Chapter 19. Laboratory: Qualitative Analysis

Laboratory 19.1: Use Flame Tests to Discriminate Metal Ions

It has been known since antiquity that the presence of certain metals causes a flame to assume characteristic colors. For example, the presence of sodium causes a flame to assume a bright yellow color, and the presence of copper tints the flame a striking blue-green. This phenomenon occurs because heat excites atoms, boosting the outer electrons of their shells to higher energy states. As those electrons return to their original, lower energy states, the excess energy is released as photons at specific, characteristic wavelengths. The flame test uses this phenomenon to determine whether specific elements are present in a sample.

In theory, the flame test can discriminate numerous elements by their flame colors, some of which are listed in Table 19-1. In practice, it’s a bit more difficult. The first problem is ambiguity. If a flame test shows a bright red color, for example, does the sample contain lithium or strontium (or both)? Does a pale violet color indicate the presence of potassium in the sample, or is it cesium? Is the flame pale green (antimony or barium), bright green (boron), yellow-green (manganese or molybdenum), or just plain green (copper)? The second problem is masking. Some elements, notably sodium, produce an intense flame coloration if they are present in even tiny amounts. Most elements produce a much subtler coloration, which may be overwhelmed by the more intense color of another element.

So is flame testing useless for real qualitative analytical work? Yes and no. In a home laboratory, flame tests are useful primarily to illustrate the principles involved. There are much more accurate methods available for qualitative analysis of metal ions in a home lab, which we’ll explore in later sessions. Conversely, in professional laboratories, flame testing is one of the primary methods used for qualitative inorganic analysis. Professional labs use flame spectrometers, expensive instruments that can detect and unambiguously identify elements at the parts-per-million level by charting the emitted spectrum of a sample.

We don’t have a flame spectrometer, and you probably don’t either, so we’ll use the traditional flame test technique in this lab.

Table 19-1. Flame test colors by species

Species	Flame color	Notes
Ammonium	Faint green	Masked by most other species
Antimony	Pale blue-green to blue	Faint and easily masked
Arsenic	Light blue	Faint and easily masked
Barium	Pale green to yellow-green
Boron	Bright green
Calcium	Red	Calcium compounds show brick-red to orange to yellow-orange; masked by barium
Cesium	Pale violet	Easily masked
Copper(I)	Blue
Copper(II)	Green	Copper(II) halides show blue-green
Iron	Yellow
Lead	Blue to blue-white
Lithium	Crimson	Masked by barium or sodium
Magnesium	Bright white
Manganese(II)	Yellow-green
Molybdenum	Yellow-green
Phosphates	Blue-green to blue	When moistened with sulfuric acid
Potassium	Pale violet to lilac	Crimson when viewed through cobalt glass; potassium compound pink to lilac to violet; masked by sodium or lithium
Sodium	Bright yellow
Strontium	Scarlet	Masked by barium
Zinc	Blue-green to green-white

Caution

Concentrated hydrochloric acid is corrosive. Wear splash goggles, gloves, and protective clothing.

Procedure

If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Fill a test tube about half full of concentrated hydrochloric acid and place it in the rack.
Light the gas burner and adjust it to produce a small, hot flame.
Hold the loop at the tip of the flame and heat it until the loop adds no color to the flame.
Test the hydrochloric acid for purity. Dip the loop into the hydrochloric acid, allow it to drain momentarily, and then hold the loop at the tip of the flame. It should add no coloration to the flame. If color appears, particularly the intense yellow color of sodium, the hydrochloric acid is insufficiently pure. Replace it with purer acid.
Before you test each sample, verify that the loop is clean by holding it in the flame. If any color appears in the flame, clean the loop by dipping it in the hydrochloric acid and heating it. If necessary, repeat the acid rinse and heat until no color appears. (Hydrochloric acid reacts with any contaminants present on the loop to form chloride salts, which have low boiling points; heating the loop vaporizes the chloride salts, removing them from the loop.)
Figure 19-1. A positive flame test for sodium
Test the first sample by touching the loop to the sample and holding it in the flame. If the sample is solid, dip the clean loop into the hydrochloric acid, allow it to drain momentarily, and then touch the loop to the sample. If the sample is liquid, dip the clean loop into the liquid. Record your observations in Table 19-2. If a sample does not produce a clear coloration, repeat the test using a larger amount of the sample.
Repeat step 7 for each of your samples.

Table 19-2. Using flame tests to discriminate metal ions—observed data

Sample	Observations
Sodium-free salt (example)	Medium yellow coloration, indicating presence of sodium; pale blue-violet potassium coloration visible through cobalt glass
A. Barium
B. Boron
C. Calcium
D. Copper(I)
E. Copper(II)
F. Copper(II) halide
G. Iron(II)
H. Iron(III)
I. Lead
J. Lithium
K. Magnesium
L. Manganese
M. Potassium
N. Sodium
O. Strontium
P. Zinc

Dr. Paul Jones Comments

You can also dissolve the salt in methanol (in a watch glass) and ignite the methanol. The flame will burn with the color of whatever is in there. This is also a nice way of testing which elements swamp others.

Disposal

Dilute the waste hydrochloric acid with a few hundred mL of water and flush it down the drain.

Review Questions

Q1: Road safety flares produce a brilliant red light. Which element do you think is used to produce that coloration?

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Q2: A pyrotechnic produces brilliant green stars. Salts of what element or elements might be included in the pyrotechnic mixture to produce this effect?

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Q3: During the Christmas season, some retailers stock fireplace logs that produce colorful flames. How might you produce such logs yourself from ordinary firewood?

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Laboratory 19.2: Use Borax Bead Tests to Discriminate Metal Ions

The borax bead test is a fast, sensitive test for the presence of chromium, copper, cobalt, gold, iron, manganese, nickel, and tungsten. The borax bead test works because very low concentrations of the oxides of these metals in their oxidized and reduced states impart characteristic colors to a transparent borax bead. (Table 19-3 lists the characteristic colors.) Although it is seldom used nowadays for serious qualitative analysis work, the borax bead test was formerly used frequently by geologists and others for field tests of ore-bearing rock. It requires little equipment, and is sufficiently sensitive to use as a screening test.

Table 19-3. Borax bead colors by element, flame type, and bead temperature

Element	Reducing flame (hot bead)	Reducing flame (cold bead)	Oxidizing flame (hot bead)	Oxidizing flame (cold bead)
Chromium	Green	Green	Yellow to yellow-green	Green
Cobalt	Blue	Blue	Blue	Blue
Copper	Colorless	Red to reddish brown	Green to light blue-green	Blue
Gold	Red	Violet	Rose to violet	Rose to violet
Iron	Pale green to green	Green	Brown to yellow	Yellow
Manganese	Colorless	Colorless	Violet	Violet
Nickel	Gray	Gray	Violet to brown	Reddish brown
Tungsten	Green	Blue	Pale yellow	Colorless

Caution

Concentrated hydrochloric acid is corrosive. Wear splash goggles, gloves, and protective clothing.

Procedure

If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Heat the loop portion of a platinum or Nichrome inoculating loop in a gas burner flame until it is red hot.
Dip the hot loop into borax powder (sodium tetraborate decahydrate, Na₂B₄O₇ · 10H₂O), a small amount of which adheres to the hot loop. Heat the adhering borax in the hottest part of the gas burner flame. At first, the borax swells as it loses its water of crystallization. As you continue to heat the anhydrous sodium tetraborate, it melts and shrinks, forming a colorless, transparent, glassy bead.
Allow the bead to cool for a moment, dip it in distilled water to moisten it, and then dip it into the powdered sample, a few grains of which will adhere to the bead.
Disposal
Dispose of all waste materials by flushing them down the drain with plenty of water. (The tiny amounts of metals contained in the beads are not hazardous.)
Reheat the bead in the reducing (inner, blue) part of the gas burner flame until the bead remelts. Observe the color of the bead, if any, while it is hot. The bead should remain transparent. If it becomes cloudy or opaque, you’ve used too much of the sample. Only a grain or two is needed.
Allow the bead to cool and again observe the color.
Reheat the bead in the oxidizing (outer, colorless) part of the gas burner flame until the bead remelts. Observe the color of the bead, if any, while it is hot.
Allow it to cool and again observe the color.
Record your observations from steps 5 through 8 in Table 19-4.
When you finish the test, reheat the bead and plunge the loop into cold water to remove the bead.

Table 19-4. Using borax bead tests to discriminate metal ions—observed data

Element	Reducing flame (hot bead)	Reducing flame (cold bead)	Oxidizing flame (hot bead)	Oxidizing flame (cold bead)
A. Chromium
B. Cobalt
C. Copper
D. Iron
E. Manganese
F. Nickel

Review Questions

Q1: How closely did the bead colors you observed for various metal samples correspond to the colors listed in Table 19-3? Propose an explanation for any significant differences.

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Q2: Copper chromate is used as wood preservative for “pressure-treated” wood. What bead colors would you expect if you applied the borax bead test to a sample of copper chromate?

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Q3: Copper ferrocyanide is used in water analysis and treatment. What bead colors would you expect if you applied the borax bead test to a sample of copper ferrocyanide?

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Laboratory 19.3: Qualitative Analysis of Inorganic Anions

Qualitative inorganic analysis is generally done in two phases. In one phase, the sample is tested to determine which anions (negatively charged ions, usually nonmetals) are present. In the other phase, the sample is tested to determine which cations (positively charged ions, usually metals) are present.

Chemists have developed specific tests for all common anions. These tests depend on chemical reactions that occur between the anion being tested for and one or more reagents. A positive test results in some detectable change—a precipitate, color change, odor, evolution of gas, and so on. Some anion tests yield positive results if any of a group of anions is present. If that test is negative, none of the anions in that group can be present. If that test is positive, a further test is done to determine which anion on anions from that group is present in the sample.

In this lab, we’ll use the following tests to identify these eight common anions:

Nitrate (NO₃^–)

The nitrate anion is identified using the brown ring test. Iron(II) (ferrous) ions are added to a portion of the sample in aqueous solution. Concentrated sulfuric acid is then added to the sample by carefully allowing it to run down the side of the test tube. Because concentrated sulfuric acid is very dense, it settles to the bottom of the test tube, under the aqueous layer that contains the sample and ferrous ions. At the phase boundary between the aqueous layer and the sulfuric acid layer, a brownish or purple-brown ring forms if nitrate ion is present in the sample. Testing for nitrate ions must be done on a separate portion of the sample, because the nitrate test introduces sulfate ions into the sample and the other anion tests introduce nitrate ions in the form of nitric acid to the sample.

Sulfite (SO₃^2–)

Acidifying a solution that contains sulfite ions causes evolution of sulfur dioxide gas, which has a characteristic sharp, choking odor familiar to anyone who has smelled the smoke from a firecracker. The presence of sulfur dioxide can be confirmed by its decolorizing effect on a dilute solution of potassium permanganate.

Carbonate (CO₃^2–)

Acidifying a solution that contains carbonate ions causes evolution of carbon dioxide gas. Although carbon dioxide is odorless, its presence can be confirmed by testing with a solution of barium hydroxide, which turns cloudy in the presence of carbon dioxide.

Chloride (Cl^–), Bromide (Br^–), and Iodide (I^–)

Adding silver nitrate solution to an acidic solution that contains halide ions causes the corresponding silver halide to precipitate. Silver chloride is white, silver bromide cream-colored, and silver iodide pale yellow, but it can be difficult to discriminate by color alone, particularly if the sample contains a mix of halide ions. The particular silver halide or halides present can be confirmed by testing with aqueous ammonia. A 6 M ammonia solution dissolves silver chloride, but not the bromide or iodide salt. A 15 M ammonia solution dissolves silver bromide, but not silver iodide. (Actually, the halides are not dissolved, but instead form soluble complexes with the ammonium ions.)

Sulfate (SO₄^2–)

Adding barium nitrate to an acidic solution that contains sulfate ions causes white barium sulfate to precipitate.

Phosphate (PO₄^3–)

Phosphate ions in solution with nitric acid are precipitated by the addition of ammonium molybdate, which reacts with the phosphate ions to form the insoluble complex salt (NH₄)₃PO₄ · 12MoO₃.

There are dozens of other common anions, each of which has a corresponding test. For example, thiocyanate (SCN^–) behaves as a pseudo-halide anion, with characteristic reactions similar to the chloride, bromide, and iodide ions. The presence of thiocyanate ions can be confirmed by adding Fe³⁺ (ferric) ions, which combine with thiocyanate ions to form the blood-red ferrithiocyanate [Fe(SCN)(H₂O)₅]²⁺ complex. Similarly, the presence of the hexacyanoferrate(III) ([Fe(CN)₆]^3–) anion (usually called ferricyanide) can be confirmed by adding a solution that contains Fe²⁺ (ferrous) ions, which combine with ferricyanide ions to form a characteristic dark-blue precipitate of the pigment Prussian Blue.

Review Questions

Q1: The precipitates produced by the anion tests take some time to settle. What alternative procedures might you use to minimize the time necessary?

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Q2: Why did we wash the precipitate in step 10?

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Q3: If silver nitrate were unavailable to you, what alternative reagent or reagents might you use to detect and discriminate among the halide anions? (Hint: look up the solubility of the various halide salts to find other insoluble halides.)

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Laboratory 19.4: Qualitative Analysis of Inorganic Cations

Qualitative analysis of inorganic cations refers to a structured method of analyzing an unknown solid or solution to determine whether specific cations are present. Such an analysis is an iterative process of separating cations into groups based on reactions characteristic to each group and then further analyzing each separated group to determine which specific cations in that group are present in the sample. Cations are conventionally grouped as follows:

Group I cations (Ag⁺, Pb²⁺, Hg⁺)

Group I cations are the only cations that produce insoluble chlorides. Group I cations can be precipitated by adding dilute hydrochloric acid, leaving all other cations in solution.

Group II cations (Cu²⁺, Bi³⁺, Cd²⁺, Hg²⁺, As³⁺, Sb³⁺, Sn⁴⁺)

Group II cations produce extremely insoluble sulfides (K_sp < 10^–30). Group II cations can be precipitated by adding sulfide ions at low concentration, which is most conveniently achieved by adding thioacetamide to the acidic solution. In aqueous solution, thioacetamide produces hydrogen sulfide gas. Because hydrogen sulfide is a very weak acid, the equilibrium in acidic solution stays far to the left, with most of the hydrogen sulfide present in molecular form rather than dissociated to form sulfide ions.

Group III cations (Al³⁺, Cr³⁺, Fe³⁺, Zn²⁺, Ni²⁺, Co²⁺, Mn²⁺)

Group III cations produce slightly soluble sulfides (K_sp > 10^–20). Group III cations can be precipitated by adding sulfide ions at high concentration, which is most conveniently achieved by adding aqueous ammonia to the acidic thioacetamide solution used to precipitate Group II cations until the solution is strongly basic. The addition of ammonia drives the equilibrium to the right, forcing more of the aqueous hydrogen sulfide gas produced by the thioacetamide to dissociate, forming sulfide ions.

Group IV cations (Mg²⁺, Ca²⁺, Sr²⁺, Ba²⁺)

Group IV cations produce insoluble carbonates. Group IV cations can be precipitated by adding carbonate ions.

Group V cations (Na⁺, K⁺, NH₄⁺)

Group V cations are not precipitated by chloride, sulfide, or carbonate ions, and so are the only cations remaining in solution after an unknown solution has been treated with those reagents.

In a standard qualitative cation analysis, the first step is to treat the unknown with chloride ions, which causes all Group I cations to precipitate out. The precipitate and supernatant liquid are then separated by filtration or centrifugation. The liquid at this point contains only Group II, III, IV, and V cations. The precipitate, which may contain any or all of the Group I cations, is subsequently treated with hot water. Lead chloride, which is insoluble in cold water, is reasonably soluble in hot water, and so can be separated from any silver chloride and/or mercury(I) chloride present in the precipitate by removing the supernatant fluid from the precipitate by filtration or centrifugation. The supernatant fluid is tested with potassium chromate solution. If lead(II) ions are present, they combine with the chromate ions to produce insoluble lead chromate, which precipitates as a solid with a characteristic yellow color, confirming the presence of lead(II) ions in the sample.

The precipitate from the first step, which may contain silver chloride and/or mercury(I) chloride, is then treated with aqueous ammonia, with which silver chloride forms a soluble complex. Any solid precipitate remaining after treatment with aqueous ammonia must be mercury(I) chloride, and so confirms the presence of mercury(I) ions in the sample. To confirm the presence of silver, the aqueous ammonia is neutralized with nitric acid. If a precipitate occurs, the presence of silver ions in the sample is confirmed.

The next step is to separate Group II ions, which form extremely insoluble sulfide salts. The supernatant liquid from the first step is made strongly acid, and thioacetamide is added to the solution. In acidic solution, thioacetamide produces a low concentration of sulfide ions, which causes the Group II ions to precipitate. Although we don’t do so in this book, that precipitate can be further analyzed to isolate and identify the specific Group II ion or ions present in the sample. The supernatant liquid can contain only Group III, Group IV, and Group V cations.

That solution is treated with aqueous ammonia until it is strongly basic, which causes the thioacetamide to dissociate further, providing a higher concentration of sulfide ions. The higher sulfide concentration causes Group III cations—whose sulfides are more soluble than Group II sulfides—to precipitate, leaving only Group IV and Group V cations in solution. Again, the Group III precipitate can be further analyzed to isolate and identify the specific Group III ion or ions present in the sample, though that is not covered in this book.

Finally, the supernatant liquid, which can contain only Group IV and Group V cations, is treated with carbonate ion, which causes Group IV cations to precipitate and leaves only Group V cations in solution. The Group IV precipitate is further analyzed to isolate and identify the specific Group IV ion or ions present in the sample. The solution, which can now contain only Group V cations, can be further analyzed to isolate and identify the specific Group V ion or ions present in the sample.

Order is critical when you separate cations. For example, the carbonate ions used to precipitate Group IV cations also precipitate all of the Group I, Group II, and Group III cations. If you treat an unknown solution with carbonate ions first, you precipitate all Group I through Group IV ions without achieving any separation.

In a real qualitative inorganic analysis, any or all of these cations might be present in the unknown, and the separation and analysis would be done using microscale or semi-micro procedures. For simplicity and to avoid using expensive reagents, in this lab we’ll instead observe the reactions of individual cations with various reagents and build a matrix of the observable changes that occur with various combinations of cation and reagent.

We’ll use only five primary reagents (six, counting distilled water): 3 M sulfuric acid, and 6 M solutions of hydrochloric acid, sodium hydroxide, aqueous ammonia, and nitric acid. We’ll also use three secondary reagents to confirm the presence of some cations: 0.25 M solutions of potassium ferricyanide, potassium ferrocyanide, and potassium thiocyanate.

Procedure

If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Use the graduated cylinder to transfer about 5 mL of 0.1 M aluminum nitrate solution to each of seven test tubes in a rack.
Add a few drops of 3 M sulfuric acid to the first test tube and swirl or stir the tube gently. If no observable change occurs, note that fact in Table 19-5 and continue to the next step. If adding sulfuric acid produces a visible change such as a precipitate, color change, bubbling, or some other action, continue adding sulfuric acid until no further change occurs. If a precipitate occurs, add aqueous ammonia to neutralize the sulfuric acid (until a drop of the solution in the test tube turns litmus paper turns blue) and then add the same volume of ammonia again to make the ammonia in large excess. Determine whether the precipitate dissolves in excess ammonia. Note your observations in Table 19-5.
Caution
Hydrochloric, nitric, and sulfuric acids and concentrated aqueous ammonia are corrosive. Some of the metal salts used in this lab session are poisons, oxidizers, corrosives, or otherwise hazardous. Some are known or suspected carcinogens. Read the MSDS for each chemical before you use it. Wear splash goggles, gloves, and protective clothing. Wear a disposable N100 respirator mask if you handle any of the hazardous chemicals in solid form. Dispose of all chemical wastes properly, in accordance with hazardous material disposal laws and regulations.
Add a few drops of 6 M hydrochloric acid to the second test tube and swirl or stir the tube gently. If no observable change occurs, note that fact in Table 19-5 and continue to the next step. If adding hydrochloric acid produces a visible change such as a precipitate, color change, bubbling, or some other action, continue adding hydrochloric acid until no further change occurs. If a precipitate occurs, continue adding hydrochloric acid with stirring until you have added twice the amount of acid required to reach the first endpoint and see if the precipitate redissolves. If the precipitate does not redissolve, neutralize the hydrochloric acid to litmus paper with aqueous ammonia and continue until the ammonia is in large excess to determine whether the precipitate dissolves in excess ammonia. Note your observations in Table 19-5.
Add a few drops of 6 M sodium hydroxide to the third test tube and swirl or stir the tube gently. If no observable change occurs, note that fact in Table 19-5 and continue to the next step. If adding sodium hydroxide produces a visible change such as a precipitate, color change, bubbling, or some other action, continue adding sodium hydroxide until no further change occurs. If a precipitate occurs, continue adding sodium hydroxide with stirring until you have added twice the amount of sodium hydroxide required to reach the first endpoint and see whether the precipitate redissolves. If the precipitate does not redissolve with excess sodium hydroxide, test it with excess aqueous ammonia and then with excess nitric acid. Note your observations in Table 19-5.
Add a few drops of 6 M aqueous ammonia to the fourth test tube and swirl or stir the tube gently. If no observable change occurs, note that fact in Table 19-5 and continue to the next step. If adding ammonia produces a visible change such as a precipitate, color change, bubbling, or some other action, continue adding ammonia until no further change occurs. If a precipitate occurs, continue adding ammonia with stirring until you have added twice the amount of ammonia required to reach the first endpoint and see whether the precipitate redissolves. If the precipitate does not redissolve with excess ammonia, test it with excess nitric acid. Note your observations in Table 19-5.
Add a few drops of 0.25 M potassium ferricyanide to the fifth test tube and swirl or stir the tube gently. If no observable change occurs, note that fact in Table 19-5 and continue to the next step. If adding potassium ferricyanide produces a visible change such as a precipitate or color change, continue adding potassium ferricyanide until no further change occurs. If no visible change occurs, add one or two drops of nitric acid to the test tube and note any visible change. If acidifying the solution causes no change, add three or four drops of sodium hydroxide and note any visible changes. Note your observations in Table 19-5. See Figure 19-4.
Add a few drops of 0.25 M potassium ferrocyanide to the sixth test tube and swirl or stir the tube gently. If no observable change occurs, note that fact in Table 19-5 and continue to the next step. If adding potassium ferrocyanide produces a visible change such as a precipitate or color change, continue adding potassium ferrocyanide until no further change occurs. If no visible change occurs, add one or two drops of nitric acid to the test tube and note any visible change. If acidifying the solution causes no change, add three or four drops of sodium hydroxide and note any visible changes. Note your observations in Table 19-5.
Add a few drops of 0.25 M potassium thiocyanate to the seventh test tube and swirl or stir the tube gently. If no observable change occurs, note that fact in Table 19-5 and continue to the next step. If adding potassium thiocyanate produces a visible change such as a precipitate or color change, continue adding potassium thiocyanate until no further change occurs. If no visible change occurs, add one or two drops of nitric acid to the test tube and note any visible change. If acidifying the solution causes no change, add three or four drops of sodium hydroxide and note any visible changes. Note your observations in Table 19-5.
Discard all solutions and precipitates in the hazardous waste container, and wash all of the test tubes thoroughly.
Repeat steps 1 through 10 for each of the other cation sample solutions.

Table 19-5. Qualitative analysis of inorganic cations—observed data

Cation	H₂SO₄ (3 M)	HCl	NaOH	NH₃	K₃[Fe(CN)₆]	K₄[Fe(CN)₆]	KSCN
Example	White ppt; not soluble in excess ammonia	White ppt; not soluble in excess HCl; soluble in excess ammonia; re-ppt’d by excess nitric acid	Black ppt; not soluble in excess NaOH; soluble in excess ammonia; re-ppt’d by excess nitric acid	Black ppt; soluble in excess ammonia; re-ppt’d by excess nitric acid
Al3+
Ag+
Ba2+
Ca2+
Co2+
Cr3+
Cu2+
Fe2+
Fe3+
Mn2+
Ni3+
Pb2+
Sr2+
Zn2+

Cautions: Cyanide Gas!

Depending on the reaction conditions, mixing strong acids with ferricyanide, ferrocyanide, or thiocyanate salts may cause deadly hydrogen cyanide gas to be evolved. The amounts we use are much too small and the solutions too dilute to present any hazard, but be very careful when handling these salts in solid form around strong acids.

Disposal

Dispose of all waste materials by pouring them into the hazardous waste 2-liter soft drink bottle. When you complete this lab session, make up a saturated solution of sodium carbonate (washing soda). Add the sodium carbonate solution to the soft drink bottle until no further precipitation occurs. Decant off the supernatant liquid and flush it down the drain with plenty of water. The precipitate contains small amounts of barium, chromium, and lead. Depending on your local environmental laws and regulations, you may be permitted to dispose of this precipitate with your ordinary solid household waste, or you may need to take it to your local hazardous waste disposal facility.

Figure 19-4. A blue precipitate upon addition of potassium ferricyanide confirms the presence of ferrous ions

Review Questions

Q1: Examine the matrix that you filled in for Table 19-5. Which, if any, of the cations cannot be unambiguously identified using only the five primary reagents?

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Q2: If any of the cations cannot be unambiguously identified using only the five primary reagents, can they be identified using one or more of the secondary reagents? If so, which secondary reagents are useful for identifying which cations?

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Q3: If you have a sample that is known to contain only iron(II) and/or iron(III) and/or copper(II) cations, what is the minimum number and identity of the reagents needed to identify unambiguously the cations present in the unknown? What characteristic reactions would serve to discriminate these three cations?

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Laboratory 19.5: Qualitative Analysis of Bone

Animal bone is a complex structure that incorporates both organic and inorganic components. Most of the inorganic component, which makes up about 70% of bone mass, is hydroxylapatite, which is produced by the body but also occurs naturally as a mineral. Hydroxylapatite has the empirical formula Ca₅(PO₄)₃(OH), but the formula is usually written as Ca₁₀(PO₄)₆(OH)₂, because the crystalline structure of hydroxylapatite is bimolecular. In bone, the hydroxyl ion is sometimes replaced by chloride, carbonate, or other anions. Other inorganic and organic components of bone include smaller amounts of other anions and cations.

In this laboratory, we’ll apply what we’ve learned in the preceding lab sessions to analyze a sample of bone for several anions and cations.

Caution

Hydrochloric acid, nitric acid, and sodium hydroxide are corrosive. Some of the salts used in this lab session are poisons, oxidizers, corrosives, or otherwise hazardous. Read the MSDS for each chemical before you use it. Wear splash goggles, gloves, and protective clothing.

Procedure

This laboratory has two parts. In Part I, we’ll prepare the bone sample for analysis. In Part II, we’ll determine whether eight specific ions are present in the sample: chloride, sulfate, phosphate, calcium, iron(III), sodium, ammonium, and potassium.

Part I: Sample Preparation

In Part I, we prepare the bone sample for analysis by digesting it in dilute nitric acid. This procedure should be done outdoors, under an exhaust hood, or in an otherwise well-ventilated area. The reaction of nitric acid with the bone sample evolves toxic and irritating fumes of nitrogen oxides. Do not breathe these fumes or allow them to contact your eyes or skin.

If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Weigh a 1.0 g sample of bone. If possible, use small pieces rather than one large chunk.
Add 25 mL of 6 M nitric acid to 25 mL of distilled or deionized water in a 150 mL beaker. Place the beaker on a support ring and wire gauze over the gas burner.
Add the bone sample to the beaker and heat the beaker gently, with stirring, to dissolve the bone sample. Do not boil the solution. As the bone sample dissolves, toxic fumes are evolved.
Continue heating and stirring the solution in the beaker for 10 or 15 minutes. As the reaction progresses, the quantity of fumes evolved rapidly diminishes. Some solid matter remains undissolved.
Remove the heat and allow the beaker to cool to room temperature.
Set up a funnel with filter paper and the second 150 mL beaker as the receiving vessel.
Filter the solution into the second beaker and discard the filter paper and solid filtrand. The filtrate is your sample.

Part II: Analysis of Ions

In Part II, we analyze the sample to determine whether chloride, sulfate, phosphate, calcium, iron(III), sodium, ammonium, and/or potassium ions are present.

If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Transfer 5 mL of the sample to the first test tube. Add several drops of 0.1 M silver nitrate solution. A precipitate confirms the presence of chloride ions. (Actually, a precipitate confirms the presence of chloride, bromide, or iodide ions, but the bromide and iodide ions are present in extremely low concentrations in animal bone, so we can assume that any precipitate is caused by chloride ions.) Note your observations on line A of Table 19-6.
Transfer 5 mL of the sample to the second test tube. Add 10 drops of 0.1 M barium chloride to the test tube, stir or swirl to mix the solution, and look for a white precipitate that confirms the presence of sulfate ions. Record your observations on line B of Table 19-6.
Transfer 5 mL of the sample to the third test tube. Add about 4 mL of 6 M nitric acid and 6 mL of ammonium molybdate to the test tube and mix the solution well. A yellow precipitate confirms the presence of phosphate ion. If no precipitate occurs immediately, continue observing the test tube for a minute or two. If there is still no precipitate, heat the test tube very gently for a minute or two. (Do not heat the solution strongly. Doing so may cause a precipitate of white molybdenum trioxide, which is not a positive test for phosphate ions.) Record your observations on line C of Table 19-6.
Transfer about 3 mL of the sample to the fourth test tube. Add about 15 mL of 0.2 M ammonium oxalate and mix the solutions thoroughly. A white precipitate or white cloudiness confirms the presence of the calcium ion. (Under the acidic conditions of this test, calcium oxalate is more soluble than it is in neutral solution. Examine the test tube carefully; even a slight white cloudiness is a positive result.) Record your observations on line D of Table 19-6.
Transfer 5 mL of the sample to the fifth test tube. Add 5 mL of 6 M nitric acid and mix thoroughly. Add 5 drops of 0.1 M potassium thiocyanate, mix thoroughly, and look for a color change. A blood red color (Figure 19-5) indicates the presence of Fe(III) ion in relatively high concentration. At lower concentrations of Fe(III) ion, the solution assumes anything from a light red color to a very pale straw yellow color. Record your observations on line E of Table 19-6.
Figure 19-5. Ferric ions react with potassium thiocyanate to produce a blood red complex
Dip the end of the platinum or Nichrome inoculating loop into concentrated hydrochloric acid and then hold it in the gas burner flame until the loop adds no color to the gas flame. Dip the clean loop into the sample, and then hold it in the gas flame. A bright yellow flame coloration confirms the presence of sodium ions in the sample. Record your observations on line F of Table 19-6.
Transfer 5 mL of the sample to the sixth test tube. Add 3 mL of 6 M sodium hydroxide solution to the test tube and mix thoroughly. Carefully sniff the test tube by using your hand to waft any vapors from the mouth of the test tube toward your nose. An odor of ammonia confirms the presence of ammonium ion in the sample. If no odor is obvious, gently warm the test tube for a minute or so, periodically testing for the scent of ammonia. Record your observations on line G of Table 19-6.
Continue heating the test tube until it comes to a very gentle boil. Continue the gentle boil for a minute or two to drive off all of the ammonium ions as gaseous ammonia. Allow the test tube to cool and then add about 3 mL of the sodium cobaltinitrite reagent. A precipitate of insoluble potassium cobaltinitrite confirms the presence of potassium ions in the sample. Record your observations on line H of Table 19-6.

Disposal

Dispose of all waste materials by flushing them down the drain with plenty of water.

Table 19-6. Qualitative analysis of bone—observed data

Ion	Reagent	Observations
A. Chloride(Cl^–)	Silver nitrate
B. Sulfate (SO₄^2–)	Barium chloride
C. Phosphate (PO₄^3–)	Ammonium molybdate
D. Calcium (Ca²⁺)	Ammonium oxalate
E. Iron(III) (Fe³⁺)	Potassium thiocyanate
F. Sodium (Na⁺)	Flame test
G. Ammonium (NH₄⁺)	Sodium hydroxide
H. Potassium (K⁺)	Sodium cobaltinitrite

Review Questions

Q1: Considering the results of your tests, name and give formulae for the compounds that you believe may be present in the bone sample.

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Q2: If a sample produces a precipitate when treated with silver nitrate, how would you proceed to establish unambiguously that the precipitate is silver chloride rather than silver bromide or silver iodide?

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Q3: Is it possible to test the dissolved bone sample to determine whether carbonate ion was present in the original solid sample? If not, explain why and devise a test procedure to determine whether

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	goggles, gloves, and protective clothing
	test tube
	test tube rack
	gas burner
	inoculating loop (platinum or Nichrome)
	cobalt glass square
	hydrochloric acid, concentrated (~10 mL)
	samples (see Substitutions and Modifications)

	goggles, gloves, and protective clothing
	gas burner
	inoculating loop (platinum or Nichrome)
	borax (~1 g)
	samples (see Substitutions and Modifications)

	goggles, gloves, and protective clothing
	gas burner
	test tubes (6)
	test tube clamp
	ring stand
	stirring rod (2)
	anion sample solution (see Substitutions and Modifications)
	sulfuric acid, 3 M (a few drops)
	sulfuric acid, concentrated (~2 mL)
	iron(II) sulfate (~2 g)
	nitric acid, 3 M (~15 mL)
	potassium permanganate, 0.02 M (one drop)
	barium hydroxide, saturated (one drop)
	silver nitrate, 0.1 M (a few mL)
	aqueous ammonia, 6 M (a few mL)
	aqueous ammonia, 15 M (a few mL)
	barium nitrate, 0.1 M (a few mL)
	ammonium molybdate, 0.5 M (a few mL)

	goggles, gloves, and protective clothing
	2-liter soft drink bottle labeled “hazardous waste”
	test tubes (7)
	test tube rack
	eye dropper or Beral pipette (5)
	graduated cylinder, 10 mL
	stirring rod (1 or more)
	litmus paper (red and blue)
	cation sample solutions (see Substitutions and Modifications)
	primary and secondary reagents (see Substitutions and Modifications)

	goggles, gloves, and protective clothing
	balance and weighing papers
	beaker, 150 mL (2)
	graduated cylinder, 10 mL
	eye dropper or Beral pipette (5)
	test tubes (6)
	test tube rack
	test tube clamp
	gas burner
	ring stand
	support ring
	wire gauze
	inoculating loop
	stirring rod
	funnel and filter paper
	litmus paper (red and blue)
	bone sample (see Substitutions and Modifications)
	nitric acid, 6 M (~30 mL)
	silver nitrate, 0.1 M (a few mL)
	barium chloride, 0.1 M (1 mL)
	ammonium molybdate, 0.1 M (5 mL)
	ammonium oxalate, 0.2 M (15 mL)
	potassium thiocyanate, 0.1 M (a few drops)
	hydrochloric acid, concentrated (~1 mL)
	sodium hydroxide, 6 M (5 mL)
	sodium cobaltinitrite reagent (see Substitutions and Modifications)