146
BUDAPEST UNIVERSITY OF TECHNOLOGY AND ECONOMICS FACULTY OF CHEMICAL ENGINEERING Dr. Tibor Pasinszki Inorganic Chemistry Laboratory Practice Department of Inorganic Chemistry 2002
BUDAPEST UNIVERSITY OF TECHNOLOGY AND ECONOMICSFACULTY OF CHEMICAL ENGINEERING
Dr. Tibor Pasinszki
Inorganic Chemistry Laboratory Practice
Department of Inorganic Chemistry2002
1
Contents
1. Introduction to the laboratory page 2
2. The Group Ia Elements (Li, Na, K, Rb, Cs) page 7 3. The Group IIa Elements (Be, Mg, Ca, Sr, Ba) page 12 4. The Group IIIa Elements (Boron and Aluminium) page 19 5. The Group IVa Elements (C, Si, Ge, Sn, Pb) page 24 6. The Group Va Elements (N, P, As, Sb, Bi) page 36 7. The Group VIa Elements (O, S, Se, Te) page 53 8. The Group VIIa Elements (F, Cl, Br, I) page 63 9. Pseudohalogens and Pseudohalides page 73 10. The Group Ib Elements (Cu, Ag, Au) page 78 11. The Group IIb Elements (Zn, Cd, Hg) page 87 12. The Group IVb Elements (Titanium) page 100 13. The Group Vb Elements (Vanadium) page 102 14. The Group VIb Elements (Chromium) page 105 15. The Group VIIb Elements (Manganese) page 111 16. The Group VIIIb Elements (Fe, Co, Ni) page 117 17. Classification of the Cations and Anions page 129 18. Testing for a Single Cation in Solution page 133 19. Testing for a Single Anion in Solution page 136 20. Separating and Identifying the Cations page 138
2
Introduction to the laboratory
SAFETY PROCEDURES
1. Chemical laboratory is a very dangerous workshop.Never work in the laboratory alone.
2. Do not eat, drink, or smoke in the laboratory. Most chemicals are poisonous.
3. Safety goggles will be worn in the laboratory any time there is laboratorywork in progress by any student.Remember that your neighbour could have any accident even though you,yourself, are not doing lab work.
4. If chemicals are spilled on the skin, immediately flush the skin with runningwater and call for the laboratory instructor. If chemicals are spilled on theclothes, remove them and flush the skin with water.
5. Never smell a reaction mixture directly. Minimise your exposure to chemicalvapours.
6. In order to avoid cuts and lacerations, protect your hands with a towel wheninserting either glass tubing or thermometers into stoppers or thermometeradapters. Fire-polish all glass tubing and stirring rods so that there are nosharp edges. Report any cuts to the lab instructor so that the injury mayreceive proper attention.
7. Restrict long hair in such a manner that it does not interfere with your work,become caught in the equipment, or catch fire.
8. Work with noxious chemicals in the hood. When in doubt, work in the hood,including rinsing equipment used in measuring such materials.Absorb escaping noxious gases in water or the suitable medium, or conductthe experiment in the hood.
9. Never heat an enclosed system.Never close completely an assembly from which a gas is being evolved.Have any equipment assembly checked by a lab instructor if this is thefirst time you have used the assembly.
10. Ordinary rubber stoppers are never used on flasks containing organic solvents.Organic solvents attack rubber and cause contamination of your product.
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11. Avoid fire. Most organic solvents are flammable. Play it safe and treat all organicmaterials as though they are flammable.NEVER heat an organic solvent over a Bunsen burner.Know the location of fire extinguishers, bucket of sand, safety showers, andfire blankets.Never attempt to extinguish an electrical fire with water. Use onlyextinguishers designated for this purpose.Report any fire regardless of how minor to the lab instructor.Report any burns to the lab instructor so that proper treatment may beadministered .
12. Avoid explosions.Never pour water into concentrated sulphuric acid. Always add concentratedsulphuric acid slowly to water.Never mix a strong oxidising agent with a strong reducing agent.Never mix nitric acid with alcohol.Never heat a flask to dryness when distilling or evaporating solvents. Smallamounts of impurities that can be explosive will be concentrated to dangerouslevels.
Always know what your neighbours are doing, be prepared for any accident.
4
REAGENTS IN THE LABORATORY
Pure chemical reagents or solutions of chemical reagents are stored in labelledbottles or dropping bottles in a convenient location in the laboratory. It is very importantto keep these reagents uncontaminated. Please obey the following rules in using thesereagents.
1. Read the labels carefully. Not only will the experiment be unsuccessful, but aserious accident may result if the wrong chemical is used.
2. Reagent bottles must be protected from contamination. You must thereforenever put spatulas, stirring rods, pipettes, or anything else into a reagent bottle.Try to avoid taking a large excess of the reagent. However, if you should errand take more than you need, do not return the excess to the bottle.In other words, you only remove material from the reagent bottle, younever put anything into it.
3. Never take the reagent bottles to the sink or to your desk. Put the bottlesback to the reagent shelf after using them.
4. Do not lay stoppers on the desk or shelf in such a way that they will becomecontaminated. Depending on the shape of the stopper, either hold it whilethe material is being removed or lay it on its flat top.
5. Glass stoppers that are stuck can generally be loosened by gently tappingthe stopper on the edge of the shelf.
6. The reagent area must be kept clean. Be sure that you clean up any chemicalsyou spill.
7. If you empty a container, take it to be refilled, as directed by your instructor.Do not return it to the reagent shelf empty.
8. Dispose the solid wastes in designated containers.Many kinds of liquid wastes must be collected and handled separately.Ordinarily acids, bases, and most inorganic liquid wastes can be flusheddown the sink with copious amounts of cold water. Check the directionsfor disposal of liquid wastes before using the sinks.
5
MOST COMMON LABORATORY EQUIPMENTS
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7
The Group Ia Elements (Li, Na, K, Rb, Cs) andTheir Principle Ions (Me+)
Lithium is the lightest of all metals, with a density only about half that of water (0.534g/ml at 0 °C). It has a low melting point of 180.5 °C. Lithium is silvery in appearance,much like Na and K.
Sodium is a soft, bright, silvery metal which floats on water (melting point: 97.8 °C). Itoxidises rapidly in moist air and is therefore kept under solvent naphtha or xylene.
Potassium is one of the most reactive and electropositive of metals. It is the secondlightest known metal, is soft, easily cut with a knife, and is silvery in appearanceimmediately after a fresh surface is exposed (melting point: 63.3 °C). Potassium israpidly oxidised in moist air, becoming covered with a blue film.
Rubidium is a soft, silvery-white metal (melting point: 38.9 °C), and is the second mostelectropositive and alkaline element.
Caesium is silvery with a golden-yellow appearance, soft, and ductile metal (meltingpoint: 28.4 °C). It is the most electropositive and most alkaline element.
Solubility in aqueous solutions
Alkali metals decompose water with the evolution of hydrogen and the formationof the hydroxide.
2 Me + 2 H2O → 2 MeOH + H2 ↑
Caesium reacts explosively with cold water. Rubidium ignites spontaneously in air andreacts violently in water, setting fire to the liberated hydrogen. Potassium catches firespontaneously on water. Sodium may or may not ignite spontaneously on water,depending on the amount of oxide and metal exposed to water. Lithium reacts with water,but not as vigorously as sodium.
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Flame test
Compounds of alkali metals (and also some of the others, see later) are volatilizedin the non-luminous Bunsen flame and impart characteristic colours to the flame, whichcan be used to identify the metal.
The physic-chemical process whichproduce the characteristic colour can besummarised as follows:
1. salt is evaporated2. molecule decompose to itsconstituents; e.g. NaCl → Na + Cl
3. thermal excitation of valence shellelectron of the metal atom;
Na → Na*4. instant relaxation of the excited states
with ejecting photons:Na* → Na + hν
If the photons ejected are in the visibleregion, coloration of the flame isobserved.
����
1s
2s
3s
4s
5s
6s
2p
3p
4p
5p
6p
3d
4d
5d 4f
������
������
���������
�������
����
Na (3p 3s)= 589 nm (yellow light����
Chlorides are among the most volatile compounds and readily decompose in the flame ofthe Bunsen burner, thus the best way to carry out the flame test is to prepare chlorides insitu by mixing the compound with a little concentrated hydrochloric acid before carryingout the tests.The procedure is as follows. Put a fewcrystals or a few drops of the solutioninto a porcelain crucible, addhydrochloric acid, and zinc chips. Thehydrogen gas liberated in the reactionbetween zinc and hydrochloric acidcarries fine drops of the solution into theflame, where the latter are volatilised.
The colours imparted to the flame:Lithium → carmine-red
Sodium → golden-yellowPotassium → violet (lilac)
Rubidium → dark-redCaesium → blue
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Characteristic reactions of lithium ions
The solubilities of lithium carbonate, Li2CO3, the phosphate, Li3PO4, and thefluoride, LiF, in water are little, definitely much less than the corresponding sodium andpotassium salts, and in this respect lithium resembles the alkaline earth metals. All otherimportant inorganic lithium salts are soluble in water.
For example: compound solubility ( g / 100 mlwater)
at 18 °C: LiFLi3PO4
0,270,039
0 °C: LiCl 63,7
To study these reactions use a 1 M solution of lithium chloride.
1. Sodium phosphate solution: partial precipitation of lithium phosphate, Li3PO4, inneutral solutions.
3 Li+ + PO43− → Li3PO4 ↓
Precipitation is almost complete in the presence of sodium hydroxide solution.
2. Sodium carbonate solution: white precipitate of lithium carbonate fromconcentrated solutions:
2 Li+ + CO32− → Li2CO3 ↓
3. Ammonium carbonate solution: white precipitate of lithium carbonate fromconcentrated solutions.
2 Li+ + CO32− → Li2CO3 ↓
No precipitation occurs in the presence of high concentration of ammonium chloridesince the carbonate ion concentration is reduced to such an extent that the solubilityproduct of lithium carbonate is not exceeded:
NH4+ + CO3
2− ↔ NH3 + HCO3−
4. Ammonium fluoride solution: a white, gelatinous precipitate of lithium fluoride isslowly formed in ammoniacal solution.
Li+ + F− → LiF ↓
5. Flame test: carmine-red colour.
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Sodium, Na+
Almost all sodium salts are soluble in water. There are, however, some specialreagents which a crystalline precipitate is formed with if it is added to a fairlyconcentrated solution of sodium salts.
To study these reactions use a 1 M solution of sodium chloride.
1. Uranyl magnesium acetate solution: yellow, crystalline precipitate of sodiummagnesium uranyl acetate NaMg(UO2)3(CH3COO)9.9H2O from concentratedsolutions:
Na+ + Mg2+ + 3 UO22+ + 9 CH3COO− → NaMg(UO2)3(CH3COO)9
2. Uranyl zinc acetate solution: yellow, crystalline precipitate of sodium zinc uranylacetate NaZn(UO2)3(CH3COO)9.9H2O :
Na+ + Zn2+ + 3 UO22+ + 9 CH3COO− → NaZn(UO2)3(CH3COO)9
3. Flame test: intense yellow colour.
Potassium, K+
Most of the potassium salts salts are soluble in water.To study the reactions which produce water insoluble or little soluble salts, use a
1 M solution of potassium chloride. Remember, the sizes of K+ and NH4+ ions are almost
identical, thus their reactions in general are very similar.
1. Perchloric acid solution (HClO4): white crystalline precipitate of potassiumperchlorate KClO4 from not too dilute solutions. You may use concentrated HClO4solution. (This reaction is unaffected by the presence of ammonium salts.)
K+ + ClO4− → KClO4 ↓
2. Tartaric acid solution (or sodium hydrogen tartrate solution): white crystallineprecipitate of potassium hydrogen tartrate:
K+ + H2C4H4O6 → KHC4H4O6 ↓ + H+
The solution should be buffered with sodium acetate. The precipitate is slightly soluble inwater (3.26 g/l). (Ammonium salts yield a similar precipitate.)
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3. Sodium hexanitritocobaltate(III) solution, Na3[Co(NO2)6]: yellow precipitate ofpotassium hexanitritocobaltate(III):
3 K+ + [Co(NO2)6]3− → K3[Co(NO2)6] ↓
The precipitate is insoluble in dilute acetic acid. In alkaline solutions a brown precipitateof cobalt(III) hydroxide is obtained. (Ammonium salts give a similar precipitate.)If larger amounts of sodium salts are present (e.g. reagent is added in excess) a mixed saltis formed:
2 K+ + Na+ + [Co(NO2)6]3− → K2Na[Co(NO2)6] ↓
The test is more sensitive if sodium hexanitritocobaltate(III) and silver nitrate solutionsare added together to halogen free solutions; the compound K2Ag[Co(NO2)6] forms,which is less soluble in water than the corresponding salt, K2Na[Co(NO2)6].
5. Flame test: violet colour.
Summarise the solubility of common inorganic salts of Li+, Na+, and K+:
CO32− PO4
3− F− Cl− NO3− SO4
2−
Li+
Na+
K+
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The Group IIa Elements (Be, Mg, Ca, Sr, Ba) andTheir Principle Ions (Me2+)
Beryllium is a steel grey, light but very hard, brittle metal, one of the lightest of allmetals, and has one of the highest melting points of the light metals (1278 °C). Berylliumobjects are oxidised on the surface, but the oxide layer protects the objects from furtheroxidisation, which is similar to that of aluminium. Beryllium resembles closelyaluminium in chemical properties; it also exhibits resemblances to the alkaline earthmetals.
Magnesium is a light, silvery-white, malleable and ductile metal with a melting point of649 °C. Magnesium objects have a protective oxide layer on the surface, similarly to thatof beryllium and aluminium. It burns upon heating in air or oxygen with a brilliant whitelight, forming the oxide and some nitride.
Calcium has a silvery colour, is rather soft, but definitely much harder than the alkalimetals (melting point: 839 °C). It is attacked by atmospheric oxygen and humidity, whencalcium oxide and/or calcium hydroxide is formed.
Strontium is a silvery-white, malleable and ductile metal (melting point: 769 °C).Strontium is softer than calcium and decomposes water more vigorously. It should bekept under kerosene to prevent oxidation. Freshly cut strontium has a silvery appearance,but rapidly turns a yellowish colour with the formation of the oxide
Barium is a silvery-white, soft, malleable and ductile metal (melting point: 725 °C). Itoxidises very easily and should be kept under petroleum to exclude air.
Solubility in water and acids Beryllium does not reacts with water at ordinary conditions.Magnesium is slowly decomposed by water at ordinary temperature, but at the
boiling point of water the reaction proceeds rapidly:
Mg + 2 H2O → Mg(OH)2 ↓ + H2 ↑
Calcium, strontium, and barium decompose water at room temperature with theevolution of hydrogen and the formation of the hydroxide.
Me + 2 H2O → Me2+ + 2 OH− + H2 ↑
Be, Mg, Ca, Sr, and Ba dissolve readily in dilute acids (unless water insoluble saltforms):
Me + 2 H+ → Me2+ + H2 ↑
Concentrated nitric acid renders beryllium passive (like aluminium).
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Characteristic reactions of magnesium ions, Mg2+
The magnesium oxide, hydroxide, carbonate, and phosphate are insoluble inwater; the other common inorganic salts are soluble.
To study these reactions use a 0.1 M solution of magnesium chloride or sulphate.
1. Ammonium carbonate solution: in the absence of other ammonium salts a whiteprecipitate of basic magnesium carbonate:
5 Mg2+ + 6 CO32− + 7 H2O → 4 MgCO3.Mg(OH)2.5 H2O ↓ + 2 HCO3
−
In the presence of ammonium salts no precipitation occurs, because the followingequilibrium is shifted towards the formation of hydrogen carbonate ions (remember,magnesium hydrogen carbonate is soluble in water):
NH4+ + CO3
2− ↔ NH3 + HCO3−
2. Sodium carbonate solution: white, voluminous precipitate of basic magnesiumcarbonate:
5 Mg2+ + 6 CO32− + 7 H2O → 4 MgCO3.Mg(OH)2.5 H2O ↓ + 2 HCO3
−
3. Ammonium hydroxide solution: partial precipitation of white, gelatinousmagnesium hydroxide, solubility product constant: Ksp(25°C)= 5.61x10−12:
Mg2+ + 2 NH3 + 2 H2O → Mg(OH)2 ↓ + 2 NH4+
The precipitate is readily soluble in solutions of ammonium salts.
4. Sodium hydroxide solution: white precipitate of magnesium hydroxide:
Mg2+ + 2 OH− → Mg(OH)2 ↓
5. Disodium hydrogen phosphate solution: a white flocculant precipitate ofmagnesium hydrogen phosphate is produced in neutral solutions:
Mg2+ + HPO42− → MgHPO4 ↓
White crystalline precipitate of magnesium ammonium phosphate MgNH4PO4.6H2O inthe presence of ammonium chloride (to prevent precipitation of magnesium hydroxide)and ammonia solutions:
Mg2+ + NH3 + HPO42− → MgNH4PO4 ↓
The precipitate is soluble in acetic acid and in mineral acids.
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6. Titan yellow reagent and magneson reagent:
Titan yellow and magneson (I and II)are water soluble dyestuffs. They areabsorbed by magnesium hydroxideproducing a deep-red colour with titanyellow and a blue colour with magneson.
Pour a little amount of the test solutioninto two test tubes, add 1-2 drops of thetitan yellow reagent to one test tube and1-2 drops of the magneson reagent to theother test tube. Render the solutions inboth test tubes alkaline with sodiumhydroxide solutions.
O2N N N OH
HO
O2N N N OH
Magneson I
Magneson II
Calcium, Ca2+
Calcium chloride and nitrate are readily soluble in water.Calcium oxide (similarly to strontium and barium oxides) readily reacts with
water producing heat and forming the hydroxide.Calcium sulphide (and also other alkaline earth sulphides) can be prepared only in
the dry; it hydrolyses in water forming hydrogen sulphide and hydroxide:2 CaS + 2 H2O → 2 Ca2+ + 2 SH- + 2 OH-
Calcium carbonate, sulphate, phosphate, and oxalate are insoluble in water.
To study the reactions of Ca2+ ions, use a 0.1 M solution of calcium chloride.
1. Ammonium carbonate solution: white amorphous precipitate of calcium carbonate,solubility product constant: Ksp(25°C)= 4.96x10−9, (the precipitate is soluble in acidseven in acetic acid):
Ca2+ + CO32− → CaCO3 ↓
2. Dilute sulphuric acid: white precipitate of calcium sulphate, solubility productconstant: Ksp(25°C)= 7.10x10−5:
Ca2+ + SO42− → CaSO4 ↓
3. Ammonium oxalate solution: white precipitate of calcium oxalate, solubilityproduct constant: Ksp(CaC2O4.H2O, 25°C)= 2.34x10−9 (insoluble in acetic acid, butsoluble in mineral acids):
Ca2+ + (COO)22− → Ca(COO)2 ↓
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4. Disodium hydrogen phosphate solution: white precipitate of calcium hydrogenphosphate is produced from neutral solutions:
Ca2+ + HPO42− → CaHPO4 ↓
5. Potassium hexacyanoferrate(II) solution: white precipitate of a mixed salt:
Ca2+ + 2 K+ + [Fe(CN)6]4− → K2Ca[Fe(CN)6] ↓
In the presence of ammonium chloride the test is more sensitive. In this case potassium isreplaced by ammonium ions in the precipitate. The test can be used to distinguishcalcium from strontium; barium and magnesium ions however interfere.
6. Flame test: yellowish-red colour to the Bunsen flame.
Strontium, Sr2+
Strontium chloride and nitrate are readily soluble in water.
Strontium carbonate, sulphate, phosphate, and oxalate are insoluble in water.
To study the reactions of Sr2+ ions, use a 0.1 M solution of strontium chloride orstrontium nitrate.
1. Ammonium carbonate solution: white precipitate of strontium carbonate,Ksp(SrCO3, 25°C)= 5.60x10−10 (the precipitate is soluble in acids even in acetic acid):
Sr2+ + CO32− → SrCO3 ↓
2. Dilute sulphuric acid: white precipitate of strontium sulphate, Ksp(SrSO4, 25°C)=3.44x10−7:
Sr2+ + SO42− → SrSO4 ↓
3. Saturated calcium sulphate solution: white precipitate of strontium sulphate,formed slowly in the cold, but more rapidly on boiling.
Sr2+ + SO42− → SrSO4 ↓
4. Ammonium oxalate solution: white precipitate of strontium oxalate:
Sr2+ + (COO)22− → Sr(COO)2 ↓
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5. Disodium hydrogen phosphate solution: white precipitate of strontium hydrogenphosphate is produced from neutral solutions:
Sr2+ + HPO42− → SrHPO4 ↓
6. Potassium chromate solution: yellow precipitate of strontium chromate:
Sr2+ + CrO42− → SrCrO4 ↓
The precipitate is appreciably soluble in water, thus no precipitate occurs in dilutesolutions of strontium ions.The precipitate is soluble in acetic acid and in mineral acids.
The addition of acid to potassium chromate solution causes the yellow colour ofthe solution to change to reddish-orange, owing to the formation of dichromate. Theaddition of acetic acid or mineral acid to the potassium chromate solution lowers theCrO4
2− ion concentration sufficiently to prevent the precipitation of SrCrO4.
The equilibria are the following:
H2CrO4 ↔ HCrO4− + H+
HCrO4− ↔ CrO4
2− + H+
2 HCrO4− ↔ H2O + Cr2O7
2−
HCr2O7− ↔ Cr2O7
2− + H+
0 2 4 6 8 10
0.00
0.02
0.04
0.06
0.08
0.10c= 0.1 M (K2CrO4)
[CrO
42-] (mol/l)
pH
7. Flame test: crimson-red colour to the Bunsen flame.
Barium, Ba2+
Barium chloride and nitrate are readily soluble in water.Barium carbonate, sulphate, phosphate, and oxalate are insoluble in water.
To study the reactions of Ba2+ ions, use a 0.1 M solution of barium chloride orbarium nitrate.
1. Ammonium carbonate solution: white precipitate of barium carbonate, Ksp(BaCO3,25°C)= 2.58x10−9 (the precipitate is soluble in acids even in acetic acid):
Ba2+ + CO32− → BaCO3 ↓
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2. Dilute sulphuric acid: white, finely divided precipitate of barium sulphate,Ksp(BaSO4, 25°C)= 1.07x10−10:
Ba2+ + SO42− → BaSO4 ↓
3. Saturated calcium sulphate solution: immediate white precipitate of bariumsulphate.
4. Saturated strontium sulphate solution: white precipitate of barium sulphate.
5. Ammonium oxalate solution: white precipitate of barium oxalate (readily dissolvedby hot dilute acetic acid and by mineral acids):
Ba2+ + (COO)22− → Ba(COO)2 ↓
6. Disodium hydrogen phosphate solution: white precipitate of barium hydrogenphosphate is produced from neutral solutions:
Ba2+ + HPO42− → BaHPO4 ↓
6. Potassium chromate solution: yellow precipitate of barium chromate, practicallyinsoluble in water, Ksp(BaCrO4, 25°C)= 1.17x10−10:
Ba2+ + CrO42− → BaCrO4 ↓
The precipitate is insoluble in dilute acetic acid (distinction from strontium), but solublein mineral acids.
7. Flame test: yellowish-green colour to the Bunsen flame.
Compare the solubility product constants of CaSO4, SrSO4, and BaSO4, andcalculate the sulphate ion concentration in saturated solutions.
CaSO4 SrSO4 SrSO4
solubility productconstant: Ksp
7.10x10−5 3.44x10−7 1.07x10−10
SO42−
concentration
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Summarise the reactions of Mg2+, Ca2+, Sr2+, and Ba2+ ions:
Mg2+ Ca2+ Sr2+ Ba2+
NH3 soln.
NaOH
Na2CO3
(NH4)2CO3+ NH4Cl
(NH4)2CO3
Na2HPO4
(NH4)2(COO)2add hot acetic acidto the precipitate
K2CrO4neutral soln. add acetic acid to
the precipitate
1. add acetic acidto the precipitate2. add mineral acidto the prec.
dilute H2SO4
satd. CaSO4 soln.
satd. SrSO4 soln.
Flame test
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The Group IIIa Elements: Boron and Aluminium, andTheir Principle Ions (B(OH)4− and Al3+)
Boron has properties that place it on the borderline between metals and nonmetals, butchemically it must be classed as a nonmetal. Boron is a hard, steel-grey solid with a highmelting point of 2079 °C. Crystalline boron is extremely inert chemically.
Aluminium. Pure aluminium is a silvery-white metal (melting point (m.p.): 660.4 °C). Itis light, malleable and ductile, can easily be formed, machined, or cast, has a highthermal conductivity, and has an excellent corrosion resistance. The aluminium powder isgrey. Exposed to air, aluminium objects are oxidised on the surface, but the oxide layerprotects the objects from further oxidisation.
Solubility in aqueous acids and alkaliBoron is unaffected by nonoxidising acids (e.g. boiling HCl or HF).
It is only slowly oxidised by hot, concentrated nitric acid, and also only slowly attackedby other hot concentrated oxidising agents (e.g. aqua regia, or a mixture of concentratednitric acid and hydrogen fluoride).Boron is soluble in alkali with the evolution of hydrogen gas.
B + HNO3 + H2O → H3BO3 + NO ↑2 B + 2 HNO3 + 4 H2F2 → 2 H[BF4] + 2 NO ↑ + 4 H2O2 B + 2 NaOH + 6 H2O → 2 Na+ + 2 B(OH)4
− + 3 H2 ↑
Aluminium is soluble in dilute or concentrated hydrochloric acid with theliberation of hydrogen:
2 Al + 6 HCl → 2 Al3+ + 6 Cl- + 3 H2 ↑
Dilute sulphuric acid dissolves the metal with the liberation of hydrogen andconcentrated sulphuric acid with the liberation of sulphur dioxide:
2 Al + 3 H2SO4 → 2 Al3+ + 3 SO42− + 3 H2 ↑2 Al + 6 H2SO4 → 2 Al3+ + 3 SO42− + 3 SO2 ↑ + 6 H2O
Concentrated nitric acid renders aluminium passive, but dilute nitric acid dissolves themetal:
Al + 4 HNO3 → Al3+ + 3 NO3− + NO ↑ + 2 H2O
Aluminium is soluble in alkali hydroxides when a solution of tetrahydroxoaluminate isformed:
2 Al + 2 OH- + 6 H2O → 2 [Al(OH)4]− + 3 H2 ↑
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Reducing power of aluminiumAluminium is a very reactive metal, in particular towards electronegative
partners, but this extreme reactivity can only be observed when the stable oxide layer atthe metal surface is destroyed or the metal is in a finely divided form.
1. Termite reaction: no reaction occurs between iron(III) oxide and aluminium powderat room temperature, but an exothermic, violent reaction takes place when it is initiatedby a thermal ignition mixture.
Fe2O3 + 2 Al → Al2O3 + 2 Fe
The reaction is extremely violent and is accompanied by the formation of a large amountof sparks.
2. Reaction with iodine: add only one drop of water to a mixture of fine aluminiumpowder and powdered iodine.
2 Al + 3 I2 → Al2I6
The reaction is induced by water, the heat which is set free at the beginning of thereaction sufficing to convert the whole mixture to dialuminium hexaiodide and tosublime the excess iodine.
Aluminium(III) ions, Al3+
Solubility: aluminium chloride, bromide, iodide, nitrate, and sulphate are solublein water.
Aluminium fluoride is hardly soluble in water. Aluminium oxide, hydroxide,phosphate, and carbonate are practically insoluble in water.
For example: compound solubility ( g / 100 mlwater)
at 15 °C AlCl3 69,925 °C AlF3
α-Al2O3AlPO4
0,5590,000098
-----
Aluminium sulphide can be prepared in the dry state only, in aqueous solutions ithydrolyses and aluminium hydroxide is formed.
Use a 0.1 M solution of aluminium chloride or sulphate to study the reactions ofaluminium(III) ions.
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1. Ammonium sulphide solution: a white precipitate of aluminium hydroxide.
2 Al3+ + 3 S2− + 6 H2O → 2 Al(OH)3 ↓ + 3 H2S ↑
2. Sodium hydroxide solution: white precipitate of aluminium hydroxide. Theprecipitate dissolves in excess reagent, when tetrahydroxoaluminate ions are formed.
Al3+ + 3 OH− → Al(OH)3 ↓Al(OH)3 ↓ + OH− → [Al(OH)4]−
The reaction is a reversible one, and any reagent which will reduce the OH- ionconcentration sufficiently should cause the reaction to proceed from right to left.
3. Ammonia solution: white gelatinous precipitate of aluminium hydroxide. Theprecipitate is only slightly soluble in excess of the reagent, the solubility is decreased inthe presence of ammonium salts.
Al3+ + 3 NH3 + 3 H2O → Al(OH)3 ↓ + 3 NH4+
4. Sodium phosphate solution: white gelatinous precipitate of aluminium phosphate,solubility product constant: Ksp(AlPO4, 25°C)= 9.83x10−21:
Al3+ + HPO42− → AlPO4 ↓ + H+
Strong acids and also sodium hydroxide dissolve the precipitate.
5. Sodium acetate solution: no precipitate is obtained in cold, neutral solutions, but onboiling with excess reagent, a voluminous precipitate of basic aluminium acetate isformed:
Al3+ + 3 CH3COO- + 2 H2O → Al(OH)2CH3COO ↓ + 2 CH3COOH
6. Sodium alizarin sulphonate (Alizarin-S) reagent: red precipitate in ammoniacal solution, which isfairly stable to dilute acetic acid.Add to the solution of Al3+ ions, dilute ammoniasolution and 2-3 drops of the solution of the reagent, andthen acidify it with acetic acid.
7. Morin reagent: add little solid sodium acetate and 1-2 drops of the reagent to the solution of Al3+ ions.Investigate the characteristic green fluorescence of thesolution in UV light.
C
C
O
O
OH
OH
SO Na3
O
OOH
OHHO
HO
OH Morin
Alizarin-S
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Oxides of Boron and Aluminium
B2O3
B2O3 is a white , hygroscopic solid.It is acidic, reacting with water to giveboric acid, B(OH)3.
Al2O3
α−Al2O3 is very hard and resistant tohydration and attack by acids.γ−Al2O3 readily takes up water anddissolves in acids.
Boric acid and borate ions in aqueous solutionBoric acid, H3BO3 or B(OH)3, is a very weak and exclusively monobasic acid that
acts not as a proton donor, but as a Lewis acid, accepting OH−:
B(OH)3 + H2O ↔ B(OH)4− + H+ KS= 6x10−10
In aqueous, concentrated borate solutions polymeric ions are also present, due to thepolymerisation between B(OH)3 and B(OH)4
−, the most important ions for example:
4 B(OH)3 + B(OH)4− ↔ B5O6(OH)4
− + 6 H2O2 B(OH)3 + B(OH)4
− ↔ B3O3(OH)4− + 3 H2O
2 B(OH)3 + 2 B(OH)4− ↔ B4O5(OH)4
2− + 5 H2O
In acidic solution (pH<4) orthoboric acid B(OH)3, in basic solution (pH>12) B(OH)4−
ions exist exclusively, and at medium pH (4<pH<12) besides B(OH)4−, polyanions
B5O6(OH)4−, B3O3(OH)4
−, and B4O5(OH)42− are also present.
The species B5O6(OH)4−, B3O3(OH)4
−, and B4O5(OH)42− are formed successively with
increasing pH.
In dilute solutions depolimerization rapidly occurs; at concentrations <0.025 M,essentially only mononuclear species B(OH)3 and B(OH)4
− are present.
Borates, BO33−, B4O72−, BO2−
The borates are formally derived from the three boric acids:orthoboric acid, H3BO3 (a well known white, crystalline solid),metaboric acid, HBO2 (not known in solution and can not be isolated) andpyroboric acid, H2B4O7 (not known in solution and can not be isolated).Most of the salts are derived from the meta- and pyroboric acids, and only very few saltsof orthoboric acid are known.
Solubility: the borates of the alkali metals are readily soluble in water.The borates of the other metals are, in general, sparingly soluble in water, but fairlysoluble in acids and in ammonium chloride solution.
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The soluble salts are hydrolysed in solution, owing to the weakness of boric acid, andtherefore react alkaline:
BO33− + 3 H2O ↔ H3BO3 + 3 OH−B4O72− + 7 H2O ↔ 4 H3BO3 + 2 OH−BO2− + 2 H2O ↔ H3BO3 + OH−
To study the reactions of borates use a 0.1 M solution of sodium tetraborate(sodium pyroborate, borax) Na2B4O7.10H2O.
1. Barium chloride solution: white precipitate of barium metaborate from fairlyconcentrated solutions:
B4O72− + 2 Ba2+ + H2O → 2 Ba(BO2)2 ↓ + 2 H+
The precipitate is soluble in excess reagent, in dilute acids, and in solutions ofammonium salts.
2. Silver nitrate solution: white precipitate of silver metaborate from fairly concentratedsolution:
B4O72− + 4 Ag+ + H2O → 4 AgBO2 ↓ + 2 H+
The precipitate is soluble in both dilute ammonia solution and in acetic acid. On boilingthe precipitate with water, it is completely hydrolysed and a brown precipitate of silveroxide is obtained.
AgBO2 ↓ + 2 NH3 + 2 H2O → [Ag(NH3)2]+ + B(OH)4−AgBO2 ↓ + H+ + H2O → Ag+ + H3BO32 AgBO2 ↓ + 3 H2O → Ag2O + 2 H3BO3
3. Hydrochloric acid: there is no visible change with dilute hydrochloric acid, but ifconcentrated hydrochloric acid is added to a concentrated solution of borax, boric acid isprecipitated:
B4O72− + 2 HCl + 5 H2O → 4 H3BO3 ↓ + 2 Cl−
4. Concentrated sulphuric acid and alcohol (flame test)
If a little borax is mixed with 1ml concentrated sulphuric acid and 5ml methanol in a small porcelainbasin, and the alcohol ignited, thelatter will burn with a green-edgedflame due to the formation of methylborate B(OCH3)3:
B4O72- + H2SO4 + 5 H2O → 4 H3BO3 + SO42-H3BO3 + 3 CH3OH → B(OCH3)3 ↑ + 3 H2O
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The Group IVa Elements (C, Si, Ge, Sn, Pb) andTheir Principle Ions
Carbon has three allotropic forms: diamond, graphite, and fullerenes.Diamond is the hardest solid known. It has a high density and the highest melting point (∼4000 °C) of any element. The chemical reactivity of diamond is much lower than that ofcarbon in the form of macrocrystalline graphite or the various amorphous forms.Diamond can be made to burn in air by heating it to 600 to 800 °C.Graphite has a layer structure and the forces between layers are relatively slight. Thus theobserved softness and particularly the lubricity of graphite can be attributed to the easyslippage of these layers over one another.Fullerenes belongs to the family of carbon-cage molecules, discovered during the lasttwo decades of the XXth century, of which C60 and C70 are the most known members.Both C60 and C70 are highly coloured crystalline solids that are sparingly soluble incommon organic solvents.
Silicon has a solid structure which is isostructural with diamond. Crystalline silicon hasa metallic lustre and greyish colour. Melting point (m.p.): 1410 °C.
Germanium is isostructural with diamond. It is a grey-white metalloid, and in its purestate is crystalline and brittle, retaining its lustre in air at room temperature. (m.p.: 937.4 °C)
Tin has two crystalline modifications: α-tin or grey tin, and β or white tin (metallicform). α-tin has the diamond structure.Tin (β-form) is a silver-white metal which is malleable and ductile at ordinarytemperatures, but at low temperatures (below 13.2 °C) it becomes brittle due totransformation into the α allotropic modification. Tin melts at 232 °C.
Lead exists only in a metallic form. It is a bluish-grey metal with a high density, is verysoft, highly malleable, and ductile. Melting point: 327.5 °C. Lead is very resistant tocorrosion.
Solubility of group IVa elements in aqueous acids and alkaliCarbon is very unreactive at normal conditions, but the reactivity of IVa group
elements is increasing down the group, from the carbon toward the lead.Carbon is not soluble in aqueous acids or alkalis.Silicon is rather unreactive. It is not attacked by acids except the mixture of HF
and HNO3; presumably the stability of SiF62- provides the driving force here. Silicon issoluble in alkalis giving solutions of silicates.
3 Si + 18 HF + 4 HNO3 → 3 H2SiF6 + 4 NO + 8 H2OSi + 2 KOH + H2O → K2SiO3 + 2 H2
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Germanium is somewhat more reactive than silicon, and dissolves in concentratedH2SO4 and HNO3, when GeO2.xH2O is formed. It is not attacked by alkalis and nonoxidising acids, soluble, however, in alkalis containing hydrogen peroxide.
3 Ge + 4 HNO3 + (x-2) H2O → 3 GeO2.xH2O ↓ + 4 NO ↑
Tin and lead dissolve in several acids. Tin is attacked slowly by cold alkali,rapidly by hot, lead only by hot, to form stannates and plumbites.
Sn + 2 NaOH + 2 H2O → Na2[Sn(OH)4] + H2 ↑
Tin dissolves slowly in dilute hydrochloric acid and sulphuric acid with theformation of tin(II) salts:
Sn + 2 H+ → Sn2+ + H2 ↑
Dilute nitric acid dissolves tin slowly without the evolution of any gas, tin(II) andammonium ions being formed:
4 Sn + 10 H+ + NO3− → 4 Sn2+ + NH4+ + 3 H2O
In hot, concentrated sulphuric acid and in aqua regia tin(IV) ions are formed atdissolution:
Sn + 4 H2SO4 → Sn4+ + 2 SO42− + 2 SO2 ↑ + 4 H2O3 Sn + 4 HNO3 + 12 HCl → 3 Sn4+ + 12 Cl− + 4 NO ↑ + 8 H2O
Tin reacts vigorously with concentrated nitric acid, and a white solid, usually formulatedas hydrated tin(IV) oxide SnO2.xH2O and also known as metastannic acid, is produced:
3 Sn + 4 HNO3 + (x-2) H2O → 3 SnO2.xH2O ↓ + 4 NO ↑
Lead readily dissolves in medium concentrated (8M) nitric acid with theformation of nitrogen oxide. The colourless nitrogen oxide gas, when mixed with air, isoxidised to red nitrogen dioxide:
3 Pb + 8 HNO3 → 3 Pb2+ + 6 NO3- + 2 NO ↑ + 4 H2O2 NO ↑ (colourless) + O2 ↑ → NO2 ↑ (reddish-brown)
With concentrated nitric acid a protective film of lead nitrate is formed on the surface ofthe metal and prevents further dissolution.Dilute hydrochloric or sulphuric acid have little effect owing to the formation of aprotective film of insoluble lead chloride or sulphate on the surface.
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Principle oxides of IVa group elements C Si Ge Sn Pb
CO, CO2colourless gases
SiO2white solid
GeO2white solid
SnO red(β)*
SnO2 whitePbO yellow (red)*
PbO2 black* white SnO.xH2O and white PbO.xH2O precipitates from aqueous solutions
Only CO2 is soluble in water; the solubility strongly depends on the pressure andtemperature. (1 litre water dissolves 0.9 litre of CO2 of 1 atm pressure at 20 °C.)SiO2 and GeO2 are hardly soluble in water; e.g. the solubility of GeO2 is 0.4 g in 100 gwater at 20 °C.SnO2 and PbO2 are insoluble in water.
CO2 is acidic, and the acidic character of oxides of the IVa group elements decreasesform the carbon dioxide toward the lead oxide. Carbon, silicon, and germanium oxidesare acidic, tin oxides are amphoteric, and lead oxide has also some basic character.
CO2, SiO2, and GeO2 are not soluble in acids, but soluble in alkalis giving carbonates,silicates, and germanates, respectively. Silicate and germanate anions are polymeric.SnO2 is not soluble in acids and alkalis, and PbO2 is only little soluble in acids.
SnO is soluble in acids and alkalis, forming tin(II) salts or stannates.PbO is soluble in acids, forming lead(II) salts.
Lead(IV) oxide, PbO2, is a strong oxidising agent (Pb2+/ PbO2= +1.455 V), thus itliberates chlorine by boiling with concentrated hydrochloric acid:
PbO2 + 4 HCl → PbCl2 + 2 H2O + Cl2 ↑
Principal ions of IVa group elements and their characteristic reactions C Si Ge Sn Pb
CO32−
HCO3−SiO32− * GeO32− * Sn2+
Sn4+Pb2+
* Does not exist in this form in aqueous solution.
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Carbonates, CO32−
All normal carbonates, with the exception of those of the alkali metals and ofammonium, are insoluble in water.
The hydrogen carbonates (or bicarbonates) of the alkali metals are soluble inwater, but are less soluble than the corresponding normal carbonates.The hydrogen carbonates of calcium, strontium , barium, magnesium, and possibly ofiron exist in aqueous solution; they are formed bye the action of excess carbonic acidupon the normal carbonates either in aqueous solution or suspension:
CaCO3 ↓ + H2O + CO2 → Ca2+ + 2 HCO3−
Hydrogen carbonates are decomposed to carbonates on boiling the solution.The following equilibria exists in aqueous solution:
CO2 + 3 H2O ↔ H2CO3 + 2 H2O ↔ H3O+ + HCO3- + H2O ↔ 2 H3O+ + CO32-
In acid solutions the equilibria shifted towards the left, while in alkaline medium they areshifted towards the right.
To study the reactions of carbonates, use a 0.5 M solution of sodium carbonate,Na2CO3.10H2O.
1. Dilute hydrochloric acid: decomposition with the evolution of carbon dioxide:
CO32- + 2 H+ → CO2 ↑ + H2O
the gas can be identified by its property ofrendering lime water or baryta water turbid:
CO2 + Ca2+ + 2 OH- → CaCO3 ↓ + H2OCO2 + Ba2+ + 2 OH- → BaCO3 ↓ + H2O
Any acid which is stronger than carbonic acidwill displace it, especially on warming. Thuseven acetic acid will decompose carbonates; theweak boric acid and hydrocyanic acid will not.
2. Barium chloride (or calcium chloride) solution: white precipitate of barium (orcalcium) carbonate:
CO32- + Ca2+ → CaCO3 ↓CO32- + Ba2+ → BaCO3 ↓
Only normal carbonates react; hydrogen carbonates do not. The precipitate is soluble inmineral acids and carbonic acid.
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3. Silver nitrate solution: white precipitate of silver carbonate, solubility productconstant Ksp(Ag2CO3, 25 °C)= 8.45x10−12:
CO32- + 2 Ag+ → Ag2CO3 ↓
The precipitate is soluble in nitric acid and in ammonia. The precipitate becomes yellowor brown upon addition of excess reagent owing to the formation of silver oxide; thesame happens if the mixture is boiled:
Ag2CO3 ↓ → Ag2O ↓ + CO2
Hydrogen carbonates, HCO3−
Most of the reactions of hydrogen carbonates are similar to those of carbonates.The tests described here are suitable to distinguish hydrogen carbonates from carbonates.
To study the reactions of hydrogen carbonates, use a freshly prepared 0.5 M solution ofsodium hydrogen carbonate, NaHCO3.
1. Boiling. When boiling, hydrogen carbonates decompose:
2 HCO3- → CO32- + H2O + CO2 ↑
carbon dioxide, formed in this way, can be identified with lime water or baryta water.
2. Magnesium sulphate. Adding magnesium sulphate to a cold solution of hydrogencarbonate no precipitation occurs, while a white precipitate of magnesium carbonate isformed with normal carbonates.Heating the mixture, a white precipitate of magnesium carbonate is formed:
Mg2+ + 2 HCO3- → MgCO3 + H2O + CO2 ↑
carbon dioxide, formed in this way, can be identified with lime water or baryta water.
3. Mercury(II) chloride. No precipitate is formed with hydrogen carbonate ions, whilein a solution of normal carbonates a reddish-brown precipitate of basic mercury(II)carbonate (3HgO.HgCO3 = Hg4O3CO3) is formed:
CO32- + 4 Hg2+ 3 H2O → Hg4O3CO3 ↓ + 6 H+
the excess of carbonate acts as a buffer, reacting with the hydrogen ions formed in thereaction.
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4. Solid test On heating some solid alkali hydrogen carbonate in a dry test tube carbondioxide is evolved:
2 NaHCO3 → Na2CO3 + H2O + CO2 ↑
The gas can be identified with lime water or baryta water.
Silicates, SiO32−
The silicic acids may be represented by the general formula xSiO2.yH2O. Saltscorresponding to orthosilicic acid, H4SiO4 (SiO2.2H2O) metasilicic acid, H2SiO3(SiO2.H2O), and disilicic acid H2Si2O5 (2SiO2.H2O) are definitely known. Themetasilicates are sometimes designated simply as silicates.
Solubility. Only the silicates of the alkali metals are soluble in water; they arehydrolysed in aqueous solution and therefore react alkaline.
SiO32− + 2 H2O → H2SiO3 + 2 OH−
To study these reactions use a 1 M solution of sodium silicate, Na2SiO3.
1. Dilute hydrochloric acid. Add dilute hydrochloric acid to the solution of the silicate;a gelatinous precipitate of metasilicic acid is obtained, particularly on boiling:
SiO32− + 2 H+ → H2SiO3 ↓
2. Ammonium chloride or ammonium carbonate solution: gelatinous precipitate ofsilicic acid:
SiO32− + 2 NH4+ → H2SiO3 ↓ + 2 NH3
3. Silver nitrate solution: yellow precipitate of silver silicate:
SiO32− + 2 Ag+ → Ag2SiO3 ↓
Precipitate is soluble in dilute acids and in ammonia solution.
4. Barium chloride solution: white precipitate of barium silicate, soluble in dilute nitricacid:
SiO32- + Ba2+ → BaSiO3 ↓
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5. Ammonium molybdate solution. Add acidified (NH4)2MoO4 solution to the solutionof the silicate; a yellow coloration of the solution is obtained due to the formation of theammonium salt of silicomolybdic acid, H4[SiMo12O40]:
SiO32- + 12 MoO42- + 4 NH4+ + 22 H+ → (NH4)4[Si(Mo3O10)4] + 11 H2O
Add tin(II) chloride to the solution; the ammonium salt of silicomolybdic acid is reducedto 'molybdenum blue'.
Tin(II) ions, Sn2+ In acid solution the tin(II) ions Sn2+ are present, while in alkaline solutions
tetrahydroxo-stannate(II) ions [Sn(OH)4]2− are to be found. They form an equilibriumsystem:
Sn2+ + 4 OH− ↔ [Sn(OH)4]2−
Use a 0.1 M solution of tin(II) chloride, SnCl2.2H2O, for studying the reactionsof tin(II) ions. The solution should contain a few per cent hydrochloric acid to preventhydrolysis.
1. Hydrogen sulphide: brown precipitate of tin(II) sulphide, solubility product constantKsp(SnS, 25 °C)= 3.25x10−28, from not too acidic solutions:
Sn2+ + H2S → SnS ↓ + 2 H+
The precipitate is soluble in concentrated hydrochloric acid (distinction from arsenic(III)sulphide and mercury(II) sulphide); it is also soluble in ammonium polysulphide, but notin ammonium sulphide solution, to form a thiostannate. Treatment of the solution ofthiostannate with an acid yields a yellow precipitate of tin(IV) sulphide:
SnS ↓ + S22− → SnS32−SnS32− + 2 H+ → SnS2 ↓ + H2S ↑
2. Sodium hydroxide solution: white precipitate of tin(II) hydroxide, Ksp(Sn(OH)2, 25°C)= 5.45x10−27, which is soluble in excess alkali:
Sn2+ + 2 OH− ↔ Sn(OH)2 ↓Sn(OH)2 ↓ + 2 OH− ↔ [Sn(OH)4]2−
With ammonia solution, white tin(II) hydroxide is precipitated, which can not bedissolved in excess ammonia.
3. Mercury(II) chloride solution: a white precipitate of mercury(I) chloride (calomel) isformed if a large amount of the reagent is added quickly:
Sn2+ + 2 HgCl2 → Hg2Cl2 ↓ + Sn4+ + 2 Cl−
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If however tin(II) ions are in excess, the precipitate turns grey, especially on warming,owing to further reduction to mercury metal:
Sn2+ + Hg2Cl2 ↓ → 2 Hg ↓ + Sn4+ + 2 Cl−
4. Bismuth nitrate and sodium hydroxide solutions: black precipitate of bismuth metal:
3 Sn2+ + 18 OH− + 2 Bi3+ → 2 Bi ↓ + 3 [Sn(OH)6]2−
5. Metallic zinc spongy tin is deposited which adheres to the zinc.
6. Iron(III) nitrate and ammonium rhodanide solutions: the red solution of Fe(SCN)3is decolorised due to the reduction of iron(III) to iron(II) by tin(II) ions.Tin(II) ions must be in excess.
7. Luminescence test (chemiluminescence of SnH4). This test is based upon the fact thatsoluble compounds of tin are reduced by'nascent' hydrogen in acid solution toSnH4:
Sn2+ + 3 Zn + 4 H+ → SnH4 + 3 Zn2+
SnH4 is decomposes to Sn and H2 whenbrought into the hot flame of a Bunsenburner, with yielding a characteristicblue light.
Tin(IV) ions, Sn4+ In acid solution the tin(IV) ions Sn4+ are present, while in alkaline solutions
hexahydroxostannate(IV) ions [Sn(OH)6]2− are to be found. They form an equilibriumsystem:
Sn4+ + 6 OH− ↔ [Sn(OH)6]2−
1. Hydrogen sulphide: yellow precipitate of tin(IV) sulphide SnS2 from dilute acidsolutions:
Sn4+ + 2 H2S → SnS2 ↓ + 4 H+
The precipitate is soluble in concentrated hydrochloric acid (distinction from arsenic(III)sulphide and mercury(II) sulphide), in solutions of alkali hydroxides, and also inammonium sulphide and polysulphide. Yellow tin(IV) sulphide is precipitated uponacidification:
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SnS2 ↓ + S2− → SnS32−SnS2 ↓ + 2 S22− → SnS32− + S32−SnS32− + 2 H+ → SnS2 ↓ + H2S ↑
2. Sodium hydroxide solution: gelatinous white precipitate of tin(IV) hydroxide, whichis soluble in excess alkali:
Sn2+ + 4 OH− ↔ Sn(OH)4 ↓Sn(OH)4 ↓ + 2 OH− ↔ [Sn(OH)6]2−
With ammonia and with sodium carbonate solutions, a similar white tin(IV)hydroxide is precipitated, which, however, is insoluble in excess reagent.
3. Mercury(II) chloride solution: no precipitate (difference from tin(II)).
4. Metallic iron: reduces tin(IV) ions to tin(II):
Sn4+ + Fe → Sn2+ + Fe2+
If pieces of iron are added to a solution, and the mixture is filtered, tin(II) ions can bedetected e.g. with mercury(II) chloride reagent.
5. Luminescence test (chemiluminescence of SnH4). (see in previous page)
Summarise the redox reaction of Sn2+ and Sn4+ :
Hg2+ Zn Fe3+ Fe
Sn2+
Sn4+
Standard electrode potentials at 25 °C:
Sn2+/ Sn4+: +0.151 V Hg22+:/ Hg2+: +0.920 V
Hg/ Hg22+: +0.7973 V
Sn/ Sn2+: −0.1375 V Fe2+/ Fe3+: +0.771 VFe/ Fe2+: −0.447 VZn/ Zn2+: −0.7618 V
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Lead(II) ions, Pb2+
A 0.2 M solution of lead nitrate or lead acetate can be used to study thesereactions.
1. Dilute hydrochloric acid (or soluble chlorides): a white precipitate in cold and nottoo dilute solution, solubility product constant Ksp(PbCl2, 25 °C)= 1.17x10−5:
Pb2+ + 2 Cl− → PbCl2 ↓
The precipitate is soluble in hot water, but separates again in long, needle-like crystalswhen cooling. (The solubility of PbCl2 in water at 100 °C and 20 °C is 33.4 g/l and 9.9g/l, respectively.)The precipitate is soluble in concentrated hydrochloric acid or concentrated potassiumchloride when the tetrachloroplumbate(II) ion is formed:
PbCl2 ↓ + 2 Cl− → [PbCl4]2−
If the PbCl2 precipitate is washed by decantation and dilute ammonia is added, no visiblechange occurs, though a precipitate-exchange reaction takes place and lead hydroxide isformed, Ksp(Pb(OH)2, 25 °C)= 1.42x10−20:
PbCl2 ↓ + 2 NH3 + 2 H2O → Pb(OH)2 ↓ + 2 NH4+ + 2 Cl-
2. Hydrogen sulphide: black precipitate of lead sulphide in neutral or dilute acidmedium, Ksp(PbS, 25 °C)= 9.04x10−29:
Pb2+ + H2S → PbS ↓ + 2 H+
Precipitation is incomplete if strong mineral acids are present. It is advisable to buffer themixture with sodium acetate.The precipitate decomposes when concentrated nitric acid is added, and white, finelydivided elementary sulphur is precipitated:
3 PbS ↓ + 8 HNO3 → 3 Pb2+ + 6 NO3− + 3 S ↓ + 2 NO ↑ + 4 H2O
If the mixture is boiled, sulphur is oxidised by nitric acid to sulphate which immediatelyforms white lead sulphate precipitate with the lead ions.
3. Ammonia solution: white precipitate of lead hydroxide, solubility product constantKsp(Pb(OH)2, 25 °C)= 1.42x10−20:
Pb2+ + 2 NH3 + 2 H2O → Pb(OH)2 ↓ + 2 NH4+
The precipitate is insoluble in excess reagent.
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4. Sodium hydroxide: white precipitate of lead hydroxide:
Pb2+ + 2 OH− → Pb(OH)2 ↓
The precipitate dissolves in excess reagent, when tetrahydroxoplumbate(II) ions areformed:
Pb(OH)2 ↓ + 2 OH− → [Pb(OH)4]2−
Hydrogen peroxide when added to a solution of tetrahydroxoplumbate(II), forms a blackprecipitate of lead dioxide by oxidising bivalent lead to the tetravalent state:
[Pb(OH)4]2− + H2O2 → PbO2 ↓ + 2 H2O + 2 OH−
5. Dilute sulphuric acid (or soluble sulphates): white precipitate of lead sulphate,solubility product constant Ksp(PbSO4, 25 °C)= 1.82x10−8:
Pb2+ + SO42- → PbSO4 ↓
The precipitate is insoluble in excess reagent. It is soluble in sodium hydroxide and inmore concentrated solution of ammonium tartarate in the presence of ammonia, whentetrahydroxoplumbate(II) and ditartaratoplumbate(II) ions are formed, respectively:
PbSO4 ↓ + 4 OH− → [Pb(OH)4]2− + SO42−PbSO4 ↓ + 2 C4H4O62− → [Pb(C4H4O6)2]2− + SO42−
6. Potassium chromate: yellow precipitate of lead chromate in neutral, ecetic acid, orammonia solution:
Pb2+ + CrO42- → PbCrO4 ↓
Nitric acid or sodium hydroxide dissolve the precipitate (reactions are reversible):2 PbCrO4 ↓ + 2 H+ ↔ 2 Pb2+ + Cr2O72− + H2OPbCrO4 ↓ + 4 OH− ↔ [Pb(OH)4]2− + CrO42−
7. Potassium iodide: yellow precipitate of lead iodide, solubility product constantKsp(PbI2, 25 °C)= 8.49x10−9:
Pb2+ + 2 I− → PbI2 ↓
The precipitate is moderately soluble in boiling water to yield a colourless solution, fromwhich it separates on cooling in golden yellow plates.
8. Sodium sulphite: white precipitate of lead sulphite in neutral solution:
Pb2+ + SO32− → PbSO3 ↓
The precipitate is less soluble than lead sulphate, though it can be dissolved by bothdilute nitric acid and sodium hydroxide.
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9. Sodium carbonate: white precipitate of a mixture of lead carbonate and leadhydroxide:
2 Pb2+ + 2 CO32− + H2O → Pb(OH)2 ↓ + PbCO3 ↓ + CO2 ↑
On boiling no visible change takes place. The precipitate dissolves in dilute nitric acidand in acetic acid and CO2 gas is liberated.
10. Disodium hydrogen phosphate: white precipitate of lead phosphate:
3 Pb2+ + 2 HPO42- ↔ Pb3(PO4)2 ↓ + 2 H+
Strong acids and also sodium hydroxide dissolve the precipitate.
11. Dithizone (diphenylthiocarbazone, C6H5-NH-NH-C(S)-NN-C6H5) reagent:brick-red complex salt in neutral, ammoniakal, alkaline, or alkalicyanide solution.
S CNH
N N
NHS C
NH
N N
NC S
NN
HNN
��2+ +2 + + 2 HPb Pb
The reaction is extremely sensitive, but it is not very selective. Heavy metals (silver,mercury, copper, cadmium, antimony, nickel, and zinc, etc.) interfere, but this effect maybe eliminated by conducting the reaction in the presence of much alkali cyanide.
Summarise the solubility of Sn, Pb, and Al metals in acids and alkali:
Al Sn Pb
HCl
H2SO4
HNO3
aqua regia
NaOH
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The Group Va Elements (N, P, As, Sb, Bi) andTheir Principal Anions and Cations
Nitrogen (N2) is a colourless, inert diatomic gas (boiling point: -196 °C). Nitrogenoccurs in Nature mainly as dinitrogen (N2) that comprises 78% by volume of theearth's atmosphere.
Phosphorus is solid at room temperature. There are three main forms of phosphorus:white, black, and red.White phosphorus is in the solid and liquid forms, and in the vapour phase below800 °C consists of tetrahedral P4 molecules. White phosphorus is the least stable solidallotrope, but all others revert to it when melted. It is highly reactive and toxic, and iscommonly stored under water to protect it from air.Orthorhombic black phosphorus, the most thermodynamically stable and leastreactive form, is obtained by heating white phosphorus under pressure. It has agraphitic appearance and consists of polymeric double layers. When subjected topressures above 12 kbar the orthorhombic form transforms successively to therhombohedral and cubic forms.Red phosphorus is of intermediate reactivity and is used commercially. Ordinarily itis amorphous. It is easily obtained by heating white phosphorus in a sealed vessel at ∼400 °C.
Arsenic, Antimony, and Bismuth. These elements have fewer allotropic forms thanphosphorus. For As and Sb unstable yellow allotropes comparable to whitephosphorus are obtainable by rapid condensation of vapours. They readily transformto the bright, "metallic" rhombohedral forms similar to rhombohedral blackphosphorus. This is also the commonest form for bismuth.Arsenic is a steel-grey, brittle solid with a metallic lustre. It sublimes on heating, anda characteristic garlic-like odour is apparent.Antimony is a lustrous, silver-white metal, which melts at 630 °C.Bismuth is a brittle, crystalline, reddish-white metal. It melts at 272 °C.
Solubility in water, aqueous acids and aqueous alkaliNitrogen is little soluble in water, but P, As, Sb, and Bi are not soluble.Phosphorus, arsenic, antimony, and bismuth are not affected by nonoxidizing
acids (e.g. HCl), but they are soluble in oxidising acids (e.g. in nitric acid, to produceH3PO4, H3AsO4, Sb2O3, and Bi(NO3)3).
Arsenic is insoluble in hydrochloric acid and in dilute sulphuric acid, but itdissolves readily in dilute nitric acid yielding arsenite ions and in concentrated nitricacid, aqua regia or sodium hypochlorite solution forming arsenate:
As + 4 H+ + NO3- → As3+ + NO ↑ + 2 H2O3 As + 5 HNO3 (conc) + 2 H2O → 3 AsO43- + 5 NO ↑ + 9 H+2 As + 5 OCl- + 3 H2O → 2 AsO43- + 5 Cl- + 6 H+
37
Antimony is not soluble in dilute sulphuric acid, but it dissolves slowly in hot ,concentrated sulphuric acid forming antimony(III) ions:
2 Sb + 3 H2SO4 + 6 H+ → 2 Sb3+ + 3 SO2 ↑ + 6 H2ONitric acid oxidises antimony to an insoluble product, which can be regarded as amixture of Sb2O3 and Sb2O5. These anhydrides, in turn, can be dissolved in tartaricacid. A mixture of nitric acid and tartaric acid dissolves antimony easily.
Aqua regia dissolves antimony, when antimony(III) ions are formed:Sb + HNO3 + 3 HCl → Sb3+ + 3 Cl- + NO ↑ + 2 H2O
Bismuth dissolves in oxidising acids such as concentrated nitric acid, aquaregia, or hot, concentrated sulphuric acid:
2 Bi + 8 HNO3 → 2 Bi3+ + 6 NO3- + 2 NO ↑ + 4 H2OBi + HNO3 + 3 HCl → Bi3+ + 3 Cl- + NO ↑ + 2 H2O2 Bi + 6 H2SO4 → 2 Bi3+ + 3 SO42- + 3 SO2 ↑ + 6 H2O
Only white phosphorus is soluble in aq. alkali with disproportionation:
P4 + 3 NaOH + 3 H2O → PH3 ↑ + 3 NaH2PO2
Trihydrides of the Group Va elements (XH3) The gases XH3 can be obtained by treating ammonium salts with alkalies, by
treating phosphides or arsenides of electropositive metals with acids, by reduction ofsulphuric acid solutions of arsenic, antimony, or bismuth with an electropositive metalor electrolytically. The stability falls rapidly down in the group, so the SbH3 andBiH3 are very unstable thermally, the latter having been obtained only in traces.PH3, AsH3, and SbH3 are extremely poisonous.Ammonia NH3
Ammonia is a colourless pungent gas with a normal boiling point of -33.4 °C.The liquid has a large heat of evaporation and is therefore fairly easily handled inordinary laboratory equipment. Liquid ammonia resembles water in its physicalbehaviour, being highly associated because of the polarity of the molecules and stronghydrogen bonding.Liquid ammonia has lower reactivity than H2O toward electropositive metals, whichmay dissolve physically giving blue solutions (e.g. Na).
Nitric acid (HNO3)The pure acid is a colourless liquid. The normal concentrated aqueous acid (∼
70% by weight) is colourless but often becomes yellow as a result of photochemicaldecomposition, which gives NO2: 4 HNO3 → 4 NO2 + 2 H2O + O2The acid has the highest self-ionisation of the pure liquid acids, and the overall selfdissociation is
2 HNO3 ↔ NO2+ + NO3- + H2O
Nitric acid of concentration below 2 M has little oxidising power.
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The concentrated acid is a powerful oxidising agent and, of the metals, only Au, Pt, Ir,and Re are unattacked, although a few others such as Al, Fe, Cr are rendered passive,probably owing to formation of an oxide film. The attack on metals generallyinvolves reduction of nitrate. Nonmetals are usually oxidised by HNO3 to oxo acidsor oxides.The ability of HNO3, especially in the presence of concentrated H2SO4, to nitratemany organic compounds, is attributable to the formation of the nitronium ion, NO2+.The so-called fuming nitric acid contains dissolved NO2 in excess of the amount thatcan be hydrated to HNO3 + NO. Red fuming nitric acid contains N2O4.Aqua regia (∼3 vol. of conc. HCl + 1 vol. of conc. HNO3) contains free chlorine andClNO, and it attacks gold and platinum metals, its action being more effective thanthat of HNO3 mainly because of the complexing function of chloride ion.
Most important oxides of the Group Va elementsN P As Sb Bi
NO colourlessNO2 red
P4O10 (P2O5)white
very hygroscopic
As4O6 (As2O3)white
Sb4O6 (Sb2O3)white
Bi2O3yellow
Principal anions and cations of N, P, As, Sb, and Bi:N P As Sb Bi
NH4+ ammoniumNO2+ nitronium
NO2- nitriteNO3- nitrate
H2PO2-
hypophosphiteHPO32- phosphitePO43- phosphate
(As3+)*
(As5+)*
AsO33- arseniteAsO43- arsenate
Sb3+
Sb5+
SbO+
[Sb(OH)4]−
[Sb(OH)6]−
Bi3+
BiO+
* They exist in aqueous solution in the form of arsenite and arsenate.
Characteristic reactions of NH4+, NO2-, NO3-, PO43-, AsO33-, AsO43-,Sb3+, and Bi3+ ions
Ammonium ion, NH4+
Ammonium salts are generally water-soluble compounds, forming colourlesssolutions (unless the anion is coloured).
The reactions of ammonium ions are in general similar to those of potassium,because the sizes of the two ions are almost identical.
39
Use a 0.5 M solution of ammonium chloride to study the reactions ofammonium ions.
1. Sodium hydroxide solution: ammonia gas is evolved on warming.
NH4+ + OH- → NH3 ↑ + H2O
Ammonia gas may be identifieda.) by its odour (cautiously smell the vapour after removing the test-tube from
the flame);b.) by the formation of white fumes of ammonium chloride when a glass rod
moistened with concentrated hydrochloric acid is held in the vapour;c.) by its turning moistened pH paper blue;d.) by its ability to turn filter paper moistened with mercury(I) nitrate solution
black (this is a very trustworthy test; arsine, however, must be absent):
2 NH3 + Hg22+ + NO3- → Hg(NH2)NO3 ↓ + Hg ↓ + NH4+
e.) filter paper moistened with a solution of manganese(II) chloride andhydrogen peroxide gives a brown colour, due to the oxidation of manganese by thealkaline solution thus formed:
2 NH3 + Mn2+ + H2O2 + H2O → MnO(OH)2 ↓ + 2 NH4+
2. Nessler's reagent (alkaline solution of potassium tetraiodomercurate(II)):brown precipitate or brown or yellow coloration is produced according to the
amount of ammonia or ammonium ions present. The precipitate is a basic mercury(II)amido-iodide:
NH4+ + 2 [HgI4]2- + 4 OH- → HgO.Hg(NH2)I ↓ + 7 I- + 3 H2O
The test is an extremely delicate one and will detect traces of ammonia present indrinking water. All metals except sodium or potassium, must be absent.
3. Sodium hexanitritocobaltate(III), Na3[Co(NO2)6]:yellow precipitate of ammonium hexanitritocobaltate(III), (NH4)3[Co(NO2)6]
, similar to that produced by potassium ions:
3 NH4+ + [Co(NO2)6]3- → (NH4)3[Co(NO2)6] ↓
4. Saturated sodium hydrogen tartrate solution, NaHC4H4O6:white precipitate of ammonium acid tartarate NH4HC4H4O6, similar to but
slightly more soluble than the corresponding potassium salt, from which it isdistinguished by the evolution of ammonia gas on being heated with sodiumhydroxide solution.
NH4+ + HC4H4O6- → NH4HC4H4O6 ↓
5. Perchloric acid or sodium perchlorate solution: no precipitate (distinction frompotassium).
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Compare the characteristic reactions of K+ and NH4+ ions:
Flame test cc HClO4 Tartaric acid [Co(NO2)6]3− Nessler`sreagent
K+
NH4+
Nitrites, NO2−
Silver nitrite is sparingly soluble in water (1.363 g AgNO2/100 ml water at 60°C). All other nitrites are soluble in water.
Use a 0.1M solution of potassium nitrite to study the reactions of nitrites.
1. Hydrochloric acid: Cautious addition of the acid to a solid nitrite in the cold yieldsa transient, pale-blue liquid (due to the presence of free nitrous acid, HNO2, or itsanhydride, N2O3) and the evolution of brown fumes of nitrogen dioxide, the latterbeing largely produced by combination of nitric oxide with the oxygen of the air.Similar results are obtained with the aqueous solution.
NO2- + H+ → HNO23 HNO2 → HNO3 + 2 NO ↑ + H2O2 NO ↑ + O2 ↑ → 2 NO2 ↑
2. Barium chloride solution: no precipitate.
3. Silver nitrate solution: white crystalline precipitate of silver nitrite only fromconcentrated solutions.
NO2- + Ag+ → AgNO2 ↓
4. Potassium iodide solution: the addition of a nitrite solution to a solution ofpotassium iodide, followed by acidification with acetic acid or with dilute sulphuricacid, results in the liberation of iodine, which may be identified by the blue colourproduced with starch solution. An alternative method is to extract the liberated iodinewith carbon tetrachloride.
2 NO2- + 2 I- + 4 H+ → I2 + 2 NO ↑ + 2 H2O
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5. Ammonium chloride. By boiling a solution of a nitrite with excess of solidammonium chloride, nitrogen is evolved and the nitrite is completely destroyed:
NO2- + NH4+ → N2 ↑ + 2 H2O
6. Urea: the nitrite is decomposed, and nitrogen and carbon dioxide are evolved,when a solution of a nitrite is treated with urea, CO(NH2)2, and the mixture isacidified with dilute hydrochloric acid.
2 NO2- + CO(NH2)2 + 2 H+ → 2 N2 ↑ + CO2 ↑ + 3 H2O
7. Sulphamic acid (H2N-SO3H). When a solution of a nitrite is treated withsulphamic acid, it is completely decomposed:
H2NSO3H + NO2- + H+ → N2 ↑ + 2 H+ + SO42- + H2O
8. Acidified potassium permanganate solution: decolourized by a solution of anitrite, but no gas is evolved.
5 NO2- + 2 MnO4- + 6 H+ → 5 NO3- + 2 Mn2+ + 3 H2O
9. Sulphanilic acid α-naphthylamine reagent. (Griess-Ilosvay test)This test depends upon the diazotization of sulphanilic acid by nitrous acid,
followed by coupling with α-naphthylamine to form a red azo dye:
NH 2
SO H3 NH2
NO2 H++ ++ 2 H O2+NH 2N NHSO3
The test solution must be very dilute, otherwise the reaction does not go beyond thediazotation stage.
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Nitrates, NO3−
All nitrates are soluble in water.The nitrates of mercury and bismuth yield basic salts on treatment with water;
these are soluble in dilute nitric acid.Use a 0.1M solution of potassium nitrate to study the reactions of nitrates.
1. Reduction to nitrite test. Nitrates are reduced to nitrites by metallic zinc in aceticacid solutions; the nitrite can be readily detected by means of the Griess-Ilosvay test.Nitrites, of course, interfere and are best removed with sulphamic acid.
2. Reduction of nitrates in alkaline medium. Ammonia is evolved when a solutionof a nitrate is boiled with zinc dust or gently warmed with aluminium powder andsodium hydroxide. solution. Ammonia is detected (i) by its odour, (ii) by its actionupon pH paper and (iii) upon mercury(I) nitrate paper.
NO3- + 4 Zn + 7 OH- + 6 H2O → NH3 ↑ + 4 [Zn(OH)4]2-3 NO3- + 8 Al + 5 OH- + 18 H2O → NH3 ↑ + 8 [Al(OH)4]-
Ammonium ions interfere and must be absent. Nitrites give similar reaction and maybe removed with sulphamic acid.
3. Iron(II) sulphate solution and concentrated sulphuric acid (brown ring test):This test is carried out in either of two ways:
a.) Add 3 ml freshly prepared saturatedsolution of iron(II) sulphate to 2 mlnitrate solution, and pour 3-5 ml concent-rated sulphuric acid slowly down the sideof the test tube so that the acid forms alayer beneath the mixture. A brown ringforms where the liquids meet.b.) Add 4 ml concentrated sulphuricacid slowly to 2 ml nitrate solution, mixthe liquids thoroughly and cool themixture under a stream of cold waterfrom the tap, or ice-water. Pour asaturated solution of iron(II) sulphateslowly down side of the tube so that itforms a layer on top of the liquid. Set thetube aside for 2-3 minutes. A brown ringwill form at the zone of contact of thetwo liquids.
The brown ring is due to the formation of the [Fe(NO)]2+. On shaking and warmingthe mixture the brown colour disappears, nitric oxide is evolved, and a yellowsolution of iron(III) ions remains.
2 NO3− + 4 H2SO4 + 6 Fe2+ → 6 Fe3+ + 2 NO ↑ + 4 SO4
2− + 4 H2OFe2+ + NO ↑ → [Fe(NO)]2+
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Orthophosphates, PO43-
Orthophosphoric acid (often referred to simply as phosphoric acid) is atriprotic acid giving rise to three series of salts:
primary orthophosphates, e.g. NaH2PO4;secondary orthophosphates, e.g. Na2HPO4:tertiary orthophosphates, e.g. Na3PO4.
Ordinary 'sodium phosphate' is disodium hydrogen phosphate, Na2HPO4.12H2O .
Solubility. The phosphates (primary, secondary, and tertiary) of the alkalimetals, with the exception of lithium and of ammonium, are soluble in water.The primary phosphates of the alkaline earth metals are also soluble in water.All the phosphates of the other metals, and also the secondary and tertiary phosphatesof the alkaline earth metals are sparingly soluble or insoluble in water.
To study the reactions of phosphates use a 0.1 M solution of disodiumhydrogen phosphate, Na2HPO4.12H2O .
1. Dilute hydrochloric acid. No apparent change.
2. Silver nitrate solution: yellow precipitate of normal silver orthophosphate,solubility product: Ksp(Ag3PO4, 25°C)= 8.88x10−17:
HPO42- + 3 Ag+ → Ag3PO4 ↓ + H+
The precipitate is soluble in dilute ammonia solution and in dilute nitric acid:
Ag3PO4 ↓ + 6 NH3 → 3 [Ag(NH3)2]+ + PO43−
Ag3PO4 ↓ + 2 H+ → H2PO4− + 3 Ag+
3. Barium chloride solution: white, amorphous precipitate of secondary bariumphosphate from neutral solutions, soluble in dilute mineral acids and in acetic acid.
HPO42- + Ba2+ → BaHPO4 ↓
In the presence of dilute ammonia solution, the less soluble tertiary phosphate isprecipitated:
2 HPO42- + 3 Ba2+ + 2 NH3 → Ba3(PO4)2 ↓ + 2 NH4+
4. Magnesium nitrate reagent or magnesia mixture: The former is a solutioncontaining Mg(NO3)2, NH4NO3, and a little aqueous NH3, and the latter is asolution containing MgCl2, NH4Cl, and a little aqueous NH3. With either reagent awhite, crystalline precipitate of magnesium ammonium phosphate,Mg(NH4)PO4.6H2O, is produced:
HPO42- + Mg2+ + NH3 → Mg(NH4)PO4 ↓
The precipitate is soluble in acetic acid and in mineral acids, but practically insolublein ammonia solution.
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5. Iron(III) chloride solution: yellowish-white precipitate of iron(III) phosphate:
HPO42− + Fe3+ → FePO4 ↓ + H+
The precipitate is soluble in dilute mineral acids, but insoluble in dilute acetic acid.
6. Ammonium molybdate reagent: The addition of a large excess of this reagent to asmall volume of a phosphate solution produces a yellow, crystalline precipitate ofammonium phosphomolybdate, (NH4)3[P(Mo3O10)4]. The resulting solution shouldbe strongly acid with nitric acid; the latter is usually present in the reagent andaddition is therefore unnecessary.(Prerare a clear solution of the reagent from (NH4)2MoO4 and concentrated HNO3 !)
HPO42- + 12 MoO42- + 3 NH4+ + 23 H+ → (NH4)3[P(Mo3O10)4] + 12 H2O
Reactions of arsenic(III) ions Arsenic(III) compounds can be derived from the amphoteric arsenic trioxide
As2O3. In strongly acid solutions the only detectable species is the pyramidalAs(OH)3. In strongly basic solutions the arsenite ion, AsO33-, appears to be present.
A 0.1 M solution of arsenic(III) oxide, As2O3, or sodium arsenite, Na3AsO3,can be used for studying the reactions of arsenic(III) ions.Arsenic(III) oxide does not dissolve in cold water, but by boiling the mixture for halfan hour, dissolution is complete. The mixture can be cooled without the danger ofprecipitating the oxide.
1. Hydrogen sulphide: yellow precipitate of arsenic(III) sulphide:
2 AsO33− + 6 H+ + 3 H2S → As2S3 ↓ + 6 H2O
The solution must be strongly acidic; if there is not enough acid present a yellowcoloration is visible only, owing to the formation of colloidal As2S3.The precipitate is insoluble in concentrated hydrochloric acid.The precipitate dissolves in hot concentrated nitric acid, alkali hydroxides, orammonia:
3 As2S3 ↓ + 28 HNO3 + 4 H2O → 6 AsO43− + 9 SO4
2− + 36 H+ + 28 NO ↑As2S3 ↓ + 6 OH− → AsO3
3− + AsS33− + 3 H2O
Ammonium sulphide and ammonium polysulphide also dissolves the precipitate,when thioarsenite (AsS3
3−) and thioarsenate (AsS43−) ions are formed, respectively:
As2S3 ↓ + 3 S2− → 2 AsS33−
As2S3 ↓ + 4 S22− → 2 AsS4
3− + S32−
45
On reacidifying these both decompose, when arsenic(III) sulphide or arsenic(V)sulphide, and hydrogen sulphide are formed. The excess polysulphide reagent alsodecomposes and the precipitate is contaminated with sulphur:
2 AsS33− + 6 H+ → As2S3 ↓ + 3 H2S ↑
2 AsS43− + 6 H+ → As2S5 ↓ + 3 H2S ↑S2
2− + 2 H+ → H2S ↑ + S ↓
2. Silver nitrate: yellow precipitate of silver arsenite in neutral solution (distinctionfrom arsenates):
AsO33− + 3 Ag+ → Ag3AsO3 ↓
3. Magnesia mixture (a solution containing MgCl2, NH4Cl, and a little NH3): no precipitate (distinction from arsenate).
4. Copper sulphate solution: green precipitate of copper arsenite, variouslyformulated as CuHAsO3 and Cu3(AsO3)2.xH2O, from neutral solutions:
AsO33− + Cu2+ + H+ → CuHAsO3 ↓
The precipitate soluble in acids, and also in ammonia solution. The precipitate alsodissolves in sodium hydroxide solution; upon boiling, copper(I) oxide is precipitated.
5. Potassium tri-iodide (solution of iodine in potassium iodide): oxidizes arseniteions while becoming decolourized:
AsO33− + I3
− + H2O ↔ AsO43− + 3 I− + 2 H+
The reaction is reversible, and an equilibrium is reached.
6. Bettendorff's test (tin(II) chloride solution and concentrated hydrochloric acid): a few drops of the arsenite solution are added to a solution made of 0.5 ml saturatedtin(II) chloride solution and 2 ml concentrated hydrochloric acid, and the solution isgently warmed; the solution becomes dark brown and finally black, due to theseparation of elementary arsenic:
2 AsO33− + 12 H+ + 3 Sn2+ → 2 As ↓ + 3 Sn4+ + 6 H2O
7. Marsh's test.This test is based upon the fact that all soluble compounds of arsenic are reduced by'nascent' hydrogen in acid solution to arsine (AsH3), a colourless, extremelypoisonous gas with a garlic-like odour.If the gas, mixed with hydrogen, is conducted through a heated glass tube, it isdecomposed into hydrogen and metal arsenic, which is deposited as a brownish-black'mirror' just beyond the heated part of the tube, particularly if it is cooled.The deposit is soluble in sodium hypochlorite (distinction from antimony).
46
AsO33− + 3 Zn + 9 H+ →
AsH3 ↑ + 3 Zn2+ + 3 H2O
4 AsH3 ↑ → heat → 4 As ↓ + 6 H2 ↑
2 As + 5 OCl− + 3 H2O →2 AsO4
3− + 5 Cl− + 6 H+
Reactions of arsenate ions, AsO43-Arsenic(V) compounds are derived from arsenic pentoxide, As2O5. This is the
anhydride of arsenic acid, H3AsO4, which forms salts such as sodium arsenate.Arsenic(V) therefore exists in solutions predominantly as the arsenate AsO43- ion.
A 0.1 M solution of disodium hydrogen arsenate Na2HAsO4 can be used forthe study of these reactions.
1. Hydrogen sulphide: no immediate precipitate in the presence of dilutehydrochloric acid. If the passage of the gas is continued, a mixture of arsenic(III)sulphide and sulphur is slowly precipitated. Precipitation is more rapid in hot solution.
AsO43− + H2S → AsO3
3− + S ↓ + H2O2 AsO3
3− + 6 H+ + 3 H2S → As2S3 ↓ + 6 H2O
If a large excess of concentrated hydrochloric acid is present and hydrogen sulphide ispassed rapidly into the cold solution, yellow arsenic pentasulphide is precipitated:
2 AsO43− + 5 H2S + 6 H+ → As2S5 ↓ + 8 H2O
Arsenic pentasulphide, like the trisulphide, is readily soluble in alkali hydroxides orammonia, ammonium sulphide, ammonium polisulphide, sodium or ammoniumcarbonate:
As2S5 ↓ + 6 OH− → AsS43− + AsO3S3− + 3 H2O
As2S5 ↓ + 3 S2− → 2 AsS43−
As2S5 ↓ + 6 S22− → 2 AsS4
3− + 3 S32−
As2S5 ↓ + 3 CO32− → AsS4
3− + AsO3S3− + 3 CO2
Upon acidifying these solutions with hydrochloric acid, arsenic pentasulphide isreprecipitated:
2 AsS43− + 6 H+ → As2S5 ↓ + 3 H2S ↑
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2. Silver nitrate solution: brownish-red precipitate of silver arsenate, Ksp(Ag3AsO4,25°C)= 1.03x10−22, from neutral solutions:
AsO43− + 3 Ag+ → Ag3AsO4 ↓
Soluble in acids and ammonia solution, but insoluble in acetic acid.
3. Magnesia mixture: white, crystalline precipitate of magnesium ammoniumarsenate Mg(NH4)AsO4.6H2O from neutral or ammoniacal solution:
AsO43− + Mg2+ + NH4
+ → MgNH4AsO4 ↓
Upon treating the white precipitate with silver nitrate solution containing a few dropsof acetic acid, red silver arsenate is formed:
MgNH4AsO4 ↓ + 3 Ag+ → Ag3AsO4 ↓ + Mg2+ + NH4+
4. Ammonium molybdate solution: when the reagent and nitric acid are added inconsiderable excess to a solution of an arsenate, a yellow crystalline precipitate isobtained on boiling:
AsO43− + 12 MoO4
2− + 3 NH4+ + 24 H+ → (NH4)3[As(Mo3O10)4] ↓ + 12 H2O
The precipitate is insoluble in nitric acid, but dissolves in ammonia solution and insolutions of caustic alkalis.
5. Potassium iodide solution: in the presence of concentrated hydrochloric acid,iodine is precipitated; upon shaking the mixture with 1-2 ml of carbon tetrachloride orof chloroform, the latter is coloured violet by the iodine.
AsO43− + 2 I− + 2 H+ ↔ AsO3
3− + I2 ↓ + H2O
The reaction is reversible.
Redox systems: I−/I2 H3AsO3/H3AsO4
Concentrations: [I−]= 0.1 M; [H3AsO4]= 0.1 Ma) [H3AsO3]= 0.001 Mb) [H3AsO3]= 0.01 Mc) [H3AsO3]= 0.1 M 0 1 2 3 4 5
0.2
0.3
0.4
0.5
0.6
0.7
cba
H3AsO3/ H3AsO4
I-/ I2
Redox potential (V)
pH
48
6. Bettendorff's test (tin(II) chloride solution and concentrated hydrochloric acid): the solution becomes dark brown and finally black, due to the separation ofelementary arsenic:
2 AsO43− + 16 H+ + 5 Sn2+ → 2 As ↓ + 5 Sn4+ + 8 H2O
7. Marsh's test: The soluble compounds of arsenic are reduced by 'nascent' hydrogenin acid solution to arsine (AsH3). If the gas, mixed with hydrogen, is conductedthrough a heated glass tube, it is decomposed into hydrogen and metal arsenic, whichis deposited as a brownish-black 'mirror' just beyond the heated part of the tube.
AsO43− + 4 Zn + 11 H+ → AsH3 ↑ + 4 Zn2+ + 4 H2O
4 AsH3 ↑ → heat → 4 As ↓ + 6 H2 ↑2 As + 5 OCl− + 3 H2O → 2 AsO4
3− + 5 Cl− + 6 H+
Reactions of antimony(III) ions, Sb3+Antimony(III) compounds are easily dissolved in acids, when the ion Sb3+ is
stable. If the solution is made alkaline, or the concentration of hydrogen ions isdecreased by dilution, hydrolysis occurs when antimonyl, SbO+, ions are formed:
Sb3+ + H2O ↔ SbO+ + 2 H+
Use 0.1 M solution of antimony(III) chloride, SbCl3, to study the reactions ofantimony(III) ions.
1. Hydrogen sulphide: orange-red precipitate of antimony trisulphide, Sb2S3, fromsolutions which are not too acidic.
2 Sb3+ + 3 H2S → Sb2S3 ↓ + 6 H+
The precipitate is soluble in warm concentrated hydrochloric acid (distinction andmethod of separation from arsenic(III) and mercury(II) sulphide), in ammoniumsulphide (forming a thioantimonite) and polysulphide (forming a thioantimonate), andin alkali hydroxide solutions (forming antimonite and thioantimonite).
Sb2S3 ↓ + 6 HCl → 2 Sb3+ + 6 Cl− + 3 H2S ↑Sb2S3 ↓ + 3 S2− → 2 SbS3
3−
Sb2S3 ↓ + 4 S22− → 2 SbS4
3− + S32−
Sb2S3 ↓ + 4 OH− → [Sb(OH)4]− + SbS33−
Upon acidification of the thioantimonate solution with hydrochloric acid, antimonypentasulphide is precipitated initially but usually decomposes partially into thetrisulphide and sulphur:
2 SbS43− + 6 H+ → Sb2S5 ↓ + 3 H2S ↑ Sb2S5 ↓ → Sb2S3 ↓ + 2 S ↓
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Acidification of the thioantimonite solution or the antimonite-tioantimonite mixtureleads to the precipitation of the trisulphide:
2 SbS33− + 6 H+ → Sb2S3 ↓ + 3 H2S ↑
[Sb(OH)4]− + SbS33− + 4 H+ → Sb2S3 ↓ + 4 H2O
2. Water: when the solution is poured into water, a white precipitate of antimonylchloride SbOCl is formed, soluble in hydrochloric acid and in tartaric acid solution(difference from bismuth). With a large excess of water the hydrated oxideSb2O3.xH2O is produced.
Sb3+ + Cl− + H2O ↔ SbOCl ↓ + 2 H+
SbOCl ↓ + HOOC-CH(OH)-CH(OH)-COOH →→ [OOC-CH(OH)-CH(OH)-COOSbO]− + 2 H+ + Cl−
3. Sodium hydroxide or ammonia solution: white precipitate of the hydratedantimony(III) oxide Sb2O3.xH2O soluble in concentrated solutions of alkalis formingantimonites.
2 Sb3+ + 6 OH− → Sb2O3 ↓ + 3 H2OSb2O3 ↓ + 2 OH− + 3 H2O → 2 [Sb(OH)4]−
4. Zinc, tin, or iron: a black precipitate of antimony is produced.
Standard redox potentials:
Sb/Sb3+: +0.24 V
Zn/Zn2+: −0.76 V; Fe/Fe2+: −0.44 V; Sn/Sn2+: −0.14 V.
2 Sb3+ + 3 Zn ↓ → 2 Sb ↓ + 3 Zn2+
2 Sb3+ + 3 Sn ↓ → 2 Sb ↓ + 3 Sn2+
2 Sb3+ + 3 Fe ↓ → 2 Sb ↓ + 3 Fe2+
5. Potassium iodide solution: yellow coloration owing to the formation of a complexsalt:
Sb3+ + 6 I− → [SbI6]3−
6. Marsh's test: This test is based upon the fact that soluble compounds of antimonyare reduced by 'nascent' hydrogen in acid solution to stibine (SbH3), a colourless,thermally unstable, and extremely poisonous gas.If the gas, mixed with hydrogen, is conducted through a heated glass tube, it isdecomposed into hydrogen and metal antimony, which is deposited as a brownish-black 'mirror' on both sides of the heated part of the tube.The deposit is not soluble in sodium hypochlorite (distinction from arsenic).
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Sb3+ + 3 Zn + 3 H+ → SbH3 ↑ + 3 Zn2+
4 SbH3 ↑ → heat → 4 Sb ↓ + 6 H2 ↑
Reactions of bismuth(III) ions, Bi3+The trivalent bismuth ion Bi3+ is the most common ion of bismuth.The hydroxide Bi(OH)3 is a weak base; bismuth salts therefore hydrolyse
readily, when the following process occurs:
Bi3+ + H2O ↔ BiO+ + 2 H+
The bismuthyl ion, BiO+, forms insoluble salts, like BiOCl, with most ions. If wewant to keep bismuth ions in solution, we must acidify the solution, when the aboveequilibrium shifts towards the left.
Use a 0.1 M solution of Bi(NO3)3, which contains 3-4 per cent nitric acid, tostudy these reactions.
1. Hydrogen sulphide: black precipitate of bismuth sulphide:
2 Bi3+ + 3 H2S → Bi2S3 ↓ + 6 H+
The precipitate is insoluble in cold, dilute acid and in ammonium sulphide.Boiling concentrated hydrochloric acid dissolves the precipitate, when hydrogensulphide gas is liberated.
Bi2S3 ↓ + 6 HCl → 2 Bi3+ + 6 Cl− + 3 H2S ↑
Hot dilute nitric acid dissolves bismuth sulphide, leaving behind sulphur in the formof a white precipitate.
Bi2S3 ↓ + 8 H+ + 2 NO3− → 2 Bi3+ + 3 S ↓ + 2 NO ↑ + 4 H2O
2. Ammonia solution: white basic salt of variable composition. The approximatechemical reaction is:
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Bi3+ + NO3− + 2 NH3 + 2 H2O → Bi(OH)2NO3 ↓ + 2 NH4
+
The precipitate is insoluble in excess reagent.
3. Sodium hydroxide: white precipitate of bismuth(III) hydroxide:
Bi3+ + 3 OH− → Bi(OH)3 ↓
When boiled, the precipitate loses water and turns yellowish-white:
Bi(OH)3 ↓ → BiO.OH ↓ + H2O
4. Potassium iodide when added dropweise: black precipitate of bismuth(III) iodide.The precipitate dissolves readily in excess reagent, when orange-coloured tetraiodo-bismuthate ions are formed:
Bi3+ + 3 I− → BiI3 ↓BiI3 ↓ + I− ↔ [BiI4]−
Heating the BiI3 precipitate with water, it turns orange, owing to the formation ofbismuthyl iodide:
BiI3 ↓ + H2O → BiOI ↓ + 2 H+ + 2 I−
5. Sodium tetrahydroxostannate(II) (freshly prepared): in cold solution reducesbismuth(III) ions to bismuth metal which separates in the form of a black precipitate.First the sodium hydroxide present in the reagent reacts with bismuth(III) ions, andthen bismuth(III) hydroxide is reduced by tetrahydroxostannate(II) ions when bismuthmetal and hexahydroxostannate(IV) ions are formed:
Bi3+ + 3 OH− → Bi(OH)3 ↓2 Bi(OH)3 ↓ + 3 [Sn(OH)4]2− → 2 Bi ↓ + 3 [Sn(OH)6]2−
Standard redox potentials:in HCl solution: in NaOH solution: Bi/[BiCl4]
−: +0.16 V Bi/BiO+: +0.32 V Sn2+/Sn4+: +0.15 V [Sn(OH)4]2−/[Sn(OH)6]2−: −0.93 V
6. Water: when a solution of a bismuth salt is poured into a large volume of water. awhite precipitate of the corresponding basic salt is produced, which is soluble indilute mineral acids, but is insoluble in alkali hydroxides.
Bi3+ + NO3− + H2O → BiO(NO3) ↓ + 2 H+
Bi3+ + Cl− + H2O → BiO.Cl ↓ + 2 H+
52
Compare the characteristic reactions of arsenites, arsenates, and phosphates
PO43− AsO3
3− AsO43−
H2S
Magnesia mixture
+ AgNO3 soln.
Ag+
(NH4)2MoO4
Summarise the reactions of AsO33−, Sb3+, and Bi3+ ions with various reagents:
AsO33− Sb3+ Bi3+
+ H2Sprecipitate + NH4OH + 1:1 HCl + (NH4)2CO3
+ (NH4)2SMarsh's test + NaOClBettendorff's test
AgNO3
KI
Fe
Na2[Sn(OH)4]
53
The Group VIa Elements (O, S, Se, Te) andTheir Anions
Oxygen occurs in two allotropic forms: O2 and O3 (ozone).Oxygen (O2) is a colourless gas (boiling point: -183 °C). O2 is paramagnetic in thegaseous, liquid, and solid states. Both liquid and solid O2 are pale blue. Oxygenforms compounds with all the elements except He, Ne, and possibly Ar, and itcombines directly with all the other elements except the halogens, a few noble metals,and the noble gases, either at room or elevated temperatures.Oxygen is moderately soluble in water, and the neutral water saturated with O2 is afairly good oxidising agent.Ozone (O3) is a diamagnetic triatomic molecule. The gas is blue, the liquid is deepblue (boiling point: -112 °C), and the solid is black violet (melting point: -193 °C).
Sulphur has a strong tendency to catenation. It forms open and cyclic Sn speciesfrom n=2 to n=20 for cycles and higher for chains. This leads to enormous complexityin the physical and chemical behaviour of the element. There are several allotropicforms of sulphur.The common, stable form of sulphur at room temperature is the solid orthorhombic αsulphur, containing cyclo-S8 molecules.
Selenium has a smaller tendency to catenation than sulphur. It has three forms, the αand β red selenium, containing sulphur type cyclo-Se8 molecules, and the greyselenium.Grey selenium (metallic) is the stable form. The structure, which has no sulphuranalogue, contains infinite, spiral chains of selenium atoms.
Tellurium: the one form of tellurium is silvery-white, semimetallic, andisomorphous with grey selenium. Like the latter it is virtually insoluble in all liquidsexcept those with which it reacts.
Solubility of chalcogens in water, aqueous acids and aqueous alkaliOxygen is moderately soluble in water, but S, Se, and Te are not soluble.Sulphur, selenium, and tellurium are not affected by nonoxidizing acids, but
they are soluble in hot concentrated oxidising acids to produce H2SO4, H2SeO4, andH6TeO6, respectively, e.g.:
S + 2 HNO3 → H2SO4 + 2 NO
Sulphur is soluble in aq. alkali with disproportionation:
4 S + 6 NaOH → 2 Na2S + Na2S2O3 + 3 H2O
54
Dihydrides of chalcogensExcept water, the dihydrides, H2S, H2Se, and H2Te are extremely poisonous
gases with revolting odors; the toxicity of H2S far exceeds that of HCN. All behaveas very weak acids in aqueous solution. H2S dissolves in water to give a solution ca.0.1 M under 1 atm pressure. The dissociation constants are
H2S + H2O ↔ H3O+ + HS- K1= 1.3x10-7
HS- + H2O ↔ H3O+ + S2- K2= 7.1x10-15
Hydrogen peroxide H2O2Pure H2O2 is a colourless liquid (bp 150.2 °C, mp -0.43 °C) that resembles
water in many of its physical properties, although it is denser (1.44 g/cm3 at 25 °C).Dilute or 30% H2O2 solutions are widely used as oxidants. In acid solutionoxidations with H2O2 are most often slow, whereas in basic solution they are usuallyfast.H2O2 can also be oxidised in aqueous solution e.g. with MnO4-.Test the redox behaviour of H2O2 by performing the following tests:
1. Potassium iodide and starch. If potassium iodide and starch are added tohydrogen peroxide, acidified previously by dilute sulphuric acid, iodine is formedslowly and the solution turns gradually to deeper and deeper blue:
H2O2 + 2 H+ + 2 I- → I2 + 2 H2O
2. Potassium permanganate solution: decolourized in acid solution and oxygen isevolved:
2 MnO4- + 5 H2O2 + 6 H+ → 2 Mn2+ + 5 O2 ↑ + 8 H2O
Principal anions and oxoanions of chalcogens:O S Se Te
O2- oxideO22- peroxide
S2- sulphideSn2- polysulphide
SO32- sulphiteSO42- sulphate
S2O32- thiosulphateS2O42- dithioniteS2O52- disulphiteS2O62- dithionateS2O72- disulphate
Sn+2O62- polythionateSO52- peroxo-monosulphate
S2O82- peroxodisulphate
Se2- selenide
SeO32- seleniteSeO42- selenate
Te2- telluride
TeO32- telluriteTeO42- tellurate
55
Characteristic reactions of sulphide, sulphite, sulphate, and thiosulphate anions
Sulphides, S2-
Sulphides of alkali metals are soluble in water; their aqueous solutions reactalkaline because of hydrolysis:
S2- + H2O ↔ SH- + OH-SH- + H2O ↔ H2S + OH-
The sulphides of most other metals are insoluble in water. Those of alkalineearths are sparingly soluble, but are gradually changed by contact with water intosoluble hydrogen sulphides:
e.g.: CaS + H2O → Ca2+ + SH- + OH-
The sulphides of aluminium, chromium, and magnesium can only be preparedin the dry, as they are completely hydrolysed by water:
e.g.: Al2S3 + 6 H2O → 2 Al(OH)3 ↓ + 3 H2S ↑
For the study of the reactions of sulphides use a 2M solution of sodium sulphide.
1. Hydrochloric acid: hydrogen sulphide gas is evolved:
S2- + 2 H+ → H2S ↑
The gas may be identified by its characteristic odour, and by the blackening of filterpaper moistened with lead acetate solution:
H2S ↑ + Pb2+ → PbS ↓ + 2 H+
2. Silver nitrate solution: black precipitate of silver sulphide, solubility product:Ksp(Ag2S, 25°C)= 1.09x10−49:
S2- + 2 Ag+ → Ag2S ↓
3. Barium chloride solution: no precipitate.
4. Lead acetate solution: black precipitate of lead sulphide, solubility product:Ksp(PbS, 25°C)= 9.04x10−29:
S2- + Pb2+ → PbS ↓
56
5. Sodium nitroprusside solution (Na2[Fe(CN)5NO]): transient purple colour in thepresence of solutions of alkalis. No reactions occurs with solutions of hydrogensulphide or with the free gas: if, however, filter paper is moistened with a solution ofthe reagent made alkaline with sodium hydroxide or ammonia solution, a purplecolouration is produced with free hydrogen sulphide:
S2- + [Fe(CN)5NO]2- → [Fe(CN)5NOS]4-
Sulphites, SO32-
Only the sulphites of the alkali metals and of ammonium are soluble in water.Sulphites of the other metals are either sparingly soluble or insoluble in water.
The hydrogen sulphites of the alkali metals are soluble in water; the hydrogensulphites of the alkaline earth metals are known only in solution.
Use a 0.5M solution of sodium sulphite Na2SO3 to study the reactions of sulphites.
1. Hydrochloric acid (or dilute sulphuric acid): decomposition, more rapidly onwarming, with the evolution of sulphur dioxide:
SO32- + 2 H+ → SO2 ↑ + H2O
The gas may be identified (i) by its suffocating odour of burning sulphur, (ii) by theblue colouration, due to the formation of iodine, produced when a filter paper,moistened with potassium iodate and starch solution, is held over the mouth of thetest-tube.
5 SO2 ↑ + 2 IO3- + 4 H2O → I2 + 5 SO42- + 8 H+
2. Barium chloride or (strontium chloride) solution: white precipitate of barium ( orstrontium) sulphite:
SO32- + Ba2+ → BaSO3 ↓
the precipitate dissolves in dilute hydrochloric acid, when sulphur dioxide evolves.On standing, the precipitate is slowly oxidized to the sulphate and is then insoluble indilute mineral acids; this change is rapidly effected by warming with bromine wateror with hydrogen peroxide.
2 BaSO3 ↓ + O2 → 2 BaSO4 ↓BaSO3 ↓ + Br2 + H2O → BaSO4 ↓ + 2 Br- + 2 H+BaSO3 ↓ + H2O2 → BaSO4 ↓ + H2O
57
3. Silver nitrate solution: first, no visible change occurs because of the formation ofsulphitoargentate ions:
SO32- + Ag+ → [AgSO3]-
on the addition of more reagent, a white, crystalline precipitate of silver sulphite,solubility product: Ksp(Ag2SO3, 25°C)= 1.49x10−14, is formed:
[AgSO3]- + Ag+ → Ag2SO3 ↓
The precipitate dissolves if sulphite ions are added in excess:
Ag2SO3 ↓ + SO32- → 2 [AgSO3]-
On boiling the solution of the complex salt, or an aqueous suspension of theprecipitate, grey metallic silver is precipitated.The precipitate is soluble in dilute nitric acid, when sulphur dioxide gas is evolved.The precipitate also dissolves in ammonia.
4. Lead acetate or lead nitrate solution: white precipitate of lead sulphite:
SO32- + Pb2+ → PbSO3 ↓
The precipitate dissolves in dilute nitric acid. On boiling, the precipitate is oxidizedby atmospheric oxygen and white lead sulphate is formed:
2 PbSO3 ↓ + O2 → 2 PbSO4 ↓
This reaction can be used to distinguish sulphites and thiosulphates; the latter producea black precipitate on boiling.
5. Potassium dichromate solution (acidified with dilute sulphuric acid before the test): a green colouration, owing to the formation of chromium(III) ions:
3 SO32- + Cr2O72- + 8 H+ → 2 Cr3+ + 3 SO42- + 4 H2O
6. Lime water: This test is carried out by adding dilute hydrochloric acid to the solidsulphite, and bubbling the evolved sulphur dioxide through lime water; a whiteprecipitate of calcium sulphite is formed.
SO32- + Ca2+ → CaSO3 ↓
The precipitate dissolves on prolonged passage of the gas, due to the formation ofhydrogen sulphite ions:
CaSO3 ↓ + SO2 + H2O → Ca2+ + 2 HSO3-
A turbidity is also produced by carbonates; sulphur dioxide must therefore be firstremoved when testing for the latter. This may be effected by adding potassium
58
dichromate solution to the test-tube before acidifying. The dichromate oxidizes anddestroys the sulphur dioxide without affecting the carbon dioxide.
7. Zinc and sulphuric acid: hydrogen sulphide gas is evolved, which may bedetected by holding lead acetate paper to the mouth of the test-tube:
SO32- + 3 Zn + 8 H+ → H2S ↑ + 3 Zn2+ + 3 H2O
Compare the characteristic reactions of sulphites and carbonates
HCl Ba2+ Ag+ Pb2+K2Cr2O7
+acid
Zn+
H2SO4
CO32−
SO32−
Thiosulphates, S2O32-
Most of the thiosulphates that have been prepared are soluble in water;those of lead, silver, and barium are very sparingly soluble. Lead and silverthiosulphate dissolve in excess sodium thiosulphate solution forming complex salts.
To study the reactions of thiosulphates use a 0.5M solution of sodiumthiosulphate.
1. Hydrochloric acid: no immediate change in the cold with a solution of athiosulphate; the acidified liquid soon becomes turbid owing to the separation ofsulphur, and sulphur dioxide is evolved (especially by heating) which can berecognized by its odour and its action upon filter paper moistened with potassiumiodate and starch solution.
S2O32- + 2 H+ → S ↓ + SO2 ↑ + H2O
2. Barium chloride solution: white precipitate of barium thiosulphate frommoderately concentrated solutions.
S2O32- + Ba2+ → BaS2O3 ↓
59
3. Silver nitrate solution: white precipitate of silver thiosulphate:
S2O32- + 2 Ag+ → Ag2S2O3 ↓
At first no precipitation occurs because the soluble dithiosulphatoargentate(I)complex is formed. The precipitate is unstable, turning dark on standing, when silversulphide is formed:
Ag2S2O3 ↓ + H2O → Ag2S ↓ + 2 H+ + SO42-
The decomposition can be accelerated by warming.
4. Lead acetate or lead nitrate solution: first no change, but on further addition ofthe reagent a white precipitate of lead thiosulphate is formed:
S2O32- + Pb2+ → PbS2O3 ↓
The precipitate is soluble in excess thiosulphate.On boiling the suspension the precipitate darkens, forming finally a black precipitateof lead sulphide:
PbS2O3 ↓ + H2O → PbS ↓ + 2 H+ + SO42-
5. Iodine solution: decolourized when a colourless solution of tetrathionate ions isformed:
I2 + 2 S2O32- → 2 I- + S4O62-
Sulphates, SO42-
The sulphates of barium, strontium, and lead are practically insoluble in water,those of calcium and mercury(II) are slightly soluble.
Most of the remaining metallic sulphates are soluble in water.
Some basic sulphates such as those of mercury, bismuth, and chromium, arealso insoluble in water, but these dissolve in dilute hydrochloric or nitric acid.
To study the reactions of sulphates use a 0.1M solution of sodium sulphate.
1. Hydrochloric acid: no visible change.
2. Barium chloride solution: white precipitate of barium sulphate, solubility product:Ksp(BaSO4, 25°C)= 1.07x10−10, insoluble in warm dilute hydrochloric acid and indilute nitric acid.
SO42- + Ba2+ → BaSO4 ↓
60
3. Lead acetate solution: white precipitate of lead sulphate, PbSO4 (solubilityproduct: Ksp(PbSO4, 25°C)= 1.82x10−8), soluble in hot concentrated sulphuric acid, insolutions of ammonium acetate and of ammonium tartarate, and in sodium hydroxidesolution.
SO42- + Pb2+ → PbSO4 ↓
4. Silver nitrate solution: no precipitate in dilute solutions, but white, crystallineprecipitate of silver sulphate (Ksp(Ag2SO4, 25°C)= 1.20x10−5) from concentratedsolutions:
SO42- + 2 Ag+ → Ag2SO4 ↓
5. Mercury(II) nitrate solution: yellow precipitate of basic mercury(II) sulphate:
SO42- + 3 Hg2+ + 2 H2O → HgSO4.2HgO ↓ + 4 H+
This is a sensitive test, given even by suspensions of barium or lead sulphate.
61
Summarise the reactions of sulphide, sulphite, sulphate, and thiosulphate anionsby filling in the following tables:
HCl Ag+ Ba2+ Pb2+ Sr2+
S2-
SO32-
S2O32-
SO42-
Cl2-water I2(soln. is
decolourised)
KI(I2
precipitates)
KMnO4acidic media(decolourised
)
S2-
SO32-
S2O32-
SO42-
62
Summarise the reactions of As3+, Sb3+, Bi3+, Sn2+, Sn4+, and Pb2+ ions:
As3+ Sb3+ Bi3+ Sn2+ Sn4+ Pb2+
HCl
H2S
precipitate+
(NH4)2S
(NH4)2Sx
HNO3
cc HCl
NaOH
NH4OH
KI
SO42-
PO43-
CO32-
NaOH
Zn +H2SO4
Fe
63
The Group VIIa Elements (F, Cl, Br, I) andTheir Anions
Because of their reactivity, none of the halogens occurs in the elemental statein Nature. Since the atoms are only one electron short of the noble gas configuration,the elements readily form the anion X- or a single covalent bond.
Fluorine (F2) is a yellow gas. It is the chemically most reactive of all the elementsand combines directly at ordinary or elevated temperatures with all he elements otherthan nitrogen, oxygen, and the lighter noble gases, and also attacks many othercompounds. It reacts with water (chemically soluble in water):
2 F2 + 2 H2O → 2 H2F2 + O2
Chlorine (Cl2) is a greenish gas. It is moderately soluble in water, with which itpartially reacts (physical and chemical solubility):
Cl2 + H2O → HCl + HOCl
Bromine (Br2) is a dense, mobile, dark red liquid at room temperature. It ismoderately soluble in water and miscible with non-polar solvents such as CCl4 andCS2.Iodine (I2) is a black solid with a slight metallic lustre. At atmospheric pressure itsublimes (violet vapour) without melting. It is only slightly soluble in water (thesolubility in water is two order smaller than that of bromine), but it is readily solublein non polar solvents such as CCl4 and CS2 to give violet solutions. Iodine solutionsare brown in solvents such as unsaturated hydrocarbons, liquid SO2, alcohols, andketones, and pinkish brown in benzene.Although iodine is slightly soluble in water, it is very soluble in aqueous solution ofpotassium iodide, because iodide ions have a pronounced tendency to interact withone (or more) molecules of I2 to form polyiodide anions:
I- + I2 → I3-
Solubility of halogens in aqueous alkaliFluorine reacts fast with aqueous alkali forming fluorides and oxygen gas.By dissolving the halogens in cold base, halide (X-) and hipohalite (XO-) ions
are produced in principle according to the general reaction:
X2 + 2 OH- → X- + XO- + H2O
This situation, however, is complicated by the tendency of the hypohalite ions todisproportionate further in basic solution to produce the halate (XO3-) ions:
3 XO- → 2 X- + XO3-
The disproportionation of ClO- is slow at and below room temperature, thuswhen chlorine reacts with base " in the cold," reasonably pure solutions of Cl- and
64
ClO- are obtained. In hot solutions (70-90 °C) the rate of disproportionation is fairlyrapid and under proper conditions, good yields of ClO3- can be secured.
The disproportionation of BrO- is moderately fast even at room temperature.Consequently solutions of BrO- can only be made and kept at around 0 °C. Attemperatures of 50 to 80 °C quantitative yields of BrO3- are obtained:
3 Br2 + 6 OH- → 5 Br- + BrO3- + 3 H2O
The rate of disproportionation of IO- is very fast at all temperatures, so that itis unknown in solutions. Reaction of iodine with base gives IO3- quantitatively:
3 I2 + 6 OH- → 5 I- + IO3- + 3 H2O
Principal anions and oxoanions of halogensF Cl Br I
F− fluoride Cl− chlorideClO− hypochlorite
ClO2− chlorite
ClO3− chlorate
ClO4− perchlorate
Br− bromide
BrO3− bromate
I− iodide
IO3− iodate
Characteristic reactions of halide ionsTo study the reactions of fluorides, chlorides, bromides, and iodides
use a 0.1 M solution of sodium fluoride (NaF), sodium chloride (NaCl), sodiumbromide (NaBr), and potassium iodide (KI), respectively.
Fluorides, F-
The fluorides of the common alkali metals and of silver, mercury, aluminium,and nickel are readily soluble in water.
The fluorides of lead, copper, iron(III), barium, and lithium are slightlysoluble in water.
The fluorides of calcium, strontium, and magnesium are insoluble in water.
E.g. at 20 °C: Compound Solubility ( g / 100 ml H2O )CaF2PbF2
0.00160.064
KFAgF
92.3185
1. Silver nitrate solution: no precipitate, since silver fluoride is soluble in water.
65
2 Calcium chloride solution: white precipitate of calcium fluoride, solubilityproduct: Ksp(CaF2, 25°C)= 1.46x10−10:
2 F- + Ca2+ → CaF2 ↓
3. Barium chloride: white precipitate of barium fluoride, solubility product:Ksp(BaF2, 25°C)= 1.84x10−7:
2 F- + Ba2+ → BaF2 ↓
4. Concentrated sulphuric acid: with the solid fluoride, a colourless, corrosive gas,hydrogen fluoride, H2F2, is evolved on warming; the gas fumes on moist air, and thetest-tube acquires a greasy appearance as a result of the corrosive action of the vapouron the silica in the glass, which liberates the gas, silicon tetrafluoride, SiF4.Gelatinous silicic acid H2SiO3 is deposited on the glass as a product of thedecomposition of the silicon tetrafluoride.
2 F- + H2SO4 → H2F2 ↑ + SO42-SiO2 + 2 H2F2 → SiF4 ↑ + 2 H2O3 SiF4 ↑ + 3 H2O → 2 [SiF6]2- + H2SiO3 ↓ + 4 H+
Chlorides, Cl-
Most chlorides are soluble in water.
Mercury(I) chloride, Hg2Cl2, silver chloride, AgCl, lead chloride, PbCl2 (thisis sparingly soluble in cold but readily soluble in boiling water), copper(I) chloride,CuCl, bismuth oxychloride, BiOCl, antimony oxychloride, SbOCl, and mercury(II)oxychloride, Hg2OCl2, are insoluble in water.
1. Silver nitrate solution: white, curdy precipitate of silver chloride, solubilityproduct: Ksp(AgCl, 25°C)= 1.77x10−10:
Cl- + Ag+ → AgCl ↓
The precipitate is insoluble in water and dilute nitric acid, but soluble in diluteammonia solution, in conc. ammonium carbonate solution, and in potassium cyanideand sodium thiosulphate solutions:
AgCl ↓ + 2 NH3 → [Ag(NH3)2]+ + Cl-AgCl ↓ + (NH4)2CO3 → [Ag(NH3)2]+ + Cl- + CO2 + H2OAgCl ↓ + 2 Na2S2O3 → 4 Na+ + [Ag(S2O3)2]3- + Cl-AgCl ↓ + 2 KCN → 2 K+ + [Ag(CN)2]- + Cl-
66
2. Lead acetate solution: white precipitate of lead chloride from concentratedsolutions, solubility product: Ksp(PbCl2, 25°C)= 1.17x10−5:
2 Cl- + Pb2+ → PbCl2 ↓
3. Concentrated sulphuric acid: decomposition of chloride occurs with theevolution of hydrogen chloride:
Cl- + H2SO4 → HCl + HSO4-
4. Manganese dioxide and concentrated sulphuric acid: chlorine evolves, when thechloride is mixed with manganese dioxide, concentrated sulphuric acid added and themixture gently warmed. Chlorine is identified by its suffocating odour, yellowish-green colour, its bleaching of moistened litmus paper, and turning of potassiumiodide-starch paper blue.
2 Cl- + MnO2 + 2 H2SO4 → Mn2+ + Cl2 ↑ + 2 SO42- + 2 H2O
5. Potassium dichromate and sulphuric acid (chromyl chloride test): the solid chloride is mixed with three times itsweight of powdered potassium dichromate in atest tube, an equal bulk of concentratedsulphuric acid is added and the mixture gentlywarmed. The deep-red vapours of chromylchloride, CrO2Cl2, which are evolved arepassed into sodium hydroxide solutioncontained in a test tube. Chromyl chloride is areadily volatile liquid (b.p. 116.5 oC) anddecompose in sodium hydroxide solution toform sodium chromate, thus a yellow solutionis obtained. The formation of chromate can be confirmed by acidifying the solution with sulphuric acid, adding 1-2 ml amyl alcoholfollowed by a little hydrogen peroxide solution. The organic layer is coloured blue(formation of CrO5).
4 Cl- + Cr2O72- + 6 H+ → 2 CrO2Cl2 ↑ + 3 H2OCrO2Cl2 + 4 OH- → CrO42- + 2 Cl- + 2 H2O
Bromides and chlorides give rise to the free halogens, which yield colourlesssolutions with sodium hydroxide. If the ratio of iodide and chloride exceeds 1:15, thechromyl chloride formation is largely prevented and chlorine is evolved.Fluorides give rise to the volatile chromyl fluoride, which is decomposed by water,and hence should be absent or removed, if this test is used for identifying chlorides.
67
Bromides, Br-Silver, mercury(I), and copper(I) are insoluble in water. Lead bromide is
sparingly soluble in cold, but more soluble in boiling water.All other bromides are soluble in water.
1. Silver nitrate solution: pale-yellow precipitate of silver bromide, solubilityproduct: Ksp(AgBr, 25°C)= 5.35x10−13:
Br- + Ag+ → AgBr ↓
The precipitate is sparingly soluble in dilute, but readily soluble in concentratedammonia solution. The precipitate is also soluble in potassium cyanide and sodiumthiosulphate solutions, but insoluble in dilute nitric acid.
AgBr ↓ + 2 NH3 → [Ag(NH3)2]+ + Br-AgBr ↓ + 2 Na2S2O3 → 4 Na+ + [Ag(S2O3)2]3- + Br-AgBr ↓ + 2 KCN → 2 K+ + [Ag(CN)2]- + Br-
2. Lead acetate solution: white crystalline precipitate of lead bromide solubilityproduct: Ksp(PbBr2, 25°C)= 6.60x10−6:
2 Br- + Pb2+ → PbBr2 ↓
The precipitate is soluble in boiling water, and reappears by cooling down the hotsolution.
3. Concentrated sulphuric acid: first a reddish-brown solution is formed, laterreddish-brown vapours containing bromine and hydrogen bromide is evolved:
Br- + H2SO4 → HBr ↑ + HSO4-2 Br- + 2 H2SO4 → Br2 ↑ + SO2 ↑ + SO42- + 2 H2O
These reactions are accelerated by warming.
4. Manganese dioxide and concentrated sulphuric acid: reddish-brown vapours ofbromine are evolved. Bromine is recognised by its irritating odour.
2 Br- + MnO2 + 2 H2SO4 → Br2 ↑ + Mn2+ + 2 SO42- + 2 H2O
5. Chlorine water: the addition of this reagent dropwise to a solution of a bromideliberates free bromine, which colours the solution orange-red; if carbon tetrachlorideis added and the liquid is shaken, the bromine dissolves in the organic solvent and,after allowing to stand, forms a reddish-brown solution below the colourless aqueouslayer.With excess chlorine water, the bromine is converted into yellow brominemonochloride (or partially into colourless hypobromous or bromic acid) and a pale-yellow solution results (difference from iodide).
2 Br- + Cl2 → Br2 + 2 Cl-Br2 + Cl2 → 2 BrCl
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6. Nitric acid: hot, fairly concentrated (1:1) nitric acid oxidises bromides tobromine:
6 Br- + 8 HNO3 → 3 Br2 ↑ + 2 NO ↑ + 6 NO3- + 4 H2O
7. Potassium dichromate and sulphuric acid: on gently warming a mixture of abromide, concentrated sulphuric acid, and potassium dichromate, and passing theevolved vapours into water, a yellowish-brown solution, containing free bromine butno chromium, is produced. A colourless solution is obtained on treatment withsodium hydroxide solution; this does not give the chromate reaction with sulphuricacid, hydrogen peroxide and amyl alcohol (distinction from chloride).
6 Br- + Cr2O72- + 7 H2SO4 → 3 Br2 ↑ + 2 Cr3+ + 7 SO42- + 7 H2O
Iodides, I-
Silver, mercury, and copper(I) iodides are insoluble in water. Lead iodide issparingly soluble in cold, but more soluble in boiling water.
All other iodides are soluble in water.
E.g.: Compound Solubility ( g / 100 ml water) 20 °C 100 °C
AgIPbI2
----- ----- 0,063 0,41
CaI2 209 426
1. Silver nitrate solution: yellow, curdy precipitate of silver iodide, solubilityproduct: Ksp(AgI, 25°C)= 8.51x10−17, readily soluble in potassium cyanide and insodium thiosulphate solutions, very slightly soluble in concentrated ammoniasolution, and insoluble in dilute nitric acid:
I- + Ag+ → AgI ↓AgI ↓ + 2 Na2S2O3 → 4 Na+ + [Ag(S2O3)2]3- + I-AgI ↓ + 2 KCN → 2 K+ + [Ag(CN)2]- + I-
2. Lead acetate solution: yellow precipitate of lead iodide, solubility product:Ksp(PbI2, 25°C)= 8.49x10−9, soluble in much hot water forming a colourless solution,and yielding golden-yellow plates on cooling:
2 I- + Pb2+ → PbI2 ↓
3. Concentrated sulphuric acid: iodine is liberated; on warming, violet vapours areevolved, which turn starch paper blue. Some hydrogen iodide is formed, but most of it
69
reduces the sulphuric acid to sulphur dioxide, hydrogen sulphide, and sulphur, therelative proportions of which depend upon the concentrations of the reagents:
2 I- + 2 H2SO4 → I2 ↑ + SO42- + 2 H2O + SO2 ↑
4. Manganese dioxide and concentrated sulphuric acid: only iodine is formed andthe sulphuric acid does not get reduced:
2 I- + MnO2 + 2 H2SO4 → I2 ↑ + Mn2+ + 2 SO42- + 2 H2O
5. Chlorine water: when this reagent is added dropwise to a solution of an iodide,iodine is liberated, which colours the solution brown; on shaking with carbontetrachloride or chloroform it dissolves in the organic phase forming a violet solution.The free iodine may also be identified by the characteristic blue colour it forms withstarch solution. If excess chlorine water is added, the iodine is oxidised to colourlessiodic acid:
2 I- + Cl2 → I2 + 2 Cl-I2 + 5 Cl2 + 6 H2O → 2 IO3- + 10 Cl- + 12 H+
6. Bromine water: iodine is liberated:
2 I- + Br2 → I2 + 2 Br-
7. Potassium dichromate and sulphuric acid: only iodine is liberated, and nochromate is present in the distillate. Difference from chloride (see chlorides).
8. Sodium nitrite solution: Iodine is liberated when this reagent is added to aniodide solution acidified with dilute acetic or sulphuric acid (difference from bromideand chloride). The iodine may be identified by colouring starch paste blue, or carbontetrachloride violet:
2 I- + 2 NO2- + 4 H+ → I2 + 2 NO ↑ + 2 H2O
9. Copper sulphate solution: brown precipitate consisting of a mixture of copper(I)iodide, CuI (solubility product: Ksp(CuI, 25°C)= 1.27x10−12), and iodine. The iodinemay be removed by the addition of sodium thiosulphate solution, and a nearly whiteprecipitate of copper(I) iodide obtained:
4 I- + 2 Cu2+ → 2 CuI ↓ + I2I2 + 2 S2O32- → 2 I- + S4O62-
10. Mercury(II) chloride solution: scarlet (red) precipitate of mercury(II) iodide,solubility product: Ksp(HgI2, 25°C)= 2.82x10−29:
2 I- + HgCl2 → HgI2 ↓ + 2 Cl-
The mercury(II) iodide precipitate dissolves in excess potassium iodide, forming atetraiodomercurate(II) complex:
HgI2 ↓ + 2 I- → [HgI4]2-
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11. Bismuth(III) nitrate solution: black precipitate of bismuth(III) iodide:
3 I- + Bi3+ → BiI3 ↓
The bismuth(III) iodide precipitate dissolves in excess potassium iodide, whenorange-coloured tetraiodo-bismuthate ions are formed:
BiI3 ↓ + I- → [BiI4]-
12. Starch test: iodides are readily oxidised in acid solution to free iodine by anumber of oxidising agents; the free iodine may then be identified by the deep-bluecoloration produced with starch solution. One of the best oxidising agent to employ isacidified potassium nitrite solution:
2 I- + 2 NO2- + 4 H+ → I2 + 2 NO ↑ + 2 H2O
or acidified hydrogen peroxide: H2O2 + 2 I- + 2 H+ → I2 + 2 H2O
Chlorates, ClO3−
All chlorates are soluble in water.
To study these reactions use a 0.1 M solution of potassium chlorate.
1. Concentrated sulphuric acid. All chlorates are decomposed with the formation ofthe greenich-yellow gas, chlorine dioxide, ClO2, which dissolves in the sulphuric acidto give an orange-yellow solution. On warming gently an explosive crackling occurs,which may develop into a violent explosion (Danger!!!!). In carrying out this test oneor two small crystalls of potassium chlorate are treated with 1 ml concentratedsulphuric acid in the cold; the yellow explosive chlorine dioxide can be seen onshaking the solution. The test-tube should not be warmed.
3 KClO3 + 3 H2SO4 → 2 ClO2 ↑ + ClO4− + 3 SO4
2− + 4 H+ + 3 K+ + H2O
2. Concentrated hydrochloric acid. All chlorates are decomposed by this acid, andchlorine, together with varying quantities of the explosive chlorine dioxide, isevolved; chlorine dioxide imparts a yellow colour to the acid. The experiment shouldbe conducted on a very small scale. The following two chemical reactions probablyoccur simultaneously:
2 KClO3 + 4 HCl → 2 ClO2 + Cl2 ↑ + 2 K+ + 2 Cl− + 2 H2O KClO3 + 6 HCl → 3 Cl2 ↑ + K+ + Cl− + 3 H2O
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3. Sodium nitrite solution. On warming this reagent with a solution of the chlorate,the latter is reduced to a chloride, which may be identified by adding silver nitratesolution after acidification with dilute nitric acid.
ClO3− + 3 NO2
− → Cl− + 3 NO3−
A solution of sulphurous acid acts similarly:
ClO3− + 3 H2SO3 → Cl− + 3 SO4
2− + 6 H+
4. Zinc and sodium hydroxide solution: the chlorate is reduced to a chloride.
ClO3− + 3 Zn + 6 OH− + 3 H2O → Cl− + 3 [Zn(OH)4]2−
The solution is acidified with dilute nitric acid after several minutes boiling andremoving the unreacted zinc with filtration, and silver nitrate is added:
Ag+ + Cl− → AgCl ↓
5. Potassium iodide solution: iodine is liberated if a mineral acid is present. If aceticacid is used, no iodine separates even on long standing.
ClO3− + 6 I− + 6 H+ → 3 I2 ↓ + Cl− + 3 H2O
6. Iron(II) sulphate solution: reduction to chloride upon boiling in the presence ofmineral acid.
ClO3− + 6 Fe2+ + 6 H+ → Cl− + 6 Fe3+ + 3 H2O
7. Silver nitrate solution: no precipitate in neutral solution or in the presence ofdilute nitric acid.
8. Barium chloride solution: no precipitate is obtained.
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Summarise the reactions of halides by filling in the following tables:
Ag+ Pb2+ Hg22+ Hg2+ Bi3+ Ba2+
F-
Cl-
Br-
I-
Cl2-water Br2-water cc H2SO4or HNO3
FeCl3 H2O2
Cl-
Br-
I-
Summarise the solubility of silver halide precipitates:
diluteHNO3 (NH4)2CO3 NH4OH Na2S2O3 KCN
AgCl
AgBr
AgI
73
Pseudohalogens and Pseudohalides
A number of nitrogen compounds are known which undergo many reactionssuggestive of the halogens. These are called the "pseudohalogens". They include suchcompounds as cyanogen (CN)2, thiocyanogen (SCN)2, oxycyanogen (OCN)2, andselenocyanogen (SeCN)2. The pseudohalogens, like the halogens, are dimeric,oxidising, and capable of forming many stable salts. They also possess outer lonepairs which become available for coordination as the atoms acquire negative charge,enabling them to form very numerous and stable complexes. Having electron donorpairs on more than one atom, they are especially able to bridge acceptors together.Like the halogens they can be prepared by electrolytic or chemical oxidation of theirsimple salts. The pseudohalides are derived from pseudohalogens, like halides fromhalogens. Azides are sometimes included also even though no corresponding dimerexists.
Pseudohalogen Pseudohalide
(CN)2 cyanogen (OCN)2 oxycyanogen (SCN)2 tiocyanogen (SeCN)2 selenocyanogen ----
CN− cyanide OCN− cyanate SCN− tiocyanate SeCN− selenocyanate NNN− azide
Characteristic reactions of thiocyanates, SCN− To study the reactions of thiocyanates use a 0.1 M solution of
potassium thiocyanate (KSCN).
The thiocyanates of silver and copper(I) are practically insoluble in water, thethiocyanates of mercury(II) and lead are sparingly soluble.
The thiocyanates of most other metals are soluble in water.
1. Silver nitrate solution: white, curdy precipitate of silver thiocyanate, solubilityproduct: Ksp(AgSCN, 25°C)= 1.03x10−12:
SCN− + Ag+ → AgSCN ↓
The precipitate is insoluble in dilute nitric acid, but soluble in ammonia solution:
AgSCN ↓ + 2 NH3 → [Ag(NH3)2]+ + SCN-
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2. Copper sulphate solution: first a green colouration, than a black precipitate ofcopper(II) thiocyanate is observed:
2 SCN− + Cu2+ → Cu(SCN)2 ↓
3. Mercury(II) nitrate solution: white precipitate of mercury(II) thiocyanate,Hg(SCN)2, readily soluble in excess of the thiocyanate solution:
2 SCN− + Hg2+ → Hg(SCN)2 ↓Hg(SCN)2 ↓ + 2 SCN- → [Hg(SCN)4]2-
4. Iron(III) chloride solution: blood-red colouration, due to the formation of acomplex:
3 SCN− + Fe3+ ↔ Fe(SCN)3
The Fe(SCN)3 complex can be extracted by shaking with ether.The red colour is removed by fluorides, when colourless, more stable fluoro-complexes are formed:
Fe(SCN)3 + 6 F− → [FeF6]3− + 3 SCN−
5. Cobalt nitrate solution: blue colouration, due to the formation oftetrathiocyanatocobaltate(II) ions:
Co2+ + 4 SCN− → [Co(SCN)4]2−
If amyl alcohol or ether is added the free acid H2[Co(SCN)4] is formed and dissolvedby the organic solvent.
2 H+ + [Co(SCN)4]2− ↔ H2[Co(SCN)4]
The test is rendered more sensitive if the solution is acidified with concentratedhydrochloric acid, when the equilibrium shifts towards the formation of the free acid.
6. Hydrochloric acid (distillation test): Free isothiocyanic acid, HNCS, can beliberated by hydrochloric acid, distilled into ammonia solution, where it can beidentified with iron(III) chloride.
Place a few drops of the test solution in atest tube, acidify with dilute hydrochloricacid, boil the solution in the test tubegently so as to distil any HNCS into theammonia solution. After distillation,acidify slightly the ammonia solutionwith dilute hydrochloric acid and add adrop of iron(III) chloride solution.A red colouration is observed.
75
7. Zinc and dilute hydrochloric acid: hydrogen sulphide and hydrogen cyanide areevolved:
SCN− + Zn + 3 H+ → H2S ↑ + HCN ↑ + Zn2+
The gas may be identified by the blackening of filter paper moistened with leadacetate solution:
H2S ↑ + Pb2+ → PbS ↓ + 2 H+
8. Sodium nitrite solution and dilute hydrochloric acid: the addition of a nitritesolution to a solution of potassium thiocyanate, followed by acidification with dilutehydrochloric acid, produces a red colour (similar to that of Fe(SCN)3) due to nitrosylthiocyanate:
SCN− + NO2− + 2 H+ → ONSCN + H2O
If carbon tetrachloride or ether is added, ONSCN is extracted by the organic solvent.Nitrosyl thiocyanate is not very stable at room temperature, especially not inconcentrate solutions or at elevated temperatures. It decomposes on heating with theformation of nitrogen oxide. The colourless nitrogen oxide gas, when mixed with air,is oxidised to red nitrogen dioxide:
2 ONSCN → 2 NO + (SCN)22 NO ↑ (colourless) + O2 ↑ → 2 NO2 ↑ (red)
9. Dilute nitric acid: decomposition upon warming, a red colouration is produced,and nitrogen oxyde and hydrogen cyanide are evolved:
SCN− + 2 NO3− + H+ → 2 NO ↑ + HCN ↑ + SO4
2−
With concentrated nitric acid a more vigorous reaction takes place, with the formationof nitrogen oxide and carbon dioxide.
10. Sulphuric acid: With the concentrated acid a yellow colouration is produced inthe cold: upon warming a violent reaction occurs:
SCN− + H2SO4 + H2O → COS ↑ + NH4+ + SO4
2−
The reaction gets slower and slower with dilution of the acid. With the 2.5 Macid no reaction occurs in the cold, but on boiling a yellow solution is formed, sulphurdioxide and a little carbonyl sulphide are evolved.
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Compare the reactions of thiocyanate ion to those of halides:
Ag+ Pb2+ Cu2+ Hg2+ Fe3+ NO2−
SCN-
Cl-
Br-
I-
Pharaoh's SnakeThe famous German chemist Friedrich Wöhler invented the trick known as
Pharaoh's snake in the XIXth century.If a pellet of Hg(SCN)2 is placed on a fireproof support or in a porcelain dish andlight is put to it, a porous mass emerges, which resembles a snake and grows everlarger.
Hg(SCN)2 may be prepared by dissolvingmercury metal in nitric acid, andprecipitating the mercury ions withpotassium thiocyanate solution as a whitepowderlike precipitate:
3 Hg + 8 HNO3 → 3 Hg(NO3)2 + 2 NO + 4 H2OHg2+ + 2 SCN− → Hg(SCN)2 ↓
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Summarise the reactions of anions:
HCl Ba2+ Ca2+ Na+ Ag+ Pb2+
CO32-
SiO32-
S2-
SO32-
S2O32-
SO42-
BO33-
PO43-
F-
Cl-
Br-
I-
NO2-
NO3-
78
The Group Ib Elements (Cu, Ag, Au) andTheir Principle Ions
. Copper is reddish coloured metal, takes on a bright metallic luster, and is soft,malleable, ductile, and a good conductor of heat and electricity (second only to silverin electrical conductivity). It melts at 1083 °C. It is only superficially oxidised in air,sometimes acquiring a green coating of hydroxo carbonate and hydroxo sulphate.
Silver. Pure silver has a brilliant white metallic luster. It is a little harder thangold and is very ductile and malleable, being exceeded only by gold and perhapspalladium. Pure silver has the highest electrical and thermal conductivity of allmetals. It melts at 962 °C. It is stable in pure air and water, but tarnishes whenexposed to ozone, hydrogen sulphide, or air containing sulphur.
Gold is a heavy metal with its characteristic yellow colour when in a mass. Inpowderous form it is usually reddish-brown, but when finely divided it may be black,ruby, or purple. It melts at 1064 °C. Gold is the most malleable and ductile metal, it issoft and a good conductor of heat and electricity, and is unaffected by air and mostreagents.
Solubility in acids and alkalis Because of their positive standard electrode potential copper, silver, and gold
are insoluble in hydrochloric acid and in dilute sulphuric acid.
Hot, concentrated sulphuric acid dissolves copper and silver, but gold isresistant against it.
Cu + 2 H2SO4 → Cu2+ + SO42− + SO2 ↑ + 2 H2O
2 Ag + 2 H2SO4 → 2 Ag+ + SO42− + SO2 ↑ + 2 H2O
Medium-concentrated (8M) nitric acid also dissolves copper and silver, butgold is resistant.
3 Cu + 8 HNO3 → 3 Cu2+ + 6 NO3− + 2 NO ↑ + 4 H2O
6 Ag + 8 HNO3 → 6 Ag+ + 6 NO3− + 2 NO ↑ + 4 H2O
Aqua regia dissolves copper and gold:
3 Cu + 6 HCl + 2 HNO3 → 3 Cu2+ + 6 Cl− + 2 NO ↑ + 4 H2OAu + 4 HCl + HNO3 → H[AuCl4] + NO ↑ + 2 H2O
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Silver and gold are resistant against alkalis.Copper is hardly soluble in strong alkalis in the presence of oxygen, but it is
soluble in concentrated ammonia solution in the presence of oxygen:
4 Cu + 8 NH3 + 2 H2O + O2 → 4 [Cu(NH3)2]+ + 4 OH−
Copper, silver, and gold is soluble in alkali cyanide solutions in the presenceof oxygen:
4 Cu + 8 KCN + 2 H2O + O2 → 8 K+ + 4 [Cu(CN)2]− + 4 OH−
4 Ag + 8 KCN + 2 H2O + O2 → 8 K+ + 4 [Ag(CN)2]− + 4 OH−
4 Au + 8 KCN + 2 H2O + O2 → 8 K+ + 4 [Au(CN)2]− + 4 OH−
Compare thestandard redox potentials
Au-Au+= +1.69 VAu-Au3+= +1.50 V
NO-NO3−= +0.96 V
Au-[AuCl4]−= +1.00 VAg-Ag+= +0.80 VCu-Cu+= +0.52 V
OH−-O2= +0.40 VCu-Cu2+= +0.34 V
H2-H+= 0.0 VCu-[Cu(NH3)2]+= -0.12 VAg-[Ag(CN)2]−= -0.31 VCu-[Cu(CN)2]−= -0.43 VAu-[Au(CN)2]−= -0.60 V
Summarise the solubility of selected elements in cold concentrated nitric acid:
Au Be Al C Si Pb
cc HNO3
80
Principal cations of copper, silver, and gold (Cu+) *
Cu2+Ag+
(Ag2+)*** (Au+) **
Au3+
* Copper(I) ions are unstable in aqueous solution. Copper(I) compounds are colourless and most ofthem are insoluble in water.** Au+ ions are exceedingly unstable with respect to the disproportionation to Au and Au3+.*** The ion is unstable to water (reduced by water into Ag+).
Oxides of copper, silver and gold
oxide Cu2O CuO Ag2O AgO Au2O3
colour red black brown black brown
Copper oxides are insoluble in water. CuO is soluble in acids, in NH4Cl orKCN solutions. Cu2O is soluble in hydrochloric acid, in ammonia, and in NH4Clsolutions, slightly soluble in dilute nitric acid.
Silver oxides are soluble in nitric acid, in sulphuric acid, and also in ammoniasolution. Ag2O is slightly soluble in water (solubility at 20 °C is 0.0013 g/100 mlwater) and its aqueous suspensions are alkaline. It is more soluble in strongly alkalinesolutions than in water, and AgOH and Ag(OH)2
− are formed. AgO is of littleimportance, it is actually AgIAgIIIO2.
Gold(III) oxide is of little importance. It is obtained in hydrated form as anamorphous brown precipitate on addition of base to AuCl4
− solutions. It is weaklyacidic and dissolves in excess strong base, probably as Au(OH)4
−. It is soluble inhydrochloric acid, in concentrated nitric acid, and in alkali cyanide solutions.
Reactions of copper(II) ions, Cu2+
Solubility of the most common copper(II) compoundsCopper(II) chloride, chlorate, nitrate, and sulphate are soluble in water.
Copper(II) acetate and fluoride are slightly soluble in water, and all the othercompounds are practically insoluble.
E.g. at 20 °C Compound Solubility ( g / 100 ml H2O)CuCl2CuF2
70,6 (0 °C-on)4,7
CuS 0,000033
To study these reactions use a 0.1 M solution of copper(II) sulphate.
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1. Hydrogen sulphide gas: black precipitate of copper(II) sulphide. Solubilityproduct: Ksp(CuS, 25 °C)= 1.27x10−36.
Cu2+ + H2S → CuS ↓ + 2 H+
The solution must be acidic in order to obtain a crystalline, well-filterable precipitate.
The precipitate is insoluble in hydrochloric acid, in boiling dilute sulphuric acid, insodium hydroxide, in ammonium sulphide, in sodium sulphide, and only very slightlysoluble in polysulphides.
Hot, concentrated nitric acid dissolves the copper(II) sulphide, leaving behind sulphuras a white precipitate:
3 CuS ↓ + 8 HNO3 → 3 Cu2+ + 6 NO3− + 3 S ↓ + 2 NO ↑ + 4 H2O
When boiled for longer, sulphur is oxidised to sulphuric acid and a clear, bluesolution is obtained:
S ↓ + 2 HNO3 → 2 H+ + SO42− + 2 NO ↑
Potassium cyanide dissolves the precipitate, when colourless tetracyanocuprate(I)ions and disulphide ions are formed (copper is reduced, sulphur is oxidised):
2 CuS ↓ + 8 CN− → 2 [Cu(CN)4]3− + S22−
2. Ammonia solution: when added sparingly, a blue precipitate of basic coppersulphate is obtained:
2 Cu2+ + SO42− + 2 NH3 + 2 H2O → Cu(OH)2.CuSO4 ↓ + 2 NH4
+
the precipitate is soluble in excess reagent, when a deep blue coloration is obtainedowing to the formation of tetraamminocuprate(II) complex ions:
Cu(OH)2.CuSO4 ↓ + 8 NH3 → 2 [Cu(NH3)4]2+ + SO42− + 2 OH−
If the solution contains ammonium salts, precipitation does not occur at all, but theblue colour is formed right away. (The reaction is characteristic for copper(II) ions inthe absence of nickel.)
3. Sodium hydroxide in cold solution: blue precipitate of copper(II) hydroxide:
Cu2+ + 2 OH− → Cu(OH)2 ↓
The precipitate is insoluble in excess reagent, but soluble in ammonia solution when adeep blue coloration is obtained:
Cu(OH)2 ↓ + 4 NH3 → [Cu(NH3)4]2+ + 2 OH−
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When heated, the precipitate is converted to black copper(II) oxide bydehydration:
Cu(OH)2 ↓ → CuO ↓ + H2O
4. Potassium iodide solution: precipitates white copper(I) iodide (solubilityproduct: Ksp(CuI, 25 °C)= 1.27x10−12), but the solution is intensely brownbecause of the formation of tri-iodide ions (iodine):
2 Cu2+ + 5 I− → 2 CuI ↓ + I3−
Adding an excess of sodium thiosulphate to the solution, tri-iodide ions are reduced tocolourless iodide ions and the white colour of the precipitate becomes visible:
I3− + 2 S2O3
2− → 3 I− + S4O62−
5. Potassium cyanide: when added sparingly forms first a yellow precipitate ofcopper(II) cyanide:
Cu2+ + 2 CN− → Cu(CN)2 ↓
The precipitate quickly decomposes into white copper(I) cyanide and cyanogen(highly poisonous gas):
2 Cu(CN)2 ↓ → 2 CuCN ↓ + (CN)2 ↑
The precipitate dissolves in excess reagent, when colourless tetracyanocuprate(I)complex is formed:
CuCN ↓ + 3 CN− → [Cu(CN)4]3−
The complex is so stable (i.e. the concentration of copper(I) ions is so low) thathydrogen sulphide cannot precipitate copper(I) sulphide from this solution.(Solubility product: Ksp(Cu2S, 25 °C)= 2.26x10−48.)
6. Potassium thiocyanate: black precipitate of copper(II) thiocyanate.
Cu2+ + 2 SCN− → Cu(SCN)2 ↓
The precipitate decomposes slowly to form white copper(I) thiocyanate (Solubilityproduct: Ksp(CuSCN, 25 °C)= 1.77x10−13):
2 Cu(SCN)2 ↓ → 2 CuSCN ↓ + (SCN)2 ↑
Copper(II) thiocyanate can be transformed to copper(I) thiocyanate immediately byadding a suitable reducing agent; e.g. a saturated solution of sulphur dioxide:
2 Cu(SCN)2 ↓ + SO2 + 2 H2O → 2 CuSCN ↓ + 2 SCN− + SO42− + 4 H+
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8. Iron: if a clean iron nail is immersed in a solution of a copper salt, a red depositof copper is obtained and equivalent amount of iron dissolves:
Cu2+ + Fe → Cu + Fe2+
The electrode potential of the copper-copper(II) system is more positive than that ofthe iron-iron(II) system.
9. Flame test: green colour is imparted to the Bunsen flame.
Compare the characteristic reactions of copper(II) and bismuth(III):
H2S NH3 soln.in excess
NaOH KI Fe flametest
Cu2+
Bi3+
Reactions of silver(I) ions, Ag+
Solubility of the most common silver(I) compoundsSilver nitrate, fluoride, chlorate, and perchlorate are readily soluble in water,
silver acetate, nitrite, and sulphate are slightly soluble, while all the other silvercompounds are practically insoluble.
E.g. at 0 °C: Compound Solubility ( g / 100 ml H2O)AgNO3Ag2SO4
1220,57
AgI 0,0000002
To study these reactions use a 0.1 M solution of silver(I) nitrate.
1. Dilute hydrochloric acid (or soluble chlorides): white precipitate of silverchloride. Solubility product: Ksp(AgCl, 25 °C)= 1.77x10−10.
Ag+ + Cl− → AgCl ↓
With concentrated hydrochloric acid precipitation does not occur. Decanting theliquid from over the precipitate, it dissolves in concentrated hydrochloric acid, when adichloroargentate complex is formed:
AgCl ↓ + Cl− ↔ [AgCl2]−
84
By diluting with water, the equilibrium shifts back to the left and the precipitatereappears.
Dilute ammonia solution dissolves the precipitate when diamminoargentatecomplex ion is formed:
AgCl ↓ + 2 NH3 ↔ [Ag(NH3)2]+ + Cl−
Dilute nitric acid or hydrochloric acid neutralises the excess ammonia, and theprecipitate reappears because of the equilibrium is shifted back towards the left.
Potassium cyanide or sodium thiosulphate dissolves the AgCl precipitate:
AgCl ↓ + 2 CN− → [Ag(CN)2]− + Cl−
AgCl ↓ + 2 S2O32− → [Ag(S2O3)2]3− + Cl−
2. Hydrogen sulphide gas: in neutral or acidic medium, black precipitate of silversulphide. Solubility product: Ksp(Ag2S, 25 °C)= 6.69x10−50.
2 Ag+ + H2S → Ag2S ↓ + 2 H+
The precipitate is insoluble in ammonium sulphide, ammonium polysulphide,ammonia, potassium cyanide, or sodium thiosulphate. (Silver sulphide can beprecipitated from solutions containing dicyanato- or dithiosulphato-argentatecomplexes with hydrogen sulphide.)
Hot, concentrated nitric acid decomposes the silver sulphide, and sulphurremains in the form of a white precipitate:
3 Ag2S ↓ + 8 HNO3 → 6 Ag+ + 6 NO3− + S ↓ + 2 NO ↑ + 4 H2O
If the mixture is heated with concentrated nitric acid for a considerable time, sulphuris oxidised to sulphate and the precipitate disappears:
S ↓ + 2 HNO3 → 2 H+ + SO42− + 2 NO ↑
3. Ammonia solution: brown precipitate of silver oxide:
2 Ag+ + 2 NH3 + H2O → Ag2O ↓ + 2 NH4+
The reaction reaches an equilibrium and therefore precipitation is incomplete at anystage. The precipitate dissolves in diluted nitric acid and also in excess of the reagent:
Ag2O ↓ + 2 HNO3 → 2 Ag+ + 2 NO3− + H2O
Ag2O ↓ + 4 NH3 + H2O → 2 [Ag(NH3)2]+ + 2 OH−
85
4. Sodium hydroxide: brown precipitate of silver oxide:
2 Ag+ + OH− → Ag2O ↓ + H2O
The precipitate is insoluble in excess reagent.The precipitate dissolves in ammonia solution or in nitric acid:
Ag2O ↓ + 4 NH3 + H2O → 2 [Ag(NH3)2]+ + 2 OH−
Ag2O ↓ + 2 HNO3 → 2 Ag+ + 2 NO3− + H2O
A well-washed suspension of the precipitate shows a slight alkaline reaction owing tothe hydrolysis equilibrium:
Ag2O ↓ + H2O ↔ 2 AgOH ↓ ↔ 2 Ag+ + 2 OH−
5. Potassium iodide: yellow precipitate of silver iodide. Solubility product:Ksp(AgI, 25 °C)= 8.51x10−17.
Ag+ + I− → AgI ↓
The precipitate is insoluble in dilute or concentrated ammonia, butdissolves readily in potassium cyanide or in sodium thiosulphate solution:
AgI ↓ + 2 CN− → [Ag(CN)2]− + I−
AgI ↓ + 2 S2O32− → [Ag(S2O3)2]3− + I−
6. Potassium chromate in neutral solution: brownish-red precipitate of silverchromate. Solubility product: Ksp(Ag2CrO4, 25 °C)= 1.12x10−12.
2 Ag+ + CrO42− → Ag2CrO4 ↓
The precipitate is soluble in ammonia solution and in dilute nitric acid:
Ag2CrO4 ↓ + 4 NH3 → 2 [Ag(NH3)2]+ + CrO42−
2 Ag2CrO4 ↓ + 2 H+ ↔ 4 Ag+ + Cr2O72− + H2O
7. Potassium cyanide solution: when added dropwise to a neutral solution of silvernitrate, white precipitate of silver cyanide is formed. Solubility product:Ksp(AgCN, 25 °C)= 5.97x10−17.
Ag+ + CN− → AgCN ↓
When potassium cyanide is added in excess, the precipitate disappears owing to theformation of dicyanoargentate ions:
AgCN ↓ + CN− → [Ag(CN)2]−
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8. Sodium carbonate solution: yellowish-white precipitate of silver carbonate.Solubility product: Ksp(Ag2CO3, 25 °C)= 8.45x10−12.
2 Ag+ + CO32− → Ag2CO3 ↓
Nitric acid and ammonia solution dissolve the precipitate:
Ag2CO3 ↓ + 2 HNO3 → 2 Ag+ + 2 NO3− + CO2 ↑ + H2O
Ag2CO3 ↓ + 4 NH3 → 2 [Ag(NH3)2]+ + CO32−
When heating, the silver carbonate precipitate decomposes and brown silver oxideprecipitate is formed:
Ag2CO3 ↓ → Ag2O ↓ + CO2 ↑
9. Disodium hydrogen phosphate in neutral solution: yellow precipitate of silverphosphate. Ksp(Ag3PO4, 25 °C)= 8.88x10−17.
3 Ag+ + HPO42− → Ag3PO4 ↓ + H+
Nitric acid and ammonia solution dissolve the precipitate:
Ag3PO4 ↓ + 3 HNO3 → 3 Ag+ + 3 NO3− + H3PO4
Ag3PO4 ↓ + 6 NH3 → 3 [Ag(NH3)2]+ + PO43−
Fill in the following table:
Fe Zn Sn Cu
Cu2+
Ag+
Standard redox potentials: Fe-Fe2+= -0.44 V; Zn-Zn2+= -0.76 V; Sn-Sn2+= -0.14 V;Cu-Cu2+= +0.34 V; Ag-Ag+= +0.80 V.
87
The Group IIb Elements (Zn, Cd, Hg) andTheir Principle Ions
Zinc is a bluish-white, lustrous metal. It is brittle at ordinary temperatures, butmalleable and ductile at 100 to 150 °C. It melts at 420 °C. It is a fair conductor ofelectricity, and burns in air at high red heat with evolution of white clouds of theoxide. Zinc is stable in air at ordinary conditions, because a protective zinc oxideand/or basic zinc carbonate layer forms on the surface.
Cadmium is a soft, malleable, ductile, and bluish-white metal which is easilycut with a knife. It melts at 321 °C. It is stable in air at ordinary conditions and it issimilar in many respects to zinc.
Mercury is a heavy, silvery-white, liquid metal; it is the only common metalliquid at ordinary temperatures. It melts at −39 °C. It easily forms alloys with manymetals, such as gold, silver, and tin, which are called amalgams.The chemistries of Zn and Cd are very similar, but that of Hg differs considerably andcannot be regarded as homologous.
Solubility in acids and alkalis The very pure zinc metal dissolves very slowly in acids and in alkalis; the
presence of impurities or contact with e.g. platinum or copper accelerates the reaction.This explains the good solubility of commercial zinc.
Zinc, owing to its negative standard electrode potential of −0.76 V, dissolvesreadily in dilute hydrochloric acid and in dilute sulphuric acid with the evolution ofhydrogen:
Zn + 2 H+ → Zn2+ + H2 ↑
With hot, concentrated sulphuric acid, sulphur dioxide is evolved:
Zn + 2 H2SO4 → Zn2+ + SO42− + SO2 ↑ + 2 H2O
Zinc dissolves in very dilute nitric acid, when no gas is evolved:
4 Zn + 10 H+ + NO3− → 4 Zn2+ + NH4
+ + 3 H2O
With increasing concentration of nitric acid, dinitrogen oxide (N2O) and nitric oxide(NO) are formed:
4 Zn + 10 H+ + 2 NO3− → 4 Zn2+ + N2O ↑ + 5 H2O
3 Zn + 8 HNO3 → 3 Zn2+ + 2 NO ↑ + 6 NO32− + 4 H2O
Concentrated nitric acid has little effect on zinc because of the low solubility of zincnitrate in such a medium.
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Cadmium dissolves slowly in dilute acids with the evolution of hydrogen,owing to its negative standard electrode potential of −0.40 V:
Cd + 2 H+ → Cd2+ + H2 ↑
It dissolves in hot, concentrated sulphuric acid with the evolution of sulphur dioxide,and in medium concentrated nitric acid with the evolution of nitrogen monoxide.
Mercury, owing to its positive standard potential (Hg/Hg2+= +0.85 V;Hg/Hg2
2+= +0.80 V), is unaffected when treated with hydrochloric acid or dilute (2M)sulphuric acid.
Hot, concentrated sulphuric acid dissolves mercury. The product is mercury(I)ion if mercury is in excess, while if the acid is in excess, mercury(II) ions are formed:
Hg in excess: 2 Hg + 2 H2SO4 → Hg22+ + SO4
2− + SO2 ↑ + 2 H2Oacid in excess: Hg + 2 H2SO4 → Hg2+ + SO4
2− + SO2 ↑ + 2 H2O
It reacts readily with nitric acid. Cold, medium concentrated (8M) nitric acidwith an excess of mercury yields mercury(I) ions, and with and excess of hotconcentrated nitric acid mercury(II) ions are formed:
Hg in excess: 6 Hg + 8 HNO3 → 3 Hg22+ + 2 NO ↑ + 6 NO3
− + 4 H2Oacid in excess: 3 Hg + 8 HNO3 → 3 Hg2+ + 2 NO ↑ + 6 NO3
− +4 H2O
Zinc dissolves in alkali hydroxides, when tetrahydroxozincate(II) is formed:
Zn + 2 OH− + 2 H2O → [Zn(OH)4]2− + H2 ↑
Cadmium and mercury is insoluble in alkalis.
Summarise the solubility of selected metals in nitric acid.
Ag Pb Zn
cold, concentratedHNO3
medium conc.(8M) HNO3
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Oxides of zinc, cadmium, and mercury:ZnO CdO Hg2O HgOwhite brown black yellow or red
The oxides are practically insoluble in water; ZnO and HgO are a very little soluble,and the solubility at 25 °C is 0.00016 g and 0.0053 g/100 ml of water, respectively.CdO and Hg2O are insoluble.They all are soluble in acids (Hg2O only in nitric acid).
Principal cations of zinc, cadmium, and mercury
Zn2+ Cd2+Hg2
2+
Hg2+
Reactions of zinc(II) ions, Zn2+
Solubility of the most common zinc compoundsZinc(II) chloride, bromide, iodide, chlorate, nitrate, sulphate, and acetate are
soluble in water. Zinc(II) fluoride is very little soluble in water, and all the othercompounds (e.g. sulphide, carbonate, phosphate) are practically insoluble.
E.g. at 20 °C: Compound Solubility ( g / 100 ml H2O)ZnBr2
Zn(NO3)2.6H2O447
184,3
ZnF2 1,62ZnS (β)Zn(CN)2
0,0000650,0005
To study these reactions use a 0.1 M solution of zinc(II) sulphate.
1. Hydrogen sulphide gas: no precipitation occurs in acidic solution (pH about 0−6),only partial precipitation of zinc sulphide in neutral solutions. In alkaline solution,e.g. adding alkali acetate, the precipitation of white zinc(II) sulphide (solubilityproduct: Ksp(ZnS, 25 °C)= 2.93x10−25) is almost complete.
Zn2+ + H2S → ZnS ↓ + 2 H+
Zinc sulphide is also precipitated from alkaline solutions of tetrahydroxozincate:
[Zn(OH)4]2− + H2S → ZnS ↓ + 2 OH− + 2 H2O
2. Ammonium sulphide: white precipitate of zinc sulphide from neutral or alkalinesolutions:
Zn2+ + S2− → ZnS ↓
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The precipitate is insoluble in excess reagent, in acetic acid, and in solutions ofcaustic alkalis, but dissolves in dilute mineral acids.
The precipitate obtained is partially colloidal; it is difficult to wash and tendsto run through the filter paper, particularly on washing. To obtain the zinc sulphide ina form which can be ready filtered, the precipitation is carried out in boiling solution. 3. Ammonia solution: white precipitate of zinc hydroxide (solubility product:
Ksp(Zn(OH)2, 25 °C)= 6.86x10−17), readily soluble in excess reagent and insolutions of ammonium salts owing to the formation of tetramminezincate(II):
Zn2+ + 2 NH3 + 2 H2O ↔ Zn(OH)2 ↓ + 2 NH4+
Zn(OH)2 ↓ + 4 NH3 ↔ [Zn(NH3)4]2+ + 2 OH−
4. Sodium hydroxide: white gelatinous precipitate of zinc(II) hydroxide:
Zn2+ + 2 OH− ↔ Zn(OH)2 ↓
The precipitate is soluble in acids and also in the excess of the reagent:
Zn(OH)2 ↓ + 2 H+ ↔ Zn2+ + 2 H2OZn(OH)2 ↓ + 2 OH− ↔ [Zn(OH)4]2−
5. Disodium hydrogen phosphate solution: white precipitate of zinc phosphate:
3 Zn2+ + 2 HPO42− ↔ Zn3(PO4)2 ↓ + 2 H+
In the presence of ammonium ions zinc ammonium phosphate is formed:
Zn2+ + NH4+ + HPO4
2− ↔ Zn(NH4)PO4 ↓ + H+
Both precipitates are soluble in dilute acids, when the reaction is reversed.Also, both precipitates are soluble in ammonia:
Zn3(PO4)2 ↓ + 12 NH3 → 3 [Zn(NH3)4]2+ + 2 PO43−
Zn(NH4)PO4 ↓ + 3 NH3 → [Zn(NH3)4]2+ + HPO42−
6. Potassium hexacyanoferrate(II) solution: white precipitate of variablecomposition; if the reagent is added in some excess the composition of theprecipitate is K2Zn3[Fe(CN)6]2:
3 Zn2+ + 2 K+ + 2 [Fe(CN)6]4− → K2Zn3[Fe(CN)6]2 ↓
The precipitate is insoluble in dilute acids, but dissolves readily in sodium hydroxide.This reaction can be used to distinguish zinc from aluminium.
91
7. Dithizone test. Dithizone (diphenyl thiocarbazone) forms complexes with anumber of metal ions, which can be extracted with carbon tetrachloride. The zinccomplex, formed in neutral, alkaline, or acetic acid solutions, is red:
S CNH
N N
NHS C
NH
N N
NC S
NN
HNNZn
���Zn2+ +2 + + 2 H
Acidify the test solution with acetic acid, and add a few drops of the reagent(dithizone dissolved in carbon tetrachloride). The organic phase turns red in thepresence of zinc. (Cu2+, Hg2
2+, Hg2+, and Ag+ ions interfere.)
Summarise the reactions of selected metal ions with various anions.
F− Cl− Br− S2− SO42− CrO4
2− CO32− OH−
Zn2+
Al3+
Ca2+
Pb2+
Reactions of cadmium(II) ions, Cd2+
Solubility of the most common cadmium compoundsCadmium acetate, sulphate, nitrate, iodide, bromide, chloride, and
chlorate are readily soluble in water, cadmium fluoride is little soluble, while all theother cadmium compounds (e.g. sulphide, carbonate, phosphate) are practicallyinsoluble.
92
E.g.: Compound Solubility ( g / 100 ml H2O)at 20 °C:
0 °C:CdCl2
Cd(NO3)2
140109
25 °C: CdF2 4,3526 °C: Cd(OH)2
Cd3(PO4)2
0,00026-----
To study these reactions use a 0.1 M solution of cadmium(II) sulphate.
1. Hydrogen sulphide gas: in acidic medium, characteristic yellow precipitate ofcadmium sulphide. Solubility product: Ksp(CdS, 25 °C)= 1.40x10−29.
Cd2+ + H2S → CdS ↓ + 2 H+
The reaction is reversible; if the concentration of a strong acid in the solution is above0.5M, precipitation is incomplete. Concentrated acids dissolve the precipitate for thesame reason.The precipitate is insoluble in potassium cyanide (difference from copper ions).
2. Ammonia solution when added dropwise: white precipitate of cadmium(II)hydroxide (Solubility product: Ksp(Cd(OH)2, 25 °C)= 5.27x10−15):
Cd2+ + 2 NH3 + 2 H2O ↔ Cd(OH)2 ↓ + 2 NH4+
The precipitate dissolves in acid when the equilibrium shifts towards the left.An excess of the reagent dissolves the precipitate, when colourlesstetramminecadminate(II) complex ions are formed:
Cd(OH)2 ↓ + 4 NH3 → [Cd(NH3)4]2+ + 2 OH−
3. Sodium hydroxide solution: white precipitate of cadmium(II) hydroxide:
Cd2+ + 2 OH− → Cd(OH)2 ↓
The precipitate is insoluble in excess reagent.Dilute acids dissolve the precipitate.
4. Potassium cyanide solution: white precipitate of cadmium cyanide, when addedslowly to the solution:
Cd2+ + 2 CN− → Cd(CN)2 ↓
An excess of the reagent dissolves the precipitate, when tetracyanocadminate(II) ionsare formed:
Cd(CN)2 ↓ + 2 CN− → [Cd(CN)4]2−
The colourless compound is not too stable; when hydrogen sulphide gas is introduced,cadmium sulphide is precipitated:
93
[Cd(CN)4]2− + H2S → CdS ↓ + 2 H+ + 4 CN−
The marked difference in the stabilities of the copper and cadmium tetracyanocomplexes serves as the basis for the separation of copper and cadmium ions, and alsofor the identification of cadmium in the presence of copper.
5. Potassium iodide: forms no precipitate (difference from copper).
Summarise the reactions of Cu2+, Cd2+, and Bi3+ ions
Cu2+ Cd2+ Bi3+
H2Sin acidic solution
NaOH
NH3in excess
KI
Fe nail
Reactions of mercury(I) ions, Hg22+
Solubility of the most common mercury(I) compounds:Mercury(I) nitrate is soluble in water and tends to decompose. Other common
inorganic salts are very slightly soluble or insoluble.
E.g. at 25 °C: Compound Solubility ( g / 100 ml H2O)Hg2SO4 0,06Hg2CO3Hg2Br2
0,00000450,000004
To study these reactions use a 0.1 M solution of mercury(I) nitrate.1. Dilute hydrochloric acid or soluble chlorides: white precipitate of mercury(I)
chloride (calomel). Solubility product: Ksp(Hg2Cl2, 25 °C)= 1.45x10−18.
94
Hg22+ + 2 Cl− → Hg2Cl2 ↓
The precipitate is insoluble in dilute acids.Ammonia solution converts the precipitate into a mixture of mercury(II)amidochloride and mercury metal, both insoluble precipitates; the mercury(II)amidochloride is a white precipitate, but the finely divided mercury makes it shinyblack (disproportionation takes place; mercury(I) is converted partly to mercury(II)and partly to mercury metal):
Hg2Cl2 ↓ + 2 NH3 → Hg ↓ + Hg(NH2)Cl ↓ + NH4+ + Cl−
This reaction can be used to differentiate mercury(I) ions from lead(II) and silver(I).
Mercury(I) chloride dissolves in aqua regia, forming undissociated but solublemercury(II) chloride:
3 Hg2Cl2 ↓ + 2 HNO3 + 6 HCl → 6 HgCl2 + 2 NO ↑ + 4 H2O
2. Hydrogen sulphide in neutral or dilute acid medium: black precipitate, which isa mixture of mercury(II) sulphide and mercury metal:
Hg22+ + H2S → Hg ↓ + HgS ↓ + 2 H+
Owing to the extremely low solubility product of mercury(II) sulphide (6.44x10−53)the reaction is very sensitive.Aqua regia dissolves the precipitate, yielding undissociated mercury(II) chloride andsulphur:
3 Hg ↓ + 3 HgS ↓ + 12 HCl + 4 HNO3 → 6 HgCl2 + 3 S ↓ + 4 NO ↑ + 8 H2O
When heated with aqua regia, sulphur is oxidised to sulphuric acid and the solutionbecomes clear:
S ↓ + 6 HCl + 2 HNO3 → SO42− + 6 Cl− + 8 H+ + 2 NO ↑
3. Ammonia solution: black precipitate which is a mixture of mercury metal andbasic mercury(II) amidonitrate, which itself is a white precipitate:
2 Hg22+ + NO3
− + 4 NH3 + H2O → 2 Hg ↓ + HgO.Hg(NH2)NO3 ↓ + 3 NH4+
This reaction can be used to differentiate between mercury(I) and mercury(II) ions.
4. Sodium hydroxide: black precipitate of mercury(I) oxide.
Hg22+ + 2 OH− → Hg2O ↓ + H2O
The precipitate is insoluble in excess reagent, but dissolves readily in dilute nitricacid.
When boiling, the colour of the precipitate turns to grey, owing todisproportionation, when mercury(II) oxide and mercury metal are formed:
Hg2O ↓ → HgO ↓ + Hg ↓
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5. Potassium chromate in hot solution: red crystalline precipitate of mercury(I)chromate:
Hg22+ + CrO4
2− → Hg2CrO4 ↓
If the test is carried out in cold, a brown amorphous precipitate is formed with anundefined composition. When heated the precipitate turns to red, crystallinemercury(I) chromate.
Sodium hydroxide turns the precipitate into black mercury(I) oxide:
Hg2CrO4 ↓ + 2 OH− → Hg2O ↓ + CrO42− + H2O
6. Potassium iodide, added slowly in cold solution: green precipitate of mercury(I)iodide:
Hg22+ + 2 I− → Hg2I2 ↓
If excess reagent is added, a disproportionation reaction takes place, solubletetraiodomercurate(II) ions and a black precipitate of finely divided mercury beingformed:
Hg2I2 ↓ + 2 I− → [HgI4]2− + Hg ↓
When boiling the mercury(I) iodide precipitate with water, disproportionation takesplace again, and a mixture of red mercury(II) iodide precipitate and finely distributedblack mercury is formed:
Hg2I2 ↓ → HgI2 ↓ + Hg ↓
7. Sodium carbonate in cold solution: yellow precipitate of mercury(I) carbonate(solubility product: Ksp(Hg2CO3, 25 °C)= 3.67x10−17):
Hg22+ + CO3
2− → Hg2CO3 ↓
The precipitate turns slowly to blackish grey, when mercury(II) oxide and mercuryare formed:
Hg2CO3 ↓ → HgO ↓ + Hg ↓ + CO2 ↑
The decomposition can be speeded up by heating the mixture.
8. Disodium hydrogen phosphate: white precipitate of mercury(I) hydrogenphosphate:
Hg22+ + HPO4
2− → Hg2HPO4 ↓
9. Potassium cyanide solution: produces mercury(II) cyanide solution and mercuryprecipitate:
Hg22+ + 2 CN− → Hg ↓ + Hg(CN)2
Mercury(II) cyanide, though soluble, is practically undissociated.
96
10. Tin(II) chloride solution: reduces mercury(I) ions to mercury metal, whichappears in the form of a greyish-black precipitate:
Hg22+ + Sn2+ → 2 Hg ↓ + Sn4+
11. Copper sheet or copper coin: deposit of mercury metal is formed on the coppersurface.
Hg22+ + Cu → 2 Hg ↓ + Cu2+
Compare the reactions of mercury(I), silver(I), and lead(II) ions.
Hg22+ Ag+ Pb2+
HCl
H2S
NH3
NH3in excess
KI
KIin excess
NaOH
Na2CO3
K2CrO4
Cu
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Reactions of mercury(II) ions, Hg2+
Solubility of the most common mercury(II) compounds:Mercury(II) nitrate is readily soluble in water; chloride, chlorate, cyanide, and
acetate are also soluble, but their solubility is much less than that of nitrate. All theother mercury(II) compounds (e.g. sulphide, carbonate, iodide) are practicallyinsoluble.
E.g.: Compound Solubility ( g / 100 ml H2O)at 20 °C:
10 °C:Hg(ClO3)2Hg(Ac)2
2525
25 °C: HgI2 (α) 0,0118 °C: HgS (α) 0,000001
To study these reactions use a 0.1 M solution of mercury(II) nitrate.
1. Hydrogen sulphide gas: black precipitate of mercury(II) sulphide. Solubilityproduct constant: Ksp(HgS, 25 °C)= 6.44x10−53.
Hg2+ + H2S → HgS ↓ + 2 H+
In the presence of dilute hydrochloric acid, initially a white precipitate of mercury(II)chlorosulphide (Hg3S2Cl2 ↓), which decomposes when further amounts of hydrogensulphide are added and finally a black precipitate of mercury(II) sulphide is formed:
3 Hg2+ + 2 Cl− + 2 H2S → Hg3S2Cl2 ↓ + 4 H+
Hg3S2Cl2 ↓ + H2S → 3 HgS ↓ + 2 H+ + 2 Cl−
The HgS precipitate is insoluble in water, hot dilute nitric acid, alkali hydroxides, orammonium sulphide.Aqua regia dissolves the precipitate:
3 HgS ↓ + 6 HCl + 2 HNO3 → 3 HgCl2 + 3 S ↓ + 2 NO ↑ + 4 H2O
Sulphur remains as a white precipitate, which however dissolves readily if thesolution is heated, to form sulphuric acid.
Sodium sulphide (2M) dissolves the HgS precipitate when thedisulphomercurate(II) complex ion is formed:
HgS ↓ + S2− → [HgS2]2−
Adding ammonium chloride to the solution, mercury(II) sulphide precipitates again.
2. Ammonia solution: white precipitate with a mixed composition; essentially itconsists of mercury(II) oxide and mercury(II) amidonitrate:
2 Hg2+ + NO3− + 4 NH3 + H2O → HgO.Hg(NH2)NO3 ↓ + 3 NH4
+
98
3. Sodium hydroxide (added in small amounts): brownish-red precipitate withvarying composition; if added in stoichiometric amounts the precipitate turns toyellow when mercury(II) oxide is formed:
Hg2+ + 2 OH− → HgO ↓ + H2O
The precipitate is insoluble in excess sodium hydroxide. Acids dissolve the precipitatereadily.
4. Potassium iodide (added slowly to the solution): red precipitate of mercury(II)iodide, Ksp(HgI2, 25 °C)= 2.82x10−29:
Hg2+ + 2 I− → HgI2 ↓
The precipitate dissolves in excess reagent, when colourless tetraiodomercurate(II)ions are formed:
HgI2 ↓ + 2 I− → [HgI4]−
5. Tin(II) chloride: when added in moderate amounts, white, silky precipitate ofmercury(I) chloride (calomel), Ksp(Hg2Cl2, 25 °C)= 1.45x10−18, is formed:
2 Hg2+ + Sn2+ + 2 Cl− → Hg2Cl2 ↓ + Sn4+
If more reagent is added, mercury(I) chloride is further reduced and black precipitateof mercury is formed:
Hg2Cl2 ↓ + Sn2+ → 2 Hg ↓ + Sn4+ + 2 Cl−
6. Copper sheet or coin: reduces mercury(II) ions to the metal:
Cu + Hg2+ → Hg ↓ + Cu2+
(Standard reduction potentials: Cu/ Cu2+= +0.3419 V; Sn2+/ Sn4+= +0.151 V;Hg2
2+/ Hg2+= +0.920 V; Hg/ Hg2+= +0.851 V; Hg/ Hg22+= +0.7973 V)
Compare the characteristic reactions of mercury(I) and mercury(II).
HCl H2S NH3 NaOH KI Cu
Hg22+
Hg2+
99
Compare the characteristic reactions of arsenic(III)(arsenite), antimony(III),tin(II), tin(IV), copper(II), cadmium(II), bismuth(III), and mercury(II) ions.
As3+ Sb3+ Sn2+ Sn4+ Cu2+ Cd2+ Bi3+ Hg2+
HCl
H2S
precipitate
+(NH4)2SX
+HCl
NaOH
KI
KCN
NH3
SnCl2
Fe
100
Titanium (Group IVb) and Its Common Ions
Titanium, when pure, is a lustrous, white metal. It has a low density, good strength,is easily fabricated, and has excellent corrosion resistance (melting point: 1660 °C).The metal is not attacked by mineral acids at room temperature or even by hotaqueous alkali.
Solubility in acids The titanium metal is not soluble in mineral acids at room temperature, but
soluble in hot, concentrated hydrochloric acid and sulphuric acid, and in hydrogenfluoride:
2 Ti + 6 HCl → 2 Ti3+ + 6 Cl− + 3 H2 ↑
Ti + 4 H2SO4 → Ti4+ + 2 SO42− + 2 SO2 ↑ + 4 H2O
The best solvents of the metal are HF and acids to which fluoride ions have beenadded; such media dissolve titanium and hold it in solution as fluoro complexes.
Titanium is insoluble in hot, concentrated nitric acid, like tin, because of theformation of titanic acid (TiO2.xH2O) on the surface of the metal, which protects therest of the metal from the acid.
Common cations of titanium in aqueous solutionTi3+
violetTi4+
colourless* Titanium(II) ions are not stable in aqueous solution; they liberate hydrogen gas from water(Ti2+/Ti3+= −0.369 V; Ti2+/Ti(OH)2
2+= −0.135 V).** Titanium(III) ions are rather unstable and are readily oxidised to titanium(IV) in aqueoussolutions, e.g. by air oxygen if exposed to air (Ti3+/Ti(OH)2
2+= +0.099 V).
The white titanium(IV) oxide, TiO2, is used as a pigment and is far the mostimportant titanium oxide, which occurs also in the nature. The solubility of TiO2depends considerably on its chemical and thermal history; strongly roasted specimensare chemically inert.
Titanium(IV) ions exist only in strongly acid solutions; they tend to hydrolyse.In strong acid the Ti4+ (aquated) ions are in equilibrium with Ti(OH)2
2+, Ti(OH)3+ and
TiO2+ ions (aquated); the main species is Ti(OH)22+ and if the acidity of the solution is
lowered titanium(IV) hydroxide is precipitated.
101
Reactions of titanium(IV) ions, Ti4+
To study these reactions use a 0.1 M solution of titanium(IV) sulphate, whichis prepared by dissolving Ti(SO4)2 in 5 per cent sulphuric acid.
1. Solutions of sodium hydroxide, ammonia or ammonium sulphide solution:white gelatinous precipitate of titanium(IV) hydroxide, Ti(OH)4 (or orthotitanicacid, H4TiO4), in the cold; this is almost insoluble in excess reagent, but soluble inmineral acids:
Ti(OH)22+ + 2 OH− → Ti(OH)4 ↓
Ti(OH)4 ↓ + H2SO4 → Ti(OH)22+ + 2 H2O + SO4
2−
Ti(OH)4 ↓ + 3 HCl → Ti(OH)Cl2+ + 3 H2O + 2 Cl−
If precipitation takes place from hot solution, white TiO(OH)2 (or metatitanic acid,H2TiO3) is formed, which is sparingly soluble in dilute acids.
Ti(OH)22+ + 2 OH− → TiO(OH)2 ↓ + H2O
2. Water: a white precipitate of metatitanic acid is obtained on boiling a solution ofa titanic salt with excess water:
Ti(OH)22+ + 2 OH− → TiO(OH)2 ↓ + H2O
3. Sodium phosphate solution: white precipitate of titanium(IV) phosphate in dilutesulphuric acid solution:
Ti(OH)22+ + 2 H2PO4
− → Ti(HPO4)2 ↓ + 2 H2O
4. Zinc or tin metal: when any of these metals is added to an acid solution of atitanium(IV) salt, a violet coloration is produced, due to reduction to titanium(III)ions:
2 Ti4+ + Zn → 2 Ti3+ + Zn2+
5. Hydrogen peroxide. An intense orange coloration is produced (yellow with verydilute solution), due to formation of stable peroxo complexes:
Ti(OH)22+ + H2O2 + OH− → Ti(O2)OH+ + 2 H2O
Ti
OH
OH2
OH2
OH2H2O
HO
4+
--�������� �������� ���������
���������
��������������������
-4+
OH2
OH2
OH2
OH
TiO
O
����������
102
Vanadium (Group Vb) and Its Common Ions
Vanadium is a bright, white metal, and is soft and ductile (melting point: 1890 °C). Ithas good corrosion resistance to alkalis, sulphuric and hydrochloric acid, and saltwaters. The metal has good structural strength.
Solubility in acids The vanadium metal is not soluble in hydrochloric, nitric, or sulphuric acids
or in alkalis at room temperature due to passivation (thin protective oxide layerforms).
It dissolves readily in aqua regia, in hot nitric acid, hot and concentratedsulphuric acid, or in a mixture of concentrated nitric acid and hydrogen fluoride:
3 V + 4 HNO3 + 6 HCl → 3 VO2+ + 6 Cl− + 4 NO ↑ + 5 H2O3 V + 10 HNO3 → 3 VO2+ + 6 NO3
− + 4 NO ↑ + 5 H2OV + 3 H2SO4 → VO2+ + SO4
2− + 2 SO2 ↑ + 3 H2O
Common ions of vanadium in aqueous solutionsoxidation state cations anions
+2 V2+
+3 V3+
+4 VO2+
+5 VO2+ VO4
3− vanadate 13<pHHVO4
2− monovanadate 8<pH<13HV2O7
3− divanadateVO3
− metavanadateV10O28
6− decavanadate 2<pH<6
* The orange decavanadate ion can exist in several protonated form, and with increasingacidity of the solution rapidly gives the dioxovanadium(V) ion, VO2
+.
Vanadium(II) and vanadium(III) ions are instable in aqueous solution, and easilyoxidised to vanadium(IV), due to their small standard reduction potentials:
[VO2(H2O)4] [VO(H2O)5] [V(H2O)6] [V(H2O)6] V 2+3+2++���� ���� ���� ���� ���� �������� ����
+0.999 V +0.359 V -0.256 V -1.186 V
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Vanadium(V) oxide, V2O5, is the most stable and is far the most important vanadiumoxide. It is an orange (or brick red) powder, which is insoluble in water, but solublein both mineral acids and alkalis.Vanadium(V) is moderately strong oxidising agent, thus if the oxide is dissolved inhydrochloric acid chlorine gas is evolved and vanadium(IV) is produced. The oxide isalso reduced by warm sulphuric acid.Vanadium pentoxide dissolves in sodium hydroxide to give colourless solutions andin the highly alkaline region, pH>13, the main ion is VO4
3−. As the basicity isreduced, a series of complicated reactions occurs with the formation of variousvanadates (mono, di, meta, deca, etc.).
Metavanadates, VO3−
To study these reactions use a 0.1 M solution of ammonium metavanadate,NH4VO3, or sodium metavanadate, NaVO3. The addition of some sulphuric acidkeeps these solutions stable.
1. Hydrogen sulphide. No precipitate is produced in acidic solution, but a bluesolution (due to the production of vanadium(IV) ions) is formed and sulphurseparates:
2 VO3− + H2S + 6 H+ → 2 VO2+ + S ↓ + 8 H2O
2. Zinc or aluminium acid solution. Zn and Al carry the reduction stillfurther than H2S. The solution turns atfirst blue (VO2+ ions),then green (V3+ ions) andfinally violet (V2+ ions).
VO2+/VO2+ = +0.999 V
V3+/ VO2+ = +0.359 V V2+/ V3+ = −0.256 V
Zn/ Zn2+ = −0.762 V
3. Ammonium sulphide solution: the solution is coloured claret-red, due to theformation of thiovanadates (VS4
3−).
VO3− + 4 S2− + 3 H2O → VS4
3− + 6 OH−
Upon acidification of the solution, brown vanadium sulphide, V2S5, is precipitated,and the filtrate usually has a blue colour:
2 VS43− + 6 H+ → V2S5 ↓ + 3 H2S ↑
The precipitate is soluble in solutions of alkalis, alkali carbonates, and sulphides.
104
4. Hydrogen peroxide: A red coloration is produced when a few drops of hydrogenperoxide solution are added dropwise to an acid (15-20 per cent sulphuric acid)solution of a vanadate; excess hydrogen peroxide should be avoided.
The red colour is due to the formation of themono- and diperoxovanadium(V) ions,VO(O2)+ and VO(O2)2
−:
VO3− + 2 H+ + H2O2 → VO(O2)+ + 2 H2O
VO3− + 2 H2O2 → VO(O2)2
− + 2 H2O
If the solution is made alkaline and morehydrogen peroxide is added, the colourchanges to yellow, due to the formation ofdiperoxoorthovanadate(V) ions:
VO(O2)+ + H2O2 + 4 OH− → VO2(O2)23− + 3
H2OVO(O2)2
− + 2 OH− ↔ VO2(O2)23− + H2O
The reaction is reversible; on acidification thesolution again turns red.
��������������������
������������ 5+OH2
OH2
VO
O
O
OV
OH2
O
5+�������������������������� ����������
����������
O
O
O
OH2
5. Lead acetate solution: yellow precipitate of lead vanadate, turning white or paleyellow on standing; the precipitate is insoluble in dilute acetic acid but soluble indilute nitric acid.
6. Barium chloride solution: yellow precipitate of barium vanadate; soluble in dilute hydrochloric acid.
7. Copper sulphate solution: green precipitate with metavanadates.
8. Iron(III) chloride:
VO3− + 4 H+ + Fe2+ ↔ VO2+ + Fe3+ + 2 H2O
The reaction proceeds from left to right in acid solution and in the reverse direction inalkaline solution. (ε°(Fe2+/Fe3+)= +0.771 V)
105
Chromium (Group VIb) and Its Common Ions
Chromium is a silver-white, lustrous, hard, and brittle metal that takes a high polish(melting point: 1857 °C). Chromium is extremely resistant to ordinary corrosiveagents, which accounts for its extensive use as an electroplated protective coating.
Solubility in acids The metal, if it is passivated (probably due to a thin protective oxide layer), is
not soluble in mineral acids, but the metal is rather active when not passivated.Redox potentials: ε°(Cr/Cr2+)= −0.913 V; ε°(Cr/Cr3+)= −0.744 V.
The chromium metal is soluble in dilute or concentrated hydrochloric acid.If air is excluded, chromium(II) ions are formed:
Cr + 2 HCl → Cr2+ + 2 Cl− + H2 ↑
In the presence of atmospheric oxygenchromium(II) gets wholly oxidised to thetervalent state:
4 Cr2+ + O2 + 4 H+ → 4 Cr3+ + 2 H2O1 2 3 4 5 6 7 8 9
-0.7
-0.6
-0.5
-0.4
-0.3
-0.2
-0.1
0.0
[Cr2+]= [Cr3+] = 0.05 M
Cr2+/ Cr3+
H2/ H+
Redox potential (V)
pH
Even if the solution is protected from air, chromium(II) ions decompose at ratesvarying with acidity, by reducing water with liberation of hydrogen (ε°(Cr2+/Cr3+)=−0.407 V).Chromium(II) ions are stable only in neutral and pure solutions at the exclusion of air.
Dilute sulphuric acid attacks chromium slowly, with the formation ofhydrogen.In hot, concentrated sulphuric acid chromium dissolves readily, when chromium(III)ions and sulphur dioxide are formed:
2 Cr + 6 H2SO4 → 2 Cr3+ + 3 SO42− + 3 SO2 ↑ + 6 H2O
Both dilute and concentrated nitric acid render chromium passive, as doescold, concentrated sulphuric acid and aqua regia.
Principal cations and anions of chromium in aqueous solution(Cr2+ chromous)*
Cr3+ chromicCrO4
2− chromateCr2O7
2− dichromate* chromium(II) ions are rather unstable, as they are strong reducing agents. Atmospheric oxygen oxidises them readily to chromium(III) ions.
106
The green chromium(III) oxide, Cr2O3, and its hydrous form, Cr2O3.nH2O, areamphoteric, dissolving readily in acids and in concentrated alkali, but if ignited toostrongly Cr2O3 becomes inert toward both acid and base.
Chromate (CrO42−) and dichromate (Cr2O7
2−) ionsIn basic solutions above pH 7, the yellow chromate ion CrO4
2− is the mainspecies; between pH 1 and 6, HCrO4
− and the orange-red dichromate ion Cr2O72− are
in equilibrium; and at pH<0 the main species are H2CrO4 and HCr2O7−.
The equilibria are the following:
H2CrO4 ↔ HCrO4− + H+
HCrO4− ↔ CrO4
2− + H+ HCr2O7− ↔ Cr2O7
2− + H+
2 HCrO4− ↔ H2O + Cr2O7
2−
0 2 4 6 8 10 12 14
0.00
0.01
0.02
0.03
0.04
0.05 co (K2CrO4)= 0.05 M
H2CrO4
HCr2O7-
HCrO4-
CrO42-
Cr2O72-
C (mol/ l)
pH
The chromates of the alkali metals and of magnesium and calcium are solublein water.Strontium chromate is sparingly soluble in water, and most other metallic chromatesare insoluble.
Sodium, potassium, and ammonium dichromates are well known and they aresoluble in water.
Reactions of chromium(III) ions, Cr3+
Chromium(III) sulphide, Cr2S3, like aluminium sulphide, can be prepared onlyin dry, because it hydrolyses readily with water to form chromium(III) hydroxide andhydrogen sulphide.
107
Hydrated chromium(III) sulphate, nitrate, chloride, bromide, iodide, andacetate are soluble in water.Chromium oxide, hydroxide, phosphate, and anhydrous halogenides (fluoride,chloride, bromide, iodide) are hardly soluble or not soluble in water.
E.g.: Compound Solubility ( g / 100 ml H2O)20 °C:25 °C:
Cr2(SO4)3.18H2OCrCl3.6H2O
12058,5
CrCl3 -----
To study these reactions use a 0.1 M solution of chromium(III) chloride CrCl3or chromium(III) sulphate Cr2(SO4)3.
1. Ammonia solution: grey-green to grey-blue gelatinous precipitate ofchromium(III) hydroxide, slightly soluble in excess of the reagent in the coldforming a violet or pink solution containing complex hexammine-chromate(III)ion; upon boiling the solution, chromium(III) hydroxide is precipitated. Hence forcomplete precipitation of chromium as the hydroxide, it is essential that thesolution be boiling and excess aqueous ammonia solution be avoided.
Cr3+ + 3 NH3 + 3 H2O → Cr(OH)3 ↓ + 3 NH4+
Cr(OH)3 ↓ + 6 NH3 → [Cr(NH3)6]3+ + 3 OH−
2. Sodium hydroxide solution: precipitate of chromium(III) hydroxide:
Cr3+ + 3 OH− → Cr(OH)3 ↓
In excess reagent the precipitate dissolves readily, when tetrahydroxochromate(III)ions are formed:
Cr(OH)3 ↓ + OH− ↔ [Cr(OH)4]−
The solution is green. On adding hydrogen peroxide to the alkaline solution, a yellowsolution is obtained, owing to the oxidation of chromium(III) to chromate:
2 [Cr(OH)4]− + 3 H2O2 + 2 OH− → 2 CrO42− + 8 H2O
After decomposing the excess of hydrogen peroxide by boiling, chromate ions may beidentified in the solution by one of their characteristic reactions.
3. Sodium carbonate solution: precipitate of chromium(III) hydroxide:
2 Cr3+ + 3 CO32− + 3 H2O → 2 Cr(OH)3 ↓ + 3 CO2 ↑
4. Ammonium sulphide solution: precipitate of chromium(III) hydroxide:
2 Cr3+ + 3 S2− + 6 H2O → 2 Cr(OH)3 ↓ + 3 H2S ↑
108
5. Chromate test.Chromium(III) ions can be oxidised to chromate, and than chromate ions can
be identified on the basis of their characteristic reactions.Oxidation of chromium(III): adding an excess of sodium hydroxide to a
solution of chromium(III) salt followed by a few ml of hydrogen peroxide.
2 [Cr(OH)4]− + 3 H2O2 + 2 OH− → 2 CrO42− + 8 H2O
The excess of H2O2 can be removed by boiling the mixture for a few minutes.
Identification of chromium after oxidation to chromate: a. Barium chloride test. After acidifying the solution with acetic acid and adding barium chloridesolution, a yellow precipitate of barium chromate is formed:
Ba2+ + CrO42− → BaCrO4 ↓
b. Chromium pentoxide test. Acidifying the solution with dilute sulphuric acid,adding 2-3 ml of ether or amyl alcohol to the mixture andfinally adding some hydrogen peroxide, a blue colorationis formed, which can be extracted into the organic phaseby gently shaking. Chromium pentoxide is formed duringthe reaction:
Cr O O
OOO
CrO42− + 2 H+ + 2 H2O2 → CrO5 + 3 H2O
In aqueous solution the blue colour fades rapidly, because chromiumpentoxide decomposes to chromium(III) and oxygen.
Reactions of chromate (CrO42−) and dichromate (Cr2O7
2−) ionsThe chromates of metal ions are usually coloured solids, yielding yellow
solutions when dissolved in water. In the presence of dilute mineral acids chromatesare partially converted into dichromates; the latter yield orange-red aqueous solutions.
2 CrO42− + 2 H+ ↔ Cr2O7
2− + H2O
To study the reactions of chromates and dichromates use a 0.1 M solution ofpotassium chromate and dichromate, respectively.
1. Barium chloride solution: pale-yellow precipitate of barium chromate, solubilityproduct constant Ksp(BaCrO4)= 1.17x10−10:
CrO42− + Ba2+ → BaCrO4 ↓
The precipitate is insoluble in water, sodium hydroxide, and acetic acid, but soluble inmineral acids.
109
Dichromate ions produce the same precipitate, but as strong acid is formed,precipitation is only partial:
Cr2O72− + 2 Ba2+ + H2O ↔ 2 BaCrO4 ↓ + 2 H+
2. Silver nitrate solution: brownish-red precipitate of silver chromate with asolution of a chromate, Ksp(Ag2CrO4)= 1.12x10−12:
CrO42− + 2 Ag+ → Ag2CrO4 ↓
The precipitate is soluble in dilute nitric acid and in ammonia solution, but isinsoluble in acetic acid. Hydrochloric acid converts the precipitate into silverchloride, Ksp(AgCl)= 1.77x10−10:
2 Ag2CrO4 ↓ + 2 H+ → 4 Ag+ + Cr2O72− + H2O
Ag2CrO4 ↓ + 4 NH3 → 2 [Ag(NH3)2]+ + CrO42−
Ag2CrO4 ↓ + 2 Cl− → 2 AgCl ↓ + CrO42−
A reddish-brown precipitate of silver dichromate is formed with aconcentrated solution of a dichromate; this passes into the less soluble silver chromateon boiling with water:
Cr2O72− + 2 Ag+ → Ag2Cr2O7 ↓
Ag2Cr2O7 ↓ + H2O → Ag2CrO4 ↓ + CrO42− + 2 H+
3. Lead acetate solution: yellow precipitate of lead chromate, Ksp(PbCrO4)=1.77x10−14:
CrO42− + Pb2+ → PbCrO4 ↓
The precipitate is insoluble in acetic acid, but soluble in dilute nitric acid and sodiumhydroxide solution:
2 PbCrO4 ↓ + 2 H+ ↔ 2 Pb2+ + Cr2O72− + H2O
PbCrO4 ↓ + 4 OH− ↔ [Pb(OH)4]2− + CrO42−
4. Hydrogen peroxide. (chromium pentoxide test; see above)
5. Hydrogen sulphide: an acid solution of a chromate is reduced by this reagent to agreen solution of chromium(III) ions:
2 CrO42− + 3 H2S + 10 H+ → 2 Cr3+ + 3 S ↓ + 8 H2O
6. Potassium iodide solution: chromate is reduced into chromium(III) in thepresence of dilute mineral acids. Iodine formed in the reaction can be extractedwith e.g. CCl4, yielding a violet organic phase.
2 CrO42− + 6 I− + 16 H+ → 2 Cr3+ + 3 I2 + 8 H2O
110
7. Iron(II) sulphate: reduces chromates or dichromates in the presence of mineralacid smoothly:
CrO42− + 3 Fe2+ + 8 H+ → Cr3+ + 3 Fe3+ + 4 H2O
8. Concentrated hydrochloric acid: on heating a solid chromate or dichromate withconcentrated hydrochloric acid, chlorine is evolved, and a solution containingchromium(III) ions is produced:
2 K2CrO4 + 16 HCl → 2 Cr3+ + 3 Cl2 ↑ + 4 K+ + 10 Cl− + 8 H2OK2Cr2O7 + 14 HCl → 2 Cr3+ + 3 Cl2 ↑ + 2 K+ + 8 Cl− + 7 H2O
9. Concentrated sulphuric acid and a chloride: (see chromyl chloride test)
O
OCr
OH
OH
+ HCl
+ HCl
��������
��������
+ HOH
+ HOH
Cl
ClCr
O
O
Summarise the reactions of chromates.
Ag+ Ba2+ Pb2+
CrO42−
solubility inacetic acid
HNO3 soln.
NaOH soln.
NH3 soln.
111
Manganese (Group VIIb) and Its Common Ions
Manganese is grey-white, resembling iron, but is harder and very brittle (meltingpoint: 1244 °C). Manganese is roughly similar to Fe in its physical and chemicalproperties, the chief difference being that it is harder and more brittle. The metal isreactive, and slowly reacts even with cold water.
Solubility in water and acids Due to its highly negative electrode potential (ε°(Mn/Mn(OH)2)= −1.56 V),
manganese reacts with water (the reaction is slow in the cold, but fast if the water iswarm) forming manganese(II) hydroxide and hydrogen (there is no protective oxidelayer, like in the case of chromium):
Mn + 2 H2O → Mn(OH)2 ↓ + H2 ↑
Dilute mineral acids and also acetic acid dissolve the metal with theproduction of manganese(II) salts and hydrogen (ε°(Mn/Mn2+)= −1.185 V):
Mn + 2 H+ → Mn2+ + H2 ↑
With hot, concentrated sulphuric acid, sulphur dioxide is evolved:
Mn + H2SO4 → Mn2+ + SO42− + SO2 ↑ + 2 H2O
Principal ions of manganese in aqueous solutionsoxidation state cations anions
+2 Mn2+
+3 (Mn3+)*+4 (Mn4+)* (MnO4
4− or MnO32−)*
+5 (MnO43−)**
+6 MnO42− ***
+7 MnO4−
* manganese(III) and manganese(IV) cations, and manganate(IV) anion are unstablein aqueous solutions, they are easily reduced to manganese(II).** unstable in aqueous solution, disproportionates to Mn(VII) and Mn(IV).*** stable in alkaline solutions, but upon neutralisation a disproportination reactiontakes place:
3 MnO42− + 2 H2O → MnO2 ↓ + 2 MnO4
− + 4 OH−
112
Permanganate ions, MnO4−, can be reduced step by step, e.g. with perborate
solution, to study the colour of the different oxidation state of manganese:
4+6+7+����� ����� ����� �����
����������
����������
+0.564 V +0.27 VMnO4 MnO4 MnO4 MnO4
2 3 45+
Alkali permanganates, MeMnO4 (Me= metal ion), are stable compounds,producing violet coloured solutions. They are all strong oxidising agents, and aresoluble in water.
Manganese(II) forms an extensive series of salts with all common anions.Most are soluble in water, although the phosphate and carbonate are only slightly so.
Five oxides of manganese are known so far:MnO Mn2O3 MnO2 Mn2O7 Mn3O4green brown black reddish oil reddish-brown
MnO is a grey-green to dark green powder, insoluble in water, but soluble inmineral acids.
Mn2O3 is insoluble in water; it produces Mn(II) ions if treated with mineralacids. If hydrochloric acid or sulphuric acid is used, chlorine and oxygen are evolved,respectively:
Mn2O3 ↓ + 6 HCl → 2 Mn2+ + Cl2 ↑ + 4 Cl− + 3 H2O2 Mn2O3 ↓ + 4 H2SO4 → 4 Mn2+ + O2 ↑ + 4 SO4
2− + 4 H2O
MnO2 is inert to most acids except when heated, but it does not dissolve togive Mn(IV) in solution; instead it functions as an oxidising agent, the exact mannerof this depending on the acid. With concentrated hydrochloric or sulphuric acid,chlorine and oxygen gas are evolved, respectively, and manganese(II) ions areproduced:
MnO2 ↓ + 4 HCl → Mn2+ + Cl2 ↑ + 2 Cl− + 2 H2O2 MnO2 ↓ + 2 H2SO4 → 2 Mn2+ + O2 ↑ + 2 SO4
2− + 2 H2O
Mn2O7 is an explosive oil, which can be extracted into CCl4 in which it isreasonably stable and safe. It solidifies at 5 to 9 °C to red crystals.
Mn3O4 is a spinel, MnIIMn2IIIO4 (MnO.Mn2O3). It is insoluble in water, but
soluble in mineral acids.
113
Reactions of manganese(II) ions, Mn2+
Manganese(II) sulphate, nitrate, chloride, bromide, and iodide are soluble inwater. Other common inorganic manganese(II) compounds (e.g. phosphate andcarbonate) are hardly soluble or insoluble in water.
E.g.: Compound Solubility ( g / 100 ml H2O)0 °C:
25 °C:MnBr2MnCl2
127,372,3
40 °C:18 °C:
MnF2Mn(OH)2
0,660,0002
To study these reactions use a 0.1 M solution of manganese(II) chloride ormanganese(II) sulphate.
1. Sodium hydroxide solution: an initially white precipitate of manganese(II)hydroxide; solubility product constant, Ksp(Mn(OH)2, 25 °C)= 2.06x10−13:
Mn2+ + 2 OH− → Mn(OH)2 ↓
The precipitate is insoluble in excess reagent, but soluble in dilute acids.The precipitate rapidly oxidises on exposure to air, becoming brown, when hydratedmanganese dioxide, MnO2.yH2O, is formed (ε°(Mn(OH)2/MnO2)= −0.05 V):
2 Mn(OH)2 ↓ + O2 → 2 MnO2.H2O ↓
Hydrogen peroxide converts manganese(II) hydroxide rapidly into hydratedmanganese dioxide:
Mn(OH)2 ↓ + H2O2 → MnO2.H2O ↓ + H2O
2. Ammonia solution: partial precipitation of (initially) white manganese(II)hydroxide:
Mn2+ + 2 NH3 + 2 H2O ↔ Mn(OH)2 ↓ + 2 NH4+
The precipitate is soluble in ammonium salts, when the reaction proceeds towards theleft.
3. Ammonium sulphide solution: pink precipitate of manganese(II) sulphide, solubility product constant, Ksp(MnS, 25 °C)= 4.65x10−14:
Mn2+ + S2− → MnS ↓
The precipitate is readily soluble in mineral acids and even in acetic acid.
MnS ↓ + 2 H+ → Mn2+ + H2S ↑
114
4. Disodium hydrogen phosphate solution: in the presence of ammonia (orammonium ions), pink precipitate of manganese ammonium phosphate:
Mn2+ + NH3 + HPO42− → Mn(NH4)PO4 ↓
If ammonium salts are absent, manganese(II) phosphate is formed:
3 Mn2+ + 2 HPO42− → Mn3(PO4)2 ↓ + 2 H+
The precipitates are soluble in acids.
5. Sodium carbonate solution: rose precipitate of manganese(II) carbonate, solubility product constant, Ksp(MnCO3, 25 °C)= 2.24x10−11:
Mn2+ + CO32− → MnCO3 ↓
The precipitate is soluble in dilute mineral acids and even in acetic acid.
6. Lead dioxide and concentrated nitric acid. On boiling a dilute solution ofmanganese(II) ions, free from hydrochloric acid and chlorides, with lead dioxideand a little concentrated nitric acid, and then diluting somewhat and allowing thesuspended solid containing unattacked lead dioxide to settle, the liquid acquires aviolet-red (or purple) colour due to permanganic acid formed.
5 PbO2 + 2 Mn2+ + 4 H+ → 2 MnO4− + 5 Pb2+ + 2 H2O
Permanganates, MnO4−
To study these reactions use a 0.01 M solution of potassium permanganate,KMnO4.
1. Hydrogen peroxide. The addition of this reagent to a solution of potassiumpermanganate, acidified with dilute sulphuric acid, results in decolourisation andthe formation of pure but moist oxygen: ε°(Mn2+/MnO4
−)= +1.507 V ε°(H2O2/O2)= +0.695 V
2 MnO4− + 5 H2O2 + 6 H+ → 5 O2 ↑ + 2 Mn2+ + 8 H2O
2. Concentrated hydrochloric acid. All permanganates on boiling with concentratedhydrochloric acid evolve chlorine. ε°(Cl−/Cl2)= +1.358 V
2 MnO4− + 16 HCl → 5 Cl2 ↑ + 2 Mn2+ + 6 Cl− + 8 H2O
115
3. Hydrogen sulphide: in the presence of dilute sulphuric acid the solutiondecolourizes and sulphur is precipitated: ε°(H2S/S)= +0.142 V
2 MnO4− + 5 H2S + 6 H+ → 5 S ↓ + 2 Mn2+ + 8 H2O
4. Iron(II) sulphate solution: in the presence of sulphuric acid permanganate isreduced to manganese(II). ε°(Fe2+/Fe3+)= +0.771 V
MnO4− + 5 Fe2+ + 8 H+ → 5 Fe3+ + Mn2+ + 4 H2O
The solution becomes yellow because of the formation of iron(III) ions. The yellowcolour disappears if potassium fluoride is added; it forms colourless complex withiron(III).
5. Potassium iodide solution: reduces permanganate with the formation of iodine,in the presence of sulphuric acid.
2 MnO4− + 10 I− + 16 H+ → 5 I2 + 2 Mn2+ + 8 H2O
In alkaline solution the permanganate is decolourized, but manganese dioxide isprecipitated. In the presence of sodium hydroxide solution, potassium iodide isconverted into potassium iodate.
2 MnO4− + I− + H2O → 2 MnO2 ↓ + IO3
− + 2 OH−
6. Sodium hydroxide solution. Upon warming a concentrated solution of potassiumpermanganate with concentrated sodium hydroxide solution, a green solution ofpotassium manganate is produced and oxygen is evolved.
ε°(MnO42−/MnO4
−)= +0.564 V ε°(OH−/O2)= +0.401 V
4 MnO4− + 4 OH− → 4 MnO4
2− + O2 ↑ + 2 H2O
When the manganate solution is poured into a large volume of water or is acidifiedwith dilute sulphuric acid, the purple colour of the potassium permanganate isrestored, and manganese dioxide is precipitated. ε°(MnO4
2−/MnO2)=+0.60 V
3 MnO42− + 2 H2O → 2 MnO4
− + MnO2 ↓ + 4 OH−
116
Summarise the reactions of CrO42−, Cr2O7
2−, and MnO4− ions.
CrO42− Cr2O7
2− MnO4−
colour(alkali metal salt)
H2Sacid solution
H2O2acid solution
KIacid solution
FeSO4acid solution
cc HCl
117
The Group VIIIb Elements (Fe, Co, Ni) andTheir Principle Ions
Iron is a relatively abundant element in the universe. Its nuclei are verystable. The metal is the fourth most abundant element on earth, by weight, making upthe crust of the earth. The use of iron is prehistoric. Iron is a vital constituent of plantand animal life, and appears in haemoglobin.The chemically pure iron is a silver-white, tenacious, and ductile metal. It melts at1535 °C. The pure metal is very reactive chemically, and rapidly corrodes, especiallyin moist air or at elevated temperatures. Iron can be magnetised. It has four allotropicforms, from which only the α-form is magnetic.The commercial iron is rarely pure and usually contains small quantities of carbide,silicide, phosphide, and sulphide of iron, and some graphite. These contaminants playan important role in the strength of iron structures. Other additives such as nickel,chromium, vanadium, etc. are also used to produce alloy steels. Iron is the cheapestand most abundant, useful, and important of all metals.
Cobalt is a steel-grey, slightly magnetic, brittle, and hard metal, closelyresembling iron and nickel in appearance. Melting point is 1495 °C.
Nickel is silvery white and takes on a high polish. It is hard, malleable,ductile, very tenacious, somewhat ferromagnetic, and a fair conductor of heat andelectricity. It melts at 1453 °C. It is quite resistant to attack by air or water at ordinarytemperatures when compact.
Solubility in acids Dilute or concentrated hydrochloric acid and dilute sulphuric acid dissolve
iron, cobalt, and nickel, when iron(II), cobalt(II), and nickel(II) salts and hydrogengas are produced.
Fe + 2 H+ → Fe2+ + H2 ↑ ε°(Fe/Fe2+)= −0.447 VCo + 2 H+ → Co2+ + H2 ↑ ε°(Co/Co2+)= −0.28 VNi + 2 H+ → Ni2+ + H2 ↑ ε°(Ni/Ni2+)= −0.257 V
Hot, concentrated sulphuric acid yields iron(III), cobalt(II), and nickel(II) ionsand sulphur dioxide:
2 Fe + 3 H2SO4 + 6 H+ → 2 Fe3+ + 3 SO2 ↑ + 6 H2OCo + H2SO4 + 2 H+ → Co2+ + SO2 ↑ + 2 H2ONi + H2SO4 + 2 H+ → Ni2+ + SO2 ↑ + 2 H2O
Cold, concentrated nitric acid and sulphuric acid renders iron passive.
118
Cold, dilute nitric acid, yields iron(II) and ammonium ions:
4 Fe + 10 H+ + NO3− → 4 Fe2+ + NH4
+ + 3 H2O
Medium concentrated nitric acid, or hot, concentrated nitric acid dissolves ironwith the formation of nitrogen oxide gas and iron(III) ions:
Fe + HNO3 + 3 H+ → Fe3+ + NO ↑ + 2 H2O
Dilute nitric acid dissolve cobalt and nickel readily in cold:
3 Co + 2 HNO3 + 6 H+ → 3 Co2+ + 2 NO ↑ + 4 H2O3 Ni + 2 HNO3 + 6 H+ → 3 Ni2+ + 2 NO ↑ + 4 H2O
Like iron, cobalt and nickel does not dissolve in concentrated nitric acidbecause it is rendered passive by this reagent.
Principal cations of iron, cobalt, and nickelFe2+
Fe3+Co2+
(Co3+)*Ni2+
* Cobalt(III) ions are unstable in water, but their complexes are stable both in solution and indry form.
Oxides of iron, cobalt, and nickel FeO* black Fe2O3 red Fe3O4 red-brown
CoO olive-green Co2O3 brown-black Co3O4 black
NiO green Ni2O3 ** Ni3O4 ***
* FeO is not stable under 560 °C; disproportionates to Fe and Fe3O4. It is also easilyoxidised by air oxygen.** There is no good evidence for Ni2O3. The black NiO(OH), however, is well known.*** Ni3O4 is not known, only a NiII-NiIII hydroxide of stoichiometry Ni3O2(OH)4.
Iron(III) oxide, Fe2O3, is soluble in dilute acids at room temperature, but if heated toostrongly it is almost insoluble even in hot concentrated hydrochloric acid.Iron(II,III) oxide, Fe3O4=FeIIFe2
IIIO4, is very resistant to attack by acids and alkalis.
Cobalt(II) oxide, CoO, and cobalt(III) oxide, Co2O3, are soluble in mineral acids.Cobalt(II,III) oxide, Co3O4=CoIICo2
IIIO4, is very slightly soluble in mineral acids.
Nickel(II) oxide, NiO, dissolves readily in acids.
119
Reactions of iron(II) ions, Fe2+ Solubility of the most common iron(II) compoundsThe chloride, bromide, iodide, chlorate, perchlorate, nitrate, sulphate, and
acetate are soluble in water.The fluoride is very slightly soluble, and oxide, carbonate, sulphide, and phosphateare practically insoluble.
E.g.: Compound Solubility ( g / 100 ml H2O)10 °C:10 °C:
FeCl2FeBr2
64,4109
18 °C:25 °C:
Fe(OH)2FeCO3
0,000150,0067
To study these reactions use a 0.1 M solution of iron(II) sulphate or iron(II)ammonium sulphate (Mohr’s salt).
1. Sodium hydroxide solution: white precipitate of iron(II) hydroxide in thecomplete absence of air, solubility product constant Ksp(Fe(OH)2, 25°C)=4.87x10−17, insoluble in excess of the reagent, but soluble in acids.
Upon exposure to air, iron(II) hydroxide is rapidly oxidised, yielding ultimatelyreddish-brown iron(III) hydroxide, Ksp(Fe(OH)3, 25°C)= 2.64x10−39 .Under ordinary conditions it appears as a dirty-green precipitate; the addition ofhydrogen peroxide immediately oxidises it to iron(III) hydroxide.
Fe2+ + 2 OH− → Fe(OH)2 ↓4 Fe(OH)2 ↓ + 2 H2O + O2 → 4 Fe(OH)3 ↓2 Fe(OH)2 ↓ + H2O2 → 2 Fe(OH)3 ↓
2. Ammonia solution: precipitation of iron(II) hydroxide occurs.
Fe2+ + 2 NH3 + 2 H2O → Fe(OH)2 ↓ + 2 NH4+
If, however, larger amounts of ammonium ions are present, precipitation does notoccur.
3. Hydrogen sulphide: no precipitation takes place in acid solution since thesulphide ion concentration is insufficient to exceed the solubility product ofiron(II) sulphide.
4. Ammonium sulphide solution: black precipitate of iron(II) sulphide FeS, Ksp(FeS, 25°C)= 1.59x10−19:
Fe2+ + S2− → FeS ↓
FeS is readily soluble in acids with evolution of hydrogen sulphide.The moist precipitate becomes brown upon exposure to air, due to its oxidation tobasic iron(III) sulphate Fe2O(SO4)2.
120
FeS ↓ + 2 H+ → Fe2+ + H2S ↑4 FeS ↓ + 9 O2 → 2 Fe2O(SO4)2 ↓
5. Potassium hexacyanoferrate(II) solution: in the complete absence of air a whiteprecipitate of potassium iron(II) hexacianoferrate(II) is formed:
Fe2+ + 2 K+ + [Fe(CN)6]4− → K2Fe[Fe(CN)6] ↓
under ordinary atmospheric conditions a pale-blue precipitate is obtained.
6. Potassium hexacyanoferrate(III) solution: a dark-blue precipitate is obtained.First hexacyanoferrate(II) and iron(III) is formed from hexacyanoferrate(III) andiron(II) ions according to the following equilibrium,
Fe2+ + [FeIII(CN)6]3− ↔ Fe3+ + [FeII(CN)6]4−
and than these ions combine to a precipitate called Turnbull’s blue:
4 Fe3+ + 3 [Fe(CN)6]4− → Fe4[Fe(CN)6]3 ↓
The structure of Turnbull's blue isbased on a three-dimensional cubicframework with FeII and FeIII atoms atthe corners of a cube and with FeII−N-C−FeIII links.
Note, that the composition of thisprecipitate is identical to that ofPrussian blue (see below at the reactionof iron(III) ions).
III
IIFe IIIFe IIFe
FeII IIIFe
Fe IIFeIIIFeII
Fe
FeII
Fe
III
Fe
II
Fe
Fe
IIIII
Fe
FeFeIIIFe II
III
Fe
IIFe IIIFe IIFe
Fe IIIII Fe
Fe IIFeIIIFeII
III
III
CN-CN-
CN-
CN-
CN-
CN-
CN-
CN-CN-
CN-
CN-CN-
CN-
CN-CN-
CN- CN-
CN-
CN-
CN-
CN-CN-
CN-
CN-
CN-
CN- CN-
CN-
CN-CN-
CN-
CN-CN-
CN-
CN- CN-CN-
CN-CN-
CN-
CN-
CN-
CN-
CN-
CN- CN-
CN-
CN-
CN-
CN-
CN-CN-
CN-
7. Ammonium thiocyanate solution: no coloration is obtained with pure iron(II)salts (distinction from iron(III) ions).
8. Dimethylglyoxime reagent: soluble red iron(II) dimethylglyoxime in ammoniacalsolution. Iron(III) salts give no coloration, but nickel, cobalt, and large quantities ofcopper salts interfere and must be absent.
121
Reactions of iron(III) ions, Fe3+ Solubility of the most common iron(III) compoundsThe chloride, bromide, iodide, chlorate, perchlorate, nitrate and sulphate are
soluble in water.The fluoride is very slightly soluble, and oxide, carbonate, sulphide, and phosphateare practically insoluble.
E.g.: Compound Solubility ( g / 100 ml H2O)0 °C:0 °C:
FeCl3Fe(NO3)3.6H2O
74,4150
Fe2O3 -----
To study these reactions use a 0.1 M solution of iron(III) chloride FeCl3.
1. Sodium hydroxide solution: reddish-brown, gelatinous precipitate of iron(III)hydroxide, Ksp(Fe(OH)3, 25°C)= 2.64x10−39, insoluble in excess of the reagent,but soluble in acids.
Fe3+ + 3 OH− → Fe(OH)3 ↓
Iron(III) hydroxide can be converted on strong heating to iron(III) oxide; the heatedoxide is soluble with difficulty in dilute acids, but dissolves on vigorous boiling withconcentrated hydrochloric acid.
2 Fe(OH)3 ↓ → Fe2O3 + 3 H2OFe2O3 + 6 H+ → 2 Fe3+ + 3 H2O
2. Ammonia solution: reddish-brown, gelatinous precipitate of iron(III) hydroxide,insoluble in excess of the reagent, but soluble in acids.
Fe3+ + 3 NH3 + 3 H2O → Fe(OH)3 ↓ + 3 NH4+
The solubility product of iron(III) hydroxide is so small (2.64x10−39) that completeprecipitation takes place even in the presence of ammonium salts.
3. Hydrogen sulphide: in acidic solution reduces iron(III) ions to iron(II) andsulphur is formed as a milky-white precipitate:
ε°(Fe2+/Fe3+)= +0.771 Vε°(H2S/ S)= +0.142 V
2 Fe3+ + H2S → 2 Fe2+ + 2 H+ + S ↓
The finely distributed sulphur cannot be readily filtered with ordinary filter papers.By boiling the solution with a few torn pieces of filter paper the precipitate coagulatesand can be filtered.
4. Ammonium sulphide solution: black precipitate, consisting of iron(II) sulphideand sulphur is formed:
2 Fe3+ + 3 S2− → 2 FeS ↓ + S ↓
122
The black iron(II) sulphide precipitate dissolves in hydrochloric acid and the whitecolour of sulphur becomes visible:
FeS ↓ + 2 H+ → Fe2+ + H2S ↑
From alkaline solutions black iron(III) sulphide is obtained:
2 Fe3+ + 3 S2− → Fe2S3 ↓
On acidification with hydrochloric acid iron(III) ions are reduced to iron(II) andsulphur is formed:
Fe2S3 ↓ + 4 H+ → 2 Fe2+ + 2 H2S ↑ + S ↓
6. Potassium hexacyanoferrate(II) solution: intense blue precipitate of iron(III)hexacyanoferrate(II) (Prussian blue):
4 Fe3+ + 3 [Fe(CN)6]4− → Fe4[Fe(CN)6]3
The precipitate is insoluble in dilute acids, but decomposes with concentratedhydrochloric acid. A large excess of the reagent dissolves it partly or entirely, whenan intense blue solution is obtained.
7. Potassium hexacyanoferrate(III) solution: a brown coloration is produced, dueto the formation of an undissociated complex, iron(III) hexacyanoferrate(III):
Fe3+ + [Fe(CN)6]3− → Fe[Fe(CN)6]
Upon adding some tin(II) chloride solution, the hexacyanoferrate(III) part of thecompound is reduced and Prussian blue is precipitated.
8. Ammonium thiocyanate solution: in slightly acidic solution a deep-redcolouration is produced (difference from iron(II) ions), due to the formation of anon-dissociated iron(III) thiocyanate:
Fe3+ + 3 SCN− → Fe(SCN)3
This molecule can be extracted by ether or amyl alcohol.Fluorides and phosphates bleach the colour because of the formation of the morestable hexafluoro and triphosphato complexes:
Fe(SCN)3 + 6 F− → [FeF6]3− + 3 SCN−
Fe(SCN)3 + 3 PO43− → [Fe(PO4)3]6− + 3 SCN−
Upon addition of SnCl2 solution in excess, the red colour disappears due to thereduction of iron(III) to iron(II):
2 Fe(SCN)3 + Sn2+ → 2 Fe2+ + Sn4+ + 3 SCN−
123
9. Disodium hydrogen phosphate solution: a yellowish-white precipitate ofiron(III) phosphate is formed, Ksp(FePO4.2H2O, 25°C)= 9.92x10−29:
Fe3+ + HPO42− → FePO4 ↓ + H+
Summarise the reactions of iron cations with hexacyanoferrates.
[Fe(CN)6]4− [Fe(CN)6]3−
Fe2+
Fe3+
Reactions of cobalt(II) ions, Co2+ Solubility of the most common cobalt(II) compoundsThe chloride, bromide, iodide, chlorate, perchlorate, nitrate, sulphate, and
acetate are soluble in water.The fluoride is very slightly soluble, and oxide, hidroxide, carbonate, sulphide, andphosphate are practically insoluble.
E.g.: Compound Solubility ( g / 100 ml H2O)7 °C:
20 °C:CoCl2CoSO4
4536,2
25 °C: CoF2 1,5Co(OH)2 0,00032
To study these reactions use a 0.1 M solution of cobalt(II) chloride orcobalt(II) nitrate.
1. Sodium hydroxyde solution: in cold a blue basic salt is precipitated:
Co2+ + OH− + NO3− → Co(OH)NO3 ↓
Upon warming with excess alkali (or sometimes merely upon addition of excessreagent) the basic salt is converted into pink cobalt(II) hydroxide precipitateKsp(Co(OH)2, 25°C)= 1.09x10−15:
124
Co(OH)NO3 ↓ + OH− → Co(OH)2 ↓ + NO3−
The hydroxide is slowly transformed into the brownish black cobalt(III) hydroxide onexposure to the air:
4 Co(OH)2 ↓ + O2 +2 H2O → 4 Co(OH)3 ↓
Cobalt(II) hydroxide precipitate is readily soluble in ammonia or concentratedsolutions of ammonium salts.
2. Ammonia solution: in the absence of ammonium salts small amounts ofammonia precipitate the basic salt:
Co2+ + NH3 + H2O + NO3− → Co(OH)NO3 ↓ + NH4
+
The excess of the reagent dissolves the precipitate, when hexamminocobaltate(II) ionsare formed:
Co(OH)NO3 ↓ + 6 NH3 → [Co(NH3)6]2+ + NO3− + OH−
The precipitation of the basic salt does not take place at all if larger amounts ofammonium ions are present, but the complex is formed in one step.
3. Ammonium sulphide solution: black precipitate of cobalt(II) sulphide fromneutral or alkaline solution:
Co2+ + S2− → CoS ↓
The precipitate is insoluble in hydrochloric or acetic acids.Hot, concentrated nitric acid or aqua regia dissolve the precipitate, when whitesulphur remaines:
3 CoS ↓ + 2 HNO3 + 6 H+ → 3 Co2+ + 3 S ↓ + 2 NO ↑ + 4 H2OCoS ↓ + HNO3 + 3 HCl → Co2+ + S ↓ + NOCl ↑ + 2 Cl− + 2 H2O
On longer heating the mixture becomes clear because sulphur gets oxidised tosulphate.
The CoS precipitate dissolves also in the 1+1 mixture of concentrated aceticacid and 30% hydrogen peroxide:
CoS ↓ + 4 H2O2 → Co2+ + SO42− + 4 H2O
4. Potassium nitrite solution: yellow precipitate of potassium hexanitrito-cobaltate(III), K3[Co(NO2)6].3H2O:
Co2+ + 7 NO2− + 2 H+ + 3 K+ → K3[Co(NO2)6] ↓ + NO ↑ + H2O
125
The test can be carried out most conviniently as follows: to a neutral solution ofcobalt(II) add acetic acid, than a freshly prepared saturated solution of potassiumnitrite.
5. Ammonium thiocyanate test: adding a few crystals of ammonium thiocyanate toa neutral or acidic solution of cobalt(II) a blue colour appears owing to theformation of tetrathiocyanatocobaltate(II) ions:
Co2+ + 4 SCN− → [Co(SCN)4]2−
If amyl alcohol or ether is added the free acid H2[Co(SCN)4] is formed and dissolvedby the organic solvent.
2 H+ + [Co(SCN)4]2− ↔ H2[Co(SCN)4]
The test is rendered more sensitive if the solution is acidified with concentratedhydrochloric acid, when the equilibrium shifts towards the formation of the free acid.
Reactions of nickel(II) ions, Ni2+ Solubility of the most common nickel compoundsThe chloride, bromide, iodide, chlorate, perchlorate, nitrate, sulphate, and
acetate are soluble in water.The fluoride is very slightly soluble, and oxide, carbonate, sulphide, and phosphateare practically insoluble.
E.g.: Compound Solubility ( g / 100 ml H2O)20 °C:0 °C:
NiCl2NiI2
64,2124,2
25 °C: NiF2 425 °C: NiCO3 0,0093
To study these reactions use a 0.1 M solution of nickel(II) sulphate ornickel(II) chloride.
1. Sodium hydroxyde solution: green precipitate of nickel(II) hydroxide, solubilityproduct constant, Ksp(Ni(OH)2, 25 °C)= 5.47x10−16:
Ni2+ + 2 OH− → Ni(OH)2 ↓
The precipitate is insoluble in excess reagent.Ammonia solution dissolves the precipitate; in the presence of excess alkalihydroxide ammonium salts also dissolve the precipitate:
126
Ni(OH)2 ↓ + 6 NH3 → [Ni(NH3)6]2+ + 2 OH−
Ni(OH)2 ↓ + 6 NH4+ + 4 OH− → [Ni(NH3)6]2+ + 6 H2O
The green nickel(II) hydroxide precipitate can be oxidised to black nickel(III)hydroxide with sodium hypochlorite solution:
2 Ni(OH)2 ↓ + ClO− + H2O → 2 Ni(OH)3 ↓ + Cl−
Hydrogen peroxide solution, however, does not oxidise nickel(II) hydroxide, but theprecipitate catalyses the decomposition of hydrogen peroxide to oxygen and waterwithout any other visible change.
2. Ammonia solution: green precipitate of nickel(II) hydroxide:
Ni2+ + 2 NH3 + 2 H2O → Ni(OH)2 ↓ + 2 NH4+
which dissolves in excess reagent:
Ni(OH)2 ↓ + 6 NH3 → [Ni(NH3)6]2+ + 2 OH−
the solution turns deep blue. If ammonium salts are present, no precipitation occurs,but the complex is formed immediately.
3. Ammonium sulphide solution: black precipitate of nickel sulphide, fromneutral or slightly alkaline solutions, Ksp(NiS, 25 °C)= 1.07x10−21:
Ni2+ + S2− → NiS ↓
If the reagent is added in excess, a dark-brown colloidal solution is formed which runsthrough the filter paper. If the colloidal solution is boiled, the colloidal solution(hydrosol) is coagulated and can than be filtered.The precipitate is insoluble in hydrochloric or acetic acids.Hot, concentrated nitric acid or aqua regia dissolve the precipitate with the separationof white sulphur:
3 NiS ↓ + 2 HNO3 + 6 H+ → 3 Ni2+ + 3 S ↓ + 2 NO ↑ + 4 H2ONiS ↓ + HNO3 + 3 HCl → Ni2+ + S ↓ + NOCl ↑ + 2 Cl− + 2 H2O
On longer heating the mixture becomes clear because sulphur gets oxidised tosulphate.
S ↓ + 2 HNO3 → SO42− + 2 H+ + 2 NO ↑
S ↓ + 3 HNO3 + 9 HCl → SO42− + 6 Cl− + 3 NOCl ↑ + 8 H+ + 2 H2O
The NiS precipitate dissolves also in the 1+1 mixture of concentrated aceticacid and 30% hydrogen peroxide:
NiS ↓ + 4 H2O2 → Ni2+ + SO42− + 4 H2O
127
4. Potassium nitrite solution: no precipitate is produced in the presence of aceticacid (difference from cobalt).
5. Potassium cyanide solution: green precipitate of nickel(II) cyanide:
Ni2+ + 2 CN− → Ni(CN)2 ↓
The precipitate is readily soluble in excess reagent, when a yellow solution appearsowing to the formation of tetracyanonickelate(II) complex ions:
Ni(CN)2 ↓ + 2 CN− → [Ni(CN)4]2−
6. Dimethylglyoxime reagent: red precipitate of nickel dimethylglyoxime fromsolutions just alkaline with ammonia or acid solutions buffered with sodium acetate:
CH3 C N OH
CH3 C N OHNi
N N
N N
C
C
C
C
OH
O
O OH
H3C
H3C
CH3
CH3
����+ 2 H +Ni + 2+ 2
Iron(II) (red colouration), bismuth (yellow precipitate), and larger amount ofcobalt (brown colouration) interfere in ammoniakal solution.
Compare the characteristic reactions of copper(II) and nickel(II):
NH3 soln.in excess
H2S KCN NaOH flametest
dimethyl-glyoxime
Cu2+
Ni2+
128
Summarise the solubility of sulphides:
MnS FeS CoS NiS
colour
HCl
hot, cc. HNO3
aqua regia
acetic acid +H2O2
Summarise the reactions of various metal ions with NaOH and NH3 solutions.
Mn2+ Fe2+ Fe3+ Co2+ Ni2+ Al3+ Cr3+ Zn2+
NaOH
NaOHin excess
NH3
NH3in excess
Summarise the solubility of selected metals in cold concentrated and 1+1 dilutednitric acid:
Ti V Cr Fe Co Ni
dilutedHNO3
cc HNO3
129
Classification of the Cations and Anions
Having become familiar with the characteristic reactions of cations and
anions, one may be able to identify an unknown material using simple chemical tests
and separations. In this process, called inorganic qualitative analysis, one deals with
the detection and identification of the elements that are present in a sample of
material. Frequently this is accomplished by making an aqueous solution of the
sample and then determining which cations and anions are present on the basis of
chemical and physical properties.
If a sample contains only a single cation and anion, their identification is a fairly
simple and straightforward process, although to distinguish between two cations (or
anions) that have similar chemical properties is not easy and in this instance
additional confirmatory tests are required. The detection of a particular ion in a
sample that contains several ions is somewhat more difficult, because the presence of
the other ions may interfere with the test. This problem can be circumvented by
precipitating, thereby removing, the disturbing ions from solution prior to testing for
the particular ion. The successful analysis of a mixture containing large number of
ions centres upon the systematic separation of the ions into groups containing only a
few ion. It is much simpler task to work with 2 or 3 ions than with 10 or more.
Ultimately, the separation of cations depends upon the difference in their tendencies
to form precipitates, or to form complex ions.
One of the best and well known separation scheme, outlined below, was first
described by R. Fresenius in 1841, though in the course of time modifications were
introduced. In this system cations are classified into five groups on the basis of their
behaviour against some reagents, called group-reagents. By the systematic use of
these reagents one can decide about the presence or absence of groups of cations, and
can also separate these groups for further examination. The group reagents used for
the classification of most common cations are hydrochloric acid, hydrogen sulphide,
ammonium sulphide, and ammonium carbonate. Classification is based on whether a
cation reacts with these reagents by the formation of precipitates or not. It can
130
therefore be said that classification of the most common cations is based on the
differences of solubilities of their chlorides, sulphides, and carbonates.
The five groups of cations and the characteristics of these groups are as
follows:
Group 1 Cations of this group formprecipitates with dilute hydrochloricacid. Ions of this group are lead(II),mercury(I), and silver(I).
Group 2 The cations of this groupdo not react with hydrochloric acid,but form precipitates with hydrogensulphide in dilute mineral acidmedium. Ions of this group aremercury(II), copper(II),bismuth(III), cadmium (II), tin(II),tin(IV), arsenic(III), arsenic(V),antimony(III), and antimony(V).The first four form the sub-group 2/aand the last six the sub-group 2/b.While sulphides of cations in Group2/a are insoluble in ammoniumpolysulphide, those of cations inGroup 2/b are soluble.
Group 3 Cations of this group donot react either with dilutehydrochloric acid, or with hydrogensulphide in dilute mineral acidmedium. However they formprecipitates with ammonium sulphidein neutral or ammoniacal medium.Cations of this group are iron(II),iron(III), cobalt(II), nickel(II),manganese(II), chromium(III),aluminium(III), and zinc(II).
Group 1AgCl, Hg2Cl2, PbCl2 (s)
DecantateGroup 2, 3, 4, and 5
DecantateGroup 3, 4, and 5
DecantateGroup 4 and 5
DecantateGroup 5
(NH4)2CO3
(NH4)2S
H2S
HCl
Group 2HgS, CuS, Bi2S3, CdS, (PbS) As2S3, As2S5, Sb2S3, Sb2S5, SnS, SnS2 (s)
Group 3Cr(OH)3, MnS, FeS, CoS, NiS, Al(OH)3, ZnS (s)
Group 4CaCO3, SrCO3, BaCO3 (s)
Ag+, Hg2 +, Pb +, Hg +, Cu +, Bi +, Cd +, 2 2 2 232
As +, As +, Sb +, Sb +, Sn +, Sn +,
Cr +, Mn +, Fe +, Fe +, Co +, Ni +, Al +, Zn +,
Ca +, Sr +, Ba +, Mg +, Li+, Na+, K+, NH4+. ��
22
2
33
4
22
22 22
5 53 3
3 2
.Mg +, Li+, Na+, K+, NH4+ 2
Group 4 Cations of this group do not react with the reagents of Groups 1, 2, and 3.They form precipitates with ammonium carbonate in the presence of ammoniumchloride in neutral medium. Cations of this group are calcium(II), strontium(II), andbarium(II).
Group 5 Common cations, which do not react with reagents of the previous groups,form the last group of cations, which includes magnesium(II), lithium(I),sodium(I), potassium(I), and ammonium(I) ions.
131
The methods available for the detection of anions are not as systematic as
those which have been described above for cations. No really satisfactory scheme has
yet been proposed which permits the separation of the common anions into major
groups, and the subsequent unequivocal separation of each group into its independent
constituents; however, it is possible to detect anions individually in most cases, after
perhaps a 1-2 stage separation. It is advantageous to remove all heavy metals from the
sample by extracting the anions through boiling with sodium carbonate solution;
heavy metal ions are precipitated out in the form of carbonates, while the anions
remain in solution accompanied by sodium ions.
The following scheme of classification of anions has been found to work well
in practice; anions are divided into four groups on the basis of their reactions with
dilute hydrochloric acid and of the differences of solubilities of their barium and
silver salts.
The four groups of anions and the characteristics of these groups are as
follows:
Group 1 Visible change, gas evolution and/or formation of a precipitate, with dilutehydrochloric acid. Ions of this group are carbonate, silicate, sulphide, sulphite, andthiosulphate.
Group 2 The anions of this group do not react with hydrochloric acid, but formprecipitates with barium ions in neutral medium. Ions of this group are sulphate,phosphate, fluoride, and borate.
Group 3 Anions of this group do not react either with dilute hydrochloric acid, orwith barium ions in neutral medium. However, they form precipitates with silver ionsin dilute nitric acid medium. Anions of this group are chloride, bromide, iodide, andthiocyanate.
Group 4 Common anions, which do not react with reagents of the previous groups,form the last group of anions, which includes nitrite, nitrate and chlorate ions.
132
Classification of the anions
I.
gro
up
II
. gro
up
II
I. gr
oup
IV. g
roup
reac
ts w
ith H
Cl
gas l
iber
atio
n or
pre
cipi
tate
Ba2
+pH
= 7
AgN
O3
+ H
NO
3---
---
prec
ipita
te w
ithA
g+pr
ecip
itate
with
Ba2
+
no re
actio
n w
ith K
I or
I 2 so
ln.
CO
32-
SiO
32-C
O32-
SiO
32-
SO42-
PO43-
F-
BO
33-
Cl-
Br- I-
SCN
−
NO
3-
deco
lour
ise I 2
soln
.S2
-
SO32-
S 2O
32-
SO32-
S 2O
32-
liber
ate
I 2 fr
om K
I sol
n.(N
O2- )
NO
2-
ClO
3-
133
TESTING FOR A SINGLE CATION IN SOLUTION
The identification of a single cation in solution is a fairly simple andstraightforward process, although without a good identification scheme it may requireso many experiments as the number of potential cations, if we know at least onespecific reaction for each cation. In order to reduce the number of tests required forthe identification, it is important to develop a good identification scheme, whichreduces the number of potential cations step by step placing them into groups. Thereare several possibilities and anyone could develop his/her own identification scheme.The scheme you find below follows the classification of cations into groups, asdescribed in the Fresenius' system. In case of a solid sample it is assumed that thesample is soluble in water or dilute nitric acid.Once the cation is found, its presence should be verified by other, characteristicreactions.
(1) Group I cationsAdd to the solution an excess of dilute HCl. If there is no change, follow (2a).
A white precipitate may contain Pb2+, Hg22+ or Ag+.Filter and wash the precipitate and then add NH3 solution to the precipitate.If the precipitate
does not change: Pb2+ present turns black: Hg22+ present dissolves: Ag+ present
(2a) Group IIA cationsAcidify the solution and add H2S in excess. If there is no change follow (3).
A precipitate may result if Hg2+, Bi3+, Cu2+, Cd2+, As3+, As5+, Sb3+, Sb5+, Sn2+,Sn4+ were originally present. (Check the colour of the precipitate !)Filter the precipitate, wash with dilute HCl, and treat with an excess of (NH4)2Sx. Ifthe precipitate dissolves, follow (2b).
If the remaining precipitate isyellow: Cd2+ present
Take a fresh sample and add dilute NaOH. If the precipitate isblue: Cu2+ presentyellow: Hg2+ presentwhite: Bi3+ present
(2b) Group IIB cationsAdd dilute HCl to the (NH4)2Sx filtrate in excess, when the precipitate
reappears. Take a fresh sample, acidify, and precipitate the sulphide.Examine its colour:
brown precipitate: Sn2+ present
134
An orange precipitate indicates Sb. To identify its oxidation state, take a fresh sample,acidify with 1:1 HCl and add KI:
no colouration: Sb3+ presentbrown colouration Sb5+ present
A yellow precipitate indicates As or Sn4+. Add (NH4)2CO3 in excess.If the precipitate remains undissolved: Sn4+ present
Perform the luminescence test for identifying Sn4+ ions.
If the precipitate dissoves: As3+ or As5+ present.To identify the oxidation state of As present in the solution, take a fresh sample,acidify with 1:1 HCl and add KI:
no coloration: As3+ presentbrown coloration: As5+ present
(3) Group III cationsNeutralise the solution with NH3 solution and add (NH4)2S in excess. If there
is no change, follow (4). Examine the precipitate.
A green precipitate indicates Cr3+. To a fresh sample, add NaOH:green precipitate which dissolvesin an excess of the reagent: Cr3+ present
A pink (flesh-like) precipitate indicates Mn2+. To a fresh sample, add NaOH:white precipitate, whichturns darker on standing: Mn2+ present
A white precipitate may be caused by Al3+ or Zn2+. To a fresh sample add NH3, firstin moderate amounts, then in excess:
white precipitate, whichdissolves in excess NH3 soln.: Zn2+ presentwhite precipitate, which remainesunchanged if excess NH3 is added: Al3+ present
A black precipitate occurs if Co2+, Ni2+, Fe2+ or Fe3+ were present originally.Filter, wash and mix the precipitate with 1:1 HCl. The precipitate dissolves if Fe2+ orFe3+ were present, otherwise it remains unchanged.
To a fresh sample add NaOH in excess:green precipitate, turningdark on standing: Fe2+ presentdark brown precipitate: Fe3+ presentblue precipitate, turning pinkif excess NaOH is added: Co2+ presentgreen precipitate, which remainsunchanged on standing: Ni2+ present
135
(4) Group IV cationsTo the solution add (NH4)2CO3 in excess, in the presence of NH4Cl. If there
is no precipitation, follow (5).A white precipitate indicates the presence of Ba2+, Sr2+ or Ca2+.
To a fresh sample add a four fold (in volume) of saturated CaSO4 solution:immediate white precipitate: Ba2+ presenta white precipitate is slowly formed: Sr2+ presentno precipitation occurs: Ca2+ present
(5) Group V cations Heat a fresh sample gently with some dilute NaOH:
characteristic odour of ammonia: NH4+ present
Carry out a flame test with the original sample:red coloration: Li+ presentyellow coloration: Na+ presentpale violet coloration: K+ present
To the solution add NaOH in excess:white precipitate, which turns red by addinga few drops of titan yellow reagent: Mg2+ present
136
TESTING FOR A SINGLE ANION IN SOLUTION
The identification of a single anion in solution is a fairly simple andstraightforward process, and anyone could develop his/her own identification scheme.The scheme below follows the classification of anions into four groups, as describedon page 131. It is assumed that the heavy metals are removed from the solution. Incase of a solid sample it is assumed that the sample is soluble in water.Once the anion is found, its presence should be verified by other, characteristicreactions.
(1) Group I anionsAdd to the solution an excess of dilute HCl. If there is no change, follow (2).
If a white precipitate or/and gas liberation is observed, one of the following anionsmay present: CO3
2−, SiO32−, S2−, SO3
2−, S2O32−.
White, gelatinous precipitate without the liberation of any gas: SiO32− present
White precipitate with the liberation of SO2. The gas is tested with a filter papermoistened with potassium iodate and starch solution.
blue coloration: S2O32− present
No precipitate, only gas liberation is observed.Test the gas with filter paper moistened with lead acetate solution.
black coloration: S2− presentTest the gas with filter paper moistened with potassium iodate and starch
solution.blue coloration: SO3
2− presentIntroduce the gas into baryta or lime water:
white precipitate: CO32− present
(2) Group II anionsNeutralise the solution and add BaCl2 solution. If there is no change follow(3).
A white precipitate may result if SO42−, PO4
3−, F−, or BO33− was originally present.
Filter the precipitate, and add HCl solution.precipitate is not soluble: SO4
2− present
Add concentrated sulphuric acid to the precipitate or to the original solid sample, andwarm the test tube:
test-tube acquires a greasy appearance: F− present
137
Add concentrated sulphuric acid and ethanol to the precipitate or to the originalsample in a porcelain basin and ignite the alcohol:
green-edged flame: BO33− present
Take a fresh sample and add ammonium molybdate reagent.yellow, crystalline precipitate: PO4
3− present
(3) Group III anionsAcidify the solution with dilute nitric acid and add AgNO3 solution. If there is
no change, follow (4).
Examine the precipitate.A yellow and yellowish white precipitate indicates I− and Br−, respectively.To a fresh sample, add 1-2 ml carbon tetrachloride and chlorine water dropweise, andshake it intensively:
violet organic layer, which turnscolourless with excess chlorine water: I− presentreddish-brown organic layer, which turnsyellow with excess chlorine water: Br− present
A white precipitate may be caused by Cl− or SCN−.To a fresh sample add FeCl3 solution:
blood-red coloration: SCN− present
Apply the chromyl chloride test:positive test: Cl− present
(4) Group IV anionsOne of the following ions may present: NO2
−, NO3−, or ClO3
−.
Acidify a fresh sample with acetic acid, add sulphanilic acid and α-naphthylaminereagents:
red coloration: NO2− present
Acidify a fresh sample with acetic acid, add sulphanilic acid and α-naphthylaminereagents and zinc chips:
red coloration: NO3− present
Add zinc chips to a fresh sample, filter the solution after a couple of minutes, acidifywith dilute nitric or sulphuric acid and add AgNO3 solution to the filtrate:
white precipitate: ClO3− present
138
Separating and Identifying the Cations
The separation scheme outlined below was first described by R. Fresenuis.Separation of the cations can be done by following the procedure outlined on pages139-145. It is important to verify the presence of a given cation by further tests.In case of a solid sample it is assumed that it is soluble in water or dilute nitric acid.
Consequences of applying the Fresenius' system for the separation of cations:
1. Ammonium ions, NH4+, must be tested in the original sample.
During the course of analysis most of the group reagents are added in the form ofammonium compounds, thus by the time Group V is reached, a considerable amountof ammonium ions will be built up in the test solution.
2. The oxidation state of As, Sb, Sn, and Fe can not be decided during theanalysis. It may be tested in the original sample.
As As2S3 AsS4 As2S5 ��������
����������
����������
(NH4)2S2 HCl3-H2S3+
5+ H2S 3-HCl(NH4)2S2
����������
����
�����As As2S5 AsS4
Sb Sb2S5 SbS4 �����
����
�����(NH4)2S2HCl
3-H2S5+
3+ H2S 3- HCl(NH4)2S2����� ���������Sb Sb2S3 SbS4 Sb2S5
Sn SnS2 SnS3 ����
�����
�����(NH4)2S2HCl
2-H2S4+
2+ H2S 2- HCl(NH4)2S2���� ����������Sn SnS SnS3 SnS2
H2S3+ 2+����Fe Fe Fe FeS ����(NH4)2S2+
3+
(NH4)2S�����
Fe
As As ����������5+ 3+H2S
139
Group I
...........................................................................................................................……..Cations of this group form precipitates with dilute (2 M) HCl.
Precipitate: AgCl, PbCl2, Hg2Cl2 Filtrate: cations of groups II-V
The precipitate is separated, and washed with hydrochloric acidic water.Transfer the precipitate to a small beaker, and boil with 5-10 ml water. Filter hot.
Residue: AgCl, Hg2Cl2 Filtrate: PbCl2
Wash the ppt. with hot water.Add NH4OH to the precipitate and filter.
Residue: black precipitate
HgNH2Cl + Hg
Solution: Ag+ complex
[Ag(NH3)2]+
Cool a portion of thesolution: a white crystallineppt. of PbCl2 is obtained ifPb is present in any quantity.
Hg22+ Ag+ Pb2+
Additional reactions for identification:
Divide the filtrate into twoparts:a.) Acidify with diluteHNO3. White ppt. ofAgCl.b.) Add a few drops of KIsolution. Pale yellow ppt.of AgI.
Ag+ present.
Divide the hot filtrate intothree parts:a.) Add K2CrO4 solution.Yellow ppt. of PbCrO4,insoluble in dilute acetic acid.b.) Add KI solution. Yellowppt. of PbI2, soluble in hotwater to a colourless solution,which deposites brilliantyellow crystals upon cooling.c.) Add dilute H2SO4. Whiteppt. of PbSO4, soluble inammonium tartrate solution.
Pb2+ present.
140
Group II
............................................................................................................................…...Cations of this group form precipitates with H2S in dilute mineral acid medium.
(They do not react with hydrochloric acid.)
Precipitate: PbS, HgS, CuS, Bi2S3, CdS, As2S3, As2S5, Sb2S3, Sb2S5, SnS, SnS2Filtrate: cations of Groups III-V
Transfer the precipitate to a porcelain dish, add about 5 ml yellow (NH4)2Sxsolution, heat to 50-60 oC, and maintain at this temperature for 3-4 minutes withconstant stirring. Filter.
Residue: Group IIA Filtrate: Group IIBMay contain HgS, PbS, Bi2S3, CuS,and CdS. Wash with small volumes of dilute(1:100) ammonium sulphide.
May contain solutions of the thio-salts(NH4)2AsS4, (NH4)2SbS4, and(NH4)2SnS3.Just acidify by adding concentrated HCldropwise, and warm gently. A yellow or orange precipitate,which may contain As2S5, Sb2S5, andSnS2, indicates Group IIB present.
141
Group IIA
...............................................................................................................…................The precipitate may contain HgS, PbS, Bi2S3, CuS, and CdS.Transfer to a beaker or porcelain basin, add 5-10 ml dilute HNO3, boil gently for 2-3minutes, filter and wash with a little water.
Residue FiltrateBlack. HgS
Dissolve ina mixture ofM NaOCland diluteHCl solutiosAdd dilute
May contain nitrates of Pb, Bi, Cu, and Cd. Test a small portion forPb by adding dilute H2SO4. A white precipitate of PbSO4 indicatesPb present. If Pb present, add dilute H2SO4 to the remainder of thesolution, concentrate in the fume cupboard until white fumes (fromthe dissociation of H2SO4) appear.Cool, add 10 ml of water, stir, allow to stand 2-3 minutes, filter andwash with a little water.
HCl, boil off Residue Filtrateexcess Cl2and cool. orDissolve inaqua regia
White PbSO4
Pour 2 ml of10%ammonium
May contain nitrates and sulphates of Bi, Cu andCd. Add concentrated NH4OH solution untilsolution is distinctly alkaline. Filter.
(add also a acetate Residue Filtrate little solidKCl).Evaporatethe solvent,dissolve theresidue indilute HCl.Add SnCl2solution.White ppt.turning greyor black.
Hg2+present.
through theppt. on thefilter paperseveral times,add to thefiltrate a fewdrops of diluteacetic acid andthenK2CrO4solution.Yellow ppt. ofPbCrO4 .
Pb2+ present.
White: may beBi(OH)3.
Wash.Dissolve in theminimumvolume ofdilute HCl andpour into coldsodiumtetrahydroxo-stannate(II).Black ppt.
Bi3+ present.
May contain [Cu(NH3)4]2+ and[Cd(NH3)4]2+.
If deep blue in colour, Cu ispresent in quantity. Confirm Cuby acidifying a portion of thefiltrate with dilute acetic acid andadd K4[Fe(CN)6] solution.Reddish-brown ppt.Cu2+ present.--------------------------------To the remainder of the filtrate,add KCN solution dropwise untilthe colour disappears, and add afurther ml in excess. Pass H2Sfor 20-30 seconds.Yellow ppt., sometimesdiscoloured, of CdS.Cd2+ present.
Hg Pb Bi Cu Cd
142
Group IIB
.................................................................................................................................Treat the yellow ammonium polysulphide extract of the Group II ppt. with dilute HCl,with constant stirring, until it is slightly acidic, warm and shake or stir for 1-2minutes.A fine white (or yellow) ppt. is sulphur only.A yellow or orange flocculant ppt. indicates As, Sb, and/or Sn present.Filter and wash the ppt., which may contain As2S5, (As2S3), Sb2S5, (Sb2S3), SnS2 ,with a little H2S water.
Transfer the precipitate to a beaker or porcelain dish, add 5-10 ml dilute HCl(1:1) and boil gently for 5 minutes. Filter.
Residue: may contain(As2S3), As2S5, and S.
Dissolve the ppt. in warmdilute NH3 solution, add3% H2O2 solution
Filtrate: may contain Sb3+, or/and Sn4+
Boil to expel H2S, partially neutralize the solution, addiron wire or nail, warm on water bath for 10 minutes.After cooling, filter the liquid
and warm for a fewminutes. Precipitate: Sb (Fe) Filtrate: Sn2+ (Fe2+)
As Sb Sn
IdentificationAdd a few ml Mg(NO3)2reagent, stir and allow tostand. White precipitateof Mg(NH4)AsO4.Filter off ppt., and pourAgNO3 solutioncontaining a few drops of2M acetic acid on toresidue on filter.Brownish-red residue ofAg3AsO4. As5+ present
May also try Bettendorff'sreaction.
Dissolve in a mixture oftartaric acid and diluteHNO3, or in aqua regia.Render the solution lessacidic by adding NH3 soln(add NH3 soln. until whiteppt. appears). Pass H2S intothe solution.Orange precipitate ofSb2S3.
Sb3+ present.
Add the filtrate to a solutionof HgCl2.White precipitate of Hg2Cl2or grey ppt. of Hg.
Sn2+ present.
May also try luminescencetest.
143
Gro
up II
I
Neu
traliz
e th
e so
lutio
n w
ith N
H3
solu
tion
and
add
(NH
4)2S
in e
xces
s.Pr
ecip
itate
: CoS
, NiS
, FeS
, MnS
, Cr(
OH
) 3, A
l(OH
) 3, Z
nSW
ash
the
prec
ipita
ted
sulp
hide
s with
a sm
all a
mou
nt o
f dilu
te (N
H4)
2S so
lutio
n. T
rans
fer t
he p
reci
pita
te to
a sm
all b
eake
r, ad
d 1
M H
Cl,
stir
wel
l, al
low
to st
and
for 2
-3 m
inut
es a
nd fi
lter.
Res
idue
: C
oS,
NiS
Filtr
ate :
Fe2
+ , M
n2+ ,
Cr3
+ , A
l3+ ,
Zn2
+
Afte
r rem
ovin
g H
2S b
y bo
iling
, 30%
NaO
H a
nd 3
% H
2O2
(1:1
) are
add
ed to
the
soln
.
Dis
solv
e th
e pr
ecip
itate
in a
mix
ture
of C
H3C
OO
H a
nd a
few
dro
ps o
f 30%
H2O
2.R
esid
ue: F
e(O
H) 3
, MnO
(OH
) 2Fi
ltrat
e:
A
l(OH
) 4- ,
Zn(O
H) 4
- , C
rO42-
The
solu
tion
is b
oile
d w
ith so
lid N
H4C
lTh
e pr
ecip
itate
is a
lso
solu
ble
in d
ilute
nitr
ic a
cid
by h
eatin
g.Pr
ecip
itate
: A
l(OH
) 3Fi
ltrat
e:
Zn(N
H3)
62+
, CrO
42-
Co
N
i
Fe
M
n
A
l
Z
n
C
r
Iden
tific
atio
nD
ivid
e th
e so
lutio
n in
to tw
oeq
ual p
arts
.a)
Add
1 m
l am
yl a
lcoh
ol, 2
gso
lid N
H4S
CN
and
shak
e w
ell.
Am
yl a
lcoh
ol la
yer c
olou
red
blue
.Co
2+ p
rese
nt.
b) A
dd 2
ml N
H4C
l sol
n., N
H4O
Hso
ln. u
ntil
alka
line,
and
then
exce
ss o
f dim
ethy
l-gly
oxim
ere
agen
t. R
ed p
pt.
N
i2+ p
rese
nt.
Div
ide
the
prec
ipita
te in
to tw
opa
rts.
a) D
isso
lve
the
ppt.
in d
ilute
HC
l(f
ilter
, if n
eces
sary
).Ei
ther
- ad
d a
few
dro
ps o
f KSC
Nso
ln.
Dee
p re
d co
lora
tion.
Or -
add
K4
Fe(C
N) 6
soln
.B
lue
ppt.
Fe
3+ p
rese
nt.
b) D
issol
ve th
e pp
t. in
cc
HN
O3,
add
PbO
2 to
the
soln
., he
at, a
llow
to se
ttle.
Pur
ple
soln
. of M
nO4- .
Mn2+
pre
sent
.
Dis
solv
e th
e pr
ecip
itate
in a
cetic
acid
, and
add
mor
in re
agen
t.Ch
arac
teris
tic fl
uore
scen
ce.
A
l3+ p
rese
nt.
If th
e so
ln. i
s yel
low
, Cr3+
isin
dica
ted.
Div
ide
the
liqui
d in
totw
o po
rtion
s.a)
Aci
dify
with
ace
tic a
cid
and
add
lead
ace
tate
soln
.
Yel
low
ppt.
of P
bCrO
4.
Cr
3+ p
rese
nt.
(May
als
o try
chr
omiu
m p
ento
xide
test
.)b.
) Aci
dify
with
ace
tic a
cid
and
pass
H2S
. W
hite
pre
cipi
tate
of
ZnS.
Zn2+
pre
sent
.(T
ry a
lso
dith
izon
e te
st.)
144
Group IV
Treat the filtrate of the Group III ppt. with concentrated HCl; boil the solutionto remove sulphide, S2-. Neutralize the solution with NH3 soln., boil, and add(NH4)2CO3 in excess.
White precipitate may contain BaCO3, SrCO3, and CaCO3.
Wash ppt. with a little hot water and reject the washings. Dissolve the ppt. in 2 Macetic acid.
Add CH3COONa to soln., heat the solution, and add a slight excess of K2Cr2O7solution (i.e. until the solution assumes a slight yellow colour and precipitation iscomplete).
Precipitate: yellow BaCrO4------------------------------
Wash the ppt. with hotwater. Dissolve the ppt. ina little concentrated HCl,evaporate almost todryness and apply theflame test. Green (or yellowish-green) flame.
Ba2+ present.
Filtrate: Ca2+, Sr2+
---------------------------------------------------------------
Add a four fold (in volume) of saturated CaSO4solution:a white ppt. slowly formed: Sr2+ indicated.
Render the soln. alkali with NH3 soln., and add(NH4)2(COO)2: white precipitate: Ca2+ indicated.
Render the solution neutral, and then addK2[Fe(CN)6] solution and a little NH4Cl:white ppt., Ca2+ indicated.
If only one of the two ions present, apply the flametest.Crimson flame: Sr2+ present.Brick-red flame: Ca2+ present.
145
Group V
The residue from Group IV contains the Group V cations, Na+, K+, Mg2+,NH4+(added during the process of separation of different cations in the form of e.g.(NH4)2S and (NH4)2CO3. Do not test for NH4+ from this solution !!!!!!!), and avery small amount of unprecipitated Ca2+.
Test for Na+: use the flame test. Persistent yellow flame: Na+ present.
Test for Mg2+: add NaOH solution until alkaline. If white precipitate forms, divide itinto two portions, and add titan yellow and magnezon to the different portions of ppt.:red colouration of ppt. with titan yellow and blue colouration of ppt. with magnezon:Mg2+ present.
Test for K+: add NaOH and boil the solution until NH3 completely disappears (testNH3 with wet pH paper). Filter the solution if necessary. Add concentrated HClO4solution: white, crystalline precipitate (of KClO4): K+ present.
Test for NH4+: use the original solution. Add NaOH and boil the solution, test thegas evolved using wet indicator paper or filter paper moistened with mercury(I)nitrate.
