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Pasinszki tibor, 2002. oc

  1. 1. BUDAPEST UNIVERSITY OF TECHNOLOGY AND ECONOMICS FACULTY OF CHEMICAL ENGINEERING Dr. Tibor Pasinszki Inorganic Chemistry Laboratory Practice Department of Inorganic Chemistry 2002
  2. 2. Contents1. Introduction to the laboratory page 22. The Group Ia Elements (Li, Na, K, Rb, Cs) page 73. The Group IIa Elements (Be, Mg, Ca, Sr, Ba) page 124. The Group IIIa Elements (Boron and Aluminium) page 195. The Group IVa Elements (C, Si, Ge, Sn, Pb) page 246. The Group Va Elements (N, P, As, Sb, Bi) page 367. The Group VIa Elements (O, S, Se, Te) page 538. The Group VIIa Elements (F, Cl, Br, I) page 639. Pseudohalogens and Pseudohalides page 7310. The Group Ib Elements (Cu, Ag, Au) page 7811. The Group IIb Elements (Zn, Cd, Hg) page 8712. The Group IVb Elements (Titanium) page 10013. The Group Vb Elements (Vanadium) page 10214. The Group VIb Elements (Chromium) page 10515. The Group VIIb Elements (Manganese) page 11116. The Group VIIIb Elements (Fe, Co, Ni) page 11717. Classification of the Cations and Anions page 12918. Testing for a Single Cation in Solution page 13319. Testing for a Single Anion in Solution page 13620. Separating and Identifying the Cations page 138 1
  3. 3. Introduction to the laboratorySAFETY PROCEDURES1. 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 laboratory work 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 running water and call for the laboratory instructor. If chemicals are spilled on the clothes, remove them and flush the skin with water.5. Never smell a reaction mixture directly. Minimise your exposure to chemical vapours.6. In order to avoid cuts and lacerations, protect your hands with a towel when inserting either glass tubing or thermometers into stoppers or thermometer adapters. Fire-polish all glass tubing and stirring rods so that there are no sharp edges. Report any cuts to the lab instructor so that the injury may receive 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 conduct the 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 the first 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. 2
  4. 4. 11. Avoid fire. Most organic solvents are flammable. Play it safe and treat all organic materials 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, and fire blankets. Never attempt to extinguish an electrical fire with water. Use only extinguishers 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 be administered .12. Avoid explosions. Never pour water into concentrated sulphuric acid. Always add concentrated sulphuric 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. Small amounts of impurities that can be explosive will be concentrated to dangerous levels. Always know what your neighbours are doing, be prepared for any accident. 3
  5. 5. 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 a serious accident may result if the wrong chemical is used.2. Reagent bottles must be protected from contamination. You must therefore never 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 err and take more than you need, do not return the excess to the bottle. In other words, you only remove material from the reagent bottle, you never put anything into it.3. Never take the reagent bottles to the sink or to your desk. Put the bottles back to the reagent shelf after using them.4. Do not lay stoppers on the desk or shelf in such a way that they will become contaminated. Depending on the shape of the stopper, either hold it while the material is being removed or lay it on its flat top.5. Glass stoppers that are stuck can generally be loosened by gently tapping the stopper on the edge of the shelf.6. The reagent area must be kept clean. Be sure that you clean up any chemicals you 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 flushed down the sink with copious amounts of cold water. Check the directions for disposal of liquid wastes before using the sinks. 4
  7. 7. 6
  8. 8. The Group Ia Elements (Li, Na, K, Rb, Cs) and Their 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. 7
  9. 9. 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: 6p 4f 6s 5d 1. salt is evaporated 5p 4d 2. molecule decompose to its 5s 4p constituents; e.g. NaCl → Na + Cl 4s 3d 3. thermal excitation of valence shell 3p electron of the metal atom; 3s Na → Na* 2p 2s 4. instant relaxation of the excited states with ejecting photons: 1s Na* → Na + hνIf the photons ejected are in the visible Na (3p 3s)= 589 nm (yellow lightregion, coloration of the flame isobserved.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-yellow Potassium → violet (lilac) Rubidium → dark-red Caesium → blue 8
  10. 10. 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 ml water) at 18 °C: LiF 0,27 Li3PO4 0,039 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, in neutral solutions. + 3− 3 Li + PO4 → Li3PO4 ↓Precipitation is almost complete in the presence of sodium hydroxide solution.2. Sodium carbonate solution: white precipitate of lithium carbonate from concentrated solutions: + 2− 2 Li + CO3 → Li2CO3 ↓3. Ammonium carbonate solution: white precipitate of lithium carbonate from concentrated solutions. + 2− 2 Li + CO3 → 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: + 2− − NH4 + CO3 ↔ NH3 + HCO34. Ammonium fluoride solution: a white, gelatinous precipitate of lithium fluoride is slowly formed in ammoniacal solution. + − Li + F → LiF ↓5. Flame test: carmine-red colour. 9
  11. 11. +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 sodium magnesium uranyl acetate NaMg(UO2)3(CH3COO)9.9H2O from concentrated solutions: + 2+ 2+ − Na + Mg + 3 UO2 + 9 CH3COO → NaMg(UO2)3(CH3COO)92. Uranyl zinc acetate solution: yellow, crystalline precipitate of sodium zinc uranyl acetate NaZn(UO2)3(CH3COO)9.9H2O : + 2+ 2+ − Na + Zn + 3 UO2 + 9 CH3COO → NaZn(UO2)3(CH3COO)93. 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 a1 M solution of potassium chloride. Remember, the sizes of K+ and NH4+ ions are almostidentical, thus their reactions in general are very similar.1. Perchloric acid solution (HClO4): white crystalline precipitate of potassium perchlorate KClO4 from not too dilute solutions. You may use concentrated HClO4 solution. (This reaction is unaffected by the presence of ammonium salts.) + − K + ClO4 → KClO4 ↓2. Tartaric acid solution (or sodium hydrogen tartrate solution): white crystalline precipitate of potassium hydrogen tartrate: + + K + H2C4H4O6 → KHC4H4O6 ↓ + HThe solution should be buffered with sodium acetate. The precipitate is slightly soluble inwater (3.26 g/l). (Ammonium salts yield a similar precipitate.) 10
  12. 12. 3. Sodium hexanitritocobaltate(III) solution, Na3[Co(NO2)6]: yellow precipitate of potassium hexanitritocobaltate(III): + 3− 3 K + [Co(NO2)6] → 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: + + 3− 2 K + Na + [Co(NO2)6] → 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+: 2− 3− − − − 2− CO3 PO4 F Cl NO3 SO4 + Li + Na + K 11
  13. 13. The Group IIa Elements (Be, Mg, Ca, Sr, Ba) and Their 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 oxideBarium 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 theboiling 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. 2+ − Me + 2 H2O → Me + 2 OH + H2 ↑ Be, Mg, Ca, Sr, and Ba dissolve readily in dilute acids (unless water insoluble saltforms): + 2+ Me + 2 H → Me + H2 ↑ Concentrated nitric acid renders beryllium passive (like aluminium). 12
  14. 14. 2+Characteristic reactions of magnesium ions, Mg 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 white precipitate of basic magnesium carbonate: 2+ 2− − 5 Mg + 6 CO3 + 7 H2O → 4 MgCO3.Mg(OH)2.5 H2O ↓ + 2 HCO3In 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): + 2− − NH4 + CO3 ↔ NH3 + HCO32. Sodium carbonate solution: white, voluminous precipitate of basic magnesium carbonate: 2+ 2− − 5 Mg + 6 CO3 + 7 H2O → 4 MgCO3.Mg(OH)2.5 H2O ↓ + 2 HCO33. Ammonium hydroxide solution: partial precipitation of white, gelatinous magnesium hydroxide, solubility product constant: Ksp(25°C)= 5.61x10−12: 2+ + Mg + 2 NH3 + 2 H2O → Mg(OH)2 ↓ + 2 NH4The precipitate is readily soluble in solutions of ammonium salts.4. Sodium hydroxide solution: white precipitate of magnesium hydroxide: 2+ − Mg + 2 OH → Mg(OH)2 ↓5. Disodium hydrogen phosphate solution: a white flocculant precipitate of magnesium hydrogen phosphate is produced in neutral solutions: 2+ 2− Mg + HPO4 → MgHPO4 ↓White crystalline precipitate of magnesium ammonium phosphate MgNH4PO4.6H2O inthe presence of ammonium chloride (to prevent precipitation of magnesium hydroxide)and ammonia solutions: 2+ 2− Mg + NH3 + HPO4 → MgNH4PO4 ↓The precipitate is soluble in acetic acid and in mineral acids. 13
  15. 15. 6. Titan yellow reagent and magneson reagent: Titan yellow and magneson (I and II) O2N N N OHare water soluble dyestuffs. They areabsorbed by magnesium hydroxide HOproducing a deep-red colour with titanyellow and a blue colour with magneson. Magneson IPour a little amount of the test solutioninto two test tubes, add 1-2 drops of the O2N N N OHtitan 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. Magneson IICalcium, Ca2+ Calcium chloride and nitrate are readily soluble in water. Calcium oxide (similarly to strontium and barium oxides) readily reacts withwater producing heat and forming the hydroxide. Calcium sulphide (and also other alkaline earth sulphides) can be prepared only inthe 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 acids even in acetic acid): 2+ 2− Ca + CO3 → CaCO3 ↓2. Dilute sulphuric acid: white precipitate of calcium sulphate, solubility product constant: Ksp(25°C)= 7.10x10−5: 2+ 2− Ca + SO4 → CaSO4 ↓3. Ammonium oxalate solution: white precipitate of calcium oxalate, solubility product constant: Ksp(CaC2O4.H2O, 25°C)= 2.34x10−9 (insoluble in acetic acid, but soluble in mineral acids): 2+ 2− Ca + (COO)2 → Ca(COO)2 ↓ 14
  16. 16. 4. Disodium hydrogen phosphate solution: white precipitate of calcium hydrogen phosphate is produced from neutral solutions: 2+ 2− Ca + HPO4 → CaHPO4 ↓5. Potassium hexacyanoferrate(II) solution: white precipitate of a mixed salt: 2+ + 4− Ca + 2 K + [Fe(CN)6] → 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. 2+Strontium, Sr 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): 2+ 2− Sr + CO3 → SrCO3 ↓2. Dilute sulphuric acid: white precipitate of strontium sulphate, Ksp(SrSO4, 25°C)= 3.44x10−7: 2+ 2− Sr + SO4 → SrSO4 ↓3. Saturated calcium sulphate solution: white precipitate of strontium sulphate, formed slowly in the cold, but more rapidly on boiling. 2+ 2− Sr + SO4 → SrSO4 ↓4. Ammonium oxalate solution: white precipitate of strontium oxalate: 2+ 2− Sr + (COO)2 → Sr(COO)2 ↓ 15
  17. 17. 5. Disodium hydrogen phosphate solution: white precipitate of strontium hydrogen phosphate is produced from neutral solutions: 2+ 2− Sr + HPO4 → SrHPO4 ↓6. Potassium chromate solution: yellow precipitate of strontium chromate: 2+ 2− Sr + CrO4 → 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 theCrO42− ion concentration sufficiently to prevent the precipitation of SrCrO4.The equilibria are the following: c= 0. M ( 2Cr 4) 1 K O Cr ] m /) [ O 42- ( oll 0. 10 − +H2CrO4 ↔ HCrO4 + H 0. 08 − 2− +HCrO4 ↔ CrO4 + H 0. 06 0. 04 − 2− 2 HCrO4 ↔ H2O + Cr2O7 0. 02 − 2− +HCr2O7 ↔ Cr2O7 + H 0. 00 0 2 4 6 8 10 pH7. Flame test: crimson-red colour to the Bunsen flame. 2+Barium, Ba 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): 2+ 2− Ba + CO3 → BaCO3 ↓ 16
  18. 18. 2. Dilute sulphuric acid: white, finely divided precipitate of barium sulphate, Ksp(BaSO4, 25°C)= 1.07x10−10: 2+ 2− Ba + SO4 → BaSO4 ↓3. Saturated calcium sulphate solution: immediate white precipitate of barium sulphate.4. Saturated strontium sulphate solution: white precipitate of barium sulphate.5. Ammonium oxalate solution: white precipitate of barium oxalate (readily dissolved by hot dilute acetic acid and by mineral acids): 2+ 2− Ba + (COO)2 → Ba(COO)2 ↓6. Disodium hydrogen phosphate solution: white precipitate of barium hydrogen phosphate is produced from neutral solutions: 2+ 2− Ba + HPO4 → BaHPO4 ↓6. Potassium chromate solution: yellow precipitate of barium chromate, practically insoluble in water, Ksp(BaCrO4, 25°C)= 1.17x10−10: 2+ 2− Ba + CrO4 → 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 product −5 −7 −10 7.10x10 3.44x10 1.07x10 constant: Ksp 2− SO4 concentration 17
  19. 19. 2+ 2+ 2+ 2+Summarise the reactions of Mg , Ca , Sr , and Ba ions: 2+ 2+ 2+ 2+ Mg Ca Sr Ba NH3 soln. NaOH Na2CO3 (NH4)2CO3 + NH4Cl (NH4)2CO3 Na2HPO4(NH4)2(COO)2 add hot acetic acid to the precipitate K2CrO4 1. add acetic acid add acetic acid to to the precipitate neutral soln. the precipitate 2. add mineral acid to the prec. dilute H2SO4satd. CaSO4 soln.satd. SrSO4 soln. Flame test 18
  20. 20. The Group IIIa Elements: Boron and Aluminium, and Their 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 alkali Boron 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 H2O + − 2 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 H2OConcentrated nitric acid renders aluminium passive, but dilute nitric acid dissolves themetal: Al + 4 HNO3 → Al3+ + 3 NO3− + NO ↑ + 2 H2OAluminium is soluble in alkali hydroxides when a solution of tetrahydroxoaluminate isformed: 2 Al + 2 OH- + 6 H2O → 2 [Al(OH)4]− + 3 H2 ↑ 19
  21. 21. Reducing power of aluminium Aluminium is a very reactive metal, in particular towards electronegativepartners, 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 FeThe 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 → Al2I6The 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 ml water) at 15 °C AlCl3 69,9 25 °C AlF3 0,559 α-Al2O3 0,000098 ----- AlPO4Aluminium 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. 20
  22. 22. 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 CH3COOH6. Sodium alizarin sulphonate (Alizarin-S) reagent: O OH C OH red precipitate in ammoniacal solution, which isfairly stable to dilute acetic acid. CAdd to the solution of Al3+ ions, dilute ammonia SO 3 Na Osolution and 2-3 drops of the solution of the reagent, and Alizarin-Sthen acidify it with acetic acid. HO OH7. Morin reagent: add little solid sodium acetate and 1-2 drops of the reagent to the solution of Al3+ ions. HO OInvestigate the characteristic green fluorescence of the OHsolution in UV light. OH O Morin 21
  23. 23. Oxides of Boron and Aluminium B2O3 Al2O3B2O3 is a white , hygroscopic solid. α−Al2O3 is very hard and resistant toIt is acidic, reacting with water to give hydration and attack by acids.boric acid, B(OH)3. γ−Al2O3 readily takes up water and dissolves in acids.Boric acid and borate ions in aqueous solution Boric 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= 6x10−10In aqueous, concentrated borate solutions polymeric ions are also present, due to the −polymerisation 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 H2O − − 2 B(OH)3 + B(OH)4 ↔ B3O3(OH)4 + 3 H2O − 2− 2 B(OH)3 + 2 B(OH)4 ↔ B4O5(OH)4 + 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 − − 2−B5O6(OH)4 , B3O3(OH)4 , and B4O5(OH)4 are also present. − − 2−The species B5O6(OH)4 , B3O3(OH)4 , and B4O5(OH)4 are formed successively withincreasing 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. 22
  24. 24. 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+ + H3BO3 2 AgBO2 ↓ + 3 H2O → Ag2O + 2 H3BO33. 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 23
  25. 25. The Group IVa Elements (C, Si, Ge, Sn, Pb) and Their Principle IonsCarbon 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 alkali Carbon is very unreactive at normal conditions, but the reactivity of IVa groupelements 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 HFand 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 H2O Si + 2 KOH + H2O → K2SiO3 + 2 H2 24
  26. 26. 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 H2OIn hot, concentrated sulphuric acid and in aqua regia tin(IV) ions are formed atdissolution: Sn + 4 H2SO4 → Sn4+ + 2 SO42− + 2 SO2 ↑ + 4 H2O 3 Sn + 4 HNO3 + 12 HCl → 3 Sn4+ + 12 Cl− + 4 NO ↑ + 8 H2OTin 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 H2O 2 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. 25
  27. 27. Principle oxides of IVa group elements C Si Ge Sn Pb * CO, CO2 SiO2 GeO2 SnO red(β) PbO yellow (red)*colourless gases white solid white solid SnO2 white PbO2 black* white SnO.xH2O and white PbO.xH2O precipitates from aqueous solutionsOnly 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− SiO32− * GeO32− * Sn2+ Pb2+ HCO3− Sn4+* Does not exist in this form in aqueous solution. 26
  28. 28. 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 ↑ + H2Othe gas can be identified by its property ofrendering lime water or baryta water turbid: CO2 + Ca2+ + 2 OH- → CaCO3 ↓ + H2O CO2 + Ba2+ + 2 OH- → BaCO3 ↓ + H2OAny 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. 27
  29. 29. 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 ↓ + CO2Hydrogen 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. 28
  30. 30. 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 NH33. 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 ↓ 29
  31. 31. 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 H2OAdd 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 solutionstetrahydroxo-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− 30
  32. 32. 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 bynascent 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 solutionshexahydroxostannate(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: 31
  33. 33. 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) 2+ 4+Summarise the redox reaction of Sn and Sn : 2+ 3+ Hg Zn Fe Fe 2+ Sn 4+ SnStandard electrode potentials at 25 °C: 2+ 4+ 2+ 2+ Sn / Sn : +0.151 V Hg2 :/ Hg : +0.920 V 2+ Hg/ Hg2 : +0.7973 V 2+ 2+ 3+ Sn/ Sn : −0.1375 V Fe / Fe : +0.771 V 2+ Fe/ Fe : −0.447 V 2+ Zn/ Zn : −0.7618 V 32
  34. 34. 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 H2OIf 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. 33
  35. 35. 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− + H2O PbCrO4 ↓ + 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. 34
  36. 36. 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. NH NH NH N N N 2+ 2 S C + Pb S C Pb C S + 2 H+ N N N N N HNThe 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 35
  37. 37. The Group Va Elements (N, P, As, Sb, Bi) and Their Principal Anions and CationsNitrogen (N2) is a colourless, inert diatomic gas (boiling point: -196 °C). Nitrogenoccurs in Nature mainly as dinitrogen (N2) that comprises 78% by volume of theearths 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 alkali Nitrogen is little soluble in water, but P, As, Sb, and Bi are not soluble. Phosphorus, arsenic, antimony, and bismuth are not affected by nonoxidizingacids (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 H2O 3 As + 5 HNO3 (conc) + 2 H2O → 3 AsO43- + 5 NO ↑ + 9 H+ 2 As + 5 OCl- + 3 H2O → 2 AsO43- + 5 Cl- + 6 H+ 36
  38. 38. 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 H2O Bi + HNO3 + 3 HCl → Bi3+ + 3 Cl- + NO ↑ + 2 H2O 2 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 NaH2PO2Trihydrides of the Group Va elements (XH3) The gases XH3 can be obtained by treating ammonium salts with alkalies, bytreating 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- + H2ONitric acid of concentration below 2 M has little oxidising power. 37
  39. 39. 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 elements N P As Sb BiNO colourless P4O10 (P2O5) As4O6 (As2O3) Sb4O6 (Sb2O3) Bi2O3 NO2 red white white white yellow very hygroscopicPrincipal anions and cations of N, P, As, Sb, and Bi: N P As Sb Bi *NH4+ ammonium (As3+) Sb3+ Bi3+ *NO2+ nitronium (As5+) Sb5+ H2PO2- SbO+ BiO+ − NO2- nitrite hypophosphite AsO33- arsenite [Sb(OH)4] NO3- nitrate HPO32- phosphite AsO43- arsenate [Sb(OH)6]− PO43- phosphate* They exist in aqueous solution in the form of arsenite and arsenate.Characteristic reactions of NH4+, NO2-, NO3-, PO43-, AsO33-, AsO43-,Sb3+, and Bi3+ ionsAmmonium 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. 38
  40. 40. 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 ↑ + H2OAmmonia gas may be identified a.) by its odour (cautiously smell the vapour after removing the test-tube fromthe flame); b.) by the formation of white fumes of ammonium chloride when a glass rodmoistened 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 solutionblack (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. Nesslers reagent (alkaline solution of potassium tetraiodomercurate(II)): brown precipitate or brown or yellow coloration is produced according to theamount 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 H2OThe 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 butslightly 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). 39
  41. 41. + +Compare the characteristic reactions of K and NH4 ions: Flame test Tartaric acid 3− Nessler`s cc HClO4 [Co(NO2)6] reagent + K + NH4Nitrites, 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+ → HNO2 3 HNO2 → HNO3 + 2 NO ↑ + H2O 2 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 40
  42. 42. 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 H2O6. 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 H2O7. 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- + H2O8. 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 H2O9. 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: NH2 + + NO2 + H+ HSO 3 N N NH 2 + 2 H2 O SO H NH 2 3The test solution must be very dilute, otherwise the reaction does not go beyond thediazotation stage. 41
  43. 43. 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+ 3+ 2− 2 NO3 + 4 H2SO4 + 6 Fe → 6 Fe + 2 NO ↑ + 4 SO4 + 4 H2O 2+ 2+ Fe + NO ↑ → [Fe(NO)] 42
  44. 44. 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: + 3− Ag3PO4 ↓ + 6 NH3 → 3 [Ag(NH3)2] + PO4 + − + Ag3PO4 ↓ + 2 H → H2PO4 + 3 Ag3. 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. 43
  45. 45. 5. Iron(III) chloride solution: yellowish-white precipitate of iron(III) phosphate: 2− 3+ + HPO4 + Fe → FePO4 ↓ + HThe 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 H2OReactions of arsenic(III) ions Arsenic(III) compounds can be derived from the amphoteric arsenic trioxideAs2O3. 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: 3− + 2 AsO3 + 6 H + 3 H2S → As2S3 ↓ + 6 H2OThe 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− 2− + 3 As2S3 ↓ + 28 HNO3 + 4 H2O → 6 AsO4 + 9 SO4 + 36 H + 28 NO ↑ − 3− 3− As2S3 ↓ + 6 OH → AsO3 + AsS3 + 3 H2OAmmonium sulphide and ammonium polysulphide also dissolves the precipitate, 3− 3−when thioarsenite (AsS3 ) and thioarsenate (AsS4 ) ions are formed, respectively: 2− 3− As2S3 ↓ + 3 S → 2 AsS3 2− 3− 2− As2S3 ↓ + 4 S2 → 2 AsS4 + S3 44
  46. 46. 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: 3− + 2 AsS3 + 6 H → As2S3 ↓ + 3 H2S ↑ 3− + 2 AsS4 + 6 H → As2S5 ↓ + 3 H2S ↑ 2− + S2 + 2 H → H2S ↑ + S ↓2. Silver nitrate: yellow precipitate of silver arsenite in neutral solution (distinctionfrom arsenates): 3− + AsO3 + 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: 3− 2+ + AsO3 + Cu + 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: 3− − 3− − + AsO3 + I3 + H2O ↔ AsO4 + 3I +2HThe reaction is reversible, and an equilibrium is reached.6. Bettendorffs 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: 3− + 2+ 4+ 2 AsO3 + 12 H + 3 Sn → 2 As ↓ + 3 Sn + 6 H2O7. Marshs test.This test is based upon the fact that all soluble compounds of arsenic are reduced bynascent 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-blackmirror just beyond the heated part of the tube, particularly if it is cooled.The deposit is soluble in sodium hypochlorite (distinction from antimony). 45
  47. 47. 3− +AsO3 + 3 Zn + 9 H → 2+ AsH3 ↑ + 3 Zn + 3 H2O4 AsH3 ↑ → heat → 4 As ↓ + 6 H2 ↑ −2 As + 5 OCl + 3 H2O → 3− − + 2 AsO4 + 5 Cl + 6 HReactions of arsenate ions, AsO43- Arsenic(V) compounds are derived from arsenic pentoxide, As2O5. This is theanhydride 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. 3− 3− AsO4 + H2S → AsO3 + S ↓ + H2O 3− + 2 AsO3 + 6 H + 3 H2S → As2S3 ↓ + 6 H2OIf a large excess of concentrated hydrochloric acid is present and hydrogen sulphide ispassed rapidly into the cold solution, yellow arsenic pentasulphide is precipitated: 3− + 2 AsO4 + 5 H2S + 6 H → As2S5 ↓ + 8 H2OArsenic pentasulphide, like the trisulphide, is readily soluble in alkali hydroxides orammonia, ammonium sulphide, ammonium polisulphide, sodium or ammoniumcarbonate: − 3− 3− As2S5 ↓ + 6 OH → AsS4 + AsO3S + 3 H2O 2− 3− As2S5 ↓ + 3 S → 2 AsS4 2− 3− 2− As2S5 ↓ + 6 S2 → 2 AsS4 + 3 S3 2− 3− 3− As2S5 ↓ + 3 CO3 → AsS4 + AsO3S + 3 CO2Upon acidifying these solutions with hydrochloric acid, arsenic pentasulphide isreprecipitated: 3− + 2 AsS4 + 6 H → As2S5 ↓ + 3 H2S ↑ 46
  48. 48. 2. Silver nitrate solution: brownish-red precipitate of silver arsenate, Ksp(Ag3AsO4,25°C)= 1.03x10−22, from neutral solutions: 3− + AsO4 + 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: 3− 2+ + AsO4 + Mg + NH4 → MgNH4AsO4 ↓Upon treating the white precipitate with silver nitrate solution containing a few dropsof acetic acid, red silver arsenate is formed: + 2+ + MgNH4AsO4 ↓ + 3 Ag → Ag3AsO4 ↓ + Mg + NH44. 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: 3− 2− + + AsO4 + 12 MoO4 + 3 NH4 + 24 H → (NH4)3[As(Mo3O10)4] ↓ + 12 H2OThe 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. 3− − + 3− AsO4 + 2 I + 2 H ↔ AsO3 + I2 ↓ + H2OThe reaction is reversible.Redox systems: 0. 7 ental V) I−/I2 Redox pot i ( 0. 6 H3AsO3/H3AsO4 - I/I2 0. 5Concentrations: 0. 4 H 3AsO 3/H3AsO 4 [I−]= 0.1 M; [H3AsO4]= 0.1 M aa) [H3AsO3]= 0.001 M 0. 3 bb) [H3AsO3]= 0.01 M cc) [H3AsO3]= 0.1 M 0. 2 0 1 2 3 4 5 pH 47