The Chemistry of Copper Heap Leaching

John DreierNovember, 1999

Introduction

Between 1980 and 2000, primary copper produced by acid sulfate leaching

has grown from about seven percent of world output to 20 percent. Driving the increase in leach production has been the widespread adoption of solvent extraction and electrowinning (SX/EW), an efficient, low-cost technology for recovering copper from leach solutions. While a few stand-alone leaching operations have been commissioned since the early 1980’s, notably, El Abra, Cerro Verde, Quebrada Blanca, El Tesoro, and Carolina-Michilla, all in Chile, most new leaching capacity has been installed at existing mines. Outside of the existing mines is a large, undeveloped oxide copper resource that is technologically amenable to acid sulfate leaching. The purpose of this chapter is to explain the geochemical principles that govern acid sulfate copper leaching and to provide a basis for understanding the testing application of acid leaching to the undeveloped oxide copper resource base.

The leaching of copper in commercial operations is similar in principal (virtually identical in most respects) to natural processes including the formation of leached cappings above copper deposits, the weathering of rocks, the generation of acid and metals by the oxidation of mine dumps (ARD), and the formation and leaching of soils. Given its importance in so many fields, natural leaching has received a great deal of attention from economic geologists, geochemists, soil scientists and others for more than a century. The approach taken by this chapter is to apply to commercial copper leaching the theories and methods used to quantitatively understand and predict the outcomes of natural leaching.

The principal message of this chapter is that leaching kinetics and acid consumption depend on the ore and gangue mineralogy. To summarize for copper minerals: (1) oxide, hydroxide, sulfate, and silicate minerals leach quickly, (2) Reduced copper oxides such as delafossite and cuprite, native copper, and copper sulfide minerals of the chalcocite series leach slower than copper oxides, sulfates, and silicates, and (3) chalcopyrite and bornite leach very slowly. Gangue composed of quartz, sericite and other non-reactive minerals (minerals that do not react with the leach solution) allows fast leach kinetics and low acid consumption. Economic recovery of copper from non-reactive gangue often requires little or no crushing or curing. By contrast, gangue composed of calcite and other reactive minerals (minerals that dissolve rapidly in acid) is typified by low leach kinetics and high acid consumption. A worst case example for reactive gangue is limestone. The commercial leaching of copper from reactive gangue requires crushing and acid curing if, indeed, leaching is feasible.

Copper leaching coupled with SX/EW consists of the following steps:

1. Application of a sulfuric acid solution to the top of a heap.2. Interaction of the leach solution with the ore and gangue minerals within the heap.3. Soichiometric exchange of one copper ion for two hydrogen ions in the SX plant.4. Recirculation of the copper depleted solution back onto the heap.

The preceding four points state that oxide copper leaching with SX/EW takes place within a system that is closed to all elements save copper and sulfur. Sulfur is added as fresh leach solution. Copper is removed from the system in the SX/EW plant. Sulfide dump or heap leaching, where all acid is derived from the dissolution of pyrite and other sulfides, is closed to all elements save copper. For all other elements and compounds, copper heap leaching is a closed system. The characterization of copper leaching as a closed system has profound implications to the chemistry of leaching.

The Mineralogy of Copper Deposits

Virtually all copper deposits amenable to acid sulfate leaching were formed at or very near the Earth’s surface by the weathering and oxidation of primary copper deposits. Most primary copper deposits form deeper in the crust under conditions of elevated temperature and pressure. Copper deposits formed by the oxidation and leaching of primary deposits are termed secondary or supergene copper deposits. Some secondary copper deposits result from the in situ oxidation of primary deposits while others were formed by the gradual accumulation of copper that was leached from its original location and transported to the new site - in some cases for distances of tens of kilometers - by acid-oxidizing ground waters.

In nature, copper combines with nearly every element in the periodic table to form minerals as carbonates, silicates, hydroxides, oxides, chlorides, sulfates, phosphates, and sulfides. While natural copper minerals number in the hundreds, if not thousands, only a few are economically important outside of a mineral collection. In a given copper deposit, substantially all of the copper may occur in only two or three minerals. The most economically important copper minerals are listed below.

Economically Important Copper Minerals

chrysocolla CuSiO3.H2O

neoticite (Cu,Fe,Mn)SiO3.H2O

tenorite CuO

pitch limonite (Fe,Cu)O2

delafossite FeCuO2

native copper Cu

cuprite Cu2O

atacamite Cu2(OH)3Cl

chalcanthite CuSO4.5H2O

antlerite Cu3SO4.(OH) 4

brochantite Cu4SO4.(OH)6

malachite Cu2CO3.(OH)2

azurite Cu3 (CO3)2.(OH)2

chalcocite Cu2S

covellite CuS

digenite Cu1.95-xS

djurleite Cu1.95-xS

bornite Cu5FeS4

chalcopyrite CuFeS2

The small number of important secondary copper minerals results from the following:

1. An overall uniformity in the chemical composition of rocks that occur in porphyry copper deposits as shown by published tables of whole rock chemistry and mineralogy for various porphyry copper deposits (Titley, 1978,1982). The whole rock chemical data in these publications suggests that weathering of primary copper deposits would tend to produce copper silicates, sulfates, and carbonates in combination with iron, and manganese.

2. The prevalence of oxygen and carbon dioxide in Earth’s atmosphere and hydrosphere suggests that oxides, hydroxides, and carbonates of copper should be common.

3. The presence of sulfur in minerals such as pyrite, chalcopyrite, and bornite in primary copper deposits suggests that copper sulfates might be common in secondary deposits.

4. That sulfuric acid is released during the weathering and oxidation of primary copper deposits suggests that common secondary copper minerals might be stable in ground water of weak to moderate acidity.

5. The thermodynamic properties of copper minerals suggests that oxidized porphyry copper deposits would tend to be rich in chrysocolla, tenorite, copper-manganese silicates, cuprite, delafossite, and various members of the chalcocite series. These copper minerals are stable relative to other secondary copper minerals because their structures permit widespread substitution of other cations for copper.

6. The kinetics of mineral precipitation suggests that copper-iron-manganese hydroxides, oxides and silicates would be common.

Eh-pH Diagrams

Eh-pH diagrams have been widely used for the past 50 years as a convenient and visually effective method to evaluate the stability of copper minerals in contact with aqueous solutions, Anderson (1982), Garrels and Christ (1965), Sato (1960), and Titley (1978). The diagrams are constructed from thermodynamic data for minerals and aqueous solution species and employ assumed or analytical data for dissolved aqueous species of interest. They are thoroughly explained as to construction, thermodynamic basis, limitations, and many other aspects by Garrels and Christ (1965).

Eh-pH diagrams show the thermodynamic relationships of minerals under specified conditions of dissolved sulfur, copper, iron, and other species, as in the following examples. In these examples, and in others throughout the text, a species enclosed in brackets [Cu+2= activity coefficient, and m = molality (moles of solute per kg of solution). Cu] signifies the activity of the species in the leach solution where a = γm, a = activity, γ2+, Al3+, Fe3+, etc. denote the valence state of the aqueous species dissolved in the solution.

The solubility of a mineral or other solid substance in a leach solution is the amount of the solid that will dissolve to produce a saturated solution. K(sp) is the solubility product of a solid substance:

For the reaction aA + bB = cC

K(sp) = [A]a[B]b

For copper sulfate where:

Cu2+ + SO4     K(sp) = [Cu2+][ SO4].

If the activity product of the reactants raised to the appropriate powers shown in the equilibrium expression exceed the value of K(sp) then the mineral will tend to precipitate. In the case of copper sulfate, if K(sp) < [Cu2+][ SO4] copper sulfate will tend to precipitate. Likewise, copper sulfate will tend to dissolve if the activity products of the reactants raised to the appropriate powers shown in the equilibrium expression are less than the value of the solubility product K(sp) >[Cu2+][ SO4].

The importance of pH in the dissolution and precipitation of copper minerals is apparent in the hydrolysis of chrysocolla.

CuSiO3.H2O + 2H+ = Cu2+ + H4SiO4

K(sp) = 10 0.8 = [Cu2+][H4SiO4]/[H+]2

log K(sp) = 0.8 = log[Cu2+] + log[H4SiO4] - 2log[H+]

According to the equation for the solubility of chrysocolla:when pH = 4 and dissolved silica = 58 ppm (0.058gms/kg sol/60gms/mole = 0.000959m)

log[Cu+2] = 0.8 - log 0.000959 - 2pH = -4.218m

If log [Cu+2] is greater than –4.218m (0.0000605m), chrysocolla will tend to precipitate. If [Cu+2] is less than 0.0000605m, chrysocolla will tend to dissolve. The calculations suggest that groundwater saturated in silica at pH = 4 could be an effective leaching agent. Over the thousands or millions of years that encompass a geological process, the ground water could slowly dissolve chrysocolla and transport the copper to some other location. By contrast, the calculations indicate that in the case of a commercial operation, which must, of economic necessity, leach and recover copper within a few months, chrysocolla would be stable and copper would be immobile.

If the pH of the leach solution is lowered to 2, then log[Cu2+] increases to –1.182m (0.65 m) or 42 gms/kg. Under these low pH conditions, chrysocolla would tend to dissolve and copper would tend to be mobile, even in commercial leaching operations.

In acid sulfate leaching, the solubility of antlerite and other copper sulfate minerals depends on the activity of sulfate in solution as well as pH.

Cu3SO4.(OH) 4 + 4H+ = 3Cu2+ + SO4

= + 4H2O

K(sp) = 10 8.29 = [Cu+2]3[SO42-]/[H+]4

log K(sp) = 8.29 = 3log[Cu2+] + log[SO4=] - 4log[H+]

log[Cu+2] = (8.29 - log[SO4-2] - 4pH)/3

at pH = 5 and log [SO4=] = -0.3 (or 0.5m), log [Cu2+] = -3.80m

or 1.57 x 10-4 m or 0.01g/l.

In nature, where leaching may go on for a very long time, ground water @pH=5 would slowly dissolve antlerite and leach copper from the rock. However, in commercial leaching operations, where leach times are short, antlerite would remain appreciably unaffected and copper would remain in the heap. If, however, the pH of the leach solution were decreased to 2, then log[Cu+2] increases to -0.59 (16.5 g/l) causing antlerite to dissolve and copper to leach from the heap.

The stability relationship between cuprite and tenorite demonstrates the effects of pH and Eh on the relative stabilities of mineral species.

Cu2O + H2O = 2CuO + 2H+ + 2e-

Eh = Eo - 0.059pH = 0.339 – 0.059pH

Calculations show that when pH = 4 then Eh = 1.1, and when pH = 2, Eh = 1.221. On a graph where Eh is the ordinate (vertical axis) and pH, the abscissa (horizontal), a plot of equilibrium calculations for all conditions of pH and Eh for the reaction Cu2O + H2O = 2CuO + 2H+ + 2e- gives a straight line. Below the line, cuprite is stable with respect to tenorite and above the line tenorite is stable with respect to cuprite.

The construction of Eh-pH diagrams such as Fig. 1 (also see Garrels and Christ, 1965 p. 232) using various values for dissolved sulfur, iron, SiO2, Mn, CO2 and other species indicates that copper minerals form, or are preserved in, oxide and chalcocite deposits as a result of geological and geochemical conditions as follows:

1. The Chemical composition of the host rocks.2. The content of pyrite and other sulfide minerals in the deposit prior to weathering.3. The Eh-pH at various parts of the deposit during weathering.

Chrysocolla is abundant in oxide copper deposits because it is stable relative to other copper minerals under oxidizing- moderate pH conditions where dissolved silica is near its solubility limit. The stability of chrysocolla relative to other secondary copper silicate minerals is enhanced by the widespread substitution of Fe, Mn, Al and other cations into the chrysocolla structure. Conditions favorable to chrysocolla deposition occur in many weathered porphyry copper deposits of low to moderate sulfide content in andesite, monzonite, or granite host rocks. The oxidation of pyrite and chalcopyrite in these deposits produces sulfuric acid sufficient to break down some silicate gangue minerals but not enough to mobilize copper from the rock. Chrysocolla deposits form in ground water environments where pH is in the range of 5 to 6 and [H4SiO4] is at or near silica saturation.

Tenorite is common in oxidized copper deposits of original low sulfide contents where dissolved silica was low. Native copper, cuprite and delafossite occur towards the bottom of oxide deposits where Eh has been lowered through the reaction of the ground waters with gangue minerals. Delafossite, as opposed to cuprite or native copper, is favored if dissolved ferrous iron is high. Malachite and azurite are common where dissolved CO2 and pH are high, as in oxidized skarn and limestone hosted copper deposits.

Antlerite, brochantite and other copper sulfate minerals occur in oxidized chalcocite blankets where sulfate was high and pH was in the range of 3 to 5. Antlerite deposits are preserved from dissolution by desert conditions or by an influx of neutral to basic ground water.

Atacamite is found in very arid regions where total chloride in the ground water is high. Not surprisingly, atacamite is a common mineral in the secondary copper deposits of northern Chile, one of the driest regions on earth.

The significance of Fig. 1 to the leaching of copper in commercial operations is that all copper minerals, including oxides, hydroxides, sulfides, phosphates, sulfates, carbonates, and silicates, are soluble under acid-oxidizing conditions as exist in sulfate leach solutions. In nature, the leaching of copper deposits commonly produces leached caps that may contain as little as a 50 ppm copper residue. This is not to say, however, that all copper minerals dissolve at the same rate or that all of the copper in a heap or dump may be recovered during the life of a given operation; far from it. Copper minerals do not all dissolve at the same rate and it may take a great deal of time to establish Eh-pH conditions conducive to leaching throughout every part of a heap. Nevertheless, all of the

copper in a copper deposit is ultimately acid soluble. In view of this fact, the purpose of leach testing is to determine if leaching is economically feasible, and, if so to optimize leach kinetics and acid consumption.

Dissolution Rates of Copper Minerals in Leach Solutions

Relative to rock forming minerals, copper minerals dissolve quite rapidly in oxidized acid sulfate leach solutions. For example, as measured in the laboratory, chrysocolla dissolves at least 1010 times faster than the common rock forming mineral albite, while chalcopyrite, a copper mineral notorious for its slow dissolution kinetics, dissolves about 103 faster than albite. Copper minerals may be grouped by dissolution rate into five categories, Very Rapid, Rapid, Moderate, Slow, and Very Slow ( Fig. 2 ).

 

In the Very Rapid category are copper sulfates, chlorides, and carbonates. Sulfates are soluble in water. They dissolve very quickly releasing metals and acid. Chlorides dissolve quickly releasing metals and chloride. Carbonates dissolve violently in sulfuric acid releasing copper and CO2 gas.

In the Rapid category are copper silicates and oxides containing Cu2+, Fe3+, and Mn4+. The high measured dissolution rates of copper silicates and oxides are assisted by three additional factors:

1. These minerals tend to have very high surface areas because they usually occur along fractures as powdery coatings (solids tend to dissolve faster as their surface areas go up).

2. Chrysocolla has a fibrous, highly porous structure that promotes the entry of solutions into the mineral.

3. The dissolution of these minerals involves breaking the relatively weak Cu-O bond (according to Brady and Walther (1989), the solubility of silicate minerals is a function of the solubility of the principal metal oxide and CuO is highly soluble.

In the Moderate category is native copper and “reduced” oxide minerals containing Cu+ and Fe2+. These minerals, like chrysocolla, commonly occur as powdery or paint-like coatings on fractures. The reduced copper and copper-iron minerals dissolve more slowly than the previous two groups because dissolution of “reduced” oxide copper minerals requires removing an electron from copper (and from iron if present):

Cu+    Cu2+ + e-

It is well known that electron transfer is a sluggish process.

In the Slow category are the copper sulfide minerals: chalcocite, covellite, digenite, and others of the chalcocite series. The dissolution of chalcocite type minerals requires the oxidation of sulfur from S= to S6+ with the removal of eight electrons. It may also involve the oxidation of copper from Cu+ to Cu2+ with the removal of an additional electron.

In the Very Slow category are chalcopyrite and bornite. The dissolution of chalcopyrite requires the oxidation of iron from Fe2+ to Fe3+ and the oxidation of sulfur from S= to S6+(with the resultant removal of 17 electrons) for each ion of copper released to the leach solution.

Of practical importance, the minerals of the Very Rapid and Rapid categories have dissolution rates that permitted vat leaching (contact times of up to 48 hours). Copper minerals in the Moderate and Slow categories are permissive of heap leaching (contact times of 30 to >300 days). Chalcopyrite dissolves so slowly that heap leaching is not a feasible recovery method. It is frequently stated that chalcopyrite cannot be recovered by acid sulfate leaching because it is insoluble. However,

equilibrium considerations given below shows that chalcopyrite is highly soluble in leach solutions; the problem with chalcopyrite is not one of solubility but of dissolution kinetics.

Eh AND pH - The Driving Forces in Copper Leaching

Chemical reaction requires a driving force. The rate at which a chemical reaction proceeds is determined by the magnitude of the driving force. In oxide copper leaching (as opposed to sulfide leaching), the driving force is the activity of hydrogen ions in the leach solution as shown in the following equation:

CuSiO3.H2O + 2H+   Cu2+ + SiO2 + 2H2O

ΔG o (reaction) = ΔGo + 2.303RT(log[Cu2+] – 2log[H+])

Where ΔGo = ΔGo(Cu2+) + ΔGo

(SiO2) + ΔGo(H2O) – ΔGo

(CuSiO3.H2O) – ΔGo(H+)

In leach solutions where [Cu+] is held constant, ΔG (reaction) is a function of 2log[H+]. According to the conventions of thermodynamics, a reaction is spontaneous if results in an overall decrease in the free energy of the system, i.e. if ΔG o 

(reaction) is negative. In the above equation it follows that as [H+] increases, ΔG(reaction) will decrease and under highly acidic conditions, the chrysocolla will spontaneously dissolve. When ΔG(reaction) = 0 the products and reactants are in equilibrium. When Δ;G(reaction) > 0 chrysocolla will tend to spontaneously precipitate.

In sulfide leaching, the driving force is electron acceptors in the leach solution (Fe3+ and oxygen are the most common) as shown in the following equation:

FeS2 + 14Fe3+ + 8H20 = 15Fe2+ + 2SO4= + 16H+

Log K = [Fe2+] 15 [SO4]2[H+]16/[Fe3+] 14

ΔGo (reaction) = ΔGo + 2.303RT(15log[Fe2+] + 2log[SO4

=] + 16log[H+] - 14log[Fe 3+])

In respect of the solubility of chalcopyrite under equilibrium conditions, (ΔG(reaction) = 0):

Cu FeS2 + 16Fe+3 + 8H20 = Cu+2 + 17Fe+2 + 2SO4-2 + 16H+

ΔGo = -164.17 = -2.303RT(log[Cu2+] +17log[Fe2+] + 2log[SO4]-16pH - 16log[Fe 3+])

120.36 = -1 +2(0) –16(2) +17log[Fe2+] - 16log[Fe 3+]

153.36 = +17log[Fe2+] - 16log[Fe 3+]

When log[Fe2+] = -1, log[Fe 3+] = -10.65

Under normal leaching conditions where log[Cu2+] = -1, log[SO4] = 1, pH =2, and log total Fe = -1,chalcopyrite would be soluble for all log[Fe 3+] values above –10.65 or, for all Eh values above 0.2.

Because copper leaching is driven by the activity of hydrogen ions and electron acceptors (ferric ions) in the leach solution, it follows that leach kinetics depend on two fundamental properties of the leach solution: (1) the activity of H+ (measured as pH) and (2) the oxidation potential (measured as Eh). At the most fundamental level, we may imagine copper leaching as a process in which we drive protons (H+) into the system while we pull electrons out of it.

The location of the copper (aq) field in the high Eh, low pH corner of in Fig. 1 is a graphical portrayal of the idea that leaching requires a solution of high Eh and low pH. The commercial leaching of copper is virtually identical to natural processes such as the corrosion of metals, the weathering of rocks and soils, and the leaching of primary ore deposits. The importance of oxidation and hydrolysis in natural processes was recognized by Pourbaix (1947) and Garrels and Christ (1965) and others who used Eh - pH diagrams to evaluate the geochemical environments of soil formation, the genesis of uranium deposits, the formation of leached cappings and chalcocite blankets and many other natural phenomena.

The Importance of Whole Rock Mineralogy in Copper Leaching

While oxide copper minerals dissolve rapidly in acid sulfate leach solutions, the amount of copper that is ultimately extracted from a heap is largely dependent on the percentage of the heap exposed to acid-oxidizing conditions. In this regard, the principal obstacle to successful leaching is the persistence within the leaching system of low Eh-high pH microenvironments. A low Eh high pH microenvironment may exist within an individual mineral grain, an ore fragment, or some large part of the heap blinded from the leach solution. Because low Eh-high pH micro environments are present within virtually all leaching systems, it is a given that the macro chemical conditions as measured in the pregnant leach solution differ from the micro conditions existing within minerals, ore fragments, or areas of the heap blinded from the leach solution.

Acid Consumption

As a leach solution flows down through a heap, it reacts with ore and gangue minerals in a great many hydrolysis reactions (exemplified by the dissolution of chrysocolla) and redox reactions (exemplified by the dissolution of pyrite).

Silicate gangue minerals comprise more than 99 percent of most oxide porphyry copper ores. Given this fact, we would expect that the behavior of silicate gangue minerals in sulfuric acid solutions would be very important to copper leaching. The most common silicate gangue minerals in porphyry copper deposits are:

Important Gangue Minerals

quartz SiO2

orthoclase KAlSi3O8

plagioclase NaAlSi3O8 – CaAl2Si2O8

biotite KFe3(AlSi3O10(OH)2

chlorite Mg5(Al,Fe)(Al,Si)4O10(OH)8

sericite KAl2(AlSi3O10)(OH) 2

montmorillonite (Mg,Ca,Fe)(Al,Mg,Fe)4(SiAl)8O20(OH)4.nH2O

vermiculite (Mg,Ca)(Al,Mg,Fe)6[(Al,Si)8O20](OH)4.8H2O

kaolinite Al4Si4O10(OH)8

hornblende Ca(Mg,Fe)3Si4O12

actinolite Ca2(Mg,Fe,Al)5Si8O22(OH) 2

muscovite K2Al4(Si6Al2O20)(OH) 2

In the laboratory, silicate minerals are shown to dissolve in hydrolysis reactions:

KAlSi3O8 + 4H+ + 4H2O = K+ + Al3+ + 3H4SiO4

K(reaction) = [K+][Al3+][H4 SiO4]/[ H+]4

The dissolution of orthoclase, as shown above, occurs by means of a hydrolysis reaction driven only by pH. Because reaction rates depend on the chemical driving force, the rate at which orthoclase dissolves in hydrolysis reactions is a function of pH ( Fig. 4 ). In the case of plagioclase, the rate of dissolution increases by about 1000 times as the pH of the leach solution is lowered from 4 to 1.

Over the past ten years or so, dissolution rates of a number of important rock-forming minerals have been measured in the laboratory. As shown in Fig. 4, Ca plagioclase dissolves at least 100 times faster than orthoclase, which dissolves 10 to 100 times faster than sericite. Brady and Walther (1989) demonstrated that in acid aqueous solutions, the dissolution rates of silicate minerals depend on crystal structure and metal – oxygen bond strength. To a first approximation, they showed that the rate of mineral dissolution is proportional to the solubility of the principal metal oxide in the mineral.

The laboratory work cited above generally corroborates field studies on the rates of weathering of various types of rocks. For example Goldich (1938) showed that (1) basalt, gabbro and diorite (rocks composed of pyroxenes, Ca plagioclase, and hornblende) break down quickly in the weathering environment, (2) monzonite, andesite, dacite, and granite (rocks composed of albite, orthoclase, and biotite) break down at moderate rates, and rocks composed of quartz, muscovite, sericite, and kaolinite such as those forming the alteration envelopes around porphyry copper deposits weather very slowly ( Fig. 3 ). In the tropics, areas of quartz-sericite alteration tend to form ridges because they are less susceptible to chemical weathering than surrounding rocks.

Field and laboratory studies, column testing ( Fig. 5 ), and operational experience show that:

1. Copper ores hosted by basalt and diabase such as Ray, Arizona or the copper deposits of the Coast Range of Chile tend to consume a great deal of acid.

2. Ores composed of monzonite, andesite, and granite, consume moderate amounts of acid.3. Ores hosted by quartz-sericite alteration consume small amounts of acid. and4. Ore in quartz sandstone consumes almost no acid.

Because porphyry copper deposits typically have a variety of rock and alteration types, they tend to contain a number of distinct ore assemblages (Fig.. 6), each characterized by unique acid consuming properties and leach kinetics.

Leaching Kinetics

Silicate gangue may be subdivided into five categories according to the rate at which the component gangue minerals react with acid. At the relatively high acid consuming end of the spectrum, basalt-diabase gangue reacts quickly with leach solution. At the lower end, quartz sandstone is almost unreactive. As gangue becomes more reactive its capacity to neutralize the leach solution increases.

Highly Reactive Gangue: Highly reactive gangue contains significant amounts of carbonate minerals, including calcite, dolomite, and siderite. These and other carbonate minerals dissolve on contact with the leach solution producing cations, water and CO2 gas. One kg of calcite will consume 0.98 kg of acid so one percent of calcite in a rock will consume 9.8 kg of acid. Depending on copper content, the presence of calcite in amounts of more than a few percent will render most oxide copper deposits unsuitable for leaching.

Reactive Gangue Reactive gangue contains significant amounts of hornblende, pyroxene, and Ca plagioclase ( upper part of Fig. 3 ), minerals that dissolve quickly relative to other silicate minerals according to dissociation reactions as follows:

CaAl2Si2O8 + 8H+ = Ca2+ + 2Al3+ + 2SiO2 + 4H2O

Ca(Mg,Fe)3Si4O12 + 8H+ = Ca2+ + 3(Mg2+ + Fe2+) + 4SiO2 + 4H2O

Leach solution descending through a column of reactive gangue becomes involved in hydrolysis and redox reactions. As these reactions consume hydrogen and ferric ions along the flow path, the leach solution will undergo progressive and continuous change such that pH will increase and Eh will decrease. This will cause the chemical driving force behind the dissolution of copper minerals to diminish, which, in turn, will result in progressively slower copper leach kinetics and a decrease in the percentage of copper recovered from each succeeding unit of depth in the column of ore (Fig.. 7 and 8). Depending on the length of the flow path and the reactivity of the gangue, at some point the Eh-pH component of the solution may leave the field of Cu2+ and enter the field(s) of native copper, cuprite, or delafossite. If this occurs, copper leaching will stop and copper precipitation might begin.

Moderately Reactive Gangue Moderately reactive gangue contains orthoclase, biotite, albite, and quartz; silicate minerals that also dissolve as a result of hydrolysis reactions:

KAlSi3O8 + 4H+ = K+ + Al3+ + 3SiO2 + 2H2O

KFe3(AlSi3O10(OH)2 + 10H+ = K+ + 3Fe2+ + Al3+ + 6H2O + 3SiO2

Biotite, albite, and orthoclase dissolve more slowly than hornblende and Ca plagioclase. For this reason, the hydrolysis reactions involving biotite, albite, and orthoclase consume fewer hydrogen and ferric ions per unit length of solution flow path than hornblende and Ca plagioclase with the result that at every point along the flow path, pH is lower and Eh is higher than in a column of reactive ore. Because fewer hydrogen and ferric ions are consumed in a unit length of moderately reactive gangue, pH remains lower and Eh higher over a greater proportion of the solution flow path through the ore column. The end result is that moderately reactive gangue has (1) faster leach kinetics for each segment of the flow path and (2) a reduced likelihood that Eh-pH conditions will move out of the Cu2+ field before the leach solution exits the ore column. Thus, the net effect of leaching moderately reactive gangue, as compared to reactive gangue, is that leaching kinetics are higher, acid consumption is lower, and the leaching of copper is more uniform throughout the ore column.

Slightly Reactive Gangue Quartz, sericite, and/or kaolinite are in equilibrium (quartz) or near equilibrium (sericite and kaolinite) with the leach solution. Because quartz is stable in leach solutions and sericite and kaolinite dissolve very slowly, the leach solution undergoes only minor depletion of hydrogen and ferric ions as it flows through the column of ore. As a result, (1) pH and Eh retain their initial values throughout the solution flow path and (2) the chemical force driving the dissolution of copper minerals remains high along the flow path. Thus, leaching rates for slightly reactive gangue are substantially higher than for ores hosted by reactive and moderately reactive gangue and acid consumption is substantially lower.

Non-reactive Gangue Quartz sandstones are sensibly unreactive to leach solutions and leaching is accompanied by little, if any acid consumption. If these rocks are porous, copper leach kinetics are high, acid consumption is low, and leaching is uniform through the ore column.

The relationship between acid consumption and leaching kinetics for three ore types is shown in Fig.. 8.

The same principles governing the depletion of hydrogen and ferric ions in the leach solution also apply to individual ore fragments. As shown in Fig.s. 9 and 10, leach solution flowing towards the center of an ore particle may become depleted in hydrogen and ferric ions. If the ore fragment is sufficiently reactive that hydrogen and ferric ions are consumed faster than they are supplied, the Eh-pH component of the solution may move out of the Cu2+field. The locus of all points around an ore

particle where Eh-pH conditions depart from the Cu2+ field defines a reaction rim. Outside of the reaction rim, copper is leached from the particle; inside of the rim copper remains unleached. Because the solubility field for iron (Fig.. 12) extends to higher pH regions than that of copper, iron oxides may be leached from the entire ore particle while copper is leached only from the outer rim. As might be expected, the more reactive the gangue, the narrower the reaction rim and the less complete the leaching (Fig..11). In practice, the more reactive the gangue and the narrower the reaction rim, the finer the ore must be crushed to maintain Eh-pH conditions conducive to copper leaching for the bulk of the ore. Paulson and Kuhlman (1989) have reported the preservation of oxide and native copper within individual biotite and plagioclase grains whereas copper was leached from the matrix surrounding the grains.

Because copper leaching kinetics depends strongly on pH and Eh, laboratory leach tests must be conducted under Eh-pH conditions that will be used in the proposed (or ongoing) operation. If the tests employ a solution of unrealistically low pH, then copper leach kinetics and acid consumption experienced in the test will be higher than those likely to be encountered in the operation. Conducting leaching tests using an unrealistically high solution pH will result in inappropriately low leach kinetics and low acid consumption. Leach tests that employ inappropriate Eh conditions may also produce misleading data.

The Geochemistry of Iron

In most oxide copper deposits iron is present in a variety of oxide, hydroxide, sulfate, sulfide and silicate minerals that react readily with the leach solution. Depending on the relative abundance of these minerals, leaching may release a preponderance of ferric or ferrous ions into the leach solution and cause Eh to go either up or down, as the case may be. The types of reactions that involve iron and the overall effects of these reactions on the chemistry of the leach solution are illustrated by the following reactions:

1. Eh, is determined by the ferric/ferrous ratio of the leach solution.

Fe+2 = Fe3++ e-

Remembering that

ΔG(reaction) = ΔG°f + 1.36log[Fe3+]/[Fe2+]

Eh = Eo + 1.36log[Fe3+]/[Fe2+] = 0.059log [Fe3+]/[Fe2+]

(n)23.06kcal/volt-gram-equiv.

(n = the number of electrons involved in the reaction)

2. The dissolution of copper sulfides and reduced oxides such as cuprite and native copper depend on dissolved ferric iron as the oxidizing agent.

CuFeS2 + 18Fe3+ + 8H2O   Cu2+ + 19Fe2+ + 16H+ + 2SO4 =

Eh-pH relations for the iron-sulfur-water system are shown in Fig. 12.

The iron content of typical porphyry copper ore is between 2% and 6%. In unoxidized deposits, the iron is present largely in the ferrous state in the following minerals:

1. Sulfides: pyrite, chalcopyrite, bornite2. Oxides: magnetite (Fe3O4)

3. Silicates: biotite, hornblende, actinolite, chlorite, and montmorillonite.

It was stated previously that the dissolution of biotite, hornblende and other iron silicate minerals (as well as iron sulfide minerals) releases ferrous iron into the leach solution. The result of iron silicate mineral dissolution is a lowering of the ferric/ferrous ratio. Because Eh is a function of Fe3+/Fe2+, lowering the ferric/ferrous ratio lowers the Eh of the leach solution. If the dissolution of iron silicate minerals is sufficiently rapid, it may cause the Eh of the leach solution to descend into the chalcocite field of Fig. 1. Should this take place, the leaching of copper would terminate and the precipitation of chalcocite (or native copper, delafossite or cuprite) might take place.

In recent years, ARD, the release of acid and toxic metals from mine workings and dumps by the oxidation and dissolution of sulfide minerals has been extensively studied (Plumlee and Logsdon, 1998). The following discussion draws heavily on the review of the geochemistry of iron summarized by Nordstrom and Alpers (1998) published in a volume devoted to the environmental geochemistry of mineral deposits.

The Oxidation of Sulfide Minerals

Sulfide minerals oxidize by means of three general mechanisms: (1) the direct application of oxygen, (2) by ferric ions, and (3) the action of microbes.

Oxidation of pyrite by the direct action of oxygen proceeds by the overall reaction:

FeS2 + 15/4 O2 + 7/2H2O   Fe(OH)3(solid) + 4H+ + 2SO4=

Laboratory rates for this reaction range from 1.1 x 10-10 to 5.3 x 10-10 moles per square meter per second.

The dissolution of pyrite minerals by the action of dissolved ferric iron has been shown to occur according to the balanced reaction:

FeS2 + 14Fe3+ + 8H2O   15Fe2+ + 2SO4= + 16H+

Rates for the dissolution of pyrite by ferric iron in the laboratory range from

9.6 x 10-9 to 1.8 x 10-8 moles per square meter per second. In this reaction, oxygen is required to oxidize the ferrous iron to ferric.

While the specific mechanisms operating at the surface of a dissolving sulfide grain are still debated, various laboratory experiments show that ferric iron dissolves sulfides 100 to 1000 times faster than oxygen. In heap leaching, the oxidation of iron by the direct action of oxygen is limited by the low solubility of oxygen in water at 25 degrees C (8.11 mg/liter or 5 x 10-4m) and the difficulty in replenishing oxygen to the leach solution in the deeper parts of heaps and dumps.

The participation of microbes in the oxidation of sulfides has been known or suspected for more than a century and has been intensively investigated for the past 40 years or so. Research has shown that microbial action enhances the oxidation of sulfides by catalyzing the ferrous to ferric transition. Although the site of the microbial mediation of ferrous to ferric oxidation has not been established, evidence suggests that much of it occurs in the solution (as opposed to the surface of the pyrite). Laboratory measurements indicate that the microbial mediated oxidation of ferrous iron occurs at rates of 5 x 10-7 to 8.8 x 10-8moles per square meter per second, or about 10 times faster than by abiotic oxidation of ferric iron (see table 6.5 of Nordstrom and Alpers (1998)).

The quantity of iron that may be held in a leach solution and the ferric/ferrous ratio of the leach solution are largely determined by of the solubility of jarosite:

KFe3 (SO4)2(OH)6 + 6H+   K+ + 3Fe3+ + 2(SO4) = + 6H2O

K(sol. Prodct) = [K+][Fe3+] 3[SO4=] 2/[ H+]6

Because Eh = 0.059log [Fe+3]/[Fe+2], we see that the solubility of jarosite depends on Eh, pH, [K+] and [SO4

-2]. The solubility of jarosite under conditions of constant total dissolved sulfate (continuous addition of acid to the heaps) and a constant supply of potassium due to the dissolution of orthoclase and biotite leads to the following conclusions:

1. If ferric iron is consumed in the dissolution of biotite or sulfides, jarosite will tend to dissolve and release ferric iron into solution.

2. If [Fe3+] increases, jarosite will precipitate and Fe3+/ Fe2+ will decrease.3. If solution pH increases, jarosite will precipitate releasing H+ (lowering pH) but consuming

Fe3+ (lowering Eh).

In fact, data from column leach tests show that Eh – pH conditions for leach solutions do tend to stabilize along the ferrous sulfate (aqueous) – jarosite phase boundary (Fig. 12) in general agreement with the preceding discussion.

If solution pH increases above 2.5 to 3, or if total dissolve sulfur decreases beyond a certain level, then the geochemistry will be controlled by the precipitation and dissolution of various FeO(OH) compounds including goethite and ferrihydrite.

FeO(OH) (solid) + 3H+   3Fe3+ + 2 H2O

If the ore is composed of unreactive gangue minerals but has a high pyrite content, the oxidation of the pyrite may produce a very low pH leach solution. Given sufficient time, pH may trend down below 1.5 or 1 (into the ferric sulfate field of Fig. 12), and total dissolved iron may become very high. In most ores, however, acid consuming silicate minerals result in Eh-pH conditions within the jarosite field of Fig. 12.

Acid Curing

In the 1920’s and 1930’s the U. S. Bureau of mines conducted leach tests at the Tucson, Arizona research center to develop an economic method to recover copper from low-grade oxide copper ores. The low copper values of the ores dictated an approach involving minimal handling. Experience with dump leaching pointed to piling the oxide ore in heaps and applying an acid solution to the top. In testing this general approach, it was found that leaching rates were low and acid consumption was high. Further work by the Bureau of Mines and by Inspiration Consolidated Copper Company in the 1970’s showed that leach kinetics and acid consumption could be considerably improved by applying a strong acid solution to the ore at the beginning of the leach cycle. Inspiration further improved performance by adding ferric iron to the initial acid solution. The initial acid dosage was termed a “cure”.

The geochemistry of acid curing is best understood by further examination of Fig. 1. The initial Eh-pH conditions of the leach solution prior to its application to the top of the heap are within the Cu2+ field while, for the ore, they are similar to the Eh-pH conditions of ground water in the mine area (in the southwest U. S. A. Eh is about 0 and pH is about 8). If the ore contains Reactive or Moderately Reactive gangue, leaching will quickly lower Eh and increase pH. If the rate of solution-mineral reaction is faster than the rate of acid addition, depletion of hydrogen and ferric ions will move the composition of the leach out of the Cu+2field and into the fields of chrysocolla, native copper, or chalcocite.

At the point shown for Arizona ground water on Fig. 1 the solubility of copper is well below 10-6m. The rapid neutralization of the leach solution caused by the dissolution of gangue that is roughly in equilibrium with Arizona ground water chemistry will result in the formation of an acid front within the ore column (Fig. 13). Behind (above) the acid front, Eh and pH are within the Cu2+ field of Fig. 1 and copper is leached. Across the acid front Eh drops from 0.6 – 0.5 to 0.1 – 0 and the solubility of copper drops from 0.1m (6 gms/l) to 10-6m, a decline of about 5 orders of magnitude. Below the acid front, the solubility of copper is very low, (perhaps 10-7m). If the acid front develops well up in the ore column, the very low solubility of copper in the neutralized solution may lead to the precipitation of copper within the column of ore. Common copper precipitates within columns and heaps include native copper, delafossite, and chalcocite. If the acid front develops near the bottom of the column, copper leaching may terminate but there might be insufficient time for copper precipitation to occur. In any case, the results of many years of testing that included Eh-pH measurements of the leach solution indicate that the function of the cure is to establish high Eh low pH conditions throughout the entire ore column at the beginning of the leach cycle.

Chemistry of the Leach Solution

The chemical composition of leach solutions results from (1) the action of a lixiviant on the ore and gangue minerals, (2) the exchange of copper for hydrogen ions in the SX/EW plant, and (3) the precipitation of new minerals from the leach solution. Considering these factors it is evident that the chemistry of the leach solution depends on the following:

1. The solubility and dissolution kinetics of ore and gangue minerals,2. The speciation of Cu, Fe, Mg, S, Na, and other solutes in the solution,3. The solubility of various potential mineral precipitates,4. The rate of removal of copper in the SX/EW plant,5. The rate of addition of new acid to the operation and,6. The rate of addition of now ore (of each compositional type) to the heap.

Because points (1) and (5), and, obliquely, (5) and (6) were discussed previously and the chemistry of the SX/EW plant is treated elsewhere in the course, this section focuses on (2) the speciation of the leach solution and its related topic, (3) mineral precipitation.

Speciation of the Leach Solution Leach solutions are typically complex solutions, containing Cu, Zn, Fe, Al, Na, K, Ca, Mg, Mn, Si, C, P, S, Cl, H, PO4 and other constituents. The behavior of the solution depends on the relative and absolute amounts of these constituents present, and on the ways in which these constituents combine (their speciation) in the leach solution.

Analyses of leach solutions indicate bulk solution compositions generally similar to the following table:

Element Composition (gms/liter)

Cu 0.5 – 15

Zn 0.05 – 0.3

Fe 0.5 – 10

Al 1.0 – 8

Na 0.01 – 0.04

K 0.02 – 0.06

Ca 0.04 – 0.07

Mg 1.0 – 4.0

Mn 0.01 – 1.0

Si 0.01 – 0.3

C 0.0

P 0.1 – 0.5

S 10 – 20

C1 1 – 10

Experimental work shows that copper may occur in aqueous solutions as: Cu+, Cu2+, CuSO4, Cu(OH)2, Cu(OH)3

1-, Cu(OH)42-, CuCl, and CuCl2 , while iron may occur as Fe2+, Fe3+, FeHPO4

+, FeSO4, FeHSO4

+, FeHPO4, FeOH+, Fe(HS)2, Fe(HS)3-, FeHPO42+, FeSO4

+, FeHSO42-, FeOH2+,

FeHPO4+, Fe(OH)3, Fe2(OH)2, Fe(OH)4

-, Fe3(OH)45+. The other dissolved solids also occur in a

multiplicity of aqueous species. Because leach solutions may potentially consist of hundreds or thousands of solution species it is necessary to perform equilibrium calculations to determine the degree to which each species is present in solution and how adjustments might occur should certain key solution parameters (such as Eh or pH) change. This exercise requires that equations for all potential solutions be written as shown below and then solved simultaneously.

Fe2+ = Fe3+: K(react) = [Fe3+]/[Fe2+]

Fe+3 + 3H2O = Fe(OH)3 + 3H+: K(react) = [Fe(OH)3][H+]3/[Fe3+]

Fe+2 + SO4=

 = FeSO4 K(react) = [FeSO4]/[Fe+2][SO4=]

In view of the large number of possible solution species, a computer is required to solve the equations. The example included with this chapter was solved using WATEQ, Ball and Nordstrom, (1991). An overview of the example solved is as follows:

Element Principal Solution Species

Cu Cu2+, Cu SO4

Zn Zn2+, Zn SO4

Fe Fe+2, FeSO4, FeH2PO42+

Al AlSO4+

Na Na+

K K+

Ca Ca2+, Ca SO4

Mg Mg2+, MgSO4

Mn Mn2+, MnSO4

Si H4 SiO4

P FeH2PO42+

S AlSO4+, MgSO4

The example solution species distribution and the above table indicate the strong predilection for iron, copper, manganese, magnesium, calcium, and zinc to complex with sulfate in leach solutions. Indeed, copper heap leaching is only possible because copper (and iron) forms a strong aqueous sulfate complex. Species distribution calculations using other values for pH will differ from the above example because the stability of solution species is strongly dependent on:

1. pH

Fe3+ + 3H2O = Fe(OH)3 + 3H+ K(react) = [Fe(OH) 3][H+]3/[Fe3+]

Fe2+ +HSO4- = FeSO4(aq) + H+

 K(react) = [FeSO4][H+]/[Fe2+][H SO4=]

HSO4-1 = H+ + SO4

= K(dissoc) = [SO4=][H+]/[HSO4

1-] = 10-1..99

And,

2. Eh:

Fe2+ = Fe3+ + e- K(react) = [Fe3+]/[Fe2+]

Eh = Eo + 0.059log[Fe3+]/[Fe2+]

Fe2O3(solid) + 4H2S = 2FeS2 + 3H2O + 2H+ + 2e-

Eh = Eo - .059 log pH – 0.118log[H2S]

FeS2 + 4H2O = Fe2+ + 2SO4= + 8H+ + 8e-

Eh = Eo - 0.059 log pH + 0.015log[SO4=] + 0.007log[Fe2+]

The Solubility of Mineral Precipitates in Leach Solutions The solute geochemistry of locked-cycle leaching is characterized by; (1) an initial phase of increasing solute concentration followed by, (2) a prolonged phase, possibly including the remainder of the life of the operation, of relatively constant solution composition. The initial phase of leaching, in which total dissolved solids (TDS) increase, occurs at the beginning of leaching and for some time thereafter as the dissolution of ore and gangue minerals releases solutes to the leach solution but, before the solubility limits of mineral salts are reached. As TDS reaches the solubility limits of silica, gypsum, jarosite, clays and other solids the second and longest phase of leaching begins. The second phase of leaching takes place when the solubility products of first one, and then all new solids are exceeded and mineral precipitation begins ( Fig. 14 ). The example solution speciation calculation referred to above and included with this chapter contains a section on a number of potential solid precipitates. The tendency for solids to precipitate from the leach solution depends on the solubility constant of the solid phase and the activities of its component species in the leach solution :

H2O + SiO2 = H4 SiO4 K(sp = [H4 SiO4] = 10-2.7

[SiO2] = 120 ppm

KFe3(SO4)2(OH)6 + 6H+ = K+ + 3Fe3+ + 2(SO4)= + 6H2O

K(sp) = [K+][Fe3+] 3[SO4=] 2/[H+]6 = 10-14.8

CaSO4.2H2O = Ca2+ + SO4

= + 2H2O

K(sp) = [Ca2+][SO4=2] = 10-4.58

In the example solution where log [SO4-2] = -2.342, log [Ca+2 ] should stabilize at –2.238 or 231 ppm.

FeO(OH) +3H+ = Fe3+ + 2H2O

K(sp) = [Fe+3]/[H+]3 = 104.891

At pH = 2: log [Fe3+] = K(sp) – 3pH = 4.891 – 6 = -1.109

At pH = 5: log [Fe3+] = K(sp) – 3pH = 4.891 – 15 = -10.811

Al2Si2O5(OH)4 + 6H+ = H2O + 2Al3+ + 2H4 SiO4

K(sp) = [Al3+][H4SiO4]2/[H+]6 = 107.435

In the example solution where H4SiO4 = 2 x 10-3 (log H4SiO4 = -1.699) and pH = 2

log[Al3+] = 7.435 – 2(-1.699) – 6pH = -1.167

At pH = 5: log[Al3+] = 7.435 – 2(-1.699) – 6pH = -19.167

As shown in the attached solution speciation example, copper leaching results in the precipitation of many new minerals. In addition to silica, jarosite, and gypsum, predicted precipitates include various sulfates and hydroxides of iron, aluminum, magnesium, and manganese. The concentrations of K, SO4, Fe3+, Fe2+, (and possibly sodium and potassium) are controlled mostly by the solubility of jarosite. Over the life of a leaching operation, tens or hundreds of thousands of tonnes of such minerals will precipitate from the leach solution within the heap. Within rock fragments, the concentrations of aluminum, silica, and potassium in solution may be fixed by the solubility of pyrophyllite (Al2Si4O10(OH)2), kaolinite, and illite (KAl4 Si7O20(OH)8). The relationship of pyrophyllite and kaolinite to orthoclase in terms of dissolved silica and the ratio: log [K+]/[H+] (solution) is shown on Fig. 15 where it can be seen that a decrease in [K+]/[H+] will promote the dissolution of orthoclase and the precipitation of kaoilnite. A further decrease in [K+]/[H+] may lead to the dissolution of kaolinite and the precipitation of pyrophyllite. In respect of theFig. 10, leach solutions might be expected to dissolve orthoclase and precipitate kaolinite or pyrophyllite but this has not been reported in the literature. From the preceding discussion it should be obvious that to predict or understand acid consumption, mineral precipitation, and most other aspects of leaching periodic total chemical analyses of the leach solution must be made.

Copper Leaching: A Geochemical Cycle

In column testing it is generally found that open cycle copper leaching leads to higher acid consumption than locked cycle leaching ( Fig. 17 ). While this might seem contrary to expectations, examination of mineral hydrolysis reactions shows that copper leaching involves not only acid consuming reactions but also, acid producing reactions. While the idea of acid consuming reactions in copper leaching is readily understood because non sulfide leaching operations must buy acid to replace that which is consumed, it is less well known that acid producing reactions also take place during leaching. (By this I am not referring to acid produced by the oxidation of pyrite during leaching or the acid generated in the SX plant by the exchange of hydrogen ions for copper). In copper leaching, acid is generated whenever a mineral having an (OH) in its formula - an example would be FeO(OH) - precipitates from the leach solution. Removing (OH) from water (H2O) leaves H+. In effect, the precipitation of a compound or mineral having (OH) in its formula is the reverse of hydrolysis because it generates acid. The mechanism for acid generation by reverse hydrolysis is illustrated in the following reactions:

K+ + 3Fe3+ + 2(SO4)-2 + 6H2O   KFe3(SO4)2(OH)6 + 6H+.

Fe3+ + 2H2O   FeO(OH) + 3H+

H2O + 2Al3+ + 2H4SiO4   Al2Si2O5(OH)4 + 6H+

These reactions show that the precipitation of minerals which contain (OH) such as jarosite, goethite, hematite, alunite, and kaolinite releases hydrogen ions to the solution (in effect the precipitated mineral strips one or more (OH) groups from H2O leaving H+ in the solution). A quantitative evaluation of the acid generating capacity of reactions of this type shows that the precipitation of 1 kg of ferrihydrite generates 1.65 kg of sulfuric acid and the precipitation of 1 kg of jarosite, 586gms of sulfuric acid.

The fact that locked cycle leaching consumes less acid than open cycle leaching results in part from the nature of precipitation reactions. Precipitation requires the formation of a stable solid reaction product. Under equilibrium and near equilibrium conditions, chemical reactions are dynamic processes involving the simultaneous formation and destruction of products and reactants. Thus, the nucleation of solid precipitates under equilibrium conditions is often found to be difficult and time consuming. Indeed, the formation of many precipitates requires super saturation or the presence of a catalyst. In open cycle leaching where the solution flows through the ore column and is then discarded, super saturation of reaction products may not occur and there may be insufficient time for the nucleation of solid reaction products.

Also, the rate of dissolution of ore and gangue minerals is not constant in time. Rather, the more easily dissolved minerals, including those in the fine ore fraction and those with fast dissolution kinetics such as carbonates, dissolve early in the leaching process, while the less reactive minerals and those blinded from the leach solution dissolve more slowly. As a result, the rate at which solutes are added to the solution decreases with time (if no new ore is added to the column). In open cycle leaching this leads to a steady decline in TDS over time. When TDS falls below the solubility products of hydroxide (OH) type minerals, mineral precipitation comes to an end and the acid generating side of leaching is terminated.

The geochemical cycle of copper leaching also has implications on the environmental side of leaching. Normally, the closure of an operation requires rinsing a heap until effluent pH is above 6 and all solution parameters are below predetermined analytical limits set by the closure permit. Unfortunately, as rinse water is run through a heap and the pH of the effluent solution increases, jarosite, ferrihydrite and other hydroxol containing minerals precipitate and release acid. The amount of acid that could be generated in this process is apparent in the following example.

Consider a 200 MMt tonne heap saturated to11% moisture by a leach solution containing 4 gms/l iron, and 4 gms/l aluminum. The leach solution in the heap will contain 88,000 tons of Fe and 88,000 tons of Al. Precipitation of the iron as jarosite will produce 154,000 tons of sulfuric acid and the precipitation of the aluminum will produce 319,407 tons of acid. Added to the acid produced by the precipitation of iron and aluminum, would be a much greater amount of acid produced by the precipitation of (OH) minerals of Na, Mg, and Mn. All together, hydroxol mineral precipitation might easily produce 1,000,000 tons of acid.

Free Acid Determination Metallurgical tests employed to estimate the leaching behavior of copper ores attempt to measure acid consumption. In bottle roll and open cycle tests acid consumption is measured by subtracting the amount of acid remaining at the end of the test from the amount added at the beginning. While the procedure sounds simple, it turns out, for practical and theoretical reasons, to be quite difficult. Acidity is defined as the activity of the hydrogen ion in solution and is measured as pH. However, because leach solutions contain high total TDS and are highly buffered, pH understates the capacity of the solution to donate hydrogen ions. Most acid determination tests employed today seek to measure “free acid” which is supposedly the quantity of sulfuric acid in the solution neglecting the amount of acid contained in the metal-salt buffers such as jarosite. Various techniques have been devised to measure “free acid”. Most involve removing cations and then

titrating for sulfuric acid with a strong base (Booman and others, 1958, Solomans, 1980), although, the Grans Plot method, Orion Research (1970) employs dilution. Application of the various methods in common use to synthetic solutions give inconsistent results that diverge from the actual amount of sulfuric acid added to the solution as the concentration of acid and TDS increase (  Fig. 16  ).

Conceptually, an acid determination method should measure the capacity of the leach solution to donate hydrogen ions in exchange for cations within the pH range of the leaching process. Whatever method is used to determine acid consumption, it must take into account the buffering capacity of metal salts such as jarosite, alunite and clay minerals.

Geological Screening of Copper Leaching Projects Prior Metallurgical Testing

It is feasible to determine the amenability of a potential project to copper leaching prior to testing by using standard geological techniques as diagrammed in Fig. 18. In the first step of the process, field geologists should classify the acid consuming properties of the deposit into one of the four categories and the leaching kinetics into one of three. The methods used are field mapping, logging of core and cuttings, and the application of standard field methods to the identification of ore and gangue minerals. Much of this work may be accomplished in the field using simple tests that employ a hand lens, a pocketknife, and small bottles of dilute HCl and H2SO4. The tests can determine the presence and quantity of carbonate minerals, the principal silicate gangue minerals and the relative abundance of copper minerals. In the second step, the mineralogy of the deposit should be refined by petrographic methods including the study of polished sections. At the end of step two the mineralogy of the deposit should have been sufficiently quantified to permit:

1. Rejection of the property.2. Identification of potential problems concerning acid consumption and leaching kinetics.3. Estimation of leach kinetics and acid consumption.4. Design of an efficient leach-testing program as opposed to the normal shotgun approach.

Further metallurgical testing is well described in another chapter.