Alum Synthesis

The Synthesis of Alum from Aluminum lab centralized on the creation of potassium aluminum sulfate dodecahydrate from a set amount of aluminum, and using quantitative analysis to then compare the actual yield to the theoretical yield. The synthesis of alum was conducted through a series of steps involving aluminum foil, potassium hydroxide, and sulfuric acid. Given that aluminum is amphoteric, meaning it can dissolve in both strong acids and strong bases, potassium hydroxide and sulfuric acid were both suitable for the dissolution of aluminum. There are multiple applications to synthesizing alum; from eliminating waste in the environment by repurposing, to utilizing alum as a medical treatment, the practicality of alum is immense.
Any chemical

compound that consists of aluminum sulfate, the sulfate of another element, and water of hydrate is considered to be alum. After synthesizing the alum from aluminum pieces, quantitative analysis was required to determine the success of the experiment. By determining how much alum could be formed from the balanced chemical equations, a theoretical yield was able to be determined. When compared to the theoretical yield, the actual yield will determine the overall efficiency of the experiment as well as the percent error.

Procedure & Observations:
To begin the experiment, first a 5 X 5 cm piece of aluminum was obtained. Using a steel wool scour, this piece of aluminum was then scraped to remove the exterior coating. Once complete, the aluminum was washed and dried, then cut up into smaller pieces, about 2 mm X 2 mm. The aluminum pieces were then transferred to a beaker with a known mass and weighed to find the mass of the aluminum pieces. Next, 25 mL of 1.4 M potassium hydroxide (KOH) was acquired and added to the aluminum pieces in the beaker. This mixture took place in the fume hood while stirring well. The solution with the aluminum was then heated on a hot plate while stirring for 20 minutes; it was imperative that the solution did not reach a boil. While the solution was heating up, a glass funnel was set up on a ring stand with a Whatman #4 filter paper was folded into quarters and placed in the funnel. The stem of the funnel was placed in a 150 mL beaker. Once all of the aluminum had reacted by all of the pieces being dissolved and no hydrogen bubbles were present, the beaker was removed from the hot plate. Once cooled enough to handle, the solution was filtered into the 150 mL flask.
Next, 10 mL of 9M H2SO4 was obtained and mixed with the filtrate while stirring steadily. If white crystal formation was visible, then the solution should be heated again to dissolve the crystals. This new solution was then added to a beaker filled with crushed ice for approximately 20 minutes to allow for crystals of alum to form. While this is occurring, a filter flask with a Buchner funnel was attached to an aspirator. Once the 20 minutes was complete, the temperature was recorded; the temperature should be below 6 C to assure crystal formation. A Whatman #4 filter paper was then weighed and place in the Buchner funnel; the filter was moistened with water to hold it in place. Next, the aspirator was turned on and the 150 mL beaker was removed from the ice bath. Before the solution was added to the Buchner funnel, it was stirred to loosen the alum crystals. The alum solution was then poured into the funnel, utilizing a rubber policeman to scrape the beaker in order to obtain all of the crystals. Two 10 mL portions of methanol were added to the beaker and then poured into the aspirator in order to wash excess H2SO4 out of the crystals. After running the aspirator for an additional 10 minutes to dry the crystals, the filter paper with the crystals were weighed on a weight boat to obtain the mass of the alum crystals synthesized.

Results:

Table 1: Data Summary
Mass of clean Aluminum strip
0.127 g
Moles of Aluminum
4.71 x 10-3 mol
Molar Mass of Aluminum
26.981 g/mol
Theoretical moles of Alum
4.71 x 10-3 mol
Molar mass of Alum AlK(SO4)2.12H2O
474.39 g/mol
Theoretical yield of Alum
2.23 g
Mass of Alum obtained
1.754
% yield of Alum obtained
78.65%
Percent Error
21.34%

Data Analysis & Calculations:
Equation 1: Moles of Aluminum Strip

Equation 2: Theoretical Moles of Alum
2Al(s) + 2KOH(aq) + 6H20(l)  2K+(aq) +2Al(OH)4-(aq) + 3H2(g)

2Al(OH)4-(aq) + H2SO4(aq)  2Al(OH)3(s) + 2H2O(l) + SO42-(aq)

2Al(OH)3(s) + 3H2SO4(aq) Al2 (SO4)3(aq) + 6H2O(l)

K2SO4 (aq) + Al2(SO4)3 (aq)+24H2O(l) 2KAl(SO4)2⋅12 H2O(s)

K+(aq) + Al3+(aq) +2SO42-(aq) +12H2O(l)  KAl(SO4)2·12H2O(s)

Equation 3: Theoretical Yield of Alum

Equation 4: Percent Yield of Alum Obtained

Equation 5: Percent Error

After weighing out the amount of aluminum obtained, the theoretical yield could then be calculated as shown in Equation 2. In the balanced chemical equations, it is shown that for every two moles of aluminum, one mole of alum would be synthesized. Given this one to one ratio, the amount of moles obtained from the aluminum, as shown in Equation 1, would be equal to the theoretical moles of alum synthesized. Once the theoretical moles of alum were known, through the molar mass of alum, 474.39 g/mol, the theoretical yield of alum could be determined by multiplying the moles and the molar mass of alum, resulting in the grams of alum as shown in Equation 3. Once the experiment was complete, the percent yield was calculated in Equation 4 by comparing the theoretical yield of alum of 2.23 g and the actual yield of 1.754 g, resulting in a yield of 78.65%. This resulted in a 21.34% error as shown in Equation 5.

Research Information:
Due to the large abundance of scrap aluminum in our society from soda cans, foil, etc., recycling aluminum is necessary to maintain the health of our environment.

Given that aluminum is the third most abundant substance as well as the most prominent metal in the Earth’s crust, repurposing it is beneficial to everyone. There are multiple ways to repurpose aluminum, however, a prominent method is converting aluminum into alum. While alum can be naturally occurring in areas of heavy weathering, that oxidize sulfide, and include potassium bearing minerals all occur, it can also be synthesized as shown in this experiment. By synthesizing alum, many practical applications arise. Medicinally, alum can be utilized to stop bleeding in minor cuts, as a treatment for gingivitis and gum bleeding, as well as a preservative for pickling fruits and vegetables. On a larger scale, alum can be used in flocculation to treat dirty water to make it drinkable. In this case, the alum binds to the heavier particles in the water, such as dirt and sand, and causes them to sink to the bottom of the container. Next, a simple filtering would need to be done to render the water

drinkable.
For this experiment, potassium aluminum dodecahydrate was synthesized, however there are multiple other forms of alum. For example, both ammonium aluminum sulfate and sodium aluminum sulfate, are alternative forms of alum that have similar applications. These sulfates are made possible given that alum is made with single-charged cations such as potassium, ammonium, cesium, and others in the place of potassium. This allows for alternative resources to be utilized, while essentially synthesizing the same product due to the similar applications that each type of alum has. All types of alum have been prevalent in society for quite some time. Ranging from Ancient Egypt to the Middle Ages, alum has been a prominent resource in society for an immense amount of time.

Conclusion/Discussion:
The purpose of this lab was to successfully synthesize potassium aluminum sulfate dodecahydrate from aluminum, and practice quantitative analysis on the results.

Overall, the experiment produced a successful percent yield of 78.65% of alum from the pieces of aluminum. However, there was a relatively large percent error of 21.34%. This error could have resulted from multiple steps in the experiment. One notable source of error could have stemmed from not obtaining all of the alum crystals from the beaker before aspirating; some crystals could have remained in the beaker, resulting in a lower yield than expected. Another potential source of error may have been only running the alum crystals through the aspirator once; the aspirator removes liquids from the sample, drying them out, however some of the crystals could have ended up in the filtered liquid. By running this solution through the aspirator a second time, a greater yield could have occurred. Finally, when adding H2SO4, white crystals could have formed, resulting in it being necessary to reheat the solution. By not reheating the solution if the crystals did form, a loss of overall alum crystals would be significant, given that they could have formed in these white crystals, rather than the desired alum. To prevent these errors, it would be necessary to ensure that all of the crystals were removed from the beaker by aspirating, as well as filtering the solution more than once. As for the white crystal formation, a

preemptive reheat could potentially lessen the probability of the H2SO4 crystals from interrupting the alum formation. Finally, impurities in the aluminum metal itself could have disrupted the yield from the synthesis.
Overall, this lab was a success. This lab was unique when compared to previous experiments, however similar calculations were necessary in this experiment as in many others. These calculations include percent error and yield, as well as the theoretical yield that a substance can produce. The results from this lab depicted results that were expected, given that no lab can produce a 100% yield due to errors in the lab. In short, alum was successfully synthesized utilizing unique methods in the lab and quantitative analysis effectively translated the results from the lab into a comprehensive form.