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The following diagrams illustrate the method: Les cartes, planches, tableaux, etc.. peuvent dtre film^s d des taux de reduction diff^rents. Lorsque le document est trop grand pour dtre reproduit en un seul cliche, il est film^ d partir de Tangle sup6rieur gauche, de gauche 6 droite, et de haut en bas, en prenant le nombre d'images n6cessaire. Les diagrammes suivants illustrent la m^thode. errata to pelure, )n it □ 32X t a » 1 2 3 4 5 6 ADVANCED CIIEMI8TKY FOR HIGH SCHOOLS. BY W. S. KLLIS, ]^.A., H.Sc, Collegiate iHililntc^ Kingston. TORONTO : THE COPP, CLAKK COMPANY, LIMITED, 1896, I Entered according to Act of the Parliament of Canada, in the year one thousand eight hundred and ninety-six, by Tiik Cop.', Clark Company. Limitkd, Toronto. Ontario, in the Ottlce of the Minister of Agriculture. INTRODUCTION. This book is intended to cover the practical work in Chemistry prescribed for the Honor Matriculation and Senior Leaving Examinations. The followin«r is the list of subjects : — Chemical Theory. The practical study of the following ele- ments, with their most characteristic compounds, in illustration of Mendelejeffs Chissirication of the Elements : Hydrouen ; Sodium, Potassium ; Magnesium, Zinc : Calcium, .Strontium, liariimi ; Boron, Aluminium ; Carbon, Silicon, Tin, Lead; Nitrogen, Phos- phorus, Arsenic, Antimony, IJismuth ; Oxygen, Sulphur ; Fluorine, Chlorine, liromine. Iodine ; Manganese, Iron. Elementary Quali- tative Analysis. Any study of the elements leading to Mendelejeff's classification is necessarily a comparative one, hence that view of the subject has been kept prominent throughout the book. Chemical theor\- has been largely omitted, because teachers will liave to give viva voce instruction in it anyway, and what may be required beyond that students can get easily and inexpensivel}- in the reference books, which should be in the library of every school in which this class of work is being attempted. I iv INTHODUCTION. The following b(jok.s, among others, are desirable ones to have for reference : — koscoe and Schorlemmer's Treatise on Chemistry, Vols. I. and II. Macmillan & Co. Kemscn's Chemistry. Henry Holt & Co. I'.loxam's Chemistry. Henry C. Lea & Co. 'lilden's Chemical Philosophy. Lon>,nnans & Co. Muir's Elements of Thermal Chemistry. Macmillan & Co. Mixter's Elementary Text Hook of Chemistry. Wiley & Sons. Ramsay's System of Inorganic Chemistry. Churchill. Wurtz' Atomic Theory. Api)lct»)n cS: Co. Remsen's Theoretical Chemistry. Henry C. Lea & Co. Cooke's New Chemistry. Appleton & Co. Richter's Inorganic Chemistry. Blakiston, Son & Co. I'rescott and Jt)hnston's Qualitative Analysis. VanNostrand & Co. Meyer's Modern Theories of Chemistry. Longmans & Co. Muir and Slater's Elementary Chemistry. Cambridge University Press. Ramsay's Chemical Theory. Macmillan & Co. Dobbin and Walkei-'s Chemical Theory. Macmillan & Co. CONTENTS. niAITKH I. l'\' here to repeat the experiments, or to refer at length to the properties and compounds dealt with in the other book 1.-— Bromine. Experiments. 1. Heat tocrether manganese dioxide and hydrobromic acid, in a test-tube fitted with stopper and delivery tube Pass the outer end of the delivery tube into a loosely corked bottle containing a little water. Do not inhale bromine vapour, nor get it on the hands. 2. Heat bromide of potassium with sulphuric acid and manganese dioxide ; collect the vapour that distils over as in the last experiment. ' 2 C'lll^OIMNK fJUolfp. 3. 1 )mi) a littli: Ijioiniiic solution on starch pasti'. Does vapour of hroniinc [)nKlucc a similar clTcct ? The reactions in tlic forc^oin^^ cx[)crimc!its arc similar to the correspondiiiL; onivs in the case of chlorine. The ecjuations nia\' l)e written from the analoijous chlorine ones by chan^nnij the symbol C.'l to Hr. The followinir experinn-nt is to delLnnine the relative c(jmbininfT powers of chlorine and bromine. 4. Drop some chlorine water into a solution of h)'dro- bromic acid, or a colourless s(r b\- the greater chemical attraction between hxdrogen and chlorine than between hydrogen and bronn'ne. Il is evident then that hydro- bromic acid must be prej^ared b\' some method in whidi free sulphuric ac id will not be used. The way in which this is done is by allow'iig bronn'ne to drop very slowly on phosphorus in presence of water. 3]k+P=PHr,. PBr, + 3n,0 = 3H15r+H,PO;, The phosphorus tribromide decomposes the water to form hydrobromic and phosphorous acids. The former may be driven off in vapour by heat, and dissolved in cold water. 3. Pci.-is a current of sulphuretted hydrogen into a solution of bromine H,S + Br,+ R,0 = 2lIHr-flI.,0-f-S rilLOUFNK r.|{OUP. •'{. — Iodine. KXI'KklMHNTS. 1. Heat a mixture of an iodide with manganese di- oxide and suli)huric acid. The vapour that comes off may be collected in an empty bottle. Is it soluble ? 2. Heat some manganese dioxide with hydriodic acid Compare the corresponding reactions under chlorine and bromine. 3. Heat a little iodine in a test-tube, observe the vapour and its condensation. 4. Add a very small bit of iodine to some starch paste. 5. Place a shaving of phosphorus on a metal plate and lay on the phosphorus a little iodine. 6. Drop a bit of iodine into a test-tube half full of water. After shaking for some time notice the colour, then drop into the solution a crystal of iodide of potas- sium, and again shake. 7. Try the solubility of iodine in alcohol, chloroform, and bisulphide of carbon. 8. Make a thin mucilage by boiling starch in plenty of water. J'our a little of this into each of three test-tubes. To the first add a drop or two of a very weak aqueous solution of iodine, to the second and third a few drops of solution of an iodide ; then into the second pour a few drops of bromine solution, and into the third a little chlorine solution. I • 1 HYDFaoniC ACID. 4.— Hydriodic Acid. I EXI'KRIMENTS. 1. Add a shaving of phosphorus to some iodine solu- tion in a test-tube, heat, j^ass the gas that conies off into cold water, and test both the gas and its solution for acid properties, also test it for free iodine and for an iodide. Compare the reaction when hydrobromic acid was pre- pared. 2. Place a shaving of phosphorus on a metal plate, and heap some powdered iodine on it (use excess of iodine). Invert a beaker over the mass, and after chemical action has ceased, quickly slip a glass plate under the mouth of the beaker, and shake the fumes which it contains with a little water. Test, as in ex. i. ^ 3- Pass a current of hydrogen sulphide through a solu- tion of iodine until the latter is decolourized I„-f-HaS = 2HI-J-S. Test, as in ex. i. Allow a portion of this solution to stand exposed to air for a few days, then repeat the test. 4- To prepare hydriodic acid gas in quantity, put a httle red phosphorus in a tube and twelve times as much (by weight) powdered iodine ; heat genth-, until a dark coloured mass is formed. Let water fail, a drop at a time, on this mass. Pass the gas that comes off into cold water. Collect a tube full of the gas and heat it 5.— Notes on the Halogens. Fluorine.— When calcic fluoride fHuor spar) CaF is warmed with sulphuric acid it yields hydrofluoric acid - b CHLORINE <;i{OUP. a colourless gas, soluble in water and active chemically. It is chiefly of use because it attacks glass, forming silicon fluoride, S\V_^. The acid may be easily prepared by heating the fluor spar with strong sulphuric acid, and it may be condensed in a freezing mixture in a lead tube. Fluorine has been prepared by the electrolytic decom- position of anhydrous hydrofluoric acid, kept at a tem- perature at which it is in the liquid form. It is a colour- less gas that unites with gi^.^at energy with most elements. It decomposes water, setting free oxygen in the form of ozone. Phosphorus, antimoi^y, boron and sulphur take fire in it. With most metals it forms fluorides, and even gold and platinum are attacked by it when heated. It does not unite directly with oxygen, nitrogen or carbon ; and no compound of fluorine and oxygen has been obtained by reduction of other compounds. Hydrogen fluoride, HF, is a very weak acid, hence fluorides are not plentiful, that of calcium being the best known. Double fluorides, such as KFHF, BiF.,3HF are not uncommcMi. They generally consist of a metallic fluoride joined with that of hydrogen. 6.— Occurrence and Preparation. Chlorine is found generally distributed in nature in combination with other elements. Bromine, in the form of bromides, occurs in the waters of many mineral springs, in some salt beds, accompanying chlorides, and with, iodine in sea weeds. Iodine is found in sea weeds that are largely collected on the north-western shores of I OCCURHENCE AND PKKPARATION. f Europe. These plants are burned, the soluble salts dis- solved out of the ashes, and the iodine and bnMTiine separated by the use of sulphuric acid and manganese dioxide. The iodine passes off first and is collected on the sides of cold vessels. Afterwards, by similar treat- ment the bromine is obtained. Fluorine is found in mineral form as calcium fluoride (fluor spar), and cryo- lite, sodium fluoride. 7.— Compounds of the Halogens. The compounds of chlorine, bromine and iodine are similar in their properties, and are generally prepared by similar methods. The hydrogen compounds offer a par- tial exception to this last statement, because hydrobromic and hydriodic acids break up in presence of sulphuric acid to form free bromine or iodine, sulphur dioxide and hydrogen sulphide. When chlorine is passed into solution of caustic potash a mixture of the chloride and the chlorate of potassium is formed (If. S. Chem., page i6i). When bromine or iodine is passed into caustic potash solution a reaction occurs according to the following equation : 3M,-f6KfIO-5MK + KMO,-f 3ir,0 when M^ either Br. or 1. The following is a tabulated statement of some of the commonest non-metallic compounds of these halogens: Ch LOR INK. HCl HCIO HC102 Bromine. HBr HBrO Iodine. HI 8 CHLORINE GKOUl'. Cm.OKINK. H CI 0.5 HCIO4 C1..0 Bromine. HBrOy HBr04? Iodine. HIO3 HIO4 cib. I2O5 S2I, SClo SgHr. Chloric acid is prepared by actipcr on a chlorate, generally of barium, with sulphuric acid. It is very unstable and readily decomposes into hydrochloric acid, oxygen and perchloric acid, thus : 4MC10,= 2HC10, + 2MCl + 30, = 2lIC10,+ 2HCl-f2CX. The perchloric acid, ITCIO,, is much more stable, is not decomposed by sulphuric acid, and forms hydrated acids of the composition ir,C10, = HC10,(II.,0), H,C10,= HC1CX(H,0),. Bromic acid is obtained from the action of chlorine on bromine solution, 5Cl,+ Hr, + 6II,0-2lIBrO,+ loHCl. Iodic acid may be obtained by decomposing an iodate or by oxidizing iodine. If iodine be boiled with strong nitric acid, the following reaction occurs : 3l,+ ioIIXO,=6IlIO,+ ioNO+2HA Periodic acid, IIIO.j, is known only from its salts. Chlorine decomposes water, setting free oxygen ; it is therefore indirectly a strong oxidizing agent. Hydriodic acid, on the other hand, is a strong reducing agent, be- cause of its ready decomposition, thus 2HI+0 = H20-f I.^. QUESTIONS AND KXKUCISES. 9 Hydrobromic acid when dissolved iti water to form a weak solution decomposes the water, as chlorine does. When in stroni; solution, however, it is decomposed by oxygen, as hydriodic acid is. 8.— Questions and Exercises. 1. Dissolve some bromine in water, then shake up with this solu- tion some chloroform. E.\|)lain the result observed. VVoukl ether or carbon bisulphide answer instead of chloroform ? 2. If a person wished to make a 2% solution (l)y weight) of bro- mine in water, how much bromide of ammonium should lie deccmi- pose, assuming that he has a liter of water, and that io% of the bromine set free is lost .'' 3- How may bromine be separated from water? 4. Some phosphorus is lowered into a jar of chlorine and after chemical action has ceased, souie water is atidcd. If you were asked to determine whether you then had a solution of chlorine or of hydrochloric acid, how would you proceed, keeping in mind that the solution also contains an arid of phosphorus .> If bromine were substituted for chlorine, in the foregoing question, how should the answer be altered ? 5. Dissolve a portion of any bromide and add chlorine water. Divide the liquid into two parts. IJoil one portion ; to the other add potassic hydrate solution, and afterwards sulphuric acid. Describe the changes observed, and write equations to explain them. 6. Compare, physically and chemically, bromine with nitrogen tetroxide. How may they be distinguished .? 7. Water containing chlorine in solution gradually acquires acid properties. Is the same true of bromine ? liow may the results be accounted for.? 8. Dissolve some potassic bromide, add sulphuric acid slowly as long as the brown colour of the solution is deepened, then shake up with chloroform. 10 CHLORINE GKOUP. 9. Fill out the followinjr schedule, so as to present a companitive view of the relations of the halo«renb n^.entioned : — CULORINK. (1) Atomic weights (2) Physical condition at 20° C (3) Boiling point (4) Melting point (5) 8p. gr. (vapour or gas, air=l) (6) 8 p. gr. in solid or li(£uid state (7) Odour (8) Colour (9) Chemical Jictivity (10) Bleaching (11) Hydrogen compounds (12) Oxy-acids (13) Basicity of acids (14) Colour of the silver salt of hydrogen acid (15) Colour imparted to starch by the element —34 2.43 1.33 Hromink. lODINF. —63 —24 5.54 200 114 8.71 i3.18(li(iui.l) 4.97 (solid) 10. Place side by side two bottles of solution of chlorine, pass hydrogen sulj)hide through one for an liour, then test both for acidity. 11. Solution of iodine was added to ammonium hydrate, when the liquid became clear, chlorine was added. Write etjuations for the reactions. What would have been the efTect of adding hydro- chloric acid instead of chlorine ? LAW OF MULTIE'I.K I'UOPORTIONS. 11 12. Try if iodine in solution will unite directly with iron, zinc and lead. Use filings of the metals. 13. Fill a test-tube with hydriodic acid gas and drop a little strong nitric aciil into it. 14. Expose a solution of hydriodic acid to strong light for some time. Repeat, using potassic iodide solution. 15. Pass some chlorine into a vessel filled with hydriodic acid gas. Lower a red hot wire into a vessel of the gas. 16. What would be the result of passing alternately hvdrogen iodide gas and oxygen into ajar containing nitrogen tetroxide.'' CHAPTER II. SOME CHEMICAL LAWS AND THEORIES. 1.— Law of Multiple Proportions. Analyses of the oxides of carbon which led Dalton to the discovery of the law of multiple proportions, ^^ave the following results : Carbon monoxide, .carbon 42.86, oxygen 57.14 Carbon dioxide. (( 72.73 From this it follows that there is relatively twice as much oxygen by weight in the dioxide as in the mon- oxide, or that in the monoxide 12 parts of carbon, by weight, unite with 16 of oxygen ; and in the dioxide, 12 of carbon with 32 of oxygen. Hence one volume of carbon unites with one of oxygen in the monoxide, and with two in the dioxide. 12 SOMK CIIKMU'AL LAWS AND THKOKIKS. Analyses of marsh gas and olcfiant gas give the follow- ing percentages : Marsh gas carbon 74.95, hydrogen 25.05 defiant gas... " S5.68, 14-32 Nitrous and nitric oxides yield respectively: Nitrogen 63.71, oxygen 36.29 and " 46.75. " 53-^5 Does the same conclusion follow from these figures as from those of the oxides ol carbon ? 2.— Isomorphism. Nearly every chemical salt has a definite form of crystallization, and it is well known that crystals of different systems will not form in combination ; thus if solutions of two .salts that crystallize in different forms be mixed and evaporated, the crystals of one salt will be quite distinct from those of the other. It is known also that if the salts crystallize in the same system and are of similar chemical composition, the two substances will sometimes be found in the same crystal, that is, in crystalline structure, their molecules appear to be inter- changeable. Such substances are said to be isomor- phous. Examples are the alums and the alkaline chlorides. The statement regarding these crystalline interchanges has been formulated as the law of isomor- phism. It has been of value as a confirmatory test in determining atomic weights and arriving at conclusions about molecular structure, thoivgh there are many excep- r>ri,oN(s AND petit's law. 13 tioiis to it. This law nia\- he staled as follows: Crys- talline structure is dependent only on the number and arrangement of atoms in the molecule, and not at all on their chemical i)n)|)erties. Hence isomorphous sub- stances probably have molecules ..lade up of c(jual num- bers of atoms. 3.— Dulong and Pe tit's Law. Early in the present century Dulong and Petit an- nounced, as a result of their investigations, that " i\\Q atoms of all simple bodies have the .same capacity for heat." While this is not exactly the case, because the same element very generally has different specific heats at different tcm|)eratures, it is true within small limits if the specific heat be taken in that range of temperature at which it is about constant. Now, in a unit weight of oxygen there is one-sixteenth as many atoms as in a unit weight of hydrogen (from Avogadro's law), and since the capacity of each atom for heat, whether oxxgen or hydrogen, is the same, it follows that the capacity for heat (specific heat) of a mass of oxygen will be one-sixteenth that of an equal mass of hydrogen, but atomic weight is sixteen times greater, hence the product of specific heat and atomic weight is the same for both. This product for all ele- ments is nearly constant, viz., 6.4. This principle has been worked out directly by the de- termination of the specific heats only in the case of solid elements, and indirectly for a few gaseous ones. The product of specific heat and atomic weight is called atomic heat. II SOMK ClfKMinAL LAWS AND THKOKIKS. The spt'cifu luMt of ^4 solid I'lcinctits li;is hccii directly (k-tennincd, jS of tlu'sc li,i\c an atomic licat of 6.4, wliilc 6 var)' between tliat and 5.5. This law is of value in determining^ atomic weights. 4. -Some Conditions of Chemical Action. It is generally true that if solutions of two soluble salts be mixed, anil if these salts by metathesis will f(jrm an insoluble compound, that C(jmpound will be formed. Sodium sul|)hate and calcium chloride give sodium chloride and calcium sulphate, the latter being nearly insoluble. Calcium sul])hate and bariiun nitrate form calcium nitrate and barium sulphate, the latter cjuite insoluble. Again if two salt solutions by their interaction form a compound volatile at ordinary temperature, that com- pound will likely be i)roduccd. The preparation of hydrochloric acid, ammonia and marsh gas furnish examples. I 5.— Influence of Mass on Chemical Action. The particular compounds formed when two substances interact chemically are frequently determined by the re- lative quantities or masses of the substances present, thus: Pj^O;{ or l^.Og is formed as the quantity of phosphorus or oxygen is in excess. Antimony and solution of hydro- chloric acid in a little water form antimony trichloride SbCl.^ ; but an excess of water throws down the precipi- tate SbOCl, thus:— SbCL-hII,0-SbOCl-h2HCl rcctly f 6.4, Its. Q. ioluljlc ,1 form ormcd. sodium nearly c form r quite form a Lt com- tion of furnish ion. bstances y the re- :nt,thus: (horus or )f hydro- ichloride I precipi- Cl. I INFf.irKNTK OF TKMI'Klf A riK'K. 15 Arsenic pentoxicie sohitioii throii-^ii which a rapid current of hychoi^rcn sulphide is passed yields arscm"c pentasulphide. As,,0,-f 5II..S As,S,+ 5lI,0. If, however, the liydro^rcn sulpln'de be passed slowly into the pentoxide solution, the followini^r reaction takes place. As/X, -f 2 H,S - As,0,, + 2 1 1,0 + 2S. The oxide of carbon formed in combustion is j)artially determined by the (piantity ofoxyL,ren as compared with the quantity of carbon available for the chemical action. 0.— Influence of Temperature. Examples are connnon enough of change of tempera- ture causing f,r hindering chenn"cal action. In some cases the compound formed is determined by the tem- perature to which the constituents were subjected. Barium oxide, 13aO, heated to dull redness in oxygen or air becomes barium dioxide, BaO.; heated a few degrees higher, however, it changes into the monoxide again. From potassic nitrate and sulphuric acid is obtained either m'tric acid, or water and oxides of nitrogen, de- pendent on whether the temperature is high enough to prevent the constituents of the acid from miiting as'^they are set free. Sulphur trioxide at a red heat breaks down into the dioxide and oxygen. Hydrogen and oxygen um'te at the temperature of a gas flame to form water, but at 2500^ it breaks up again into the original gases. 16 HOMK CIIKMK'Ar, LAWS AM» TIIKOUI I'.S. Oxyijfii .ind mercury do not coinhiiR' at onliuiiry tem- peratures. When heated they do unite .sh)\vl\', but at a httle lu'L;her tcmi)eralure the compound is broken down into its elements ; so tliat tlierc is but one narrow ran^^e of temperature at which tlic imi(jn is possible. Compare ammonium cliloride and potassic chlorate. f 7.— Influence of Pressure. In some cases chemical combination can only be brouLjht abcnit by ap[)l\'ini;- pressure to the mixture of j^ases that for'm the union. In many other cases the temperature at which coml)ination takes place is lowered by tiie application of pressure to the mixture of the con- stituents. This is a matter of some importance as pointinij to chemical action being promoted by other forms of energy than heat. I 8.— Catalytic Action. The influence wliich a third substance has in promoting chemical action between two otliers is called catalysis. The effect of manganese dioxide in forming oxygen and potassic chloride is an example. Recent investigations seem to show that this is a very connnon thing in chemiccd actions. Commonly carbon monoxide and oxygen explode when mixed and brought into presence of a flame. If both are perfectly dry no union whatever will take place. A very minute trace of water vapour will cause the combination to occur. i iry tcm- 1 )Ut at a m down < w raii^^c "> omparc only be ecomes ly with ) a few ydrides ron and ith the , in the of the joined rm the ^ — dis- 2 sohd a posi- )ecause In the the hst lis. THE ALKALINE METALS. 19 CHAPTER IV. THE ALKALINE METALS. 1. They are sodium, potassium, h'thium, rubidium, cae- sium and the hypothetical metal ammonium. Of these, only the first two and ammonium require to be studied here. Sodium and potassium are both white metals, so soft as to be easily cut, they rapidly oxidize in the air, are readily fusible and burn to white oxides. They both decompose water at ordinary temperatures, forming free hydrogen and the hydroxides of the metals. The oxides and hydroxides are both strongly alkaline, and both are readily soluble. The oxides in solution form hy- droxides. Both the oxides and hydrates are strongly basic and never form acids. They are both monads forming compounds of the type MCI, but they form hydrides of the composition M,H. Each metal com- bines with its own hydrate, when fused, to yield the oxide and free hydrogen, thus,— M, + 2MHO- 2M.,0 + Ho. The hydrates are not decomposed by heat (com- pare calcium group), but they are fusible and are very easily soluble in water. The chemical energy of the group increases with the atomic weight. MKTAii. At. Wkioiit. Lithium 7 Sodium 23 Potassium. . . 39 Rubidium ... 85 Caesium 133 Sp. Gr. Mklt'o Pt. Sp. Hkat .594 180° •941 (gas) .972 95.6° .293 .865 62.5° .166 1.52 38.5° 20 THE ALKALINK METALS, 2.— Sodium. Experiments. 1. Hold a loop of platinum wire which has been dipped in a solution of any sodium salt in a non-lumin- ous flame. 2. Try if sodium nitrate gives an alkaline reaction with litmus or forms a precipitate with silver nitrate. Then strongly ignite some of this nitrate either on mica, or in a hard glass tube. Dissolve the remainder and repeat the tests. 2NaNO, = 2NaN02 + 0, = Na202 + 2NO + 02. When taken together, these may be written 2NaN03=:Na202 + 2NO2. 3. Drop some bits of sodium on a little water in an evaporating dish ; after a strongly alkaline solution has been obtained, neutralize it with hydrochloric acid, then pour the clear liquid into its own bulk of strong ammonia, and lead a current of carbon dioxide into the mixture. Na,+ 2H.O--2NaHO l-H... NaHO + HCl- NaCl+H20 NaCl+NH,0H + C0,-NaHC03+NIT,Cl. The precipitate is the bicarbonate of sodium (baking soda). 4. Heat some of the bicarbonate of soda strongly and lead the gas that comes off into lime water. 2NaHCO,= Na2CO, + C02+H20. 5. Try if the carbonate can be changed into the oxide by heating, as in the case of calcium. The resulting substance may be tested by dissolving it with water and I NOTES ON SODIUM. 21 s been -lumin- on with Then :a, or in repeat When ^a202 + Dr in an :ion has ic acid, f strong into the (baking igly and trying if it will throw down an insoluble hydroxide of any of the heavy metals such as silver, iron or tin. See that the precipitate is not a carbonate. 6. Dissolve some sodium carbonate and boil it with lime, keeping the whole stirred. After boiling for some time, decant the liquid and test for sodium hydroxide. Na2C03 + Ca(OH)2=CaC03 + 2NaOH. 7. If sodium dioxide Na.jO.^ is formed when sodium is burned in air, it should yield hydrogen dioxide with an acid. Try it, by burning a piece of sodium on mica in the tip of a gas flame, then scrape the oxide into a tube and add some dilute hydrochloric acid. Test with iodide of potash and starch paper. he oxide resulting ^ater and 3.— Notes on Sodium. Occurrence and Preparation. — On account of its ready union with other elements, sodium is not found native. Its chief compounds that are found as minerals are the chloride, carbonate nitrate and sulphate. Sodium is prepared either by electrolysis of the fused chloride; by heating together sodium carbonate and car- bon to a high temperature and collecting the sodium vapour in naphtha, — NaXO.{+2C — Na,^-f 3CO; or by igniting the hydroxide with carbon, — 2NaOH + 2C — 2CO + H2 + Na2. The carbonate, however, is obtained either as in ex. 3, — t//t' ainmofiia process ; or by altering the chloride, — Leblanc method. The latter consists of three operations: 22 TIIK ALKALINE METALS. {a) 2NaCl + H2S04--Na,S04 + 2HCl. The hydro- chloric acid is condensed for other uses. The sulphate is known technically as saltcake. ill) The saltcake is heated stron<^ly with broken coal and limestone. Na.^SO,^ + 4C==Na2S + 4CO. {c) The sulphide acts on the limestone to form sodium carbonate and calcium sulphide. Na^S + CaCO, - Na,CO,+CaS. The substance now obtained is called black ash from the mixture of coal. By dissolving the soluble parts these are obtained separately, and consist of the carbon- ate with a little mixture of the hydroxide. 4.— Compounds. Oxides. — Sodium forms two compounds with oxygen, NagO, sodium oxide, and Na20.„ sodium peroxide or di- oxide. The monoxide, Na^O, is obtained pure by heating sodium with the fused hydroxide. When sodium burns in air a mixtui"e of the monoxide and dioxide is formed, and when the metal is ignited in oxygen or nitrous oxide the dioxide is produced. It may also be prepared as in ex. 2. When either of the oxides is dissolved in water, — and they dissolve very readily, — the hydroxide, NaOH, is formed, in the case of the dioxide, oxygen is set free at the same time. The hydroxide is a strongly alkaline monad base. Its ready solubility makes it a valuable reagent. Chloride. — Sodium chloride, common salt, NaCl, is an extensively occurring mineral. It is prepared either by SODIUM COMPOUNDS. le hydro- 23 •ken coal 1 sodium ish from )le parts carbon- oxygen, le or di- heating n burns formed, s oxide -d as in water, ^aOH, set free Ikaline aluable 1, is an her by % direct union of the elements or by the action of hydro- chloric acid on the oxide, hydroxide or carbonate. Carbonates.— ^\\(i bicarbonate, NallCO,, (baking soda) is the first product obtained. From this the neutral car- bonate, NaXO;j, (washing soda) is r ot by heating. Other reactions that are of theoretical interest for the prepara- tion of these salts are {a) by passing carbon dioxide into a solution of the hydrate, thus : 2NaHO-f CO,,--Na^CO.^ + H.,0; (/;) by passing CO,, into a solution of the normal carbonate : Na,CO,,+ H,,6 -f 2CO0 = 2NaHCO,. When the bicarbonate is heated it loses water and carbon di- oxide, the exact reverse of the reaction with CO2. In the acid carbonate, the hydrogen may be replaced by alkalies, thus producing NaKCO,, NaNH,CO„ etc. Nitrate. — The nitrate of sodium may be prepared by the action of nitric acid on the oxides or carbonates. It is found in great quantities in the nitrate beds of Peru, and is known as Chili saltpetre. Sulphur Compounds.— The following sulphides of sodi- um^ are known : Na,S, Na,S„ Na.S^, x\a,S,. They are of interest chiefly because of their resemblance in com- position to the oxides. When the metal is heated with sulphur, a mixture of sulphides is formed. Hydrogen or carbon heated with the sulphate reduces it thus :— Na,SO, + 4H,=:Na,S + 4H,0 Na,SO,-f4C -Na,S4-4CO. When sodium chloride is fused in hydrogen sulphide, the following reaction occurs : H,S-[-2NaCl Na^S-i-2HCl. 24 THE ALKALINE METALS. Sulphuretted hydrogen forms with the hydroxide a hydro- sulphide, NaOH + H.^S = NaHS + H^O. This compound can also be obtained by the action of sodium on hydrogen sulphide, Na.^-|-2ll2S = 2NaIIS+ H^ (compare sodium on water). With sulphuric acid, the oxide or carbonate forms a sulpJiate^ Na2S04 (Glauber's salt). It is also produced extensively in the manufacture of soda ash, by Leblanc's process, when the chloride is treated with sulphuric acid. The tJdosulphate^ commonly called hyposulphite, Na^SoCj. This is the sodium salt of thio-sulphuric acid, H^S^Og. If sulphur dioxide be passed into a solution of caustic soda, sodium sulphite Na^SOg is formed, and when this is boiled with sulphur, the thiosulphate is produced. Hydrides. — When sodium is heated in a current of hydrogen the hydride Na^.H is produced. It is decom- posed readily by heat, and, in presence of mercury, breaks up to form an amalgam of sodium and free hydrogen. The very strong basic properties of sodium oxide almost make the number of the salts into which it enters to be limited only by the number of acids known. 5.^-Tests. {(I) The yellow colouration of the flame. (/;) Sodium salts, when acidified with hydrochloric acid, form red crystals with platinic chloride. ide a hydro- is compound )n hydrogen are sodium )r carbonate I^ is also loda ash, by -eated with /posulphite, phuric acid, solution of )rmed, and julphate is current of t is decom- f mercury, I and free :ide almost Qters to be POTASSIUM — COMPOUNl)8. 6.— Potassium. 25 loric acid, This metal resembles sodium very much. The ex- periments with sodium may be repeated, but in each case the corresponding potassium compound is to be substituted for the sodium one. The flames of potassium and sodium should be compared, when looked at through a piece of blue glass. This cuts off the yellow sodium rays. 7.— Notes on Potassium. Occurrence and Preparation.— Potassium is a con- stituent of wood ashes, in which it exists mostly as a carbonate. It is also one of the substances contained in felspar ; and from the disintegration of felspathic rocks it becomes an ingredient of clay. It occurs as a chloride forming the mineral syh'/te, and as the nitrate it forms an incrustation on dry soils in some parts of the world. The metal is prepared just as sodium is. 8.— Compounds. The hydride, K.^H, the monoxide, K,,0 and dioxide, KgOg, the hydroxide, KOH, the nitrate, KNO.j, chloride,' KCl, and sulphate, K,SO„ are prepared in ways similar to those employed for the corresponding sodium com- pounds. Oxzdes.—Fom oxides of potassium are known. They are monoxide, K,0, formed (along with the dioxide) when the metal is burned in air, or when the metal is fused with the hydrate ; dioxide, K,0,, produced by com- 26 THE ALKALINE METALS. bustion in air, or in oxygen, in the former case it is mixed with monoxide, in the latter, with trioxide K^.O^, and tetroxide K^,04. These higher oxides, with the hydrox- ide, are also produced by the decomposition of K.X),, in air in presence of water vapour. A possible reaction is indicated thus : — 3K,0,+ H,0-fO,-K,0, + KoO, + 2KIIO. All the other oxides are reduced to K^O at high temper- atures. Iodide. — This important compound, KI, is obtained mixed with the iodate KIO.<, when iodine is introduced into the hydrate and heated. (Compare High School Chemistry, p. 170.) 6I-f6KOH-:5KI + KIO,+3H20. The iodide is soluble in alcohol, but the iodate is not, hence they can be separated. The iodide may also be prepared by the direct action of the elements, or by acting on the hydroxide, oxide or carbonate of the metal with hydriodic acid. Iodate, KIO3, and periodate, KIO^, are two other potassic compounds of iodine. CJdorine Cojnponnds. — The most important of these are the cJdoride and chlorate. The former is prepared in ways similar to those in which the iodide is obtained, or by decomposing the chlorate by heat. The chlorate has been partly described in High School Chemistry, page 170. It is valuable chiefly as an oxi- dizing agent, and as a source of oxygen. COMPOUNDS. 27 is mixed ^20;^, and hydrox- K,(), in taction is 3. temper- obtained troduced School :e is not, also be s, or by tie metal /o other :hese are Dared in lined, or 1 School an oxi- '4. Potassic hypochlorite, KCIO, results from the action of chlorine upon caustic potash, 2KOH-I-CI, KClO-f KCI + H2O. On boiling the mixture the hypochlorite breaks up into the chloride and chlorate, 3KC10 = KC10.5-f-2KCl. Sulphur Compounds. — The monosulphide, K.,S, may be prepared by fusing the sulphate with carbon, 2C + K2S04 = K,S + 2C02. The polysulphides K,S,„ KoS„ K^Sj and K.S, are all obtained by boiling the monosulphide with sulphur. A hydrosulphide, KHS, is formed when caustic potash solution is saturated with hydrogen sulphide. It is probable that the monosulphide is converted into the hydrosulphide and hydroxide by the action of water, thus:— K,S + H,0 = KHS + KHO. This is analogous to the reaction of the oxide with water. The polysul- phides, in presence of acids, break down into hydrogen suphide and sulphur. KA + 2HCl-.H,S + 2KCl-f2S. Other Co7npounds.-~?o\^?>?:\Mm unites with the cyano- gen radical to form potassium cyanide, KCN. This is obtained from the decomposition of ferrocyanide of potash by heat. K4Fe(CN),^4KCN+FeC, + N2. If an oxidizing substance or sulphur be present, the cyanate, KCNO, in the former case, or the sulphocyan- ide, KCNS, in the latter case is formed. I 28 THE ALKALINE METALS. 9.— Tests. {a) The flame colour of pale violet or lilac, but appear- ing red through blue glass or indigo solution. {b) With tartaric acid, II.,C.,H,0,., or with its acid sodium salt, potassium forms a white granular crystalline precipitate. (c) With platinic chloride, PtCl,, potassium salts, in solution, form a yellow precipitate, but often it is neces- sary to stir the mixture for some time. 1 0. — AmmoD ium. The hypothetical metal NH^ belongs to the group of the alkalies. Its supposed amalgam with mercury has been referred to in High School Chemistry, page 109. When the amalgam is cooled below freezing point it forms a brittle metallic mass. The chief ground, how- ever, for believing that there is a metal having the mole- cular composition NH4 is the resemblance of its salts in number and character to those of the alkaline metals, particularly potassium. The following is a partial list of compounds of potassium and ammonium, not by any means complete, but sufficient to show the marked similarity. Each unites with the groups: HO, S, S.^, S3, S4, S„ SO,, HSO3, SO,, HSO„ NO,, NO.,, CI, Br, I, CN, SCN, A1(S0,)2, CO,, HCO„ AsO„ AsS,. It will be observed that these are precipitation compounds, or else those formed by double displacements by heat- ing. The compounds of ammonium corresponding to those of potassium obtained by heating the metal are (^IJKSTIONS AND KXKIUMSKS. 29 appear- its acid ystalline salts, in is ncces- :| 1 TOUp of :ury has Lge 109. ;• point it id, how- le mole- salts in metals, d list of by any i marked .5 >, 0.2, 03, ■ Br, I, < It will ^ jnds, or ; y heat- -■ ding to = not formed because the ammonium hydroxide breaks down into ammonia and water. Another difference is that when the ammonium halides are heated they break up into ammonia and the corresponding halogen com- pound of hydrogen. This is not the case with the other alkalies. Ammonium salts are isomorphous with those of potas- sium (see page 12). Ammonium salts, when heated, are decomposed ; thus the chloride breaks up into ammonia "nd the acid, but on cooling these again unite. Sodium or potassium hydrate displaces ammonium hydrate from its com- pounds and sets ammonia free. 11.— Tests. (a) Pungent odour and effect on litmus and turmeric. (^) Solution of an ammonium salt boiled with a caustic alkali sets free ammonia. (c) Nessler's test solution. ^tal are 12.— Questions and Exercises. 1. Potassium and sodium may be obtained by ij,miting some of their compounds with carbon or hydrogen. Why can not the ammonium radical be got in a similar way .'' 2. "The acid character of hydrogen sulphide and its basicity would lead one to expect that with caustic alkalies two replacements 30 MKTALH OF TIfR (!Al,(irM fSKoHP. would he possible, resulting fin.iJly in M.^S wlu'ie M is the metal of the alkali." l''.x|)Iain this stateuient. Does the actual result cor- respond with the theoretical ? 3. Boil sonic solution of sodium hydrate in a test-tulie with a little powdered sulphur, after it has become dark coloured adil a few (Ir()|)s of hydrochloric acid. Smell the escaping gas. Write equations for the reactions. 4. Sodium hydrate is heated with sodium, some dilute nitric acid is added, the result evaporated to dryness and ignited ; the residue treated with a solution of sulphur dioxide, again evai)orated and ignited with carbon in the air. Trace the chemical changes, with e(|uations. 5. (rt) Steam is passed through a tube containing some pieces of sodium. (/>>) Hydrogen is passed over heatetl sodium oxide. Is there a chemical action in either case ? In both .'' CHAPTER V. METALS (JF Tllh. CALCIUM CROUP (ALKALINE EAKTIIS): CALCIUM — STRONTIUM— BARIUM. 1.— Group Characteristics. All are yellowish metals, obtained by electrolysis of the fused chlorides. They form oxides of the composi- tion MO, peroxides, MO^„ hydrates, M(HO).^, which are sparingly soluble forming weak alkaline solutions. The hydrates, when strongly heated, are decomposed into the oxide and water, thus M(HO)2=MO + H.p. CAMMITM. 31 V: metal of esult cor- 36 with a red add a s. Write litric acid le residue a ted and Kcs, with pieces of xide. Is KTIIS): ysis of )mposi- ich are n posed ,0. The carbonates are decoiiiposed by heat into the oxides and carbon dioxide, MCOy^MO + CO.j. The oxides are stron^dy basic, that of calcium beinj^ tiie weakest and of barium the strongest ; thus chemical eneri^y increases with atomic weight. The carbonates, phosphates and sulphates are almost, or entirely, insoluble. 2.— Calcium. Experiments. 1. l^cnd tiie end of a bit of platinum wire into a loop, with this hold a little of some calcium salt in the non- luminous flame of a Bunsen burner. 2. Powder a piece of marble, and try if it is soluble in water. Drop some of the powder into some hydro- chloric acid diluted one-half with water. When gas ceases to come off, evaporate a part of the liquid to dry- ness, and keep the powder obtained. To a second portion of the liquid add soine dilutfe sulphuric acid, and filter. To another portion add excess of sodium car- bonate solution, and filter. The salts obtained are calcic chloride, calcic sulphate and calcic carbonate. Write equations for the entire reactions. 3. Try if the sulphate is at all soluble. To do this shake it up with a large quantity of water, and let it stand for somr time, then test the clear fluid with barium nitrate solutioi 4. Test a littk of the solution of the chloride with a solution (not too dilute), of any soluble sulphate to see 32 METALS OF TUK CALCIUM (SKOUP. if there is calcic sulphate formed. What reason is there for the direction "not too dilute "? 5. Heat strongly some of the carbonate in a small crucible for some time. Test with moist red litmus paper. Does the unheated carbonate give a similar result? Try if a piece of marble, a piece of chalk (not common crayon), and a bit of oyster shell will act as the carbonate does. CaCOa when heated is decomposed into CaO and CO^. CaO is lime ; try if it will dissolve in water. To do this shake it up with water and after the sediment has settled test the liquid with red litmus paper. 6. Pour into a solution of chloride of calcium a little solution of sodium phosphate. 7. Repeat the last experiment, but use an alkaline oxalate instead of the phosphate. CaCl., + (NH^).,Co04 = CaC^,04H-2NH^Cl. Examine the crystals when they form. 8. Will other soluble carbonates act like that of sodium in ex. 2 ? 9. Let some lime water (solution of calcic oxide) stand in an open vessel exposed to the air of a schoolroom for a couple of days, then pour out the liquid, leaving the white scum on the side of the vessel. Drop a little dilute acid on this scum. Whai is it ? Has the alka- linity of the liquid been affected ? 10. Place a bit of slaked lime in a test-tube and heat it strongly. Water should be formed on the cold part of the tube. § NOTliS ON CALCIUM. 33 on is there n a small ed litmus a similar :halk (not act as the composed 11 dissolve and after ed litmus im a little 1 alkaline 4)2Co04 = hen they Df sodium de) stand Iroom for aving the 3 a little the alka- id heat it I part of .1. Notes on Calcium. Occurrence. — Calcii:'n never occurs free, but in com- bination it exists in vast quantities as limestones, chalk, gypsum, marble and apatite. Calcium and Oxygen.— Calcic oxide, CaO, quick- lime, is formed by i^i;niting the carbonate in presence of air. CaCO.j=^CaO + COo. If the carbon dioxide be allowed to surround the carbonate, decomposition does not occur. With water it forms tlie hydrate Ca(OH)^„ slaked lime, which is slightly soluble, the solution being li}ne ivatcr. The undissolved hydroxide suspended in water is Dtilk of lime. Mortar is a mixture of calcic hydroxide and sand with water ; the hardening of it is due to the hydrate absorb- ing carbon dioxide from the air to form the carbonate, and at the same time acting on the sand to produce a silicate of calcium. Calcium forms a second oxide, CaO.,, b\' the action of calc'c bvcirate solution on hydrogen dioxide. It is a true peroxide, that is, loses oxygen when acted on by ■i. . .is, because it is then reduced to the basic oxide CaO. The i^eroxide is a hydrat-^d compound, its full formula be ng CaO,+8H,0. Calcium and the Halogens.— The most important of these is calcic chloride CaCl.,. It is prepared by act- ing on calcic carbonale wirh hydrochloric acid, or, w hen chlorine is v-asscd invo cold dry slaked lime, a mixture of the chloriuc and the hypochlorite is formed, this is bleaching poiVcLr Tins substance is coiumonly supposed to be CaClo-|-Ce. Clv))^„ but there is reason for believing 34 MKTALS OF THE CALCIITM GIJOUP. it to be C<'iCl(OCl). Sec H. S. Chem., p. 169. Calcium fluoride (fluor spar), CaFa is an insoluble mineral, the source of fluorine. (Compare potassium.) Calcium Carbonate. — This occurs in nature in many- forms ; some of the most important arechaP', limestone, marble, calc spar, arragonite, do'omite, coral rocks and shells of animals. It may be prepared by treating any soluble calcium salt with a solution of a carbonate. When heated it yields the oxide and carbon dioxide. It is insoluble in water, but soluble, to soir.o extent, in solution of carbon dio.xide, probably formii'.;^^ Ca(HC03).2. See H. S. Chem., p. 129. Other Compounds. Calcium Chlorate, Ca(C10.5)^, formed by passing chlorine into a hot solution of the hydrate. See H. S. Chem., pp. 161, 170. Calcic Sulphate, gypsum, CaSO^ is found in nature. When heated to iio°— 140*^ and powdered it form-: plaster of Paris, which readily absorbs water and harden'. {sets). Calcic Phosphate, Ca.,(P04)2 is the chief ingredient 01 apatite {phosphates), and is the mineral constituent of bones. This hr-: already been referred to under the pre- jjaration of phosphorus. Calciiiin Sulphide^ CaS is obtained by he-iting the sul- phate with powdered charcoal, CaS04 + 2C=CaS-{-2C02. 4.— Strontium. Experiments. I. Hold a bit of strontium nitrat'.' -'O a platinum loop in a non-luminous flame. m Calcium era], the in many Tiestone, )cks and ting any rbonate. dioxide, ^tent, in :hIorine Chem., nature. ; form- larden', ient 01 itrt of lie pre- he sul- •2CO.. 00pm 1 NOTES ON STRONTIUM. 35 2. Drop a little solution of calcic sulphate into a solu- tion o{ nitrate of strontium. If necessary let this stand for some time until a white precipitate appears ; filter and dry the precipitate. By the flame test determine whether it is a calcium or strontium salt. 3. Try to repeat this experiment, but use instead of calcic sulphate solution, water in which barium sulphate has been shaken, and has become clear by settling. 4. Mix a little solid nitrate of calcium and nitrate of strontium. Try to dissolve the nn'xture in absolute alcohol. After shaking for some time, filter, evaporate the filtrate and test in the flame both the part dissolved and that left. 5.— Notes on Strontium. Strontium is found in nature chiefly as two minerals, strontianite and celcstine, the carbonate and sulphate, respectively, of strontium. Compounds. — Strontium resembles calcium very much, but chemically it occupies a place between calcium and barium. The two oxides SrO and SrO.,, and the hydroxide are prepared in the same manner as the corres- ponding calcium compounds. The sulphate SrSO,, the nitrate Sr(NO,),„ and the chloride, SrCl,, are obtained by treating the carbonate SrCO^ with the proper acids. 6. — Barium. Experiments. I. Treat a little of any barium salt with strong hydro- chloric acid, then :;Tnite it on a platinum loop E>^ 36 METALS OP THE CALCIUM OROUP. 2. To a solution* of barium nitrate, add a solution of ammcMiium carbonate, filter the precipitate, dry it. Try if it is a barium salt, by the flame test. Test it also with a few drops of strong acid for a carbonate. 3. Heat some barium carbonate as calcium carbonate was heated to drive off the c irbon dioxide. Try with strong acid if the carbonate has been decomposed. If it i s not, raise it tc a high temperature (this will require a b; . aip flame) for some time. This should at least partly u ompose the salt into BaO and CO.^. 4. Try if the white solid left from the last experiment is at all soluble in water. This may be done by testing the liquid with litmus and with a drop of dilute sulphuric acid. 5. Heat a small quantity of barium nitrate in a crucible. After all effervescence has ceased test the residue for barium oxide by treating with water and using litmus or turmeric. This really tests for the hydroxide Ba(0H)2, which is formed in the same way that Ca(OH)2 is. 6. Dissolve some of a barium salt, say the chloride, in dilute hydrochloric acid. Divide the solution into two, parts. Add to one part strong nitric acid, to the other strong hydrochloric acid until a precipitate forms. (The fact that a precipitate does form with strong hydrochloric acid should be kept in mind else errors will occur in s\'stematic testing and barium will be mistaken for one of the elements whose chlorides are insoluble.) After a dis- tinct precipitate has formed add water. Is the chloride soluble ? Note. — Hi'riiim salts that are not readily soluble in water may be dissolved in weak hydrochloric acid. I NOTES ON HARllJM. n? Lition of ■fi' t. Try it also .' i-i; rbonate M ry with "% jed. If require 1 i at least - .■ >; ^rirnent J testing ■ Iphuric 1 rucible. « iue for '■ ?*■ mus or (OH),, ride, in .- '\ o two, .■^- 2 other (The chloric cur in ■v- one of " a dis- .■V- iloride i in weak 7. Mix solutions of any barium salt and bichromate of potash K.,Cr.,Oy. Repeat, but use instead of barium a solution of a strontium, then of a calcium salt. 7.— Notes on Barium. Occurrence. — Barium is chiefly found in the form of two of its insoluble salts, the carbonate BaCO.,, (wither- ite) and the sulphate BaSC)4 (hcav}'spar). Compounds. — The combinations of barium are very similar to those of calcium, except that they are more stable. The oxide, when heated in oxy^^en or air to dull red- ness, becomes the dioxide ; and when heated still higher changes again to the monoxide yielding free oxygen. This method has been used for the preparation of oxy- gen. When barium oxide is brought into contact with water, barium hydroxide Ba(C)II)o, is formed ; and this, when dissolved in water, yields an alkaline solution, baryta water, similar to lime water. Baryta water is affected by carbon dioxide just as lime water is. BariiiJfi carbonate is much more difficult of decomposi- tion by heat than is calcium carbonate ; indeed it is hardly possible to drive off all the carb:)n dioxide. In this particular the carbonate has some resemblance to those of the alkaline metals, which are not at all broken up by heating. Barium sulphide is prepared by passing hydrogen sulphide into barium hydrate. Ba(OH)o -f-H,S = BaS-|-2HA 38 MKTAr.H OF THE CALCIUM (JKOUP. Barium sulphide with water forms the hydroxide and hydrosulphide of barium, thus :—2BaS + 2H.,0 = Ba(OH),, + Ba(HS),,. (Compare alkah'ne metals.) 8— Tests for the Calcium Group. (0 Alkaline carbonates, such as (NHJ^CO,, give a white precipitate in neutral or alkaline solutions of the salts. {^) A soluble sulphate throws down a white precipitate. (^) A solution of calcic sulphate yields a white precipi- tate with salts of Sr and Ba. (4) Alkaline oxalates give white precipitates in pre- sence of ammonia. Acetic acid, HC,,H,0,, dis- solves barium and strontium oxalates slightly, but that of calcium scarcely at all. (5) Sodium phosphate yields a white precipitate with each metal in presence of ammonia. The elements of the group are distinguished from each other (a) By the flame test. (^) By solution of CaS04 forming precipitate with Sr and Ba. (c) By bichromate solution, see ex. 7. 9— Questions and Exercises. I. If bleaching powder, in solution, were boiled, chlorine would be set free, what would be the effect of this on the other com- ponents of the mixture .? ^ (QUESTIONS AND EXKRCISKH. 30 What 2. What compound is left when eggshells are burned ? 3. Drop a lump of calcic carbonate into sulphuric acid happens, and why? 4. (a) Dissolve a little lead acetate; decant the clear solution, then drop into it some hydrochloric acid. (/-') Drop a little saturated solu- tion of barium nitrate into hydrochloric acid, (c) Pour gradually strong hydrochloric acid into a little of the saturated solution of a barium salt. These experiments show that in testing with hydro- chloric acid there is danger of mistaking barium for lead. How may this be avoided ? 5. Make, separately, solutions of barium, strontium and calcium sulphate, add some ammonium hydrate to each, then some solu- tion of ammonium carbonate. After standing, filter and drop a little nitric acid on the powder on the filter paper. Does this give a method of distinguishing between the elements? If the three sulphates had been mixed would this enable you to make any separation? To the filtrate add calcic sulphate solution ant' let it stand for some time. What now is precipitated ? Does this aid in separating the substances ? 6. Make solutions of the chlorides of the metals, add strong alcohol (absolute alcohol is much the best) to each. Observe their solubilities, dissolved. Add a large volume of water to the ones that have 7. Compare the metals of the calcium group and of the potas- sium group in regard to the following : — (a) their chemical action with water ; (d) the effect of heat upon their hydrates ; (6) the effect of heat upon their carbonates ; {(f) the solubility of their sulphates ; ((?) their distinctive flame colours. 40 'I'irK MAliNKSIUM <;H()UP. CHAPTER VI. TITK MACNESIUM (iKOUP. ^ The metals of this -roup are berylh-um, magnesium, zinc and cadmium. 1— Group Characteristics. ^ All are white metals, not oxidized by exposure to the air. ^ Ma^rncsium and zinc are both oxidized when heated in air, they will both decompose water at high tempera- tures, both form basic oxides, that of magnesium slightly alkaline, both unite with the halogens. They form one insoluble (nearly, in the case of magnesium) oxide of the composition MO. The carbonates are insoluble. 2— Magnesium. Experiments. 1. Burn some magnesium under a bell-jar, collect the white powder, try to dissoKe it, and test the liquid for alkalinity. MgO + IT,0 = Mg(MO),. 2. Pass steam through a hard glass tube in which is a piece of magnesium wire, heat the tube under the wire and collect the gas that comes off while the metal is burning. Test this gas. Compare the substance in the tube with that obtained in the last experiment. 3- Add to a solution of magnesic sulphate (Epsom salts) some caustic soda, or caustic potash solution. MAUNKSIUM. •il Add to this, after the precipitate is formed, ammonium hydrate. Try if ammonium hydrate may be substituted for that of potassium or sodium in the first operation. Add it slowly at first, then in excess. Will solutions of the ammonium salts cause the precipitate to dissolve? Try the chloride, nitrate and sulphate. 4. Add to a solution of magnesic sulphate, a solution of potassic or sodic carbonate, until the mixture ceases to give off gas. 2MgSO, + 2K,CO,+ H,0-MgCO,Mg(HO),,+ 2k,SO, + C02 -Mg,(CO,)(MO),, etc. The precipitate is basic carbonate of magnesium. This salt is quite variable in its composition, depending appar- ently on mass and temperature of the substances used, thus : — 3MgS04 + 3K,C03+H,0^MgCO3Mg0MgrOH)..+ 3K,S(), + 2CO, -Mg.^(CO,)(HO),(0),etc., is another form of the reaction. This basic carbonate is usually spoken of as carbonate simply. Pass a current of carbon dioxide into tlv original solution. 5. If the carbonate be strongly heated it will give off carbon dioxide, as may be shown by its ceasing to effervesce with acids. 6. Treat some magnesic carbonate with hydrochloric acid, evaporate to dryness, then heat. MgCO,-f2HCl-MgCl.^ + HX03. 42 NOTKS ON MACNKSIUM. The chloride contains 6 molecules of vvate: and is there- fore a hydrated compound. When heated to drive off the water the chloride is decomposed and forms the oxide, MjtO. MgCl, + 6H20-:Mg0 + 2HCl + 5H,O. 7. Treat some solution of a magnesiiwn salt with disodic phosphate. An acid phosphate of magnesium, MgHPO^, is formed. Is it soluble in ammonia.'* 3.— Notes on Magnesium. Occurrence. — Magnesium is found as a neutral car- bonate, magnesite, MgCO.j ; as the sulphate, Epsom salts, MgS04 y ^^ ^ double carbonate with calcium in dolomite ; as silicates in serpentine, asbestos, talc, soap- stone; and as a chloride, sometimes associated with those of sodium and potassium. Preparation. — Magnesium is obtained in the metallic state by heating the chloride with sodium. The sodium unites with the chlorine and the magnesium is set free. Other Properties.— Magnesium is nearly related to the metals of the calcium group, as is shown by the formation of similar oxides and hydroxides which are slightly soluble and slightly alkaline, by the decomposi- tion of the carbonate by heating, by the chlorides crys- tallizing with the same number of molecules of water. Magnesium unites with the alkaline metals to form double salts, such as MgCUKCl or MgKCl,, NH^ClMgClg, MgS04K.,S04, etc. It decomposes water at about 100° C, though somewhat slowly. I ZINC. 43 is there- drive off )rms the alt with ^niesium, tral car- Epsom :ium in c, soap- h those netallic sodium : free. Ltcd to by the ch are nposi- 3 crys- ter. ' form 30° C, 4.— Zinc. EXPKKIMENTS. 1. Heat a cutting of zinc on charcoal in the oxidizing flame. Observe the colour of the oxide when hot, and when cold. 2. Cut a strip of zinc to a point at one end ; hold this end in a gas flame until it melts, continue holding it in the edge of the flame but without letting the melted part drop off. Observe the colour of the flame produced, and the appearance of the zinc oxide formed. Try if a bit of zinc clipping when heated on mica will oxidize and burn in a similar way. 3. Pour on some granulated zinc, separately, some hydrochloric acid, nitric acid, sulphuric acid, diluted in each case. Evaporate the fluid to dryness, wash the salts, dry again and test for a chloride, a sulphate and a nitrate. 4. Make a solution of zinc sulphate and add sodium carbonate. Basic carbonate of zinc is formed, the com- position depending largely on accompanying conditions, high temperature and excess of water increasing the basicity. Probably two of these carbonates are formed as follows: — 4ZnS0,+ 4NaX0,+ 3H.,0 - ZnCO;,3Zn(OHX4- 3CO,+4Na,SO, = Zn,(CO.0(OH),, etc. 3ZnS04+3Na.,C03+H,0 -ZnCO,Zn(OH),ZnO + 2CO,4-3Na,S04 = Zn,(C0,)^0H)2(0;, etc Compare corresponding salts of magnesium. ^ 44 /IN<*. 5. I'illcr out some of the j)rcci|)itatc of the last cxperi- inciit, heat it stroiii^ly on mica or in a h^ird j^dass tube. After it turns yellow, try if it effervesces with an acid ZnCO.p— ZnO+ClO^,, when h .li. 6. Mix some zinc oxide with powdered charcoal and heat it on charcoal in the reducing flame. This should give metallic zinc. 7. Try if zinc oxide is soluble. Is it alkaline? Use both turmeric and litmus in testing it. 8. Make a solution of zinc sulphate, divide it into three parts. Add to one of these very gradually a s(;lutioti of sodium hydrate luitil it is largely in er'^ess. Do the same with the other parts, but use foi c potassium hydrate and for the other ammonium liydrate. The pre- cipitate formed at first, in each case, was zinc h\-drate Zn(OH),, thus :— ZnSO,+2NaOH Zn(()H),+Na,SO,. When the prccij)itate dissolves, probably the following reaction occurs : Zn(OII),,-f2Na(OH) = Na,ZnCX+2H,(). Na.,ZnO.,, called zincate of sodium, has the composition of the sodium salt of ^n acid H.,ZnO^. 9. Filter out some of the hydrate of zinc and heat it strongly. It changes to the oxide. Zn(OH)o — ZnO+ H.,0. Compare the similar operations in the cases of the calcium group and of magnesium. 10. Try if a precipitate is formed in a solution of a zinc salt by hydrogen sulphide or ammonium sulphide ; make the former test in both acid and alkaline solutions. After adding the sulphuretted hydrogen, pour in gradu- I * t'xpcn- ss tube, acid )al and should I NO'IKS oV 7.\Sf\ 4.S ,'ill\- aminoniuin li)'drati'. Tlu: whilt* prcc ipitatc is zinc sul|)hidc and is the (>///}> u'/zi/t' >u\[)\\'\i\c oi a metal known. Test it for sohihiiity in nitric, h.j'drocliloric, sulphuric and acetic acids. II. Will hydrogen suli)hide precij)itate zinc sulphide in presence of excess of potassic h)(lroxide? Use to three Lit ion of Do the tassiuni lie j)re- lydrate S'a.,SO.. II ow- ing it ion of heat it ZnO+ ises of n of a phide ; Litions. jradu- I n. — Notes on Zinc. Occurrence. — The chief sources of the metal are the carbonate, ZnCC)., (calamine), and the sulphide, ZnS (blende). From thc-,c, metallic zinc is obtained by con- verting the ore into the oxide and heating it in closed retorts with charcoal. The metal distils out and collects in cold tubes. Properties. — Zinc is a white metal, rather brittle at ordinary temperatures, but malleable and ductile at 150'' — 200', above 200^ it becomes very brittle. It inells at about 430", boils at about 1000', hence may be distilled. Zinc burns with a biight somewhat greenish flame. Like magnesium, zinc forms with the alkalies double salts, such as ZnCI.,2KCl, ZnCl.,2NaCl, etc. The zincate of potash, K.^ZnO.,, of ex. 8, may be compared in constitu- tion to this double salt if th.e chlorine is exchanfjed for oxygen. This compound corresponds (n) to the salt of an acid having the composition IT.ZnO^, {/>) to a double salt of zinc and the alkali in which the acid radical, or the halogen is replaced by oxygen, (c) to the hydroxide of zinc in which t metal. hydi tC1,;4-6H20, can be obtained by treating the hydroxide with hydrochloric acid. (Compare magnesium.) When an aluminium salt in solution has an alkaline hydrate added to it, the hydroxide is first precipitated, but this dissolves in excess of the reagent, forming an almnitiatc of the metal Al(OH),4-3M(OH)- Al(OM).j-}-3H,OwhenM- KorNa. Similar compounds are formed with calcium, b.'rium and strontium. Many of these aluminates occur native NOTES ON ALUMINIUM. 63 !). The ^ The h them :idic to oxides ng the iter by though -^e with ition is chloric of the ler the alone. tained acid. kaline tated. and are known as spinels. Ortho-aluminates have the composition M,A10;,; meta-aluminates, MAIO.. Com- pare boron, page 48. Double Salts. — When aluminium hydroxide and an alkaline hydrate or carbonate are melted together an aluminate of the composition AlOMO is obtained when M=^K or Na. This may be represented thus, — 0=:Al-0-Na. Similarly, aluminium chloride melted with a chloride of a strong base-forming element yields a double chlo- ride, thus : — A1X1,+2KC1 == A1,KX1, == 2 AlKCl,,or CI, == Al - CI, - K. Of these double salts the sulphates form the most important compounds of aluminium. They have the composition A1,M,(S0J, + 24H20 -Al(S04),MoS04 + 24HA These are the alums. This name is given to a group of double sulphates, resembling each other in composition, in crystallization and in their chief physical properties. The type is R2M2(S04)4 4- 2411.^0, and in ordinary alums. R = Al, Fe, Cr or Mn, and M-K, Na or NM,. Silicates. — Among metamorphic rocks silicates of aluminium are abundant, two of the most common being feldspars, that have the composition KAlSiO,), and NaAl(SiOj.2. From the decomposition of these through atmospheric influence arises the insoluble aluminium silicate, clay. 54 TESTS KOI{ liOKON ANI> ALUMINIUM. illustnitc til )f all ^experiments 9 and 10 illustrate the use ot aluminium salts in fixin|^ dyes. One of the common industrial uses of these salts is as mordants, that is, substances fixing or making fast colouring matters. 6.— Tests. Iterates. (i) Treat with any strong acid, which displaces the boric acid, dissolve this in alcohol, ignite the solution, — flame green. (2) Moisten the salt first with sulphuric acid, then with glycerine, after heating to drive off excess of the acid. When ignited, green flame appears. (3) Borates impart a red colour to turmeric paper, when dry. AlHinmiuin. (i) Alkaline hydroxides precipitate Al.,(OH),;, slightly gelatinous, soluble in excess of the alkali except when ammonium is used (compare zinc). (2) Sulphuretted hydrogen does not precipitate the sulphide. (3) Ammonium sulphide precipitates Al/OH)g with free H.^S. (4) Ignite the powdered salt on charcoal, then moisten the mass with cobaltous nitrate, Co(NO.)o ; again ignite and a bright blue colour will appear. QUESTIONS AND EXKRCISES. 55 7.— Questions and Exercises. I. Prepare a bead of borax on a platinum loop. Do this by bending a loop about one-eighth of an inch in diameter on the end of a bit of wire. Heat this, and while still red, dip it in some powdered borax, fuse the borax, until it forms a transparent bead the size of the loop. Take up on this bead just a trace of copper oxide and hold it in the oxidizinj,-- tlame (the outer tip of a bunsen flame will do). Note the colour the bead becomes when hot, when cold. Clean the wire by crushin-,^ the cold bead by a blow on a flat plate. Repeat the experiment, using instead of cojiper, iron rust, manganese dioxide, and bichromate of pota%, sei)arately. Take only the smallest portion of the salt in each case. 2. Try if other alkaline sulphides than that of ammonium will precipitate AL(OH)g from a solution of a salt of aluminium. 3. Add sodium carbonate to aUnVi solution, filter off the precipi- tate, and test it for a carbonate. Compare the result with that similarly obtained from magnesium and zinc. 4. When a soluble aluminate is treated with a chloride of one of the metals of the calcium group an insoluble aluminate is thrown down, thus: sNaAlO.^ + CaCl2 = 2NaCl + Ca(A10o)2. What is the valence of the radiciU AlOo .'' 5. Try if a barium salt, when wet with sulphuric acid will give the characteristic green flame. 6. Aluminium hydrate when ignited yields the oxide ; what other metals have been found capable of the same reaction.^ Have any been found in which this was impossible } 7. Compare the action of strong alkalies on alumina and on zinc. 8. When an alkaline carbonate is added to an aluminium salt the hydroxide is precipitated and the carbon dioxide is given off. Com- pare the corresponding results in the cases of magnesiura and zinc. 56 m;ai> gi'.ouf. CHMTl^R VIII. LKAI) C.kOlJI'. 1.— Oarbon, Silicon, Tin, Lead. Carbon is a non-mctalUc element existing in three entirely clifTcrent forms, gra[)hitc, amorphous carbon and diamond. It forms oxides CO and CO.,, the latter only being acidic. With hydrogen it forms a great number of volatile compounds. Sil^icon is either an amorphous grey powder, or a crystalline metallic looking substance. It is non-metallic in so far as its oxide and hytlroxide are acidic, never basic, so that salts of silicon are unknown. The only oxide is SiO,,, and the hydride is SiH^. Tin, a soft white metal, oxides basic and with strong bases weakly acidic. It forms both stannous and stannic salts, having the type composition SnClg and SnCl^. Lead, a soft, slightly bluish coloured metal, oxides are generally basic and form two series of salts, plumbous and plumbic. The former are far commoner than the latter. With strong bases the peroxide PbO.„ acts acid. As the acid properties of the oxides of these elements decrease, so does the stability, thus : COg, SiOg, SnO<„ PbOg. The first is not at all basic and is very stable, the last is scarcely at all acidic and is very easily decomposed. CARBON- - COMPOUNDS. 67 2.— Carbon. Occurrence and Preparation.— The three forms of carbon are found as minerals. Two of these arecrystal- hne, the diamond and graphite (plumbago). The other is amorphous and in the mineral form is coal ; prepared artificially it is charcoal, lampblack, etc. Charcoal is made from organic substances by heating under con- ditions which will not permit of combustion except of the gaseous constituents. Graphite has been obtained by mixing charcoal with molten iron. Charcoal burns easily when heated with oxygen ; the product is the dioxide, CO.^. When the supply of air is limited, the lower oxide, CO, is sparingly formed (compare phos- phorus). Diamond burns, but with difficulty, the product again is the dioxide, with a trace of ash. Graphite burns, to carbon dioxide, but only at a very high tem- perature. Oxides. 3.— Compounds. Carbon forms two well known compounds with oxy- gen, the monoxide, CO, obtained by direct union, or by the reduction of some higher compound; and the di- oxide, also formed either by reduction, or by direct union of its elements. The trioxide, C^Oij is unknown in the free state, but oxalic acid may be a union of this oxide with water, as shown by the decomposition with sulphuric acid, thus: H2C204 = H,0+C,03-^H,0+CO-f CO,. The monoxide, CO, is obtained [a) by reducing the di- oxide by red hot carbon, C + C02=2CO; ib) by reducing 58 LKAI) (;K(>ni'. iTictalUc oxides by charcoal, CuO + C CO-]-Cu. ZiiO-j- C = Zn+CO. l(CN),. and K.,Fe(CN),., or K,.Fe.,{CN)j^ are respectively the ferrocyanide and ferri- cyanide of potassium ; that is, they are both double cyanides of potassium and iron ; but in the former, iron is a diad ; in the latter, a triad See under Iron. Since cyanogen has a valency of one it is believed to have the composition — C rz N. In its combinations with hydrogen and the alkalies cyanogen resembles chlorine. Carbon and nitrogen do not unite directly, cyanogen being obtained as a reduction product. 4.— Silicon. Experiments. I. Powder some broken glass in a mortar (powdered quartz will answer as well), mix this with an equal quan- I SILICON'. 63 tity of fluor spar, also powdered. Put the mixture into a flask fitted with a cork and deUvcry tube which dips into some mercury in a beaker. Pour water into the beaker above the mercury, and sulphuric acid into the flask. Meat the flask. A gas should come off wiiich in the water forms a gelatinous solution. Filter this solution ; save both the filtrate and the solid on the paper. The clear fluid is hydro-fluosilicic acid, H^SiF,., and the pasty substance on the paper is silicic acid H^SiOp The reactions are : — (i) CaF,+ H,SO,=:CaSO,-f-2HF. (2) 4HF+SiO._,(or Si in the glass) = SiF^+2H.A (3) SiF,+4H,0=H,SiO, + 4HF. (4) SiF,+2HF=H,SiF,=2(HF)SiF,. These may be summed up thus : — 4CaF,-f4H,SO, + 2SiO,-H,SiF,.+ H,SiO, + 2HF4- 4CaS04. 2. To the hydro-fluosilicic acid of the last experiment add potassic hydrate until it is neutralized, then evapo- rate. Fluosilicate of potassium K,,SiF,;^K.,F.,SiF re- sults. 3. In a hard glass tube, heat to bright redness some potassic fluosilicate mixed with metallic sodium, about equal weights of the two should be used. After coolin-stals are left in the form of prismatic needles ; (r) when the solvent is molten aluminium in- stead of zinc the silicon crystallizes in scales much resembling graphite. Combinations. O.v/de^s and Hydroxides.— ThcYG is said to be a m.on- oxide of silicon SiO, but its existence does not seem to be established beyond question. The sesquioxide Si,0.; is known and is of interest as showing the relations of silicon to other elements. Silicon dioxide SiOo, occurs in nature as quartz. The hydrate corresponding to the dioxide is Si(OH) = H^SiO^, silicic acid. Oflicr Cowpoiinds.—Thc monosulphide SiS, is obtained from the disulphide SiS... The disulphide is prepared by passing vapour of car- bon disulphide over heated silicon. The tetrachloride SiCl,, is formed when chlorine is passed over a mixture of silicon and carbon. Sdidc Add and Sdicates.~S\\\c\c acid, like carbonic acid, is known only in solution, for attempts to isolate it always have ended in decom|)()sing it. Its salts arc both very numerous, important and stable. Silicates of sodium potassium, aluminium and calcium form many valuable minerals, such as the feldspars, micas, serpen- tine, hornblende, the zeolites, etc. In the arts, silicates are of importance becau.se a mix- ture of them makes glass. The principal ones employed 9 66 LEAD (J ROUP. for this purpose are those of sodium, potassium, lead and calcium. Alkaline silicates are soluble, others are not. Silicon hydride, SiH4, a spontaneously inflammable gas, is prepared by acting on a mixture of silicon and magnesium with hydrochloric acid. Silicon nitride SL^N.,, is formed by heating silicon in a current of nitrogen. Compare boron and carbon. 6.— Tin. Experiments. 1. Examine a piece of tin, — not the sheet of the tin- smith's shop, — but a block of the metal. Cut it. Heat it on charcoal in the oxidizing flame. After the oxide cools, heat it again, either on the charcoal or on mica. 2. Try if a shaving of tin will dissolve in nitric acid, in hydrochloric acid, either when hot or cold. Try aqua regia. 3. If a little tin be melted and dropped into cold water, and then boiled with strong hydrochloric acid, it will dissolve very slowly and give off hydrogen. It is better to drop into the acid with the granulated tin some bits of platinum, which will be recovered unchanged when the operation is completed. 2HCl + Sn-=SnCl.j+Ho. SnClg is stannous chloride. 4. Make a mixture of finely divided tin and mercuric chloride in a hard glass tube, and arrange that the vapour passing off may be received in a cooled condenser ; then heat. The liquid obtained is stannic chloride, SnCl^. NOTKS ON TIN. 67 5. Add liydr(JL,ren sulphide to the chlorides obtained in the last two experiments. 6. Make a solution of any tin salt and hang in it a piece of zinc. Repeat, using aluminium instead of zinc 7. Heat together in a hard glass tube some fine tin filings, sulphur and ammonium chloride. 8. Add potassic chlorate to a solution of stannous chloride, and heat. Test for a stannic salt with hydro- gen sulphide. 7.— Notes on Tin. Occurrence and Preparation. — Tin is very rarely found native. The chief ores are the dioxide, SnO.„ called cassiteritc, as a mineral ; and tin pyrites, a mix- ture of sulphides of copper, iron and tin. The ore is first roasted to drive off sulphur, then reduced by heat- ing with carbon. Compounds. — Tin forms two series of compounds, — the stannous and the stannic. The types of these are SnCl,, stannous chloride, and SnCl^, stannic chloride. Oxides. — When dichloride of tin, SnCK, is treated with an alkaline carbonate, an oxyhydrate of tin Sn.,0(HO)o is formed; this when dried out of contact with oxvinrn. l"'. Antimony and Sulphur. oa When sulpliidc of hydrogen is passed into a solution of an antimony salt, trisulj)hide of antimony, an orange precipitate, is thrown down. 2SbCl, + 3lI,S = Sb,S,-|-6lICl. This suli)hidc forms salts of an acid that has the formula SbS(SH) or Sb(Sn)., The trisulphide (stibnite) occurs in nature as a black compound ; when prepared artificially it is orange-coloured. The followinjr chemical changes are produced by proper treatment : — 2Sb..S.j + 90,, when heated = Sbp^. + 6S02 Sb,S, + 6HCl = 2SbCl,-}-3ir,S 2Sb,S, + 4HNO, = Sb,0,. + 6S-f-4NO + 2H,0 Sb,S.,+6H,, when heated = 2Sb-f 3 H,S. A pentasulphide SK,Sr^ is also known. IG.— Tests. (1) Orange precipitate with hydrogen sulphide, soluble in strong hydrochloric acid. (2) The chloride gives white precipitate with water, soluble in tartaric acid. (3) Antimony is distinguished from arsenic by these tests : — (a) The spots formed on the plate by burnin stibine are sooty looking ; CT il m 94 NITU0(;KN JMtoUl'. (/;) They turn orant^c when treated with animo- nic sulpliide ; (c) They are not dissolved by alkahne hypo- chlorites. 1 7.— Bismuth. Experiments. 1. Compare bismuth with arsenic and with antimony as to (r) physical properties, (2) burning- in air, (3) solu- tion in acids. 2. I3issolve some bismuth in stron^r nitric acid, and evaporate to drj'ness, without unnecessary heat ; result- ant powder is Bi(X().,).,, trinitrate of bismuth. Bi+ 4HN(),= Hi(\0,.)^j-|.NOH-2lI.,0. 3. Heat some nitrate of bismuth strongly, compare the result with that obtained by burning the metal in air, and with the coating on charcoal when bismuth is heated in the oxidizing flame. 4. Treat some of the nitrate with solution of potassic hydrate, Hi(i\0.,),+ 3KOH = Bi(OHj3+3KNO., If the hydrate be strongly heated it changes into the oxide, thus: 2Bi(OH).=Bi,0,;+3H,0. When Bi(()H)3 is heated to 100°, it changes into BiO(OH), bismuthyl hydroxide. 5. Drop some of the nitrate into a large quantity of cold water. This gives basic nitrate of bismuth (sub- nitrate), Bi(OM).,NO^. This basic nitrate when heated yields the yellow bismuthous oxide 2Bi(OH)2NOa= NOTKS OM lUHMUTII. 9n 6. Pass a current of suIpluirt.'ttcMl h\(lio^^cti tliroii^h a solution of any bismuth salt. Try the prccijjitatc for solubility in ammonia and in nitric acid. The black precipitate is I^i.,S.,. 7. Try if the sulphide is soluble in potassic sulphitlc or amnionic sulphide. Compare this with arsenic and antimony. 8. Treat a solution of bismuthyl liydroxidc with a solu- tion of bichromate of potash, K.X>^,()7-|-2Bi()(()II)= (BiO)Xr,0,-f2KOIi. 9. Dissolve some bismuth in aqua rcj^ia. Dissolve some of the trioxide in hydrochloric acid. In each case there is formed bismuth trichloride, I^iCl.,, which may be obtained by eva|)oration. When the solution of the trichloride is thrown into excess of water an insoluble white precipitate is formed, having the compositi(jn BiOCl, oxychloride of bismuth or bismuthyl chloride. 10. Add a solution of the nitrate to solution of an alkaline carbonate, 2Bi(NO,),+3KXO,=(BiO),CC\+6KNO,+2CO,. 18. — Notes on Bismuth. Bismuth is the distinctly metallic element of the nitro- gen group. Its chief compounds are similar in form to those of the other elements, but it does not unite with hydrogen, and its trioxide is not acidic, while the pent- oxide is only weakly so. The most important compound of bismuth is the nitrate, which is theoretically interesting because of its !!' (»rt NiTi{()«J»:V <:i(OiJi'. forming tlu; h.isii: salt hy oiu- (»r l\\(» hydroxyl {^TOiips displiiciii}^ ail n[un\ miinbcr of acid radicals, thus, Hi] N(),-f II.. i NO, [NO, Hi \()., [ HO 4- UNO., f N^O, r NO, or Hi N(),-f-2H,() Hi II() + 2 Ino" ' I HO UNO. Definition. — A basic salt is one in which the acid part of a molecule does not completely saturate the basic part. Nitrate of bismuth in solution with an alkaline hydrate yields a white powder, Hi(OH).,. On heating, tiiis succes- sively changes into HiO(OII) and Hi.,0.,. This oxide when dissolved in nitric acid yields Bi(NO,).{ again ; and when treated with chlorine and potassic hydrate forms IH^iOjj, ared powder. The acid when heated changes first into Hi./X,, then into Hi,0^. With an alkaline car- bonate, the nitrate forms what is commonly called car- bonate of bismuth, but what is in reality bismuthyl carbonate, (i^iO).^CO.j. 19.— Questions and Exercises. 1 . Roil some arsenic trioxide in the solution of an alkaline carbon- ate, K.^CO^ for example ; divide into two p'Mts, to one add some copper sulphate solution and a few drops of ammonia ; to the other some silver nitrate solution with ammonia. The former yields Scheele's green, CuHAsOg, cupric hydric arsenite ; the latter arsen- ite of silver, Ag-jAsOy. AS2O3 -f- 2K.CO3 + HjO = 2K.^H AsOa -I- 2CO0 K.HA'sOa-f-CuSO^^CuHAsOa-l-KaSO^. Write the reaction for the silver salt. QlTKSTIONM ANf) KXKKCISF.M. 97 car- car- ithyl ,S04. 2. Prrp.in' a s«»liiliMii (»f arscnioiis arid, by hoilinj^ arsenic trioxidc in wain. Duidf ii into two parts, tlwnii^'li (tne pass a ciirrnit of cliloriuf, into tlu- other drop some powdered iodine. Test both with nitrate of silver. Write e(|iiali()ns. 3. To a sohition of arsenic acid, add some sid|)hiiroiis acid, or pass a cmrent of snl|)hur chosidt; thronj^h it ; hoil. to expel the sulphur ^as, then test with silver nitrate solution. 4. Try if fusing t();;cther arsenious oxide and a nitrate n'ill give arsenic acid. 5. To a solution of an antinion ' salt, add liydrogen sidphide, heat the precipitate with hydrocMcric acid, ililute the solution largely. 6. Construct .1 table to show whet' -r the elements of the nitro- gen gr(»up change directly or invcisely with the increase of atomic weights in the following particidais : — (i) State of rigiilily at ordinary temjjeratines, (j) melting points, N - ; P, 45° ; As, 600.' : Sb 450" ; I'.i, 270. (3) Metallic appearance, (4) specific gravity, N =-.97, air= i, V i.g, water - i ; As, 5.7 ; Sb, 6.7 ; Hi, io. (5) Hasicity of oxides, (6) colour of the sulphide precipitated by H,S. 7. Do the conditions favour the formation of arsine, when a liydrochloric acid solution of arsenic is being decomposed by electrolysis ? 8. A current of chlorine is passed into solutions of ammonia, arsine and stibinc. What will take place in each case .'' 9. Some arsenical pyrites is roasted in an open tube, the white suljstance that collects on iIil cold part of the tube is scraped off and divided into two p.;rts, one part is boiled with water, the other with nitric acid. After cooling, both solutions are tested with silver nitrate. Chlorine is then passed through the aipieous solution and sulphur dioxide gas through the other. After standing for some time, they are again tested with silver nitrate. Trace the changes, giving equations. 7 98 NITHOCKN CUOUP 10. Write forimil.L' tor the following compounds : -disod'c phos- pliHle, iiinnioniuin nicta|)liospluitc (neutral), tlipiHassic hydiic sulph- antinioniate, antinionyl pota^sic tartrate (!'= liaC^H^O^). 11. (X())C1 is nitrosyl cliloride, (N02)C1 is nitroxyl chloride, (P())Cl;j is phosphoryl chloride, (SbO)Cl is antimonyl chloride, (r>iO)Cl is bisinuthyl cliloride. Write graphic fornuihe for each which will satisfy the valencies of the elements whose symbols stand first in each group. 12. A piece of phosphorus is acted on by strong nitric acid ; another piece is lowered into a vessel of chlorine and after a lime water is added ; a third piece is placed in acjua regia (see High School Chemistry). What will be the reactions in each case ? Give equations. 20. — Notes on the Nitrogen Group. The members all have the valencies three and five. Nitrogen and johosphoriis are non-metallic both in physical and chemical properties, arsenic is somewhat metallic in appearance, but its oxides are acidic. Anti- mony is distinctly metallic in appearance, and while its oxides arc generally acid forming, they are less strongly so than those of arsenic, and the trioxide acts as a base with strong acids, thus, — Sb(NO..).. is antimony n'trate and SbONO.. antimoii}! nitrate. Bismuth is still more metallic than antimony ; the trioxide is basic, and the pentoxide is feebly acidic, forming salts with strong bases, such as the alkalies. All except bismuth form hydrides having the compo- sition XII.,; these are soluble in water. All except arsenic form oxychlorides from the inter- action of the chlorides and water. All except arsenic form monad radicals of the composit^ion MO. NoTKs ov nn: mtrockv fjuotip. 99 } The follow ill- is a tahiilar stakMiiciit oftlu' oxides ami oxyacids of tlu* iiKMiihcrs (jf tliis j^roup : — Nitrogen. . . . - OXIDK. N()vN,()._, NO., :- (X,(),) N,0, riiosphorus . - P,0, U\.o., Arsenic . . . Antimony Bisnuitli 0, As.,(>. S}).,0, nii,o„ I Bi,0, 15i.,(), I'.iJ). Acid. HNO HN<),, coiresponds to tlu; nieta-Jicids of other nieniht'i's. H,P,(), HyPO,, at L'OU' H ,1V,(),, at 40(»" JIIM), HyAsO;,, known by its salts. IL,AsO„ ll,As,().,HAs03. HySbO., (oxide botli ])asic and acidie). H.,S1)0,, H,Sb,/)., JfSl)(),. HP.iO,., H,Hi,0- (f mm its salts). Amoii<^ tlic oxides of nitroj^^en N.,() when heated rcach'ly breaks up and acts as an oxich'zer, NO witli ox\- {^en chan<^es spontaneously into NO._. and N„C).,, while, on the other hand, these hi-her oxitles are readily- reduced to NO, and act as oxidizini^ agents. The oxides that are anhydrides form acids hy dissolviuL;" in water. 100 NITHOOKV OROI^P. Ill the case of oxides of phosplionis, tlie compoiuul funiiccl {lci)cnds on the masses takiiii;- part in tlie chemi- cal action. P.O., heated in presence of oxygen bums to l',0,. Both are anhydrides. When arsenic and antimony are burned in air they oxidize to the oiis- oxide only, and this does not combine with more ox\gen. To form the pentoxide, the lower oxide is treated with some strori^ oxidizing- atrent, such as chlorine or nitric acid. A simdar statement applies to bismuth. All the hydrides are oxidizable through the action of heat, both elements uniting with oxygen in each case, hence all are reducing agents. All the elements form sulphides of the same type as the ous and ic oxides, simply changing oxygen for sul- phur. All form sulphides, and antimony and bismuth form sulphates and nitrates. CHAPTER X. SULPHUR GROUP. Oxygen, Sulphur, Selenium, Tellurium. Only the first tw^o are to be considered. 1. — Notes on Oxygen. - Preparation.- Oxygen may be obtained by three general methods. NOTKS ON OXVGKN. 101 Or) hy electrolysis of oxides or lixdrates. Im)p this purpose the compound must be liquefied by fusion, if it is solid at ordinary temperatures. (<^) By the decomposition of oxyiren comjiounds by heat. Those more commonly employed f(^r this purpose are chlorates, nitrates, chromates, peroxides, and, in rare cases, simple oxides. (c) By displacement, as in the case of suli)huric acid and man^ranese dioxide, chlorine and water, fluorine and water. Compounds.— The simple compound formed by the union of oxygen with an clement is called an oxide. Oxygen thus unites with all elements so far as known, except fluorine and bromine. When tlie oxygen of an oxide is replaced by hydroxyl, HO, an hydroxide or hydrate is formed. It sometimes happens that an ele- ment of well-defined valency unites with oxygen in a proportion greater than that required by the valency of the element, thus H.,0.., PbQ.., Na.,().. are examples. Such compounds are called peroxides, because with acids they do not form salts of the type of the oxide and water ; but oxygen is freed, and the salts corresponding to the lower oxide are produced ; thus,— 2Na..O ,+j MCl^ 4NaCl+0,+2H,0. The peroxides are, in general, somewhat unstable, parting readily with oxygen, hence their value as oxidiz- ing agents. The oxides are sometimes classed as toVr, (;p witli water, make acids. I'L-roxidcs do not iiciitralizc acids, thou<'h thcv ma\' form In'dratcs with water. This division is not a satisfactory one because they over-lap one anotlier, freciuentl)' the same oxide bein<^ botli aci(h*c and basic, depending; on the otlier substance takini^ ])art in the action. Tlie other cliief compounds of oxyi;en are tlic oxyacids and tlicir saUs. Some of tliese, sucli as tlie nitrates, clilorates and permanL;anates, are easily decom- posable and )'ield free oxyi^en, but most of the com- pounds arc not readily broken up in that way. Ozone. — This allotropic form of oxygen is supposed to be ^\\(^ to a different atomic arrangement of the molecule from that in free oxygen. The symbol O.^ exjiresscs the supi)()sed fict that the molecule is triatomic. The alteration in \-olume when oxygen is changed into ozone or vice versa is in accord with this hypothesis. Any divalent ijascous element freed from its combinations might take on an allotropic form due to a similar group- ing, but with monad elements it is not theoretically possible. 2.— Notes on Sulphur. Preparation. — Besides the sulphur which is found native, this element is separated in many chemical re- actions, such as by the decomposition of hydrogen sul- phide cither in presence of oxygen, chlorine or sulphur- dioxide, thus, — . 2lLS + 2rip4-0,-4lLC)-f2S ILS+CL 2IICI + 3S so.,+2ii.,s :2n..o-f3S NOTKS ON .srr-IMIl K. 103 by the brcakinnr down of the polysulphidcs of alkali metals by acitls, sec \y,v^Q 27, and by the decomposition of thio-sulphates, thus, — K,,s,c), -f- 1 r,so, K„SO, + 1 i._.o+so,+s. (See H. S. Chem., p. 180.) This is the basis of one method for the recovery of the sulphur in the alkali waste of the soda works. (Sec under sodium.) The calcium suli)hi(le is chan.<,red to thio-sulphatc by blowini^r air throu«Th it, then with hydrochloric acid it is decern'^ posed, thus, — CaSA+2HCl-:CaCl,+ H.,0+2S. Like the oxides, many sulphides and polysul|)hides arc decomposed by heat, but frecjuently the sulphur at hiijh temperature unites with oxygen to form SO.,. Properties.- -Sulphur, in its combinations, very much resembles ox)'gen. It also is known in a number of allotropic forms, which are probably due to various atomic groupings in the molecules. The well-known modifica- tions in the fusing and suddenly cooling of ordinary .sulphur can be easily ob.served by very gradually melting sulphur in a tube or dish, stirring it from time to time and dropping a little of it into cold water. The vapour density of sulphur also indicates that the molecule up to a temperature of 500 C. is hexatomic, that abo\e this temperature the molecules begin to break down into diatomic ones, and at 1000° this decomposition is com- plete, S,.---^3S^,. This would .seem to indicate that the atoms composing the molecule are arranged in groups, and that the increased kinetic energy cau.ses the attach- ment between the groups to be broken. lOi THK SI 1,1'IIUU (SKOUP. 3. —Compounds of Sulphur. Sulphur and hydrogen, ZnO + 2lI(1 - ZnCl.,+ n,() ) ^^ FcS + H,SO,-FcS(\-f H,S / ' Na..(),+2HC1 2XaCH-H..()., | j^ K,S,4-2HC1 -2KCl-)-n-.sJ+3S/ 2lI + 0-II,0 ] 2H+S-H,S J The equations A and 1^ sliow the similarity of the methods by which oxides of hydrogen and the sulphides of hydro<;en are obtained. When hydr()<;en is i<^nited in oxy,i;en j^as the oxide is produced, but when the ignition occurs in sulphur vapour the suljohide is formed, as in C. The sulphide is easily decomposed, as when the gas is passed through a red hot tube, or ignitt-d, or passed into a solution of most metallic Scdts, or in presence of a halo- gen. When the gas is ignited, if the sup[)ly of air be limited, the hydrogen burns, but part, at least, of the sulphur remains unconsumed, as a deposit on the vessel. The persulphide is of somewhat doubtful composition, and as prepared from potassium pentasulphide, contains free sulphur. It is mentioned here simply for com- parison with the oxide. 4.— Oxides of Sulphur. These are the dioxide, SO.,, the trioxide, SO.^ the ses- quioxide, S.,0.,, and heptoxide S.,07. All may be produced by union of the elements. Sulphur ignited in oxygen forms SO.,. SO., heated with ox)'gen over a platinum sponge gives SOjj. SO., cannot be obtained In- burning ACIDS OF SULl'HUR. ion C sulphur in air, because the trioxidc, at a temperature below that of burning sulphur, breaks down into the dioxide and oxygen. SO3 fused with sulphur produces S.P; is formed when electric sparks pass through sulphur dioxide and ox\-gen. The oxides are mostl>' prepared, however, by reducing higher compounds. The dioxide and trioxidc alone are important. The former is obtained by the reducti(jn of sulphuric acid by the nascent hydrogen formed when a metal acts on part of the acid, thus : Cu + H„SO^=:CuSC),, + 2H Hydrogen sulphide acts similarly, H.,S + 1 i.,SO, -- 2H O-f SO2+S. " " 5. —Acids of Sulphur. Hydrogen sulphide shows acid properties. Its hydro- gen is replaceable by metals, hence the formation of metallic sulphides and hydrosulphides, as CuS and KHS. Ox}'crct^s.-~The two most important of these are obtained by dissolving the dioxide and trioxide in water. They are sulphurous, 11,80.^, and sulphuric, H.,SO,. The preparation of the latter depends on the propert\- that sulphur dioxide possesses of reducing nitrogen tetroxide in presence of water, thus forming sulphur trioxide and nitrogen dioxide. Sulphuric acid, when boiled, loses water. The hy- drated acid having the formula H.,S(), + 4H..O has been separated ; but, by heating, it parts with successive portions of water, as indicated by the formula: H^SO^-j- Km; THR SULl'HUR (WiOUP. 3ll,(), M,S(),4-2lIA ii,so,H-n,A H_.so„ n,s,07= 2lI._,SC)., — 1 1./). Phosplioric anhjdridc, l^.O^, when mixed vvitli sulphuric acid and distilled absorbs water and allows the trioxide to pass over. It is advantageous here, of course, to use acid already largely dehydrated by heat. II,SO,+ PA-2lIPO,-f SO3. The method of preparing the commercial acid has alreatly been described (sec 1 1. S. Chcm., page 183). If sulphur be oxidized to the trioxide, this, with water, will give the acid. The trioxide cannot be produced by heating sulphur, but it may be obtained by (a) oxidizing the dioxide as in the case where nitrogen trioxide is reduced, or when a halogen is dissolved in an aqueous solution of the dioxide, (/;) by oxidizing sulphur directly, as when it is boiled with nitric acid, or when it is added to a solution of chlorine in water. S + 6HNO, SO,+ 3H,0 + 6N02 or S + 4HNO,-SO, + 2}I,0+N,03+2NO, S+3C1,4-3H.,0-S03+6HCI. It will be remembered that hydrochloric acid is pre- pared by heating a chloride with sulphuric acid, but with a bromide or iodide, a further reaction occurs by which the halide acid acts on the sulphuric acid and decomposes it, at the same time depositing the bromine or iodine. Water is also formed, along with sulphur dioxide, free sulphur or hydrogen sulphide, depending on the quantity of the halide acid present. Probably all three reactions generally go on together. A similar decomposition occurs with hydrogen sulphide. A piece of zinc dropped TIIIO AfMDS. 101 into hot concentrated sulphuric acid brinL;;s ahout such reactions as the following, — 5Zn + 5lI,SO,, = 5ZnSO,+ ioH 2H4-2ir,SO, = 2lI,0+H,S4-SOo H,S + H,S0j-2n,0 + S0,-fS. Suljihurous acid II.,SO.< has not been separated, as it breaks up into its anh\dride and water when either evaporated or displaced from its compounds by a stronger acid. Its salts are important. 6. - Thio-Acids. Sulphurous acid oxidizes to sulphuric. Similarly, if a salt of sul})hurous acid, Na^SO.j, be boiled with sulphur it becomes NaoS._,0.., thio-sulphatc of sodium, — the sodium salt of thio-sulphuric acid II^SoO.j. This acid has not been isolated either. In popular language it is hypo- sulphurous acid and Na^S.^O-j is hyposulphite of soda, but the names just quoted should belong to H^SCj and Na.^,SO,,. There is a series of the sulpho-acids of sulphur, which may be mentioned here. They are, — HO thio-sulphuric acid, H„S.,0., = SO.,< while sulphuric HS, HO acid is SO.,<[ HO. H.,S..O, H,S30, dithionic acid, trithionic acid tetrathionic acid, ILS40,j pentathionic acid, ILSsOy 108 TIIK SI LI'IIUK GKOUP. 7.— Sulphur and the Halogens. When dry chlorine is passed over heated sulphur, as powder, sulphur monochloride S._,Clo is formed. By satu- ratin<( this monochloride with chlorine gas at low tem- peratures, the dichloride SCl.j and tetrachloride SC\^ are obtained. S.,Bi'.j and S.,I._, are also known. It will be noticed that in these lower halide compounds, sulphur has a valency wiiich it does not exhibit in the oxygen compounds ; the SCl^ corresponding in this respect to SO.,. Just as in the oxygen compounds, one atom of oxygen may be replaced by two hydroxyl groups, — S ]^ becoming S \ ,\ - = lI.,SO,j, so the oxygen may give place to chlorine, thus S .^ changes to S p, thi- onyl chloride. Another somewhat similar compound is sulphuryl chloride, which contains the group SO.,, and may be looked upon as sulphuric acid in which hydroxyl is replaced by chlorine, thus SO.,| ..^x by substitution becomes SO., /CI _ ^ - [CI SOXU sulphuryl chloride. Natu- rall)', if we have SO/HO), changing to SO.,Cl„one looks for the intermediate compound, when one hydroxyl group is replaced. This compound is also known. It is SO. fOM \C1 SO./OH)Cl, chlorsulphuric acid. IRON. lOO c:haiti-:r \i, TFIK IRON C.kOUl'. 1. The elements rcstMnbliii^^ one another sufficiently to be put in this ^n,up are chromium, manj^aiu-se, iron, nickel, cobalt. Of these onl)- man^^inese and iron are to be considered here. Both are greyish white metals, capable of bein^ re- duced from their oxides by the action of carbon at hi^di temperatures. I^oth are fusible and ductile ; man<4anese is much harder than iron though. ]^oth are base forminL-- as well as acid formini,^ elements ; in iron, however, the basic properties are much stron^^^er than the acid ones, while in the case of manj^anese, tiiough the basic proper- ties predominate, they are not so marked as with iron. These elements, as bases, form two series of com- pounds, — the -OHS compounds in which the metals are bivalent, and the -ic compounds in which the)- arc tri- valent. The types of these are MO and M.,0,j. 2.— Iron. Experiments. A — The Ferrous Compounds. I. Put some cuttings of iron wire, or some coarse iron filings, into a beaker and cover with dilute hydrochloric acid ; keep iron in excess until gas ceases to come off from the liquid when heated nearly to boiling. The no TiiK \nos rjKOUP. fluid slujiiM Ik- hliiisli ^ncni in colour. I^'ill a well stoppered hotlle with this licjiiid and preserve U)r future use. The fluid is a solution of ferrous chloride, I'cClo, FeH-2lK:i = FeCl,+ n,. 2. Set a portion of tne solution of ferrous chloride contaim'n*^ some free acid apart for a few days in an open vessel. It changes colour and, by taking up oxy- gen from the air, becomes ferric chloride, I^'e^Cl,., thus : — 4FeCl, + 4HCl4-(), = 2Fe,Cl, + 2M,0. 3. Pass throu<^h ferrous chloride solution a current of chlorine, 2FeClo+CK,= FeXl,j. 4. Prepare a strong solution of ferrous chloride, in ex- cess of hydrochloric acid, in a lary^e beaker; add nitric acid, and heat until tho mixture foams up and turns reddish brown. 6FeCl, + 2H N(),4- 6IIC1- sFeXl, + 2NO+4H2O. The foaming at the end of the operation is caused by the escape of nitric oxide gas which is formed but is not held in solution by the ferric chloride. Save this chloride for further experiments. 5. To ferrous chloride add a drop or two of silver nitrate solution. The greenish substance is ferrous nitrate. Can it be prepared by acting on iron wire with nitric acid ? 6. Heat some iron with dilute sulphuric acid, evaporate to dryness. The green salt is ferrous sulphate FeSO^, (copperas). 7. Dissolve some ferrous sulphate, add ammonium hydrate. The precipitate is ferrous hydrate. Try if other alkaline hydroxides will yield the same compound. Try lime water. IIKIN. Ill 8. Pass hydrogen siil|)hi(l(r iiito a solution of any ons salt of iron in presence of free acid. Repeat, hut make the solution alkaline. 9. Add a solution of an alkaline carbonate to ferrous chloride. Will the sulpiiate or the hydrate ^ive similar results? The greenish salt is ferrous carbonate. /> — I'/ic I^'ernc Compounds. 10. To some ferric chloride add iron filings and let the . ilxture stand, but shake it frccpieiitly. ('omi^are the result with ferrous chloride, Fe.,Cl,.-j-Fc -3FeCl.,. 11. Try if chlorine or nitric acid will oxidize ferrous sulphate to the ferric compound. 12. In one vessel put some pieces of iron wire with dilute nitric acid ; in another some iron wire with strong nitric acid. The ferric nitrate, Fe.j(NO.j)g is formed in the first case. 13. Heat strongly on a piece of mica, a crystal of copperas. 14. Prepare some ferric hydrate, Fe./OIT),;, by treating ferric chloride with ammonium hydrate. Filter out the precipitate, divide it into two parts and ignite one of them strongly on mica. Fe.^(( ) 1 1 ),. = Fe^O.j -}- 3 H.,0. Test the hydrate and oxide for solubility in hydrochloric acid. 15. Drop into a solution of potassic ferrocyanidc, (yellow cyanide), K4FeCy,i, some ferric chloride. Prussian blue is formed. Repeat the experiment, but use a ferrous salt. Again repeat both parts, but use ferricyanide, red cyanide, K^FcgCyj.^. ' ' -^ THK IRON GROUP. 16. Try if ferric chloride can be reduced by (^) pass in p" throu-h it a current of h)'drooe„, (/;) FeO^ has not been separated, but the potassium sodium, and barium salts of it are known, K^FeO^, Na^FeO^ and BaFeO^. (See ex. 20.) Chlorides. — Both the ferrous and ferric chlorides are easily prepared, the former, in solution as a bluish green liquid, by the action of dilute hydrochloric acid on iron ; the latter as a yellow or reddish compound formed by the action of oxidizing agents, such as chlorine or nitric acid on the ferrous chloride, or by passing chlorine over hot finely divided iron. This latter process may be reversed, and by means of reducing agents, such as sulphurous acid and nascent hydrogen, the ferric salt may be changed to the ferrous. (Compare arsenic and antimony.) SuIpJiates. — The ferrous sulphate, FeS04H-«HoO is COMPOUNDS OP rilON. lin lonncd when ir„„ in excess is dissolve,! in sulplniric acid. Tlie quantity of water of crystallization taken up abso- lve alcohol. This salt is formed in nature by the action of damp air on pyrites, FeS,. Ferric sulphate is obtained by oxidation of the ferrous salt, or by dissolving ferric hydroxide or oxide in sul- phuric acid. Both sulphates form double salts with the alkalies, thus Fe(Nal(SOj.,+6II..O and FeXa'SOJ + I2H„0. The latter series (ferric) form the iron alums". Sulphides. -T:\,s ferrous sulphide is obtained either by roasting together iron and sulphur, or by the action of hydrogen sulphide on an alkaline solution of a ferrous salt. I-rom a ferric salt the hydrogen sulphide precipi- tales the ferrous sulphide and sulijiuir. FeCL+(NH,XS = FeS + 2NHCl FeXl„+3(NH,),S--2FeS + S-|-6\II,Cl. Ferric sulphide Fe.,S, is an unimportant comijound. Iron disulphide FeS, is iron pyrites, a very common constituent of rocks. When heated to redness ,t is re- duced to the monosuiphide. Arsenical pyrites, mispickel, is a sulphide of iron and arsenic, FeAsS. Cyanides.— C■.^.xhoxx in combination with nitrogen forms cyanogen, C.\, (Vy). The acid of this is HCN, hydro- no THE IHON GROUP. cyanic acid. There arc two double cyanides of iron and potassium. Ferrocyanidc, Kjl^'eCy,., and ferric)iinide K.T cCy,3 = K,T e.,( y,j. The former is a yellow cry.stalline solid, the latter is red, (respectively yellow and red prus- siate of potash). With a ferric salt the ferrocyanide yields Prussian blue. With a ferrous salt, the ferricya- nide yields TurnbuU's blue (li^ht coloured). 5.— Tests. (A) — Ferrous Couipomids. (i) Alkaline hydrates precipitate I'^efOII).,, which soon chani^es from white or grey to reddish brown, Fe./OH),,. (2) Alkaline carbonates precipitate ferrous carbonate FeCO.,, white or grey in colour, changing to brown Fe,(oi^),, (3) Ammonium sulphide precipitates black sulphide of iron, FeS ; hydric sulphide does not, except in alkaline solutions. (4) Ferricyanide precipitates a dark blue solid, in- soluble in acids but decomposed by alkalies. (B) — Ferric Compounds. (i) Alkaline hydrates precipitate ferric hydroxide Fe,(OH),, (2) Alkaline carbonates also precipitate the hydroxide. (3) Mydrogen sulphide reduces the ferric solutions to ferrous with the precipitation of free sulphur. MANOANKSK. 117 (4) Ammonium sulphide reduces the ferric sokition to ferrous and deposits free sulphur, but at the same time throws down the ferrous sulphide as a black precipitate. (5) Ferrocyanide precipitates Prussian blue, insoluble in acids, but decomposed by alkalies. (6) Sulphocyanides, for example KCXS, give blood red colour to ferric solutions, but not to ferrous. Ferrous compounds and their solutions are generally green in colour. Ferric compounds and their solutions are reddish, yellow or brown. 6.~Manganese. Experiments. Manganous compounds, type MnCl,. I. Manganese dioxide heated with sulphuric acid yields oxygen. When this gas ceases to come off, the solution contains manganese sulphate. Filter and add to the filtrate slowly an alkaline hydrate solution. If a precipitate is thrown down filter again, and evaporate some of the filtrate, which should be pink or reddish in colour. The precipitate with the alkali consists of iron contained as an imi)urity in the dioxide 2]\InO..+ 2H,SO^--2AInSO,-f3lIX) + 0,. J'rcserve the rest "of the filtrate. 2. If the residue from the prcjjaration of chiorino, by the action of h)-drochI()ric acid on manganese dioxide, be filtered, and the iron, if any precipitated, as in ex. i, the solution will be one of manganous chloride, MnCl.. lift TUF IKON GROUP. 3. Add U) cither the suljihatc or the chloride solu- tion some sodium carbonate, M nSO, + K^,CO;;= iMnC().j4- K0SO4. The carbonate is at first white, but soon becomes dark coloured by oxidation. 4. Repeat the last experiment, but use caustic soda instead of the carbonate. MnS04+2Na(OII) = Na,SO,+ Mn(OII),. This also oxidizes, thus,— 2Mn(0H),,-f 0,.= Mn,0,(dH),. 5. Add ammonium sulphide to manganous sulphate solution, reddish manganous sulphide is formed. Manganic compounds, type Mn.,Cl,;. 6. Melt on a piece of mica, or in a small iron vessel, some solid caustic potash, stir into this, when fused, a mixture of manganese dioxide with a little powdered chlorate of potash. When heated to redness, the mixture should turn green in colour. 3KOH + KC10,+ MnO, = K,MnO, + KCl + 0,-f H,0. K2Mn04 is potassic manganate, the potassium salt of manganic acid, IL.MnO^, which has not been isolated. Compare ferric acid. 7. Repeat the last experiment, but use sodium carbon- ate, manganese dioxide and potassic nitrate. 8. Dissolve some of the careen substance and heat it to boiling. Let the liquid then stand for a day. The red liquid is solution of KMnOj, permanganate of potash ; and the powder is manganese dioxide. 3K,Mn04+2H,0 = 2KMnO,i+MnO,H-4KOH. NOTKs ON manganp:se. 119 7.— Notes on Manganese. Manganese is a grey metai, like cast iron, very hard and oxidizes spontaneously in the air. It is fusible, and' soluble in acids. It is prepared from the oxide MnO by a process of reduction similar in principle to that employed in the case of iron. Occurrence.-The chief compounds of manganese found m nature are the black oxide or pyrolusite, MnO, • braunite, Mn,03 1 hausmannite, Mn^O, ; and mangane^^e spar, MnCOg. Compounds. ar/^r^. —Manganese forms the following compounds with oxygen : Manganous oxide MnO, manganic oxide Mn,0,, mangano-manganic oxide, Mn.,0, = MnOMn O manganese dioxide or peroxide MnO.„ and manganese heptoxide Mn.O;. Of these Mn..O, and Mn.,0, are respectively the anhydrides of manganic and perman- ganic acids. These oxides connect manganese, on the one hand, with chlorine by the similarity of the heptoxides and the per ic acids ; and, on the other hand, with lead which has an almost analogous series of oxides, PbO, PU,0., Pb,0, and PbO.,. When treated with hydrochloric acid the actions with lead oxides, too are similar to those with the oxides of manganese' There is, however, no Pb.O- known. Manganous oxide, MnO, is prepared from any of the higher oxides by reducing them in a current of hydro- gen (compare iron). ^ Permanganates.—OiWxGsi,, the most common is potas- sium permanganate, formed by acting on potassium or sodium manganate in some way to remove one atom of 120 THK IKON OKOUP. the alkaline metal. This may be done as indicated in ex. 9, or by passing a current of carbon dioxide into the man<^anate solution. Potassium permanganate is a most useful oxidizing agent on account of the readiness with which it yields up oxygen and becomes reduced to the manganous salt. Ferrous chloride, for instance, will decolorize a solution of the permanganate, and at the same time become oxidized to ferric chloride. The uses of the perman- ganate as a test for the purity of water, and as a disin- fectant are dei)cntlent on this property. The organic substances in the water deoxidize the salt, thus destroy- ing its characteristic colour, and noxious germs are similarly oxidized and burned up by it. A piece of soft paper dipped in the solution of permanganate will soon turn brown, owing to oxidation of the organic in- gredients. The following are some examples of this re- duction :—2KMnO,+ 2FeCl,+ 8 IiCl= 2 MnCU+2KCl+ 8H,0 + Fe,Cl,, 4KMnO,-f 4FeSO,-f-8n,SO,-:4MnS04-f 2K,SO,+2Fe,(SO,),H-4lI„0. As a disinfectant, its action may be illustrated by the equation, 4KMnO,-|-6II.,SO.j -4MnSO^-l-2KoSO.j + 6H._,0-f 5O.,. The ox)'gen is, of course, nascent, and the large number of atoms set free from each molecule probably exi^lains their grouping to form ozone, as well as oxygen in the molecular condition. The permangam'c acid may be obtained in solution by treating the barium salt with sulphuric acid. BaMnO,4- H.,S04=BaS0, -}- 1 1.MnO,. NOTKS (>\ MAN(3ANK,SIC 121 When the j)cmi.inc,mnatc of i)()tash is (h-oppcd into a concentrated s(;lution of a caustic alkah" and wanned it is reduced to the man<,^anate, 4KMnO, + 4K()n = 4K,MnO, + 211,0 -fO,. After the sohition becomes quite clearly green, if it be poured into excess of water it decomposes, as in ex. 8. So lonir as the alkali is strong the decomposition of the manganate does not occur. Other Compounds. MaugiXJious chloride, MnCl,, is obtained by the action of hydrochloric acid on any of the oxides, as in prepara- tion of chlorine. Manganous sulphide is formed in alkaline solutions by hydrogen sulphide. It is characteristically pink coloured, as are many of the manganous compounds. The hydrates.— 1\<. in the case of most other heavy metals the h\'drates are formed by treating the soluble salts with soluble hydroxides. The tw^o most important ones are Mn(OH),, manganous hydroxide and iMn.(OH)^ the manganic hydroxide. Manganous sulphate is formed by the action of con- centrated suli)huric acid on manganese dioxide. The ordinary experiment for the preparation of oxygen from these substances leaves as a residue a solution of the sulphate which may be filtered out and purified by precipitating the iron with calcic carbonate. ■ppm 122 THE iHON <;uour. 8.— Tests. (i) Alkaline hydrates precipitate a white hydroxide that soon turns brown. (2) A manganous salt boiled wMth nitric acid and lead dioxide gives red permanganic acid. (3) Any salt of manganese ignited with caustic potash and a little chlorate of potash gives a bright green mass. 9.— Questions and Exercises. 1. I'ass a current of hydrogen over red hot manganese dioxide in a hard glass tube. 2. Pass a current of carbon dioxide over hot iron filings. Test the escaping gas for carbon monoxide. 3. Boil a little ferrous chloride with some hydrochloric acid, and drop in a crystal of chlorate of potash. 4. Drop into some ferric chloride solution a little sulphuric acid and a piece of zinc. After a time test for a ferrous salt. 5. Mix with a purple solution of permanganate of potash a newly made solution of sulphur dioxide. 6. Pass a current of carbon dioxide into a green solution of manganate of potash. , 7. Add a little barium dioxide to an acidulated solution of per- manganate of potash. Ba02 + 2HCl = BaCl2 + H20o and 5H202+6HCl + 2KMn04=2MnCl2 + 2KCl + 8H20 + 50.. CALCl'LATIOV op ATOMIC WKKillTS. 123 CHAPTKR XII. CALCULATION OF ATOMIC WKIGIITS. There are several methods used in the determination of atomic weights ; the particular ones to be applied in any case depend on the physical and chemical pro- perties of the element. These ways may be briefly summed up as follows :— (a) By determining the masses of equal volumes of elements in the gaseous state and under like conditions of temperature and pressure. This cannot be relied upon for accuracy. (/O By determining the proportions by weight in which elements replace one another. This is valuable, because with proper precautions it gives reliable results. (c) By determining the proportions in which the ele- ment whose atomic weight is being found unites with a large number of other elements. For example, carbon in the following compounds yields from Carbonic oxide 12 of C and i6ofO. Carbon dioxide 12 of C Carbon tetrachloride 12 of C Methane 12 of C Carbon disulphide 12 of C Alcohol 24 of C 32ofO. 142 of CI. 4 of H. 64 of S. 22 of other elements. The smallest proportion by weight in which carbon enters into these combinations is ,2, hence the smallest 121 CAL('ULATI()N OV ATOMIC WKKJIITH. part liy \\'ui;4ht in a molecule is I3 times the \vciL,Hit of an atom cjf h)'(lr()L,rcii. It docs not follow fr(jm this that the atomic vvci;^ht is 12, because there is no evidence that there is only one atom of carbon in any of these molecules, but it docs follow that the atomic weight cannot exceed 12. Hy comparin*^ as many compounds as possible, it is prol)able that the atomic weight will be the least found in any compound. (d) By Dulong and Petit's law. This applies only to solid elements. Atomic heat (6.4) divided by specific heat gives atomic weight. (e) By the principle of isomorphism. This is of use only as a check, and not a very reliable one, upon other methods. (/) }ly comparing analogous elements and their posi- tions in the tabulated lists of the periodic law. (£•) By the spark spectrum of the substance (De Bois- baudrun). (//) By the observed lowering of the freezing point of a solvent in which a weiglied portion of the element is dissolved (Raoult's method). These last two methods are scarcely suited for des- cription in a book of this kind, hence are only alluded to in order to complete the list. .iUi.- PFItroi>tf' LAW. 125 cii.\n"i:K Mil. I'l'lklODK I.WV. 1. When the elements arc written in the order of tlieir atomic wei-^hts iln-y fall into certain rcL^ular groups, such that tiiose in the same ^rroup are closely related to one another in their properties and in the character of their combinatior.s; hut are sharply sejiarated from those of nei-^hbouring L^roiips. For convenience tliese <^n-oups may he written as in Table 1. 2 The elements as now arran<;cd consist of two periods of se\en each, two others, almost complete, (jf seventeen each, and three more of the lar<^er i;roui)s, but with a considerable number of omissions. The larages.) lHTPKR,o,,_Li7. B9, Bll, C12, N 14, O 16. F 19 2N.>PKK,o»-Na23. Mg 24. Al 27, Si 28, I> 31. S 32. CI 35.5 3Ri.PKK,on-K39. Ca 40. S. 44, Ti 48, V 51, Cr 52, Mn 55 4TH Period -Kb 85, Sr 87. Y 89, Zr 90, N,. 04, Mo 96. -100 5Tn Pkriod-Cs 133. Ba 137. La 139, Ce 140. Di 142, _ OTiiPKRiOD- - _ Yhl73, - Tal82, W 184. - 7th Period— — Th 232. Ur 240. Fe 5C Rul04 Os 191, siicli the •h a \ a ^ '*^/"^^' • ^-V-- »^-- — •>' -ill cause a few of the rarer eU.n.ents, h as La. Ki. V h a.ul I), lo he .•han.^e.l sliyhtly fro,,, their prese.il p..itio,.s 1.,., . . ohjeet here is snnply to explai,. the tables, they are «ive,. as usually printed' nieiits, iMit as PERIODIC LAW. 127 Fe 50 Rul04 3s 191, Co r.8, Ni 58 Rh 104, Pd 106 Ir 102.5, I't 194 C"63. Zn05. Oa TO, Oe 72. A.;,. «, ,,, „, ,, Agios. CdU2. I„n3. s„ll7. Shl20 Te 126, I 127 ■— — — Er 166, — _ _. Au 107, Hg 202. Ti 204, I'l, 20(5, J{i 210, - _ 1-28 Pkiuodic law, r CO i-H 00 o O -f S ^ 5 ? = »— 1 X 1— ( 1 M CC OJ ■^ 1> z'. 1- •/i a; •M XI o r-i s -M 1- ;^ •^ • I^ 1—* ■■£) ■"■ ••*• ^ :) TO 3J I— ' 1-H ■M 5 ?; 1-H -M -f s; X I-H U: o. ?1 O L. a> j= H N '-J H *— 1 1- ■^ ^ zt ^M TI 1- I-H T-H 51 0. d' -< ^ T^ H o 3i ?^ ? r-l I-H O X >" > iC 'M ^ — » r-t ■ ^ I-H 71 fi. - 05 ■M N U (^ 1 ' 5 »^, K o ^ « f. -<" r) l-H 1 O cq 1- »^ r 1 I- » * 1 ■55 1 *— , hJ ^ 1-H ^^ i. ^ ^ ;j be io 1 S '■— :) ^ I-H "5 ^ ^ ?5 I-H y^ :i K _ K , >— \ , /^ ,^-v A -^*v ^^ c; ^..^ c ^^-^ c • c ■^ '^ — 0/ '^ K '^ c X X 5 5 01 5 0) 5 75 E V. 1. •r. •JI 73 •c 73 1) ■c 73 O' 'ill D ■r. •c 01 72 0> O «— ' 'C A X J= Si SS "w •^ ^T^ *-> 4-t *.-> 1—* ■M t 10 •*' I- iJO - I-H ""• I PKK IODIC LAW. 3. 129 sho^ s t f ' ; ': ""^^^^" arran.cme,.t which shou.s the relationship of the groups of elements in a somewhat better way, because of the positions in the schedule. Hydroiren is omitted because it is the only representative of its series ^ I i^ ■ii t _ I. II. III. IV. V. VI. VII. VI 11. TABLE III. Cr ( Mn Mo Ru Fe K Rb Cs Ca Sr Ba i Sc Y La iTi Zi Ce V Nb Di I. Lf Na Cu Ag C(l II. . He M- ;Zn III. n . A\ iGa In IV. c N Si • P ;Ge Sn V. . As • Sb VI. S Sc Te , vir. F CI Hr I 1 2 3 4 o Yb Ta W Os All Hg TI I^b JJi Th U o a n> n' to < 3 n> J 9 ^IWHMWI 130 PERIODIC LAW. 4. —Comparison of Properties. Mere the vertical columns show the periods (horizontal rows included between the lines) of Table II. The two rows having the same Roman numeral make up one of Mendelejeff's groups (vertical columns) in Table II. In th' table fill.) the following are important points to be oc ) : From element tr element within each period (verii ' column) there is but slight change of properties, yet from the first element to the last one in the period there is all the change from a strong basic to a strong acidic substance. The first element of each period is an alkali, the last is a halogen. The last element in any period differs entirely in its properties from the first of the next. The sub-groups of the large periods begin in the even series with strong basic elements, and end with others that are both basic and acidic (Mn, etc.); while in the odd series the first elements are ordinary metallic bases, and the last ones strongly acid forming. On account of the relationship existing between the members of the groups (horizontal rows) and the small differences between adjoining members of the periods, it might be expected that groups of similar elements would be found together. That this is really the case is shown by the heavy metals being enclosed in the lower dotted space, while in the upper one is a group of metals all rare, occurring together, and very much alike in properties. Immediately adjoining the lower rectangle are some of the light metals, and metallic like elements ; while the other alkaline metals are next the upper rect- VALKNCY IN TlfK (iKOtJPS. 131 anrrlo. The non-mctallic elements arr. in fi, ♦• J'i » ir:sr:rrn;;r.°,:,,;';: '"- and sfmn— P--' of non-metals, but metalhc elements mcrea.se in number. 5. -Valency in the Groups teln I' '"''""'^ '■'•""'^'"^■^ '' maximum with the tetrad carbon L'rouo Th,. „.f. VI is chad .nfvu- ^'" S"-™'!'. v., is triad, -.ff:i,:;^d:o;r,-^^ ployed. "°"" ™°"=''^ "'<'<^«1^ "--e em- The o.xides show a regular increase fron, the first croup t tne lou.th group, on account of two .se-ies of oxides bem, Jrmed^by the elements of Groups V^C^^ G-up I., ' II., !„., IV,. v., vr., vn., viii Oxides M.,0, MO, M.,0„ MO /'"^-'Ov M0„ M.,0-, MO, ' - ' "IMA, MO, M.;o; MO. Rfd 132 PEHlOniC LAW. The first series of oxides shows the maximum of the elements of each group. • eiice 6.— Some Numerical Comparisons. Each element of a series (horizontal row, Table II.) is nearly the arithmetic mean of the ones before and after it. In cases of small groups of three or four closely related elements, the middle one has its atomic weight often '. < ny nearly equal to the mean of the other two. By taking the successive groups of three in a series and noticing liov. nmch the middle one of each group differs from the mean of the other two, it is possible to very closely approximate the atomic weight of an element in that series, even if it is not known. The differences between the atomic weights of mem- bers in the same groups, but in adjoining periods is nearly constant. The blank in Group VII., period 6, has been filled in with lOO, by comparing the atomic weights, of periods 5 and 6, periods 4 and 6, and taking the mean of the weights of adjacent elements in 6th period. 7.— Physical Properties. The following table gives the numerical values for some of the physical constants of the elements. The numbers in the density and atomic volume columns are from Meyer's Modern Theories of Chemistry. PHYSICAL PHOPKHTTKS. 1 33 TABLE IV. Nl'MKHR'Al. VAI.IF.S OK CnNSI'A NTS OK I'lIK KI.E.MKNTS. Elkmknt. Ato.mic Wkioht. Dknsity. At. Vol,. MKI,TINtiiie,f. KliKMKNT. Atomic Wkkjut. Dkssity. At. Vol. MKI/I'INO I'dlNT. (.\li.sohilo ti'iiip.) Cu ti3.4 8.8 7.2 1330 Zri t)r..7 7.15 0.1 6'JO Gil 6!>.!) 5.06 11.7 303 An 75.(1!) 5.67 13.2 773 Se 79 4.6 17.1 490 Br 7!1.»r. 2.07 26.0 ■ice, Rb 85. /i 1.52 56.1 311 Sr 87.5 •2.5 34.9 1 Ik'Iow (,'a (^ above Ha Y Zr 89 00 4.15 21.7 very hi|jfh Nb Mo !»4 05.7 6. -27 8.6 16 11.1 very hi<;h Ru 101.65 12.-26 S.4 -2fi70 Rh lOK.O 12.1 8.6 -2-270 Pd 100.35 11.5 9.2 1775 ■ Ag 107.03 10.5 10.2 1-230 Cd iri.i 8.65 12.9 .500 In 113.7 7.42 16.3 440 Stt 110.1 7.29 16.1 503 Sb 1-20.3 6.7 17.9 710 T6 1-25 (?) 6. -25 20.2 7-28 I 1-26.85 4.94 25.6 387 Cs 13-2.0 1.88 70.6 300 Ba 137.00 3.75 36.5 48 PfrVSin.VL PROPERTIES. TABLE IV Co„ihmnf. 135 Elkmknt. Atomic VVkioi IT. I>BN8iTV. At. Vol. 1 — (At >solnh" tfiiip.) Ce Dl Yb 142.3 140.3 Nr..() 173 6.2 •i.7 6.r, 22.9 20.!) 22.3 710 1273 Ta 1 !«,., j,^ 16.0 0.6 \V ^^•* 10.13 1 1 Os Ir I't 101.3 103 104.3 22.48 22.42 21.50 8.7 8.6 9.1 2770 2220 2050 All Tl IM) Bi Th 197.22 2(»0.2 204.2 2(H). 03 208.10 232.4 10.3 IS.fiO 11.86 11.38 0.82 11.1 10.2 14.7 18.1 21.1 20.0 1310 234 563 500 540 U 240.0 -— _ 18.60 12.8 rerl hfat The numbers expressing the densities of the element, generally pass, with slight exception, from a m inrmul mddle then decrease towards the end. The atomic volume (quotient of atomic weight by density) varies te reverse way. having high values at the beginning and end of the period but a smaller value near the middle de JsUies"'"'"^ •'°'"'= ''-y '" ^ ^--' -y. as do the 136 pkkiodk; law. Other physical properties, which he outside of the scope of tliis book, such as ductility, malleabilit)', refrac- tion of h\<^ht, magnetic properties, conduction of heat and electricity, also vary in a rejjular way throughout each period. 8.~0hemical Properties. The valency of the i^roups has already been alluded to, viz., — beginning with Group I., the hydrogen valence of the groups increase until I V. (carbon) is reached, then a decrease takes place, ending with the monad halogens in VIII. In addition to the similarity in properties of the groups of seven each in the vertical columns, there is a remarkable sub-grouping into threes, the elements of each small group showing a very close connection, and generally the atomic weights of these differ by a nearly constant quantity. Examples are Li, Na, K ; Ca, Sr, Ba; P, As, Sb ; O, S, Se ; CI, Br, I. « The general formulae for the compounds in the various groups are as follows : — Group I., MR where M is any element in Group I., and R a unit equivalent of an element or radical, thus, LiCl, Na(OH), K A AgNO, NaClO,. Group II., MR,, thus, MgCL, Ca(OH),, SrO, Hg(SO,),. Group III., MR,, thus, ALO„ B(HO)„ BCl,. Group IV., MR, and MR,, thus, CO, CCl^, SrO„ Sr(OH),, PbO, PbO,. APPLICATIONS OF TlllO PKKIODU; LAW. i.i; Group v., MR,, and MR.^, thus, I>C1,, PCI, NO/OH), Hi(NO,)„ i>l I, I>0(01 1)„ 1>._.0, Group VI., MR, and MR,, thus, SO,., SO,, CrO„ SO,(OH),. Group VII., MR., MR,., MR,, MR,, CIH, C10„ Cl,0„ MnO, MnO,. Group VIII., MR.,, MR.., MR., FeO, Fe,0„ PtCl,, NiSO,, NiO,. The first three c^roups do not forn. hydrides, except rarely, as Na.JI, KJI and possibly Lill. 9.— Applications of the Periodic Law. Since each element, as placed in the Tables II and III., resembles the ones both before and after it in the hori- zontal line, and those above and below it in the vertical one, It ,s possible to tell its properties pretty accurately from those of the four surroundino- it, once its atomic weight IS known, for then its position in the table is assured. This arrangement oft!., elements also servos as a check in determining doubtiui atomic weights It IS probable that elements will yet be found to occupy some of the blank spaces in the table, and their pro- perties can be known with a considerable degree of accuracy even before the elements are discovered. Group V. may be taken to illustrate some of the relationships among the elements comprising it. 13S PFIUorifr LAW. (1) The clctiuMits pas^ from i^.iseoiis iiilto^^cn through waxy pliosphorus, then to hritllc scjiids, ciuliti^ vvitli metallic bismuth. (2) The upper members in the ^^roup are entirely acid forming. Towards the middle (V. Sb, Nb) ti -xides are botii basic and acidic. Farther down they arc basic only. (3) Some members of the ^roup (N, \\ As, Sb,) form hydrides of the composition MM.j. Hydrides of about half the elements in the j^roup are unknown. The hy- drides are deconn)osed (into the elements forming them), at hi<;h temi)eratures. (4) All forin compounds with halogens, particularly CI, having the composition MC1.( and MCL. (5) All except N burn directly to M,0.( aiiu J.fl . In the acid forming elements these are the anhydrides of the -ous and -I'c acids respectively. (6) The acids generally lose one or more molecules of water when heated. (7) All except N form sulphides, thus, M.^S^ (8) The haloid compounds are all decomposed Ly water to form the haloid acid and the oxide of the element. If, then, an element were discovered whose atomic weight were about 168, so that it would fall in Group V., 9th series, it would be reasonable to suppose that its compounds would be of the character of those given above. rni.MK ,\|, ACTION \M» KNKI{(;V. 131) of CIIAPTKR XIV^ 1.— Chemical Action and Energy. In every chemical action there is tntjie than a nierc joining togetlier, separation, or interchanj^e of atoms amon<^ the molecules of the substances affected. 'I'his may be shown as follows : — 1. Make a strong solution of caustic j^otasli, and measure 50 cc. of it into a beaker. Into another beaker pour 50 cc. of water. Dilute some hydrochloric acid about one-half Place all the vessels in a tank of water imtil they reach a common temperature, then pour 50 cc. of the acid into the beaker of water, having previously placed a thermometer in it, and watch for any change of tempera- ture. Rei)eat the experiment, but use the potash solution instead of the water. 2. Dilute some sulphuric acid, and after it has cooled to the temperature of the surrounding air place in it a thermometer, drop in some bits of zinc, and observe the temperature. 3. Prepare some sodium amalgam and drop it into a beaker containing a little water. Does the temperature change ? 4. Make a mixture of sulphuric acid and water, about one tenth acid ; jjlace in this, but not in contact with each other, a strip of zinc and another of copper ; then pass a wire from one strip to the other outside of the beaker. Does the temperature change? 5. Try if the temperature of water changes while being decomposed by electricity. 140 CMKMICAL ACTION AND KNKKfiY. These experiments, which are typical chemical actions, make it clear that accompanying the atomic changes of the reaction, there is in eacli cas6 a development of free energy appearing here, either as a rise of temperature or as electrification. Now the j^rinciple of the *' Conserva- tion of Energy " has taught us that energy may be transformed but cannot be created ; hence the chemical action means a rearrangement of the energy of the system as well as of its atoms. If there was no develop- ment of heat, electrification or any other form of energy, we should have to deal simply with transferrence of matter, but as it is, any explanation which is complete, must take account of the energy either used up or freed in the change. Those reactions in which energy (generally as heat) is set free are called exothermic ; and those in which energy in some form has to be constantly supplied that the action may continue, are known as endothermic. In the fonricr, the chemical action may take place spontaneously ; or, as is more generally the case, means must be taken to bring the molecules of the constituents into ver)' close proxinn'ty to each other before the action begins. Once the action has begun, however, the energy developed is sufficient to keep it going. In endothermic reactions, there has to be a constant suppl}' of external energy, else the chemical action at once ceases. The union of hydrogen and oxygen to form water is an exothermic combination; while the breaking up of water by making it part of an electric circuit carrying a current is an endothermic decomposition. It is customary, in chemistry, to treat of this energy Law of maximum wohk. 141 in ternis of heat units. (The tlunii.il unit of physics is H caloric, and is that (luanlily of heat Ruiuircd to raise the temperature of one oram of water tlirou-h one degree. For convenience this unit is sometimes multi- pUed by looo, and is then written Calorie, thus i Cal.= looo cal.) •2.— Law of Maximum Work. When two substances by their union set energy free, It always requires the use of a quantity of energy exactly equal to ihat freed to bring about decomposition into the original substances, provided that particular decomposi- tion is possible. This, of course, is a deduction from the principle of the Conservation of I'jiergy, and it has been experimentally demonstrated as well. A chemical action which is either spontaneous, or which goes on without a continual ap[)lication of external energy, results in the formation of substances of a stable character. This may be otherwise expressed as follows : generally exothermic compounds are stable, while endo- thermic ones are somewhat unstable. This principle ),as been expressed as a general law, known as the law of maximum work, or the law of maximum heat development, thus, when Lwo substances unite spon- taneously, they form that comi)ound by which the greatest quantity of energy is set free. 3.— Quantitative Heat Results. By using such a calorimeter as is required for deter- mining the latent heat of the liquidit)- of water or the 112 CHKMICAL ACTION AND ENEROY. il ! m specific heat of a piece of metal, the student can obtain an a])proximate result for the heat of some common chemical actions. Experiments. 1. Weii^h out 10 grams of strong sulphuric acid, and dilute it to lOO cc. with water. Then weigh out 5 grams of zinc. After the acid has cooled '^ut it in a glass vessel in a calorimeter (a glass beakv.r in a box packed with cotton wool will do very well), insert a thermometer and observe the rise of temperature. From this calcu- late the number of calories freed by 65 grams of zinc when similarly treated. 2. Dissolve 2 grams of solid caustic soda, treat it with an excess of dilute hydrochloric acid, and measure the rise of temperature. Calculate how many calories would be developed from 40 grams of the hydrate. In order that comparisons may be made between vari- ous chemical actions in regard to heat, it has been agreed to take as the unit of the substance a " gram-molecule,' that is, that number of grams expressed by the combin- ing weight (atomic weight of elements, and molecular weight of radicals) of the substance. The following are the values in heat units of a few of these reactions: 2H + 0-H20-f 68300 cal., C-|-20"COo + 97000 cal., H-f-1- HI-6000 cal., H-f-Cl-HCl-f-22000 cal., Hg-fO - HgO + 30600 cal., 20^^-302-1-72000 cal., or-l-68.3 Cal. ()r-j-97 Cal. or — 6 Cal. or 4-22 Cal. or 30.6 Cal. or 2?. Cal. QUANTITATIVE HKAT KKSULTS. I4;i I loO., - HoO + O + :J 3000 cal, or 23 Cal. also 30^v 2 0.5 — 72000 cal., or — 72 Cal. HgO + O^H.Po — 23000 cal., or — 23 Cal. These symbols mean that when two grams of hydrogen unite with sixteen of oxygen heat enough is set free to raise 68300 grams of water through one degree, and when one gram of hydrogen combines with 127 of iodine heat enough is absorbed to raise 6000 grams of watei one degree. When hydrogen dioxide breaks down into water and oxygen 23000 calories are set free, but it requires 68300 calories to decompose water into its constituents, hcnv . taking the two stages of the decomposition into account the resultant energy stored up is 45300 calories. In the combustion of methane the thermal results may be calculated as follows: — CH4 + 202 = C02 + 2H.,0. From the table above, C + Oa = CO.^-f- 97000 cal. 2H., + Oo = 2H,0-f 2(68300) cal. Then C H4 + 2O2 = CO2 + 2 H^OH- 233600 cal. Just as in mechanics a system in a stable position cannot of itself move into one of instability, but must have external energy used up in bringing about such a change ; so a chemical system (of atoms) which requires the application of external energy to cause a rearrange- ment of its parts (atoms) is in a condition of stability and will not of itself become unstable, that is, decom- posed. Substances that act chemically tend to form those combinations in which the greatest amount of energy is set free, hence, those that are most stable (maximum work). 144 QUALITATIVK ANALYSIS. CHAPTER XV. 1.— Qualitative Analysis. Throughout the body of the book directions have been given for distinguishing tlie elements from one another. It is necessary, however, to know how to systematically apj)ly tests to an unknown salt until the acid and base have been determined. 2.— Bases. There are certain substances, known as group reagents which, when properly applied, enable the operator to divide the bases into groups, then the members of any group may be further separated into smaller divisions, until finally any particular one may be traced out. 3.— First Group Reagent.— HCl. ICXPERIMKNT. Dissolve separately some lead nitrate, silver nitrate and mercurous nitrate, {a) Add to each dilute hydro- chloric acid, {b) After settling, decant most of the water and add ammonia in excess, (c) To some of the original solution (O.S.) add solution of bichromate of potassium, K.,Cr.,0-. {(i) Boil the three original precipitates. Hydrochloric acid is the first group reagent and it precipitates lead, silver and mercury (ous) salts from acpieous solutions. The acid must be dilute when used. GROUP REAGENTS. 145 :rate iro- iater [inal fum, d it rom sed. 4.— Second Group Reagent, --H^S. Experiment. (a) Dissolve separately salts of the following metals, add hydrochloric acid, then pass into the solution hydro- gen sulphide : — Copper, mercury (ic), bismuth, antimony arsenic, tin (ous and ic), gold. (/;) Try if the precipitate (sulphide) is soluble in ammonia or caustic potash. The second group reagent is hydric sulphide, which from acid solutions precipitates the salts of I. Copper, mercury, bismuth, cadmium, lead. II. Antimony, arsenic, tin, gold, platinum. The sulphides of the metals in division I. are insoluble in ammonia, those in II. are soluble. By using a solution of the salt acidified with HCl, the metals of the first group will be eliminated, except a small trace of lead, which may appear in this group because lead chloride is very slightly soluble in water. 5.— Third Group Reagent.— (NH^X.S. Experiment. Salts of iron, aluminium, zinc, nickel, manganese and cobalt may be used, (a) Dissolve the salt and add ammonia, ammonium chloride and ammonium sulphide. Note the colour of the precipitates. (d) To some of the original solution, add ammonia or caustic potash. The third group reagent is ammonium sulphide in presence of ammonium chloride in an alkaline solution. The metals of the group are I. Aluminium, chromium, iron, which form hy- drates with alkaline hydroxides. II. Zinc, manganese, cobalt, nickel. 10 146 GROUP REAGENTS. 6.— Fourth Group Reagent.— (NHJaOOs. Experiment. Soluble salts of barium, calcium and strontium, may have added to them some ammonium carbonate. Test the precipitate for flame colouration. The fourth group rca^rent is ammonium carbonate. Barium, calcium, strontium, are precipitated by it. .i«. 7.— The Fifth Group. No Reagent. The members of this group are not precipitated by any reagent. They are potassium, sodium, ammo- nium and magnesium. To sum up, there is the, First, the Hydrochloric Acid or the Silver Group, consisting of those metals whose Chlorides are insoluble. The Second, the Hydrogen Sulphide or the Copper Group, made up of those metals whose sulphides are insoluble in water. The group consists of two sections, known as the Copper and Arsenic divisions. The sulphides of the former are insoluble in ammonium hydrate, those of the latter are soluble. The Third, or Ammonium Sulphide or Iron Group, consists of those metals whose sulphides are insoluble in presence of ammonium hydrate and ammonium chloride. This group is also in two divisions, the metals of the first or iron section, have their hydrates precipitated by ammonium hydrate. The metals ol QUALITATIVE ANALYSIS. ^47 The Fourth, or Ammonium Carbonate or Barium Group consists of metals whose carbonates are precipi- tated when formed in aqueous solution. The Fifth or Potassium Group, i, made up of metals that do not form insoluble salts with any one reagent. 8.-Some Directions for Working in Making an Analysis. I. Keep all vessels, brushes, stirring rods, etc., perfectly clean ; and use pure water only. ^^2.^ Examine the substance to be tested before dissolv- 3. Keep accurate notes madeai the time of your obser- vations and inferences. 4. Solutions of the substances are always used for analysis, except where other^M'se directed. If the salt is soluble m water, use aqueous solution for testing If no so ub e in water try first, strong hydrochloric acid; if not soluble m that use nitric acid ; if not in nitric acid use "= .»». be ir)0 (^UALITATIVK ANALYSIS. Precipitate, black. Precipitate, yellow. Precipitate partly soluble in HNO,, dilute. .....Pb. Confirm as in Table II. Precipitate soluble in boiling HNO3. Add NaOH to O.S. white precipitate. Add ex- cess of water to O.S. white precipitate Bi. Confirm by dissolving last precipitate in HCl. Soluble in hot dilute HgSO^ Cd. Confirm by adding NH4OH and NaOH slowly to separate parts of O.S. ; white precipitate, soluble in excess. B— Arsenic Section of Second Group. %: Precipitate with H2S, black or brown. Precipitate, yellow. Precipitate, orange. Sn(ous). Confirm by adding to O.S. slowly KOH, white precipitate soluble in excess of alkali. Add to O.S. oxalic acid, white precipitate. Soluble in N H4OH. Turns green with am. solution of CUSO4 As. Confirm with am. solution of AgNO., (yel- low). Marsh's test with O.S. (page 89, ex. 9). White precipitate with NH4OH. Yellow sulphide insoluble in (NM4)2C03. .Sn(ic). Reduce to ous salt and confirm as above. Insoluble in NH4OH. Soluble in hot con- centrated HCl Sb. Confirm — white precipitate with NH4OH in O.S. Test for stibine (pages 93 and 90, ex. 6). VIMLITATIVK ANALYSIS. TABLE IV.-Metals of the Third Group. Treat some of O.S. with KOH. Precipitate . . No precipitate..". ,; ^;"" ''"^•^tion. ^^- -^inc iJcciion. 151 Precipitate with (NHJaS, black. A-Iron Section of Third Group. Precipitate, white. Precipitate, green- ish. Precipitate dissolved by HCI. O..S. with ^aOH gives whitish precipitate, turning J,-oen, then brown p^^^^^ Conhrm by adding to O.S. K,FeCy„, hght '^ue precipitate. K,Fe,Cy, ,, dark bh.e. ^a^LOa, white, then green, then brown. Precipitate dissolved by HCI. O.S with NaOH gives reddish precipitate. ...Fe(ic ) Confinn with K,FeCy,„ dark blue; ^si'eaCy.o, brown. KCyS, bright red. Precipitate soluble in HCI. With NaOH O.S. gives white precipitate sohible in ex- ^^^^ Al Confirm by adding to last solution Hc'l until acid, then NH.OH until alkaline (gelatinous precipitate). Heat the salt on charcoal, then moisten with CoCl^ and heat again, fine blue. Dissolves in HCI, forming green solution /^_ Confirm thus,-O.S. forms greenish" pre- cipitate with NaOH, soluble in excess Add a solution of lead salt to this (yellow precipitate). 152 yUALlTATlVli ANALYSIS. B— Zinc Section of Third Group. Precipitate, with Precipitate with (NH4)aS, white. Precipitate with (NHJ^S, //«/'. Precipitate sohible in HCl. O.S. with NaOH gives hUie precipitate, insohihlc in excess, sohihle in NII4OH Co. Contirni. Morax bead turns bhie when Co is dissolved in it. O.S. gives green pre- cipitate with K4 FeCy,,, rechUsh precipi- tate with K(jFc.jCy,o, both insoUible in HCl. Precipitate soluble in HCl, greenish preci- pitate with NaOH, insoluble in excess. Sohible in NH.OH Ni. Confirm. K4FeCy„ in O.S. forms greenish precipitate. Na^COg gives greenish precipitate. With borax it forms violet bead in O.fl. when hot. Precipitate soluble in HCl, insoluble in HC._.H..jO. Zn. Confirm. O.S. yields white precipitate with alkaline hydrate, soluble in excess. Na.jCOg in O.S. forms white precipitate soluble in NH4OH. With NaOH, O.S, gives white precipitate, becoming brown in air, soluble in * ^' , CI Mn. Confirm thus, — heat sa' K.jCOg n d KNO3. It turns gr> Borax K.id with salt, amethyst wiu n h'U, reddish when cold after heating in O.fl. QUAl,ITATIVi: ANALYSIS. If).} TABLE V. -Metals of the Fourth Group. white Precipitate. Add CaSO, to O.S., white precipiinte ,ft '''!"'•: Ba. Conhi 111 by adding acetic acid and K.Xr.^O^ to O.S., yellow by tlanie test (page^js, ex. i). Precipitate with CaSO^ forms slowly. . .Sr. Confirm by flame test (page 34, ex. i). With CaS04, no precipitate Ca. Conlirm thiis,~(N 1I,,)._.C.,0, gives while precipitate with O.S. Ca gives orange- red colour to flame (page 31, ex. i). TABLE VI. -Metals of the Fifth Group. Salt on platinum loop colours the flame crimson t : Confirm by treating concentrated O.S. with NaoCOg, white precipitate. Salt on platinum loop colours flame bright y^""^^' Na. Salt on platinum h)op colours the flame violet XT Confirm by testing with tartaric acid (white precipitate). O.S. gives white precipitate with NaoCO,, • Mg. Confirm by treating O.S. with Ca(OH).,, white precipitate. Also with (N 114)2 HI'o'^ or NaoHPO^ in presence of NH4OH Twhite precipitat.;. Solution heated with KOH gives off am - '"»"''* .NH4. Confirm by Xessler's test. 154 QUALITATIVK ANAF.VSF.S. To test for the more commonly occurring acid radicals in salts containing one acid : — First Group Reagent, Barium Nitrate. This reagent forms a white precipitate with a sulphate, carbonate, oxalate, silicate, phosphate, borate and sul- phite. Second Group Reagent, Silver Nitrate. With silver nitrate a precipitate is formed by a chloride, bromide, iodide, cyanide, sulphide, nitrite, ar;d thiosul- phate (hyposulphite). Third Group No general reagent. This group may contain a chlorate, nitrate, tartrate, or acetate. To test for the acid, make a solution of the salt. Add solution of BaCNO^X-* If no precipitate appears at once, let the mixture stand for a few minutes. White Precipitate Table I. No Precipitate. Add AgNOs to O.S. Precipitate Table 1 1. No Precipitate Table III. TABLE I. Shake up a little of the precipitate with water and add HCl. Observe if a gas is given off, if so test it with Ca(H0)2. A — White precipitate not dissolved A Sulphate. Confirm by heating the salt with NajCOg on charcoal in reducing flame, then apply- ing test with silver. * After dissolving the nitrate of barium, add nitric acid until no more precipitate is formed, then filter and use the clear solution as the reagent. QUALITATIVE ANALYSIS. 155 B— Precipitate dissolved. Add H...SO4, and heat, if necessary. (i) CO2 given off. (It was also given off when H CI was added) a Carbonate. Confirm by adding HCI to the salt, then pr.ssing gas into Ca(OH)„. (2) SO^givenoff ^ Sulphite. Confirm by adding AgNO., to O.S. White precipitate, AgoSOg, turns dark on heating. (3) CO and CO^ given off An Oxalate. Confirm by adding CaCL to O.S. White precipitate, soluble in HCI, in- soluble in HC2H3O2. If this precipitate be heated just to redness it effervesces with HCI. O.S. with AgNo..j gives white precipitate, soluble in HNO3. O.S. with stannous chloride gives white precipitate. (4) A deposit of Silica separates A Silicate. Confirm by heating a little of the powder in NagCOg on platinum loop, when a colourless glassy bead is formed with evolution of COg. (5) No apparent change. Test O.S. with Ag-NO;,. (6) A yellow precipitate a Phosphate. Confirm with amnionic molybdate. (7) A brownrsh precipitate a Borate. Confirm by flame test (page 54). 156 QUALITATIVIO ANALYSIS. TABLE II. Group reagent — AgNO,,. A solution of the salt, with silver nitrate added, gives A— A white precipitate. Add H0SO4 to O.S. and heat. (f) Hydrochloric acid gas given off A Chloride. Confirm by testing for chlorine witli H2SO4 and MnOa. (2) SO a and S separated A Thiosulphate. Confirm by adding any othei" strong acid to O.S. (3) Nitrous fumes (reddish) given off A Nitrite. Confirm by adding some O.S. to a little solution of KMnO^ acidified with HgSO^, also by adding to O.S. some starch paste and KI. B — A yellowish white precipitate. Add H 2 SO 4 to O.S^and heat. ( i) Red fumes of bromine given off. A Bromide. Confirm by adding CI solution to O.S. with starch. {2) Violet vapours of iodine An Iodide. Confirm Ijy adding solution of a lead salt or of mercuric chloride to O.S. ; also by adding CI to O.S. with starch. C — A black precipitate. Add H2SO4. (1) \h^ given off A Sulphide. Confirm by silver test. SELKfri-KI^ QIIKSTFONS. 157 late. TABLE III. No precipitate with Ba(NO,)o or AgNOa. Add H,SO, to a little of the solid, and heat, if necessary. (i) Brownish acid fumes A Nitrate. Confirm by treating sokition with HgSO^ and ^^6804. (2) Colourless acid vapour, with smell of vi"egar An Acetate. Confirm by neutralizing the solution of the salt, then adding F>2 <-''.;, this gives bright red colour without precipitate. (3) A yellow gas that explodes in air A Chlorate. Confirm by heating for free oxygen and a chloride. (4) Mixture turns black or brown A Tartrate. Confirm by adding slowly K2CO3 with constant stirring, while precipitate. Re- peat, using ammonia. In both cases the solution must be acid. SELECTED CUES TIONS. The followiiinr questions arc selected from examination papers on the work covered by this book :— I. One gram of a certain metal when dissolved in dilute sul- phuric acid, liberates 200 c, bining weight of the metal. c. of hydrogen gas. Find the com- 2. Hy what experiments would you dist oletiant gas inguish marsh gas and 5 c.c. of a mixture of marsh wi th gas and olefiant gas are exploded 14 c.c. ofoxy-en: 9 c.c. of gas remain, of uhi-j, 7 c.c. are 158 SELECTKn QURSTIONS. I aljsorhecl by c.uistir potash. I'iiul the volume of each of the .Ljases in the original mixture. 3. (a) How would you prove that the gas obtained by pouring sulphuric acid upon ferrous sulphide contains both S and H .'' (/f) A solid substance contains lioth a carbonate and an easily dissolved sulphide. How would you prove the presence of these two bodies .'' (c) A piece of sodium was completely converted into chloride by uniting with 200 c.c. of CI at the standard temperature and pressure. What was the weight of the sodium? 4. State Avogadro's hypothesis and give the evidence in support of it. Deduce the general staten\ent that if A/ be the density (Hydrogen= i) of any substance in the gaseous state, Jf grams of that substance in the gaseous state will occupy approximately 1 1. 16 litres at o°C. and 760"""- Bar. 5. 0.6 grams of a certain metal when dissolved in dilute sul- phuric acid liberates 558 c.c. of hydrogen at oX. and 76o""»- Bar. A determination of its specific heat gave .25. Find the atomic weight of the metal. 6. (a) Describe experiments illustrating the difference between (i) nascent hydrogen and hydrogen, (ii) nascent oxygen and oxygen. (/j) Write equations illustrating the action of nascent hydrogen upon (i) nitric acid, (ii) solution of arsenious oxide, (iii) nitric oxide. 7. Describe experiments illustrating how you would detect (d) Potassium nitrite in presence of potassium nitrate ; (/>) Potassium sulphite in presence of potassium sulphate ; (c) I'otassium chloride in presence of potassium bromide ; ((/) Arsenic hydride in presence of antimony hydride. 8. (a) How much oxalic acid (C.H.jO^) must be heated with sulphuric acid to prepare 145 litres of carbon monoxide at I7°C. and 8oo"""- Piar. (//>) Explain what occurs when a mixture of carbon monoxide and chlorine is exposed to sunlight and the product shaken up with water. SELECTKI) QUESTIONS. 159 {c) If loo c.c. of a mixture of carbon monoxide and hydrogen gas were given yon for analysis, describe how you would proceed to determine the volume of each gas in the mixture. 9. Describe, giving equations, what will occur in each of the following experiments : — (a) A piece of yellow pliosphorus is suspended in Chlorine in a bottle. {b) The product of reaction {a) is shaken up with water. {c) An electric spark is passed through the air in a closed flask containing also a small quantity of a solution of potassium iodide. id) Hydrochloric acid is added to a few crystals of potassium chlorate in a bottle and hydrogen sulphide is then passed into the bottle. {e) Lead nitrate is heated and the gas given off is gradually cooled down to — 2o°C. 10. Describe the relations which sodium and its compounds bear to potassium and its compounds. 11. (a) In the periodic arrangement of the elements, manganese IS placed in the same group as chlorine. Write the formuU« of the compounds of these elements that illus- trate this relation. {/>) Write equations showing the action of (i) concentrated sulphuric acid on manganese dioxide, and (ii) di'ute sulphuric acid and potassium permanganate on ferrous sulphate, 12. {a) The specific gravity of hydrogen is 0.0692, that of ammonia is 0.595 (:u'- O- What is the molecular weight of ammonia.? Explain how you arrive at your conclusion. (/>) The vapour density of sulphur is said to be abnormal at 5oo°C. and normal at loooT. Explain the meaning of these statements. {c) Describe experiments to show that ammonium chloride is decomposed by heat into ammonia and hydrogen chloride gas. 160 SKLE(!TRn QUKSTfONS. 13, A gas gives, on analysis, Cail)()n J lydrogen 85.78 14.32 100. 300 c.c. of the gas at lo^C. and 750 mm. Bar. weigh 0.35 grams. Write the formula of the gas. 14. {ii) How can barium oxide be used as a means of preparing oxygen from the air ? {li) Compare the chemical properties of oxygen and ozone. (r) If the molecule of oxygen consists of two atoms, then the molecule of ozone consists of more than two atoms. State facts in s"pport of this statement. 15. Chlorine, bromine and iodine are said to l^elong to the same natural family of elements. Explain and illustrate this statement. 16. {(i) How could you prepare the trihydride of arsenic (arsine).? Compare its properties with ammonia and phosphine. (/>>) Write equations showing the reactions which occur when arsinc, phosphine, and ammonia respectively burn in air. 17. Write equations illustrating the action of hydrogen sulphide upon : {a) Chlorine gas. (/-') Ammonium hydrate. {c) Nitric acid. {d) Antimonous chloride. (<') Solution of ferric chloride. 18. A solution is known to contain silver nitrate, ferrous nitrate, ferric nitrate or arsenious oxide. How would you determine which it contains.'' 19. {a) Give an account of the chemistry of calcium. Illustrate its relations to barium and strontium. {b) What is the cause of the so-called temporary hardness of water .-^ Explain the chemical reactions which occur in the different methods adopted for its removal. HKLECTKD QUEvSTIONS. 161 20. (a) " The temperature of a burning match is far above that of a red hot iron poker." Describe an experiment which proves the correctness of the statement. (d) What voUniie, calculated at standard temperature and pressure, will be occupied by the products of the combus- tion in oxygen of one gram of each of the following substances: ammonia, phosphoretted hydrogen, carbon disulphide, and arsenic ? 21. Dalton's analysis of two compounds yielded the following results: Carbon - 42.86. Carbon - 27.27. Oxygen - 57.14. Oxygen - 72.73. Show the relation of these data to (a) the relative (or propor- tional) weights of these elements, (/;) the law of multiple proportion, and (c) the atomic theory. 22. Find the atomic weights of the elements from the following data : ^ (a) A gram of a certain metal dissolved in dilute sulphuric acid yielded 344 c.c. of hydrogen at standard temperature and pressure. Its specific heat was found to be -0956. (/') 35-5 parts of chlorine unite v/ith 100 parts of mercury to form mercurous chloride, and with 200 to form mercuric chloride. The specific heat of mercury is 0*32. 23. (a) Make a solution of barium nitrate ; add to it, drop by drop, a solution of sodic carbonate until no further preci- pitate forms ; filter off the liquid portion ; collect, dry and heat precipitate just to redness ; add hydrochloric acid until all action ceases ; dip a platinum wire into the solu- tion which forms and place in the non-luminous flame of a Bunsen burner. Explain the whole series of changes and phenomena. Give equations. (d) What other metals exhibit a similar series of changes under somewhat similar conditions ? 24. («) You are furnished with sulphate of magnesia and all ^^ necessary reagents and apparatus. Describe how you 162 SELECTED QUESTIONS. I would prepare the oxide, the chloride, and the carbonate of magnesia. (/;) Of what other elements could you form similar compounds in a somewhat similar manner ? 25. Sketch the chemistry of sodium. 26. (iive a practical definition of a dibasic acid. Name one, write its formula, and show how your definition applies to the acid. 27. The vapour density of hydriodic acid is 4.43 (air=^ 1) ; that of phosphorus tri-iodidc is 14.27 ; the percentage weight of iodine in these compounds is respectively 99.2 and 92.5. Calculate the atomic weight of iodine. 28. Compare the hydrides of the members of the nitrogen group. 29. {a) Two-tenths of a gram of a compound having the com- position CigHooOj I is burnt in air. Explain the chemi- cal changes that take place, using equations. Calculate the products of the volume of the products of combustion at ioo°C. and 740 mm. {f?) .18 gram of a compound containing carbon, hydrogen and oxygen, yields, on analysis, .072 gram of carbon, .012 gram of hydrogen, .096 gram of oxygen. Calculate the simplest formula of the substance. 30. Define the terms "oxidizing agent "and "reducing agent," and illustrate your definitions by reference to the experiments : — {a) .Sulphuretted hydrogen gas is passed into a solution of ferric chloride. ib) Carbon is heated to a high temperature with ferric oxide. ■'• {c) .Sulphur dioxide is passed into a solution of permanganate of potash. C.ive equations. 31. {a) Dalton's gravimetric analysis of two compounds yielded the following results : — Nitrogen 63.64 Nitrogen 46.67. Oxygen 36.36 Oxygen 53.33. Show the relation of these data to Dalton's formulae for these substances (NO and NO.^). SRLKCTKI) QUKSTIONS. ins (/') (Jay LussHc's volimietric analysis of the inixtme resulting from the decomposition of these same compounds j^ave the following results : — Nitrogen 66? vols. Nitrogen 50 vols. Oxygen ^3!^ vols. Oxygen 50 vols. Show the relation of these data to the present formula? for these substances, and to Avogadro's law. 32. (a) Describe what takes i)lace when : (i) Iron is immersed in a solution of sulphate of copper. (ii) Copper, in a solution of bichloride of mercury, (iii) Zinc, in a solution of acetate of lead, (iv) Magnesium, in a solution of nitrate of silver. (/^) Explain how quantitative results in these experiments can be used as an aid in determining atomic weights. 33. Describe simple laboratory methods of preparing small quantities of (a) metallic arsenic from the trioxide, {/>) trichloride of antimony, (c) ferrous sulphate. 34. Sketch the chemistry of lead. 35. What is meant by the " Periodic Law".? Illustrate its signi- ficance by reference to the members of group iv. (Carbon -12, silicon = 28, tin = 1 1 8, lead = 207. ) 36. The average composition of coal gas is : hydrogen, 45% ; methane, 35% ; carbon monoxide, 7"^ ; oleHant gas, 4/ ; Inityienc (C4H8), 2.4%; sulphuretted hydrogen 0.3%; nitrogen 2.5%; carbon dioxide, 3.8%. What volume will the products of the com- bustion of 100 litres of such a gas occupy at 20 C. and 750 mm. pressure ? 37. Explain, using equations, what changes take place in the following experiments :— {a) Dry sulphuretted hydrogen is passed over iron filings in a glass tube. (^) Sulphur dioxide is passed into a vessel containing nitro- gen peroxide. {c) Carbon dioxide is passed over ignited sodium. TT 104 SF.I.KCI'KI) (/I'KSTloNS. 3. Analysis, tables for, 148. Antimony, 8!). Antimony, compounds of, i»2. Antimony, tests ft>r, i)3. Antimonyl, 1)2. Applit ations of Periodic Law, 137. Arstnic, 84. " notes on, 8(5. " and oxygen, 8(5. '* and sulphur, 88. Arsine, 87. Atomic heat, 13. Atomic weight, calculation of, 123. B. Barium, 30, 35. " couii-MJunds of, 37. " notes on, 37. " test« for, 38. Basic carbonate of magnesium, Jl. Basic salt, !H). Bismuth, 04. liismutli, notes on, !t4. liliick ash, 22. l^orax, 49. ]3oron, 47. *' notes on, 49. " tests for, 54. Bromine, 1. Cahiuni group, 30. Calcium, 31. compoimds, 34. tests foi', 38. Calculation of atomic weights, 123. Carbon, 5(i, 57. Carbonic acid, 58. Carbon, hydrides of, 59. oxides of, 57 and nitrogen, 02. " sulphide of, (51. T'atalysis, 1(). Chi'mical action and energy, 139. '• two laws of, 14. ' and mass, 14. ' and temperature, 15. ( 'liloric acid, 8. Chlorides of sulphur, 108. Chlorine group, 1. (.'oinpounil-i of the halogens, 7. antimony, !il. iron, 113. lead, 72. manganese, 119. n (I (1 U tl 1()9J 170 INDKX. Com|KMni(i.s of oxygen, 101. phoFpliorus, SH. sulphur, 104. D. Dithionic Jioirl, 107. Dulong and Petit'.s Law, 13. E. Kndothcnnic, 104, Energy and chemical action, 130. Kpsom salts, 40. Exothermic, 140. F. Ferric compounds. 111. F«;rric acid, 115. Ferrous compounds, lOO. Fluorine, 5. a- ^ ( Jroup reagents, 144. H. Halogens, 1. Halogens, notes on, 6. Hydride of arsenic, 87. Hydride of carlion, 5IJ. Jlydriodic acid, 5. Hydrobromic acid, 2. Hydrogen, 17. ■ J.. Iodine, 4. Iron, 10!). Iron compound"', 113. " group, 100. " notes oa, 112. Isomorphism, 12. L. Law of Dulong and I'etit, 13. maxinnnn work, 141. nndtiple i)roportions, 11. Leblanc's process, 22. Lead, 5f), (iO. Lead, comjjounds of, 72. " notes on, 71. " tests of, 70. M. Magnesium, 40. Magnesium, group, 40. Magnesium, notes on, 42. Manganese, 17. Manganese, com|xmnds of, liy. Ma*s and chemical action, 14. Maximum heat development, 141. Maximum work, law of, 141. Mordant, 54. Multiple proportions, law of, 11. N. Nitrogen group, 77, 98. Nitrogen and carbon, 02. Notes on aluminium, .51. antimony, 01. bismuth, 04. halogens, 0. iron, 112. lead, 71. magnesium, 43. mangane-ie, 119. nitrogen group, 98. oxygen, 100. phosphorus, 80. silicon, 05. tin, 67. zinc, 45. (t LNDKX. 171 O. Oxides of airsenic, 57. carbon, 57. pliospliorus, 82. sulphur, 104. Ozouo, 102. Oxyjren, KKI. Oxygen, coniiMjunds of, 101. P. Periodic law, 125. Periodic law, apjilications of, 137. Permanganates, 119. Phosphine, S.S. Phosphorus, 7S. Phosphorus, allotropic forms of, 81. notes on, 80. oxides and acids of, 81. vapor density of, 84. Potassium, 19, 20. Potassium, notes on, 25. compounds of, 25. tests for, 28. * it Qualitative Analysis, 144. Reagents, group, 144. Salt cake, 22. Silicon, .50, 02. Silicon, notes on, 0.5. Sodium, 19, 20. Sodimn, tests for, 25, " com|Kiunds, 22. " notes on, 21. Stannic acid, 08. Stannous coni|)ouiids, 07. Strontium, ',M>, 34. Strontium, notes on, 35. Strontium, t«^sts for, 38. Sulphacids of arsenic, 89. Sulphide of arsenic, 88. Sidphide of carbon, 01. Sulphur, 102. Sulphur, comjK)unds of, 104. Sulphur, density of, 103. T. Tables for aiialysis, 148. Temperatiire and chemical action, 15. Tests for anthnrmy, 93. arsenic, 87. boron and aluminium, 54. iron, 117. " had, 70. " mangane-ie, 122. magnesium group, 40. Thio-acids, 107. Thio sulphuric acid, 107. Tin, 57, 00. Tin, nf)tes on, 07. Trithionic acid, 107. Vapour density of phosphorus, 84. Vapour density of sulphur, 103. Zinc, 43. Zinc, notes on, 45.