#7 The Chemistry and Qualitative Analysis of Cations: Group Separations and Separation of Group I Cations

PURPOSE OF THE EXPERIMENT

Deadop the Crest and ion Grag, unandle onstrate the

a procedure for qualitatively verifying the presence of Agt, Hg

2t, and

Pb2+ ions in an unknown solution.

BACKGROUND INFORMATION

A selected group of chemically important and commonly encountered cations appears in Table 1 on the next page. Although the chemistry of all

ment only involves five of these cations: Ag+

ment his cavilve is de oused hate the A oratory, prion than rei.

ions. The unique chemical properties of each cation makes it possible to separate and verify its presence, even in complex mixtures.

One of the simplest but most successful ways to separate one metal cation from another in a mixture is selective precipitation. Thus, one or more cations in solution can be precipitated from other cations in the same solution by the addition of various reagents followed by relatively simple manipulations.

When this approach is adopted, two questions immediately arise.

Which one of several possible reagents is best? How much of this reagent should be used? The answer to the first question requires a knowledge of the relative solubilities, hence, the solubility product constants (Ksp), of potential precipitated compounds. The answer to the second question depends on the extent to which the separation must be quantitative.

However, even if only a qualitative separation is desired, as in this

I. Separating Group I Cations from Group Il, Ill, and IV

Cations

Table 1 Nomenclature and formulas of 12 important cations

cation

common name

barium

bismuth(III)

calcium

cadmium(II)

chromium(III)

copper (ll)

iron (lI)

lead (II)

mercury()

nickel(II)

silver(I)

tin (IV)

chromic cupric ferric plumbous mercurous nickelous

formula

Ba?+

Bi+

Ca?+

cd+

Cr+

Cu2+

Fe3+

Pb2+

Hg22+

Ni +

Ag Sn*t

stannic

experiment, a calculation is required involving solubility product constants and concentrations of various species in solution.

Before carrying out this experiment, we must first find answers to these two questions. Then, we must design a set of procedures and manipulations that will allow the application of this knowledge in the laboratory.

The cations listed in Table 1 are commonly divided into four groups on the basis of their chemical behavior to the chloride ion (CI) and the sulfide ion (S2-). First we will discuss the separation of Group I, II, III, and IV cations as groups of cations. Then, we will consider the separation of individual Group I cations.

Of the several anions that could be added as precipitating reagents, only two, Cl and S-

, will be considered initially. Chloride ion is the anion of

the strong acid hydrochloric acid (HCI). It is therefore a weak Brønsted base and undergoes negligible hydrolysis in aqueous solution. The only equi-libria that need to be considered, then, are solubility equilibria of the type

Ag*(aq) + CI (aq) = AgC(s, white)

(Eq. 1)

and, under appropriate conditions, complex ion equilibria of the type

Agt (aq) + 2 Claq) = [AgCh,| (aq)

(Eq. 2)

Sulfide ion is the anion of the weak acid HS. It is a strong Brønsted base and undergoes extensive hydrolysis in aqueous solution, as shown in Equation 3.

s2 (aq) + HOH(I) = HS (aq) + OH (aq)

(Eq. 3)

In discussing the more complicated sulfide acid-base equilibria, it is helpful to consider first the source of Sion and then how the various equilibria involved can be manipulated to obtain the desired S2- ion concentration.

When thioacetamide (CHCSNH) is dissolved in water, the following

equilibrium is established:

CH-C-NH_(aq) + 2 HOH(I) -

CH3-C-0 (aq) + NHa* (aq) + H,S(g)

(Eq. 4)

The forward reaction in this equilibrium is endothermic. The application of heat causes the reaction to produce increased concentrations of products, in accordance with Le Châtelier's principle. A warm solution of thioaceta-mide serves as a convenient in situ source of hydrogen sulfide (HS). In aqueous solution, HS is involved in the following acid-base equilibria:

H2S(aq) + H2O(l) = H3O+(aq) + HS- (aq)

(Eg. 5)

Ka1 = [H3O^+][HS^-] / [H2S] = 1.0 x 10-7

(Eg. 6)

HS^- (aq) + H2O(l) -><- H3O^+ (aq) + S^2- (aq) (Eg. 7)

Ka2 = [H3O^+][S^2-] / [HS^-] = 1.3 x 10^-13(Eq. 8)

Application of Le Châtelier's principle allows us to predict that an increase in the hydronium ion (HOT) concentration in a solution containing HS will result in a decrease in the Sion concentration, while a decrease in the H3O+ ion concentration will result in an increase in the S2- ion equilibrium concentration. Thus, the S'- ion concentration in a solution saturated with HS can be controlled by adjusting the acidity of the solution. When Equation 5 and 7 are combined, the overall dissociation of HaS may be described with a single chemical equation and a single equilibrium expression. Addition of Equation 5 and 7 gives

HS(aq) + 2H2O(l) = 2H3O+(aq) + S^2-(aq)

(Eq. 9)

and multiplication of the correponing equilibrium expressions gives

([H3O^+][HS^-] / [H2S] ) ([H3O^+] [S^2-] / [HS^-] ) ([H3O^+]^2[S^2-] / [H2S] = ka1 ka2 = kH2S(Eq. 10)

In conventional notation schemes, the square brackets indicate concentration in moles per liter.

overall equilibrium constant.

By using the equilibrium constants given above, we can obtain the overall equilibrium constant.

KHS = Ka1 Ka2

= (1.0 × 10^-7) (1.3 x 10^-13)

(Eq. 11)

= 1.3 × 10^-20

The overall equilibrium expression in Equation 10 may be used to calculate the HO* or 52 ion concentration in a solution saturated with H2S when either concentration is known. Both H3O* and S- ion concentrations cannot be solved for simultaneously by using this relation because there would be one equation and two unknowns. To illustrate the use of this equilibrium expression, we can determine the Sº- ion concentration in a solution that is 1.0 × 10-2 M in H,S. By rearranging the overall equilibrium expression (Equation 10) as follows

[S^2-] = [H2S] (1.3x10^-20) / [H3O^+]^2 =

1.3 x 10^-22 /[H3O^+]^2

1.3 × 10-22

(Eq. 12)

[H3O+2

we can see that the equilibrium S- ion concentration is inversely proportional to the square of the HOt ion concentration. By controlling the HOT ion concentration, that is, by controlling the pH of the solution, we can control the S- ion concentration. This concentration, in turn, will determine which metal sulfide will precipitate from solution. Table 2 gives the 5°- ion concentration in equilibrium with 1.0 × 10-2M HaS at various HOT ion concentrations and pH values.

Now we will consider the problem of deciding which of these two reagents, CI or S2- ion, will more easily and efficiently effect a separation of the larger group of cations into smaller subgroups. Examination of the solubility product constants given in Table 3 is revealing. Of the 12 cations being considered, only three, Agt

+, Hgzt, and Pb2t ions, form reasonably

insoluble chlorides. These cations are commonly referred to as Group I cations or insoluble chlorides.

Therefore, a procedure of separation becomes evident. Hydrochloric acid serves as an ideal source of CI ion because H3O+ ion is not involved in any interfering equilibria with Agt ion, Hgz? ion, Pb2t ion, or any of the other cations remaining in solution. First, we add HCl to a solution containing the 12 specified cations. Agt ion, Hg22t ion, and/or Pb2+ ion will precipitate as insoluble silver chloride (AgCI), mercury(I chloride (Hg2Cl, mercurous chloride), and/or lead (II) chloride (PbCl, plumbous chloride). The precipitate can be separated from the solution containing the other cations by filtration or centrifugation. The mixture of insoluble chlorides can then be further separated and individually identified.

Table 2 Equilibrium concentrations of S-ion in 1.0×10-2M HS solutions of various hydrogen ion concentrations

solution

52- ion concentration, mol L-

1.0 × 10-2M H2S

4.5

1.3 × 10-13

1.0 × 10-2M HS + 1.0 × 10-3 M H3O+

1.3 × 10-16

1.0 × 10-2M H2S + 1.0 × 10-2M H30+

1.3 × 10-18

1.0 × 10-2M H,S + 0.10M H,O+

1.3 × 10-20

1.0 × 10-2M H_S + 1.0M H3O+

1.3 × 10-22

Table 3 Solubility product constants at 25°C

ion

chloride

sulfide

minimum [S'-] needed to precipitate 0.1M cations

Ag*

1.8 × 10-10

6.8 × 10-50

Hg2?+

1.3 × 10-18

5.8 × 10-44

Pb?+

1.6 × 10-5

8.4 × 10-28

Cu+

>10 (soluble)

8.7 × 10-36

8.7 x 10-35

Cd?+

>10 (soluble)

7.8 × 10-27

7.8 × 10-26

Bi3+

>10-1 (soluble)

6.8 × 10-97

4.1 × 10-32

Sn*+

>10-1 (soluble)

SnS2=1.0 x10

3.2 × 10-35

Fet

>10- (soluble)

FeS = 4.9 x 10-18

4.9 × 10-17

Fe3+

>10- (soluble)

Fe2S3 = 1 × 10-88*

2.2 x 10-29*

Ni?+

>10-1 (soluble)

1.8 × 10-21

1.8 × 10-20

Cr +

>10- (soluble)

not stable in H2O

Ba2+

>10-1 (soluble)

not stable in H20

Ca?+

>10-1 (soluble)

not stable in H2O

*FezS3 in acid decomposes to FeS and S.

The cations remaining in the filtrate can be separated and individually identified.

Although these Group I chlorides are very insoluble, their solubilities vary quite a bit. If we add enough chloride to precipitate virtually all of the most soluble of the three chlorides, PbCl, then we may be assured the other two more insoluble Group I chlorides are even more completely precipitated.

How much CI ion should be used for the precipitation? Consider the most soluble chloride of the three Group I cations, PbCl, as an example:

PC(s, white) = Pb+ (aq) + 2CI (aq)

[Pb'+] [CI)? = Ksp = 1.6 x 10-5

(Eq. 13)

(Eg. 14)

The left side of Equation 14 is called the ion product, while the right side is the Ksp. In a given solution, precipitation will begin to occur only when the ion product exceeds the solubility product constant. If the object is to precipitate the Pb't ions, a high CI ion concentration, according to Le Châtelier's principle, will force the reaction shown in Equation 13 to the left.

Thus, the Pb-+ ion will be virtually completely precipitated.

Addition of Cl ion will have no noticeable effect until enough has been added. Because of the Pb2t ion already present when enough CI has been added, the ion product exceeds the solubility product and precipitation occurs. It is important to note that prior to the formation of solid precipitate, addition of CI ion does not shift the equilibrium of Equation 13.

Indeed, there can be no equilibrium situation unless some PbCs) is present. After some solid PCh is present, further addition of Clion shifts the equilibrium to the left in what is known as the common-ion effect. This equilibrium shift reduces the solubility of PC and the concentration of Pb?+ ion remaining in solution. If Pb+ ion is initially present at a concentration of 0.10M, precipitation will occur when the Cl ion reaches a II. Separating Group I and I!!

Cations

concentration such that the ion product of Pb+ ion and CI- ion just exceeds the solubility product constant 1.6 × 10-5

• This precipitation occurs when

[Cl^-] = square root (1.6 x 10^-5) / 1x 10^-1 =1.3 x 10^-2 M

(Eq. 15)

If you want the precipitation to be 99.9% complete, then the Pb+ ion concentration remaining in solution will be 1.0 × 10-4M and the Cl ion concentration will have to be raised to

[Cl^-] = square root (1.6 x 10^-5) / 1x 10^-4 =0.4 M

(Eg. 16)

With AgCl and Hg2Cl, which are even less soluble, the precipitation of Ag* ion and Hg22+ ion as the chlorides will be even more than 99.9% complete.

In view of the common-ion effect, it may be tempting to add 12M HC instead of 0.40M HCl to assure complete precipitation of the Group 1 cations. Careful! Not only might the soluble cation chlorides begin to precipitate, but the following subtle aspect of chemistry will become important in such a case.

Silver ion reacts with excess CI ion to form a complex anion,

dichloroargentate(I, [AgCh, as shown in Equation 2.

Ag^+ (aq) + 2CI^- (aq) = [AgCl2]^- (aq)

(Eg. 2)

K = 1.0 × 10^5

If the CI ion concentration is raised too high, a considerable amount of the Agt ion will be complexed as soluble [AgC ion. This process is most conveniently considered as the redissolving of the AgCl precipitate.

AgCI(s) is in equilibrium with its ions:

AgCI(s, white) = Ag^+(aq) + CI^- (aq)

(Eq. 17)

Ksp = 1.8 x 10^-10

The formation of [AgCh ion does not occur appreciably at CI ion concentrations of less than 5.0M. We will use 6.0M HCI, which, through dilution and reaction, will yield a Cl ion concentration of 1-2M in the solution.

A further subdivision of the nine remaining cations in solution might be considered after CI ion addition precipitates the Group I cations. A look at Table 3 reveals a rather large variation in the solubility product constants of the sulfides. The solubility product constants for these sulfides are small, indicating they are all very insoluble. The constants do fall into two groups on the basis oftheir magnitude. Separation of these cations is possible if we control reaction conditions, for example, by varying the S- ion concentra-tion. The data in the last column, the S2- ion concentration necessary to precipitate 0.10M cation, suggest that a separation might be possible by adjusting the Sion concentration through control of the pH of the reaction.

We must first consider the situation involving Fe+ ion. In qualitative analysis schemes, iron is present as Fest ion instead of iron(II ion (Fet, ferrous ion) to avoid the oxidation reduction reaction with Agt ion as shown in Equation 18

Ag^+ (aq) + Fe^2+(aq) = Ag(s, gray-black) + Fe^3+(aq)

(Eq. 18)

After the precipitation of Ag^+ ion as AgCl, the addition of H2S in acidic solutions results in the reduction of Fe^3+ ion to Fe^2+ ion as shown in

Equation 19.

2 Fe^3+ (aq) +S^ 2- (aq) = 2 Fe^2+ (aq) + S(s)

(Eq. 19)

If we refer to Table 3 and neglect Fest ion, which is converted to Fe?t ion (See Equation 19), we can see that there is a large break in the S2- ion concentration required to precipitate cations whose concentrations are 0.10M. Cut, Cd +

+, Bit , and Sn't ions fall into one category and Fet and

Ni2+ ions fall into a second. The latter two cations require a much higher 5'-ion concentration for precipitation than the first four cations. The first four cations are from Group II, the acid-insoluble sulfides, and Fet and Ni?+ ions are placed in the Group III category. When the S2- ion concentration is controlled, the more insoluble metal sulfides (Group II cations) can be selectively precipitated while the Group III cations, the acid-soluble sulfides, remain in solution.

While the interesting details of this separation will be covered in other experiments, the general chemical principles involved will be illustrated in this experiment. To do so, we will use Cut ion from Group II and Fet ion from Group III. We will take advantage of the large difference in the solu-zbility product constants of their sulfides to effect a separation of the two ions.

carrying out this separation, we will use three important chemical reactions to identify the cations: (1) dithionite ion (S2042-) as a strong reducing agent,

(2) hexacyanoferrateII) ion ([Fe(CN),]*-, ferrocyanide ion) as a precipitating agent, and (3) thiocyanate ion (SCN) as a complexing agent. Dithionite ion is a strong reducing agent that reduces Agt ion to elemental silver, Ag(s), and sulfite ion (SO;) is formed as the oxidation product. Equation 20 shows the reduction of silver(I) in the ammonia complex ion diammine-silver (I), [Ag(NH)2It, to Ag(s).

2[Ag(NH3)2]^+ (aq) + S2O4^2-(aq) + 2H2O(l) = 2 Ag(s, gray-black) + 2SO3^2-(aq) + 4NH4^+ (aq) (Eq. 20)

Elemental silver formed in this way appears gray-black in color.

The copper! sulfide (CuS, cupric sulfide) precipitate is dissolved in

nitric acid (HNO), forming Cut ion as shown in Equation 21.

3 CuS(s) + 2HNO3(aq) + 6H3O+(aq) = 3 Cu^2+ (aq) + 3S(s) + 2 NO (aq) + 10H2O(l) (Eq. 21)

Then, Cut ion forms a dark blue complex ion, tetramminecopper(II, [Cu(NH3)4}, with ammonia (NHa) in basic solution as shown in Equation 22.

Cu^2+ (aq) + 4 NH3 (aq) = [Cu(NH3)4]^2+(aq, dark blue)

(Eq. 22)

IlI. Group I Cations (Insoluble

Chlorides)

Adding acid shifts the equilibrium shown in Equation 22 by converting

NH3 to ammonium ion (NH,). The presence of Cut ion is detected by the addition of [Fe(CN).l* ion, forming a reddish purple precipitate of copper hexacyanoferrate(I, [Cu,Fe(CN)6l, as shown in Equation 23.

2 Cư (aq) + [Fe(CN),/* (aq) = CuzFe(CN),(s, reddish purple) (Eq. 23)

Iron(II sulfide (FeS, ferrous sulfide) is also dissolved in HNO, during which process Fet ion is oxidized to Fest ion.

Thiocyanate ion (SCN) forms a stable complex with Fet ion, as shown

in Equation 24.

Feet (aq) + SCN (aq) = [FeSCN)2+ (aq, deep red)

(Eq. 24)

K = 1.0 × 102

The thiocyanatoiron(III complex ion ([FeSCN)+, ferrithiocyanate ion)

imparts a deep red color to an aqueous solution. The formation of this complex upon addition of SCN ion is a sensitive test for the presence of Fe3+ ion.

The problem remains of how to confirm definitely the presence of Agt, Pb-+ and Hg22t ions, the Group I cations, in the chloride precipitate.

A. Separating and Confirming Lead II)

Figure 1 shows the temperature dependence of the solubilities of the three chlorides. These curves suggest a procedure for further separation. In the

140

temperature range 20-100°C, the solubilities of each of these chlorides increases, but to varying degrees. The approximate threefold increase in the solubility of PCl that occurs between 20 to 100 °C is such that at 100 °C appreciable amounts of PbCl2 will dissolve, while AgCl and Hg2C12 remain insoluble. Therefore, if the combined residues of the three chlorides are treated with hot water, the PC should dissolve preferentially.

After any dissolved Pb? ion is separated by filtering the hot solution, the presence of Pb'+ ion may be confirmed by cooling the filtrate, resulting in the reprecipitation of any Pb't ion present as PCI. If only a turbid solution results, the addition of HCI will precipitate more PCl if Pb+ ion is present.

B. Separating Silver(I and Mercury and Confirming Mercury(T)

When separation of the remaining residue is attempted, a bit of chemical luck is encountered. A chemical reagent exists that preferentially acts to dissolve AgCl and at the same time confirms the presence of Hg2Clz. This reagent is aqueous NH. Silver ion forms a complex cation with NH

Ag^+ (aq) + 2 NH3(aq) = [Ag(NH3)2]^+(aq)

(Eg. 25)

K = 1.5 × 10^7

By combining Equation 17 and 25, the dissolution of solid AgCl to form

[Ag(NH3)2]* ion may be described with a single chemical equation

AgCI(s, white) +2 NH3(aq) = [Ag(NH3)2]^+(aq) + CI^-(aq)

Eq. 26)

for which the equilibrium constant is

K = (1.5 × 10^7)(1.8 × 10^-10) = 2.7 × 10^-3

(Eq. 27)

This equilibrium constant is a reasonably large one. Applying Le Châtelier's principle, we can shift the equilibrium to the right by adding excess aqueous NH3, and virtually all of the AgCl in the precipitate can be dissolved.

On the other hand, Hg2Cl2 reacts quite differently with aqueous NH3.

The following is an overall stoichiometric equation describing this reaction:

Hg2Cl2 (s, white) + 2NH3(aq) = HgNH2CI(s, white) + NH4^+(aq) + CI^- (aq) + Hg(l, black) (Eq. 28)

Greater insight can be gained as to what is involved here if this equation is rewritten in a sequence of three equilibria, as shown in Equation 29, 30, and 31.

Hg2Cl2(s, white) = Hg2^2+ (aq) + 2Cl^- (aq)

(Eq. 29)

Hg2^2+(aq) = Hg(l, black) + Hg^2+ (aq)

(Eq. 30)

Hg^2+ (aq) + 2NH3 (aq) + CI^- (aq) = HgNH2Cl(s, white) + NH4^+ (aq) (Eq. 31)

neg i Equation 30, the 19 ac on free prese read to shown. in

Equation 29 undergoes an autoredox or disproportionation reaction, in which Hg.?+ ion is both oxidized to mercury ion (Hg, mercuric ion)

and reduced to elemental mercury, Hg(s). The addition of aqueous NH3 in the presence of CI ion converts the Hg?+ ion formed in Equation 29 to insoluble amidochloromercury(II (HgNH,CI), as shown in Equation 31.

Hence, the addition of aqueous NH, to solid Hg,Cl2 will shift the equilibrium in Equation 29, 30, and 31 to the right, resulting in the conversion of Hg?C/½ to Hg(s) and HgNH,C1. The very low solubility of HgNH, Cl is the driving force behind this sequence of reactions. Elemental mercury appears black when finely divided, and HgNH,Cl is white. The total precipitate appears to have a black or gray color, thus confirming the presence of Hgz ion in the original solution.

C. Confirming Silver (I)

Silver ion is now present as [Ag(NH3)2lt ion in solution with CT ion, the latter formed from the dissolution of AgCl, as shown in Equation 26.

In aqueous solution, NH3 is a Brønsted base.

NH3 (ag) + H20(l) = NH4^+ (aq) + OH^- (aq)

(Eq. 32)

Kb = 1.8 x 10^-5

Equations can be written for the simultaneous equilibria relevant to the

Ag ion confirmation in the following sequence of two reactions:

[Ag(NH3)2]^+ (aq) = Ag^+(aq) + 2NH3 (aq)

(Eq. 33)

K = 6.7 × 10^-8

OH^- (aq) + H3O^+ (aq) = 2H2O(l)

(Eq. 34)

K = 1.0 × 10^14

As a result of the equilibrium in Equation 33, some NH3 is present in solution. The reaction of NHz formed in Equation 33 with water forms NH4 and OH ions, as shown in Equation 32. The OHion formed in Equation 32 will react with HOt ion to form water, as shown in Equation 34. Thus, the addition of acid to a solution containing [Ag(NH3)2]* ion will shift all three equilibria to the right, resulting in the release of Ag ion into the solution.

With the CI ion already present, if the volume of the solution has not been increased too much, the ion product of Ag ion and Clion will exceed the solubility product constant of AgCl. White AgCl will precipitate; the formation of this precipitate confirms the presence of Agt ion in the original solution.

In this experiment, you will separate Agt, Cut, and Fest ions from a known solution and verify the presence of each cation by a confirmation test. Then, you will separate Agt, Hgz+ , and Pb'+ ions (Group I cations)

from a second known solution and verify the presence of each cation by a confirmation test. Finally, you will determine the cations present in an unknown solution containing one or more of the Group I cations, using the same procedure as you did with the known cation solution.

PROCEDURE

Some Notes on Semimicro

Technique

1. Precipitations: The precipitating reagent should be added dropwise with stirring until precipitation is complete.2. Completeness of Precipitation: So that interfering cations are not left in solution, it is often necessary, where noted, to allow the precipitate to settle in the test tube, or to centrifuge the supernatant liquid and the precipitate, and then carefully add 1 additional drop of reagent. If more precipitate forms, a few more drops of reagent should be added and the above procedure repeated until no further precipitation is observed.

3. Washing Precipitates: To ensure removal of interfering ions from moist precipitates, it is necessary, where noted, to wash the precipitate.

This procedure involves decanting the supernatant liquid, adding the required amount of specified wash liquid to the tube containing the moist precipitate, mixing the precipitate and wash liquid thoroughly with a clean glass stirring rod, centrifuging, and decanting and discarding the wash liquid.

CAUTION

A solution in a small test tube cannot be heated safely over a direct flame.

4. Heating a Solution: Placing the test tube containing the solution to be heated in a hot or boiling water bath is generally a preferred method.

5. Removing Supernatant Liquids: Liquids over precipitates are most often removed by careful decantation. However, the supernatant liquid can carefully be withdrawn with a small medicine dropper or pipet, if necessary.

Occasionally a precipitate does not centrifuge completely. To fully separate a supernatant liquid in such a situation, twist a small piece of cotton batting to a point and insert it partway into the tip of a clear eyedropper, leaving a tuft extending outside the tip. Draw up the supernatant liquid through the cotton, which will filter out the floating particles of precipitate. Enough precipitate should remain in the test tube for testing of those cations. Carefully remove the cotton and release the clear supernatant liquid into a clean 10 x 75-mm test tube.

CHEMICAL ALERT

6M acetic acid-toxic and corrosive

6M ammonia-toxic, corrosive, and irritant

0.1 M copper(il) nitrate in 0.1M HNO—toxic, irritant, and oxidant

6M hydrochloric acid-toxic and corrosive

0.1M iron(lI) nitrate in 0.1M HNO—toxic, corrosive, and oxidant

0.1M lead(l) nitrate in 0.1M HNOtoxic, irritant, and oxidant

0.1M mercury) nitrate in 0.1 MHNOhighly toxic and oxidant

6M nitric acid-toxic, corrosive, and strong oxidant

0.1M potassium ferrocyanide-irritant

0.1M potassium thiocyanate-irritant

0.1M silver nitrate in 0.1M HNO—toxic, corrosive, and oxidant sodium dithionite-toxic and irritant

1M thioacetamide toxic and carcinogen 2. Completeness of Precipitation: So that interfering cations are not left in solution, it is often necessary, where noted, to allow the precipitate to settle in the test tube, or to centrifuge the supernatant liquid and the precipitate, and then carefully add 1 additional drop of reagent. If more precipitate forms, a few more drops of reagent should be added and the above procedure repeated until no further precipitation is observed.

3. Washing Precipitates: To ensure removal of interfering ions from moist precipitates, it is necessary, where noted, to wash the precipitate.

CAUTION

A solution in a small test tube cannot be heated safely over a direct flame.

4. Heating a Solution: Placing the test tube containing the solution to be heated in a hot or boiling water bath is generally a preferred method.

CHEMICAL ALERT

6M acetic acid-toxic and corrosive

6M ammonia-toxic, corrosive, and irritant

0.1 M copper(il) nitrate in 0.1M HNO—toxic, irritant, and oxidant

6M hydrochloric acid-toxic and corrosive

0.1M iron(lI) nitrate in 0.1M HNO—toxic, corrosive, and oxidant

0.1M lead(l) nitrate in 0.1M HNOtoxic, irritant, and oxidant

0.1M mercury) nitrate in 0.1 MHNOhighly toxic and oxidant

6M nitric acid-toxic, corrosive, and strong oxidant

0.1M potassium ferrocyanide-irritant

0.1M potassium thiocyanate-irritant

0.1M silver nitrate in 0.1M HNO—toxic, corrosive, and oxidant sodium dithionite-toxic and irritant

1M thioacetamide toxic and carcinogen

Wear departmentally approved eye protection while doing this experiment.

1. Separating Selected Group

I Cations from Group ll and I!!

Cations

A. Precipitating Group I Cations (Insoluble Chlorides)

In a clean appropriate container, obtain from your laboratory instructor a known solution that is approximately 0.1M in each of the three cations, Ag^+, Cu^2+, and Fe^3+ ions. With a clean dropper, transfer 15 drops of this mixture

to a 75-mm test tube.

CAUTION

6M hydrochloric acid is a corrosive, toxic solution that can cause burns.

Prevent contact with your eyes, skin, and clothing. Avoid inhaling vapors and ingesting the solution.

From a second clean dropper, add dropwise 6M HCI to the mixture in the test tube, while stirring, until the precipitation is complete.

Record all observations on Data Sheet 1.

Place the tube containing the mixture in the centrifuge and balance it with another tube containing an amount of water equal to the volume of the mixture in the tube. Centrifuge for 2 min. Decant the supernatant liquid, containing Group II and III cations, from the precipitate into a clean, dry test tube. Label this tube and retain it for future use.

Wash the precipitate in the test tube with 2 mL of distilled or deionized water. Thoroughly stir the contents of the tube using a clean glass stirring rod. Centrifuge the mixture in the tube for 2 min.

NOTE: In all parts of this experiment, follow the directions of your laboratory instructor for discarding reaction mixtures and unused reagents.

Decant the supernatant liquid from the precipitate and discard the super-natant liquid.

CAUTION

6M ammonia is a corrosive, toxic solution. Prevent contact with your eyes, skin, and clothing. Avoid inhaling vapors and ingesting the solution.

Add 15 drops of 6M NH3 to dissolve the precipitate. Add a small amount, about the size of a grain of rice, of solid sodium dithionite to the solution in the tube. Record all observations on Data Sheet 1.

B. Precipitating Group II Cations as Sulfides

10.3 pH unit of a pH of 0.5, adjust the pH of the solution to approximately

0.5 by adding dropwise from a clean dropper either 0.5M HCl or 0.5M NH3, as appropriate. Prepare the 0.5M solutions by diluting 1 drop of a 6M solution with 11 drops of distilled water and thoroughly mixing with a clean glass stirring rod. This pli adjustment is important. If the pH is too low, that is, if the hydrogen ion concentration is too high, the concentration of S2- ion will be insufficient to permit the precipitation of CuS. If the pH is too high, CuS and FeS will both precipitate.

CAUTION

Thioacetamide has been demonstrated to be a carcinogen in animal feeding studies. It is safe to use when handled prudently. Avoid skin contact with the thioacetamide solution. Wash your hands thoroughly with soap or detergent after using the solution.

When you have properly adjusted the pH of the supernatant liquid, use a clean dropper to add 10 drops of 1M thioacetamide solution to the solution in the test tube. Heat the tube and its contents in a boiling water bath for 5 min.

Record all observations on Data Sheet 1.

Centrifuge the tube and its contents for 2 min. Decant the supernatant liquid into another clean 75-mm test tube. Label these tubes. Retain the tube containing the CuS precipitate and the tube containing the supernatant liquid.

Test the supernatant liquid for completeness of precipitation by adding 5 drops of 1M thioacetamide to the solution. Heat the tube and its contents in a boiling water bath for 5 min. If additional dark-colored precipitate forms, the first CuS precipitation was not complete. A light-colored precipitate may form, which is elemental sulfur formed from the decomposition of thioacetamide.

Centrifuge the tube containing the supernatant liquid. Decant the supernatant liquid into a clean 75-mm test tube and discard any precipitate that may have formed. Label the tube and retain the supernatant liquid for Group II! tests.

CAUTION

6M nitric acid is a corrosive, toxic solution that can cause severe burns.

Prevent contact with your eyes, skin, and clothing, and combustible material.

Avoid inhaling vapors and ingesting the solution.

Add 20 drops of 6M HNO to the tube containing the original CuS precipitate. Carefully heat the tube and its contents in a boiling water bath while carefully stirring the mixture with a clean glass stirring rod to suspend the CuS precipitate in the 6M HNO. Continue heating for 3 min.

Centrifuge the warm tube and its contents, if necessary, to facilitate the removal of any undissolved material. Decant the supernatant liquid into a

75-mm test tube and discard any precipitate. Let the tube and its contents cool. Add 6M NH dropwise until the solution is just basic to litmus paper.

Then add 5 more drops of 6M NH3 in excess.II. Precipitating Group I Cations, Silver),

Mercury(I), and Lead(Il)

CAUTION

6M acetic acid is a corrosive solution that can cause burns. Prevent contact with your eyes, skin, and clothing. Avoid inhaling the vapors and ingesting the solution.

Add 6M acetic acid (HOAc) dropwise until the solution is just acidic to litmus paper. Add 5 drops of 0.1M potassium hexacyanoferrate(I),

[K,Fe(CN)6], solution. Stir the solution with a clean glass stirring rod.

Record all observations on Data Sheet 1.

C. Precipitating Group III Cations as Sulfides from Alkaline HS Solution Select the labeled test tube containing the supernatant liquid from the Group II precipitation. Add 6M NH dropwise with stirring until the solution is just basic to litmus paper. Then add 7 drops of 6M NH3 in excess.

The solution should now be strongly basic. If there is residual HaS in solution and Fe't is present, a black precipitate of FeS will form as the solution becomes more basic.

To complete the formation of the FeS precipitate, add 10 drops of 1M thioacetamide to the basic solution. Thoroughly stir the mixture with a clean glass stirring rod. Warm the tube and its contents in a boiling water bath for 5 min.

Centrifuge the tube and its contents for 2 min. Decant the supernatant

liquid from the precipitate and discard the supernatant liquid.

Using a clean dropper, add 15 drops of 6M HNO to the precipitate.

Thoroughly stir the mixture with a clean glass stirring rod. Heat the tube and its contents in a boiling water bath for 5 min. If the solution is cloudy, centrifuge the tube and its contents for 2 min. Using a clean dropper, add 3 drops of 0.1M potassium thiocyanate (KSCN) to the clear solution in the tube.

Record all observations on Data Sheet 1.

A flowchart is given in Fish and or the separation and identification of the Group I cations Agt,

Obtain a known mixture containing the Group I cations from your laboratory instructor. Using a clean dropper, transfer 15 drops of the known solution to a clean 75-mm test tube. Using another clean dropper, add sufficient 6M HCI dropwise with shaking to effect complete precipitation.

Centrifuge the tube and its contents for 2 min. Add 1 drop of 6M HCI to the mixture. If additional precipitate forms, centrifuge the tube and its contents again for 2 min. Retest the mixture for completeness of precipitation. When precipitation is complete, decant the supernatant liquid from the precipitate.

If only Group I cations were present in the original sample solution, discard the supernatant liquid. If the sample solution is to be analyzed for Group II, III, and/or IV cations, retain and label the supernatant liquid.

Retain the precipitate in the tube for use in the next section.

Record all observations on Data Sheet 2.

A. Separating and Confirming Lead(II

Add 2 mL of distilled water to the Group I precipitate in the test tube. While stirring with a clean glass stirring rod, heat the tube and its contents in a

boiling water bath for 3 min. While heating the Group I precipitate, heat 4 mL of distilled water in a separate test tube in the boiling water bath. Heat another 10 mL ofdistilled water for later use. Pour 4 mL of hot distilled water through the filter paper immediately prior to filtering the precipitate.

Discard this water. By heating the filter paper and the funnel, any dissolved Pb+ ion will probably not reprecipitate during filtration. Thoroughly stir the Group I precipitate in the hot water. Rapidly filter the hot solution through the warmed funnel and filter paper. Retain the filtrate in a clean test tube.

Place the funnel in a large test tube. Wash the residue with two 5-mL portions of hot distilled water. Discard the washings. Leave the residue on the filter paper in the funnel and retain the residue for further analysis in the next section of this experiment.

Cool the filtrate in the test tube by immersing the tube in cold water.

If there is any question as to the results, add 2-3 drops of 6M HCI to the cooled filtrate. Thoroughly stir the mixture in the tube with a clean glass stirring rod.

Record all observations on Data Sheet 2.B. Separating and Confirming Mercury (I*

Place a small test tube under the funnel containing the residue from the previous filtration. Using a clean dropper, add 2 mL of 6M NH3 dropwise to the residue on the filter paper. Retain the filtrate for confirmation of the presence of Agt ion.

Record all observations on Data Sheet 2.

C. Confirming Silver (I)

Add 6M HNO, from a clean dropper to the test tube containing the ammonia filtrate, with stirring, until the solution is acidic when tested with litmus paper.

Record all observations on Data Sheet 2.

Obtain an unknown solution containing one or more of the Group I cations from your laboratory instructor. Record the unknown identification number on Data Sheet 3.

Repeat the procedure in Part II with this solution. List on Data Sheet 3 the cations you find in your unknown solution. Cite evidence on Data Sheet 3 proving the presence of the cations you report.

CAUTION

laboratory.

Wash your hands thoroughly with soap or detergent before leaving the