usp 191 identification tests-general_chemical equations_by_jude daval-santos

DAVAL-SANTOS, Jude A. Ph Ch 121 USP ID Tests for Common Ions (arranged according to their classification in Qualitative Analysis) of 17

Separation of the Metals into Analytical Groups (based on Holtzclaw & Robinson scheme and Svehla scheme)

Group 1 Cations (Group I

Chlorides)

Description of the Tests

(from USP 30 – NF 25) Corresponding Equations

Perceptible Result(s)

Principle Involved/ Notes

Lead (II)

(1) With 2 N sulfuric acid, solutions of

lead yield a white precipitate that is Pb(NO3)2 + H2SO4 PbSO4 + 2 HNO3 (Svehla, 1996, p.64)

White precipitate (Svehla, 1996, p.64)

Double displacement precipitation reaction (Holtzclaw & Robinson, 1988, p.945).

a. insoluble in 3 N hydrochloric

acid or 2 N nitric acid, but is

PbSO4 + HCl no reaction (Soine & Wilson, 1967, p.515) No dissolution of precipitate;

white precipitate persists (Soine

& Wilson, 1967, p.515)

The sulfate ion is converted largely to HSO4- by H3O

+ in high

concentration (Holtzclaw & Robinson, 1988, p.962), specifically Pb(HSO4)2 (Gilreath, 1954, p.194, 199). The concentrations may

not be weakly acidic not enough to provide enough H3O+ to form

the soluble bisulfate. PbSO4 + HNO3 no reaction (Soine & Wilson, 1967, p.515)

b. soluble in warm 1 N sodium hydroxide and

PbSO4 + 2 NaOH Na2SO4 + Pb(OH)2 (or H2PbO2)

Pb(OH)2 + 2 NaOH Na2PbO2 + 2 H2O

(Soine & Wilson, 1967, p.515)

Dissolution of precipitate forming

a clear, colorless Solution (Svehla, 1996, p.64)

Pb(II) is essentially basic, although with alkalies Pb(OH)2 will form plumbites. That is to say it will act like an acid (Parks et al, 1949, p.209). The fact that its hydroxide is amphoteric is an indication of its nonmetallic character. The solubility of lead chloride increases with temperature (Holtzclaw & Robinson, 1988, pp.833, 941).

PbSO4 + 4 NaOH Na2SO4 + Na2[Pb(OH)4] (Svehla, 1996, p.64)

Precipitate dissolves though complex ion formation of tetrahydroxoplumbate(II) (Holtzclaw & Robinson, 1988, p.943). The higher valence state (+4) is primarily acidic but also amphoteric (Parks et al, 1949, p.941).

c. in ammonium acetate TS.

PbSO4 + 4 CH3COONH4 (NH4)2[Pb(CH3COO)4] + (NH4)2SO4 (Svehla, 1996, p.65)

Dissolution of precipitate forming a clear, colorless solution (Svehla, 1996, p.65)

Precipitate dissolves though complex ion formation of tetraacetoplumbate(II).

PbSO4 + 2 CH3COONH4 Pb(CH3COO)2 + (NH4)2SO4 (Holtzclaw & Robinson, 1988, p.932).

The formation of a weak electrolyte lead acetate (Holtzclaw & Robinson, 1988, p.943).

(2) With potassium chromate TS,

solutions of lead salts, free or nearly

free from mineral acids, yield a yellow precipitate that is

Pb(NO3)2 + K2CrO4 PbCrO4 + 2 KNO3 (Svehla, 1996, p.65)

Bright yellow precipitate (Soine & Wilson, 1967, p.515)

Double displacement precipitation reaction. The solution be free from mineral acids because in acidic media, PbCrO4 will not be formed but the dichromate ion, and the latter yield orange-red aqueous solutions;

2 CrO42- + 2 H

+ Cr2O7

2- + H2O (Svehla, 1996, p.225).

a. insoluble in 6 N acetic acid but it PbCrO4 + CH3COOH no reaction


No dissolution of precipitate; bright yellow precipitate persists


CH3COOH is a weak acid thus it could not provide enough H3O+ to

form the dichromate, which could latter form an orange-red

aqueous solution (Svehla, 1996, p.225).

b. soluble in 1 N sodium hydroxide.

PbCrO4 + 2 NaOH Na2CrO4 + Pb(OH)2 (or H2PbO2)

Pb(OH)2 + 2 NaOH Na2PbO2 + 2 H2O (Soine & Wilson, 1967, p.515) ------------------------------------------------------------

PbCrO4 + 4 NaOH Na2CrO4 + Na2PbO2 + 2 H2O (Soine & Wilson, 1967, p.515)

Dissolution of precipitate forming a yellow solution (Lee, 1996, p.728)

Pb(II) is essentially basic, although with alkalies Pb(OH)2 will form

plumbites. That is to say it will act like an acid (Parks et al, 1949, p.209).

PbCrO4 + 4 NaOH Na2[Pb(OH)4] + Na2CrO4 (Svehla, 1996, p.65)

Formation of the soluble complex, which suppresses the Pb2+

ion concentration to such an extent that the solubility product of PbCrO4 is no longer exceeded (Svehla, 1996, p.226)


Mercury

General tests: (1) When applied to bright copper foil, solutions of mercury

salts, free from an excess of nitric acid, yield a deposit that upon rubbing,

becomes bright and silvery in appearance.

Hg (I): Cu + Hg2(NO3)2 Cu(NO3)2 + 2 Hg Deposit of mercury on a copper foil and then becomes bright and silvery upon rubbing (USP CI,

2006, p.140)

Single replacement reaction since copper is higher in the electromotive series than Hg (Soine & Wilson, 1967, 672). The solution of Hg should be free from an excess of HNO3 since the Hg2(NO3)2 tend to hydrolyze in aqueous solutions to form basic salts thus stabilizing it but preventing it from reacting Hg2(NO3)2 +

H2O Hg(OH)2NO3 + HNO3 (Soine &Wilson, 1967, p.424). Hg (II): Cu + HgCl2 CuCl2 + Hg (Soine & Wilson, 1967, pp.426-7)

(2) With hydrogen sulfide, solutions of

mercury compounds yield a black precipitate that is

Hg (I): Hg2(NO3)2 + H2S HgS + Hg + 2 HNO3

(Soine & Wilson, 1967, pp.424)

Black precipitate (Soine & Wilson, 1967, pp.424)

Disproportionation reaction (Holtzclaw & Robinson, 1988, p.931).

Hg (II): 3 HgCl2 + 2 H2S HgCl2 2 HgS + 4 HCl

HgCl2 2 HgS + H2S 3 HgS + 2 HCl (Soine & Wilson, 1967, pp.426)

White precipitate, which turns yellow to brown and finally to black (Soine & Wilson, 1967, pp.426)

White in the presence of HCl : HgCl2 2 HgS (a complex of the original mercuric salt with mercuric sulfide); yellow (complex having

greater proportions of HgS) (Soine & Wilson, 1967, pp.426). When heated, HgS becomes bright red; the red sulfide is isomeric with the black and has in the past has been used as the pigment vermilion (Holtzclaw & Robinson, 1988, p.845).

a. insoluble in ammonium sulfide TS and

HgS + (NH4)2S no reaction (Svehla, 1996, p.78)

Black precipitate (Svehla, 1996, p.78)

Colorless (NH4)2S; HgS is one of the least soluble precipitates known but soluble in solutions of soluble sulfides in the presence of an excess of hydroxide ions and forms the thiomercurate ion

[HgS2]2-

(Holtzclaw & Robinson, 1988, p.845).

b. in boiling 2 N nitric acid. HgS + HNO3 Hg(NO3)2 + H2S

Dissolution of precipitate forming a solution with rotten egg odor (Holtzclaw & Robinson, 1988, p.667)

Double displacement reaction.

Mercurous

(1) Mercurous compounds are decomposed by 1 N sodium hydroxide, producing a black color.

Hg2(NO3)2 + 2 NaOH Hg2(OH)2 + 2 NaNO3 (Soine & Wilson, 1967, p.425)

Hg2(OH)2 Hg2O + H2O

Hg2(NO3)2 + 2 NaOH Hg2O + H2O + 2 NaNO3

Black-brown precipitate (Soine & Wilson, 1967, p.425)

Double displacement reaction. Hg2(OH)2 is very unstable toward heat and light, thus it decomposes (Soine & Wilson, 1967, p.425). Stable dinuclear metal ion (Hg-Hg)

2+, bonded using 6s orbitals. The

only other metals which form dinuclear ions are (Zn-Zn)2+

and (Cd-Cd)

2+ , which are unstable and have only been detected

spectroscopically (Lee, 1996, p.846).

(2) With hydrochloric acid, solutions of

mercurous salts yield a white precipitate

Hg2(NO3)2 + 2 HCl 2 Hg2Cl2 + 2 HNO3 (Svehla,

1996, p.98)

White, crystalline precipitate (Holtzclaw & Robinson, 1988,

p.842)

Double displacement reaction. “Calomel” : Hg2Cl2 (Holtzclaw & Robinson, 1988, p.842).

(hv)

Hg2Cl2 HgCl2 + Hg (Holtzclaw & Robinson, 1988, p.842)

Grayish to black precipitate Decomposition. “Corrosive sublimate”: HgCl2 (Holtzclaw & Robinson, 1988, p.842).

that is blackened by 6 N ammonium hydroxide.

Hg2Cl2 + 2 NH3 Hg + Hg(NH2)Cl + NH4Cl (Svehla, 1996, p.68)

Black precipitate distributed

among white precipitate, which looks as shiny black (Svehla, 1996, p.68)

Disproportionation. This reaction can be used to differentiate mercury(I) ions from lead(II) and silver(I). Hg(NH2)Cl is a white

precipitate but the finely divided Hg makes it shiny black. The name calomel (Greek = nice black) refers to this characteristic of the originally Hg2Cl2 (Svehla, 1996, p.68). Ammonia water serves as a medium for the auto-redox of Hg2Cl2 (Gilreath, 1954, p.186).

(3) With potassium iodide TS, a yellow

precipitate, that may become green upon standing, is formed.

Hg2(NO3)2 + 2 KI Hg2I2 + 2 KNO3 (Soine & Wilson, 1967, p.425)

Yellow precipitate which turns green upon standing (Soine & Wilson, 1967, p.425)

Double displacement reaction. ! When KI is in excess of the Hg2(NO3)2, a black precipitate of metallic Hg is obtained, thus KI solution is always added to the Hg2(NO3)2 with constant agitation (Parks et al, 1949, p.57). When boiling it with H2O, disproportionation takes place, and a mixture of HgI & finely distributed black is formed (Svehla, 1996, p.70).


Silver (I)

(1) With hydrochloric acid, solutions of

silver salts yield a white, curdy precipitate that is

AgNO3 + HCl AgCl + HNO3 (Svehla, 1996, p.72)

White, curdy precipitate (Svehla, 1996, p.72)

Double displacement precipitation reaction. AgCl is soluble in dilute NH4OH, AgBr dissolves in 0.880 ammonia, and AgI is insoluble even in 0.880 ammonia (Lee, 1996, p.825).

(hv)

2 AgCl 2 Ag + Cl2

(Svehla, 1996, p.73)

Grayish or black precipitate

(Svehla, 1996, p.73)

Sunlight or uv irradiation decomposes silver halides, which turn to grayish or balck owing to the formation of silver metal. The reaction is slow though (Svehla ,1996, p.73).

a. insoluble in nitric acid, but is AgCl + HNO3 no reaction (Svehla, 1996, p.72)

No dissolution of precipitate; white, curdy precipitate persists (Svehla, 1996, p.72)

b. readily soluble in 6N ammonium

hydroxide. AgCl + NH4OH [Ag(NH3)2]Cl + H2O (Svehla, 1996, p.191)


Formation of the complex silver ammonia cation. Silver chloride ionizes to a greater extent than does silver ammonia complex cation (in an excess of NH4OH) and, therefore, silver ions are used up in the formation of complex cation (Soine & Wilson, 1967, p.329). Caution! See Bromide.

(2) A solution of silver salt to which 6 N

ammonium hydroxide and a small quantity of formaldehyde TS are added

deposits, upon warming, a mirror of metallic silver upon the sides of the container. “Tollen’s Test”; “Silver mirror test”

AgNO3 + NaOH AgOH + NaNO3 ;

2 AgOH Ag2O + H2O ------------------------------------------------------------

2 AgNO3 + 2 NaOH Ag2O + H2O + 2 NaNO3


Brown precipitate (Soine & Wilson, 1967, p.329)

Double displacement reaction between AgNO3 + NaOH while AgOH goes decomposition. †Silver fulminate Ag-C2N2O2 (Most, 1988, p.519) VS Ag3N from Svehla.

Ag2O + 4 NH4OH 2 Ag(NH3)2OH + 3 H2O (Soine & Wilson, 1967, p.329)

Dissolution of precipitate forming a solution (Soine & Wilson, 1967, p.329)

Precipitate dissolves though complex ion formation (Holtzclaw & Robinson, 1988, p.943).

HCHO + 2 Ag(NH3)2OH HCOONH4 + 2 Ag + 3

NH3 + H2O (Soine & Wilson, 1967, p.329)

Silver mirror on the walls of the test tube with ammonia odor (Soine & Wilson, 1967, p.329)

Redox reaction. Tollen’s reagent is an oxidizing agent while the reducing agents can be formaldehyde, glucose, hypophosphites,

eugenol, hydrazine sulfate from Vogel’s (Soine & Wilson, 1967, p.329).

Group 2 Cations (Group II Sulfides)





Antimony (III)

(1) With hydrogen sulfide, solution of

antimony (III) compounds, strongly acidified with hydrochloric acid, yield an orange precipitate of antimony sulfide that is

2 SbCl3 + 3 H2S Sb2S3 + 6 HCl (Svehla, 1996, p.100)

Orange-red precipitate (Svehla, 1996, p.100)

Double displacement precipitation reaction. Antimony (III) acts both

as a metal and non-metal while the pentavalent condition is almost exclusively as a nonmetal (Soine & Wilson, 1967, p.545). Sb2S3 is stibnite or antimony glance (Parks et al, 1949, p.149). The solution should be strongly acidified to prevent hydrolysis of SbCl6

- giving a

white precipitate of antimony oxychlorides (Gilreath, 1945, p.198) or kept as the complex ion [SbCl4]

- (Soine & Wilson, 1967, p.545).

a. insoluble in 6 N ammonium hydroxide Sb2S3 + NH4OH no reaction (USPCI, 2006, p.139)

No dissolution of precipitate; orange precipitate persists(USPCI, 2006, p.139)

NH4OH neutralizes HCl.

b. but is soluble in ammonium sulfide TS Sb2S3 + (NH4)2S 2 NH4SbS2 (Soine & Wilson, 1967, p.547)

Dissolution of precipitate forming a solution (Soine & Wilson, 1967, p.547)

Bismuth (III)

(1) When dissolved in a slight excess of nitric acid or hydrochloric acid,

Bi + 4 HNO3 Bi(NO3)3 + NO2 + H2O (Soine & Wilson, 1967, p.551)

Reddish brown fumes (Lee, 1996, p.501)

It forms compounds with oxidation state of +5 only when treated with strong oxidizing agents (Holtzclaw & Robinson, 1988, p.833).

BiOCl + 2 HCl BiCl3 + H2O (Gilreath, 1954, p.187)

Dissolution of white solids to form a clear, colorless solution H2O (Gilreath, 1954, p.187)

The reaction is reversible and the presence of the oxychloride precipitate is dependent upon a low concentration of hydronium ions (Gilreath, 1954, p.187).

bismuth salts yield a white precipitate upon dilution with water.

BiCl3 + 2 H2O Bi(OH)2Cl + 2 HCl

upon standing: Bi(OH)2Cl BiOCl + H2O (Soine & Wilson, 1967, p.551)

White precipitate (Soine & Wilson, 1967, p.551)

Hydrolysis allowed precipitation of the oxychloride (Gilreath, 1954,

p.187). Basic bismuth chloride loses water and is converted into the


oxychloride (Soine & Wilson, 1967, p.551).

Bi(NO3)3 + H2O Bi(OH)2NO3 + 2 HNO3



The same goes with the oxynitrate.

This precipitate is colored brown by hydrogen sulfide,

2 BiCl3 + 3 H2S + 6 H2O Bi2S3 + 6 HCl (Soine & Wilson, 1967, p.551)

Brownish-black precipitate Soine & Wilson, 1967, p.551)

Double displacement precipitation reaction

and the resulting compound dissolves in a warm mixture of equal parts of nitric acid and water.

Bi2S3 + 4 HNO3 2 Bi(NO3)3 + NO + 2

H2O + 3 S (Soine & Wilson, 1967, p.551)

2 NO + O2 2 NO2 (Soine & Wilson, 1967, p.62)

Dissolution of precipitate forming a solution and leaving a white precipitate (Svehla, 1996, p. 80), bubbling and reddish brown fumes (Lee, 1996, p.501)

Bismuth nitrate is not soluble in water unless an excess of nitric acid is present because of the formation of insoluble basic bismuth nitratesby the hydrolytic action of water (Parks et al, 1949, p.183). Oxidation of NO to NO2.

Copper (II)

(1) Solutions of cupric compounds, acidified with hydrochloric acid, deposit a red film of

metallic copper upon a bright, untarnished surface of a metallic iron.

CuSO4 + Fe FeSO4 + Cu (Soine & Wilson, 1967, p.320)

Red film deposited on the iron nail (Soine & Wilson, 1967, p.320)

Single replacement reaction. The position of copper in the

electromotive series is responsible for its ability to plate out on metallic iron. Since iron is higher in the series than is copper, the copper is displaces from solution by the iron, which becomes ionized (Soine & Wilson, 1967, p.320).

(2) An excess of 6 N ammonium hydroxide, added to a solution of cupric

salt, produces a bluish precipitate

2 CuSO4 + 2 NH3 + 2 H2O Cu(OH)2

CuSO4 + (NH4)2SO4 (Svehla, 1996, p.84)

Blue precipitate (Svehla, 1996, p.84)

If the solution contains ammonium salts (or it was highly acidic and larger amounts of ammonia were used up for its neutralization), precipitation does not occur at all, but the blue color is formed right away. The reaction is characteristic for copper (II) ions in the absence of nickel (Svehla, 1996, p.84). When OH

- ions are added to

cold solutions of Cu(II) salts, Cu(OH)2 is formed while with hot solutions CuO is obtained (Holtzclaw & Robinson, 1988, p.819).

and then a deep blue-colored solution. Cu(OH)2 CuSO4 + 8 NH3 [Cu(NH3)4](OH)2 + [Cu(NH3)4]SO4 (Svehla, 1996, p.84)

Dissolution of blue precipitate forming a deep-blue solution (Svehla, 1996, p.84)

Complexation. The blue precipitate dissolves because the complex removes Cu

2+ from the solution, thus disturbing the equilibrium

between the slightly ionized Cu(OH)2 and its constituent ions, and the Cu(OH)2 keeps dissolving in an effort to supply the demand for Cu

2+ (Soine & Wilson, 1967, p.319). The hydrated ion [Cu(H2O)6]

2 is

formed when the hydroxide or carbonate are dissolved in acid, or when CuSO4 or Cu(NO3)2 are dissolved in water (Lee, 1996, p.827). Copper possesses the property of forming both complex cations and anions that are quite stable (Holzclaw & Robinson, 1988, p.821).

The blue color of aqueous solutions of Cu2+

salts may be due to Cu

2+ but some exhibit green or brown, which may be due to the

undissociated molecules (Soine & Wilson, 1967, p.318).

(3) With potassium ferrocyanide TS,

solutions of cupric salts yield a reddish-brown precipitate,

2 CuSO4 + K4Fe(CN)6 Cu2Fe(CN)6 + 2 K2SO4 (Soine & Wilson, 1967, p.320)

Reddish-brown precipitate (Soine & Wilson, 1967, p.320)


a. insoluble in diluted acids. Cu2Fe(CN)6 + HCl no reaction (Soine & Wilson, 1967, p.320)

No dissolution of precipitate; reddish-brown precipitate persists (Soine & Wilson, 1967, p.320)

Lead (II) See pp.1-2.

Mercury See pp.2-3 for the reactions of mercuric ion with the general test for mercury.

Mercuric

(1) Solutions of mercuric salts yield a yellow precipitate with 1 N sodium hydroxide.

HgCl2 + 2 NaOH Hg(OH)2 + 2 NaCl

Hg(OH)2 HgO + H2O (Soine & Wilson, 1967, p.427)

Brownish-red precipitate which

turns yellow (Svehla, 1996, p.78)

The hydroxide that is first formed, spontaneously decomposes into

the oxide and water (Soine & Wilson, 1967, p.426)

(2) they yield also, in neutral solutions with potassium iodide TS, a scarlet precipitate that is

HgCl2 + 2 KI HgI2 + 2 KCl (Svehla, 1996, p.78)

Yellow precipitate which quickly turns red (Soine & Wilson, 1967, p.426)


a. very soluble in an excess of the reagent.

HgI2 + 2 KI K2[HgI4] (Svehla, 1996, p.78)

Dissolution of precipitate forming a clear, colorless solution (Svehla,

1996, p.78)

Alkalined potassium tetraiodomercurate(II) serves as a selective and sensitive reagent for ammonium ions (Nessler’s reagent) (Svehla,

1996, p.78); Valser’s reagent (Soine & Wilson, 1967, p.426)


Group 3 Cations (Group III

Sulfides and Hydroxides)





Aluminum(III)

(1) With 6 N ammonium hydroxide,

solutions of aluminum salts yield a gelatinous, white precipitate, that is

AlCl3 +3 NH3 + 3 H2O Al(OH)3 + 3 NH4Cl (Svehla, 1996, p.119)

Gelatinous, white precipitate Double displacement precipitation reaction.

a. insoluble in an excess of 6 N ammonium hydroxide

Al(OH)3 + NH4OH no reaction (Svehla, 1996, p.119)

No dissolution of precipitate; gelatinous, white precipitate persists (Svehla, 1996, p.119)

The solubility is decreased in the presence of ammonium salts (Svehla, 1996, p.119).

(2) 1 N sodium hydroxide or AlCl3 + 3 NaOH Al(OH)3 + 3 NaCl (Svehla, 1996, p.119)

Gelatinous, white precipitate

(Svehla, 1996, p.119) Double displacement precipitation reaction.

sodium sulfide TS produces the same precipitate, which

Na2S + H2O NaSH + NaOH ; NaSH + H2O H2S + NaOH (Svehla, 1996, p. 174)

2 AlCl3 + 3 Na2S + 6 H2O 2 Al(OH)3 + 3 H2S + 6 NaCl (Svehla, 1996, p. 120; Soine & Wilson, 1967, p. 451)

Gelatinous, white precipitate (Svehla, 1996, p. 120) with rotten egg odor (Holtzclaw & Robinson, 1988, p.667)

The acid, normal and polysulfides of alkali metals are soluble in water; their aqueous solutions exhbit an alkaline reaction because of hydrolysis (Svehla, 1996, p. 174).

a. dissolves in an excess of either of these reagents (1 N sodium hydroxide)

Al(OH)3 + NaOH NaAlO2 + 2 H2O (Soine & Wilson, 1967, p. 451) OR

Al(OH)3 + NaOH Na[Al(OH)4] (Svehla, 1996, p. 119)

Dissolution of precipitate forming a clear, colorless solution (Svehla, 1996, p. 119)

NaOH is a strong base and thus can give the

necessary OH- for the soluble [Al (OH)4]

- to be formed.

Al(OH)3 shows acidic properties, forming sodium aluminate. With very weak acidic properties, it is precipitated by CO2 (Lee, p.382)

b. sodium sulfide TS

Al(OH)3 + Na2S + 2 H2O Na [Al(OH)4] + H2S + 6 NaCl (Svehla, 1996, p. 118)

Dissolution of precipitate forming a clear, colorless solution with

rotten egg odor (Holtzclaw & Robinson, 1988, p.667)

The hydrolysis of S2-

is strong enough to provide the necessary OH

- for the complex to be formed (Svehla,

1996, p.174).

Cobalt (II)

(1) Solutions of cobalt salts (1 in 20) in 3 N hydrochloric acid yield a red precipitate when heated on a steam bath with an equal volume of a hot, freshly prepared

solution of 1-nitroso-2-naphthol (1 in 10) in 9N acetic acid. CoCl2 + Co(C10H6O2N)3 (Svehla,

1996, p.130)

Reddish-brown precipitate (Svehla, 1996, p.130)

Formation of a chelate complex (Svehla, 1996, p.130)

(2) Solutions of cobalt salts, when saturated with potassium chloride and treated with potassium nitrite and acetic acid, yield a

yellow precipitate.

CoCl2 + 6 KNO2 + HNO2 + CH3COOH Co(NO2)3 3

KNO2 + 2KCl + CHOOK + H2O + NO (Soine & Wilson, 1967, p.633)

2 NO + O2 2 NO2 (Soine & Wilson, 1967, p.62)

Yellow precipitate (Soine & Wilson, 1967, p.633) and reddish brown fumes (Lee, 1996, p.501)

b. in 6 N ammonium hydroxide. AgI + NH4OH no reaction (Svehla, 1996, p.196)

No dissolution of precipitte, which persists (Svehla, 1996, p.196)

Iron General test: Ferrous and ferric compounds in solution yield a black precipitate with ammonium sulfide TS.

Fe (II): FeCl2 + (NH4)2S FeS + 2NH4Cl (Svehla, 1996, p.110)

4 FeS + 6 H2O + 3 O2 4 Fe(OH)3 + 4 S (Svehla, 1996, p.114) OR

4 FeS + 9 O2 2 Fe2O(SO4)2 (Svehla, 1996, p.110)

Black precipitate, which becomes brown upon exposure to air (Svehla, 1996, p.110)

Double displacement precipitation reaction between FeCl2 + (NH4)2S. Followed by redox, and lastly the oxidation of FeS to basic iron(III) sulfate (Svehla, 1996,

p.110). Sulfide ppt: The reaction is exothermal. So much heat can be produced and may catch fire, thus never placed in a waste bin but washed way under running water; only the filter paper should be thrown away (Svehla, 1996, p.114).

Fe (III):

(a) 2 FeCl3 + 3 (NH4)2S Fe2S3+ 6 NH4Cl (Soine & Wilson, 1967, p.606) OR

Black precipitate (Svehla, 1996, p.114)

Fe3+

is not reduced by sulfide. Reaction is obtained in

alkaline solutions (Svehla, 1996, p.114).


(b) 2 FeCl3 + 3 (NH4)2S 2 FeS + S + 6 NH4Cl (Svehla,

1996, p.114)

4 FeS + 6 H2O + 3 O2 4 Fe(OH)3 + 4 S (Svehla, 1996, p.114)

Redox reaction. This reaction occurs in acidic conditions (Svehla, 1996, p.114).

a. This precipitate is dissolved by cold 3 N hydrochloric acid with the evolution of hydrogen sulfide.

From Fe (II): FeS + 2HCl FeCl2 + H2S (Svehla,

1996, p.110)

[4 FeS + 6 H2O + 3 O2 4 Fe(OH)3 + 4 S (Svehla, 1996, p.114) ]


a clear, colorless solution with rotten egg odor (Holtzclaw and Robinson, 1988, p.667)


From Fe (III) (a): Fe2S3+ 4 HCl 2 FeCl2 + 2 H2S + S (Svehla, 1996, p.114)

Dissolution of black precipitate forming a green solution while white precipitate becomes

visible (Svehla, 1996, p.114) with rotten egg odor (Holtzclaw and Robinson, 1988, p.667)

Double displacement reaction / redox.

From Fe (III) (b): FeS + 2HCl FeCl2 + H2S (Svehla,

1996, p.110) Double displacement reaction.

Ferric

(1) Acid solutions of ferric salts yield a dark blue precipitate with potassium

ferrocyanide TS (yellow prussiate of potash)

4 FeCl3 + 3 K4[Fe(CN)6] Fe4[Fe(CN)6]3 + 12 KCl (Svehla, 1996, p.115)

Intense blue precipitate (Svehla, 1996, p.115)

(Prussian blue)

Double displacement reaction. Sensitive test for ferric ion (Soine & Wilson, 1967, p.605). If FeCl3 is added to an excess of potassium hexacyanoferrate(II), a product with the composition of KFe[Fe(CN)6] is formed and this tends to form the colloidal solutions (Soluble Prussian bue) and cannot be filtered (Svehla, 1996, p.115)

(2) With an excess of 1 N sodium hydroxide, a reddish-brown precipitate.

FeCl3 + 3 NaOH Fe(OH)3 + 3 NaCl (Svehla, 1996, p.113)

Reddish-brown precipitate; brown gelatinous precipitate (Soine & Wilson, 1967, p.605)


Fe(OH)3 is insoluble in excess of the reagent (distinction from aluminum and chromium) (Svehla, 1996, p.113).

(3) With ammonium thiocyanate TS, solutions of ferric salts produce a deep red color

3 FeCl3 + 3 NH4SCN + 5 H2O 3 [Fe(H2O)5SCN]Cl2 + 3 NH4Cl (Holtzclaw and Robinson, 1988, p.812)

Deep-red solution (USP CI, 2006, p.140)

Fluorides and mercury (II) ions bleach the color (more stable hexafluoroferrate (II) [FeF6]

3-; non-dissociated

Hg(SCN)2). The presence of nitrites should be avoided since they also produce red color (nitrosyl thiocyanate, NOSCN), which disappears on heating (Svehla, 1996, p.116). Vogel reaction, especially for Co

2+ (Svehla,

p.129).

that is not destroyed by dilute mineral acids Fe(SCN)3 + HCl non-dissociated Fe(SCN)3 (Svehla, 1996, p.116)

Blood-red solution (Holtzclaw and Robinson, 1988, p.812)

Ferrous

(1) Solutions of ferrous salts yield a dark blue precipitate with potassium ferricyanide TS (red prussiate of potash)

3 FeSO4 + 2 K3[Fe(CN)6] Fe3 [Fe(CN)6]2 + 3 K2SO4 (Soine & Wilson, 1967, p.605)

Dark blue precipitate (Soine &

Wilson, 1967, p.605) (Turbull’s blue)

Double displacement reaction. Sensitive test for ferrous ion. !! The identical

composition and structure of Prussian blue and Turnbull’s blue has recently been proved by Mössbauer spectroscopy.

a. This precipitate is insoluble in 3 N hydrochloric acid but

Fe3 [Fe(CN)6]2 + HCl no reaction (Soine & Wilson, 1967, p.606)

No dissolution of precipitate; dark blue precipitate persists (Soine & Wilson, 1967, p.606)

Turnbull’s blue formation: Fe

2+ + [Fe(CN)6]

3- Fe

3+ + [Fe(CN)6]

4- First

hexacyanoferrrate (III) ions oxidize ferrous to ferric, when hexacyanoferrate(II( is formed and these ions combine to a precipitate called Turnbull’s blue: 4 Fe

3+ + 3 [Fe(CN)6]

4- Fe4[Fe(CN)6]3 <same

composition with Prussian blue!> (Svehla, 1996, p.111).


b. is decomposed by 1 N sodium hydroxide.

Fe3 [Fe(CN)6]2 + 8 NaOH 2 Na4[Fe(CN)6] + 2

Fe(OH)3 + Fe(OH)2 (Soine & Wilson, 1967, p.606)

Dark blue precipitate turned gelatinous, brown and green; the green precipitate turned brown on shaking forming a dirty-green appearance (Svehla, 1996,

p.110)

Decomposition of Fe3 [Fe(CN)6]2 (Soine & Wilson, 1967, p.606)

4 Fe(OH)2 + 2H2O + O2 4 Fe(OH)3 (Svehla, 1996, p.110)

Redox.

(2) With 1 N sodium hydroxide, solutions of

ferrous salts yield a greenish-white precipitate, the color changing rapidly to green and then to brown when shaken.

FeSO4 + 2 NaOH Fe(OH)2 + Na2SO4 (Svehla, 1996, p.110)

Greenish-white precipitate, which rapidly changes to green (Soine & Wilson, 1967, p.606) and then to brown, forming a dirty-green appearance (Svehla, 1996, p.110)


4 Fe(OH)2 + 2H2O + O2 4 Fe(OH)3 (Svehla, 1996, p.110)

Redox.

Manganese (II)

(1) With ammonium sulfide TS, solutions of

manganese salts yield a salmon-colored precipitate that

MnCl2 + (NH4)2S MnS + 2 NH4Cl (Svehla, 1996, p.137) Pink precipitate (Svehla, 1996, p.137)


a. dissolves in acetic acid. MnS + 2 CH3COOH Mn(CH3COO)2 + H2S (Svehla, 1996, p.137)


a solution (Svehla, 1996, p.137) with rotten egg odor (Holtzclaw & Robinson, 1988, p.667)

Double displacement reaction. Distinction from nickel, cobalt and zinc.

Zinc (II)

(1) In the presence of sodium acetate,

solutions of zinc salts yield a white precipitate with hydrogen sulfide.

ZnSO4 + H2S ZnS + H2SO4 (Svehla, 1996, p.140)

ZnSO4 + 2 H2S + 2 CH3COONa ZnS + H2SO4 + 2 CH3COOH + Na2SO4


ZnS is the only white sulfide precipitate. Sodium acetate is added to the solution because it reduces H

+

as it forms the feebly dissociated acetic acid, thus the sulfide-ion concentration is correspondingly increased,

and precipitation is almost complete (Svehla, 1996, p.140).

a. This precipitate is insoluble in acetic acid,

ZnS + CH3COOH no reaction (Soine & Wilson, 1967, p.408)

No dissolution of precipitate; white precipitate persists (Soine & Wilson, 1967, p.408)

Acetic acid is a weak acid and not to provide enough

H3O+ to solubilize the sulfide: ZnS + 2 H3O

+ Zn

2+ +

H2S + 2 H2O (Gilreath, 1954, p.210).

b. but is dissolved by 3 N hydrochloric acid.

ZnS + 2 HCl ZnCl2 + H2S (Gilreath, 1954, p.210).


a clear, colorless solution (Gilreath, 1954, p.210) with rotten egg odor (Holtzclaw & Robinson, 1988, p.667)

Acetic acid is a weak acid and not to provide enough

H3O+ to solubilize the sulfide: ZnS + 2 H3O

+ Zn

2+ +

H2S + 2 H2O (Gilreath, 1954, p.210).

(2) Ammonium sulfide TS produces a

similar precipitate in neutral and alkaline solutions.

ZnSO4 + (NH4)2S ZnS + (NH4)2SO4 (Soine & Wilson, 1967, p.407)



(3) With potassium ferrocyanide TS, zinc

salts in solution yield a white precipitate that is

2 ZnSO4 + K4[Fe(CN)6] Zn2[Fe(CN)6] + 2 K2SO4 (Soine & Wilson, 1967, p.408)



a. insoluble in 3 N hydrochloric acid. Zn2[Fe(CN)6] + HCl no reaction (Soine & Wilson, 1967, p.408)

No dissolution of precipitate; white precipitate persists(Soine & Wilson, 1967, p.408)

Group 4 Cations (Group IV

Carbonates)





Barium (II)

(1) Solutions of barium salts yield a white precipitate with 2 N sulfuric acid

BaCl2 + H2SO4 BaSO4 + 2 HCl (Svehla, 1996, p. 145)

Heavy, white precipitate (Svehla, 1996, p. 145)

Double displacement precipitation reaction.

a. this precipitate is insoluble in hydrochloric acid and nitric acid

BaSO4 + HCl no reaction

BaSO4 + HNO3 no reaction (Svehla, 1996, p. 145)

No dissolution of precipitate; white precipitate persists (Svehla, 1996, p. 145)

The solubility of very small crystals is greater than that of larger ones with pronounced effect in hard crystals, like BaSO4, due to surface


energy (Gilreath, 1954, p.96).

(2) Barium salts impart a yellowish-green color to a nonluminous flame that appears blue when viewed throughgreen glass.

Δ

BaCl2 Ba + 2 Cl, Ba Ba*, Ba*

Ba + hv (Petrucci and Harwood, 1997, p.766)

Yellowish-green flame (Svehla, 1996, p.7) or weak and fleeting yellowish-green flame (Holtzclaw & Robinson, p.986)

The energy from the flame excites an electron to a higher energy level, and when it falls back to the lower energy level the extra energy is given out as light (Lee, 1996, p.287).

Calcium (II)

Solutions of calcium salts form insoluble oxalates when treated as follows. (1) To a solution of the calcium salt (1 in 20) add 2 drops of methyl red TS, and neutralize with 6 N ammonium hydroxide. Add 3 N hydrochloric acid, dropwise, until the solution is acid to the indicator. Upon the addition of ammonium

oxalate TS a white precipitate is formed.

CaCl2 + (NH4)2 (COO)2 Ca(COO)2 + 2 NH4Cl (Svehla, 1996, p.150)

White, crystalline precipitate (Svehla, 1996, p.234)

Double displacement reaction. Calcium oxalate tends to form supersaturated solutions and may not precipitate unless the solution is heaed (Gilreath, 1954, p.217).

a. This precipitate is insoluble in 6 N acetic acid Ca(COO)2 + CH3COOH no reaction (Svehla, 1996, p.150)

No dissolution of precipitate; precipitate persists (Svehla, 1996, p.150)

b. but dissolves in hydrochloric acid. Ca(COO)2 + 2 HCl CaCl2 + H2O + CO +

CO2 (Svehla, 1996, p.234)

Dissolution of precipitate forming a clear, colorless solution, which bubbles (Svehla, 1996, p.234)

Decomposition of oxalate (Svehla, 1996, p.234)

(2) Calcium salts moistened with hydrochloric acid impart a transient yellowish-red color to a nonluminous flame.

Δ

CaCl2 Ca + 2 Cl, Ca Ca*, Ca*

Ca + hv (Petrucci and Harwood, 1997, p.766)

Transient yellowish-red flame (Soine & Wilson, 1967, p.375) or brick-red (yellowish) flame (Svehla, 1996, p.7)

The energy from the flame excites an electron to

a higher energy level, and when it falls back to the lower energy level the extra energy is given out as light (Lee, 1996, p.287).

Group 5 Cations Soluble Ions)





Ammonium(I)

(1) Ammonium salts are decomposed by the addition of an excess of 1 N sodium hydroxide, with the evolution of ammonia, recognizable by its odor and by its alkaline effect upon

Δ

NH4Cl + NaOH NH3 + H2O + NaCl (Soine & Wilson, 1967, p.238)

Ammonia odor (Soine & Wilson, 1967, p.302)

Neutralization reaction: NH4

+ is essentially an acid which is capable of ionizing into

ammonia and a proton, and OH- is a strong base capable of

combining with proton.

NH4+ NH3 + H

+ (Soine & Wilson, 1967, p.302) moistened red litmus paper exposed to the vapor.

Warming the solution accelerates the decomposition. NH3 + H2O NH4

+ + OH

- (Hotzclaw, 1988,

p.994)

red litmus turns blue

(Hotzclaw, 1988, p.994)

Magnesium (II)

(1) Solutions of magnesium salts in the presence of ammonium chloride yield no more than a slightly hazy precipitate when neutralized with ammonium carbonate TS,

NH4OH NH4+ + OH

-

NH4Cl NH4+ + Cl

- (Soine & Wilson, 1967,

p.355)

Slightly hazy precipitate (USP CI, 2006, p.140)

The chemical reactions of interest that involve Mg2+

are based upon solubility since the divalent ion is neither reduced nor oxidized. NH4OH incompletely precipitates Mg(OH)2 from solution of Mg salts. In the presence of an NH4

+ salt, NH4OH does not precipitate Mg

2+, because the

concentration of OH- produced by the NH4OH is depressed

by the excess of NH4+ from the NH4Cl by the law of mass

action. When the concentration of Mg

2+is multiplied by the square of

the concentration of the OH-, there is obtained a product that

is less than the ion-product constant (solubility product) of Mg(OH)2 and hence no precipitate will be formed (Soine &

Wilson, 1967, p.355).

MgCl2 + 2 NH4OH Mg(OH)2 + 2 NH4Cl (Svehla, 1996, p.153)

but on the subsequent addition of dibasic sodium phosphate TS, a white crystalline

precipitate, which is

Na2HPO4 2 Na+ + HPO4

2-

HPO42-

+ NH3 NH4+ + PO4

3-

MgCl2 + NH4Cl + Na3PO4 + 6 H2O

MgNH4PO4 6H2O

MgCl2 + Na2HPO4 + NH3 + 6 H2O

MgNH4PO46H2O + 2 NaCl (Soine & Wilson, 1967, p.355)

White, crystalline precipitate (Soine & Wilson, 1967, p.355)

NH4Cl – prevents precipitation of Mg(OH)2 (Soine & Wilson, 1967, p.356)

insoluble in 6 N ammonium hydroxide, is formed. MgNH4PO4 + NH4OH no reaction (Soine & White, crystalline precipitate NH4OH – lessens the hydrolysis and decreases the


Wilson, 1967, p.356) (Soine & Wilson, 1967, p.355) solubility of the salt

MgNH4PO4 + H2O Mg2+

+ HPO42-

+ NH4OH (Soine & Wilson, 1967, p.356)

Potassium (I)

(1) Potassium compounds impart a violet color to a nonluminous flame, but the presence of small quantities of sodium masks the color unless the yellow color produced by sodium is screened out by viewing through a blue filter that blocks at 589 nm (sodium) but is transparent to emission at 404 nm (potassium). Traditionally, cobalt glass has been used, but other suitable filters are commercially available.

Δ

KCl K + Cl, K K*, K* K + hv (Petrucci and Harwood, 1997, p.766)

Violet (lilac) (Svehla, 1996, p.7)

Color arises because the energy absorbed or emitted in electronic transition corresponds to a wavelength in the

visible region. Alkali Metal ions all have noble gas configurations in which all the electrons are paired. Thus the required energy in promoting an electron is large since some for unpairing an electron, some for breaking a shell of electrons and some to promote to a higher energy level. Hence no transition and the compounds are typically white,

except those where the anion is colored (Lee, 1996, p.288).

(2) In neutral, concentrated or moderately

concentrated solutions of potassium salts (depending upon the solubility and the potassium content), sodium bitartrate TS produces a white crystalline

precipitate that is

KCl + NaHC4H4O6 KHC4H4O6 + NaCl (Svehla, 1996, p.156)

White, crystalline precipitate; granular (Soine & Wilson, 1967, p.270)

Double displacement reaction. Ammonium salts yield a similar precipitate and must be absent.

a. soluble in 6 N ammonium hydroxide and in solutions of alkali hydroxides and carbonates.

The formation of the precipitate, which is usually slow, is accelerated by stirring or rubbing the inside of the test tube with a glass rod. The addition of a small amount of glacial acetic acid or alcohol also promotes the precipitation.

Dissolution of precipitate forming a clear, colorless solution (USP CI, 2006, p.141)

Sodium (I)

(1) Unless otherwise specified in an individual

monograph, prepare a solution to contain 0.1 g of the sodium compound in 2 mL of water. Add 2 mL of 15% potassium carbonate, and heat to boiling. No precipitate is formed. Add 4 mL of potassium pyroantimonate TS, and heat to boiling. Allow to cool

in ice water, if necessary, rub the inside of the test

tube with a glass rod. A dense precipitate is formed.

K2CO3

2 NaCl + 2 KH2SbO4 Na2H2Sb2O7 + 2 KCl + H2O Δ (Soine and Wilson, 1967, p.193)

White, crystalline precipitate (Soine and Wilson, 1967, p.193)


(2) Sodium compounds impart an intense yellow color to a nonluminous flame.

Δ

NaCl Na + Cl, Na Na*, Na* Na + hv (Petrucci and Harwood, 1997, p.766)

Persistent golden-yellow flame (Svehla, 1996, p.7)


Lithium (I)

(1) With sodium carbonate TS, moderately

concentrated solutions of lithium salts, made alkaline with sodium hydroxide, yield a white precipitate on

boiling.

Li2SO4 + NaOH 2 LiOH + Na2SO4

(Soine and Wilson, 1967, p.188) Δ

2 LiOH + Na2CO3 Li2CO3 + 2 NaOH (Soine and Wilson, 1967, p.188)

White precipitate on boiling (Soine and Wilson, 1967, p.188)

Li2CO3 is not very souble and precipitates from solution more readily than the other alkali metal carbonates (Holtzclaw & Robinson, 1988, p.375).

a. The precipitate is soluble in ammonium

chloride TS. Li2CO3 + 2 NH4Cl (NH4)2CO3 + 2 LiCl (Svehla, 1996, p.311)


No precipitation occurs in the presence of high concentrations of NH4Cl since the carbonate-ion

concentration is reduced to such an extent that the solubility product of Li2CO3 is not exceeded (Svehla, 1996, p.311).

(2) Lithium salts moistened with hydrochloric acid

impart an intense crimson color to a nonluminous flame.

Δ

LiCl Li + Cl, Li Li*, Li* Li + hv (Petrucci and Harwood, 1997, p.766)

Intense crimson flame (USP Ci, 2006, p.140)


(3) Solutions of lithium salts are not precipitated by 2 N sulfuric acid or soluble sulfates (distinction from

strontium)

LiCl + H2SO4 no reaction No white precipitate (Svehla, 1996, p.147)


Anions may be classified according to the processes employed and the corresponding results or reactions. Class A anions: those involving the identification by volatile products obtained on treatment with acids. It is subdivided into:

(1) gases evolved with dilute HCl or H2SO4, and (2) gases or vapors evolved with concentrated H2SO4, includes also anions in (1)

Class B anions: those dependent upon reactions in solution. It is subdivided into: (1) precipitation reactions (2) oxidation and reduction in solution (Svehla, 1996, p.163).

Anions may also be separated into groups based upon the solubilities of Calcium, Barium, Cadmium, and Silver salts. 1Group I includes ions whose calcium salts are insoluble in neutral or slightly basic solution. Anions under this group are AsO4

2-, AsO2

-, CO3

2-, F

-, C2O4

-, PO4

3-, and SO3

2-.

2Group II includes ions whose calcium salts are soluble but whose barium salts are insoluble slightly basic solution. Anions under this group are CrO4

- and SO4

2-.

3Group III is composed of anions whose calcium and barium salts are soluble but whose cadmium salts are insoluble slightly bas ic solution. Anions under this group are [Fe(CN)6]

3-, [Fe(CN)6]

4- andS

2-.There are no any official

identification tests for these Group III anions. 4Group IV is composed of anions whose calcium, barium and cadmium salts are soluble but whose silver salts are insoluble in a solution slightly acid with nitric acid. Anions under this group are Br

--,, Cl

-, I

-, CNS

- and S2O3

2-.

5Group V contains anions whose calcium, barium, cadmium and silver salts are soluble in water and acids. Anions under this group are B4O7

2-, ClO3

-, NO3

-, and NO2

- (Gilreath, 1954, p.225).

*not included in this scheme.

ANION DESCRIPTION OF THE TESTS CORRESPONDING EQUATIONS PERCEPTIBLE RESULTS PRINCIPLE INVOLVED/ NOTES

Class A (1) Gases evolved with dilute HCl or H2SO4

Carbonate1

(1) Carbonate and bicarbonates effervesce with acids evolving a colorless gas that,

CaCO3 + 2 HCl CaCl2 + CO2 + H2O (Svehla, 1996, p.164)

Bubbling of the solution (Svehla, 1996, p.164)

Decomposition od the carbonate (Svehla, 1996, p.164)

when passed into calcium hydroxide TS, produces a

white precipitate immediately. CO2 + Ca(OH)2 CaCO3 + H2O (Svehla, 1996, p.164)

The clear, colorless solution becomes turbid (Svehla, 1996, p.164)

Neutralization reaction.

(2) A cold solution (1 in 20) of a soluble carbonate is colored red by phenolphthalein TS,

Red solution (USP CI, 2006, p.140)

(Bicarbonate)* while a similar solution of BICARBONATE remains unchanged or is slightly only colored.

Slightly pinkish solution (USP CI, 2006, p.140)

Sulfite1

(1) when treated with 3 N hydrochloric acid, sulfites

and bisulfites yield sulfur dioxide, Na2SO3 + 2 HCl SO2 + H2O + 2 NaCl (Svehla, 1996, p.168)

Suffocating odor of burning sulfur or acrid odor (Svehla, 1996, p.168)

Decomposition (Svehla, 1996, p.168)

which blackens filter paper moistened with mercurous nitrate TS.

SO2 + Hg2(NO3)2 + 2 H2O 2 Hg + 2 HNO3 + H2SO4 (Soine and Wilson, 1967, p.575)

Filter paper turns black (Soine and Wilson, 1967, p.575)

Redox with the mercurous nitrate reduced

to metallic mercury (Soine and Wilson, 1967, p.575)

Thiosulfate4

(1) With hydrochloric acid, solutions of thiosulfates

yield a white precipitate that turns yellow, and sulfur dioxide,

Na2S2O3 + 2 HCl 2 NaCl + H2S2O3 ; H2S2O3 S+

SO2 + H2O ------------------------------------------------------------

Na2S2O3 + 2 HCl 2 NaCl + S+ SO2 + H2O (Svehla, 1996, p.171)

White precipitate that turns yellow with acrid odor or suffocating odor of burning sulfur (Svehla, 1996, p.171)

Disproportionation reaction. Sodium thiosulfate is unstable in the presence of acids because free thiosulfuric acid is liberated (Parks et al, 1949, p.144).

which blackens filter paper moistened with mercuruous nitrate TS.

SO2 + Hg2(NO3)2 + 2 H2O 2 Hg + 2 HNO3 + H2SO4 (Soine and Wilson, 1967, p.575)

Filter paper turns black (Soine and Wilson, 1967, p.575)

Redox with the mercurous nitrate reduced to metallic mercury (Soine and Wilson, 1967, p.575)

(2) The addition of ferric chloride TS to solutions of

thiosulfates produces a dark violet color 2 Na2S2O3 + FeCl3 Na[Fe(S2O3)2] + 3 NaCl (Svehla, 1996, p.173)

Dark-violet solution (Svehla, 1996, p.173)


that quickly disappears. Na[Fe(S2O3)2] + FeCl3 Fe(S4O6) + FeCl2 + NaCl (Svehla, 1996, p.173)

Upon standing the violet color disappears and a green solution may be formed (Svehla, 1996, p.173)

The overall reaction can be written as the

reduction of iron(III) by thiosulfate: 2 S2O3

2- + 2 Fe

3+ S4O6

2- + 2 Fe

2+

(Svehla, 1996, p.173)

Nitrite5 (1) When treated with dilute mineral acids or with 6

N acetic acid, nitrites evolve brownish-red fumes. NaNO2 + HCl HNO2 + NaCl (2 HNO2 H2O + N2O3)

3 HNO2 HNO3 + 2 NO + H2O Transient, pale-blue solution

Disproportionation reaction: NO2

- + H2O NO3

- + 2 H

+ + 2 e

-

2 [NO2- + 2 H

++ e

- NO + H2O]

3 NO2- + 2 H

+ NO3

- + 2 NO + H2O


NO + O2 2 NO2 (Svehla, 1996, p.177) with reddish-brown fumes (Svehla, 1996, p.177)

Redox reaction with NO oxidized to NO2. Combination of NO and O2 to form nitrogen dioxide (Svehla, 1996, p.177)

The solution colors starch-iodide paper blue.

2 HNO2 + 2 KI + 2 HCl I2 + 2 NO + 2 KCl + 2 H2O (Svehla, 1996, p.177)

Blue paper (USP CI, 2006, p.140)

Redox reaction with HNO2 acting as an

oxidizing agent.

I2 + starch-iodide starch-iodo complex Formation of adsorption complex

Sulfide3

No official ID test Hypochlorite*

Cyanide*

Cyanate*

Class A (2) Gases or acid vapors evolved with concentrated H2SO4, including anion of Class A (1)

Chloride4

(1) With silver nitrate TS, solutions of chlorides yield

a white, curdy precipitate that is NaCl + AgNO3 AgCl + NaNO3 (Svehla, 1996, p.191)

White, curdy precipitate (Svehla, 1996, p.191)


(hv)

2 AgCl 2 Ag + Cl2 (Svehla, 1996, p.73)

Grayish or black precipitate (Svehla, 1996, p.73)

Sunlight or uv irradiation decomposes silver halides, which turn to grayish or balck owing to the formation of silver metal. The reaction is slow though (Svehal

,1996, p.73).

a. insoluble in nitric acid but AgCl + HNO3 no reaction (Svehla, 1996, p.191)

No dissolution of precipitate; Grayish or black precipitate persists (Svehla, 1996, p.73)

b. is soluble in a slight excess of 6 N ammonium hydroxide.

AgCl + NH4OH [Ag(NH3)2]Cl + H2O (Svehla, 1996, p.191)

Dissolution of precipitate forming a clear, colorless solution

Formation of the complex silver ammonia cation. Silver chloride ionizes to a greater

extent than does silver ammonia complex cation (in an excess of NH4OH) and, therefore, silver ions are used up in the formation of complex cation (Soine and Wilson, 1967, p.329). Caution! See Bromide.

(2) When testing amine (including alkaloidal hydrochlorides) that do not respond to the above test, add one drop of diluted nitric acid and 0.5 mL of silver nitrate TS to a solution of the substance being

examined containing, unless otherwise directed in the monograph, about 2 mg of chloride ion in 2 mL; a

white, curdy precipitate is formed

White, curdy precipitate (USP CI, 2006, p.140)

Centrifuge the mixture without delay, and decant the supernatant layer. Wash the precipitate with three 1-mL portions of niric acid solution (1 in 100), and discard the washings. Add ammonia TS dropwise to this precipitate. It dissolves readily.

Dissolution of precipitate; white precipitrate persists (USP CI, 2006, p.140)

(3) When a monograph specifies that an article responds to the test for dry chlorides, mix the solid to be tested with an equal weight of manganese dioxide, moisten with sulfuric acid, and gently heat the mixture: chlorine,

Δ

2 NaCl + MnO2 + 2 H2SO4 MnSO4 + Cl2 + 2 H2O +

Na2SO4 OR Δ

2 NaCl + MnO(OH)2 + 2 H2SO4 MnSO4 + Cl2 + 3 H2O + Na2SO4 (Schela, 1996, p.191)

Formation of yellow-green gas with suffocating odor (Schela, 1996, p.191)

Redox reaction.

which is recognizable by the production of a blue color with moistened starch iodide paper, is

evolved.

Cl2 + KI KCl + I2 I2 + starch starch-iodo complex

Paper turned blue (USP CI, 2006, p.140)

Redox Formation of adsorption complex

Bromide4 (1) Solutions of bromides, upon addition of chlorine

TS, dropwise, liberate bromine, 4 NaClO + 4 HCl 4 NaCl + 2 Cl2 +O2 + 2 H2O (Soine and Wilson, 1967, p.243)

Greenish yellow gas having characterisitc, unpleasant and

Redox. In practice, it is more convenient to use


suffocating odor (Soine and Wilson, 1967, p.82)

dilute sodium hypochlorite solution, acidified with dilute hydrochloric acid to generate chlorine (Svehla, 1996, p.193)

2 NaBr + Cl2 Br2 + NaCl (Svehla, 1996, p. 194)

Orange-red solution

Redox. A halogen of lower atomic weight will always displace one of higher atomic

weight from its binary hydrogen compounds of from the salts thereof. Cl2 is an oxidizing agent, thus it is reduced as chloride. (Soine and Wilson, 1967, p.76

Bromide

4

(continued)

which is dissolved by shaking with chloroform,

coloring the chloroform red to reddish brown. Br2 + CHCl3 (Br2)CHCl3 (Shriner et al, 1998, p. 88)

Reddish brown solution below the colorless aqueous layer (Svehla,

1996, p. 193)

Solvent extraction

(2) Silver nitrate TS produces in solutions of bromides a yellowish-white precipitate that is

SrBr2 + 2 AgNO3 2 AgBr + Sr(NO3)2 (Svehla, 1996, p.193)

Curdy, pale-yellow precipitate (Svehla, 1996, p.193)


(hv)

2 AgBr 2 Ag + Br2 (Svehla, 1996, p.73)


Sunlight or uv irradiation decomposes silver halides, which turn to grayish or black owing to the formation of silver metal. The reaction is slow though (Svehla ,1996, p.73).

a. insoluble in nitric acid and AgBr + HNO3 no reaction (Svehla, 1996, p.193)

No dissolution of precipitate; precipitate persists (Svehla, 1996, p.193)

b. is slightly soluble in 6 N ammonium hydroxide. AgBr + 2 NH4OH [Ag(NH3)2] Br + 2 H2O (Svehla, 1996, p.191)


Formation of soluble complex. ! The solution obtained dissolving in ammonia

should be discarded quickly to avoid explosion or acidified quickly with 2 M nitric acid and disposed of. Otherwise, when set aside, a precipitate of silver nitride Ag3N (fulminating silver†) is formed slowly, which explodes readily even in a wet form (Svehla, 1996, p.72).

Iodide4

(1) Solutions of iodides, upon the addition of chlorine TS, dropwise, liberate iodine, which colors the solution

yellow to red. 2 NaI+ Cl2 I2 + 2 NaCl (Soine and Wilson, 1967, p.128)

Yellow to red solution (USP CI, 2006, p.140)

Redox

When the solution is shaken with chloroform, the

latter is colored violet. I2 + CHCl3 (I2)CHCl3 (Shriner et al, 1998, p. 88)

Violet solution below the colorless aqueous layer (Svehla, 1996, p. 196)

Solvent extraction

The iodine thus liberated gives a blue color with starch TS.

I2 + starch starch-iodo complex (Soine and Wilson, 1967, p.96)

Blue solution (Soine and Wilson, 1967, p.96)

Formation of adsorption complex

(2) Silver nitrate TS produces, in solutions of iodides,

a yellow, curdy precipitate that is NaI + AgNO3 AgI + NaNO3 (Svehla, 1996, p.196)

Yellow, curdy precipitate (Svehla, 1996, p.196)


(hv)

2 AgI 2 Ag + I2 (Svehla, 1996, p.73)


Sunlight or uv irradiation decomposes silver halides, which turn to grayish or balck owing to the formation of silver metal. The reaction is slow though (Svehal

,1996, p.73).

a. insoluble in nitric acid and AgI + HNO3 no reaction (Svehla, 1996, p.196) No dissolution of precipitate; Yellow, curdy to grayish or black precipitate persists (Svehla, 1996, p.196)

Ag+ ion is small and highly polarizing and

the anion, I-, is large and highly

polarizable. This leads to some covalent character (Lee, 1996, p.825).

b. in 6 N ammonium hydroxide. AgI + NH4OH no reaction (Svehla, 1996, p.196) No dissolution of precipitate; Yellow, curdy to grayish or black precipitate

persists (Svehla, 1996, p.196)


Nitrate5

(1) When a solution of nitrate is mixed with an equal volume of sulfuric acid, the mixture is cooled, and a solution of ferrous sulfate is superimposed, a brown

color is produced at the junction of the two liquids. “Brown ring test”

6 FeSO4 + HNO3 + 3 H2SO4 2 NO + 3 Fe2(SO4)3 + 4 H2O

FeSO4 + NO FeSO4 NO (Soine and Wilson, 1967, p.151)

FeSO4 + NO [Fe(NO)]SO4 (Svehla, 1996, p.200)

Formation of purple ring changing to brown at the junction of two liquids (USP CI, 2006, p.140)

Redox. Reduction of nitrate to nitric oxide. On shaking and warming the mixture the brown color disappears, nitrogen dioxide is evolved, and a yellow solution of iron(III) ions remains (Svehla, 1996, p.200).

(2) When a nitrate is heated with sulfuric acid and metallic copper, brownish-red fumes are evolved.

Δ

2 NaNO3 + 4 H2SO4 + 3 Cu Na2SO4 + 3 CuSO4 + 2 NO + 4 H2O (Svehla, 1996, p.200)

2 NO + O2 2 NO2 (nitric oxide nitrogen dioxide)

Blue solution with reddish-brown fumes (Svehla, 1996, p.199)

Redox.

Nitrate (continued)

(3) Nitrates do not decolorize acidified potassium permanganate TS (distinction from nitrites).

HNO3 + KMnO4 + H2SO4 no reaction Solution remains purple (USP CI, 2006, p.140)

NO3 cannot be further oxidized. It only acts as an oxidizing agent wherein it could

be reduced to NO or NO2, NH3 or N2.

5 HNO2 + 2 KMnO4 + 3 H2SO4 5 HNO3 + 2 MnSO4 + 3 H2O + K2SO4 (Soine and Wilson, 1967, p.250)

Purple solution becomes colorless with no fumes (Svehla, 1996, p.197)

Redox: NO2- acts as RA.

Borate5

(1) To 1 mL of a borate solution, acidified with hydrochloric acid to litmus,

Na2B4O7 + 2 HCl + 5 H2O 4 H3BO3 + 2 NaCl (Svehla, 1996, p.209)

Solution remains clear and colorless

but white precipitate may form on standing (Svehla, 1996, p.209)

Double displacement reaction

add 3 or 4 drops of iodine TS and 3 or 4 drops of polyvinyl alcohol solution (1 in 50): an intense blue

color is produced.

2 I- I2

I2 + (C2H4O)n adsoprtion complex Intense blue color

Due to the borate present in the solution, iodide was oxidized to iodine which forms the blue complex with the alcohol.

(2) When a borate is treated with sulfuric acid,

methanol is added, and the mixture is ignited, it burns

with a green-bordered flame.

Na2B4O7 + H2SO4 + 5 H2O 4 H3BO3 + 2 Na2SO4 H3BO3 + 3 CH3OH (CH3)3BO3 + 3 H2O (Soine and Wilson, 1967, p.210)

green-bordered flame (Soine and Wilson, 1967, p.210)

Esterification of the boric acid to form

trimethyborate (Soine and Wilson, 1967, p.210)

Acetate*

(1) When acetic acid or an acetate is warmed with sulfuric acid and alcohol, ethyl acetate, recognizable by its characteristic odor is evolved

Δ Zn (CH3COO)2 + 2 H2SO4 + 2 CH3CH2OH

CH3COOCH2CH3 + Zn (HSO4)2 + 2 H2O (Svehla, 1996, p.230)

Characteristic pleasant, fruity odor of ethyl acetate (Svehla, 1996, p.230)

Esterification

(2) With neutral solutions of acetates, ferric chloride TS produces a deep red color that is destroyed by the addition of mineral acids

3 Zn (CH3COO)2 + 3 FeCl3 + 2 H2O [Fe3(OH)2(CH3COO)6]Cl + 3 ZnCl2 + 2 HCl (Svehla, 1996, p.231)

Deep-red solution (Svehla, 1996, p.231)

Complexation

Oxalate1

(1) Neutral and alkaline solutions of oxalates yield a white precipitate with calcium chloride TS.

(COONa)2 + CaCl2 (COO)2Ca + 2 NaCl (Svehla, 1996, p.234)



a. This precipitate is insoluble in 6 N acetic acid

but is (COO)2Ca + CH3COOH no reaction (Svehla, 1996, p.234)

No dissolution of precipitate; white,

crystalline precipitate persists(Svehla, 1996, p.234)

Acetic acid is too weakly acidic to provide

enough H3O+ to form the soluble oxalic

acid (Gilreath, 1954, p.230).

b. dissolved by hydrochloric acid. (COO)2Ca + HCl CaCl2 + H2(COO)2 (Gilreath, 1954, p.230)

Dissolution of precipitate forming a clear, colorless solution (USP CI, 2006, p.140)


(2) Hot acidified solutions of oxalates decolorize potassium permanganate.

(COONa)2 + 2 KMnO4 + 3 H2SO4 2 CO2 + 2 MnSO4 + 3 H2O + K2SO4 + Na2SO4 (Svehla, 1996, p.235)

Purple solution becomes colorless

with bubbling (Svehla, 1996, p.235) Redox.

Chlorate5

(1) Solutions of chlorates yield no precipitate with silver nitrate TS.

KClO3 + AgNO3 no reaction (Svehla, 1996, p.203) No precipitate (Svehla, 1996, p.203)

(2) The addition of sulfurous acid to this mixture

produces a white precipitate that is

2 KClO3 + H2SO4 2 HCl + K2SO4 + 3 O2

HCl + AgNO3 AgCl +HNO3 (Gilreath, 1954, p.243)

White precipitate (Gilreath, 1954, p.243)

The chlorate ion is reduced by sulfuric acid to the chloride ion; the latter is then precipitated as AgCl (Gilreath, 1954, p.243)

a. insoluble in nitric acid, AgCl +HNO3 no reaction (USP CI, 2006, p.140) No dissolution of precipitate; white precipitate persists (USP CI, 2006,


p.140)

b. but is soluble in 6 N ammonium hydroxide. AgCl + NH4OH [Ag(NH3)2]Cl + H2O (Svehla, 1996, p.191)

Dissolution of precipitate forming a

clear, colorless solution (Svehla, 1996, p.191)

Formation of the complex silver ammonia cation. Silver chloride ionizes to a greater extent than does silver ammonia complex

cation (in an excess of NH4OH) and, therefore, silver ions are used up in the formation of complex cation (Soine and Wilson, 1967, p.329). Caution! See Bromide.

(3) Upon ignition, chlorates yield chlorides, recognizable by appropriate tests.

2 KClO3 2 KCl + 3 O2 (Svehla, 1996, p.204) Crackling sound (Svehla, 1996, p.202)

Decomposition reaction (Svehla, 1996, p.204)

When sulfuric acid is added to a dry

chlorate, decrepitation occurs, and a greenish yellow-gas is evolved. [caution – use only a small amount of chlorate for this test, and exercise extreme caution in performing it.]

3 KClO3 + H2SO4 ClO2 + ClO4 + K2SO4 + H2O + KCl (Svehla, 1996, p.202)

Orange-yellow solution with greenish-yellow fumes (Svehla, 1996, p.202)

Decomposition reaction (Svehla, 1996, p.202)

Per-manganate*

(1) Solutions of permanganate acidified with sulfuric

acid are decolorized by hydrogen peroxide TS and by

2 KMnO4 + 3 H2SO4+ 5 H2O2 K2SO4 + 2 MnSO4 + 8 H2O

+ 5 O2 (Soine & Wilson, 1967, p.286)

Purple solution becomes colorless

with bubbling (Soine & Wilson, 1967, p.286)

Redox.

sodium bisulfite TS, in the cold, 5 NaHSO3 + 2 KMnO4 + 3 H2SO4 K2SO4 + 2 MnSO4 + 5 NaHSO4 + 3 H2O (Soine & Wilson, 1967, p.287)

Purple solution becomes colorless (Soine & Wilson, 1967, p.287)

Redox.

and by oxalic acid TS, in hot solution. 5 H2C2O4 2 H2O + 2 KMnO4 + 3 H2SO4 K2SO4 + 2

MnSO4 + 18 H2O + 10 CO2 (Soine & Wilson, 1967, p.287)

Purple solution becomes colorless with bubbling


Redox.

Thiocyanate4

(1) With ferric chloride TS, solutions of thiocyanates yield a red color that is not destroyed by moderately concentrated mineral acids.

3 KSCN + FeCl3 Fe(SCN)3 + 3 KCl (Svehla, 1996, p.184)

Blood-red solution


6 KSCN + FeCl3 K3[Fe(CNS)6] + 3 KCl (Gilreath, 1954, p.239)

Complexation (Gilreath, 1954, p.239)

Tartrate*

(1) Dissolve a few mg of a tartrate salt in 2 drops of

sodium metaperiodate solution (1 in 20). Add a drop of 1 N sulfuric acid, and after 5 minutes add a few drops of sulfurous acid followed by a few drops of fuchsin-sulfurous acid TS: a reddish-pink color is

produced within 15 minutes.

Reddish-pink solution within 15 minutes (USP CI, 2006, p.141)

Citrate*

(1) To 15 mL of pyridine add a few mg of a citrate salt, dissolved or suspended in 1 mL of water, and

shake. To this mixture add 5 mL of acetic anhydride,

and shake: a light red color is produced.

Light red solution (USP CI, 2006,

p.140)

Fluoride1

No official ID test

Formate*

Hexa-fluorosilicate*

Perchlorate*

Bromate*

Hexacyanoferrate (II)

3

Hexacyano-ferrate (III)

3

Class B (1) Precipitation Reactions

Sulfate2

(1)With barium chloride TS, solutions of sulfates

yield a white precipitate that is BaCl2 + H2SO4 BaSO4 + 2 HCl (Svehla, 1996, p. 145)

Heavy, white precipitate (Svehla, 1996, p. 145)



a. insoluble in hydrochloric acid and nitric acid, BaSO4 + HCl no reaction

BaSO4 + HNO3 no reaction (Svehla, 1996, p. 145)

No dissolution of precipitate; white precipitate persists (Svehla, 1996, p. 145)

The solubility of very small crystals is greater than that of larger ones with pronounced effect in hard crystals, like BaSO4, due to surface energy (Gilreath, 1954, p.96).

(2) With lead acetate TS, neutral solutions of sulfates

yield a white precipitate that is (CH3COO)2Pb + H2SO4 PbSO4 + 2 CH3COOH (Soine and Wilson, 1967, p.157)

White precipitate (Soine and Wilson, 1967, p.157)


a. soluble in ammonium acetate TS. PbSO4 + 2 CH3COONH4 (CH3COO)2Pb + (NH4)SO4 Dissolution of white precipitate forming a clear, colorless solution


(3) Hydrochloric acid produces no precipitate when

added to solutions of sulfates (distinction form thiosulfates).

H2SO4 + HCl no reaction No precipitate (USP CI, 2006, p.141)

Phosphate1

[Note – Where the monograph specifies the identification test for phosphate, use the tests for orthophosphates, unless the instructions specify the use of pyrophosphate tests or indicate that the product to be ignited before performing the test.]

Ortho-phosphate

(1) With silver nitrate TS, neutral solutions of

orthophosphates yield a yellow precipitate [p.141], that is

Na2HPO4 + 3 AgNO3 Ag3PO4 + 2 NaNO3 + HNO3 (Svehla, 1996, p.219)

Yellow precipitate (Svehla, 1996, p.219)


a. soluble in 2 N nitric acid and Ag3PO4 + 2 HNO3 AgH2PO4 + 2 AgNO3 (Svehla, 1996, p.219)



b. in 6 N ammonium hydroxide. Ag3PO4 + 6 NH4OH [Ag(NH3)2]3PO4 + 6 H2O (Svehla, 1996, p.219)


Formation of soluble complex. ! The

solution obtained dissolving in ammonia should be discarded quickly to avoid explosion or acidified quickly with 2 M nitric acid and disposed of. Otherwise, when set aside, a precipitate of silver nitride Ag3N (fulminating silver†) is formed slowly, which explodes readily even in a

wet form (Svehla, 1996, p.72).

(2) With ammonium molybdate TS, acidified

solutions of orthophosphates yield a yellow precipitate that is

Na2HPO4 + 12 (NH4)2MoO4 + 23 HNO3

(NH4)3[P(Mo3O10)4] + 12 H2O + 2 NaNO3 (Svehla, 1996,

p.219)

Yellow, crystalline precipitate Svehla, 1996, p.219)

Mo3O10 group replaces each oxygen atom in the phosphate. Commercial ammonium molybdate, a heptamolybdate,

(NH4)6Mo7O244H2O (Svehla, 1996, p.220). Phosphates yield the precipitate in the cold or upon gentle warming.

Distinction from arsenate wherein the yellow, crystalline precipitate is formed upon boiling, and from arsenite, where no precipitate forms (Svehla, 1996, p.95).

a. soluble in 6 N ammonium hydroxide. This

precipitate may be slow to form. (NH4)3[P(Mo3O10)4] + NH4OH dissolution of precipitate

Dissolution of precipitate forming a solution (USP CI, 2006, p.141)

Pyro-phosphate P2O7

2-

(1)With silver nitrate TS, pyrophosphates obtained by

ignition yield a white precipitate that is Na2P2O7 + AgNO3 Ag4P2O7 + NaNO3

White precipitate (USP CI, 2006, p.141)


a. soluble in 2 N nitric acid and Ag4P2O7 + 4 HNO3 H4P2O7 + 4 AgNO3 Dissolution of precipitate forming a solution (USP CI, 2006, p.141)

Double displacement reaction

b. in 6 N ammonium hydroxide. Ag4P2O7 + 8 NH4OH [Ag(NH3)2]4P2O7 + 8 H2O Dissolution of precipitate forming a solution (USP CI, 2006, p.141)

Formation of soluble complex. !

(2) With ammonium molybdate TS, a yellow

precipitate that is

Yellow precipitate (USP CI, 2006,

p.141)


a. soluble in 6 N ammonium hydroxide is formed. Dissolution of precipitate forming a solution (USP CI, 2006, p.141)

Salicylate*

(1) In moderately dilute solutions of salicylates, ferric chloride TS produces a violet color.

Violet-red solution (Svehla, 1996, p.241)

(2) The addition of acids to moderately concentrated

solutions of salicylates produces a white, crystalline precipitate of salicylic acid that melts between 158° and 161°C.


Hypo-phosphite*

(1) When strongly heated, hypophosphites evolve

spontaneously flammable phosphine.

4 NaH2PO2 Na4P2O7 + 2 PH3 + H2O (Svehla, 1996, p.224)

Colorless, extremely toxic, smells slightly of garlic or bad fish, highly reactive (Lee, 1997, p.481)

Disproportionation reaction

(2) Hypophosphites in solution yield a white precipitate with mercuric chloride TS.

HPH2O2 + 4 HgCl2 + 2 H2O 2 Hg2Cl2 + H3PO4 + 4 HCl (Svehla, 1996, p.224)


Redox

This precipitate becomes gray when an excess of hypophosphite is present.

2 Hg2Cl2 + HPH2O2 + 2 H2O 4 Hg + H3PO4 + 4 HCl (Svehla, 1996, p.224)

White precipitate turned gray (Svehla, 1996, p.224)

Mercurous salts are easily reduced to free mercury (Soine and Wilson, 1967, p.138)

(3) Solution of hypophosphites, acidified with sulfuric acid, and warmed with cupric sulfate TS yield a red

precipitate.

Δ

2 CuSO4 + 2 HPH2O2 + 4 H2O 2 H3PO4 + Cu2H2 + 2

H2SO4 + H2 (Soine and Wilson, 1967, p.138)

Red precipitate; bubbling (Soine and Wilson, 1967, p.138)

Redox.

Δ

4 CuSO4 + 3 HPH2O2 + 6 H2O 3 H3PO4 + 4 CuH + 4 H2SO4 (Svehla, 1996, p.224)

No precipitate in the cold, but on warming red precipitate is formed (Svehla, 1996, p.224)

Redox.

Benzoate*

(1) In neutral solutions, benzoates yield a salmon-colored precipitate with ferric chloride TS

3 C6H5COOK + 2 FeCl3 + 3 H2O (C6H5COO)3Fe

Fe(OH)3 + 3 KCl + 3HCl (Svehla, 1996, p.242)

Buff (pale yellowish beige)-colored precipitate (Svehla, 1996, p.242) or salmon-colored precipitate (USP CI, 2006, p.139)

(2) In moderately concentrated solutions, benzoates yield a precipitate of benzoic acid upon acidification with 2 N sulfuric acid. This precipitate is readily

soluble in ethyl ether.

2 C6H5COOK + H2SO4 2 C6H5COOH + K2SO4 (Svehla, 1996, p.242)


Substitution reaction.

Succinate*

No official ID tests

Arsenate1

Chromate2

Silicate*

Peroxo-

disulphate*

Phosphite*

Arsenite1

Dichromate

Hexa-flurosilicate*

Class B (2) Oxidation/ reduction in solution

Permanganate*

Manganate*

No official ID test Chromate2

Dichromate*

Other unclassified common anions


Acetate – Carbonate (acids) p. 139 ; Carbonate (phenolphthalein) – Phosphate (o-, yellow ppt) p. 140 ; Phosphate (soluble) – Zinc p.141

References: GENNARO, A. R., 2000. Remington: The Science and Practice of Pharmacy. 20

th ed. Philadelphia: Lippincott Williams & Wilkins, 402, 1020, 1032, 1212.

GILREATH, E. S., 1954. Qualitative Analysis Using Semimicro Methods. Quezon City: Ken, Inc., 96, 186-7, 194, 198-9, 210, 217, 225, 230, 239.

HOLZCLAW, H. F., Jr., and W. R. ROBINSON, 1988. College Chemistry With Qualitative Analysis. 8th ed. Massachusetts: D. C. Heath and Company, 375, 667, 812, 819, 833, 842, 931, 941, 943, 945, 962, 994.

LEE, J. D., 1996. Concise Inorganic Chemistry. 5

th ed. London: Chapman & Hall, 287-8, 382, 481, 501, 728, 825, 827, 846.

MOST, C. F., Jr., 1988. Experimental Organic Chemistry. New York: John Wiley & Sons, Inc., 519.

PARKS, L. M., JANNKE, P. J., and L. E. HARRIS, 1949. Inorganic Chemistry in Pharmacy. Philadelphia: J. B. Lippincott Compnay, 57, 138, 144, 149, 183, 209, 941. PETRUCCI, R. H., and W. S. HARWOOD, 1997. General Chemistry: Principles and Modern Applications. 7

th ed. New Jersey: Prentice-Hall, Inc., 766.

SHRINER, R. L., HERMANN, C. K. F., MORILL, T. C., CURTIN, D. Y., and, R. C. FUSON, 1998. The Systematic Identification of Organic Compounds. 7

th ed. New York: John Wiley & Sons, Inc., 88.

SOINE, T. O., and C. O. WILSON, 1967. Rogers’ Inorganic Pharmaceutical Chemistry. 8th ed. Philadelphia: Lea & Febiger, 62, 76, 82, 96, 128, 138, 151, 157, 193, 210, 220, 238, 243, 250, 286-7, 302, 318-20, 329, 355-6, 375, 407-8, 424-7,

451, 515, 545, 547, 551, 575, 605-6, 633, 672.

SVEHLA, G., 1996. Vogel’s Qualitative Inorganic Analysis. 7th ed. England: Longman Group Limited, 7, 64-5, 68, 70, 72-3, 78, 80, 84, 95, 98, 100, 110-1, 113-6, 118-20, 125, 129-30, 137, 140, 145, 147, 150, 153, 156, 163-4, 168, 171, 173-4,

177, 184, 191, 193-4, 196, 199, 200, 202-4, 209, 219-20, 224-6, 230-1, 234-5, 241-242, 311. UNITED STATES PHARMACOPEIAL CONVENTION, 2006. United States Pharmacopeia 30

th Revision and the National Formulary 25

th Edition. Asian Edition. Maryland: USPC, 139-141.

Lactate*

(1) When solutions of lactates are acidified with sulfuric acid, potassium permanganate TS is added, and the mixture is heated, acetaldehyde is

evolved.

Odor resembling that of bitter almond (Gennaro, 2000, p.1020)

This can be detected by allowing the vapor to come into contact with a filter paper that has been

moistened with a freshly prepared mixture of equal volumes of 20% aqueous morpholine and sodium

nitroferricyanide TS: a blue color is produced.

Paper turns blue (USP CI, 2006, p.140).

Peroxide*

(1) Solutions of peroxides slightly acidified with sulfuric acid yield a deep blue color upon the addition of potassium dichromate TS.

2 K2Cr2O7 + 2 H2SO4 + 2 H2O2 2 CrO5 + H2Cr2O7 + 2 K2SO4 + H2O (Svhela, 1996, p.226). H2SO4 H2Cr2O7 H2O + 2 CrO3 (Parks et al, 1949, p.138)

Evanescent blue color (Soine and Wilson, 1967, p.53) then a green solution (Svhela, 1996, p.226)

“Chromium pentoxide test”; (Svehla, 1996, p. 245). Dichromic acid, with more sulfuric acid, is dehydrated to chromium trioxide (Parks et al, 1949, p.138)

K2Cr2O7 + H2SO4 H2Cr2O7 + K2SO4

H2Cr2O7 + H2O2 HOCrO2OOCrO2OH + H2O (Soine and Wilson, 1967, p.53)

Perchromic acid (Soine and Wilson, 1967, p.53).

On shaking the mixture with an equal volume of ethyl ether and allowing the liquids to separate, the

blue color is found in the ethyl ether layer.

4 CrO5 + 6 H2SO4 2 Cr2(SO4)3 + 7 O2 + 6 H2O (Svehla, 1996, p.125)

Formation of beautiful blue in the upper organic layer (Svehla, 1996, p.245) and green lower aqueous layer, with bubbling (Svehla, 1996,

p.125).

Chromium pentoxide (chromium peroxide, peroxochromic acid) test; Decomposition of CrO5 (Svehla, 1996, p.125).

usp 191 identification tests-general_chemical equations_by_jude daval-santos

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