View metadata, citation and similar papers at core.ac.uk brought to you by CORE provided by Spiral - Imperial College Digital Repository THE IRON ELEPHANT A Brief History of Hydrometallurgists’ Struggles with Element No. 26. A.J. Monhemius Emeritus Professor Department of Earth Science and Engineering Royal School of Mines Imperial College London SW7 2BP UK (email: [email protected]) ABSTRACT The paper reviews the history of modern iron control processes in hydrometallurgy. Particular attention is paid to the processes developed in the 1960s by the electrolytic zinc industry for the removal of iron from zinc sulphate leach liquors - namely the Jarosite, Goethite and Hematite processes. Problems faced by the industry in the safe disposal of the iron residues produced by these processes are discussed. The lessons learnt over the past half century by the zinc industry are projected on to the current efforts to develop hydrometallurgical processes for the treatment of copper sulphide concentrates. It is argued that any attempt to create a “hydrometallurgical copper smelter” is doomed to ultimate failure, unless and until methods are developed that can recover iron from solution in a form that can be directly utilised by the steel or pigment industries. KEYWORDS History of iron control; jarosite process; goethite process; hematite process; copper sulphide concentrate leaching. 1 INTRODUCTION You suddenly realise that age is catching up with you and that technology is moving on apace when you are asked to talk about the history of a process or a technology that you’ve grown up with. Such was my reaction when, out of the blue, I received an email last year from the chairman of the CIM’s Historical Metallurgy Committee, asking me if I would give a talk on the early developments of iron control in hydrometallurgy. Once I’d got over the initial shock of realising that I had now become living history, I was of course flattered that the invitation was to give a talk at the prestigious IMPC, a conference series that has been going for as long as I have been in the business. That this particular IMPC was to include the 4th Symposium on Iron Control in Hydrometallurgy, made the invitation even more appealing and appropriate, because I was the co-chairman of the First Symposium on Iron Control, together with my old friend, John Dutrizac (known by many as Dr Jarosite), which was held 30 years ago in Toronto. I first became involved with iron at the very start of my career in hydrometallurgy – the research topic for my Masters degree thesis at UBC, back in the mid 1960s, was “Reductive leaching of Goethite with SO2”. My thesis supervisor was the late Professor Ian Warren, who himself was a protégé of Professor Frank Forward, one of the founding fathers of modern hydrometallurgy. Little did I realise at that time, now 50 years ago, that the problem of iron and how to deal with it in hydrometallurgical processes would be a theme that followed me throughout much of my 40-year career in academia. IRON MINERALOGY AND CHEMISTRY So what is it about iron in hydrometallurgy that has merited four decennial, international conferences – 1986 in Toronto, 1996 in Ottawa, 2006 in Montreal and now 2016 in Quebec, not to mention the thousands of scientific and technical research papers and patents that have been published worldwide outside these conferences? As we all know, iron is a pretty common element and it is one that we certainly couldn’t live without. It’s an essential part of our biology, where it plays a critical role in the transference of oxygen from the air into our bloodstreams, and it is indispensible for our modern way of life in the form of metallic iron and steels, from which most of the things around us are manufactured. Let’s start with a few facts about the geological occurrences of iron, element number 26 in the Periodic Table. To source these, I turned to the very first paper in the Proceedings of the First Iron Control Symposium, which was by Chen and Cabri (1986) from CANMET, entitled “Mineralogical overview of iron control in hydrometallurgical processing”. In this excellent review, we read that iron is the second most abundant element in planet earth, after aluminium, and it is the fourth most abundant element in the earth’s solid crust. Iron is an essential component of the crystal structure of over 600 minerals, most of which are classed as oxide minerals (i.e. oxides, silicates, phosphates, arsenates, sulphates, hydroxides and so on) and relatively few are sulphide minerals. Some of this large number of minerals are considered to be ore minerals, but most are gangue minerals, as far as the mineral processor is concerned. Chen and Cabri listed the most important of the iron-bearing minerals in a series of tables in their paper, giving either their chemical formulae, or, where minerals form solid solutions, their compositional ranges. They classified the occurrence of iron in non-ferrous ores into three different modes: i. Ore minerals; ii. Gangue minerals; and iii. Solid solution minerals. i. Ore minerals: where iron is an essential constituent of a mineral used as a source of another metal, e.g. chalcopyrite, CuFeS2, a copper ore. Other examples are pentlandite, (Fe,Ni)9S8, a nickel ore, and ilmenite, FeTiO3, a titanium ore. ii. Gangue minerals: where iron is an essential constituent of a mineral that is undesirable in the ore, e.g. pyrite, FeS2, often the most abundant species in sulphide ores. Iron oxides, carbonates and silicates are common gangue minerals in many oxidised non-ferrous ore bodies. iii. Solid solution minerals: this is the substitution of iron for an essential element in the crystal structure of either an ore or a gangue mineral. The most common example of a solid solution ore mineral is sphalerite, (Zn,Fe)S, the principal source of zinc, where up to 17% of the zinc can be substituted by iron. Much of the early work on iron removal from hydrometallurgical solutions was centred on the treatment of zinc leach liquors and we shall 2 consider this work in some detail shortly. Other examples of solid solution ore minerals include tetrahedrite, (Cu,Fe,Zn)12(Sb,As)4S13, and cobaltite, (CoFe)AsS. Given that non-ferrous metals have been produced from ores for hundreds of years, or in the cases of copper and lead, for thousands of years, one might be forgiven for asking why, suddenly, has iron become the focus of so much attention by non-ferrous metallurgists? The simple answer is that it is due to the rise in importance of hydrometallurgy compared with pyrometallurgy for the production of non-ferrous metals. In the traditional high temperature smelting processes used for the production of copper, lead, zinc and other non-ferrous metals up to the beginning of the 20th century, iron in the smelter feeds ended up in the slags, together with many other deleterious impurity elements. Iron in silicate slags is present mainly as fayalite, ferrous silicate, Fe2SiO4. This is a chemically and environmentally inert material, which can be safely discarded without detrimental effects to the local surroundings. This benign method of discarding unwanted iron is one of the main advantages that pyrometallurgy has over hydrometallurgy. To understand the problems faced by hydrometallurgists who want to discard iron, we need to remind ourselves of the chemical behaviour of dissolved iron. The most concise way of doing this is to use an Eh-pH, or Pourbaix-type diagram, as shown in Figure 1. +1.0 Fe3+ Eh (volts) +0.5 Fe(OH)3 Fe2+ 0 -0.5 Fe(OH)2 Fe -1.0 0 2 4 6 8 10 12 14 pH Figure 1. Eh – pH diagram for the Fe – H2O system at 298°K This diagram, which illustrates the behaviour of the simple Fe-H2O system, shows that iron in the ferrous oxidation state (Fe2+) is soluble across a broad pH range from highly acidic to mildly alkaline, whereas ferric iron (Fe3+) is soluble only in acid solutions (pH<3) and oxidising conditions (Eh>0.77V). Outside these conditions, iron is insoluble and, in the absence of other anions, at ambient temperatures exists as ferric hydroxide, Fe(OH)3, or in alkaline reducing conditions, as ferrous hydroxide, Fe(OH)2. The relatively small area of solubility of ferric iron, up in the top left hand corner of the Pourbaix diagram, shows us that, theoretically, it is quite easy to get iron out of solution by oxidising it to the ferric state and making sure the pH is greater than about 3. This will swiftly cause iron to precipitate from solution as ferric hydroxide. Most other common base metals (Cu, Zn, Pb, etc) remain soluble under these conditions and so hydroxide precipitation provides in principle a ready method of separating dissolved iron from other more valuable metals in leach liquors. The precipitation of so-called ferric hydroxide is a very complex process, which even today is not fully understood and it was even less so, back in the 1960s. Ferric iron is a powerful complexant, 3 which will form complex ions with many anions, such as chloride, sulphate, fluoride, or phosphate, and the presence of different anions affects the form and composition of the hydrolysis product of ferric iron. Kinetic factors are also important, thus temperature, rate of change of pH and even the type of neutralising agent used, e.g. lime or sodium hydroxide, can all affect the outcome of hydrolysis. The usual result of rapid hydrolysis at ambient temperatures is the formation of a ferric hydroxide gel, with an open network structure, producing a soft, low density, material with a very high internal surface area.