The Lead and Iron Oxides

Here is a chemistry paper on the lead and iron oxides that I wrote in September of 2018. It provides definitions for the terms “compound” and “oxidation”, as well as giving physical and chemical descriptions of each of the three oxides of lead and the three oxides of iron. I hope you learn something, enjoy, and, if you have a question, ask in the forums!


Lead and Iron are two elements most ancient in their utilization by humans for a vast array of tools and products from swords to ceramics. Quite often, however, they were used in impure forms as oxides. Here a brief explanation of oxidation might prove useful: oxidation is an example of a chemical reaction, which is any interaction between atoms of one or more elements in specific ratios to form a new substance, called a compound (“Definition of Compound”, 2017, pp. 1). The constituents of the compound become chemically bonded and cannot be separated by physical means (“Definition of Compound”, 2017, pp. 1). The resulting compound may have entirely different properties from the constituents (Chemical Reactions, n.d.). One way to form such compounds is by oxidation, the process by which one element, the oxidizer, accepts electrons from another element thus becoming bonded to it (Clark, 2016). Oxidation was named after the gaseous element oxygen, because oxygen is an oxidizing element, as is it highly electronegative, eager to steal electrons. In fact, oxidation was originally understood as in terms of oxygen transfer, rather than the more accurate model of electron transfer (Clark, 2016). Lead and iron are both more electropositive than oxygen, so they will be oxidized in a reaction with oxygen. Depending upon the conditions in which this reaction takes place, it can lead to several different compounds with different properties and uses (Winn, 2004).

Beginning with iron, the first possible compound is a mineral called hematite. Hematite, or Fe2O3, is one of the most common minerals in the world, and is present, at least in small amounts, in many rocks, e.g. sandstone, that have a reddish or brownish coloration, caused by the presence of hematite, although the mineral itself can vary greatly in color, from gray to silver-gray, black to brown and reddish brown (Winn, 2004, pp. 1). In fact, hematite was used until recently to make a dye of the latter color, before cheaper alternatives were developed. It is also responsible for the coloration of Mars, the Red Planet (Winn, 2004, pp. 2). Although it is only paramagnetic under normal conditions, it becomes strongly magnetic when heated, similar to another iron oxide, magnetite. Its hardness ranges from 6-7 on the Mohs Scale, and it may contain small amounts of Titanium. As the principle ore of iron, hematite is mined for the industrial production of iron and is the source of approximately ninety percent of all iron. Fine mineral specimens can be found in several localities, including Minas Garais (Brazil), Cumberland (Cumbria, England), and Ria Marina (island of Elba, Italy). (Friedman, “The Mineral Hematite”, 2018). The chemical reaction that forms hematite looks like this:

4Fe+3O2→2Fe2O3 (Winn, 2004)

That is what happens if there is ample oxygen available, but a different result occurs if the oxygen is less plentiful: Fe3O4, or magnetite. Notice that magnetite has higher iron to oxygen ratio than its cousin, hematite. (Winn, 2004, pp. 4)As referenced before, magnetite earns its name for being a natural magnet and the only mineral with this property. In coloration it is dark gray to black with a hardness slightly greater than hematite at 5.5-6.5. It also differs from its duller cousin in luster, having a metallic luster. Like hematite, though less widely used, it is an important ore of iron. It is of scientific interest due its pronounced magnetic properties. Magnetite can be found almost anywhere around the world, but there are a few noteworthy sources, such as Binn Tal (Wallis, Switzerland), Parachinar (Pakistan), and Cerro Huanaquino (Potosi, Bolivia). (Friedman, “The Mineral Magnetite”, 2018). The chemical reaction that forms magnetite looks like this:

6Fe+4O2→2Fe3O4 (Winn, 2004)

The final oxide of iron is known wüstite. This compound was named after geologist and paleontologist Ewald Wüst (1875-1934) of the University of Kiel in Germany. Wüstite has a hardness of 5-5.5 and occurs mainly in meteorites and anthropogenic slags. (“Wüstite”, 2018). Although its chemical formula is often given as FeO, it breaks the law of definite proportions; the ratio of iron to oxygen ranges between 0.85-0.95/1. Because of this, it is known as a nonstoichiometric compound. This technically allows for an almost infinite number of iron oxides, but all the non-stoichiometric oxides of iron are categorized as wüstite. (Winn, 2004, pp. 9)

Despite this anomaly, most compounds do have distinct stoichiometries, like the lead oxides. first lead oxide is lead monoxide, or PbO. It forms when is heated in the presence of oxygen and can take one of two forms, litharge or massicot, differentiated by their crystal structure. Both are yellowish solids, litharge has a tetragonal crystal structure and massicot has an orthorhombic crystal structure. (“Lead”, 2018, pp. 12). They both have a hardness of 2 on the Mohs scale and have dull, greasy lusters. Litharge has a variety of uses, including in lead acid batteries, glazing pottery, pigments, lead glass, and oil refining. Litharge mines occur on every continent of the world, with an especially high concentration in European countries, such as Sweden, the United Kingdom, and Germany. (“Litharge”, 2018). Massicot mines can be found in many countries around the world, including Madagascar, Namibia, Australia, and Germany (“Massicot”, 2018).

The second lead oxide is known as minium, after the Minius river located in the Northwest of Spain. The chemical formula is Pb3O4, lead tetroxide. Another name for it is red lead, because it can be made into a beautiful read pigment that has been used in paintings since the time of the ancient Romans. Paintings made with minium are called miniatures. (“Red Lead”, n.d.). The hardness of minium is 2.5, and it has a tetragonal crystal structure just like litharge, with a similar luster. Mines are concentrated in Europe but can be found on every continent. (“Minium”, 2018).

The final oxide of lead is plattnerite, otherwise known as lead dioxide (PbO2). It was named in honor of Karl Friederich Plattner (1800-1858) who served as professor of metallurgy and assaying at the Bergakademie of Freiburg in Saxony, Germany, by Karl Wilhelm von Haidinger. It is a brown to black mineral that is commercially produced in a process involving the oxidation of the lead oxide previously discussed, minium, by chlorine (“Lead”, 2018, pp. 13). Plattnerite is used in curing polysulfide rubbers, matches and pyrotechnics, and dyes (“Lead”, 2018, pp. 13). The hardness of plattnerite is 5.5 and it has a dull, metallic luster. The plurality of plattnerite mines are in North America, and of those approximately half are in Mexico and half are in the United States, concentrated in the Western side of the country. (“Plattnerite”, 2018).

These three oxides of iron and three oxides of lead are all very useful and very different from each other. This demonstrates the power of chemical reactions, to take the same two elements in different proportions and create new substances with different properties. However, as was touched on briefly, the chemical composition of a compound is not the sole determining factor of the properties of a substance; other factors, such as crystal structure, also play very important roles, as seen in the two forms of lead monoxide, litharge and massicot (“Lead”, 2018, pp. 12). Regardless of their composition and crystal structure, human beings have used the six oxides discussed above for a long time, some for millennia (“Red Lead”, n.d.), and will probably continue utilizing these useful compounds long into the future.


References

“Chemical Reactions”. (n.d.). Retrieved September 4, 2018, from ric.edu: 

Clark, J. (2016, May 1). Definitions of Oxidation and Reduction. Retrieved September 3, 2018, from chem.lbretexts.org: 

“Definition of Compound”. (2017). Retrieved September 3, 2018, from chemicool.com: 

Friedman, H. (2018). The Mineral Hematite. Retrieved September 3, 2018, from minerals.net: 

Friedman, H. (2018). The Mineral Magnetite. Retrieved September 3, 2018, from minerals.net: 

‘Lead”. (2018). Retrieved September 3, 2018, from brittanica.com: 

“Litharge (Lead(II) Oxide), Lead Monoxide”. (2018). Retrieved September 3, 2018, from reade.com: 

“Litharge”. (2018). (Hudson Institute of Minerology) Retrieved September 3, 2018, from mindat.org: 

“Massicot”. (2018). (Hudson Institute of Minerology) Retrieved September 3, 2018, from mindat.org: 

“Minium”. (2018). (Hudson Institute of Minerology) Retrieved September 3, 2018, from mindat.org: 

“Plattnerite”. (2018). (Hudson Institute of Minerology) Retrieved September 3, 2018, from mindat.org: 

“Red Lead”. (n.d.). Retrieved September 2018, 2018, from webexibits.org:

Winn, J. S. (2004, January 6). Stoichiometry of Iron Oxides. Retrieved September 3, 2018, from dartmouth.edu: 

“Wüstite”. (2018). (Hudson Institute of Minerology) Retrieved September 3, 2018, from mindat.org: 

A Short Exposition of Photosynthesis

Here is a short research paper on photosynthesis, that most wonderful and complex phenomenon that makes life possible. It deserves much deeper treatment that a short, high-school level report, but I hope this will provide a decent starting point on your learning journey. Don’t forget to ask your questions in the forum!


Photosynthesis is the process by which energy from the sun is used to chemically combine carbon dioxide (CO2) with water (H2O) to make oxygen and glucose. Green plants, i.e. plants with chlorophyll, and some other organisms utilize this chemical reaction to make food. Plants are primary producers occupying the lowest trophic level. They support all higher trophic levels and thus their level has the highest biomass. Without photosynthesis, most life on Earth would not exist. (1).

            The majority of photosynthesis in plants takes place in the middle layer of the leaves, or the mesophyll. The cells in the mesophyll are equipped with organelles called chloroplasts, specifically designed for carrying out photosynthesis. Inside the chloroplasts are what resemble stacks of coins. Each coin is called a thylakoid and has a green pigment called chlorophyll in its membrane. The entire stack is a called a granum. The grana are occupy a fluid-filled space called the stroma. (1).

            Photosynthesis is actually a complex series of chemical reactions, some being light-dependent and others being light-independent. The light-dependent reactions take place in the thylakoid membranes where the chlorophyll absorbs light which is converted to adenosine triphosphate (ATP), an energy carrying molecule, and NADPH, an electron carrying molecule. It is here that the oxygen we breathe is created from water as a byproduct and diffuses out through the stomata, tiny pores in the surface layer of leaves letting oxygen diffuse out and carbon dioxide diffuse in. (1).

            Then begin the light-independent reactions that are collectively known as the Calvin cycle. They occur in the stroma and use the ATP and NADPH to fix carbon for use in constructing cells and form three-carbon sugars, glyceraldehyde-3-phosphate, or G3P, molecules,  that link up to make glucose. (1).

            In summary, the heat energy from sunlight ends up stored as chemical energy in the bonds of the sugar molecules that can be metabolized by plants and other organisms. (1). The reaction absorbs heat so it can be described as endothermic. (2).

References


  1. “Intro to Photosynthesis.” (2018). Khan Academy. https://www.khanacademy.org/science/biology/photosynthesis-in-plants/modal/a/intro-to-photosynthesis. Date-accessed: 5/14/2018.

2. Helmenstine, Anne Marie. (2018). “Endothermic Reaction Examples.” ThoughtCo. https://www.thoughtco.com/endothermic-reaction-examples-608179

Why the Statue of Liberty is Green

The Statue of Liberty with its iconic green color.
The Statue of Liberty, image retrieved from the National Park Service (4).

The Statue of Liberty is an iconic national monument on Liberty Island that was dedicated on October 28, 1886 (1). Back then, it was not as we see it today; then the exterior was copper colored, because, of course, it was made of copper! But today it is tiffany blue (2) to mint to seafoam green (3), depending upon the lighting (1, see image) (4). This was the result of a series of chemical reactions that took place over the first thirty years after the statue was assembled (5) and provide the reason why the Statue of Liberty is green.

            The first reactions involve a concept called reduction in chemistry. Reduction occurs when an atom that is being oxidized donates electrons to the oxygen atoms. Chemists say that the oxygen atoms have been reduced. (6). They assign an oxidation number to the atom being oxidized that indicates the number of electrons gained or lost by that atom: a positive oxidation number means electrons have been donated and a negative oxidation number means electrons have been gained. Since oxygen is highly electronegative, it is eager to steal electrons. So oxygen tends to take electrons, which reduces its charge. That is the origin of the term “reduction.” Oxygen is more electronegative than most other elements, so in a reaction it is generally the atom that takes electrons and the it is generally the other atoms that give electrons, since they are more electropositive. (7). Since most elements tend to donate electrons to oxygen, the losing of electrons to another element is called “oxidation.” Reduction and oxidation are opposites, but they always go together. Thus, a reaction involving the giving and taking of electrons is called a redox reaction. (6).

            In the first reaction, the copper is oxidized, by oxygen (which is reduced), to form CU2O. This compound is pink or red. Then the copper cation continues to react with oxygen to form copper oxide, 4CuO, which is black to brown. In the first years after Libertas’ figure was erected near New York City, much coal was burned in that city. The resulting air pollution wafted over the Statue of Liberty, bringing with it sulfur. This reacted with the copper to form the compound 4CuS, which is black. Three final compounds form from these initial compounds with the addition of carbon dioxide and hydroxyl ions: CuCO3(OH)(green), Cu3(CO3)2(OH)2 (blue), and Cu4SO4(OH)6(green). (8).

            These three compounds form the iconic blue-green verdigris that encases the Statue of Liberty today. The Statue of Liberty provides a great lesson in chemistry about redox reactions and successive reactions.

References

  1. “Liberty Enlightening the World.” (n.d.).National Park Service. https://www.nps.gov/stli/index.htm. Date-accessed: 4/10/2018
  2. Knapton, Sarah. (2017). “First new shade of blue discovered for 200 years to be turned into Crayola crayon.” See image at end of article. The Telegraph. https://www.telegraph.co.uk/science/2017/05/12/first-new-shade-blue-discovered-200-years-turned-crayola-crayon/. Date-accessed: 4/10/2018.
  3. Morris, Brian. (2015). “50 Shades of Green…and One Shade of Blue.” PsPrint. https://blog.psprint.com/de(signing/50-shades-green-one-shade-blue/. Date-accessed: 4/10/2018.
  4. “Plan Your Visit.” (n.d.). National Park Service. https://www.nps.gov/stli/planyourvisit/index.htm. Date-accessed: 4/10/2018
  5. “Why is the Statue of Liberty Green?” (2018). Wonderopolis. https://wonderopolis.org/wonder/why-is-the-statue-of-liberty-green. Date-accessed: 4/10/2018.
  6. Clarck, Jim. (2016). “Definitions of Oxidation and Reduction.” LibreTexts. https://chem.libretexts.org/Core/Analytical_Chemistry/Electrochemistry/Redox_Chemistry/Definitions_of_Oxidation_and_Reduction. Date-accessed: 4/10/2018.
  7. “Oxidation-Reduction (Redox) Reactions.” (2018). Khan Acandemy. https://www.khanacademy.org/science/chemistry/oxidation-reduction/modal/a/oxidation-reduction-redox-reactions. Date-accessed: 4/10/2018.
  8. Helmenstine, Anne Marie. (2018). “Why is the Statue of Liberty Green?” Thought co. https://www.thoughtco.com/why-statue-of-liberty-is-green-4114936. Date-accessed: 4/10/2018.

The Chemical History of Aluminum

An example of aluminum in use
Aluminum electric line, used for light weight and decent conductivity.

Aluminum is an essential component in a myriad of modern conveniences from airplanes to pop cans, prized for its high strength to weight ratio and resistance to atmospheric corrosion. This silvery-white metal was even more highly valued in the time of Napoleon III, more even than gold. However, this was merely for its extreme rarity rather than for its applications in manufacturing. This only changed when two young chemists, American Charles Martin Hall and Frenchman Paul Héroult, standing on the shoulders of other notable scientists before them, discovered a chemical method to economically extract pure aluminum from its ores. Their method was only the latest in a long line of attempts in the history of aluminum, but it was the first to be commercially utilized on a large scale and is still in use today. (“Commercialization of Aluminum”).

However, for thousands of years aluminum was not even known to exist, despite its use in compounds predating 5000 BC. Ancient Mesopotamians used aluminum-rich clays to craft fine pottery. In addition, aluminum compounds were utilized by Ancient Egyptians and Babylonians as medicines almost 4,000 years ago. And from the ancient world to the medieval period, an aluminum compound, known as alum today, was used to bind dyes to textiles. However, it was not until the eighteenth century that anyone suspected that a metal could be found in these useful compounds. (“Hall-Heroult”).

Aluminum is Christened

Humphry Davy, an English chemist, made the first attempt to extract this metal in 1807. It was made after a long string of successes in isolating pure metals from compounds, such as potassium from potash and sodium from salt, using a method called electrolysis. (Pizzi). Electrolysis is the process of running a direct electrical current from a battery or other source through an ionic solution called an electrolyte using two metal bars as electrodes. The electrons flow from one electrode to the other making one, the cathode, negative and the other, the anode, positive. Positive ions, cations, in the solution are attracted to the cathode and negative ions, anions, are attracted to the anode. When the ions reach their respective electrodes, electron exchange occurs causing a chemical reaction. In this way pure elements can be separated from compounds. For example, if two copper electrodes connected to a power source were inserted into a solution of molten salt, sodium chloride, the negative chloride ions would be attracted to the anode and the positive sodium ions would be attracted to the cathode. At the cathode, the sodium ions would transfer their excess electrons to the cathode and would become neutral. And the same thing would happen to the chlorine ions at the anode, except here there would be a gain of electrons for the chloride ions. The reactions would look like this: (“Electrolysis”).

​At the cathode: Na++ e- → Na

At the Anode: 2Cl- → Cl2 + 2 e-

​Although Davy failed to extract aluminum from alum in this way, he satisfied himself that the metal existed and named it alumium, afterwards rechristening it as aluminum. (Pizzi).

Aluminum is Isolated

​The first sample of aluminum was obtained in 1825 by Danish scientist Hans Christian Oersted who heated a mixture of aluminum chloride and potassium-mercury amalgam under reduced pressure. This caused the mercury to boil away leaving an impure sample of aluminum. (Ashby). This chemical reaction was as follows (Caroll, 5):

AlCl3+3k→Al+3KCl

Using a similar process but with metallic potassium instead of potassium-mercury amalgam, German chemist Friederich Wohler distilled aluminum pieces up to the size of pinheads by 1840. From these samples, he determined the properties of aluminum such as ductility, color, and specific gravity. This made aluminum available, but only at the hefty premium of approximately 545 dollars per pound (1852 dollars). (“The Element Aluminum”).

The Deville Process

Thus, aluminum remained a mere curiosity until 1854, by which time French chemist Henri Saint-Claire Deville had successfully implemented his improvements on the methods of Wohler, namely the substitution of sodium for the more expensive potassium, to produce globules of aluminum the size of marbles using the following method: (“Deville-Castner Process”).

​Deville’s goal in aluminum manufacture was to reduce sodium and aluminum’s double chloride, 2NaClAl2Cl, using heated metallic sodium. The natural first step, then, was to manufacture the double chloride. This process was begun by taking powdered aluminum oxide, or hydrate of alumina (Al2O3+water), and combining it with lamp-black, salt, and charcoal. The resulting mixture was then moistened and processed in a pug mill before being extruded through dies, cut into three-inch cylinders, and dried. These cylinders were then precisely heated in an atmosphere of chlorine gas which causes the desired double chloride to evaporate from which gaseous form it was condensed into a pale-yellow, deliquescent material pungent in odor. The reaction that produced this most important ingredient was as follows: (“Deville-Castner Process”).

2Al2O3+3C2+4NaCl+6Cl2→2(Al2Cl62NaCl)+6CO

Translated into English, the reaction is alumina plus carbon plus salt plus chlorine gas equals double chloride plus carbon dioxide.

Now all that had to be done was the reduction of the double chloride using sodium. To this end, the double chloride was pulverized and mixed with slices of metallic sodium before being heated in a furnace along with cryolite, Na3AlF6, as flux. This resulted in the following reaction: (“Deville-Castner Process”).

2 (Na Cl) Al2Cl6 + 3 Na2 = 8 Na Cl + Al2.

​This simply means that the double chloride chemically reacted with the sodium to produce the desired aluminum and a byproduct of salt. This alone reduced the price to 115 dollars per pound, but the price did not solely depend on the quantities of aluminum that could be obtained.

​The cost of the materials used in aluminum manufacture, a very important one of which was sodium, was a very important factor influencing the price of aluminum. Until 1886 when a man named Hamilton Y. Castner began developing a safe, inexpensive method, sodium production had been very arduous and perilous. Involving electrolysis, the Castner process (completed in 1888) reduced the price of sodium five-fold. Although this development made aluminum more affordable, it was still prohibitive enough to keep aluminum from widespread use. (“Deville-Castner Process”).

The Hall-Heroult Process

​While these achievements represented giant leaps forward in aluminum production, the greatest and yet unsurpassed method was still to come. This discovery was twice made independently by two twenty-two-year-old chemists during the same year that Hamilton Y. Castner developed his sodium production process, 1886. Charles Martin Hall, an American graduate of Oberlin College in Oberlin, Ohio,, worked with Oberlin College professor Frank Fanning Jewett to develop the following process: (“Hall-Heroult”).

First, alumina is dissolved in a vat of molten cryolite at a temperature of 982 degrees Celsius (Ashby), acting as a flux as in the Deville process. Then, it, the electrolyte, is channeled into a cell with a carbon-lined cast-iron shell. Carbon anodes are suspended in the electrolyte and the carbon lining acts as a cathode. An electric current is passed through the cell and the dissolved alumina separates into its constituents, oxygen and aluminum. The molten aluminum sinks to the bottom of the cell and the oxygen remains around the anodes. This involves two half-reactions, reduction of the aluminum at the cathode and oxidation of the oxygen at the anode: (“Extraction of Aluminum”).

Reduction: Al3+  +  3e-       Al

Oxidation: 2O2-  –  4e-       O2

These two half reactions can be combined into one whole reaction which is as follows (“Extraction of Aluminum”):

2Al23+O32-(l)                 4Al(l)   +     3O2(g)

​Paul Louise Toussaint Heroult independently discovered this same process just two months after Charles Martin Hall. He applied for and received a patent for it in France and applied for one in the United States in May of 1886. Although this was before Hall applied for his patent in July, Hall was able to prove that he discovered the method before Heroult made his patent application. Two years later, Hall founded the Pittsburg Reduction Company with financial assistance from six industrialists involved in Pittsburgh’s metallurgical market, including MIT graduate Alfred Hunt. That same year, the price of aluminum plummeted to $4.86. By 1993, the price had dropped to seventy-eight cents per pound. And by the 1930s, it aluminum was valued at just over twenty cents per pound. (“Hall-Heroult”). This remarkable figure was not, however, achieved by Hall and Heroult alone.

Two other notable developments aided in the commercialization of aluminum, namely the dynamo and the Bayer Process. Without the former, none of what Hall and Heroult achieved would have been possible. Invented by Siemens, Hopkinson, and Edison in 1881, it provided the power source necessary for electrolysis. Later, aluminum companies moved to places where abundant hydroelectric power could be found to drive the dynamos. While the dynamo and hydroelectricity made inexpensive electricity abundant, the Bayer process cheapened the raw material for aluminum production, alumina. As the Castner process cheapened sodium, the Bayer process cheapened the production of alumina. (Ashby).

Summary

​Use of aluminum began several millennia ago in pottery and medicine, but it was not until relatively recently that it was discovered. Over the years, better and better methods were devised for aluminum production and the production of ingredients involved in its production. This continued until the price, once greater than that of gold, was reduced sufficiently so that the full industrial potential of aluminum could be realized. This was achieved in large part by the work of Charles Martin Hall and Paul Louise Toussaint Heroult in 1886. (“Hall-Heroult”). These young chemists were heirs to the work of several other notable chemists who made important contributions to the field of aluminum including Humphrey Davy, Hans Christian Oersted, Friedrich Wohler, Henri Saint-Claire Deville, Frank Fanning Jewett, Hamilton Y. Castner, and Karl Joseph Bayer. Standing on the shoulders of these intellectual giants and using new technologies of their day, Hall and Heroult were able to make aluminum available for widespread use.

Works Cited

​”Aluminium and Its Manufacture by the Deville-Castner Process.” (1889). Science, 260-262.

​Ashby, J. (1999). “The Aluminium Legacy: the History of the Metal and its Role in Architecture.” Construction History, 79-90. Retrieved from: https://www-jstor-org.www.libproxy.wvu.edu/stable/pdf/41613796.pdfrefreqid=excelsior%3Af258906f5aaf47b4f874f03071206939                                                                              

“Commercialization of Aluminum.” (2001, November 2). Retrieved November 11, 2018, from acs.org:                                                                                https://www.acs.org/content/dam/acsorg/education/whatischemistry/landmarks/aluminumprocess/commercialization-of-aluminum-commemorative-booklet.pdf

Dr. William F. Caroll, Jr. (2012, April). “From Garbage to Stuff: How we Recycle Plastics.” The Alembic, 39(3), p. 5. Retrieved from                                 https://www4.uwsp.edu/chemistry/acscws/39%20-%203%20April%202012.pdf

Education, T. J.-O. (n.d.). “The Element Aluminum.” (G. Steve, Editor) Retrieved from Jefferson Labs:                                                                                   https://education.jlab.org/itselemental/ele013.html

“16.7: Electrolysis: Using Electricity to Do Chemistry.” (2018, May 4). Retrieved November 11, 2018, from LibreTexts.

“Extraction of Aluminium – Hall (Electrolytic) Cell.” (n.d.). Retrieved November 11, 2018, from mypchem.com:                                                                  http://mypchem.com/myp10/myp10_htm/al_ext.htm

Pizzi, R. A. (2004). “Humphry Davy, Self Made Chemist.” Chemistry Chronicles, 49-51.

“Production of Aluminum: The Hall-Héroult Process.” (2018). Retrieved November 11, 2018, from American Chemical Society:                                  https://www.acs.org/content/acs/en/education/whatischemistry/landmarks/aluminumprocess.html