CHAPTER I INTRODUCTION
1.1 Nickel
Nickel is mined from two types of ores, laterites and sulfides. Although about 70% of the ore reserves are found in laterite ores, only about 40% of the nickel production is from laterites. As shown in Figure 1, laterites are mostly used to produce ferronickel, which is used directly in steelmaking. Some laterite ores are used to make melting-grade nickel and nickel matte. Sulphides are refined to produce high-grade nickel.
Figure 1. The extraction of nickel and ferronickel from laterite and sulfide ores.
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1.2 Laterite Ores Laterite ores are a heterogeneous mixture of hydrated iron oxides and
hydrous magnesium silicates. These deposits were formed by weathering of peridotite rocks. Peridotite consists mainly of olivine [(Mg,Fe)2SiO4] with a small amount of pyroxene [(Mg,Fe) 2Si2O6]. Water, containing organic acids and carbon dioxide, percolates down through the weathered material. Iron, nickel, magnesium and silica dissolve in this water. Toward the top of the deposit, iron is oxidized by air and precipitates as hydrated iron oxides, such as goethite. Nickel and cobalt co precipitate with the iron, substituting for iron in the structure of goethite. The term ‘limonite’ is usually is usually used to describe this part of the laterite deposit. Closer to the bedrock, magnesium and silica precipitate, forming magnesium silicates of the serpentine group, such as Mg3Si2O5(OH)4.
Figure 2. An idealized profile of a laterite deposit.
Nickel precipitates as nepouite [Ni3Si2O5(OH)4]. Mixtures of these two minerals are called garnierite. Other minerals found in this layer are as follows: (i) talc [Mg3Si4O10(OH)2] and willemseite [(Ni,Mg)3Si4O10(OH)2]; (ii) clinochlore [(Mg,Fe)5Al(Si3Al)O10(OH)8] and nimite [(Ni,Mg,Al)6(Si,Al)4O10(OH)8] and (iii) sepiolite [Mg4Si6O15(OH)26H2O] and falcondite [(Ni,Mg) 4Si6O15(OH)2 6H2O]. This layer of the laterite is frequently referred to as saprolite. There may be a third
2
layer of clay material. These clays generally belong to the group of minerals called smectites. An example of a clay mineral found in nickel laterites is nontronite. Clays can be present in either the limonite or saprolite layers or may be present as a separate and distinct layer. An idealized profile of a laterite deposit is shown in Figure 2.It is emphasized that these layers are not necessarily distinct. Rather, there is a continuous variation with depth. The limonite layer is generally uniform, composed mainly of goethite. In contrast, the saprolit e layer is very heterogeneous. It is composed of a variety of silicates, such as serpentines, talcs, chlorites and sepiolites.
1.2.1 Laterite Ores Processing A schematic diagram of the flowsheet of the process is shown in Figure 3. The four main steps are the following: 1.
Dewatering: the removal of mechanically entrained water from the concentrate;
2.
Calcination: the removal of chemically bonded water from the dried ore;
3.
Reduction: the removal of oxygen from the nickel and iron oxide in the calcine; and,
4.
Refining: the removal of impurities, such as sulfur and phosphorus, from the molten ferronickel. Dewatering, or drying, is typically performed in a rotary kiln that i s 4
m diameter x 30 m long. The kiln is heated with pulverized coal, fuel oil or natural gas. The moisture of the dried ore is typically about 22%, although in some operations this is significantly lower. The dried material is either sent directly to the calcination kiln or first to a crushing and screening plant and then to a calcination kiln. The calcination kilns are much larger than the dewatering kilns, typically 5 m diameter by 100 m long. The purpose of the kilns is to produce a bone dry, partially reduced ore. The kiln is heated with pulverized coal, fuel oil or natural gas. The purpose of ferronickel smelting is to reduce the nickel and iron to metal and to reject the MgO, SiO2and Al2O3to slag. Ferronickel electric
3
furnaces are round or rectangular. Round furnaces are 15 – 20 m in diameter, while rectangular furnaces are 10 m wide and 25 – 35 m long. These furnaces produce 100 – 200 tonnes of ferronickel per day. Refining is performed in electrically heated ladles in batch mode in four steps: charging, phosphorus removal, sulfur removal and casting.
Figure 3. Schematic diagram of the flowsheet for smelting moist 1.5% – 2.5% Ni, low-iron laterite (saprolite) to ferronickel. 1.3 Industrial Applications of Nickel
The primary use of nickel is in the preparation of alloys such as stainless steel, which accounts for approximately 67% of all nickel used in manufacture. The greatest application of stainless steel is in the manufacturing of kitchen sinks but it has numerous other uses as well.
4
Other nickel alloys also have important applications. An alloy of nickel and copper for example is a component of the tubing used in the desalination of sea water. Nickel steel is used in the manufacture of armour plates and burglar proof vaults. Nickel alloys are especially valued for their strength, resistance to corrosion and in the case of stainless steel for example, aesthetic value. Electroplating is another major use of the metal. Nickel plating is used in protective coating of other metals. In wire form, nickel is used in pins, staples, jewellry and surgical wire. Finely divided nickel catalyses the hydrogenation of vegetable oils. Nickel is also used in the colouring of glass to which it gives a green hue. Other applications of nickel include: 1. Coinage 2. Transportation and construction 3. Petroleum industry 4. Machinery and household appliances 5. Chemical industry.
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CHAPTER II THERMODYNAMICS ASPECTS
2.1 Basic of Thermodynamics
Thermodynamics is a Greek word which means flow of heat in physical and chemical reactions. It is a branch of science which deals with study of different forms of energy and their interconversions. It deals with energy changes in physical and chemical processes . 2.1.1 System and Surrounding in Thermodynamics System
can be defined as the part of universe selected for
thermodynamic consideration i.e. to study the effect of temperature, pressure etc.It may also be defined as specified part of universe in which energy changes are taking place. Surrounding is The remaining portion of universe excluding the system. The System And Surrounding can be separated by real and imaginary boundary as shown in Figure 4.
Figure 4. System And Surrounding Types of system, there are: 1. Open System. Mass and energy can be exchanged with surroundings, eg. if some water is kept in open vessel or hot tea in open cup. 2. Closed System. In a closed system, there is only the exchange of energy with surroundings, no exchange of mass takes place. For example, if water is placed in closed metal vessel or hot tea placed in closed tea pot.
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3. Isolated System. There is neither exchange of mass nor energy with surrounding. Eg. Water placed in a vessel which is closed as well as insulated or tea placed in a thermos flask Classification of system on the basis of nature of constituents: 1. Homogenous System. All the constituents are present in the same phase and composition of system is uniform throughout 2. Heterogenous System. It contains two or more phases and the composition is uniform throughout. And the properties of system conditions are diveded into two, they are: 1. Intensive Properties. They do not depend on the size of the system or quantity of matter present in it. They are dependent on the nature of substance present in it. Example: pressure, temperature, density and surface tension. 2. Extensive Properties. They depend on the Quantity of matter present in the system. Examples: volume, energy, heat capacity and entropy. Furthermore, the thermodynamics processes are: 1. Reversible Process. Such a process is carried out infintesimally slowly so that all changes occuring in the direct process can be reversed and the system and the surrounding remain in state of equilibrium. It is an ideal process and cannot be realized in actual process. 2. Irreversible process. Change is brought about rapidly and the system does not attain equilibrium. The force which drives the reactants towards products is greater than opposing force which is to carry reverse process.
2.2 The Law of Thermodynamics
2.2.1 Zeroth Law of Thermodynamics
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There is a state function, called temperature which has the s ymbol T, which has the following relationship to heat, q : 1.
Addition of heat to a system will increase the temperature of the system.
2.
If two closed system (together isolated), with different temperatures are brought into thermal contact, then the temperatures of the two systems will change to approach the same temperature. That is, the temperature of the system which is at a higher temperature will decrease and the temperature of the system with the lower temperature will increase. They will eventually have the same temperature.
3.
The zeroth law leads to the general idea of heat capacity.
The
symbols C p and C v are used for this (constant pressure and constant volumn) but for solid there is usually little difference between these two. Using the relationship at constant volume (and therefore Cv ) between a change in temperature, ΔT , of a substance and the amount of heat transferred, q, to this substance is given by: q = Cv ΔT
(1)
2.2.2 First Law of Thermodynamics It is said that energy can neither be created nor destroyed in a system of constant mass, although it may be converted from one form to another. There is a state function, the internal energy E (in some texts U), which has the following properties: 1. In an isolated system E remains constant. 2. Addition of work, symbol w, to a closed system will increase the internal energy by the amount of work expended. This can be express by the following relation ship for a change in internal energy and work, w, done on a closed system: ΔE = q + w
(2)
If heat is transferred from the system (heat is negative) and if heat is transferred to the system (heat is positive). Internal energy increases when heat is added toa system or work is done on a system. Internal energy
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decreases when heat is removed from a system or work is done bya system.Use of internal energy or change in internal energy, ΔE , is not very convenient in chemistry. The reason for this is that when chemical reactions occur or samples are heated, the volume does not stay constant. If one is therefore interested in only q, the ΔE is complicated by an additional w. To avoid this a new quantity called enthalpy is defined, given the symbol H. H = E + PV
(3)
or ΔH = ΔE + PΔV
(4)
If no other external form of work is present, then: ΔH = w + q + PΔV
(5)
and ΔH = q
(6)
2.2.3 Second Law of Thermodynamics The second law of thermodynamics stated that the entropy of any isolated system not in thermal equilibrium almost always increases. There is a state function, entropy S, which has the following properties: 1. For a very small incremental addition of heat to a system, δq, one will obtain a very small increment of entropy (dS) according to the relationship: dS = δq/T, where T is the absolute temperature at the time and place of the heat transfer. 2. For an isolated system, any change over time in S is either positive or zero, that is: ΔS > or = 0 Another way of saying this is to assume one can add heat to a system in such a way as to not change the temperature. This might seem impossible but someone could be inside the system and balance the heat input with a chemical reaction that would take up the heat. Alternative system would be
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one in which a phase change, e.g.i ce to water, is taking place. In such a system the change in entropy would be: ΔS = δΔq /T
(7)
For those who have calculus in your future, an increment of entropy designated by d S is related to a small increment of added heat, dq, by: dS = δq /T
(8)
where dS is now an exact differential, but δq is not. Thus 1/T is the integrating factor. If there is no net change in the state inside the isolated system then ΔS =0. This then is the thermodynamic criterion for equilibrium. Inside an isolated system, in order for a process to proceed, Δ S > 0. Such a process is said to be spontaneous. A process for which ΔS < 0 is called non-spontaneous and
is
impossible
for
an
isolated
system.
Mathematically one can derive the following conclusion for a closed system with movable boundaries to keep the internal pressure constant. To do this, a new state function is defined which combines the entropy with enthalpy. This is the Gibbs' free energy (G) defined by: ΔG = ΔH - T ΔS
(9)
Table 1. The Criteria for Equilibrium and Spontaneity. Condition
For An Isolated
For A Closed System at
System
Constant Pressure
Spontaneous process
ΔS > 0
ΔG < 0
Equilibrium
ΔS = 0
ΔG = 0
Non Spontaneous process
Impossible
ΔG > 0
2.2.4 Third Law of Thermodynamics Nernst stated that at 0 K the entropy increment of reversible reactions among perfect crystalline solids is zero. And also Planck stated that the entropies of all perfect crystalline solids at 0 K are zero. And the
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last, Lewis & Randall stated that if the entropy of each element in some (perfect) crystalline state be taken as zero at the absolute zero of temperature, every substance has a finite positive entropy, but at the absolute zero of temperature the entropy may become zero, and does so become in the case of perfect crystalline substance. The important point about the third law is that entropy is an absolute quantity which depends upon temperature. This is in contrast to ΔH for reactions which have as a reference the elemental state. Thus, when one looks up the ΔHof of an elements, the answer is 0. In contrast, So for an element (note difference in symbols as well) has a value for temperature above 0 K. The entropy change with respect to temperature can be thought of a continuous summation of all the increments of heat added to the system divided by the temperature at the time of the addition. Or symbolically: ΔS =
(dq/T) dT
(10)
Thus, to calculate a change in S one simply adds up the little increments of heat added divided by temperature. For a pure component in the most stable condition, S = 0 at T = 0 K. This leads to the assumption needed above, that the S o for pure components are absolute values and are not referenced against some arbitrary initial condition like the ΔH o are. 2.3 Ellingham Diagram
The Gibbs free energy (∆G) of a reaction is a measure of the thermodynamic driving force that makes a reaction occur. A negative value for ∆G indicates that a reaction can proceed spontaneously without external inputs, while a positive value indicates that it will not. The enthalpy (∆H) is a measure of the actual energy that is liberated when the reaction occurs (the “heat of reaction”). If it is negative, then the reaction gives off energy, while if it is positive the reaction requires energy. The entropy (∆S) is a measure of the change in the possibilities for disorder in the products compared to the reactants. For example, if a solid (an ordered state) reacts with a liquid (a somewhat less
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ordered state) to form a gas (a highly disordered state), there is normally a large positive change in the entropy for the reaction. An Ellingham diagram is a plot of ∆G versus temperature. Since ∆H and ∆S are essentially constant with temperature unless a phase change occurs, the free energy versus temperature plot can be drawn as a series of straight lines, where ∆S is the slope and ∆H is the y-intercept. The slope of the line changes when any of the materials involved melt or vaporize. Free energy of formation is negative for most metal oxides, and so the diagram is drawn with ∆G = 0 at the top of the diagram, and the values of ∆G shown are all negative numbers. Temperatures where either the metal or oxide melt or vaporize are marked on the diagram. The Ellingham diagram shown is for metals reacting to form oxides (similar diagrams can also be drawn for metals reacting with sulfur, chlorine, etc., but the oxide form of the diagram is most common). The oxygen partial pressure is taken as 1 atmosphere, and all of the reactions are normalized to consume one mole of O 2. The majority of the lines slope upwards, because both the metal and the oxide are present as condensed phases (solid or liquid). The reactions are therefore reacting a gas with a condensed phase to make another condensed phase, which reduces the entropy. A notable exception to this is the oxidation of solid carbon. The line for the reaction C + O2
CO2
is a solid reacting with a mole of gas to produce a mole of gas, and so there is little change in entropy and the line is nearlyhorizontal. For the reaction 2C + O2
2CO
we have a solid reacting with a gas to produce two moles of gas, and so there is a substantial increase in entropy and the line slopes rather sharply downward.
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Similar behavior can be seen in parts of the lines for lead and lithium, both of which have oxides that boil at slightly lower temperatures than the metal does.
Figure 5. Ellingham Diagram There are three main uses of the Ellingham diagram: 1. Determine the relative ease of reducing a given metal oxide to metal. The position of the line for a given reaction on the Ellingham diagram shows the stability of the oxide as a function of temperature. Reactions closer to the top of the diagram are the most noble metals (for example, gold and
13
platinum),and their oxides are unstable and easily reduced. As we move down toward the bottom of the diagram, the metals become progressively more reactive and their oxides become harder to reduce. A given metal can reduce the oxides of all other metals whose lines lie above theirs on the diagram. For example, the 2Mg + O 2 below the Ti + O2
2MgO line lies
TiO2 line, and so magnesium can reduce titanium oxide
to metal titanium. Since the 2C + O 2
2CO line is downward-sloping, it
cuts across the lines for many of the other metals. This makes carbon unusually usefull as a reducing agent, because as soon as the carbon oxidation line goes below a metal oxidation line, the carbon can then reduce the metal oxide to metal. So, for example, solid carbon can reduce chromium oxide once the temperature exceeds approximately 1225°C, and can even reduce highlystable compounds like silicon dioxide and titanium dioxide at temperatures above about 1620°C and 1650°C, respectively. For less stable oxides, carbon monoxide is often an adequate reducing agent. 2. Determine the partial pressure of oxygen that is in equilibrium with a metal oxide at a given temperature. The scale on the right side of the diagram labelled P O2 is used to determine what partial pressure of oxygen will be in equilibrium with the metal and metal oxide at a given temperature. The significance of this is that, if the oxygen partial pressure is higher than the equilibrium value, the metal will be oxidized, and ifit is lower than the equilibrium value then the oxide will be reduced. To use this scale, you will need a straightedge. First, find the temperature you are interested in, and find the point where the oxidation line of interest crosses that temperature. Then, line up the straightedge with both that point, and with the point labelled “0” that is marked with short radiating lines (upper left corner of the diagram). Now, with the straightedge running through these two points, read off the oxygen partial pressure (in atmospheres) where the straightedge crosses the PO2 scale, and this is the
14
equilibrium partial pressure.It is possible to reach the equilibrium oxygen partial pressure by use of a hard vacuum, purging with an inert gas to displace the oxygen, or using a scavenger chemical to consume the oxygen. 3. Determine the ratio of carbon monoxide to carbon dioxide that will be able to reduce the oxide to metal at a given temperature. When using carbon as a reducing agent, there will be a mi nimum ratio of CO to CO2 that will be able to reduce a given oxide. The harder the oxide is to reduce,the greater the proportion of CO needed in t he gases.To determine the CO/CO2 ratio to reduce a metal oxide at a particular temperature, use the same procedure as for determining the equilibrium pressure of oxygen, except line up the straightedge with the point marked C (center of the left side of the diagram), and read the ratio off of the scale marked CO/CO 2.
2.4 Thermodynamics of Nickel Smelting
2.4.1 Rotary Kilns for Calcination and Reduction The calcination of laterite ores is performed in rotary kilns. The kilns are long, up to 185 m and are fired b y hydrocarbon fuels. The kilns are typically sloped about 4° from horizontal and rotate at approximately 1 rotation/minute. Upgraded and dewatered ore is fed continuously to the upper end of the rotary kiln. Coal is also added continuously along with recycled dust from the calcination kiln that has been pelletized. The feed rate of coal is about 5% of the feed rate of upgraded ore. These materials flow slowly down the rotary kiln and out of the hot discharge end of the kiln. Reducing gas is continuously supplied to the rotary kiln by partiall y combusting the hydrocarbon fuel at the discharge end of the kiln. This combustion produces hot reducing gas which (i) heats and dries the feed; and, (ii) partially reduces the nickel oxide and iron minerals in the feed as it travels in a counter-current direction up the kiln. The reactions that occur in the rotary kiln are as follows: 1.
Evaporation of the remaining mechanically entrained water in the ore, given as follows:
15
{H2O}
(H2O)
The evaporation of water is endothermic. 2.
Thermal dissociation of laterite minerals to oxides and H2O(g), given by reactions such as the following: 700°C
3 + 3 + 4 + 4(H2O)
garnierite 700°C
2
+ (H2O)
geothite
These reactions are endothermic. 3.
Reduction of the resulting oxides by reducing gas by the following reactions. 800°C
+ (CO)
+ (CO2) ΔG°1073 = -15,42 kcal ΔG° = -RT ln ( PCO2/PCO) 800°C
+ (CO)
2 + (CO2)
ΔG° = -RT ln ( PCO2/PCO) About 20% of NiO is indirectly reduced to Ni. At the same time that these reactions occur, the solid products are heated to approximately 900°C for hot charging to an electric furnace. 2.4.2 Smelting in Electric Furnace The product of the calcination operation, which was described is bone dry and partially reduced calcine at 900°C. Ferronickel smelting makes this calcine into hot, molten ferronickel that is suitable for refining and use in making stainless steel and other ferrous alloys. The calcine is fed continuously into an electric furnace where it is reduced and melted. The reactions that occur during ferronickel smelting are the following: 1. The reduction of the nickel oxide in the calcine to nickel, which is given as:
16
When NiO and FeO are reduced by C (these reaction occur very slow). 900°C
900°C
1450°C
+ 900°C
1450°C
{Ni} + (CO)
900°C
1450°C
2 +
ΔG°1723 = -263 Kj
1450°C
2{Ni} + (CO2)
ΔG°1723 = -180 Kj
At 1723K, NiO can be reduced by C when CO is more stable than CO 2. So, actually the reaction will produce CO2 first, but because CO is more stable than CO 2 hence CO2 will be reduced by CO. That’s why that the reaction will produce CO insted of CO2.
Figure 6. Free energy for the formation of oxides
900°C
900°C
+
1450°C
1450°C
{Fe} + (CO)
SiO2 and MgO are also reduced by CO which is given as: 900°C
900°C
+ 2
1450°C
1450°C
{Si} + 2(CO2)
ΔG° = -RT ln ( PCO2/PCO) 900°C
900°C
+ 2
1450°C
1450°C
{Mg} + 2(CO2)
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ΔG° = -RT ln ( PCO2/PCO) CO which is produced by the reactions above and also the CO produced by boudouard reaction, will indirectly reduce NiO to Ni and FeO to Fe. Boudouard reaction
(CO2) + (C)
2(CO)
ΔH° = 172,46 kJ/mol
Hence, NiO and FeO will be indirectly reduced by CO. Because the reduction occurs with C as the reductant agent is very slow, so that the reductant agent will be CO. +
{Ni} + (CO2)
ΔG° = -RT ln ( PCO2/PCO) +
{Fe} + (CO2)
ΔG° = -RT ln ( PCO2/PCO) 2. The melting and alloying of nickel and iron to form molten ferronickel by the following reaction: 1450°C
1450°C
{Ni} + {Fe}
1450°C
{Ni, Fe}
2.5 Study Case
Calculate the equilibrium PCO2/ PCO ratio for the reaction: + (CO)
+ (CO2)
At 800°C from the following data: + (CO)
+ (CO2)
ΔG° = 81.150 – 2,25T cal
Solution: At 800°C (1073K), ΔG°1073= 81.150 – 2,25 (1073) cal ΔG°1073 = 81.150 – 2.736,15 cal ΔG°1073 = 78.413,85 cal
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At equilibrium, ΔG° = -RT ln K 78.413,85 = -1,987 (1073) ln K ln K = 78.413,85 -2.132,051 ln K = -36,77 K
= 1,065.10-16
Since
K = PCO2 . a Ni P
CO . a NiO
and a Ni and a NiO are unity, hence K = PCO2 P
CO
So that, 1,065.10-16= PCO2 P
CO
Thus, the equilibrium PCO2/ PCO ratio for reaction at 800°C is 1,065.10 -16.
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CHAPTER III CONCLUSION
Nickel is mined from two types of ores, laterites and sulfides. Laterite ores are a heterogeneous mixture of hydrated iron oxides and hydrous magnesium silicates. Laterite ores processing are dewatering (the removal of mechanically entrained water from the concentrate), calcination (the removal of chemically bonded water from the dried ore), reduction (the removal of oxygen from the nickel and iron oxide in the calcine) and refining (the removal of impurities, such as sulfur and phosphorus, from the molten ferronickel). The calcination of laterite ores is performed in rotary kilns. Energy and reducing gas are continuously supplied to the rotary kiln by partially combusting the hydrocarbon fuel at the discharge end of the kiln. This combustion produces hot reducing gas which (i) heats and dries the feed; and, (ii) partially reduces the nickel and iron minerals in the feed as it travels in a counter-current direction up the kiln. The product of the calcination operation, which was described in is bone dry, partially reduced calcine at 900°C. Ferronickel smelting makes this calcine into hot, molten ferronickel that is suitable for refining and use in making stainless steel and other ferrous alloys. The complete reduction will occur in the electric furnace. The reduction occurs 80% directly and 20% indirectly reduction up to 1650°C.
Ni3S2
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