Reduction of Metals

Submitted by admin on Tue, 01/18/2011 - 02:16

The ease with which a metalAn element characterized by a glossy surface, high thermal and electrical conductivity, malleability, and ductility. may be obtained from its oreNaturally occurring minerals from which metals and other substances can be extracted commercially. varies considerably from one metal to another. Since the majority of ores are oxides or can be converted to oxides by roastingHeating in the presence of air or oxygen; most commonly used in converting natural metal sulfide ores to oxides as first step in recovery of the metal from its ore., the free-energyA system's capacity to do work. change accompanying the decomposition of the oxide forms a convenient measure of how readily a metal may be obtained from its ore. Values of the free energyGibbs energy: a thermodynamic function corresponding to the tendency for spontaneous change in a system; represented by the symbol G. change per mol O2 produced are given in the table for a representative sample of metals at 298 and 2000 K. A high positive value of ΔGm° in this table indicates a very stable oxide from which it is difficult to remove the oxygen and obtain the metal, while a negative value of ΔGm° indicates an oxide which will spontaneously decompose into its elements. Note how the value of ΔGm° decreases with temperatureA physical property that indicates whether one object can transfer thermal energy to another object. in each case. This is because a gasA state of matter in which a substance occupies the full volume of its container and changes shape to match the shape of the container. In a gas the distance between particles is much greater than the diameters of the particles themselves; hence the distances between particles can change as necessary so that the matter uniformly occupies its container. (oxygen) is produced by the decomposition, and ΔS is accordingly positive.


Free-Energy Changes for the Decomposition of Various Oxides at 298 and 2000 K.

Reaction ΔGm° (298 K)/kJ mol–1 ΔGm° (2000 K)/kJ mol–1
\tfrac{2}{3}Al2O3\to\tfrac{4}{3}Al + O2
+1054
+691
2MgO → 2Mg + O2
+1138
+643
\tfrac{2}{3}Fe2O3\to\tfrac{4}{3}Fe + O2
+744
+314
SnO2 → Sn + O2
+520
+42
2HgO → 2Hg + O2
+118
–381
2Ag2O → 4Ag + O2
+22
–331
Formation of CO2
C(s) + O2(g) → CO2(g)
–394
–396


The two metals in the table which are easiest to obtain from their oxide ores are Hg and Ag. Since the ΔGm° value for the decomposition of these oxides becomes negative when the temperature is raised, simple heating will cause them to break up into O2 and the metal. The next easiest metals to obtain are Sn and Fe. These can be reduced by coke, an impure form of C obtained by heating coal. Coke is the cheapest readily obtainable reducing agentA chemical species that donates electrons in order to reduce another species. In the process the reducing agent is itself oxidized. which can be used in metallurgyThe study of metals and their extraction from ores.. When C is oxidized to CO2, the free-energy change is close to – 395 kJ mol–1 over a wide range of temperatures. This fall in free energy is not quite enough to offset the free-energy rise when Fe2O3 and SnO2 are decomposed at 298 K, but is more than enough if the temperature is 2000 K. Thus, for example, if Fe2O3 is reduced by C at 2000 K, we have, from Hess’ law,


{}_{\text{3}}^{\text{2}}\text{Fe}_{\text{2}}\text{O}_{\text{3}}\text{(}s\text{) }\to \text{ }{}_{\text{3}}^{\text{4}}\text{Fe(}l\text{) + O}_{\text{2}}\text{(}g\text{)}                   ΔGm° = +314 kJ mol–1


\text{C(}s\text{) + O}_{\text{2}}\text{(}g\text{)}\to \text{ CO}_{\text{2}}\text{(}g\text{)}                             ΔGm° = –394 kJ mol–1
                                                                                                                    

{}_{\text{3}}^{\text{2}}\text{Fe}_{\text{2}}\text{O}_{\text{3}}\text{(}s\text{) + C(}s\text{) }\to \text{ }{}_{\text{3}}^{\text{4}}\text{Fe(}l\text{) + CO}_{\text{2}}      ΔGm° = –82 kJ mol–1


Thus ΔGm° for the reductionThat part of a chemical reaction in which a reactant gains electrons; simultaneous oxidation of a reactant must occur. is negative, and the reaction is spontaneous.

The two metals in the table which are most difficult to obtain from their ores are Mg and Al. Since they cannot be reduced by C or any other readily available cheap reducing agent, they must be reduced electrolytically. The electrolytic reduction of bauxite to yield Al (the Hall process) is used to produce aluminum.


Reduction of Iron
Schematic Diagram of a Blast Furnace for Smelting Iron 1. Hot blast from Cowper stoves 2. Melting zone 3. Reduction zone of ferrous oxide 4. Reduction zone of ferric oxide 5. Pre-heating zone 6. Feed of ore, limestone and coke 7. Exhaust gases 8. Column of ore, coke and limestone 9. Removal of slag 10. Tapping of molten pig iron 11. Collection of waste gases.

Since iron is the most important metal in our industrial civilization, its reduction from iron ore in a blast furnace (Fig. 1) deserves a detailed description. The oxides present in most iron ores are Fe2O3 and Fe3O4. These oxides are reduced stepwise: first to FeO and then to Fe. Ore, coke, and limestone are charged to the furnace through an airlock-type pair of valves at the top. Near the bottom a blast of air, preheated to 900 to 1000 K, enters through blowpipes called tuyères. Oxygen from the air blast reacts with carbon in the coke to form carbon monoxide and carbon dioxide, releasing considerable heat. The blast carries these gases up through the ore, coke, and limestone, and they exit from the top of the furnace.

By the time the ore works its way to the lower part of the furnace, most of the Fe2O3 has already been reduced to FeO. In this region, temperatures reach 1600 to 2000 K, high enough to melt FeO and bring it into close contact with the coke. Most of the FeO is reduced by direct reaction with carbon, the latter being oxidized to carbon monoxide:


2C(s) + 2FeO(l) → 2Fe(l) + 2COΔGm° (2000 K) = –280 kJ mol–1      (1)


The molten iron produced by this reaction drips to the bottom of the furnace where it is collected and occasionally tapped off.

Higher in the furnace, temperatures fall below the melting points of the iron oxides. Because there is little contact between solidA state of matter having a specific shape and volume and in which the particles do not readily change their relative positions. chunks of ore and of coke, direct reduction by solid carbon is rather slow. Gaseous carbon monoxide contacts all parts of the ore, however, and reacts much more rapidly:


CO(g) + Fe2O3(s) → CO2(g) + 2FeO(s)      (2)

CO(g) + FeO(s) → CO2(g) + Fe(s)      (3)


Thus much of the “carbon reduction” in ironmaking is actually carried out by carbon monoxide.


The gangue in iron ore consists mainly of silicates and silica, SiO2. These impurities are removed in slag. Limestone added with coke and ore is calcined (decomposed to the oxide) by the high temperatures of the blast furnace:


CaCO3(s) \xrightarrow{\text{1100 K}} CaO(s) + CO2(g)


Lime (CaO) serves as a flux, reducing the melting points (mp) of silica (SiO2) and silicates:


\underset{\text{mp = 2853 K}}{\mathop{\text{CaO(}s\text{)}}}\,\text{ + }\underset{\text{mp = 1986 K}}{\mathop{\text{SiO}_{\text{2}}\text{(}s\text{)}}}\,\text{ }\to \text{ }\underset{\text{mp = 1813 K}}{\mathop{\text{CaSiO}_{\text{3}}\text{(}l\text{)}}}\,


The liquidA state of matter in which the atomic-scale particles remain close together but are able to change their positions so that the matter takes the shape of its container silicates flow rapidly down through the hottest part of the furnace. This helps to prevent reduction of silica to silicon, hence yielding purer iron. The slag is less dense than molten iron and immiscibleDescribes two liquids that cannot be dissolved in one another in all proportions. with it. Therefore the slag floats on the surface of the iron and can easily be tapped off.

Although most blast-furnace iron now goes directly to a steelmaking furnace in molten form, much of it used to be run into molds where it hardened into small ingots called pigs because of their shape. Consequently blast-furnace iron is still referred to as pig iron. A single large blast furnace may produce more than 106 kg iron per day. For each kilogramThe SI unit for mass. of iron, 2 kg iron ore, 1 kg coke, 0.3 kg limestone, 4 kg air, 63 kg water, and 19 MJ of fossil-fuel energy are required. The furnace produces 0.6 kg slag and 5.7 kg, flue gas per kg iron. Nearly 5 percent of the iron ore is lost in the form of small particles suspended in the flue gas unless, as in the furnace shown in Fig. 1, air-pollutionThe contamination of the air, water, and earth by personal, industrial, and farm waste. controls are installed. The latter trap FeO particles for recycling to the furnace and also make the flue gas (which contains about 12% CO and 1% H2) suitable as a fuel for preheating air fed to the tuyères. Thus control of blast-furnace air pollution (a major contributor to the one-time “smoky city” reputations of Pittsburgh, Pennsylvania and Gary, Indiana) also conserves ore supplies and energy resources.