Aluminum: A Modern Nontransition Metal

James Richard Fromm


If, in the history of mankind, the ages may be classified by a characteristic metal, modern history would have to be called the Aluminum Age. From beer cans to jetliners, much of our modern technology employs this light metal.

Aluminum compounds, primarily the oxide in forms of various purity and hydration, are fairly widely distributed in nature. The feldspars, the most common rock-forming silicates, make up nearly 54% of the earth's crust; in these, aluminum has replaced up to half the silica atoms in SiO2. The major ore of aluminum is bauxite, a hydrated aluminum(III) oxide (Al2O3.xH2O). It normally contains significant quantities of iron as an impurity. This iron can be removed by the Bayer process, which is treatment with a hot strong aqueous NaOH solution:

Al2O3.xH2O + 2OH-(aq) 2AlO2-(aq) + (x + 1)H2O

The oxide Fe2O3 and other salts of Fe(III) react to precipitate solid Fe(OH)3 and iron silicates. The solution containing AlO2- is then filtered, cooled, and acidified with CO2 to precipitate the solid Al(OH)3:

2AlO2-(aq) + CO2(g) + 3H2O 2Al(OH)3 + CO32-(aq)

Under strong heating (calcining), Al(OH)3 loses water at 1000oC to produce a pure aluminum(III) oxide (alumina):

2Al(OH)3(c) Al2O3(c) + 3H2O(g)

Reduction of aluminum oxide with carbon is not possible, since the free energy of reaction is so highly unfavorable that it cannot be overcome.

2Al2O3 + 3C 3CO2 + 4Al; DG0 = +399.2 kJ/mole

The reduction of Al2O3 with carbon is unfavorable by 399.2 kJ/mole while the analogous reduction of Fe2O3 is unfavorable by only 301.3 kJ/mole. Reduction with carbon monoxide, which is possible for Fe2O3 where DG0 is -29.4 kJ/mole, is not possible for Al2O3:

Al2O3 + 3CO 3CO2 + 2Al; DG0 = +810.7 kJ/mole

Chemical Methods of Aluminum Production

Prior to around 1720, the probable existence of the element aluminum had been recognized in its compounds, chiefly alumina (Al2O3), which is referred to as argill by Lavoisier (1790). The reduction of alumina, like reduction of magnesia (MgO), baryta (BaO), and others, had been attempted using carbon but had never been successful -- as we now know, because the free energy for the reduction is unfavorable as shown above. The oxide can be converted to the chloride, AlCl3, which is a volatile white solid, by treatment with chlorine. Reduction of AlCl3, like reduction of Al2O3, is not possible by use of carbon.

The first impure metallic aluminum powder was prepared by H.C. Oersted in 1824 by reduction of AlCl3 with sodium,

AlCl3 + 3Na Al + 3NaCl,

or reduction using the amalgam of potassium, a solution of potassium in mercury, according to the reaction

AlCl3 + 3K(Hg) Al + 3KCl + Hg.

Such an amalgam can be prepared by electrolysis of solutions of potassium chloride. The electrolysis of sodium and potassium salts was also studied by Humphrey Davy, who produced metallic sodium in 1807. In 1827, Friedrich Woehler used a similar reaction in a sealed vessel to prepare the first pure aluminum:

AlCl3(g) + 3K(g) Al(s) + 3KCl(s).

The lightness and malleability of aluminum had been demonstrated by 1845. Although some commercial preparation was done, the cost of some $140/lb was prohibitive.

In 1854-1855, Henri-Etienne Saint-Claire Deville, under the patronage of Emperor Napoleon III of France, began the true commercial production of aluminum. His process began with treatment of bauxite to yield pure alumina, which is the same as the present commercial Bayer process. This was followed by conversion to aluminum chloride:

2Al2O3 + 6Cl2(g) + 3C 4AlCl3 + 3CO2, DG0 = -533.7 kJ/mole

The reduction step was

AlCl3 + Na 3NaCl(s) + Al, DG0 = -523.6 kJ/mole

The sodium used was produced from sodium carbonate by the reactions:

Na2CO3 Na2O + CO2, DG0 = +668.98 kJ/mole

2Na2O + C CO2 + 4Na, DG0 = +356.56 kJ/mole

Na2O + CO CO2 + 2Na, DH0 = +131.24 kJ/mol

In the presence of water, there is also the reaction

2NaOH + C 2Na + CO2 + H2,

where

Na2O + H2O 2NaOH

easily. All of these sodium-producing reactions are energetically unfavorable but can be forced to proceed by supplying enough external heat from

C + O2 CO2,

which yields -394.359 kJ/mole, and by continuous removal of the carbon dioxide. The usual temperature of operation is 1000oC. By these procedures aluminum could be produced and sold for around $5.00 per pound at the best efficiencies obtained. The sodium could also be produced by electrolysis of molten NaCl but this was not a practical procedure at the time.

On a modern manufacturing scale, these processes were small; Deville's production was perhaps one tonne/year while others, using the same process, produced up to one tonne/month. The most costly material required by Deville was sodium metal, which he obtained by the process of reducing liquid sodium hydroxide-sodium oxide mixtures with carbon at 1000oC:

NaOH-Na2O(l) + C NaOH + CO2 + Na(g).

This process is very slow and inefficient since carbon is much lighter than molten sodium hydroxide and simply floats on top. The cost of the sodium Deville produced by this method was about $1.75 per pound,

Improvements in the production of sodium were made by the American chemical engineer Hamilton Y. Castner (1858-1899). The idea was very simple: weight down the carbon with iron filings so it sinks into the liquid sodium hydroxide-sodium oxide mixture. This could be done by mixing iron with pitch, then coking the pitch to carbon, and throwing in the carbon so weighted. This simple improvement reduced the cost of manufacturing sodium from $1.75 to $0.09 per pound.

With the price of sodium thus drastically reduced, Castner together with English financial backers built a large plant for sodium production. The sodium so produced was intended solely for the manufacture of aluminum, since no other significant market for that much sodium existed. Unfortunately, the development of the Hall-Heroult process described in the sections on applied electrochemistry came just as the plant was completed. Since hydroelectric power was not available at Castner's plant and since the Hall-Heroult process was even cheaper than his own, he could not compete and thus had the choice of creating new major markets for sodium or go bankrupt.

With characteristic ingenuity, Castner set out to develop new processes and new markets for sodium. One of these was sodium peroxide, in some demand as a bleaching material, which could be made by reaction of oxygen from air with the sodium:

Na + O2(g) Na2O2(s).

Another was sodium amide,

2NH3(g) + 2Na(l) 2NaNH2(l) + H2,

whose production takes place on heating ammonia in contact with sodium. Sodium amide is a solid at room temperature but is molten at high temperatures, where it reacts with carbon to produce sodium cyanide. The reaction proceeds via a hypothetical intermediate of

NaH2CN: NaNH2(l) + C [NaH2CN] NaCN(s) + H2(g).

Since cyanide had a major market in the gold-mining industry economic success was assured. For some time Castner's sodium cyanide was marketed as "130% potassium cyanide" since the potassium salt, produced by a different route,was already well known in the mining industry. In 1894, Castner discovered thatthe production of sodium cyanide could be carried out directly using the reaction

2Na + 2NH3 + 2C 2NaCN + 3H2.

STUDY PROBLEMS

A certain sample of galena (PbS) was found to be 2.3% argentite (Ag2S) by weight. Compute the mass of galena which would have to be processed to obtain one metric tonne (Mg) of silver from this ore. Compute also the number of cubic kilometers of sulfur dioxide produced in the course of obtaining this silver at 25oC and 1 atm pressure. Sulfur dioxide affects plants at a concentration or level of about 1 ppm; how many cubic kilometers of the atmosphere would be contaminated to this level?

A chemical engineer is designing a plant to produce 20 tonnes/day of sulfuric acid by:first burning sulfur to sulfur dioxide in air, then converting the sulfur dioxide to sulfur trioxide by heating it to 450oC in air over a catalyst, and finally reacting the sulfur trioxide with water to give the acid.
(a) Write the complete balanced equations for each step. In the first two steps,indicate what is oxidized and what is reduced. Is the third step an oxidation? Explain why or why not.
(b) How many tonnes of sulfur will this plant use each day?

Write the complete chemical reactions for the production of metallic copper from malachite,

CuCO3.Cu(OH)2. How much charcoal is necessary to smelt one tonne (Mg) of copper from malachite ore?

Calculate the percentage composition by mass of the mineral malachite, CuCO3.Cu(OH)2.

The Canadian export price for lead in 1968 was $222/ton.If your company processed 1,000 tonnes of lead ore which was 70% galena by mass, what would you have been paid for the lead produced? Galena is the geological name for lead(II) sulfide, PbS.

Compute the net enthalpy change in the reaction of one mole of magnetite, Fe3O4, with carbon and oxygen to form iron metal.


Copyright 1997 James R. Fromm