-date there are no satisfactory theoretical treatments of this system

8 Energy consumption This is often the most significant operating cost in the recovery of metals by electrolytic processes.

To examine the energy efficiency of such processes, it is useful to breakdown the total cell current and voltage into their individual components as follows:

Voltage Components

Equilibrium Cell Voltage (Ee(anode) - Ee(cathode)) VeCathodic overpotential cAnodic Overpotential aOhmic potential drop in electrolyte VsOhmic potential drop in cell hardware VhVoltage loss in power supply Vp Current Components

Stoichiometric current nF Current in-efficiency due to side-reactions CIeCurrent in-efficiency due to shorts CIsStray currents in tankhouse SC

The energy components can be obtained from the product of the corresponding voltage and current components. If the current (or, more accurately, the charge/mol of metal deposited) is plotted against the voltage as shown in Figure 8.1, the resulting areas of the rectangles are proportional to the components of the total energy requirement as shown below for the electrowinning of copper,

43

0 1 2 3 4

1

2

F/mol

Figure 8.1 Components in the overall energy consumption.

Thus for the electrowinning of copper,

Thermodynamic Energy Requirement = (1.23 0.34)V . 2 . 96500/63.5

= 2690 Ws/g

= 0.75 kWh/kg

Actual Energy Requirement = 1.97 kWh/kg

( Cell voltage of 2.1V and 90% current efficiency.)

i.e. energy efficiency is about 30%.

9 he materials involved in the construction and operation of an electrolytic cell are obviously of critical importance. Clearly, materials which corrode or are otherwise incompatible with the electrolyte cannot be used as they would require frequent replacement. However,

Materials for cells and electrodes

The electrowinning of copper is carried out at a current density of 300 A/m2 with a cell voltage of 2.1V at a current efficiency of 90%. It was found that the current efficiency increased to 94% if the current density was increased to 350 A/m2 but the cell voltage increased to 2.4V. You are asked to assess with justification which option requires the least electrical energy/unit of copper produced.

T

V

c aVa + Vc Vs Vh+ VP

CIe + CIsstray current, (SC)

44

electrorefining relies on the corrosion and dissolution of an impure anode for subsequent plating onto a pure cathode. Any material used in an electrochemical cell should have a range of properties depending upon it function within the cell.

9.1 Anodes and cathodes

Table 9.1 shows which materials are used as anodes and cathodes for a range of industrial electrolytic processes, it is worth noting that sometimes the metal can be recovered electrolytically using oxidation (e.g. to MnO2) rather than the more common reduction to metal.

Table 9.1 Materials for anodes and cathodes

Starter sheet Metal Electrowin blank Anode

()-electrode for starter sheet

Cu (win) none, stainless Cu (stainless, Ti) Pb Ca - Sn Cu (refine) stainless Cu (stainless, Ti) impure Cu Zn Al none Pb-0.5%Ag(Ca) Co stainless none Pb - 0.5% Sb Ni none Ni (Ti) Pb - Sb Mn stainless, Ti none Pb MnO graphite, titanium none stainless (cathode) 2Au (win from CN- steel wool none stainless Au (refine from Cl- none Au (rolled) impure Au

In general, the materials used for anodes and cathodes are chemically resistant to the solution, except in the case of electrorefining where the anode is made of impure metal which is deliberately dissolved. In cases where the anode does not dissolve, it is important to minimise the overpotential of the reaction occurring at the anode (typically oxygen evolution)to reduce the cell voltage and minimise power. The choice of electrode material in this circumstance depends on the composition of the solution due to the dependence of oxygen evolution overpotential on pH, electrolytes present etc. Lead is used in most systems involving sulfate electrolytes as it is chemically inert and other elements are alloyed with the lead to strengthen the anode and reduce the rate of corrosion.

A number of the processes use starter sheets, these are simply thin sheets of the metal to be won. The sheets are made on inert electrodes in a separate plating bath and after they achieve the desired thickness (around 3mm) are stripped. Most modern copper tankhouses now make use of so-called permanent cathode technology in which copper is plated onto stainless steel sheets or blanks from which it is stripped as shown in Figure 9.1.

45

Fig. 9.1 Stainless steel cathode blank and automatic stripping of copper cathodes

The starter sheets are typically larger than the anodes to minimize edge effects on the cathode and plastic edge strips are fitted to the vertical edges and sometimes to the bottom of the blank. In the case of the blank shown in Figure 9.1, a thin section of wax is applied to the bottom of the blank to prevent copper deposition around the bottom of the sheet and to facilitate stripping which is generally carried out using mechanical stripping machines.

The stainless steel used is generally 316L which has a reasonably good corrosion resistance, particularly to pitting corrosion in the presence of chloride ions which is added to the electrolyte as a grain refining agent. The stainless steel plate is bonded to a solid copper or copper-plated steel hangar bar.

The complexity of a tankhouse flowsheet can be assessed by the schematic of a copper refinery shown in Figure 9.2.

46

ANODECASTING

STARTERSHEET

MACHINE

SHEETSTRIPPING

ANODE FURNACE

WASHING SPACING

TANKHOUSE

STRIPPERSECTION

BLANKSANODE SCRAP

BLISTERCOPPER

RECTIFIERCATHODEWASHING

REFINEDCATHODESTO MELTING

ELECTROLYTETREATMENT

STEAM

ACIDPLANT

LIBERATORCELLS

NiSO4

SLIMETREATMENT

PRECIOUSMETALS

Se and Te

ANODECASTING

STARTERSHEET

MACHINE

SHEETSTRIPPING

ANODE FURNACE

WASHING SPACINGWASHING SPACING

TANKHOUSE

STRIPPERSECTION

BLANKSANODE SCRAP

BLISTERCOPPER

RECTIFIERRECTIFIERCATHODEWASHING

REFINEDCATHODESTO MELTING

CATHODEWASHINGCATHODEWASHING

REFINEDCATHODESTO MELTING

ELECTROLYTETREATMENT

STEAM

ELECTROLYTETREATMENT

ELECTROLYTETREATMENT

STEAM

ACIDPLANT

LIBERATORCELLS

NiSO4

ACIDPLANT

LIBERATORCELLS

ACIDPLANT

LIBERATORCELLS

NiSO4

SLIMETREATMENT

PRECIOUSMETALS

Se and Te

SLIMETREATMENT

SLIMETREATMENT

PRECIOUSMETALS

Se and Te

Figure 9.2 Flowsheet of a typical copper electro-refinery

10 he basic design of cells hasn't really changed much since the first application of electricity to metal recovery and still consists of what is little more than a rectangular box. Depth is larger than width as it is simpler to increase the cell height than width to give a larger

working area. The greater depth is also required to avoid electrical short circuiting due to the build-up of anode slime at the bottom, a small clearance results in more frequent cleaning of the cells and loss of production. The length of the cell is dictated by the number of anode / cathode pairs to be accommodated and their spacing. The trend is to decrease the interelectrode spacing thus giving more pairs in a fixed cell length, this necessitates the installation of higher current capacity systems to maintain the required current density. A larger electrode area also requires a larger volume of pregnant solution to be pumped through each cell to maintain efficiency. Even distribution of fresh solution is also important to ensure that the solution is not depleted of metal in some parts of the cell. For this reason, the solution is not usually completely depleted of metal but retains a moderate concentration by only plating part of the content on each pass through the cell.

Cell design

T

The electrolyte level is maintained by adjustable weirs at either end of the cell, this system allows for differences in the heights of the cell and selective removal of the spent solution from one end or the other. The electrolyte can be fed at one end or via distribution pipes (typically a complex plumbing system) which evenly feeds the cell with fresh solution.

Most large plants use overhead cranes to install and remove the electrode pairs. The picture shows a crane picking up pre-spaced electrodes for installation in a cell.

47

Figure 10.1 Typical modern cell construction and cathode handling equipment

The cells themselves are typically made from reinforced concrete with a chemically resistant lining, such as lead, welded PVC, epoxy resin, rubber, neoprene, silica-loaded-asphalt or, more recently, a sprayed polymer layer.

11 Tankhouse current distribution An image of a typical copper tankhouse is shown in Figure 11.1. A large number of cells can be seen on both sides of the walkway with each consisting of a large number of anodes and cathodes.

Figure 11.1 Typical copper electrowinning tankhouse

48

As the cost of providing DC power and the size and cost of the distribution system (bus-bars) is determined largely by the magnitude of the current, modern tankhouses have adopted the procedure of operating the electrodes in each cell (22-60 cathodes and anodes) in parallel but the cells in series as shown schematically in Figure 11.2 which also shows the electrolyte distribution to the cells.

Figure 11.2 Current and electrolyte distribution A typical electrical contact arrangement in which cathodes of a cell are connected to the anodes of the adjacent cell by the use of a triangular contact bar is shown in Figure 11.3

In this arrangement, the anodes of one cell are connected to the cathodes of the next and so on. Thus, for N cells in series,

Total Voltage Vt = N x Vc

where Vc is the voltage/cell.

Spent manifold

PLS manifold

Anodes Cathodes

49

Figure 11.3 Typical arrangement of contacts between adjacent cells

If It is the total current requirement i.e. if all the electrodes were operated in parallel across the power supply (V

1234567891011121314151617181920212223242526272829303132

373839404142434445464748495051525354555657585960

33343536

t=Vc) , and Ic is the current requirement with N cells in series, then

Ic = It /N

i.e. the same total power W = Vc .It = Vt .Ic but at a reduced current which reduces the cost of the power supply and the size and cost of the bus-bars.

Short circuiting of current between electrodes that are touching is a loss of current efficiency. Sometimes this is the result of improper positioning of the electrodes in the cell, warped anodes and / or starter sheets, or nodular or dendritic growth on the cathodes. Various techniques such as the use of hand-held gaussmeters, infra-red sensors mounted on the overhead cranes and, more recently, continuous individual cell voltage measurements can be used to detect poorly operating cells and electrodes.

A well designed and operated tankhouse will have a minimum number of electrodes whose current den