Nice Job Halfapint!!!
This is the shit for home phos production.
Here is some info.
Reduction of metallic oxide to metal
Reduction
The conversion of metal oxide into metal (by removal of oxygen) is called reduction.
Generally the 3 methods used are:
Reduction by heating the oxide
Chemical reduction
Electrolytic reduction.
Chemical Reduction Processes
Under this process the oxides of metals that are in the middle of the reactivity series are reduced to free metals using chemical reducing agents such as carbon, aluminium, sodium or calcium.
1. Reduction by carbon process
Oxides of metals like zinc, iron, copper, nickel, tin and lead are reduced using carbon as the reducing agent.
(note: phosphorus is done this way) Carbon can be used only if it has greater affinity for oxygen than the metal. For example, carbon can reduce copper oxide to copper, but it cannot reduce calcium oxide. It can reduce zinc oxide.
ZnO + C -----> Zn + CO
Zinc metal Carbon
monooxide
2. Reduction with aluminium by thermite process
Metals which are too active to be obtained by reduction of their oxides with carbon are reduced using aluminium, which is a more powerful reducing agent.(note: It will reduce Phosphorus!!!) Chromium and manganese oxides are reduced using aluminum. This reaction is highly exothermic.
Cr2O3 + 2Al -----> 2Cr + Al2O3
Chromium metal
3MnO2 + 4Al -----> 3Mn + 2Al2O3
Manganese metal
……………………………………………………………………………………………
http://www.compassbox.com/cbse/demo/chemistry/chapter10/3704061004.asp
And More.
Thermite reaction.
Soon after metallic aluminium was first isolated, both Sainte-Claire Deville and Wöhler noted its property when mixed in the form of powder or granules with the oxides of some of the metals, of reducing them with the evolution of sufficient heat to melt the metal and the alumina produced [1].
On this was based the process applied by Dr Hans Goldschmidt of Essen, who in 1898 was developing commercially the promising work of Claude Vauntin and Hugh Picard in London, following the change in the economics of aluminium supply which had taken place when production by the Héroult and Hall electric furnace method superseded the expensive sodium reduction process; prior to that time the use of aluminium as a reducing agent was little more than a scientific curiosity [2]. Goldschmidt has obtained either the metal, or an alloy of the metal, with aluminium, from the oxides of chromium, manganese, iron, copper, titanium, boron, tungsten, molybdenum, nickel, cobalt, zirconium, vanadium, niobium, tantalum, cerium, thorium, barium, calcium, sodium, potassium, lead and tin. He said:
In a thermite reaction, a metallic compound is reduced by one of several metals or metallic alloys in such a way that when the mixture is ignited at one place, the reaction continues of its own accord, so that under complete oxidation of the reducing element, a fluid slag is formed, while the reduced metal is obtained as a homogeneous uniform regulus; if the oxide is used in excess, the reduced metal is free, or practically free, from the element used as a reducing agent.
Goldschmidt’s significant innovations comprised the means of starting the reaction by a fuse, instead of heating the mixture until ignition took place, and the various controlling procedures which made the process practicable on a plant scale.
This is the first of two solid/solid reactions, see also experiment 18a. This reaction, uses powdered aluminium metal to reduce metallic oxides and is especially useful for the reduction of those metal oxides which are difficult to reduce such as titanium and molybdenum. It is also known as the Goldschmidt process, and the Aluminothermic process. This is a highly exothermic reaction and the metal emerges in its molten state often very much above its melting point. A mixture of finely divided aluminium with ferric oxide (or the oxide of some other metal), approximating to 2Al+Fe2O3, is sold under the registered name Thermit (no final ‘e’), and is used in joining or welding iron and steel rails, pipes, etc. Sections of railway lines are usually welded in this way and the reaction is mostly associated with the reduction of iron oxide as in our demonstration. Like the first part of the next demonstration this reaction takes place between solids and is, consequently, difficult to start.
Two equivalents of finely divided aluminium and one of iron oxide are mixed together and ignited. The reaction takes place as follows:
Fe2O3 + 2 Al ==> Al2O3 + 2 Fe + 848.54 kJ (17.1)
Safety. This is a hazardous procedure, but is a safe demonstration in experienced hands. However, there must always be a first time, and it is to those persons that we offer the following advice. Read the safety recommendations of say reference 7 below.
First perform this experiment outdoors a couple of times, preferably use a table as you would indoors. There are very few demonstrations that can safely be performed for the first time in front of an audience and, in this case, we recommend that you practice at least twice more indoors before regarding yourself as sufficiently experienced and competent. It is always a good idea to get someone else to work with you as an equal partner, in that you should both be equally familiar with the procedure, and both sense the burden of responsibility. In these circumstances the partners should assume individual and collective responsibility and so will ‘watch out for each other’. For example a teacher might involve his technician in this way.
It is important that the Thermite mixture is perfectly dry otherwise extremely hot material might be violently ejected from the crucible. Dry the components separately at ca. 125C. Do not use heat to dry the mixture use a desiccant.
Metallic iron is produced at a temperature significantly above its melting point at 1535C. Once the reaction has started it is almost impossible to stop it. A CO2 fire extinguisher should to hand; water should not be used because potentially explosive hydrogen can be produced. No one should be closer to the reaction than two metres.
References
1. Thorpe’s Dictionary of Applied Chemistry, 4th ed., vol. XI, London - New York - Toronto, Longmans, Green and Co., 1954, pp. 562-565.
2. H. Goldschmidt, German Pat., D.R.P. 96317, 1895; Z. Elektrochem., 1898, 4, 494; Ibid., 1899, 6, 53; Electrochem. Metall. Ind., 1908, 6, 360; Stahl Eisen, 1898, 18, 408; H. Goldschmidt and C. Vautin, J. Soc. Chem. Ind., 1898, 17, 543.
3. J.W. Mellor, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, vol. 5, new impr., London, Longman, Green and Co., 1956, p. 218.
4. Tested Demonstrations in Chemistry, ed. G.L. Gilbert, et al., vol. 1, Granville, OH, Denison University, 1994, p. H-9.
5. P.J. Roebuck, Educ. in Chem.(Britain), 1979, 16, 178.
6. J.R. Crellin, Educ. in Chem. (Britain), 1980, 17, 93.
7. B.Z. Shakhashiri, Chemical Demonstrations, A Handbook for Teachers of Chemistry, vol. 1, Madison, The University of Wisconsin Press, 1983, p. 88.
8. G. Fowles, Lecture Experiments in Chemistry, 3-rd ed., London, G.Bell & Sons, Ltd., 1947, pp. 285-287.
http://www.chem.leeds.ac.uk/delights/texts/Demonstration_17.htm
More References
Microscale thermite reactions.
J. Chem. Educ. (1998), 75(12), 1630-1631.
How to get the most from the dichromate volcano demonstration: aluminothermy.
J. Chem. Educ. (1984), 61(10), 908.
The thermite lecture demonstration.
J. Chem. Educ. (1979), 56(10), 675-6.
Abstract
A lecture demonstration of the thermite reaction is described. The mixt. of Al powder and Fe2O3 is ignited by the heat produced on oxidn. of several milliliters of glycerin added to KMnO4 crystals placed in a small depression on the surface of the reaction mixt. The reaction can be carried out in a small clay flower pot with a sand bath to collect the molten Fe, with the whole setup being surrounded by a proper safety shield to contain sparks and provide protection from this highly exothermic reaction.
A modified thermit lecture demonstration.
J. Chem. Educ. (1981), 58(10), 802.
Abstract
A lecture demonstration using Thermit with a KMnO4-glycerin ignition mixt. is described.
Do Your Part To Win The War
