Note: After eight hellish weeks of general chemistry two, I survived with a C+ in lecture and a B in the lab.

How I managed to pass, let alone get decent grades, I don't even know. To be honest, it's a miracle. Thank God it's over. 

Below is my final lab project report that I got a 90% on. Again, how I did that, I don't know, especially since I still don't know shit about basic chemistry. As such, I can't say that I'm proud of what I did. But, according to Word, I did spend 14 hours actively working on my final lab report, so I'm gonna put it in my library for that reason alone. Even though I don't like what I wrote (because... ew, chemistry). 

 

Identifying an Unknown Compound Via Qualitative Analysis

 

Introduction

Lewis acids and Lewis bases can be used in a variety of ways in Chemistry, especially when they’re combined with ligands. After a Lewis Acid-Base reaction occurs involving a ligand, either complex ions or coordination compounds form. When a neutral metal atom becomes surrounded by and bonded to other neutral atoms, it is known as a coordination compound, and typically ends up as a solid precipitate after a reaction. A complex ion is a coordination compound that is either positively or negatively charged, and typically shows up as a brightly colored, aqueous, clear solution after a reaction. This makes distinguishing these two types of compounds very easy in the lab.

Not only can coordination compounds and complex ions be distinguished from each other, but each coordination compound and complex ion has its own unique properties, such as different colors, textures, and melting points. Solutions containing known cations can be mixed with different Lewis acids and bases to make a product of a specific color and texture. The findings from that experiment can then be used to identify the cations of an unknown solution.

Procedure

Eight bottles of aqueous, 0.2 M solutions of eight different cations were obtained, along with a variety of acids and bases. A few drops of each cation solution were put inside of their own test tube. Then, exactly five drops of an aqueous 3.0 M Sodium Chloride (NaCl) solution were added to each test tube, and observations were recorded. After that, eight new test tubes were obtained, and each were filled with a few drops of one of the cations. Five drops of 6 M Hydrochloric Acid (HCl) were added to each test tube, and once again, observations were recorded. In both instances, the samples that contained solutions of Silver Nitrate and Lead (ii) Nitrate were kept and exposed to heat, and observations were recorded.

Similar experiments using the cation solutions and various other Lewis acids and bases were done. Sometimes, after a reaction was performed, a solution was exposed to heat to determine whether or not a precipitate would melt. Other times, the solution was washed with DI water and centrifuged to obtain a purer solid after a reaction resulted in a cloudy solution.

Finally, confirmatory tests specific to each cation were performed, producing results that were entirely unique to each of the eight cations. These numerous experiments were done so that known cations mixed with known Lewis acids and bases could be catalogued and later used to identify a solution containing four unknown cations.

Discussion

3 M NaCl:

The cation solutions used throughout the entire experiment were barium nitrate, copper nitrate, ferric nitrate, lead (ii) nitrate, manganese nitrate, nickel nitrate, silver nitrate, and aluminum nitrate, each of which were an aqueous solution with a concentration of 0.2 M. Eight bottles each containing a specific cation were obtained, along with eight test tubes and an aqueous, 3.0 M solution of Sodium Chloride (NaCl).

A few drops of each cation were placed into their own test tube, along with five drops of NaCl, and observations were recorded. Barium nitrate, copper nitrate, manganese nitrate, nickel nitrate, ferric nitrate, and aluminum nitrate did not react to the 3.0 M NaCl solution. However, when the NaCl was added to each test tube, lead (ii) nitrate reacted to create a solid white precipitate that completely separated from the solution, and silver nitrate reacted to create a milky white solution. The net ionic equations of these chemical reactions go in order as follows:

The same experiment was run, except instead of NaCl being added to each test tube containing a specific cation, five drops of an aqueous solution of 6 M Hydrochloric Acid (HCl) were added to each test tube. Only lead (ii) nitrate and silver nitrate reacted with the 6 M HCl, each of which produced a white solid precipitate. The net ionic equations for these chemical reactions go in order as follows:

Evidently, both lead (ii) nitrate and silver nitrate react similarly to the chloride ion  found within both NaCl and HCl to create lead (ii) chloride and silver chloride.

 

Table 1a. Summary of what happened after 5 drops of 3 M NaCl were added to each cation solution.

Cation	Drops of 3 M NaCl added	Color	Texture	Product
Barium	5	Clear	liquid	none
Copper	5	Clear	liquid	none
Ferric	5	Clear	liquid	none
Lead (ii)	5	clear/white	solid	PbCl2
Manganese	5	Clear	liquid	none
Nickel	5	Clear	liquid	none
Silver	5	milky white	aqu.solid	AgCl
Aluminum	5	Clear	liquid	none

 

Table 2a. Summary of what happened after 5 drops of 6 M HCl were added to each cation solution.

Cation	Drops of 6 M HCl	Color	Texture	Product
Barium	5	clear	liquid	none
Copper	5	clear	liquid	none
Ferric	5	clear	liquid	none
Lead (ii)	5	solid white	solid	PbCl2
Manganese	5	clear	liquid	none
Nickel	5	clear	liquid	none
Silver	5	solid white	solid	AgCl
Aluminum	5	clear	liquid	none

 

However, when the precipitates formed using NaCl were submerged in a hot bath and exposed to heat, the silver chloride melted into a clear solution, while the lead (ii) chloride did not. After that, a few drops of an aqueous solution of 6 M Nitric Acid ( were added to each solution, resulting in the silver chloride returning as a milky white precipitate, while the lead (ii) chloride did not react.

Table 1b. Summary of what happened after Silver and Lead (ii) were exposed to heat and 6 M HNO₃.

Cation	Drops of NaCl	Heated	Melted	Drops of 6 M HNO3	Precipitate
Lead(ii)	5	yes	no	3	Solid
Silver	5	yes	yes	3	Solid

 

When the precipitates created using HCl were exposed to heat, the lead (ii) chloride did not dissolve, while the silver chloride separated from the HCl to form a solid form of silver nitrate.

Table 2b. Summary of what happened after Silver and Lead (ii) were exposed to heat and 6 M HCl.

Cation	Drops of HCl	Heated	Melted	Drops of 6 M HCl	Precipitate
Lead (ii)	5	yes	no	5	solid
Silver	5	yes	yes	5	solid

 

• 1 M Na₂SO₄

 

Aqueous, 0.2 M samples of barium nitrate, copper nitrate, ferric nitrate, lead (ii) nitrate, manganese nitrate, nickel nitrate, silver nitrate, and aluminum nitrate were gathered, along with a dropper bottle full of 1.0 M sodium sulfate (Na₂SO₄). A few drops of each cation solution was placed into its own test tube, along with five drops of the Na₂SO₄. Silver nitrate, copper nitrate, ferric nitrate, manganese nitrate, aluminum nitrate, and nickel nitrate did not react with Na₂SO₄. However, Barium nitrate reacted with Na₂SO₄ to create a milky white solution, and lead (ii) nitrate reacted to create a gelatinous, milky white solution. The net ionic equations representing each reaction go as follows, in order:

Table 3a. Summary of what happened after 5 drops of 1.0 M Na₂SO₄ were added to each cation.

 

Cation	Drops of 1 M Na2SO4	Color	Texture	Product
Barium	5	milky white	solid	BaSO4
Copper	5	clear	liquid	none
Ferric	5	clear	liquid	none
Lead (ii)	5	milky white	gelatinous	PbSO4
Manganese	5	clear	liquid	none
Nickel	5	clear	liquid	none
Silver	5	clear	liquid	none
Aluminum	5	clear	liquid	none

 

From there, 1 drop of 6 M Nitric Acid (HNO₃) was added to both the barium sulfate (BaSO₄) and lead sulfate (PbSO₄), creating a milky white solution in each test tube. Then, 10 drops of 6 M sodium hydroxide (NaOH) were added to the test tube containing BaSO₄, and 5 drops of 6 M NaOH were added to PbSO₄. The BaSO₄ did not dissolve in the NaOH, but the PbSO₄ did. The equations for lead’s reaction to nitric acid are as follows:

Table 3b. Summary of what happened after 1 drop of 6 M HNO₃ and various amounts of NaOH were added to lead (ii) and barium.

Cation	Drops of 6 M HNO3	Precipitate	Drops of 6 M NaOH	Precipitate
Lead (ii)	1	white solid	5	none
Barium	1	white solid	10	white solid

 

The reaction took place in two parts because the lead sulfate did not react with the nitric acid to directly make a new product. Rather, lead sulfate reacted with nitric acid to first make a new solid, then that solid reacted with the excess  ions to create an aqueous, clear solution of .

Lastly, 5 drops of 3 M sulfuric acid (H₂SO₄) was added to the clear solution of , causing the solution to turn warm and white, then partially release into the air as a stinky gas. In other words, the sulfuric acid caused an exothermic reaction.

3 M NaOH:

Once again, the same eight aqueous, 0.2 M cation solutions were gathered, and a few drops of each reagent were put into its own test tube so that there were eight test tubes, each containing one of the cation solutions. Then, 3 drops of 3 M NaOH were added to each test tube. Aluminum nitrate, barium nitrate, nickel nitrate, and lead (ii) nitrate did not react to the NaOH. However, silver nitrate reacted with NaOH to create a dark green/brown precipitate, copper nitrate reacted to create a goopy, blue precipitate, ferric nitrate reacted to create a thick, rusty precipitate, and manganese nitrate reacted to create a creamy yellow precipitate. The net ionic equations representing each reaction, in order, are as follows:

Table 4a. Summary of what happened when 3 drops of 3 M NaOH were added to each cation.

Cation	Drops of 3 M NaOH	Color	Texture	Precipitate
Silver	3	brown/green	aqu. solid	AgO2
Aluminum	3	clear	liquid	None
Barium	3	clear	liquid	None
Copper	3	toothpaste blue	aqu. solid	Cu(OH)2
Ferric	3	rusty red	aqu. solid	Fe(OH)3
Manganese	3	yellow/white	solid	Mn(OH)2
Nickel	3	clear	liquid	None
Lead (ii)	3	clear	liquid	None

 

After that, the samples from each test tube were split such that there were eight test tubes, each of which contained one of the four precipitates. 4 more drops of 3 M NaOH were added to four of the test tubes, resulting in no changes to the precipitates. To the other four test tubes, 4 drops of 6 M HNO₃ were added, resulting in all four samples dissolving into clear liquid solutions. The equations representing these reactions in order are as follows:

Table 4b. Summary of what happened after 4 more drops of 3 M NaOH were added to the solid cations.

Cation	drops 3 M NaOH	Change	Precipitate
Silver	4	none	AgO2
Copper	4	none	Cu(OH)2
Ferric	4	none	Fe(OH)3
Manganese	4	none	Mn(OH)2

 

 

 

 

Table 4c. Summary of what happened after 4 drops of HNO3 were added to the solid cations.

Cation	drops of 6 M HNO3	Change	Precipitate
Silver	4	clear liquid	Ag+
Copper	4	clear liquid	Cu2+
Ferric	4	clear liquid	Fe3+
Manganese	4	clear liquid	Mn2+

 

3.0 M NH₄OH:

After eight more samples of eight different aqueous, 0.2 M cation solutions were gathered into eight different test tubes, 3 drops of 3.0 M aqueous ammonia (NH₄OH) were added to each tube. The silver nitrate, barium nitrate, and copper nitrate solutions stayed clear, indicating that no reaction had occurred. However, the aluminum nitrate reacted with the NH₄OH to form an opaque white precipitate, ferric nitrate reacted to form a slightly chunky, red-brown precipitate, manganese nitrate reacted to create a precipitate that looked like curdled milk, nickel nitrate reacted to form a cloudy teal precipitate, and lead (ii) nitrate reacted with NH₄OH to create a milky white precipitate. The net ionic equations for each, in order, are as follows:

Table 5a. Summary of what happened when 3 drops of 3 M aqueous ammonia were added to each cation.

Cation	Drops of NH4OH	Color	Texture	Precipitate
Silver	3	clear	Liquid	none
Aluminum	3	white	Solid	Al(OH)3
Barium	3	clear	Liquid	none
Copper	3	cobalt blue/clear	Liquid	complex ion
Ferric	3	goopy brown	aq. Solid	Fe(OH)3
Manganese	3	curdled milk	aq. Solid	Mn(OH)2
Nickel	3	cloudy blue	aq. Solid	Ni(OH)2
Lead (ii)	3	white	aq. Solid	Pb(OH)2

 

The precipitates were washed and centrifuged until most of the liquid supernatants had been removed and disposed of. Then, a 15 M aqueous ammonia solution was added to the precipitates. The iron (iii) hydroxide, aluminum (iii) hydroxide, manganese (ii) hydroxide, and lead (ii) hydroxide precipitates did not react to the 15 M NH₄OH. However, the nickel (ii) hydroxide reacted with 15 M NH₄OH to form a clear, blue aqueous solution, though it didn’t fully react because some of the nickel (ii) hydroxide remained as a solid. It is represented by the following net ionic reaction:

Table 5b. Summary of what happened when 3 drops of 15 M aqueous ammonia were added to each solid cation precipitate.

Cation	Drops of 15 M NH4OH	Change	Precipitate
Ferric	3	none	Fe(OH)3
Lead (ii)	3	none	Pb(OH)2
Nickel	3	clear blue	[Ni(NH3)6]+2
Aluminum	3	none	Al(OH)3
Manganese	3	none	Mn(OH)2

 

S^2- Reactivity:

A near-boiling hot bath was prepared while the same eight cation solutions were put into its own test tube, totaling eight test tubes. 2 drops of aqueous 6 M HNO₃ were added to each cation, before each cation solution was submerged in the hot bath for exactly four minutes. After the solutions were taken out of the hot bath, they were allowed to cool for five minutes. The silver nitrate solution produced a solid bronze precipitate at the bottom of the test tube, the copper nitrate solution produced a solid black precipitate, the ferric nitrate solution produced a solid white precipitate, and the lead (ii) nitrate produced a solid grey precipitate. The net ionic equations for these reactions, in order, are as follows:

 

Table 6a. Summary of what happened after 2 drops of 6 M HNO3 were added to each cation solution, which was then submerged in a hot bath.

Cation	Drops of 6 M HNO3	Minutes in hot bath	Precipitate	Color	Texture
Silver	2	4	Ag	bronze	solid
Copper	2	4	CuO	black	solid
Ferric	2	4	Fe(OH)3	white	solid
Lead (ii)	2	4	Pb(OH)2	grey	solid

 

The cations that did not produce a precipitate the first time were heated in the hot bath for another five minutes, but none of them produced a precipitate. To those solutions, 3 drops of 15 M NH₄OH were added. Upon doing so, the barium nitrate solution produced a solid black precipitate, the aluminum nitrate solution produced a solid bronze precipitate, the manganese nitrate solution produced a solid white precipitate, and the nickel nitrate solution produced a solid grey precipitate.

Table 6b. Summary of what happened after 3 drops of aqueous ammonia were added to the solutions that didn’t produce a precipitate the first try.

Cation	Drops 15 M NH4OH	Precipitate	Color	Texture
Barium	3	Ba(OH)2	black	solid
Aluminum	3	Al(OH)3	bronze	solid
Manganese	3	Mn(OH)2	white	solid
Nickel	3	Ni(OH)2	grey	solid

 

Cation Confirmatory tests:

To positively confirm the identities of all eight cations, each cation solution was subjected to a specific reagent that caused it to react in a unique way.

To start, two drops of 1.0 M aqueous potassium thiocyanate (KSCN) were added to three drops of 0.2 M aqueous ferric nitrate inside of a test tube, resulting in an aqueous solution that turned dark, blood red. The net ionic equation for this reaction is as follows:

A scoopula tip of sodium bismuthate (NaBiO₃) was mixed with 2 drops of 6 M HNO₃ to create a solution that was added to three drops of manganese nitrate inside of a test tube. The resulting reaction created a purple-tinted dark brown precipitate inside of an aqueous solution. The liquid was washed and centrifuged out of the precipitate. The net ionic equation for this reaction is as follows:

A few drops of 0.2 M aqueous lead (ii) nitrate was submerged in a hot bath. Once the lead (ii) nitrate solution had warmed up, 2 drops of 1 M potassium iodide (KI) were added to the solution, resulting in the formation of a thick, egg yolk yellow precipitate. The net ionic equation for this reaction is as follows:

2 drops of 1 M KI were added to a few drops of 0.2 M aqueous copper nitrate solution, resulting in a muddy, green-brown precipitate. The net ionic equation for this reaction is as follows:

Three drops of a 0.2 M aqueous silver nitrate solution were put inside of a test tube. Then, three drops of 15 M aqueous ammonia, three drops of 6 M aqueous HNO₃, and three drops of aqueous 1.0 M sodium thiosulfate (Na₂S₂O₃) were added to the solution. Each time a reagent was added to the silver nitrate solution, the resulting reaction was a smoky, exothermic reaction, but the aqueous solution within the test tube stayed clear the whole time. The net ionic equations for each reaction, in order, are as follows:

Three drops of 0.2 M aqueous barium nitrate were put into a test tube along with three drops of potassium chromate (K₂CrO₄), resulting in a milky, yellow precipitate. The net ionic equation for this reaction is as follows:

A drop of 0.2 M aqueous aluminum nitrate was put on a piece of litmus paper, to get an initial pH of 4. Four drops of aluminon  were added to three drops of the aluminum nitrate solution in a test tube. Then, a solution of 15 parts deionized water to one part 6 M HNO₃ was made and then used to bring the aluminum nitrate solution’s pH to 8. The resulting precipitate was an orange-red color. The net ionic equation for this reaction is as follows:

Lastly, three drops of all eight cation solutions, including 0.2 M aqueous nickel nitrate, were adjusted to a pH between 8-9, before three drops of dimethylglyoxime (C₄H₈N₂O₂) were added. The results and equations are as follows:

Nickel nitrate:  = gelatinous red/pink precipitate.

Barium nitrate: Clear solution, no reaction.

Copper nitrate:  = thick, camo green precipitate.

Silver nitrate:  = muddy lake brown precipitate.

Lead (ii) nitrate:  = banana shake yellow.

Manganese nitrate:  = chocolate brown.

Ferric nitrate:  = rusty orange.

Aluminum nitrate: Clear solution, no reaction.

The results from the cation confirmatory tests were organized in Table 7 below.

Table 7. Summary of results from cation confirmatory tests.

Cation	Reagent	Color	Texture	Precipitate
Ferric	1 M KSCN	blood red	aq. Solid	[FeSCN]2+
Manganese	NaBiO3	purple/brown	aq. Solid	2MNO4-
Lead (ii)	1 M KI	egg yolk	aq. Solid	PbI2
Copper	1 M KI	dark brown	aq. Solid	2CuI
Silver	Na2S2O3	smokey clear	liquid/gas	[Ag(S2O3)2]3-
Barium	K2CrO4	milky yellow	aq. Solid	BaCrO4
Aluminum	C22H23O9-	orange red	aq. Solid	AlC22H23
Nickel	C4H8N2O2	hot pink	gelatinous	Ni(C4H8N2O2)
Barium	C4H8N2O2	clear	liquid	none
Copper	C4H8N2O2	camo green	aq. Solid	Cu(C4H8N2O2)
Silver	C4H8N2O2	muddy brown	aq. Solid	Ag(C4H8N2O2)
Lead (ii)	C4H8N2O2	banana yellow	aq. Solid	Pb(C4H8N2O2)
Manganese	C4H8N2O2	chocolate brown	aq. Solid	Mn(C4H8N2O2)
Ferric	C4H8N2O2	rusty orange	aq. Solid	Fe(C4H8N2O2)
Aluminum	C4H8N2O2	clear	liquid	none

 

Separation Scheme:

Based on the results of the previous experiments, a separation scheme was devised to act as a guide when attempting to identify four cations in an unknown cation solution.

 

Figure 1. Separation Scheme that was created based on the known cation experiments.

 

 

Unknown identification:

A tube containing roughly 8 mL of an aqueous, green unknown solution identified only as 1246 was obtained, along with a dropper bottle containing 3 M NaCl. 3 drops of the unknown solution was mixed with 3 drops of 3 M NaCl in a test tube, resulting in a goopy, green precipitate being formed. The precipitate was washed with deionized water and centrifuged until only a white precipitate beneath a green supernatant remained. The supernatant was removed, and the white precipitate was exposed to heat, causing it to melt. It was hypothesized that the cation solution contained lead.

To identify the second cation in the unknown solution, three drops of 1246 were put into a test tube, and mixed with 5 drops of 1 M sodium sulfate (Na₂SO₄). The solution into a milky white precipitate, which indicated either lead, silver, barium, and/or aluminum based on the previous experiments. One drop of 6 M HNO₃ was added to the precipitate, which did not seem to cause a reaction. The sample was washed and centrifuged until only a solid white precipitate remained. Again, this did not help to narrow down any potential cations, so something else had to be done.

Within the fume hood, eight test tubes were each filled with three drops of the unknown solution 1246. 2 drops of 6 M HNO₃ were added to all eight test tubes, which were then submerged in a hot bath for several minutes. Upon cooling the test tubes, some precipitates had formed while others had not. The solutions that had formed precipitates were washed and centrifuged, leaving behind a solid mix of white, grey, and black precipitates. Those precipitates were set aside for later.

To the solutions that did not form any precipitates, 3 drops of 15 M aqueous ammonia were added, causing all of the test tubes to form a black precipitate.

To positively identify the cations within each sample of the unknown precipitate, 3 drops of dimethylglyoxime were added to each sample. The sample that turned black after being reacted with 15 M aqueous ammonia became pink and slightly gelatinous, confirming that nickel nitrate was one of the four cations present in the unknown solution 1246.

However, the other three samples that had produced precipitates were harder to identify due to them being mixed together. The assumption was made that the three remaining cations were lead (ii) nitrate, ferric nitrate, and copper nitrate due to the white, grey, and black mixture of precipitates.

5 drops of 6 M HCl were added to one test tube containing the grey, black, and white precipitates, resulting in a white precipitate being formed, while the supernatant turned slightly gold. The test tube was then warmed up, but the precipitate did not dissolve, confirming that it was lead (ii) nitrate within the solution, which was the second cation.

3 drops of 3 M NaOH were added to a second test tube containing the grey, black, and white precipitate, immediately resulting in a thick, rusty precipitate being formed, along with a blue supernatant. It was determined that the third cation was ferric nitrate.

Finally, 3 drops of 3 M ammonium hydroxide were added to a third test tube containing the grey, black, and white precipitate, resulting in a green/brown precipitate being formed, along with a cobalt blue, clear supernatant. This confirmed that copper nitrate was the fourth and final cation within the unknown solution.

In summary, unknown solution 1246 contained lead (ii) nitrate, copper nitrate, nickel nitrate, and ferric nitrate.

Below is the final separation scheme made based on these results.

 

Figure 2. The final separation scheme based on the results of the unknown solution identification strategy.