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Qualitative Analysis

Chemical analysis can be either qualitative or quantitative in character. A qualitative analysis enables us to find out what elements or chemical species are present, while a quantitative analysis tells us how much of each component is present. Quantitative analysis comprises a one- semester course usually taken during the sophomore year. Qualitative analysis is usually offered as part of the laboratory requirement for the second semester of General Chemistry. In this course, six weeks will be devoted to “qual” in lab.

The methods and techniques you will learn are the classical “wet” methods, whereas, modern analytical labs rely on instrumentation almost exclusively. Technicians with little background in chemistry can be readily trained to do routine instrumental analyses without knowing anything about the chemical principles underlying the technique. You, however, upon completion of qualitative analysis will have learned a lot of descriptive chemistry and applied many of the important chemical principles that were introduced in lecture. “Qual” presents a perfect opportunity to study the properties of metal ions in aqueous solution, chemical equilibrium, precipitation reactions, acid-base chemistry, complex ion formation, oxidation-reduction reactions, and much more. You will also sharpen your powers of observation and enhance your ability to critically evaluate data. This may sound like a monumental task, but the great majority of students find “qual” to be their most enjoyable experience in chemistry lab. In fact, no other laboratory experience in general chemistry comes as close to providing a feeling for the flavor of scientific research.

A good research chemist must have a clear understanding of the problem he or she is trying to solve and the methods used to approach the problem. Good chemists are usually good cooks, but the reverse is not necessarily true. Don't take a “cookbook” approach to qualitative analysis, simply following a set of instructions by rote. You most probably will encounter unexpected results, ambiguity, and uncertainty. Only a real familiarity with the chemistry involved in your analysis will allow you to work through puzzling results and be successful.

Your actual laboratory work in qualitative analysis will consist of analyzing a “known” solution containing ten different metal cations. You will also analyze three unknown cation solutions. Following this you will carry out tests to identify the anion in nine “known” sodium salts, as well as two unknown sodium salts. Lastly, all of the laboratory technique and analytical chemistry that you have learned will be integrated when you correctly identify both the anion and cation in an unknown salt. An approximate schedule is:

Week 1 Cation group 1 and 2 knownWeek 2 Cation group 1 and 2 unknownWeek 3 Cation group 3 and 4 known and cation group 3 and 4 unknownWeek 4 General cation unknownWeek 5 Anion knowns and two anion unknownsWeek 6 Salt unknown

Qualitative analysis 4-1

The Qualitative Analysis Notebook

The notebook must have the pages sewn to the binding. Papers are never removed from a lab notebook. All information must be recorded in ink. If you make a mistake in the notebook, simply cross through it with a single line. If the mistake is still readable, you can see where you goofed and avoid the same mistake later. Record observations in detail. Except for the reactions, all information must be put in the notebook as you go along. Do not write on loose paper.

The notebook is more than a record of what you do in the lab; it can help when doing an unknown. Ignoring the notebook during an unknown causes many errors. For example; the lab manual says that the precipitate confirming the presence of Al3+ is red (brown shades). You will find that description is not accurate. Many students have missed the presence of Al3+ in their unknowns because they look only at the flow chart and not their notebook. When doing an unknown, if you are unsure of something, check your notebook. It will contain the same test carried out on a known solution. Believe your notebook!

It's to your advantage that you record your observations in detail. When describing colors, don't just say “green”. Is it “seasick green”, “light green”, “dark green”, “olive” or “pond scum green”? Estimation as to the amount of a precipitate is also helpful. This can help when faced with very small quantities of precipitates.

When starting a known or an unknown, the first thing to record is the appearance of the solution or solid. If it is an unknown, make some conclusion as to the components of the unknown based on color. Of the ten cations in this scheme, only four are colored. A yellow unknown must contain Fe3+! What else could give the unknown this color?

Set up your notebook as shown below for cation analysis, with the four columns spread across two pages.

Procedure Observations Conclusions ReactionTake 1 mL of known solution

Dark blue-green liquid

Known contains Ag+, Pb2+, Cr3+, Al3+, Fe3+, Mg2+, Ba2+, Cu2+, Ni2+, Zn2+

Add 2d HCl to known; stir & centrifuge

Soln 2A: dark green. Ppt 1A: white crystalline

Ag+ and/or Pb2+ indicated in ppt 1A.Soln 2A may contain any of the other 8 cations and PbCl4

2-.

Ag+ + Cl- AgCl(s)Pb2+ + 2Cl- PbCl2(s)Pb2+ + 4Cl- PbCl4

2-(aq)

Important guidelines for keeping a notebook: Before your first day of Qual Lab do steps 1, 2, and 3!

1. Number at least 30 pages in the upper right-hand corner.


2. Include a Title Page with course name and section, your name, phone number and email address (for returning lost notebooks), and your instructor’s name.3. Include a Table of Contents with the following topics; add page numbers as you do each part.

Group 1 and 2 known Group 1 and 2 unknownGroup 3 and 4 known Group 3 and 4 unknownGeneral cation unknownAnion known Anion unknown 1Anion unknown 2 Salt unknown

4. Refer to all solutions and precipitates by number and letter, i.e., solution 2A or precipitate 1A.5. Save lots of room for the reactions – there may be 7 or 8 of them when only 1 reagent is added. Use a ruler to draw a horizontal line after each step and its reactions!6. Always stir and centrifuge before recording observations about precipitates. Colorless and clear do not mean the same thing. (A blue solution may be clear—i.e., transparent.)7. Do not recopy anything in your notebook. Get in the habit of recording notes neatly. The notebook must be an original contemporaneous record. Copying observations into the notebook from scratch paper is definitely forbidden--poor science. On the other hand, summarizing results after completing a known and unknown series is encouraged. 8. Your notebook may be collected and graded at any time (unannounced) during the semester.

Colors In The Qualitative Analysis Scheme For Cations:

It is useful to memorize the colors used for identifying ions and precipitates. When two or more colors are present, the darker one will mask the lighter one.

Simple ions: Ag+, Al3+, Mg2+, Ba2+, Zn2+, & Pb2+ colorlessFe3+ yellow Cr3+ dark blueCu2+ pale blue Ni2+ pale green

Complex ions: Ag(NH3)2+, Zn(NH3)4

2+, Zn(OH)42+, & Al(OH)4

- colorlessCu(NH3)4

2+ deep blue CrO42- yellow

I3- golden FeSCN2+ blood red

Ni(NH3)62+ light violet

Precipitates: AgCl, PbCl2, Al(OH) 3, Mg(OH)2, BaSO4 WhitePb(OH)2 tannish-white, very pale tanMg(NH4)PO4 white, crystallineCuI tan or white Fe(OH) 3 brownCr(OH) 3 blue-gray Ni(OH)2 pale-greenPbCrO4 bright yellow Al(OH) 3 . aluminon1 cherry-red lake

BaCrO4 pale yellow Mg(OH)2 . magneson2 blue lakeNi(DMG)2

3 pinkish red

1 Aluminon is the triammonium salt of aurintricarboxylic acid, C22H23N3O9·2 Magneson is 4-(4-nitrophenylazo)resorcinol, NO2C6H4N=NC6H3-1,3-(OH) 23 Dimethylglyoxime is (CH3) 2C2 (NOH)2


Laboratory Techniques:

The objective of qualitative analysis is the identification of the ions present in a sample. The usual procedure is to run a carefully selected series of reactions on a solution containing all the ions of a particular group. This solution is called a “known” solution or “control.” The behavior of an unknown, which contains one or more of the ions of the group, is then compared with that of the known. The occurrence of certain reactions and the absence of others allows us to determine which ions are present in the unknown. Such identification of ions requires careful laboratory work.

The procedures that we will use in the laboratory are on a semimicro scale. Semimicro procedures involve considerably smaller volumes of solutions than are normally employed in ordinary laboratory work; solution volumes are normally on the order of 1 ml, and solution concentrations are on the order of 0.1 M. Semimicro methods have the advantages of requiring only small amounts of chemicals and of permitting rapid separation procedures. When working on the semimicro level, most reactions can be performed in small test tubes that are about 3-inches high (10 x 75 mm) and hold 3 milliliters. Solutions are separated from precipitates by centrifuging the solid to the bottom of the tube and then carefully removing the liquid with a dropper. Tedious filtrations are avoided.

Standard operating procedures for semi-micro qualitative analysis are discussed below. These procedures and techniques are generally used throughout the qualitative analysis scheme and will not be described in detail when they occur.

1. Cleanliness: If a known or unknown sample comes in contact with any surface that has not been thoroughly cleaned, it will be contaminated and the analytical results will be unreliable. It is important that all glassware be kept clean (not necessarily dry) during qual lab. Bring in a clean dish towel to cover your work area. Label two beakers, “clean” and “dirty”, and use these for stirring rods, droppers, test tubes, etc. Always rinse glassware twice with deionized water and use only deionized water whenever instructions call for water. When a sample is contaminated it should be discarded.

2. Labeling: Label all solutions and precipitates immediately with the designations used in the lab procedure; i.e., 2A, 3B, etc. Use these same labels in the lab notebook.

3. Dispensing Solutions: Volumes of solutions can usually be estimated. You should have a plastic dropper with a 1 mL mark etched into it for your convenience. It is also helpful to determine how many drops from your glass droppers are equal to 1 mL (usually around 20 drops). You should also be aware of the level of a milliliter of liquid in your 10 x 75 mm test tubes so that you can readily estimate that volume.

When adding reagents from dropper bottles (either your own or those in the hoods) be careful not to contaminate the dropper: do not put the tip into your solution or touch the inside of the test tube with the dropper tip. And always place the dropper back into the reagent bottle immediately. Never lay the dropper on a contaminated bench top.


Clear liquid

Centrifuged mixture Add 1 drop of reagent

Formation of more precipitate shows that precipitation was not complete

If a reagent bottle does not have a dropper, pour a small amount into a clean test tube and use one of your own droppers to dispense it from the test tube. Never place your own dropper into a reagent bottle that is used by others.

4. Testing Acidity/Basicity: Usually when instructions call for the addition of acid or base to a solution it is important to add enough to change the solution from acidic to basic or vice versa. A litmus test is used to detect this change. Blue litmus changes to red in acid, and red litmus changes to blue in base. When adding acid or base to your sample, stir with a clean stirring rod. Do not dip litmus paper into your solution in the test tube since this may contaminate your sample. Use the stirring rod to transfer a drop of the solution to a small piece of litmus paper and note when the color change occurs.

5. Heating Solutions: Keep a small beaker of water on a hot plate with a boiling stone in the beaker to prevent bumping. To heat a sample, place your test tube in the beaker of hot water.

6. Precipitation of Solids: To accomplish a precipitation, add the indicated amount of precipitating agent to the solution and stir well with a clean stirring rod. Heat the solution in a water bath if so directed. Some precipitates form slowly and must be given time to form completely. When precipitation is believed to be complete, centrifuge the sample. Generally, before the liquid above the solid is removed, it is tested for complete precipitation. This test is accomplished by adding another drop of the precipitating reagent. If the precipitation is not complete, additional precipitate will form in the liquid. In such a case, add a few more drops of precipitating reagent, stir thoroughly, centrifuge, and then retest for complete precipitation.

Testing a solution for completeness of precipitation

7. Centrifugation: This is the principle method used in qualitative analysis to separate a solid from a liquid. A centrifuge spins the sample at high speeds which forces the precipitate to the bottom of the tube. It is very important to balance the centrifuge contents before turning it on. Tubes placed in the centrifuge must be symmetrically distributed. If necessary, use a tube with a similar volume of water to balance another tube. Most precipitates will settle when centrifuged for 30-60 seconds at medium speed.


8. Removing Supernatants: The clear solution above the precipitate after centrifugation is called the supernatant. Removal of the supernatant can be accomplished by decanting or by using a capillary pipet. Decantating involves merely pouring the liquid off into another test tube without disturbing the precipitate. However, if the precipitate is very light it may not be possible to decant successfully. In this case, the liquid may be transferred to a clean test tube using a long, thin dropper called a capillary pipet. It is usually not necessary to remove the liquid entirely.

9. Washing Precipitates: Precipitates will always be contaminated with ions from the supernatant liquid, which may interfere with subsequent tests run on them. Therefore, it is necessary to wash precipitates with deionized water before proceeding. Small amounts of water (1 ml) are added to the precipitate and stirred well. The mixture is centrifuged and the wash water discarded. The washing procedure is then repeated. Two washings are usually sufficient.

10. Reprecipitation: It is sometimes necessary, especially with gelatinous precipitates to dissolve the precipitate and then cause it to form again. This will result in a purer precipitate. This is commonly done with the gelatinous hydroxides of metal ions. The precipitate is first dissolved completely in acid. Then the clear solution is made basic to litmus paper and the hydroxides will precipitate out of solution once again.

Qual Equipment:

Obtain the following equipment for your locker:10 small test tubessmall test tube racksmall stirring rodsmall test tube brushnew 1-mL plastic droppers

Make sure the test tubes and stirring rod are clean! If not, wash with soap and water and rinse twice with deionized water.


Part I - Cation Analysis

Reagents

Reagents used in the cation analysis are listed below. Acids and the concentrated ammonia are located in the hoods. Other reagents are available on the side benches.

BaCl2 (0.2 M) KSCN (0.2 M)NaOH (dilute, 6 M) Na2HPO4 (0.2 M)KI (2 M) K4Fe(CN)6 (0.2 M)K2CrO4 (0.2 M) NH4NO3 (5 M)HCl (concentrated, 12 M) H2SO4 (dilute, 6 M) NH3 (same as NH4OH)(concentrated, 15 M) Aluminon reagentHNO3 (conc., 15 M) Magneson reagentH2O2 (3%) Dimethylglyoxime (1%)Na2SO3 (saturated)

Overview

Ten cations will be studied in qualitative analysis. The cations can be separated into four groups, based on solubilities of their various salts:

Group 1, the silver group: Ag+, Pb2+

Group 2, the aluminum group: Al3+, Fe3+, Cr3+

Group 3, the alkaline earth group: Ba2+, Mg2+

Group 4, the copper group: Cu2+, Ni2+, Zn2+

You will begin cation analysis by working with a “known” solution of all cations. This known solution actually is a 0.05 M solution of the nitrates of all 10 cations. The NO3

- ion, therefore, is present but does not interfere in any way with the analysis and can be ignored.

The known solution will first be analyzed for the Group 1 and Group 2 cations. It is important that the known be run carefully and that all results be recorded immediately in the lab notebook. Descriptions of the relative amounts of precipitates as well as accurate colors of precipitates and solutions are important! You will next be given an unknown containing between one and five of the cations in Groups 1 and 2. Repeat the same procedure, running a new known sample and an unknown sample in parallel. Caution! The analysis of unknown solutions sometimes leads to observations of very small quantities of precipitates and/or faint positive color tests for ions. Quite often these “trace” amounts are just that: small amounts of impurities or precipitates resulting from contamination or from failure to separate interfering cations adequately. In deciding whether the test is positive for an ion, you should compare the quantity of precipitate (or intensity of color) obtained with the unknown and that obtained with the known; these may not be exactly the same, but should be comparable.

Upon completion of the Group 1 and 2 unknown, you will proceed to analyze the remainder of the original known solution for Groups 3 and 4 cations. Following this you will analyze your


second unknown containing between 1 and 5 of the Group 3 and 4 cations. You will then receive a general unknown which may contain any of the 10 cations.

“Divide and Conquer” will be your plan of attack for the cations. You will separate the four groups of ions from each other and then separate and identify the ions within each group. The separations used in this qualitative analysis procedure are based on the following:

a. the solubility rules found in your textbookb. the fact that some insoluble metal hydroxides are amphoteric (dissolve in base as well

as in acid)c. the fact that some transition metal ions form soluble ammine complexes with NH3.

Therefore the principles of acid-base-, solubility-, and complex-ion- equilibria are important in understanding the separations that follow. Let us look at the amphoteric hydroxides and the ammine complexes before taking a detailed look at the chemistry of each group in the cation separation scheme.

Aluminum, chromium, and zinc hydroxides are amphoteric. The following reactions occur when these hydroxides are placed in acid solution:

Al(OH)3(s) + 3H+(aq) Al3+(aq) + 3H2O(l)Cr(OH)3(s) + 3H+ (aq) Cr3+(aq) + 3H2O(l)Zn(OH)2(s) + 2H+ (aq) Zn2+(aq) + 2H2O(l)

In basic solutions, these same hydroxides react as follows:

Al(OH)3(s) + OH-(aq) Al(OH)4-(aq)

Cr(OH)3(s) + OH- (aq) Cr(OH)4- (aq)

Zn(OH)2(s) + 2OH- (aq) Zn(OH)42-(aq)

The silver ion and the three Group 4 cations form soluble ammine complexes with NH3:

Ag+(aq) + 2NH3 (aq) Ag(NH3)2+ (aq)

Cu2+(aq) + 4NH3 (aq) Cu(NH3)42+(aq)

Ni2+ (aq) + 6NH3 (aq) Ni(NH3)62+(aq)

Zn2+ (aq) + 4NH3 (aq) Zn(NH3)42+(aq)

After completing the cation analysis, a different approach will be taken to identify anions. The analysis of anions will not be based on a separation scheme. Rather, a series of simple tests will be done in parallel on each known and unknown anion and the results recorded. Each anion unknown will then be matched to one of the known anions.


The Chemistry Behind The Cation Qualitative Analysis Scheme

Group 1: Ag+, Pb2+

Ag+ and Pb2+ are separated from the remaining eight cations in the known solution by precipitation as the chlorides:

Ag+ + Cl- AgCl(s)Pb2+ + 2Cl- PbCl2(s)

Pb2+ + 4Cl- PbCl42-(aq)

The chlorides of the other ions in this scheme are soluble. Not all of the lead will precipitate, since PbCl2 is slightly soluble and the complex ion, PbCl4

2-, forms. Therefore, there will be some Pb2+ in the solution containing the other eight ions.

Separation of silver and lead is based on the fact that PbCl2 becomes appreciably soluble in hot water (0.99 g/100 mL at 20 °C and 3.34 g/100 mL at 100 °C); AgCl does not (2 x 10-4 g/100 mL at 20 °C and 1 x 10-3 g/100 mL at 100 °C). Thus, hot water will extract PbCl2 from the precipitate.

After the separation, Pb2+ is confirmed by the bright yellow precipitate of PbCrO4.

Pb2+(aq) + CrO42-(aq) PbCrO4(s)

The AgCl residue from the extraction is soluble in aqueous ammonia, producing the diamminesilver ion, Ag(NH3)2

+.

AgCl(s) + 2NH3(aq) Ag(NH3)2+(aq) + Cl-(aq)

When HNO3 is added to the above solution the equilibrium is shifted to the left because of the removal of NH3.

NH3 + H+ NH4+

The result is that AgCl is reprecipitated and this precipitate confirms the presence of Ag+ ion.

Group 2: Fe3+, Cr3+, Al3+

After removing Ag+ and most of the Pb2+ from the original solution it still contains Fe3+, Cr3+, Al3+, Ba2+, Mg2+, Cu2+, Ni2+, Zn2+ and a trace of Pb2+. The addition of ammonia to this solution raises the pH to 9 and causes several of the ions to precipitate out as hydroxides while others preferentially complex with the ammonia (see the reactions in the Overview). Recall that NH3 is a weak base due to the following equilibrium:

NH3 + H2O NH4+ + OH-


Al3+, Fe3+, Cr3+, and the remaining Pb2+ precipitate as insoluble hydroxides.

Al3+ + 3OH- Al(OH)3(s)Fe3+ + 3OH- Fe(OH)3(s)Cr3+ + 3OH- Cr(OH)3(s)Pb2+ + 2OH- Pb(OH)2(s)

Table 2 shows the Ksp values of metal hydroxides. (A small Ksp value indicates that the compound is less soluble (or, more insoluble).) Note that the Ksp values for the hydroxides of Cu2+, Zn2+, and Ni2+ are smaller than that for Pb2+. Why do these metal hydroxides not precipitate out also? The answer lies in the fact that the transition metal ions, Cu2+, Zn2+ and Ni2+ form stable ammine complexes in the presence of excess NH3 rather than forming the hydroxide precipitates. The formulas of the complex ions are Cu(NH3)4

2+, Zn(NH3)42+, Ni(NH3)6

2+. The reactions for forming these ions were given in the overview above.

Table 2. Ksp Values of Metal HydroxidesAl(OH)3 5 x 10-33

Fe(OH)3 6 x 10-33

Cr(OH)3 7 x 10-31

Cu(OH)2 1.6 x 10-19

Zn(OH)2 5 x 10-17

Ni(OH)2 1.6 x 10-16

Pb(OH)2 4 x 10-15

Mg(OH)2 8.9 x 10-12

Ba(OH)2 5 x 10-3

At pH 9, the [OH-] is 1 x 10-5 M. Mg(OH)2 will precipitate only if the [OH-] is greater than about 1.3 x 10-5 M; Ba(OH)2 is also too soluble to precipitate in this solution. This solution from the precipitation of Group 2 therefore contains Mg2+, Ba2+, and the ammine complexes of Cu2+, Ni2+, and Zn2+. This solution (deep blue) is saved for analysis for Group 3 and 4 cations later.

Now, back to the hydroxide precipitates:

First, H2SO4, is added to this mixture of iron, aluminum, chromium, and lead hydroxides. This results in the dissolution of the hydroxides of Fe3+, Al3+, and Cr3+, and precipitation of PbSO4 which is removed. Then aqueous ammonia is added to the solution, precipitating Fe(OH)3, Al(OH)3, and Cr(OH)3 again. The overall effect of these two steps is the removal of lead, and the reprecipitation of Fe(OH)3, Cr(OH)3, and Al(OH)3 with fewer impurities.

Al(OH)3(s) + 3H+ Al3+ + 3H2OCr(OH)3(s) + 3H+ Cr3+ + 3H2OFe(OH)3(s) + 3H+ Fe3+ + 3H2OPb(OH)2(s) + 2H+ Pb2+ + 2H2O

Pb2+ + SO42– PbSO4(s)


The equations for the reprecipitation of the hydroxides are the same as was shown on the previous page. Next, iron, chromium, and aluminum must be separated. The mixture of Fe(OH)3, Al(OH)3, and Cr(OH)3 is treated with H2O2 and NaOH. Fe(OH)3 is not amphoteric and is recovered unchanged. Al(OH)3 dissolves forming Al(OH)4

–; chromium is oxidized to soluble CrO4

2–.

Al(OH)3(s) + OH- Al(OH)4–

Cr(OH)3(s) + OH- Cr(OH)4–

2Cr(OH)4– + 3HO2

- 2CrO42– + 5H2O + OH-

Recall that Al(OH)3 and Cr(OH)3 are amphoteric, dissolving in both acid and base as the above reactions show.

The iron(III) hydroxide precipitate is separated and dissolved in acid, KSCN is added, and iron is confirmed by the blood-red thiocyanatoiron(III) ion, FeSCN2+.

Fe3+ + SCN- FeSCN2+

To separate chromium and aluminum, NH4NO3 is added to the solution of Al(OH)4– and CrO4

2-. Al(OH)4

– is converted to Al(OH)3 by the weak acid NH4+. (A strong acid would have produced

Al3+. Addition of the NH4+ ion reduces the [OH-] and shifts the Al(OH)3/Al(OH)4

– equilibrium to the left to reprecipitate Al(OH)3). Chromate remains in solution unchanged.

NH4+ + OH- NH3 + H2O

Al(OH)3(s) + OH- Al(OH)4-(aq)

The aluminum hydroxide is removed and chromium is confirmed by addition of BaCl2 precipitating the pale yellow BaCrO4.

Ba2+ + CrO42– BaCrO4(s)

The Al(OH)3 is dissolved in acid, aluminon and aqueous ammonia are added to form a “red lake”. A lake is formed when a dye is adsorbed from solution by a gelatinous precipitate. A chemical reaction does not occur between the Al(OH)3 and the aluminon dye. In this case, the dye, aluminon, is preferentially adsorbed by Al(OH)3. It is not adsorbed by Cr(OH)3.

Group 3: Ba2+, Mg2+

At the beginning of the Group III analysis the known solution is deep blue and contains Ba2+, Mg2+, Cu(NH3)4

2+, Ni(NH3)62+, and Zn(NH3)4

2+. Sulfuric acid is added to precipitate BaSO4, confirming the presence of the Ba2+ ion.

Ba2+ + SO42- BaSO4(s)


The three transition metal complexes decompose also.

Cu(NH3)42+ + 4H+ Cu2+ + 4NH4

+

deep blue pastel blue

Ni(NH3)62+ + 6H+ Ni2+ + 6NH4

+

violet pale green

Zn(NH3)42+ + 4H+ Zn2+ + 4NH4

+

The solution, therefore, after removal of the Ba2+ by precipitation, is pale blue-green and contains Mg2+, Cu2+, Ni2+, and Zn2+. Addition of an excess of aqueous ammonia reforms the ammine complexes (see reactions in the Overview above). It is important to add enough ammonia to convert all of the Ni2+ into Ni(NH3)6

2+. Otherwise Ni(OH)2 may precipitate out in the next step along with Mg(OH)2 and a separation will not have been achieved. The strong base NaOH is added and a gelatinous white precipitate of Mg(OH)2 forms.

Mg2+ + 2OH- Mg(OH)2(s)

HCl is added to dissolve the precipitate.

Mg(OH)2 + 2H+ Mg2+ + 2H2O

NH3 and Na2HPO4 are added to the solution next and a white, finely divided precipitate of MgNH4PO4 slowly forms.

NH3 + HPO42- NH4

+ + PO43-

Mg2+ + NH4+ + PO4

3- Mg(NH4)PO4(s)

This precipitate is then dissolved in HCl.

Mg(NH4)PO4(s) + 2H+ Mg2+ + NH4+ + H2PO4

-

To confirm magnesium, an organic dye (known as Magneson or magnesium S & O reagent) is added and the solution is made basic with NaOH. Mg(OH)2 precipitates and the dye is adsorbed by the Mg(OH)2 forming a blue precipitate that is referred to as a “blue lake”. The supernatant solution should be pinkish-purple at this stage.

Group 4: Cu2+, Ni2+, Zn2+

The known cation solution now contains only the ammine complexes of Cu2+, Ni2+, and Zn2+. The solution is deep blue. This basic solution is acidified with HCl in order to decompose the three ammine complexes once again. Addition of KI causes the solution to turn a golden color due to the formation of the I3

- ion. A tan precipitate of CuI also forms. The following oxidation-reduction reaction occurs, in which Cu2+ is reduced to Cu+ and I- is oxidized to I2.


CH3

C=N

OH

C=N

H3C OH

2 + Ni(NH3)62+

CH3

C=N

O

C=N

H3C O–H

H–O

N=C

CH3

N=C

O CH3

Ni

2Cu2+ + 4 I- 2CuI(s) + I2

tan

The elemental iodine thus formed reacts with excess iodide ion to form the ion, I3-, which is

golden-brown in water. I- + I2 I3

-

Next the CuI precipitate is oxidized back to Cu2+ using H2O2, in the presence of ammonia. The deep blue color of Cu(NH3)4

2+ confirms the presence of Cu2+ in the known.

2CuI + H2O2 2Cu2+ + 2OH- + 2I-

Cu2+ + 4NH3 Cu(NH3)42+

The solution that was separated from the CuI precipitate is golden in color and contains Ni2+, Zn2+, and I3

-. Na2SO3 acts as a reducing agent to reduce the I3- back to the colorless iodide ion, I–.

SO32- + I3

- + H2O SO42- + 3I- + 2H+

The I- ion will not interfere in any subsequent reactions.

The pale green solution containing Ni2+ and Zn2+ is treated with excess strong base, NaOH. Pale green Ni(OH)2 precipitates out. The Zn(OH)2 that also forms, being amphoteric, dissolves in the excess base to form the Zn(OH)4

– complex ion.

Ni2+ + 2OH- Ni(OH)2(s)Zn2+ + 2OH- Zn(OH)2(s)

Zn(OH)2(s) + 2OH- Zn(OH)42-

In order to confirm the presence of Ni2+, the pale green Ni(OH)2 is first dissolved in HCl.

Ni(OH)2 + 2H+ Ni2+ + 2H2OThe addition of NH3 forms the ammine complex.

Ni2+ + 6NH3 Ni(NH3)62+

An organic reagent, dimethylglyoxime (abbreviated DMG), is added and a pinkish-red precipitate forms. The precipitate is a coordination compound of Ni2+ and dimethylglyoxime.

The solution containing Zn(OH)42- is acidified with HCl.


+ 2NH4+ + 4NH3

Zn(OH)42- + 4H+ Zn2+ + 4H2O

Potassium ferrocyanide solution, K4Fe(CN)6, is added and a precipitate of Zn3K2[Fe(CN)6]2

forms.

3Zn2+ + 2K+ + 2Fe(CN)64- Zn3K2[Fe(CN)6]2(s)

This precipitate is white when pure, however, other cations present in trace amounts can form colored precipitates with this reagent. Frequently the Zn3K2[Fe(CN)6]2 formed from the known will appear gray-green in color.


Part II - Anion Analysis

Nine anions will be considered in this analysis. All are colorless in solution. Cl– Br– I–

SO32– SO4

2– CO32–

C2H3O2– PO4

3– NO3–

Unlike the cation analysis you will not be working with solutions containing mixtures of anions. The anion knowns are simply the sodium salts of the above nine anions in solid form. Each of your unknowns will consist of only one of the above salts, not a mixture. Your “salt” unknown will also contain only one of the above anions.

The anion analysis involves treating each of the nine knowns and your three unknowns with several different reagents in a series of spot tests. More than one anion may give a similar reaction with a given reagent, but by considering the results of all the tests, each of the anions included in the scheme for analysis can be distinguished.

Set up a table in your notebook covering two facing pages as shown below, listing the ions being tested (in the order given) in the far left-hand column. You must have 12 horizontal rows going across the two pages and 4 vertical columns on each page. Head each successive column with the test performed.

Anion 6M H2SO4 18M H2SO4

AgNO3 HNO3 Ba(OAc)2 HCl Special reactions

Cl-

Br-

I-

SO32-

SO42-

CO32-

C2H3O2-

PO43-

NO3-

Unk #1Unk #2Salt unk

Procedure for Anion Analysis:

You will be working down each column (not across each row), i.e., test each of the 11 anions (save the unknown salt anion until later) with 6 M H2SO4 first, recording all observations


immediately after testing each one. Then test each anion with 18 M H2SO4, etc. This strategy permits you to make direct comparisons of the unknowns with the knowns on each test.

Observations: After adding a reagent to a tube, thoroughly stir the contents of the tube before recording your observations in the appropriate block. Your observations must be as complete and detailed as possible. Write small!

a. Note whether the solid dissolves or not.b. Note any color changes in the solid or solution.c. Are any bubbles formed? Is a gas evolved? Note its color and odor. d. Does the test tube become warm? Is heat given off? e. Record the color of any precipitate that forms when two clear solutions are mixed. If you are testing solids, as in part A, precipitate formation is not a possibility! f. If no change occurs with a particular reagent write NR.

A. Begin with 11 small test tubes, cleaned and labeled for each anion and the two anion unknowns. (You will repeat these tests for the anion of your salt unknown later.) Place about 1/8 inch of each solid in a micro test tube.

1. 6 M H2SO4: Add 2 drops of 6 M H2SO4. Stir and record your observations immediately after each test.

Caution: Never place your nose directly over the mouth of a test tube to determine odors. Gently wave your hand over the top of the test tube, wafting any gas toward your nose.

Important: For C2H3O2-, the acetate ion, check the odor of the solution after adding

H2SO4 by sniffing the stirring rod used to mix the two.

2. 18 M H2SO4: Empty the tubes from A1 in the sink, rinse with deionized water and repeat the previous procedure using concentrated, 18 M H2SO4. B. Tests B1 and B3, and C1 and C2 are run using a solution of each anion. Prepare the 11 anion solutions by placing about 1/8 inch of solid on the tip of a spatula into a micro test tube and adding 2-3 mL of distilled water. Stir. If the salt does not all dissolve, centrifuge, decant the clear solution and use it for subsequent tests, discarding the undissolved solid. These solutions will be enough for the remaining tests. Do not use all of the solution for one test!

1. 0.2 M AgNO3: Place 2 drops of the anion solution in a micro test tube and add 3-4 drops of 0.2 M AgNO3. Mix well and record the results. If a precipitate forms, centrifuge and decant, saving the precipitate for the next step.

2. 3.0 M HNO3: To the precipitates formed in step 1 add 10 drops of 3.0 M HNO3 and stir well. Record whether or not the precipitate dissolves.

3. 0.2 M Ba(C2H3O2)2: Place 5 drops of anion solution in a micro test tube. Add 2-3 drops of 0.2 M Ba(C2H3O2)2. Mix well and record your results. If a precipitate forms, centrifuge and decant, saving the precipitate for the next step. 4. 2.0 M HCl: Add several drops of 2.0 M HCl to the precipitates formed in step 3. Mix well


and record whether or not the precipitates dissolve.

C. Special Tests: These are to be performed only on the particular known anions for which each test was designed and on your unknowns if you think you may have one of these anions.

1. Test for NO3-: Since all the tests for NO3

- are negative, the following positive test is recommended. Place 2 drops of the salt solution in a micro test tube. Carefully add 10 drops of 18 M H2SO4. Mix thoroughly and cool. Carefully add 3-4 drops FeSO4 solution, allowing the latter to float on top of the H2SO4 solution. Allow to stand for two minutes. A brown coloration at the junction of the two layers confirms the presence of the nitrate anion. The reactions which occur in this test are first reduction of NO3

– to NO by Fe2+ at high [H+],

NO3– + 4H+ + 3Fe2+ NO + 3Fe3+ + 2H2O

followed by formation of the brown FeNO2+ complex by reaction of NO with unoxidized Fe2+,

Fe2+ + NO FeNO2+

2. Test for Cl-: Place 2 drops of salt solution in a micro test tube. Add 0.2 M AgNO3 dropwise to the solution until no further precipitate forms. Centrifuge the mixture, discarding the decantate. Wash the precipitate twice with 5-drop portions of water, with 6 M discarding the wash solutions. Add 2 mL of H2O, 2 drops of 6 M aqueous ammonia, and 2 drops of 0.20 M AgNO3. Stir thoroughly for 2 minutes. To test the solution for Cl-, remove any remaining precipitate by centrifugation and decantation. Acidify the decantate with 6 M HNO3. The formation of a white precipitate (AgCl) indicates the presence of chloride ion. Dispose of all of these tubes in the silver waste.

3. Test for C2H3O2-: Place a sample of solid about the size of a grain of rice in a test tube.

Add 3 drops of 18 M H2SO4 and mix thoroughly. If effervescence is evident, heat the mixture in a boiling water bath until gas evolution ceases. Then add 4 drops of ethyl alcohol, C2H5OH, stir, and heat for 1 minute in the water bath. Fan the vapors toward your nose. An odor (ethyl acetate) similar to that of the solvent in airplane glue indicates the presence of acetate ion. Compare the odor with that of an authentic specimen of ethyl acetate found on the side bench with the other reagents.

4. Test for SO32-: Place a drop of deep purple KMnO4 on a Q-tip. Place a small amount of

solid salt in a test tube and add 2 or 3 drops of 6 M H2SO4. Carefully insert the Q-tip into the test tube without touching the sides of the tube or the solution in the tube. The evolving SO2 gas is a good reducing agent that will reduce the purple MnO4

- ion to the pale pink Mn+2 ion. Hint: Many students mistakenly report SO4

2- when the unknown is actually SO32-. Run this

test on SO42- also to see what a negative test looks like.


Part III - Analysis of a Soluble Salt

We consider here the analysis of a pure salt to determine the identity of the cation and anion present. The samples contain only one type of cation and one type of anion, each being a member of the corresponding set of ions already considered in Parts I and II. Further, all the salts that will be analyzed are soluble in water.

Primary observations of the solid sample should include noting its color.

1. Analysis for the anion. Carry out the anion analysis just as directed in Part II. Record your results. Your test results for the anions may differ somewhat due to the presence of a different cation in your salt than was present in the known anion samples. For example, if a transition metal ion is present in the unknown salt, it may impart color not seen in the knowns containing sodium salts.

2. Analysis for the cation. Dissolve 150 to 200 mg of solid sample (a volume equal to about two rice grains) in 10 mL of water. Use a 1 mL portion of this solution to analyze for the cation present using the procedure given in Part I. Record your result.

What salt is indicated by your anion analysis and cation analysis? Is the final result a reasonable one? (Some years ago, a student reported that a water soluble unknown salt had to be AgCl. In Part I, we observed that AgCl is insoluble in water, so the report was unreasonable!) All of the unknown salts given out are water soluble!


Add 2d HCl

Ppt: AgCl, PbCl2

Ppt: Al(OH)3, Fe(OH)3, Cr(OH)3, Pb(OH)2 Solution: Ba2+, Mg2+, Ni2+, Cu2+, Zn2+

Add 3d H2SO4

Ppt: BaSO4 Solution: Mg2+, Ni2+, Cu2+, Zn2+

Add 8 d NH3, 12d NaOH

Ppt: Mg(OH)2Solution:

Ni2+, Cu2+, Zn2+

Ag+, Pb2+, Al3+, Fe3+, Cr3+, Mg2+, Ba2+, Ni2+, Cu2+, Zn2+

Add 5d NH3

Solution: Pb2+, Al3+, Fe3+, Cr3+, Mg2+, Ba2+, Ni2+, Cu2+, Zn2+

Cation Flow ChartsThe information on this page is an overview of the

separation of the four groups of cations.Directions for separation of the ions in each group are given in the flow charts for the individual groups. These flow charts begin on the next page.

1. Separation of the Silver Group (Group 1)

2. Separation of the Aluminum Group (Group 2)

3. Separation of the Alkaline Earth Metals (Group 3)

4. The Copper Group (Group 4) remains in solution


Ag+, Pb2+, Al3+, Fe3+, Cr3+, Mg2+, Ba2+, Ni2+, Cu2+, Zn2+

Add 2d HCl, centrifuge

Ppt: 1A: AgCl (white), PbCl2 (white) Solution 2A: Pb2+, Al3+, Fe3+, Cr3+, Mg2+, Ba2+, Ni2+, Cu2+, Zn2+

Go to next flow chartWash once or twice with 1 mL of cold water, centrifuge, discard wash. Add 1 mL H2O, heat to near boiling, stir centrifuge & separate while hot.*

Ppt 1B: AgCl

Add 3d NH3, 1 mL H2O, stir, centrifuge. Save solution, discard ppt.

Soln 1D: Ag(NH3)2+

Add 5d HNO3

White ppt confirms Ag+

Solution 1C: Pb2+

Add 1d K2CrO4

Yellow ppt, PbCrO4 confirms Pb2+(see note 2 below)

Cation Analysis Procedure

Start with 1 mL of the known or unknown solution.

1. Separation and identification of the silver group.

_______________* If the solution cools before it is separated from the precipitate, PbCl2 will reprecipitate and contaminate the AgCl.** PbCrO4 will dissolve in a few drops of concentrated HCl.


Go to 3

Ppt 2B: Al(OH)3 (white),Fe(OH)3 (brown), Cr(OH)3 (blue-gray),

Pb(OH)2 (tan or white)

Add 3d H2SO4 & 1 mL H2O

Ppt. Tan or white PbSO4 or BaSO4. Discard. Soln 2C: Al3+, Fe3+, Cr3+

Add NH3 drop wise with cooling, as before (~2-4d). Centrifuge.

Soln 3A: Mg2+, Ba2+, Ni(NH3)62+, Cu(NH3)42+, Zn(NH3)42+

Solution. Add this to alkaline earth & copper groups, solution 3A above. BaSO4 (white) may ppt at this point. Do not be concerned.Ppt 2D: Al(OH)3, Cr(OH)3, Fe(OH)3

Wash once with 1 mL H2O & discard wash.Add 1 mL H2O, 5d NaOH, stir.Add 3d H2O2, stir, heat 2 min.

Ppt 2E: Fe(OH)3 (brown)

Dissolve in 3d HCl. Add 1d KSCN

Red solution of FeSCN2+ confirms Fe3+

Soln 2F: Al(OH)4- (colorless), CrO42- (yellow)

Add 1 mL NH4NO3, heat for 2 min.

Soln 2H: CrO42-

Add 1d BaCl2

Yellow ppt of BaCrO4 confirms Cr3+

Ppt 2G: Al(OH)3

Wash with 1 mL H2O, discard wash. Dissolve ppt in 2d HCl. Add 2d Aluminon reagent. Add 2d NH3.

Red (or pink) ppt Al(OH)3 confirms Al3+

Place a few ice cubes in 100 mL of water and cool Soln 2A while stirring. Add NH3 one drop at a time until an excess is present (~5-10 drops). Mix with a glass rod. Centrifuge and separate.

Solution 2A from 1: Pb2+, Al3+, Fe3+, Cr3+, Mg2+, Ba2+, Ni2+, Cu2+, Zn2+

2. Separation and identification of Aluminum Group.


Go to 4

Ba2+, Mg2+ ,Cu(NH3)42+ (deep blue), Ni(NH3)62+ (light violet), Zn(NH3)42+ (colorless)

Add H2SO4 until acidic (~1-2d)

Soln. 3C Mg2+, Cu2+, Ni2+, Zn2+Ppt. 3B white, BaSO4 confirms Ba2+

Add NH3 until an excess is present (~8d). Add 12d NaOH. Check the pH of the solution and continue adding NaOH drop by drop until the pH is 10 or greater.

Ppt. 3D Mg(OH)2 (white gelatinous)

Ppt. 3E Mg(NH4)PO4 (white crystalline)

A blue lake (flocculent ppt) confirms Mg2+

Soln. 4A Cu(NH3)42+ , Ni(NH3)62+ , Zn(NH3)42+

Wash twice with 1 mL H2O; discard wash. Add 1d HCl, 1d NH3, & 1 mL H2O. Solution should be almost neutral (~pH 8). Centrifuge and discard any ppt at this point. Add 3d Na2HPO4 soln. Wait 5 min.

If the ppt. is blue (due to the presence of Cu2+) add 1d HCl and 2d NH3. Centrifuge and discard supernatant liquid.Wash ppt. once with 1 mL H2O. Dissolve ppt. in 3d HCl, then add 3d of magneson reagent (magnesium reagent). Add NaOH with stirring until distinctly alkaline pH 10, centrifuge.

3. Separation of Alkaline Earth Metal Cations from Copper Group.

Begin with solution 3A from 2. Do not discard any white ppt. at this point.


Soln 4A from 3. Cu(NH3)42+ (deep blue), Ni(NH3)62+ (light violet), Zn(NH3)42+ (colorless)

Add HCl until the soln is acidic. Add 3d KI, stir. Brown color due to I3- develops. Wait 5 min for ppt to develop; A tan ppt of CuI will form if Cu2+ is present. Centrifuge. Test for completeness of pptn by adding 1d KI. Continue until no further ppt forms.

Ppt 4B: CuI (white or tan)

Wash once with 1 mL H2O, discard wash.Add 3d NH3, add 2d H2O2.

Deep blue solution, Cu(NH3)42+ confirms Cu2+

Soln 4C: Ni2+, Zn2+

If CuI pptd above, add Na2SO3 drop by drop, with stirring until brown color due to I3- disappears.Add 10d NaOH. Heat soln to boiling for at least 10 min. At this point, the pH of the solution should be 10 or greater. Repeat the NaOH heating treatment until the pH is 10 or greater.

Ppt 4D: Ni(OH)2 (green)

Add 1 mL H2O, 1d HCl, 2d NH3 and 1d dimethylglyoxime (DMG) reagent

Pink-red ppt of Ni(DMG)2 (nickel–dimethylglyoxime complex) confirms Ni2+

Soln 4E: Zn(OH)42-

Add 5d HCl, or until acidic. Add 2d K4Fe(CN)6.

Gray-green* ppt of Zn3K2[Fe(CN)6]2 confirms Zn2+

4. Separation and identification of Copper Group.

_______________* The ppt. formed from the known solution appears gray-green in color. Zn3K2[Fe(CN)6]2 appears white when pure. Other cations can form colored precipitates with this reagent. A pink tinge may be present due to Cu2+. In this case, wash the ppt. with 2d NH3 and 1 mL H2O to remove the trace of Cu2+


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