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The Extraction of Copper from
Malachite and an adaption of Copper
Extraction to a Classroom Setting
By Jimmy Frisbie
Partner: Olivia Forney
Chem 113M, Thursday 8 A.M. – 12 P.M.
TA: Rob Johnson
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Introduction:
Metals have long been important to humans ever since they roamed the Earth. Metals
have been used in industry, consumer products, and architecture.1 However, most metals are not
found in the Earth in their pure form. Instead, ores (naturally occurring materials that contain
extractable valuable metals or minerals) have to be extracted from the earth. Then, these metals
have to be extracted from the ores. Extraction of these metals can occur in numerous ways. One
way extraction can occur is through physical methods, such as floating. In this method, oil can
be coated on the ores which will create hydrophobic interactions when introduced to water.
When air is blown through a mixture of water and the ore, the ore will float to the top.
Additionally, metal can be extracted from reduction reactions by stripping the ore of oxygen and
combining it with another element (often carbon).2
Evidence shows that multiple metals were used by ancient civilizations. For example,
many Egyptian artifacts contain purified metals. Although some believe the Egyptians
deliberately knew how to purify the metals, most believe they obtained their knowledge of metal
extraction serendipitously. Quite possibly, their discovery was due to large pottery that was put
into kilns and then hammered to produce household goods.3
Copper, one of the world’s most used metals, was first used in coins and ornaments
starting around 8000 B.C.E. Throughout the ages, copper was also used in tools and eventually
combined (alloyed) with tin to produce Bronze.2 Extraction, the process of getting copper from
different ores,4 has become an important topic of conversation as copper supplies have become
more limited. Porphyry copper deposits yield about two-thirds of the world’s copper because of
its accessibility and its low cost. These porphyry deposits are found in western mountainous
regions of South America (especially the Andes).5 Today, the U.S. has the second largest copper
mines in Arizona, New Mexico, Michigan, and Montana.1 Copper can be extracted in multiple
ways. Normally, copper is obtained by taking ore (which contains the metal) and processing it
(or sorting it) by flotation, smelting, and refining. Specifically, copper can be extracted from
either sulfide or oxide ores.1
Copper is a very popular metal for numerous reasons. It has a high level of thermal and
electrical conductivity. Most household appliances contain copper. In addition to using it for
electrical wiring (due to the high electrical conductivity), it is often used in roofing, decorations,
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and even gutters due to its durability.6 Finally, copper is resistant to corrosion.7 As a result, it is
used in many roofs (which is made from sheets of copper).8
To extract copper from sulfide, sulfide ores are crushed into small pieces and then usually
mixed with water to help sulfide oxides float to the surface (although sometimes other leach
solutions are used). Then, the oxide can be dried and “smelted” (fusion of metal ores) to create a
relatively pure (99%) amount of copper. Copper can become even more pure through
electrolysis.1 Electrolysis occurs through oxidation and reduction actions. During electrolysis,
copper will be reduced at the cathode. Specifically, copper ions will become solid copper (Cu2+
(aq) + 2e-  Cu (s)). At the anode, oxygen is oxidized with carbon to form CO2. For copper
oxides, crushed ores go through a process where chemical reactions help separate the copper
from the other parts of the ore. Specifically, the ore is reacted with acids such as sulfuric acid.
Then, the resulting copper (II) sulfate solution will be introduced to an organic solvent that binds
to the copper ions. Finally, the new solution containing the bonded copper ions will be put
through electrolysis.16
In this experiment, we used malachite samples (Cu2(CO3)(OH)2) to extract our copper.15
The malachite will first be roasted to help isolate copper oxide from carbon dioxide gas and
water vapor. Then, the copper oxide will be heated with carbon and a reduction reaction will
allow the copper to become isolated while the carbon will form carbon dioxide with the oxide
anion.1
Procedure and Results:
The purpose of our experiment was to extract pure copper from malachite samples. Our
procedure came from the experiment “Copper Metal Extraction from Malachite.”17 I, along with
my partner, Olivia Forney, got a small amount of malachite and put them in a crucible. Then, the
crucible was heated by a Bunsen burner. The top part of the orange flame touched the bottom of
the crucible to ensure that the entire malachite sample could be heated. The mass of the
malachite sample was 0.5804 grams. The sample was heated for fifteen minutes. After waiting
five minutes for the sample to cool, we opened the crucible and let the sample cool for the next
ten minutes. The chemical process that took place during this step of the experiment consisted of
the following:9
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Cu2CO3(OH)2(s)  2CuO(s) + CO2(g) + H2O(g)
Next, the (now) copper (II) oxide was heated with carbon covering the entire copper (II)
oxide sample (which acts as the reducing agent). After being loaded into the crucible, the
crucible was heated for ninety minutes. The chemical process that occurred during this step is
shown below:
2CuO(s) + Cu(s)  2Cu(s) + CO2(g)
Then, after cooling for five minutes, the crucible lid was removed and the sample cooled
for another ten minutes. After the ten minutes, the carbon and copper was dumped out. The
sample was weighed in a weigh boat and the weight was recorded. The mass of the final copper
was 0.2995. The theoretical yield was calculated was obtained from the following calculation:10
0.5804 g (Cu2CO3(OH)2)*(1/222.1 g/mol)*(2 mol CuO/1mol Cu2CO3(OH)2)
= .3336 grams of Copper
The percent yield was 89.8%.8 I obtained the percent yield from the following calculation:
(0.2995 Grams/0.3336 Grams)*100=89.8%
After we obtained our final copper sample, we hit the copper with a hammer to test its
malleability. The sample bent and did not break. After hitting the sample with our hammer, we
conducted an electrical current through the sample and measured its resistance. When we tested
our sample, 0.01 Ω was recorded9&10.
Discussion:
For both of the reactions described above, the energy used to drive the reaction was given
in the form of heat. By heating the sample of malachite and the copper (II) oxide, we were able
to drive the endothermic reaction (we had to provide heat to make the reaction occur) to
completion. This energy was needed to not only release the carbon dioxide and water vapor in
the first reaction, but also to help the reduction reaction occur. Our percent yield was 89.9%.
This low percent yield is surprising since most error would lead to a percent yield of over 100%.
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During the heating steps of the experience, not all of the CO2 or H2O may have been burned off.
In addition, it was very difficult to remove all the carbon from the copper sample at the end of
the second heating process. I believed that the final mass would be heavier than the theoretical
yield. I believe that the low percent yield of copper was due to weighing inaccuracies.
However, there is also a chance that during the first and second heating, some of the copper may
have broken off from the malachite sample. If pieces of the copper sample broke off and were
too small to see using only the naked eye, it is possible that all of the copper was not gathered
after the burning process.
After we had extracted the copper, we hit it with a hammer. It was very malleable. This
proved that most of the impurities had been burned off of the copper. If the impurities were still
present, the sample would be much more brittle and may have broken. Also, an electrical current
was run through the metal sample and resistance was recorded. Our resistance was extremely
low (0.01 Ω). This low resistance proved that our sample was mainly copper since copper is an
extremely good conductor.
Conclusion:
Overall, our procedure did not yield any surprises. After weighing the initial malachite
sample, we expected a decrease in mass after the burning due to the release of carbon dioxide
and water vapor. Also, since the copper (II) oxide was reduced to copper and carbon dioxide,
there was an additional decrease in mass. To improve the percent yield of pure copper, I believe
it would be beneficial to look through a microscope to find any microscopic pieces of copper (II)
oxide that might have been present after the first heating. If possible, I would weigh all of these
small pieces to see if their mass would improve the percent yield after the second heating. In
addition, if a method was developed to clean the carbon off the pure copper sample, it would
help make the final mass reading much more accurate. I would try to accomplish these accurate
readings by hitting the purified sample with a hammer more than prescribed in the lab manual.
Also, I would roll the sample(s) in a clean paper towel to try to remove any excess carbon.
However, I would not soak my sample in water since the water would alter the final mass of the
sample.
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Aluminum Extraction and Adaptation to a Lab Setting:
Unlike copper that can be readily extractible by carbon reduction reactions in a simple
laboratory setting, aluminum is much more difficult to extract from ore. In today’s society,
aluminum is very important for many consumer products. In fact, aluminum is the second-most
used metal in the United States (right behind steel).12 Aluminum is used in large quantities in the
auto industry, including license plates and lightweight car parts. In addition, they are used in
construction, especially in roofs and packaging. As many engineers quickly figured out,
aluminum does not rust like a traditional metal because it reacts with air to form a protective
layer of aluminum oxide around the metal. Finally, aluminum is used in many household items,
including cooking utensils, sporting equipment, and even lawn furniture.12
Unfortunately, aluminum cannot be directly extracted from the ground. Instead, it is
found in ore deposits called bauxites. Right now, there are an estimated 29 billion metric tons of
bauxite which could last for more than one hundred years.14 Bauxites are mostly found in the
tropics and are found a few meters down in the Earth’s surface. As a result, there are multiple
environmental issues caused by bauxite mining. Erosion, excess run off, and disturbance of
hydrology are all impacted wherever bauxites are mined. Bauxites are mainly red because of the
iron present in the bauxite. After they are mined, they are crushed and then are sent to
processing plants.14
Bauxites are more or less impure aluminum oxides12 that combine with iron oxides,
silicon dioxides and titanium dioxides. Therefore, it becomes necessary to extract the aluminum
before it can be used for commercial goods. Bauxites are purified using the Bayer Process.12 On
an industrial scale, the Bayer Process involves mixing the bauxites with extremely hot
concentrations of sodium hydroxide. The sodium hydroxide not only dissolves the bauxite, but
prevents the other impurities (the iron oxides and titanium oxides) from dissolving. Depending
on the form of the aluminum oxide, different temperatures are needed to effectively dissolve the
bauxite (between 140 and 240 degrees Celsius, and 35 atmospheres).11 The insoluble oxides are
then removed from the solution via filtration. Sometimes, silicon oxides dissolve in the sodium
hydroxide solution. When this happens, the silicon oxides are dealt with in the next step of the
process.13
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Next, the solution containing the dissolved bauxites is treated with carbon dioxide. The
carbon dioxide neutralizes the sodium hydroxide and precipitates the aluminum oxides out of the
solution. However, the silicon oxides are left in the solution. After this step is complete, the
solution with the silicon oxides is discarded.
For many years, scientists were unable to move past this step in purifying aluminum.
However, in 1850, a method was devised that helped reduce aluminum chloride with sodium
metal to produce aluminum and sodium chloride.13 Unfortunately, this method was largely
inefficient and costly. In 1886, Charles Hall invented the Hall-Heroult Process that allowed
large extractions of aluminum. The aluminum oxides extracted from the Bayer Process were
mixed with cryolite (sodium fluoride and aluminum fluoride) and heated until both the cryolite
and the aluminum oxides were melted. Then, the liquid was electrolyzed. Although a voltage of
only 4-5 volts was needed to electrolyze the liquid, the current needed to be between 50,000 and
150,000 amps. At the anode, carbon in the form of graphic was used which aided in oxidation
(so the aluminum could be reduced). The cathode allowed the aluminum ions to be reduced to
aluminum. The anode oxidized oxygen anions to produce carbon dioxide.13
Lab Procedure:
To extract aluminum in a laboratory setting, small pieces of bauxites must be transported
to the classroom. Then, one would need to determine the percentage of aluminum in the bauxite.
For instance, if only 10% of the bauxite was aluminum, you would have to measure out ten
grams of the bauxite to ensure that you could get at least a full gram of aluminum. Then,
bauxites would need to be put in a crucible with a fixable lid (so pressure would increase as
temperature increased). Then, sodium hydroxide solution would be added to the bauxite. The
amount of sodium hydroxide needed would be dependent upon the amount of moles of
aluminum in the sample. If there was one gram of aluminum (which would be equal to 0.037
moles (1/26.982 amu) of aluminum), you would have to use 0.037 moles of sodium hydroxide
(which would equal 1.48 grams (0.037 mol*(22.9+35.50), or 1.48 ml). Regardless of the amount
of sodium hydroxide required, excess solution should be added to the bauxite to make sure it
completely dissolved.
Next, the bauxite would be placed on a heating pad that would be able to reach at least
140 degrees Celsius. Then, after a fifteen minute period (or a period appropriate through testing
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and error), the solution would be allowed to cool. After cooling, any precipitates would be
removed. Then, the solution would be put into an air tight tube. Next, carbon dioxide would be
added to the tube to help precipitate out the aluminum oxide. Carbon dioxide would be run
through the tube until no more precipitate formed. The rest of the solution would be discarded.
Finally, cryolite would be made by mixing sodium fluoride with aluminum fluoride. It would be
added to the aluminum oxide in a crucible. Once again, sodium fluoride and aluminum fluoride
solution would have to be equal or greater than the amount of aluminum oxide. Using the
example of 1 gram of aluminum in the original bauxite sample, 0.037 moles of aluminum oxide
would be needed. Therefore, 0.037 moles of sodium fluoride and aluminum fluoride would be
necessary (0.0185 moles of sodium fluoride and 0.0185 moles of aluminum fluoride). This
would be equivalent to 0.779 grams (0.799 ml) of sodium fluoride and 1.59 grams (1.59 ml) of
aluminum fluoride. For both sodium fluoride and aluminum flouride, grams and milliliters in the
previous sentence were found by multiplying 0.037 moles by the molar mass of sodium fluoride
and aluminum fluoride, respectively. Like the step earlier, excess should be added to make sure
all of aluminum oxide is dissolved.
The crucible would then be heated over a Bunsen burner flame (temperatures need to
reach approximately 900 degrees Celsius). After all solids had dissolved, electrolysis would be
run on the solution. Carbon would be placed at the anode and 4-5 volts would be used with a
current of 15,000 amps.13 The electrolysis would be run and aluminum ions would be reduced to
aluminum at the cathode (Carbon would have to be constantly and carefully analyzed to see if it
needed replaced at the anode).
Although it is very possible to create aluminum in a lab setting, I believe it would be very
difficult to use the correct amounts of temperature, pressure, and solutions for the various steps
of the extraction process. Since aluminum is somewhat costly to produce due to high electricity
requirements, there are no readily available stats and figures on how much sodium hydroxide and
cryolite is needed to successfully extract aluminum. Also, to get high percent yields, trials would
have to be conducted to see how long the bauxites would have to be heated with the sodium
hydroxide. An instrument would also have to be constructed that could introduce carbon dioxide
into a solution without letting atmospheric air in. Finally, I do not know how long the cryolite
and aluminum hydroxide would need to be heated until all of aluminum hydroxide dissolved and
was ready for electrolysis.
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If the method above did not end up being feasible in a laboratory setting, pressure acidleaching may allow aluminum extraction to occur in a classroom setting. A recent study came
out where scientists were able to extract aluminum by pressure acid-leaching. In this process,
coal fly ash and calcium oxide were mixed at a high temperature and pressure to produce
aluminum. At temperatures between 200-210 degrees Celsius, a leaching time of 80 minutes,
and a ratio of acid to coal fly ash of 5:1, it is possible to yield relatively high amounts of
aluminum. By mixing the coal fly ash and calcium oxide, aluminum oxides can be created.
Then, the aluminum oxide can be electrolyzed, and efficiencies can reach as high as 87%.
However, extraction efficiencies are not as high for this method as other conventional methods.
This acid-leaching process also produces much more solid waste. Finally, it requires a large
amount of energy.16
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LITERATURE CITED
1. Wilts, E., & Ehret, C. (2013). Copper metal extraction from malachite. Chemistry 113M:
Laboratory Manual. 17 March 2014.
2. Clark, J. (2005). The extraction of metals – an introduction. Chemguide.co.uk. Web. 17
March 2014. Retrieved from:
http://www.chemguide.co.uk/inorganic/extraction/introduction.html
3. Ancient Egyption Raw Materials. (2005). Metals: sources, technologies, uses. Web. 17
March 2014. Retrieved from:
http://www.reshafim.org.il/ad/egypt/trades/metals.htm#rem12
4. Merrian-Webster. (2014). Extraction. Merriam-webster.com. 17 March 2014. Retrieved
from: http://www.merriam-webster.com/dictionary/extraction
5. Viking Minerals Incorporated. (2014). Importance of copper. Vikingmineralsinc.com.
Web. 17 March 2014. Retrieved from: http://www.vikingmineralsinc.com/copper.html
6. Clark, J. (2005). Copper. Chemguide.co.uk. Web. 17 March 2014. Retrieved from:
http://www.chemguide.co.uk/inorganic/extraction/copper.html
7. Lagowski, J. (2005). Copper: chemistry: foundations and applications. Vol. 1. (266267). New York: Macmillan Reference.
8. HSPG. (2012). Copper: characteristics, uses and problems. Gsa.gov. Web. 25 March,
2014.
9. Frisbie, J. (2014). Chem 113 Lab Notebook. 17 March 2014.
10. Forney, O. (2014). Chem 113 Lab Notebook. 17 March 2014.
11. Discovery. (2014). What is aluminum used for? Curiosity.discovery.com. Web. 20 March
2014. Retrieved from: http://curiosity.discovery.com/question/what-is-aluminum-usedfor
12. Clark, J. (2005). Aluminum. Chemgide.co.uk. Web. 20 March 2014. Retrieved from:
http://www.chemguide.co.uk/inorganic/extraction/aluminium.html
13. Ophardt, C. (2003). Conversion of bauxite ore to aluminum metal. Elmhurst.edu. Web.
20 March 2014. Retrieved from:
http://www.elmhurst.edu/~chm/vchembook/327aluminum.html
14. Hydro. (2014). Bauxite mining. Hydro.com. Web. 20 March 2014. Retrieved from:
http://www.hydro.com/en/About-aluminium/Aluminium-life-cycle/Bauxite-mining/
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15. Mineral News. (No Year). Malachite mineral data. Webmineral.com. Web. 25 March
2014. Retrieved from: http://webmineral.com/data/Malachite.shtml#.UzEJFvldVXk
16. Wu, C., Yu, H., & Zhang, H. (2011). Extraction of aluminum by pressure acid-leaching
method from coal fly ash. Science Direct: 22(2012). Pages 2282-2288. Web. 25 March
2014. http://www.tnmsc.cn/down/upfile/soft/20120929/34-p2282.pdf
17. Wilts, E., & Ehret, C. (2014). Copper metal extraction from malachite. Chemistry 113:
Labortory Manuel. Haydn McNeil: 2014. Print. 26 March 2014.