the group ia elements (li, na, k, rb, cs) and their principle ions
TRANSCRIPT
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
9
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+
12
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.
23
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
24
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