GIFT F 
 Harry East Miller 
 
A SHORT COURSE 
 
 IN 
 
 QUALITATIVE ANALYSIS 
 
 BY 
 
 J. M. CRAFTS, 
 
 FORMERLY PROFESSOR OF GENERAL CHEMISTRY IN THE CORNELL UNIVERSITY. 
 
 THIRD EDITION. 
 
 REVISED BY 
 
 CHARLES A. SCHAEFFER, A.M., PH.D., 
 
 PROFESSOR OF GENERAL AND ANALYTICAL CHEMISTRY IN 
 THE CORNELL UNIVERSITY. 
 
 NEW YORK: 
 
 JOHN WILEY & SONS, 
 
 15 ASTOR PLACE. 
 
 1880. 
 
COPYRIGHT, 
 
 1879, 
 BY JOHN WILEY & SONS. 
 
 GIFT OF 
 
 i-ffo 
 
This Book is inscribed by the Author 
 
 M. ADOLPH WURTZ, 
 
 AS A TOKEN OF AFFECTIONATE REGARD, TO A FRIEND AND TEACHER, 
 
 AND AS A TRIBUTE OF RESPECT AND ADMIRATION, TO THE 
 
 MASTER OF A SCHOOL IN MODERN CHEMISTRY. 
 
 M81896 
 
PREFACE TO THE THIRD EDITION. 
 
 IN preparing a new edition of this work, the editor has made 
 no essential alteration of the plan adopted by Professor Crafts 
 in the previous editions. He has endeavored, however, as far 
 as possible, to embody the practical results of the great ad- 
 vance made in both theoretical and practical chemistry within 
 the past ten years. The few important changes thereby ren- 
 dered necessary may be briefly specified. 
 
 Chapter II., on chemical nomenclature, has been entirely 
 rewritten. The usefulness of the work has been somewhat ex- 
 tended by the introduction of several elements not included 
 in the previous editions, viz. : strontium, cadmium, iodine, and 
 bromine. 
 
 Further changes have been made in the methods of separa- 
 tion and detection of several of the elements ; more reliable 
 processes being substituted for those which in the hands of 
 students were found not to be entirely satisfactory. 
 
 The editor desires especially to acknowledge his obligation 
 to his friend and colleague, Dr. G. C. Caldwell, for numerous 
 valuable suggestions of which he has availed himself. 
 
 ITHACA, December 17, 1879. 
 
PREFACE TO THE FIRST EDITION. 
 
 THIS little work was written for the use of a class of stu- 
 dents in the Cornell University, who take a year's course of 
 chemistry, including four hours a week of laboratory practice ; 
 reference was also had to the requirements of the scientific 
 students in Union College, whose course is nearly equivalent 
 to that mentioned. The author is indebted to the kindness 
 of Professor Perkins for many valuable suggestions, and for 
 the compilation of Tables IV. and V. at the end of the book. 
 
 A considerable portion of the introductory part of this book 
 is devoted to an explanation of the theory of chemical reac- 
 tions and nomenclature. Many of the standard works on an- 
 alytical chemistry still use the old notation, and the formulas 
 to be found in them do not correspond to those used in the 
 best text-books on general chemistry ; and none of the ele- 
 mentary works for the laboratory-student supply the want, 
 often felt by him, of a system of rules, at hand for use at the 
 moment when he most requires them, namely, when he is 
 writing the formula of a reaction at his desk in the laboratory 
 with his tests before him. The author has had these points 
 in view in writing the two introductory chapters. 
 
 It might be objected that the theory of chemical notation 
 should be found in text-books on general chemistry ; but even 
 when the student has mastered the rudiments of the science, 
 as given in any of the best modern works, he will find the 
 
PREFACE. v ii 
 
 arrangement of many of them inconvenient for reference, 
 although it is excellent for instruction, and, moreover, it is by 
 no means necessary that the study of those works should pre- 
 cede that of analytical chemistry. 
 
 It is quite feasible to commence a course of experiments in 
 the laboratory concurrently with the study of any of the ele- 
 mentary text-books on general chemistry, or with the attend- 
 ance upon lectures, illustrated with the usual experiments. 
 Chemistry certainly becomes a more attractive study when 
 the practical and the theoretical present themselves side by 
 side, so that while the theory explains the experiment, the ex- 
 periment awakens an interest in the theory ; and no course of 
 study is more apt to interest the beginner in chemistry than 
 that of the admirably simple and delicate tests of qualitative 
 analysis ; tests which illustrate the general laws of the science, 
 while they have a very direct bearing upon some of the prob- 
 lems of every-day life. 
 
 Analytical chemistry, besides its immediate value as an im- 
 portant branch of knowledge, cannot be too highly prized as 
 affording a convenient introduction to the methods of investi- 
 gation used in an experimental science, and as offering a 
 means of education of many faculties, which are not easily 
 developed by school or university training. The importance 
 of laboratory experiments is awakening every day increased 
 attention, and the time is fast passing by when chemistry is 
 taught to persons, who suppose that they have a vocation for a 
 scientific profession, only by lectures and recitations. 
 
 The system of analysis, given in Part III., is founded upon 
 that of Fresenius, and includes a minute description of all the 
 steps to be taken in performing tests and directions for pass- 
 
viii PREFACE. 
 
 ing from one test to another ; and since these details, which 
 are required by the beginner, not only become unnecessary, 
 after a certain familiarity with analytical work has been ac- 
 quired, but also render the scheme less convenient for refer- 
 ence, tables similar to those of Will have been added, which 
 indicate the important tests and leave the experience of the 
 student to suggest the proper mode of applying them. 
 
 The Tables IV. and V. are intended to record a number of 
 facts in analytical chemistry in a compact form, and to give 
 an exact conception of what is meant by the insolubility of a 
 precipitate, and a means of judging of the advantages of dif- 
 ferent methods of precipitation. 
 
 ITHACA, August 19, 1869. 
 
 PREFACE TO THE SECOND EDITION. 
 
 A FEW changes have been made in the second edition, 
 which have been suggested by the use of the previous one in 
 the laboratory and some practical directions have been added, 
 which are intended to mark out more closely the course of 
 study which it seems most advisable to follow. 
 
 One of the principal applications of analytical chemistry is 
 to the investigation of minerals, and it may not be out of place 
 here to recommend to the consideration of instructors and 
 students the combination of mineralogical work with labora- 
 tory practice. A convenient order of study is, first to acquire 
 
PREFACE. i x 
 
 familiarity with the methods of detection and separation of 
 bodies in solution, which are given in Part III., and then, be- 
 fore taking up the analysis of solid bodies, to study their crys- 
 tallographical and general mineralogical characters from any 
 elementary work treating of the subject. Turning again to 
 the general analysis of solid bodies (minerals and technical 
 products of all kinds), the course should be made as varied 
 and extended as possible ; and it will be found that the drill 
 in blowpipe practice and the study of the physical characters 
 of minerals make an excellent preparation for the preliminary 
 testing of solids, Part III., and often afford hints which are of 
 great value in the subsequent performance of the complete 
 analysis in the wet way. 
 
 It is quite possible for a class to complete a course like the 
 one proposed by working four hours a week for a year in the 
 laboratory. 
 
 ITHACA, 1870^ 
 
A SHORT COURSE 
 
 IN 
 
 ANALYTICAL CHEMISTRY. 
 
 
 PART L INTRODUCTION. 
 
 CHAPTER I 
 
 THE analytical chemist is not generally required to investi- 
 gate a large range of substances, and the theory of the com- 
 position of those bodies which he usually deals with can be 
 briefly explained. 
 
 Elements. The ultimate result of the analysis or de- 
 composition of all matter is the discovery of a number of sub- 
 stances which cannot be decomposed, and are therefore called 
 simple or elementary substances. The table on the next page 
 contains the names of thirty-six elements. They are chosen 
 from among the seventy or more elements which have been 
 discovered, because they are more frequently met with than 
 the others, and because they are the only ones treated of in 
 this work.* 
 
 * Within the past four years ten new elements have been reported by 
 their discoverers. How many of these will be substantiated it is at 
 present impossible to say. On the other hand, very recent investiga- 
 tions seem to show that at least several of our well-known elements are 
 really compound bodies. The exact number of elements is, therefore, 
 especially at the present crisis, very difficult to fix. 
 I 
 

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SYMBOLS. 3 
 
 The task of the analytical chemist is to recognize these ele- 
 ments when they occur alone, or to separate them from the 
 compounds of which they form a part, or most frequently to 
 resolve a compound substance into simpler compounds, and 
 to isolate the latter in such form that they can be easily re- 
 cognized. When the different elements are known which 
 constitute such simpler compounds, the composition of the 
 complex body from which they were obtained can be deter- 
 mined. 
 
 All compound substances are formed by the union of chem- 
 ical elements, and a previous study of the manner in which 
 the elements combine with each other is essential to the suc- 
 cessful pursuance of analytical investigations. 
 
 The language in which chemists express their ideas regard- 
 ing the chemical constitution of bodies comprises certain sym- 
 bols or abbreviations, to which a conventional meaning is 
 attached, and formulas, which are made by grouping symbols 
 together. 
 
 Symbols* The symbols which stand in the table after 
 the names of the elements are abbreviations, which are used 
 instead of the names in writing chemical formulas. 
 
 Chemical Formulas describe by means of symbols 
 the chemical constitution of bodies. 
 
 Example : HC1 is the formula of hydrochloric acid, and 
 signifies that it is composed of hydrogen and chlorine. 
 
 The Combining Weights are the numbers standing 
 in the table after the symbols. They are peculiar to each 
 element, and denote the proportion by weight in which it 
 unites with other elements. In formulas the symbols stand 
 for these weights, as well as for the names of the elements ; 
 thus, in the formula of hydrochloric acid, HC1 signifies that i 
 part by weight of hydrogen is combined with 35.5 parts by 
 weight of chlorine. 
 
 ELEMENTS COMBINE WITH EACH OTHER IN NO OTHER 
 
 PROPORTIONS BY WEIGHT THAN THOSE EXPRESSED BY THE 
 
4 PART /. 
 
 COMBINING NUMBERS (OR BY VERY SIMPLE MULTIPLES OF 
 THEM). 
 
 This statement comprises the law of definite proportions 
 and also that of multiple proportions. 
 
 USUALLY THE MULTIPLES OF THE COMBINING NUMBERS, 2, 
 3, 4, 5, 6, 7, EXPRESS THE PROPORTIONS IN WHICH THE ELE- 
 MENTS COMBINE WITH EACH OTHER. 
 
 Example : Nitrogen combines with oxygen only in the pro- 
 portions expressed by the formulas : N 2 O, NO, N 2 O 3 , NO 2 , 
 and N 2 O 6 , or, referring to the table for the combining num- 
 bers for which the symbols stand, the following proportions 
 appear : 
 
 2 x 14 parts N unite with 16 parts O. 
 
 14 " " " " 16 " " 
 
 2 x 14 " " " " 3 x 16 " " 
 
 14 " " " " 2 x 16 " " 
 
 2 x 14 " " " " 5 x 16 " " 
 
 Chemical Atoms* It is supposed that the utmost limit 
 to which the division of matter could be carried would lead 
 to its separation into a great number of particles, so small as 
 to be incapable of further division. With reference to their 
 quality of indivisibility, such particles of matter are called 
 atoms (from d, privative, and refAvoD, I cut). Atoms, there- 
 fore, are the indivisible constituents of matter. It is further 
 supposed that chemical combination consists in the union of 
 atoms, or groups of atoms, and chemical decomposition in the 
 separation of atoms, or groups of atoms ; and a chemical 
 change supposes a change in the arrangement or grouping of 
 the atoms of a body involving the destruction of the previous 
 arrangement. 
 
 Atomic Weights. The atomic theory attaches a new 
 meaning to the combining weights of the elements, and defines 
 them as the relative weights of atoms. Thus, if the weight of 
 an atom of hydrogen is i x w (w = a very small unknown 
 
NOTATION. 5 
 
 quantity), the weight of an atom of chlorine is 35.5 x w. 
 Not the absolute weights of atoms, but their relative weights, 
 have been discovered. The weight of an atom of hydrogen is 
 taken as a standard, and called i ; hence the weight of the 
 atoms of other elements are expressed in terms of this unit. 
 Example, 35.5 for chlorine, 16 for oxygen. With reference 
 to the above theory, the combining weights are usually called 
 atomic weights. The full meaning, therefore, of the chemical 
 formula HC1 is, that the body which it represents consists of 
 compound atoms, each one containing an atom of hydrogen 
 and an atom of chlorine. An atom of chlorine weighs 35.5 
 times as much as an atom of hydrogen, so that the proportion 
 by weight of each constituent of the body is expressed by its 
 formula. 
 
 Chemical Notation. Any change in the constitution 
 of bodies, as well as their formation and decomposition, in- 
 volves what is called a chemical reaction. 
 
 A CHEMICAL REACTION may be described as a change in the 
 arrangement, or the state of combination of the atoms of bodies. 
 Such a change can be denoted by combining formulas together 
 in the same way that quantities are combined in common alge- 
 braic calculations. The signs used are +, , and ='. Coeffi- 
 cients are only used to multiply the symbols to which they are 
 joined. When placed on the line, they multiply all the sym- 
 bols which follow. When placed below or above the line, 
 they are used to multiply only the symbol, or the group of 
 symbols in brackets, that immediately precedes them. Brackets 
 are used to distinguish certain groups of atoms in a compound 
 from the remaining atoms. The combination of atoms is ex- 
 pressed by writing their symbols side by side, or by grouping 
 them together without + or . The sign + expresses that 
 the bodies connected by it are brought in contact with each 
 other by addition, but that they are not combined. 
 
 Example : Ba(NO 8 ) 2 + CaSO 4 . The formula indicates 
 that a compound, Ba(NO 8 ) 2 , containing barium, nitrogen, and 
 
6 PARTI. 
 
 oxygen, in the proportion of i atom of barium combined with 
 twice i atom of nitrogen and twice 3 atoms of oxygen, is 
 brought in contact with a compound [CaSOJ containing cal- 
 cium, sulphur, and oxygen, in the proportion of i atom of cal- 
 cium combined with i atom of sulphur and 4 atoms of oxygen. 
 A reaction is denoted by combining the formulas of the bodies 
 which take part in it, as in an ordinary equation. The formulas 
 before the sign ( = ) indicate the state of combination of the 
 atoms before the reaction ; those after the sign of equality (=) 
 show the state of combination of the atoms after the reaction. 
 
 Example : Ba(NO 3 ) a + CaSO 4 = BaSO 4 + Ca(NO 3 ) 2 . 
 
 The equation expresses the result of bringing in contact 
 the bodies described above, viz. : the formation of new com- 
 pounds containing barium, sulphur, and oxygen, and calcium, 
 nitrogen, and oxygen. 
 
 CHEMICAL AFFINITY. 
 
 The force which impels atoms to unite with other atoms is 
 called chemical affinity. The quantity or the nature of the 
 force inherent in the atoms of every substance determines 
 the chemical properties of the substance. The study of the 
 results of the action of chemical affinity is the province of 
 chemistry. 
 
 The phenomena which the action of chemical affinity gives 
 rise to can best be studied under several heads. 
 
 Firstly. Chemical affinity may cause the atoms of an ele- 
 mentary body to unite among themselves. Only the cases of 
 such action in which the element is capable of assuming the 
 simplest physical condition of matter, namely, the form of a 
 gas, have been studied satisfactorily. 
 
 The following conclusions in regard to the state of com- 
 
CHEMICAL AFFINITY. 7 
 
 bination of the atoms of elementary bodies have been ar- 
 rived at : 
 
 The atoms of mercury and zinc remain separate in the 
 gaseous state. 
 
 The atoms of hydrogen, oxygen, chlorine, bromine, iodine, 
 nitrogen (and sulphur at a temperature higher than 1000 
 centigrade) are united in groups of two atoms. 
 
 The atoms of phosphorus and arsenic in the gaseous state 
 unite in groups of four atoms. Sulphur, at a temperature of 
 500 centigrade, in groups of six atoms. 
 
 The formulas for these bodies in a gaseous state are : 
 
 Hg HHorH 2 PPPP or P 4 SSSSSS or S 6 
 
 Zn OO or O 2 AsAsAsAs or As 4 
 
 C1C1 or C1 2 
 
 BrBr or Br 2 
 
 NN or N 2 
 SS or S 2 
 
 Molecule. At this point a definition of the term mole- 
 cule is required. A molecule is the smallest particle of a body 
 which can exist alone. 
 
 The molecules of elementary bodies contain one or more 
 atoms of the same kind. The molecules of compound bodies 
 contain two or more atoms of different kinds. 
 
 A molecule of mercury, hydrogen, hydrochloric acid, or 
 water is represented by the respective formulas, Hg, H 2 , HC1, 
 or H 2 O. 
 
 Atom. An atom may be further defined as the smallest 
 particle of matter which can take part in a chemical reaction. 
 Atoms, therefore, appear while a chemical reaction is going 
 on, although it is impossible to suppose that a physical sub- 
 division of matter could be carried further than the isolation 
 of molecules. 
 
 Thus in the reaction : NasS + CuCl 2 = CuS -f 2 NaCl, the 
 force of chemical affinity breaks up the molecules (Na^) and 
 
8 PART I. 
 
 (CuCl 2 ) to form the new ones (CuS) and (NaCl), and during 
 this reaction the atoms Na, S, Cu, and Cl must be set free 
 from their combinations, and therefore must exist as atoms. 
 
 Secondly. Chemical affinity may combine atoms of a single 
 element, or groups containing atoms of one or more elements, 
 with the atoms of another element, or with groups containing 
 atoms of one or more other elements. 
 
 Examples : Hg -f C1 2 = HgCl 2 . 
 
 NH 3 + HC1 = NH 4 C1. 
 
 The inverse action frequently takes place through the 
 agency of heat or of some other force, and groups of atoms 
 (molecules) break up into other groups, which are usually 
 simpler in constitution than the primitive ones. 
 
 Example : Hg(CN) 2 when heated becomes Hg + (CN) 2 . 
 
 Thirdly. Chemical affinity may cause compound bodies, 
 brought in contact with each other, to mutually exchange 
 some of their constituents ; or an atom or a group of atoms 
 may substitute itself for another atom or for a group of atoms 
 in a compound body. 
 
 Examples : Na 2 CO 8 + BaCl 2 = BaCO 8 + 2NaCl. 
 CuSO 4 + Zn = ZnSO 4 + Cu. 
 
 THE QUALITY OF THE CHEMICAL AFFINITY inherent in the 
 atoms of each element determines the part which the element 
 will play in the different chemical changes mentioned above. 
 It is usually necessary to study each particular case, in order 
 to determine the exact result of bringing in contact any two 
 substances. Empyrical rules, however, defining the nature of 
 the chemical affinity of the elements and the consequences of 
 its action, can be given in a few cases. These rules are not 
 capable of a very strict application, but they serve to indicate, 
 in most cases, when a number of bodies are brought together 
 in a reaction, those which will probably combine with each 
 other. Gold is attacked by acids less readily than the metals 
 
CHEMICAL AFFINITY. 9 
 
 of the arsenic group. The difference between the affinities of 
 the metals which are ranged in the same group in the table 
 (page 2) is too slight to be of consequence in the application 
 of rule second. 
 First. COMBINATION USUALLY OCCURS BETWEEN METALS 
 
 AND NON-METALLIC ELEMENTS, LESS READILY BETWEEN DIF- 
 FERENT NON-METALLIC ELEMENTS, AND LEAST READILY BE- 
 TWEEN DIFFERENT METALS. See table of the elements, page 2. 
 
 The following non-metallic elements, or oxygen compounds 
 of non-metallic elements, combined with hydrogen form acids, 
 and combined with metals form salts. STRONG ACIDS sul- 
 phuric, H 2 SO 4 ; nitric, HNO 3 ; chloric, HC1O 3 ; chlorhydric, 
 HC1. WEAK - ACIDS sulphurous, H 2 SO 3 ; chromic, H 2 CrO 4 ; 
 phosphoric, H 3 PO 4 ; boracic, H 3 BO 3 ; oxalic, H 2 C 2 O 4 ; acetic, 
 HC 2 H 8 O 2 ; fluorhydric, HF ; sulphydric, H 2 S ; cyanhydric, 
 HCN ; carbonic, H 2 CO 3 ; silicic, H 4 SiO 4 . A metal has a ten- 
 dency to substitute itself for the hydrogen in a strong acid, to 
 form a salt with it, in preference to a weak one ; so that a 
 strong acid usually displaces a weak one from its salts. 
 
 Second. THE METALS WHICH STAND LOWEST IN THE TABLE 
 
 (page 2) HAVE THE GREATEST TENDENCY TO COMBINE WITH 
 THE STRONGEST ACIDS. 
 
 Example : CuSO* 4- Zn = ZnSO 4 + Cu. 
 
 Third. Another rule which sometimes takes precedence of 
 the second is, that WHEN FROM SOME OF THE CONSTITUENTS 
 
 OF DIFFERENT COMPOUNDS IN SOLUTION, AN INSOLUBLE BODY 
 CAN BE FORMED, THE ELEMENTS WHICH WOULD COMPOSE 
 SUCH A BODY GENERALLY UNITE WITH EACH OTHER. 
 
 Example: Na 2 CO 3 + Ca(HO) 2 = 2NaHO + CaCO 8 . The 
 compound CaCO 3 is formed in virtue of its insolubility when 
 the reaction takes place in an aqueous solution. 
 
I0 PART I. 
 
 QUANTIVALENCE. 
 
 The quality of the chemical affinity of the elements deter- 
 mines the nature of reactions. The quantity of their chemical 
 affinity determines the proportions in which the elements com- 
 bine with each other. The result of the combination of two 
 or more atoms with each other is the neutralization of the 
 chemical force which brings them together, so that usually no 
 force is left in the compound tending to combine other bodies 
 with it. Thus a molecule of HC1 has no power to combine 
 further with atoms of H or of Cl. The atoms of some ele- 
 ments are animated with greater quantities of chemical force 
 than those of other elements. Thus an atom of oxygen may 
 unite with an atom of hydrogen, and still be capable of com- 
 bining with another atom of hydrogen, or with an atom of 
 chlorine. If an atom of hydrogen has one unit of chemical 
 force, an atom of oxygen has two units, carbon has four, and 
 nitrogen has five. The number of units of chemical force re- 
 siding in an atom is called its QUANTIVALENCE. The table 
 of the elements, page 2, classifies them, according to their 
 quantivalence, into monatomic elements, or monads ; and 
 diatomic elements, or dyads, etc. 
 
 Some of the units of chemical force of an element may lie 
 dormant until developed by the approach of some other force, 
 which awakens them. Iron, for instance, acts as a dyad when 
 no element is present to call forth all of its four units of 
 chemical force. It is worthy of note that in such cases two 
 units of force generally disappear together, as if they became 
 dormant by each one neutralizing the effect of the other. It 
 will be noticed that in the table, page 2, sulphur is placed in 
 the column of dyads, and also in that of the hexads ; this is 
 because in many cases four of the six units of chemical force 
 in sulphur lie dormant, and the element plays the part of a 
 dyad. In nickel, cobalt, and iron, two of the four units of 
 
Q UANTIVA LENCE. ! x 
 
 chemical force frequently lie dormant, and hence these metals 
 often act as dyads. Only the valence of each element, which 
 is usually displayed in the kind of reactions which we have to 
 consider, is shown in the table. 
 
 The knowledge of these facts is essential to enable the stu- 
 dent to write formulas correctly. 
 
 Examples: The compound of barium and chlorine must 
 contain two atoms of chlorine, combined with one atom of 
 barium. Its formula is BaCl 2 . The greatest amount of oxy- 
 gen that an atom of carbon can unite with is two atoms. 
 The formula of the compound is CO 2 . When zinc is substi- 
 tuted for silver in a compound, one atom of zinc, with its two 
 units of chemical force or affinity, takes the place of two 
 atoms of silver, because each atom of silver has only one unit 
 of chemical force. Thus, in writing formulas, an atom of one 
 element is equivalent to or takes the place of another element 
 of the same class. In comparing elements of different classes, 
 their value in an equation depends upon the number of units 
 of chemical force which they contain. 
 
 TWO MONADS ARE EQUIVALENT TO A DYAD ; THREE MO- 
 NADS TO A TRIAD, ETC. THREE DYADS ARE EQUIVALENT TO 
 TWO TRIADS, ETC. 
 
 Examples : PbO + 2 HC1 = PbCl 2 + H 2 O. 
 SiCl 4 + 2H 2 O = SiO 8 + 4HC1. 
 
 QUANTIVALENCE OF GROUPS OF ATOMS. When part of the 
 affinities or units of chemical force of a polyatomic group are 
 satisfied or neutralized, the residual valences, or those which 
 remain free, determine the quantivalence of the group. 
 
 Example : Nitrogen, a pentad, when combined with three 
 atoms of hydrogen, is capable of uniting with two other mo- 
 nads or with one dyad NH 3 can unite with H and Cl to make 
 NH 4 C1. Moreover, certain groups of elements take part in 
 many chemical reactions without being broken up, and, so far 
 as this is the case, they may be considered as playing the part 
 
12 PART L 
 
 of elementary bodies in the reactions. Such groups are fre- 
 quently called radicals, with reference to a theory that they 
 are the roots of compounds. 
 
 The following table, showing the quantivalence of such 
 groups, will be found convenient in writing formulas : 
 
 Monads. Dyads. Triads. Tetrads. Hexads. 
 
 HO* SO4 PO 4 SiO 4 Fe 2 
 
 N0 8 f S0 3 B0 3 Fe(CN) 6 Cr 2 
 
 C10 3 Cr0 4 A1 2 
 
 C 2 H 3 O 2 CO 3 Fe 2 (CN) 12 
 
 NH 4 C 2 O 4 
 
 CN C 4 H 4 O 6 
 H g2 
 
 Examples : Na(HO); H(NO 3 ); (NH 4 ) (HO); Pb(C 2 H 8 O 2 ) 2 ; 
 
 * The group HO is a monad, because one of the two affinities of the 
 oxygen atom is satisfied and nullified by the affinity of the hydrogen with 
 which it is combined, leaving one affinity free in the oxygen atom. 
 
 f The existence of the group NO 3 might seem inconsistent with the 
 five-atomic character of nitrogen. The explanation is, that four affinities 
 in the nitrogen atom are used to unite to it two atoms of oxygen, and the 
 fifth combines with only one of the affinities of a third atom of oxygen, 
 leaving one oxygen affinity free, which determines the monatomic char- 
 acter of the group. 
 
 \ A similar argument applies to the case of the group SO 4 . Here, sul- 
 phur being hexatomic, four affinities are used to combine with those of two 
 atoms of oxygen, while the two remaining affinities of the sulphur atom 
 are each combined with one in each of the two remaining atoms of oxy- 
 gen, leaving two oxygen-affinities free. 
 
 In the case of Fe 2 , etc., two of the eight affinities, belonging to two 
 atoms of iron, are used to bind the two atoms together, leaving six free, 
 and for this reason the group Fe 2 is six-, and not eight-atomic. It may 
 be noticed that groups of this nature, having free chemical affinities, can 
 only be found in the case of polyatomic elements, because, when two 
 monatomic elements combine with each other, no chemical affinity can 
 be left free. For fuller explanations of the laws of chemical combination, 
 see Frankland's Lecture Notes for Chemical Students. 
 
QUANTIVALENCE. ! 3 
 
 H 2 (S0 4 ) ; Ba (CrO 4 ) ; Hg 2 (C1 2 ) ; H 3 (PO 4 ) ; Fe 2 Cl 6 ; Cr 2 O, ; 
 A1 2 (HO).. 
 
 IN WRITING FORMULAS, THE SYMBOL OF THE ELEMENT, OR 
 GROUP OF ELEMENTS, HAVING MOST DECIDEDLY THE CHARAC- 
 TER OF A METAL, IS PLACED FIRST. 
 
 Example : NaNH 4 HPO. 
 
I4 
 
 CHAPTER II. 
 
 CHEMICAL NOMENCLATURE. 
 
 WHEN metallic sodium comes in contact with water a reac- 
 tion takes place, as indicated by the following equation : 
 
 Na* + 2 H 2 O = 2NaHO + H 2 . 
 
 Each molecule of water is broken up, setting free an atom 
 of hydrogen and receiving in its place an atom of sodium. 
 The new body thus formed may then be considered as a mole- 
 cule of water in which one atom of hydrogen has been replaced 
 by sodium. Or it may be considered as an atom of a monad 
 metal united to the monovalent radical HO, known as hy- 
 droxyl. Furthermore, it may be taken as a type of a numerous 
 class of bodies, known as METALLIC HYDRATES, all of which 
 consist of one or more atoms of metal united to an equivalent 
 amount of hydroxyl, as may be seen from the following 
 
 Examples: KHO ; Ca(HO) 2 ; Fe(HO) 2 ; Fe 2 (HO), ; Al, 
 (HO), 
 
 All of the metallic hydrates, when acted upon by a certain 
 other class of bodies, are similarly affected, as can be shown 
 by the following equations : 
 
 HC1 + KHO = KC1 + H 2 ; 
 
 2 HC1 + Ca(HO) 2 = CaCl 2 + 2 H 2 O ; 
 
 6HC1 + Fe 2 (HO) 6 = Fe 2 Cl 6 + 6H 2 O ; 
 
 HNO 8 + KHO = KNO 3 + H 2 O ; 
 
 6HN0 8 + Fe 2 (HO) 6 = Fe 2 (NO 8 ) 6 + 6HO ; 
 
 H 2 S0 4 + KHO = KHS0 4 + H 2 O ; 
 
 H 2 SO 4 + 2KHO = K 2 SO 4 + 2H 2 O ; 
 
 H 2 SO< + Ca(HO) 2 = CaS0 4 + 2H 2 O. 
 
CHEMICAL NOMENCLATURE. ! 5 
 
 By inspection of the above equations it will be seen that we 
 have in the left-hand member of each a metallic hydrate, and 
 in the right-hand member of each water. Furthermore the 
 substance indicated first in each equation contains hydrogen, 
 and when this hydrogen is replaced by the metal of the metal- 
 lic hydrate the formula of the substance indicated first in the 
 right-hand member of the equation is obtained. 
 
 The original substances containing hydrogen are called 
 acids, and the compounds occurring first in the right-hand 
 members of the above equations are salts. 
 
 The above we may formulate, then, in the following defi- 
 nitions : 
 
 Acid. An acid is a compound which, when brought into 
 contact with a metallic hydrate, exchanges the whole or a part 
 of its hydrogen for the metal of the metallic hydrate ; water 
 being formed at the same time. 
 
 The acids may be divided into two groups. 
 
 Firstly, Those which consist of hydrogen united directly to 
 another non-metallic element as HC1 ; HBr ; HF ; H 2 S. 
 
 Secondly, The oxygen acids, or those in which the hydrogen 
 is united with another non-metal, or with a group, by means of 
 oxygen, as in HC1O ; HNO a ; H 2 SO 4 . Here the arrangement 
 of the chemical forces may be shown by the following graphic 
 formulas : H - O - Cl ; H - O - N O 2 ; H 2 = O 2 --= SO 2 . 
 
 Salt . A salt is a compound formed by replacing the hy- 
 drogen of an acid by a metal. 
 
 All compound bodies are susceptible of decomposition by a 
 current of electricity, and when thus decomposed resolve them- 
 selves into a positive atom or group, and a negative atom or 
 group. In writing the ordinary formulas of compounds the 
 electro-positive atom or group is always placed first, as 
 
 + - + - + - + - + 
 
 HC1 ; H 2 ; Na(HO) ; Ba(SO 4 ) ; (NH 4 ) 2 (SO 4 ). 
 
 The names of chemical compounds, which are used in works 
 on analytical chemistry, do not describe their constitution as 
 
16 PART I. 
 
 fully as formulas do ; they serve, however, to identify the 
 bodies and to recall certain principles of classification. 
 
 In the ordinary operations of analysis, the only bodies which' 
 are dealt with are binary compounds (i.e., compounds which 
 only contain two elements, as KC1 and H 2 O) and ternary com- 
 pounds, which contain a non-metallic element and oxygen 
 combined with a metal or with hydrogen, as BaSO 4 , H 3 PO 4 . 
 
 In the statement of the rules for the naming of compounds 
 the subject will be divided under these two heads : 
 
 I. BINARY COMPOUNDS. * 
 
 a. ACIDS. Two distinct methods are in common use for 
 naming binary acids. According to the one the name consists 
 of hydro preceding the root of the name of the other element 
 and terminates in ic; as, HC1, hydrochloric acid ; HBr, hydro- 
 bromic acid ; H 2 S, hydrosulphuric acid. By the other method 
 the characteristic element is indicated first, and the names of 
 the three acids just mentioned would be, chlorhydric acid, 
 bromhydric acid, and sulphydric acid. 
 
 b. ALL OTHER COMPOUNDS. In determining the names of 
 binary compounds not acids, if there exists but one compound 
 of the two elements in question, the positive element takes the 
 termination ic and the negative ide; as, NaCl, sodic chloride ; 
 K 2 S, potassic sulphide ; BaO, baric oxide. Or, inverting the 
 order, they may also be called, respectively, chloride of sodium, 
 sulphide of potassium, and oxide of barium. 
 
 Still another method, and one which has of late been very 
 generally adopted, is to leave the name of the positive element 
 unchanged, while giving the termination ide to the negative. 
 By this method the examples just given would be sodium 
 chloride, potassium sulphide, and barium oxide. 
 
 When two distinct compounds exist, consisting of the same 
 two elements, as SO 2 , SO 3 , that one which contains the greater 
 proportion of the negative element is named in accordance 
 
CHEMICAL NOMENCLATURE. ^ 
 
 with the first method as given above, while the other simply 
 changes the termination ic of the positive element to ous. SO 2 
 would then be sulphurous oxide and SO 3 sulphuric oxide. 
 
 When a compound exists which contains still less of the 
 negative constituent than that indicated by the termination 
 ous y as above, it receives the prefix hypo ; as, N 2 O 3 being ni- 
 trous oxide, N 2 O is called hyponitrous oxide. Similarly the 
 prefix per or hyper is used to indicate that the compound con- 
 tains more of the negative element than the one whose positive 
 constituent terminates in ic. This prefix is, however, usually 
 given to the negative element ; as, Mn 2 O 3 being manganic oxide, 
 MnO 2 is called manganic peroxide. 
 
 The prefix sesqui is sometimes used, as in Fe 2 O 3 , the sesqui 
 oxide of iron. 
 
 The relative positions of the several compounds of a series 
 may be illustrated by the following schedule, in which the plus 
 sign indicates the root of the positive element and the minus 
 sign that of the negative : 
 
 hypo + ous ide, as XO ; 
 
 4- ous ide, as XO 2 ; 
 
 hypo + ic ide, as XO 3 ; 
 
 + ic ide, as XO 4 ; 
 
 + ic per ide, as XO 5 . 
 
 Unfortunately for the rule, however, the practice of recent 
 authors varies very widely, so much so that it is the exception, 
 for example, to find the whole series of oxides of nitrogen 
 similarly named in any two works on chemistry. 
 
 In accordance with the above schedule the following would 
 be the names of the five members of that series : 
 
 N 2 O, hyponitrous oxide ; 
 NO, nitrous oxide ; 
 N 2 O 3 , hyponitric oxide ; 
 NO 2 , nitric oxide ; 
 N 2 O 5 , nitric peroxide. 
 
l8 PART I. 
 
 Instead of these, however, each member of the series has 
 several different names in common use. For example, the first 
 compound is known commonly as nitrous oxide ; it is also 
 called nitrogen monoxide. The second is variously termed 
 nitric oxide, nitrosyl, nitrogen dioxide, since the molecule is 
 by some considered to be indicated by the formula N 2 O 2 , and 
 so on through the whole of the series. 
 
 Another striking example of exception to the rule will be 
 found in the two compounds of carbon and oxygen, CO and 
 CO 2 . Clearly the former should be called carbonous oxide, 
 and the latter carbonic oxide. Formerly, however, CO 2 was 
 considered to be an acid, and was called carbonic acid ; and 
 in fact it is still so termed in most metallurgical and technical 
 works, while CO was known as carbonic oxide. Since the 
 adoption of the present views of the constitution of acids 
 chemists no longer call CO 2 carbonic acid, and it is therefore 
 usually named carbonic dioxide, or carbonic anhydride, from 
 the fact that it may be considered as carbonic acid from which 
 a molecule of water has been removed (H a CO 8 H 2 O = CO 2 ). 
 The name carbonic oxide would be inapplicable, since it would 
 not be perfectly clear whether the compound indicated was 
 CO or CO 2 . As to CO, it still retains its old name carbonic 
 oxide in some works, while in others it is termed carbonous ox- 
 ide, or again carbon monoxide. 
 
 Still another source of difficulty to beginners exists in the 
 fact that many compounds are known by irregular names ; as, 
 H 2 O, water ; NaCl, common salt ; NH 8 , ammonia, etc. 
 
 For the sake of uniformity it would perhaps be well if the 
 system as shown in the following examples could be adopted : 
 
 N 2 O, nitrogen monoxide'; 
 N 2 O 2 , nitrogen dioxide ; 
 N 2 O 3 , nitrogen trioxide ; 
 N 2 O 4 , nitrogen tetroxide ; 
 N 2 O 5 , nitrogen pentoxide ; 
 or, nitric monoxide, etc. 
 
TERNARY COMPOUNDS. ! 9 
 
 Mn 8 O 4 would, according to this plan, be called triman- 
 ganese tetroxide. 
 
 II. TERNARY COMPOUNDS. 
 
 a. ACIDS. The elements hydrogen and oxygen are common 
 to all of this class of acids. They therefore receive their speci- 
 fic names from the third or characteristic element ; as, HNO 8 
 is called nitric acid. In case two or more acids occur, consist- 
 ing of the same three elements united in different proportions, 
 the prefixes and terminations, as explained under binary com- 
 pounds, are applied, and in a similar manner ; as, 
 
 HNO 2 , nitrous acid ; 
 HNO 8 , nitric acid ; 
 H 2 SO 3 , sulphurous acid ; 
 H 2 SO 4 , sulphuric acid ; 
 HC1O, hypochlorous acid ; 
 HC1O 2 , chlorous acid ; 
 HC1O 3 , chloric acid ; 
 HC1O 4 , perchloric acid. 
 
 b. SALTS. The names of salts express both the metal con- 
 tained and the acid from which they are derived. The metal 
 takes the termination ous or ic. The termination of the acid 
 is changed from ous to tie, and from ic to ate j as, 
 
 KNO 2 , potassic nitrite ; 
 NaNO 3 , sodic nitrate ; 
 BaSO 4 , baric sulphate ; 
 FeSO 4 , ferrous sulphate ; 
 Fe 2 (SO 4 ) 8 , ferric sulphate ; 
 KC1O, potassic hypochlorite ; 
 NaClO 2 , sodic chlorite ; 
 KC1O 8 , potassic chlorate ; 
 KC1O 4 , potassic perchlorate. 
 
20 PART I. 
 
 c. METALLIC HYDRATES. The metal generally takes the 
 termination /V, as NaHO, sodic hydrate. Where two hydrates 
 of the same metal exist the one which has the lesser relative 
 amount of the negative radical takes the termination ous; as, 
 Fe(HO) 2 , ferrous hydrate ; Fe 2 (HO) 6 , ferric hydrate. 
 
 In cases where only one hydrate is formed some authors 
 prefer to retain the name of the metal unchanged ; as, sodium 
 hydrate. 
 
 CLASSIFICATION OF ACIDS. 
 
 The hydrogen of an acid which is replaceable by a metal is 
 called basic hydrogen. 
 
 Those acids which contain in the molecule but one atom of 
 basic hydrogen are said to be monobasic ; those which con- 
 tain two such atoms are dibasic ; those containing three are 
 tribasic, and those containing four, tetrabasic. In hypophos- 
 phorous acid, H 3 PO 2 , we have an example of a monobasic 
 acid, since only one of the three atoms of hydrogen is replace- 
 able. In phosphorous acid, H 3 PO 8 , we have a dibasic acid, 
 and in phosphoric acid, H 8 PO 4 , we have a tribasic acid, all 
 three atoms of hydrogen being basic. Pyrophosphoric acid, 
 H 4 P 2 O 7 , is an example of a tetrabasic acid. 
 
 CLASSIFICATION OF SALTS. 
 
 Salts are either normal, acid, double, or basic. 
 
 A normal salt is one in which the whole of the basic hydro- 
 gen is replaced by the metal ; as, NaNO 3 ; Na 2 SO 4 ; NaH 2 PO 2 ; 
 Na, 2 HPO 8 ; Na 8 PO 4 . 
 
 An acid salt is one in which some of the basic hydrogen 
 still remains in the molecule ; as, NaHSO 4 ; NaH 2 PO 3 ; 
 Na 2 HPO 4 . There can of course be no acid salt of a mono- 
 basic acid, since, as the molecule contains but one atom of hy- 
 drogen, if any portion of that element is removed the whole 
 must be. 
 
CHEMICAL REAGENTS. 21 
 
 A double salt is a complex molecule in which either the 
 basic hydrogen is replaced by two distinct metals, or else the 
 same metal replaces the basic hydrogen in two different acids. 
 As examples of the first variety of double salt may be men- 
 tioned (KCl) 2 PtCl 4 , in which the two metals, potassium and 
 platinum, replace the hydrogen in hydrochloric acid ; CaMg 
 (CO 3 ) 2 ; K 2 A1 2 (SO 4 ) 4 . Examples of the second form, 
 
 PbS0 4 , 3PbC0 8 ; PbCl 2 , PbCO 3 . 
 
 Basic salts are complex molecules which may be consid- 
 ered as consisting of a normal salt united to the oxide or the 
 hydrate of the same metal ; as, BiCl 3 , Bi 2 O 3 ; Fe 2 (SO 4 ) 3 , 
 Fe 2 (HO) 6 ; (CuCO 8 ) 2 , Cu(HO) 2 . 
 
 It will be seen by the foregoing examples how much more 
 fully the formulas express the composition of bodies than the 
 names. 
 
 The following list of chemical compounds contains the 
 names and formulas of the substances most frequently used 
 for testing in the laboratory. These names and formulas 
 should be committed to memory, and the rules of nomencla- 
 ture may be studied in their application to them : 
 
 Chlorhydric, or hydrochloric Potassic ferrocyanide, 
 
 acid, HC1 ; K 4 (FeCy 6 ), or K,Ye(CN) 6 ; 
 
 Nitric acid, HNO 3 ; Potassic ferricyanide, 
 
 Sulphuric acid, H 2 SO 4 ; K 6 (Fe 2 Cy, 2 ), or K 6 Fe 2 (CN) 12 ; 
 
 Sulphydric acid, H 2 S ; Potassic sulphocyanate, 
 
 Acetic acid, HC 2 H 3 O 2 ; K(CyS), or KCNS ; 
 
 Ammonic hydrate, NH 4 HO ; Calcic hydrate, Ca(HO) 2 ; 
 
 Ammonic sulphide, (NH 4 ) 2 S ; Calcic chloride, CaCl 2 ; 
 
 Ammonic carbonate, Calcic sulphate, CaSO 4 ; 
 
 (NH 4 ) 2 CO 3 ; Baric chloride, BaCl 2 ; 
 
 Ammonic chloride, NH 4 C1 ; Baric nitrate, Ba(NO 3 ) 2 ; 
 
 Ammonic oxalate, Baric carbonate, BaCO 3 ; 
 
 (NH 4 ) 2 C 2 O 4 ; Magnesic sulphate, MgSO 4 ; 
 
22 
 
 PART /. 
 
 Ferrous sulphate, FeSO 4 ; 
 Ferric chloride, Fe 2 Cl 6 ; 
 Cobaltic nitrate, Co(NO 8 ) 2 ; 
 Plumbic acetate ; 
 
 Pb(C 2 H 3 2 ) 2 ; 
 Argentic nitrate, AgNO 3 ; 
 Mercuric chloride, HgCl 2 ; 
 Platinic chloride, PtCU ; 
 Stannous chloride, SnCl 2 ; 
 Alcohol, C 2 H 6 O. 
 
 After having made himself familiar with the formulas of the 
 substances with which he tests, the student should write in the 
 form of an equation the result of the action of each test which 
 he performs upon a compound under examination. 
 
 Ammonic molybdate, 
 
 (NH 4 ) 2 Mo0 4 ; 
 Sodic hydrate, NaHO ; 
 Sodic carbonate, Na 2 CO 8 ; 
 Disodic hydric phosphate, 
 
 Na 2 HPO 4 ; 
 
 Sodic acetate, NaC 2 H 8 O 2 ; 
 Potassic dichromate, 
 
 K 2 Cr 2 O 7 ; 
 
CHEMICAL OPERATIONS. 
 
 CHAPTER III. 
 
 CHEMICAL OPERATIONS. 
 
 Reaction with Test Paper. A small piece of red 
 or blue litmus paper is dipped in the solution. If it is acid, 
 blue paper is turned red ; if it is alkaline, red paper is turned 
 blue. 
 
 Turmeric paper is turned brown by alkalies. 
 
 Precipitation. When an insoluble body is formed in 
 a solution and separates (falls) from it, precipitation is said to 
 take place. Precipitates are gelatinous, as aluminic hydrate ; 
 flocculent (consisting of flakes), as sulphide of zinc ; or pul- 
 verulent, as baric sulphate. Usually the particles are less finely 
 divided, and filtration is easier with precipitates which form in 
 dilute solutions, particularly in boiling solutions. With some 
 precipitates, as magnesic phosphate, in very dilute solutions, 
 the act of precipitation is a slow process of crystallization, and 
 the formation of a precipitate does not take place until after 
 several hours. 
 
 Filtration is a process by which an insoluble body, 
 usually a precipitate, is separated from a liquid. It is usually 
 important to allow a precipitate to settle before filtration ; and 
 frequently after the clear liquid has been poured upon the 
 filter it is best to add more water to the precipitate, and to 
 wait again until it has settled. When the precipitate requires 
 to be washed, this process may be repeated many times before 
 the precipitate is brought upon the filter. The precipitate 
 usually clogs the pores of the filter, so that it retards the flow 
 of the liquid. A precipitate which has to be washed is finally 
 brought on the filter by rinsing with the wash-bottle, and it is 
 
24 PART I. 
 
 washed by repeatedly filling the filter with water and allowing 
 it to empty itself. The process should be intermittent, and 
 the filter should never be kept constantly full during the latter 
 part of a filtration, when the object is to wash a precipitate 
 with pure water. Sometimes precipitates which take the form 
 of a powder are so finely divided as to pass through the pores 
 of a filter. This can usually be avoided by precipitating in a 
 dilute solution, and particularly by boiling the solution. Sul- 
 phur cannot be prevented from going through the filter. Fil- 
 tration is more rapid with hot than with cold water. Before 
 commencing a filtration, the filter should be made to fit closely 
 to the side of the funnel, and it should always be moistened 
 with pure water. 
 
 Decatfltatiovi consists in allowing a precipitate to settle 
 and in pouring off the liquid above it, in the manner already 
 mentioned under filtration. This process may or may not be 
 united with that of filtration. For instance, argentic chloride 
 settles so completely and quickly that it can usually be washed 
 simply by repeated decantations. 
 
 ^Evaporation is usually performed in a porcelain dish. 
 A few drops of liquid can be evaporated by heating them on 
 platinum foil. Bits of broken glass or porcelain vessels are 
 very useful for the same purpose. 
 
 The Use of the Blowpipe. An olive oil or kerosene 
 lamp, with a wick $ in. long and J in. broad, or a Bunsen's 
 lamp, with the regulator turned to shut off the draught of air, 
 may be used. If the Bunsen's lamp has no regulator for the 
 draught, a smaller tube may be introduced into the lamp-tube, 
 until it rests upon the piece from which the gas issues and 
 excludes the air. After considerable practice a continuous 
 stream of air can be forced through the blowpipe by making 
 the cavity of the mouth the reservoir, into which air is forced 
 at intervals from the lungs, and is prevented from escaping by 
 a peculiar contraction of the throat and hanging palate, which 
 is easily learned ; the breathing goes on through the nose un- 
 
THE USE OF THE BLOWPIPE. 25 
 
 interruptedly. The chief difficulty that beginners usually ex- 
 perience is fatigue of the muscles of the cheeks, which pre- 
 vents a long-continued effort. Two kinds of flames can be 
 produced with the blowpipe one containing an excess of air, 
 consequently of oxygen ; the other containing an excess of 
 combustible gases, consequently of gases capable of consuming 
 oxygen or reducing. 
 
 The Oxidizing Flame is produced by introducing 
 the point of the blowpipe one-third through the lamp-flame. 
 It is a clear blue cone, surrounded and continued at the point 
 by a colorless flame, intensely hot, and capable of producing 
 oxidation. The substance should be heated beyond the point 
 of the blue cone. 
 
 TJie Reducing Flame is produced by holding the 
 point of the blowpipe at the outside of the lamp-flame, and 
 by blowing somewhat more gently. The flame is much less 
 pointed and is more luminous than the oxidizing flame. The 
 substance should be heated at a distance from the point of the 
 flame equal to one-third of its length, and should be completely 
 enveloped in the flame. The position of the blowpipe and 
 the force of the blast regulate the quality of the flame. It is 
 a difficult matter to produce the true reducing flame, which 
 should not deposit carbon on the substance heated in it, and 
 at the same time should contain no excess of oxygen. 
 
 The Borax Bead is formed by making a loop -J- in. 
 in diameter in a piece of platinum wire, heating it red hot in 
 the blowpipe flame, and touching it to a small piece of borax 
 while it is hot. The borax, which adheres to the hot wire, is 
 heated in the blowpipe flame. It at first swells while losing 
 its water of crystallization, and finally it melts to a clear glass 
 bead. A finely divided substance can be taken up by touch- 
 ing the hot bead to it, and it can then be tested as to its solu- 
 bility in the borax bead, coloring properties, etc., in the blow- 
 pipe flame. 
 
 Heating on Charcoal. Select a good piece of char- 
 
26 PART I. 
 
 coal, at least 4 in. long and i in. broad and thick, and smooth 
 a plane surface in a direction at a right angle with that of the 
 year-rings. (If heat is applied to a surface parallel to the 
 planes of the year-rings, the charcoal is more liable to snap 
 from the expulsion of moisture.) In many cases the charcoal 
 serves as a convenient support for a substance to be heated ; 
 in others the reducing agency of the charcoal comes in play. 
 Substances are also evaporated at a high temperature from the 
 surface of the charcoal. 
 
 Ductility, Malleability, Urittleness are charac- 
 teristic properties of metals, and metals can be tested with re- 
 gard to them by pounding with a hammer or by rubbing with 
 the pestle of a mortar. When a substance which appears to 
 contain a higher metal is reduced by sodic carbonate on char- 
 coal, unless metallic globules are at once apparent, the portion 
 of the charcoal which has been heated should be cut out and 
 pulverized with water in a mortar, and washed by decantation. 
 If metallic globules have been formed, they will sink to the 
 bottom, and after thorough washing, during which the sodic 
 carbonate is dissolved, and the light particles of charcoal are 
 floated away, they will appear as globules, if they are hard 
 like copper ; as brittle grains, if they are brittle like bismuth ; 
 or as 'flattened disks, when they are ductile like lead, and when 
 they have been pressed by the pestle against the mortar. 
 
 Color of the Flame. If the substance to be tested is 
 a solid, a small piece of it is brought on a loop of fine platinum 
 wire, or in a pair of forceps, into the flame of the alcohol or 
 Bunsen lamp, and the color imparted to the flame observed. 
 If a substance in solution is to be tested, the platinum wire is 
 dipped in the solution and is then introduced into the lamp- 
 flame. If the solution is too dilute to afford a distinct test in 
 this way, it must be evaporated, and it is usually best to evapo- 
 rate nearly to dryness, and to take some of the solid residue 
 for the test. 
 
 Tlie Manipulation of Glass Tubing. Glass 
 
CHEMICAL OPERATIONS. 27 
 
 softens when a small piece is heated in the flame of an alcohol 
 lamp, or when a larger piece is heated in a Bunsen lamp, or 
 with the blowpipe flame or in a blast lamp. 
 
 A tube can be bent easily as soon as the glass softens. 
 
 It is best only to bend gently at first, then to .heat the adja- 
 cent part of the tube, and to bend again, and so on, in order 
 that the sides of the tube may not fall together in bending. 
 
 When a glass tube is heated for some time, it contracts and 
 the sides thicken. By drawing out a tube either immediately 
 after it has become soft, or after the sides have thickened, a 
 tapering point of any desired calibre and thickness of glass can 
 be obtained. 
 
 To close the end of a glass tube, draw the tube off while the 
 glass is as thin as possible, and hold the tapering point in the 
 flame, and draw the end off again ; in this way a tube with a 
 pointed closed end is obtained ; by heating the closed end of 
 the tube, removing it from the flame, rotating it and blowing 
 in it, while the glass is still red hot, it expands, and a more 
 rounded end or a bulb can be produced. Glass tubing can be 
 cut by making a mark at the required place by a few file 
 strokes, and then by breaking the glass. The ends of glass 
 tubes cut in this way should be held in the flame till they 
 become red hot ; in this way the sharp edges become rounded. 
 
PART II. 
 
 PART II. is preparatory to Part III., which contains a general 
 scheme of analysis applicable to compounds of all the elements. 
 The most important tests are those which are described in Part 
 III., and a knowledge of them would suffice alone for the pur- 
 poses of analysis, if the liability to error in chemical manipu- 
 lations did not make it expedient to employ a variety of tests, 
 as corroborative evidence, before coming to a conclusion in 
 regard to the composition of a substance. 
 
 It is important that the student should turn to Part III., and 
 commit to memory the general features of the scheme of 
 separation for each group, at the time that he is performing 
 the reactions of the members of the group as they are described 
 in the following pages. By this means he will make himself 
 familiar with the important points in which the compounds 
 with which he has to deal differ from each other, and the 
 manner in which these differences can be used in analysis ; also 
 at this stage of his progress it is advisable for him to make 
 mixtures of compounds of several elements, and to analyze 
 them according to the directions given in Part III. 
 
 The student, keeping in view the reasons for learning the 
 characteristic reactions of the compounds of each element, 
 should perform carefully the tests described in Part II., supple- 
 menting the description by the closest observation of the phe- 
 nomena as they pass before his eyes. By practice of this kind 
 he will soon acquire the skill in manipulation necessary for 
 analytical work. ALWAYS, WHEN A REACTION is PERFORMED, 
 
 THE EQUATION DESCRIBING IT SHOULD BE WRITTEN. The for- 
 
 28 
 
PART II. 29 
 
 mulas of the reagents and the compound operated upon* form 
 the first half of the equation ; the formula of the precipitate, 
 which is given in the book, enters into the second half of the 
 equation and determines the formulas of its other members. 
 
 Thus it is known that baric chloride and calcic sulphate give 
 a precipitate of BARIC SULPHATE, BaSO 4 . (See page 34.) 
 From the formulas on the labels of the bottles we can con- 
 struct the equation : BaCl 2 + CaSO 4 = BaSO 4 + X, and by 
 inspecting the equation we find the unknown quantity X can 
 only be CaQ 2 . The following case is more complicated : di- 
 sodic hydric phosphate, ammonic hydrate, and magnesic sul- 
 phate form a precipitate of MAGNESIC AMMONIC PHOSPHATE, 
 MgNH 4 PO 4 ; or, putting the statement into formulas, Na^ 
 HPO 4 + NH 4 HO + MgSO 4 = MgNH 4 PO 4 + X. Here X = 
 Na 2 + H 2 4- O -f SO 4 , and the question arises : How are these 
 bodies combined ? A slight experience will teach that the rule 
 2d (page 9) brings the SO 4 and the Na together, and conse- 
 quently the H 2 and the O ; while an inspection of the tables 
 of quantivalence (pages 2 and 12) shows that SO 4 combines 
 with 2Na, and that O combines with 2H ; hence X becomes 
 Na 2 SO 4 + H 2 O. 
 
 The grouping together of the elements or groups of elements 
 appearing in reactions is not usually a difficult matter, and is 
 soon learned with practice. 
 
 The formulas of the compounds which are most frequently 
 used in the laboratory stand after the names of the metals and 
 acids, and can be used in writing equations. 
 
 * These formulas should be given in full upon the labels of the bottles 
 containing the compounds used and the reagents. 
 
TESTS FOR METALS. 
 
 GROUP I. 
 
 SODIUM, POTASSIUM, AND AMMONIUM. 
 
 THERE is no reagent which precipitates all the metals of this 
 group. The salts of metals of Group I. have a neutral reac- 
 tion when they contain strong acids like chlorhydric, nitric, 
 and sulphuric acids. They have an alkaline reaction when 
 they contain weak acids like sulphydric, boracic, and carbonic 
 acids. 
 
 SODIUM. 
 NaCl ; Na 2 CO 8 ; NaHO. 
 
 Sodium compounds can be recognized by heating them in 
 the loop of a piece of fine platinum wire in the flame of a 
 lamp. Sodium colors the flame yellow, and can be recognized, 
 even when mixed with much larger quantities of other ele- 
 ments, which alone impart other colors to the flame. 
 
 When a liquid is to be tested, it may be evaporated and the 
 residue brought on the platinum wire, or frequently it is suffi- 
 cient to dip the wire in the liquid and to bring it into the flame 
 of the lamp. 
 
 No reagent * precipitates sodium compounds. 
 
 * Only the reagents spoken of in this book are referred to. 
 
 30 
 
POTASSIUM. 31 
 
 POTASSIUM. 
 
 KC1; K 2 SO 4 . 
 
 Potassium compounds color the flame of a lamp violet. A 
 small admixture of sodium obscures the color of the flame of 
 potassium, but the sodium color disappears, and that character- 
 istic of potassium can be observed when the flame is viewed 
 through a thick glass colored blue with cobalt. 
 
 Mix together a sodium and a potassium salt, observe the 
 yellow color imparted to the flame by the mixture, showing 
 the presence of sodium, and then examine the flame for the 
 potassium color through a piece of cobalt glass thick enough 
 to exclude the sodium flame. Examine a pure sodium flame 
 with cobalt glass to ascertain that the glass does not allow the 
 color of the sodium to pass through it, or until the blue color 
 of the sodium flame can be easily distinguished from the violet 
 of the potassium flame.* 
 
 flatinic Chloride precipitates concentrated solutions 
 of potassic chloride as a DOUBLE CHLORIDE OF PLATINUM and 
 POTASSIUM, (KC1) 2 PtCl 4 . No other salt of potassium can be 
 used for this test. It is best to evaporate the solution to dry- 
 ness in a water-bath, with a large quantity of platinic chloride, 
 and to wash the residue several times with alcohol. The 
 double chloride is left as a yellow crystalline powder, which 
 gives the potassium flame. The double chloride is slightly 
 soluble in water, but insoluble in alcohol. 
 
 No other reagent precipitates potassium compounds. 
 
 * If the glass is sufficiently* thick and intense in color, the light from 
 the sodium flame is completely excluded. A thinner glass allows a part 
 of the light to pass through, but it then has a blue color which can be dis- 
 tinguished after practice from the violet color of the potassium flame, 
 which passes through the glass with little of its brilliancy diminished. 
 
32 PART H. 
 
 AMMONIUM. 
 
 NH 4 C1; NH 4 HO. 
 
 Ammonium compounds do not color the flame of a lamp. 
 jPlatinic Chloride precipitates ammonic chloride as a 
 
 DOUBLE CHLORIDE OF PLATINUM AND AMMONIUM (NH 4 C1) 2 
 
 PtCl 4 . The precipitate forms under the same circumstances, 
 and has the same aspect and properties as that obtained with 
 potassic chloride, but it can be distinguished from the latter 
 by the absence of color imparted to the flame of a lamp when 
 it is heated in it. It is destroyed at a dull red heat, and a 
 residue of metallic platinum is left. 
 
 Sodic Hydrate, added in excess to ammonic com- 
 pounds, causes them to give off the smell of AMMONIA GAS, 
 NH 3 , especially when the solution is heated. AMMONIA GAS 
 colors moist turmeric paper brown (a delicate test). 
 
 No other compound interferes with the application of this 
 test. 
 
BARIUM. 
 
 GROUP II. 
 
 BARIUM, STRONTIUM, CALCIUM, AND 
 MAGNESIUM. 
 
 THE chlorides and nitrates of metals of Group II. are soluble 
 in water. The sulphates of calcium and magnesium are also 
 soluble. The solutions have a neutral reaction with test- 
 paper. 
 
 Ammonic and Sodic Carbonates precipitate the 
 metals of Group II. in neutral solutions as CARBONATES. The 
 carbonate of magnesium is very soluble in solutions of ammo- 
 nic salts, particularly in ammonic chloride ; therefore no pre- 
 cipitate of magnesic carbonate is produced when these salts 
 are present in considerable quantity. 
 
 Central "Phosphates (as disodic hydric phosphate) 
 precipitate the metals of Group II. in neutral or alkaline solu- 
 tions as PHOSPHATES. 
 
 The carbonates and phosphates of metals of Group II. are 
 soluble in dilute acids, unless the acids themselves are capable 
 of precipitating the metals. 
 
 Sulphydric Acid, Ammonic Sulphide, and 
 Ammonic Hydrate do not precipitate the metals of 
 Group II. 
 
 BARIUM. 
 
 BaCl 3 ; Ba(NO 3 ) 2 ; Ba(HO) 2 . 
 
 Sulphuric Acid (dilute) precipitates baric compounds, 
 as BARIC SULPHATE, BaSO 4 , white powder. 
 3 
 
34 
 
 PART II. 
 
 Calcic Sulphate and other soluble sulphates give the 
 same precipitate with baric compounds. 
 
 Baric sulphate is insoluble in acids. 
 
 Ammonic Oxalate precipitates baric compounds from 
 concentrated neutral or alkaline solutions, as BARIC OXALATE, 
 BaC 2 O 4 , white powder, soluble in acids. 
 
 Barium compounds, particularly when moistened with chlor- 
 hydric acid, color the flame yellowish green. 
 
 STRONTIUM. 
 
 SrCl 2 ; Sr(N0 3 ) 2 . 
 
 Sulphuric Acid (dilute) and soluble sulphates precipi- 
 tate strontium from its solutions, as STRONTIUM SULPHATE, 
 SrSO 4 , white powder. The precipitate will not appear imme- 
 diately unless the solution be concentrated. In very dilute 
 solutions it may not appear at all, since strontium sulphate is 
 slightly soluble in water (i part in 6900). 
 
 Ammonic Oxalate precipitates strontium compounds 
 from neutral or alkaline solutions, as STRONTIUM OXALATE, 
 SrC 2 O 4 , white powder, soluble in acids. 
 
 Strontium compounds, particularly when moistened with 
 chlorhydric acid, color the flame brilliant red. 
 
 CALCIUM. 
 
 CaCl*; CaS0 4 ; Ca(HO) 2 . 
 
 Sulphuric Acid (dilute) and soluble sulphates, except 
 calcic sulphate, precipitate concentrated solutions of calcium 
 compounds, as CALCIC SULPHATE, CaSO 4 , white powder. Cal- 
 cic sulphate is soluble in a considerable quantity of water, 
 therefore no precipitate is produced in very dilute solutions by 
 sulphuric acid. It is insoluble in dilute alcohol, therefore 
 sulphuric acid (dilute), with the addition of a large quantity of 
 alcohol, precipitates calcic sulphate from even dilute solutions 
 of calcium compounds. 
 
MAGNESIUM. 2$ 
 
 Ammonic Oxalate precipitates calcic compounds from 
 neutral or alkaline solutions, as CALCIC OXALATE, CaC 2 O 4 , 
 white powder. The precipitate forms best after standing some 
 time in a solution to which ammonic hydrate has been added 
 in excess. Calcic oxalate does not dissolve in acetic acid. 
 
 Calcium compounds, particularly when moistened with chlor- 
 hydric acid, color the flame yellowish red. 
 
 MAGNESIUM. 
 
 MgCl 2 ; MgS0 4 . 
 
 Hydro-Disodic JPJiosphate precipitates magnesic 
 compounds, to whose solution ammonic hydrate and ammonic 
 chloride have been added, as MAGNESIC AMMONIC PHOSPHATE, 
 MgNH 4 PO 4 , white crystalline powder, or white flakes if the solu- 
 tion is concentrated. This precipitate only appears after the 
 lapse of some time in very dilute solutions. It is then crys- 
 talline. 
 
 Sodic Hydrate, in excess, precipitates magnesic com- 
 pounds, as MAGNESIC HYDRATE, Mg(HO) 2 , white powder, when 
 it is boiled with their solutions. In case ammonic chloride is 
 present, the boiling must be continued until the odor of ammo- 
 nia is no longer perceptible. 
 
 Mix together the chlorides of barium, strontium, calcium 
 and magnesium. Add a small quantity of ammonic chloride, 
 then ammonic hydrate, and finally ammonic carbonate. The 
 barium, strontium, and calcium will be precipitated as carbon- 
 ates. Filter, and to the filtrate add hydro-disodic phosphate. 
 The magnesium will be precipitated as magnesic ammonic 
 phosphate. Wash the precipitate, consisting of the carbonates 
 of barium, strontium, and calcium, with water, and pour over it 
 a small quantity of dilute hydrochloric acid. This will decom- 
 pose the carbonates, and the chlorides of the three metals will 
 
36 PART II. 
 
 be formed, which being soluble will pass through the filter and 
 can be caught in a test-tube or small beaker. Dilute the solu- 
 tion with considerable water and add dilute sulphuric acid. 
 Barium and strontium will be precipitated as sulphates, while 
 the calcic sulphate will remain in solution. Filter, add am- 
 monic hydrate to the filtrate until the reaction becomes alka- 
 line, and then ammonic oxalate. The calcium will be precipi- 
 tated as oxalate. Wash the sulphates of barium and strontium 
 on the filter, and then moisten a small particle of the mass with 
 chlorhydric acid and examine it on platinum wire in the flame. 
 The yellowish-green color of barium will first appear. After 
 holding the wire in the flame for a few seconds dip it into 
 chlorhydric acid and again bring it into the flame. By repeat- 
 ing this operation several times the barium flame will disappear 
 and the crimson color, characteristic of strontium, will be very 
 plainly visible.* 
 
 * Processes similar to the above are used for the separation of all the 
 metals from each other, and care must always be taken to ascertain whether 
 enough of a reagent has been added to completely effect a precipitation 
 before the next test is proceeded with in the filtrate, and when a precipi- 
 tate is to be examined, it must be thoroughly freed by washing from the 
 solution which adheres to it. 
 
ALUMINIUM. 37 
 
 GROUP III. 
 
 ALUMINIUM AND CHROMIUM. 
 
 THE sulphates, chlorides, and nitrates of metals of Group III. 
 are soluble in water, and the solution has an acid reaction with 
 test-paper. Aluminic and chromic alum solutions have a 
 neutral reaction. 
 
 Ammonic Hydrate, Carbonate, and Sulphide 
 precipitate the metals of Group III. as HYDRATES. 
 
 Neutral Phosphates precipitate the metals of Group 
 
 III. as PHOSPHATES. 
 
 Sulphydric Acid does not precipitate the metals of 
 Group III. 
 
 ALUMINIUM. 
 
 A1 2 (S0 4 ) 3 ; A1 2 C1 6 . 
 
 Ammonic Hydrate precipitates alum inic compounds 
 as ALUMINIC HYDRATE, A1 2 (HO) 6 , gelatinous, white flakes. 
 
 The precipitate is insoluble in ammonic hydrate and in am- 
 monic chloride, but dissolves in acids and in sodic hydrate. It 
 forms best on boiling. 
 
 Sodic Hydrate precipitates aluminic compounds like 
 ammonic hydrate, but an excess of sodic hydrate dissolves the 
 precipitate so quickly that its formation easily escapes notice. 
 
 No precipitate is formed when the solution in sodic hydrate 
 is boiled. 
 
 Solid compounds of aluminium (except silicates), when mois- 
 tened with cobaltic nitrate solution, and heated with the oxid- 
 izing blowpipe flame (see page 25) on charcoal, or on a plat- 
 inum wire, take a blue color. 
 
38 PART II. 
 
 CHROMIUM. 
 
 Cr 2 (S0 4 ) 3 ; Cr 2 Cl 6 . 
 
 Solutions of chromic oxide compounds are green. 
 
 Atnmotlic Hydrate precipitates chromic compounds 
 as CHROMIC HYDRATE, Cr 2 (HO) 6 , gelatinous, dirty green flakes. 
 The precipitate is insoluble in ammonic hydrate after boiling, 
 and in ammonic chloride, but dissolves in acids and in sodic 
 hydrate. 
 
 Sodic Hydrate precipitates chromic compounds like 
 ammonic hydrate, but dissolves them when an excess of sodic 
 hydrate is present. CHROMIC HYDRATE is precipitated from 
 its solution in sodic hydrate when the dilute solution is boiled 
 for some time. 
 
 ^Blowpipe Reactions. Compounds of chromium 
 color the borax bead green. 
 
 If chromic hydrate, or any solid chromic compound, is 
 mixed with equal parts of sodic carbonate and sodic nitrate, 
 and heated red hot on the platinum foil, chromate of sodium 
 is formed by the oxidation of chromic oxide. Chromate of 
 sodium dissolves in water with a yellow color. The color is 
 heightened when an acid is added, and an acid chromate is 
 formed in the solution. 
 
 This reaction is a characteristic test for chromium compounds. 
 
 Mix together solutions of chromium and aluminium salts, 
 add sodic hydrate until the reaction becomes very strongly 
 alkaline (the precipitate which first forms will dissolve), dilute 
 with a considerable quantity of water in a small flask, and boil 
 for several minutes after a dirty green precipitate has formed. 
 Filter from the precipitate. This precipitate contains all the 
 chromium. Test it according to Part III. (102). The fil- 
 trate contains all the aluminium. Test it according to Part 
 III. (104). 
 
ZINC, MANGANESE, AND IRON. 39 
 
 GROUP IV. 
 ZINC, MANGANESE, IRON, NICKEL, AND COBALT. 
 
 THE sulphates, chlorides, and nitrates of metals of Group 
 IV. are soluble in water. The solutions have an acid reaction 
 with test-paper. 
 
 Ammonic Sulphide precipitates the metals of Group 
 IV. as sulphides ; if the solution is not neutral, it should be 
 made so with ammonic hydrate. 
 
 Sodic Hydrate and Ammonic Hydrate pre- 
 cipitate the metals of Group IV. as HYDRATES. The HYDRATE 
 OF ZINC is soluble in an excess of the precipitant, and the HY- 
 DRATES OF NICKEL AND COBALT are soluble in ammonic 
 hydrate. 
 
 Sodic Carbonate precipitates the metals of Group IV. 
 as CARBONATES (ferric compounds as ferric hydrate). 
 
 Neutral Phosphates precipitate the metals of Group 
 
 IV. as PHOSPHATES. 
 
 Sulphydric Acid does not precipitate the metals of 
 Group IV. when they are in an acid solution. (See Zinc, 
 page 40.) 
 
 SECTION I. 
 
 ZINC, MANGANESE, AND IRON. 
 Metals whose sulphides are soluble in cold dilute chlorhydric acid. 
 
 ZINC. 
 
 ZnSO 4 ; ZnCl 2 . 
 
 Metallic zinc dissolves readily in dilute chlorhydric acid 
 and sulphuric acid, with evolution of hydrogen. 
 
40 PART II. 
 
 The metal melts readily when heated on charcoal with the 
 blowpipe, and at a high temperature it distils, and the vapor 
 burns with a bluish-white flame, depositing an incrustation on 
 the charcoal of OXIDE OF ZINC, ZnO, which is white when hot 
 and yellow when cold. If the incrustation is moistened with 
 cobaltic nitrate and heated in the oxidizing flame, it turns 
 dirty green. By this test zinc can often be recognized in 
 alloys. 
 
 Ammonic SulpJiide precipitates zinc compounds as 
 the SULPHIDE OF ZINC, ZnS, white, flocculent precipitate. The 
 sulphide of zinc is soluble in dilute chlorhydric acid, but 
 not in acetic acid. It is the only white insoluble sulphide. 
 
 SulpJiydric Acid precipitates zinc as the SULPHIDE OF 
 ZINC only when the metal is combined with acetic acid. To 
 obtain the precipitate, if a stronger acid is present, add sodic 
 hydrate until the solution has a strongly alkaline reaction, and 
 then add acetic acid, until the reaction becomes acid, before 
 treating with sulphydric acid. 
 
 Sodic and Ammonic Hydrates precipitate zinc 
 compounds as the HYDRATE OF ZINC, Zn(HO) 2 , white flakes. 
 The precipitate dissolves easily in an excess of the precipi- 
 tants ; and the solution in sodic hydrate is not reprecipitated 
 when it is boiled. 
 
 MANGANESE. 
 
 MnSO 4 ; MnCl 2 . 
 
 Solutions of manganese compounds have a faint pink color. 
 
 Ammonic SulpTlide precipitates manganese com- 
 pounds as the SULPHIDE OF MANGANESE, MnS, flesh-colored 
 flakes. The sulphide of manganese is soluble in dilute acids. 
 
 Sodic and Ammonic Hydrates precipitate man- 
 ganese compounds as MANGANOUS HYDRATE, Mn(HO) 2 , white 
 flakes, which turn brown on exposure to the air. Manganous 
 
IRON. 41 
 
 hydrate is insoluble in an excess of the precipitant, but it is 
 soluble in a large quantity of ammonic chloride. 
 
 ^Blowpipe Reactions. Manganese compounds color 
 the borax bead amethyst in the oxidizing flame. If a com- 
 pound of manganese is heated with a soda bead (which can be 
 made in the same way as a borax bead in the loop of a plati- 
 num wire) in the oxidizing flame, it colors it green, in conse- 
 quence of the formation of manganate of sodium. The same 
 color is produced when a compound of manganese is heated 
 on the platinum foil with sodic carbonate and nitrate. On 
 boiling the green salt with water containing a little alcohol it 
 is destroyed, the color disappears, and brown flakes of man- 
 ganic hydrate are precipitated. 
 
 IRON (ferrous salts). 
 FeS0 4 ; FeCl 2 . 
 
 Metallic iron dissolves readily in dilute acids, with evolution 
 of hydrogen. 
 
 Ferrous salts in solution have a pale green color. 
 
 Oxidizing Agents (as nitric acid and potassic chlo- 
 rate), when heated with acid solutions of ferrous salts, oxidize 
 them to FERRIC SALTS, whose color is brownish red, or reddish 
 yellow, and is more intense than the green color of ferrous salts. 
 
 Ammonic Sulphide precipitates ferrous salts as FER- 
 ROUS SULPHIDE, FeS, black flakes. Ferrous sulphide is soluble 
 in dilute acids. 
 
 Sodic and Ammonic Hydrates precipitate ferrous 
 compounds as FERROUS HYDRATE, Fe(HO) 2 , at first nearly 
 white, then bluish green, and, finally, by absorption of oxygen, red- 
 dish brown. Ferrous hydrate is insoluble in an excess of sodic 
 hydrate. The presence of a large quantity of ammonic salts 
 in a solution prevents its precipitation. 
 
 Potassic Ferrocyanide precipitates ferrous com- 
 pounds as POTASSIC FERROUS FERROCYANIDE, K 2 Fe(FeCv 6 ),* 
 
 * Cy, cyanogen, is used as a symbol for the group CN. 
 
42 PART II. 
 
 bluish white ', turning quickly dark blue, through absorption of 
 oxygen from the air. 
 
 Potassic Ferricyanide precipitates ferrous com- 
 pounds, as TURNBULL'S BLUE, Fe 8 (Fe 2 Cyi 2 ),* deep blue. This is 
 the best test for ferrous compounds. 
 
 The last two precipitates are insoluble in dilute acids. 
 
 JPotassic Sulphocyanate gives no coloration with 
 ferrous compounds. 
 
 IRON (ferric salts). 
 Fe 2 Cl 6 . 
 
 Reducing Agents, as sulphurous and sulphydric acids, 
 metallic zinc and iron, reduce ferric salts in solution to ferrous 
 salts when a free acid is present. 
 
 The reaction with sulphydric acid is accompanied by a pre- 
 cipitation of sulphur, Fe 2 Cl 6 + H 2 S = 2FeCl 2 + 2HC1 + S. 
 
 A similar reaction takes place with ammonic sulphide and 
 the other salts of sulphydric acid. 
 
 Ferric salts in solution have & yellow color, and they possess 
 a much stronger coloring power than ferrous salts. 
 
 Ammonic Sulphide reduces ferric salts to ferrous 
 salts, and then precipitates FERROUS SULPHIDE, FeS, black 
 flakes. 
 
 Sodic and Ammonic Hydrates precipitate ferric 
 compounds as FERRIC HYDRATE, Fe 2 (HO) 6 , red gelatinous 
 flakes, insoluble in an excess of the precipitants and in am- 
 monic salts. 
 
 Potassic Ferrocyanide produces a precipitate of 
 PRUSSIAN BLUE, Fe 4 (FeCv 6 ) 3 , deep blue, in a solution of ferric 
 salts. 
 
 JPotassic Ferricyanide colors solutions of ferric salts 
 deep reddish brown, but produces no precipitate. On the ad- 
 dition of a reducing agent a deep blue precipitate forms. 
 
 . * See note on preceding page. 
 
NICKEL AND COBALT. 43 
 
 "Potassic Sulphocyanate gives a deep red color* with 
 the smallest traces of ferric compounds in acid solutions. 
 
 The different behavior with the last three reagents of ferrous 
 and ferric salts serves to distinguish between them. 
 
 ^Blowpipe Heaction. Iron colors the borax bead 
 green in the reducing flame, and reddish yellow while hot, and 
 yellow while cold, in the oxidizing flame. 
 
 Mix together solutions of zinc, manganese, and iron salts ; if 
 a ferrous salt is taken, add a little chlorhydric acid and boil 
 the solution for a few minutes with one or two crystals of po- 
 tassic chlorate, in order to convert the ferrous into ferric salt. 
 Add sodic hydrate to the solution until the reaction is strongly 
 alkaline, boil for a few minutes and filter. All the manganese 
 and iron will be precipitated. Test the precipitate for iron 
 according to Part III. (101), and for manganese according 
 to Part III. (102). All the zinc will be contained in the fil- 
 trate ; test it for zinc according to Part III. (103). 
 
 SECTION II. 
 
 NICKEL AND COBALT. 
 
 Metals whose sulphides are insoluble in cold, dilute chlorhydric acid. 
 
 NICKEL. 
 NiS0 4 ; Ni(N0 3 ) 2 ; NiCl* 
 
 Solutions of nickel salts are green. 
 
 Ammonic Sulphide precipitates nickel compounds as 
 the SULPHIDE OF NICKEL, NiS, black flakes. The sulphide of 
 nickel is insoluble in cold, dilute chlorhydric acid. It dis- 
 solves readily on boiling, or in a strong acid. 
 
 * Potassic sulphocyanate gives the same color with a solution containing 
 a large quantity.of free nitric acid. 
 
44 PART II. 
 
 Sodic and Ammonic Hydrates precipitate nickel 
 compounds as the HYDRATE OF NICKEL, Ni(HO) 2 , apple green. 
 The hydrate of nickel is insoluble in an excess of sodic hy- 
 drate. It dissolves in ammonic hydrate, and the solution has 
 a blue color. 
 
 ^Blowpipe Reactions* Nickel colors the borax bead 
 in the oxidizing flame violet, when it is hot, and a faint reddish- 
 brown when it is cold. By long-continued reduction in the 
 reducing flame, or on charcoal, the bead may be obtained col- 
 orless, but with gray specks of reduced metal in it. 
 
 COBALT. 
 
 Co(N0 3 ) 2 ; CoCl 2 . 
 
 Solutions of cobalt salts, when dilute, are red. 
 
 Ammonic Sulphide precipitates cobalt compounds as 
 the SULPHIDE OF COBALT, CoS, black flakes. The sulphide of 
 cobalt is insoluble in cold dilute chlorhydric acid. It dissolves 
 readily on boiling or in a strong acid. 
 
 Sodic and Ammonic Hydrates precipitate cobalt 
 compounds at first as a blue basic salt, which changes to the 
 pale red COBALTOUS HYDRATE, Co(HO) 2 on boiling. On ex- 
 posure to the air it becomes brown, through the formation of 
 cobaltic hydrate. Cobaltous hydrate is insoluble in an excess 
 of sodic hydrate, but it dissolves in ammonic hydrate, and the 
 solution is red, tinged with brown. 
 
 Blowpipe Reactions. Cobalt compounds color the 
 borax bead blue, and the color is so intense that a small quan- 
 tity of cobalt eclipses the color produced by a much larger 
 quantity of nickel, when the latter is mixed with it. The blue 
 color does not disappear on reduction, so that when sufficient 
 nickel is present to hide the color of a small quantity of cobalt 
 in a bead, the color of the nickel may be made to disap- 
 pear by a thorough reduction, either on the platinum wire or 
 on charcoal, so that the blue color characteristic of cobalt can 
 
COBALT. 45 
 
 be detected in the bead. If the bead was reduced on char- 
 coal, it is advisable to remove it from the charcoal and to melt 
 it on the platinum wire in order to examine its color. 
 
 Mix together sulphates of zinc, manganese, and nickel, and 
 the nitrate of cobalt and ferrous sulphate, add ammonic hy- 
 drate until a permanent precipitate begins to form, and then 
 ammonic sulphide * until the metals are completely precipi- 
 tated as sulphides. Wash the precipitate on a filter and treat 
 it with cold, dilute chlorhydric acid, in order to separate the 
 sulphides of nickel and cobalt from the other sulphides. See 
 Part III. (96). 
 
 Test for nickel and cobalt according to Part III. (97) and 
 (OS). 
 
 Test for the other metals in the chlorhydric acid solution 
 according to Part III. (99), (WO), (101), (102), and 
 (103). 
 
 * In order to precipitate the nickel completely, the solution must not 
 contain free ammonic hydrate, and the ammonic sulphide must not con- 
 tain an excess of ammonic hydrate. If these precautions are not observed 
 a part of the sulphide of nickel dissolves, imparting a brown color to the 
 solution, 
 
46 PART II. 
 
 GROUP F. 
 
 SILVER, MERCURY, LEAD, COPPER, BISMUTH, 
 AND CADMIUM. 
 
 THE salts of metals of Group V. with chlorhydric, nitric, and 
 sulphuric acids, which are soluble in water, have an acid re- 
 action with test-paper. 
 
 Sulphydric Acid and Ammonic Sulphide pre- 
 cipitate metals of Group V. in neutral or acid solutions as 
 SULPHIDES. The sulphides of metals of Group V. are insolu- 
 ble in AMMONIC SULPHIDE and in SODIC HYDRATE, and in dilute 
 acids, even when heated ; but, with the exception of mercuric 
 sulphide, they are all dissolved when boiled with strong NITRIC 
 ACID. 
 
 Sodic and Ammonic Hydrates precipitate metals 
 of Group V. as OXIDES or HYDRATES. The hydrate of lead is 
 somewhat soluble in an excess of sodic hydrate, and the oxide 
 of silver and the hydrates of copper and cadmium are readily 
 soluble in an excess of ammonic hydrate. 
 
 Sodic Carbonate precipitates the metals of Group V. 
 
 as CARBONATES. 
 
 Neutral Phosphates precipitate metals of Group V. 
 in neutral solutions as PHOSPHATES, 
 
 SECTION 7. 
 SILVER, MERCUROUS COMPOUNDS, AND LEAD. 
 
 Metals whose compounds are precipitated as chlorides by chlor- 
 hydric acid. 
 
MERCUROUS COMPOUNDS. 47 
 
 SILVER. 
 AgN0 3 . 
 
 Metallic silver and its alloy with copper dissolve readily in 
 nitric acid. 
 Sulphydric Acid and Ammonic Sulphide 
 
 precipitate compounds of silver as the SULPHIDE OF SILVER, 
 Ag 2 S, black flakes. 
 
 Chlorhydric Acid precipitates compounds of silver as 
 the CHLORIDE OF SILVER, AgCl, white flakes, which settle readily 
 after boiling or prolonged shaking. The chloride of silver dis- 
 solves readily in ammonic hydrate. It is insoluble in concen- 
 trated nitric acid, even when it is boiled with it. 
 
 Sodic and Ammonic Hydrates precipitate com- 
 pounds of silver as the OXIDE OF SILVER, Ag 2 O, a grayish 
 brown powder. 
 
 The oxide of silver dissolves very readily in ammonic hy- 
 drate. 
 
 Blowpipe Reactions. Compounds of silver, mixed 
 with sodic carbonate, are easily reduced on charcoal, by heat- 
 ing in the blowpipe flame, and a hard globule of metallic silver 
 is obtained. 
 
 MERCUROUS COMPOUNDS. 
 
 H g2 (N0 3 ) 2 ; Hg 2 Cl 2 . 
 
 Metallic mercury dissolves readily in nitric acid 'to form 
 mercurous nitrate. 
 
 Copper in the form of a thin sheet or wire can be coated 
 with mercury by immersing it in a solution of mercury con- 
 taining a free acid. The coating is deposited in a longer or 
 shorter time, according to the strength of the solution. It 
 takes the color of mercury when it is rubbed gently with a bit 
 of cloth or paper. The copper should be cleansed by immers- 
 ing it in dilute nitric acid solution before it is put in the solu- 
 tion containing mercury. 
 
48 PART II. 
 
 Oxidi&ing Agents, as chlorine, aqua regia, or strong 
 nitric acid, transform mercurous into mercuric compounds. 
 Sulphydric Acid and Ammonic Sulphide 
 
 precipitate mercurous compounds as MERCUROUS SULPHIDE, 
 Hg 2 S, black. Mercurous sulphide is not dissolved when it is 
 boiled with moderately strong nitric acid. It dissolves readily 
 in aqua regia. 
 
 Chlorhydric Acid precipitates mercurous compounds 
 as MERCUROUS CHLORIDE, Hg 2 Cl 2 , white powder. Mercurous 
 chloride turns black, but does not dissolve when ammonic hy- 
 drate is added to it. It is insoluble in dilute acids. It dis- 
 solves in aqua regia. 
 
 Sodic and Ammonic Hydrates give with mer- 
 curous compounds black precipitates, insoluble in an excess of 
 the precipitant. (For tests by heating mercurous compounds, 
 see page 50.) 
 
 LEAD. 
 
 Pb(C 2 H 8 2 ) 2 ; Pb(N0 3 ) 2 . 
 
 Metallic lead dissolves readily in nitric acid. 
 
 Sulphydric Acid and Ammonic Sulphide pre- 
 cipitate lead compounds as the SULPHIDE OF LEAD, PbS, 
 black. Sulphide of lead is oxidized by strong nitric acid, with 
 formation of sulphate of lead, white powder y which is insoluble, 
 unless a very large quantity of nitric acid is present. 
 
 Chlorhydric Acid precipitates lead compounds as 
 the CHLORIDE OF LEAD, PbCl 2 , white. When the solution is 
 dilute no precipitate is produced. The chloride of lead dis- 
 solves entirely on boiling with a large quantity of water. 
 
 Sulphuric Acid precipitates lead compounds as the 
 SULPHATE OF LEAD, PbSO*, white powder. Sulphate of lead 
 is soluble to some extent in chlorhydric and nitric acids, and 
 it is slightly soluble in water. It is insoluble in a mixture of 
 alcohol and water. When complete precipitation is required, 
 it is best to evaporate with an excess of sulphuric acid until 
 
LEAD. 49 
 
 all the other acids are driven off, then to dilute with water, 
 and to add an equal bulk of alcohol. 
 
 LEAD, BARIUM, STRONTIUM, CALCIUM, and STANNIC COM- 
 POUNDS are the only ones precipitated by SULPHURIC ACID. 
 
 Sodic and Ammonic Hydrates precipitate com- 
 pounds of lead as BASIC SALTS OF LEAD, white precipitate. 
 Basic salts of lead are somewhat soluble in sodic hydrate. 
 
 blowpipe Reactions. Solid compounds of lead give, 
 when heated with sodic carbonate on charcoal, globules of 
 metallic lead, recognizable by their softness and ductility. An 
 incrustation of OXIDE OF LEAD, PbO, deep yellow when hot, light 
 yellow when cold, is formed upon the charcoal, not far from the 
 place where the substance is heated. 
 
 Pure lead, or alloys containing a large amount of lead, when 
 heated on charcoal, without soda, burn with a blue flame, and 
 give the PbO incrustation. 
 
 Mix together acetate of lead, nitrate of silver, and mercurous 
 nitrate, add dilute chlorhydric acid until a precipitate ceases 
 to form on further addition, and filter the liquid. The pre- 
 cipitate contains all the silver and mercury and a part of the 
 lead as chlorides. Add to the filtrate an equal bulk of alcohol 
 and a small quantity of dilute sulphuric acid. All the lead 
 which it contains will be precipitated. Filter, wash the pre- 
 cipitate, and test it for lead on charcoal. 
 
 Make a hole in the bottom of the filter and wash the pre- 
 cipitate obtained with chlorhydric acid into a small flask. 
 Test the precipitate for the remainder of the lead, and for 
 silver and mercury, according to Part III. (#7), (08), and 
 (69). 
 
 4 
 
PART II. 
 
 SECTION II. 
 
 MERCURIC COMPOUNDS, COPPER, BISMUTH, AND CADMIUM. 
 Metals which are not precipitated by chlorhydric acid. 
 
 MERCURIC COMPOUNDS. 
 
 HgCl 2 . 
 
 Snlphydric Acid and Ammonic Sulphide pre- 
 cipitate mercuric compounds, at first as double salts, which 
 appear first white, then yellow, orange, and brown, and finally as 
 MERCURIC SULPHIDE, HgS, black. Mercuric sulphide does not 
 dissolve when it is boiled with moderately concentrated nitric 
 or chlorhydric acid. It dissolves readily in aqua regia. 
 
 Sodic Hydrate precipitates mercuric compounds at 
 first as basic salts, reddish brown, finally as MERCURIC OXIDE, 
 HgO, yellow ; insoluble in an excess of the precipitant. 
 
 Ammonic Hydrate precipitates mercuric compounds 
 as SALTS CONTAINING AMMONIA, white. The precipitate is 
 insoluble in an excess of ammonic hydrate. 
 
 StannoMS Chloride reduces mercuric compounds, and 
 precipitates them as MERCUROUS CHLORIDE, Hg 2 Cl 2 , white. 
 After the metals of Group V., Section I., if they are present 
 have been removed from a solution by the addition of chlorhy- 
 dric acid, mercuric compounds are the only ones which give a 
 precipitate in acid solution with stannous chloride. 
 
 Reactions with the Aid of Heat. Dry mercurous 
 and mercuric chlorides form white sublimates when they are 
 heated in a closed glass tube. All dry compounds of mer- 
 cury, when they are mixed with dry sodic carbonate, and 
 heated in a closed tube, give a sublimate of metallic mercury. 
 The sublimate is at first a faint metallic film, which augments 
 until drops of mercury appear. If the quantity of mercury is 
 
COPPER. 
 
 small, the film may be made to take the form of metallic drops 
 by rubbing it with a copper wire. 
 
 COPPER. 
 CuSCV, CuCl 2 . 
 
 Metallic copper dissolves readily in dilute nitric acid. It 
 dissolves with difficulty in chlorhydric acid. All copper solu- 
 tions are blue or green. 
 
 Iron OT Zinc precipitates copper from its acid solutions, 
 either as a metallic coating or as brownish red metallic grains. 
 If a strip'of zinc and one of platinum are placed in a dilute 
 acid solution of copper, so that they touch each other, the 
 platinum is plated with copper. 
 
 Sulphydric Acid and Ammonic Sulphide 
 precipitate cupric compounds as CUPRIC SULPHIDE, CuS, black. 
 Cupric sulphide is insoluble in dilute acids, but dissolves readily 
 in strong acids. It is somewhat soluble in an excess of am- 
 monic sulphide. 
 
 Sodic Hydrate precipitates cupric compounds as 
 CUPRIC HYDRATE, Cu(HO) 2 , light blue. On boiling, CUPRIC 
 OXIDE, CuO, black, is formed. The precipitate is insoluble in 
 an excess of sodic hydrate. 
 
 Ammonic Hydrate precipitates cupric compounds as 
 CUPRIC HYDRATE, which dissolves immediately in an excess of 
 ammonic hydrate. The solution has a very intense blue color. 
 
 A valuable test for cupric compounds. 
 
 Ferrocyanide of Potassium precipitates cupric 
 compounds in acid solutions, as FERROCYANIDE OF COPPER, 
 Cu 2 (FeCy 6 ), reddish brown. This is a delicate test for very 
 small quantities of copper. 
 
 Blowpipe Reactions* Cupric compounds color the 
 borax bead blue when cold, and green when hot, in the oxidiz- 
 ing flame. The bead is colored red, and becomes opaque in 
 
ij2 PART II. 
 
 the reducing flame. No other metal produces this color- 
 ation. 
 
 When cupric compounds are heated with sodic carbonate 
 on charcoal, metallic copper is reduced in the form of small 
 globules, which can be easily recognized by their hardness and 
 red color. 
 
 Copper Flame. Compounds containing copper (alloys 
 and salts) color the flame of a lamp green, or if chlorine is 
 present, blue. By moistening a cupric compound with chlorhy- 
 dric acid the blue color can easily be detected. 
 
 BISMUTH. 
 Bi(N0 3 ) 3 . 
 
 Metallic bismuth dissolves readily in moderately concen- 
 trated nitric acid. It dissolves with great difficulty in chlor- 
 hydric acid. 
 
 Water, Solutions of bismuth, particularly those contain- 
 ing chlorhydric acid, are remarkable for giving a precipitate 
 consisting of a BASIC SALT, when water is added to them, 
 unless they contain a large quantity o free acid. The basic salt 
 can be dissolved by the addition of an acid. CHLORHYDRIC 
 ACID precipitates nitrate of bismuth solution as a BASIC CHLOR- 
 IDE because the basic chloride of bismuth is more insoluble 
 than the other basic salts. . The precipitate is soluble on the 
 further addition of chlorhydric acid. 
 
 Sulphydric Acid and Ammonic Sulphide pre- 
 cipitate bismuth compounds as the SULPHIDE OF BISMUTH, 
 Bi 3 S 8 , black. 
 
 Sodic and Ammonic Hydrates precipitate bis- 
 muth compounds as the HYDRATE OF BISMUTH, Bi(HO) 3 , white. 
 
 The hydrate of bismuth is insoluble in an excess of sodic 
 and ammonic hydrates. 
 
 Blowpipe Reactions. Bismuth compounds, mixed 
 with sodic carbonate, and heated on charcoal, give brittle 
 
CADMIUM. 
 
 53 
 
 metallic globules, and an incrustation of OXIDE OF BISMUTH, 
 Bi 2 O 3 , on the charcoal. 
 
 The incrustation is orange yellow when hot, and bright yellow 
 when cold. 
 
 CADMIUM. 
 
 Cd. 
 
 Metallic cadmium dissolves readily in nitric acid, and the 
 solution contains nitrate of cadmium, Cd(NO 8 ) 2 . 
 Sulphydric Acid and Ammonic Sulphide 
 
 precipitate cadmium compounds as the sulphide, CdS, lemon 
 yellow. This is insoluble in ammonic sulphide and potassium 
 cyanide, but readily soluble in hot dilute sulphuric or nitric 
 acid. 
 
 Sodic and Ammonic Hydrates precipitate com- 
 pounds of cadmium as the hydrate, Cd(HO) 2 , very readily solu- 
 ble in a slight excess of ammonic hydrate, but insoluble in 
 sodic hydrate. 
 
 When compounds of cadmium are mixed with sodic carbon- 
 ate and exposed on charcoal to the inner blowpipe flame, the 
 charcoal becomes covered with an incrustation of yellow or 
 brownish-yellow oxide of cadmium, CdO. 
 
 Mix together mercuric chloride, cupric sulphate, and the 
 nitrates of bismuth and cadmium, and add sulphydric acid 
 until the metals are completely precipitated. Filter, and wash 
 the precipitates with a little water ; spread the filter out ; 
 scrape the precipitate from it, and heat the precipitate gently 
 in a porcelain dish with strong nitric acid until red fumes 
 cease to be given off ; then add a little water and boil for a few 
 minutes. The sulphide of mercury remains insoluble. Filter 
 it off and test according to Part III. (89). Test the filtrate 
 for bismuth, copper, and cadmium according to Part III. (91), 
 (92), and (93). 
 
54 PART II. 
 
 GROUP VI. 
 TIN, ANTIMONY, ARSENIC, AND GOLD. 
 
 THE metals of Group VI. sometimes act as acids, uniting 
 with metals, and sometimes as metals, uniting with acids. 
 Their combinations with acids have an acid reaction ; those 
 with metals which are soluble in water have an alkaline re- 
 action. 
 
 Sulphydric Acid precipitates the metals of Group VI. 
 as SULPHIDES. The precipitation takes place slowly, particu- 
 larly in the case of arsenic acid, and it should never be con- 
 sidered complete unless a current of sulphydric acid gas is 
 passed through the solution for some time, and it is left to 
 stand twenty-four hours. 
 
 The sulphides of metals of Group VI. are insoluble in dilute 
 acids, but they are decomposed or dissolved by boiling with 
 concentrated acids. They are soluble in sodic hydrate and 
 in ammonic sulphide (the sulphide of gold with difficulty). 
 
 TIN. 
 
 Sn. 
 
 Metallic tin dissolves readily in strong chlorhydric acid on 
 boiling, and the solution contains STANNOUS CHLORIDE, SnCl 3 . 
 Tin is oxidized to STANNIC OXIDE, SnO 2 , white powder, by 
 strong nitric acid. Stannic oxide is insoluble in nitric acid, 
 but dissolves readily in hot concentrated chlorhydric acid, 
 and the solution contains STANNIC CHLORIDE, SnCl 4 . Stannic 
 chloride is also formed by the action of aqua regia on metallic 
 tin. 
 
STANNOUS COMPOUNDS. 
 
 55 
 
 Zinc precipitates METALLIC TIN from its solu- 
 tions in acids as crystalline metallic particles. 
 
 ^Blowpipe ^Reactions. Compounds of tin are re- 
 duced, when they are heated with sodic carbonate and potas- 
 sic cyanide on charcoal ; and METALLIC TIN may be discov- 
 ered in flattened globules by rubbing the fused mass, taken 
 from the charcoal, in a mortar with water, and washing several 
 times by decantation. When the globules are large, they may 
 be observed on the charcoal during the fusion. 
 
 Oxidation by Sodic Nitrate. When a sulphide of 
 tin is oxidized at the lowest possible temperature by a mixture 
 of sodic nitrate and carbonate, STANNIC OXIDE, SnO 2 , which 
 is insoluble in water, is formed. If the oxidation is carried on 
 at too high a temperature, stannate of sodium is formed, which 
 is soluble in water. 
 
 STANNOUS COMPOUNDS. 
 
 SnCl 2 . 
 
 Snlphydric Acid and Ammonic Sulphide pre- 
 cipitate stannous compounds from acid solutions as STANNOUS 
 SULPHIDE, SnS, dark brawn. Stannous sulphide dissolves with 
 difficulty in a colorless solution of ammonic sulphide (the 
 mono- sulphide), and it is scarcely soluble in ammonic hydrate 
 and carbonate. It is converted into stannic sulphide by a 
 yellow solution of ammonic sulphide (the poly-sulphide), and 
 dissolves readily when warmed with the solution. Stannous 
 sulphide is soluble in sodic hydrate. The sulphides of tin are 
 precipitated from these solutions, when a dilute acid is added 
 gradually, until the reaction becomes strongly acid. 
 
 Sodic and Ammonic Hydrates and Carbo- 
 nates precipitate stannous compounds as STANNOUS HY- 
 DRATE, Sn(HO) 2 , white. 
 
 Stannous hydrate dissolves in an excess of sodic hydrate, 
 
56 PART 77. 
 
 but it is insoluble in an excess of sodic carbonate and of am- 
 monic hydrate and carbonate. 
 
 Mercuric Chloride changes stannous chloride, SnCl 2 , 
 into STANNIC CHLORIDE, SnCl 4 , and a white precipitate of MER- 
 CUROUS CHLORIDE, Hg 2 Cl 2 , is formed. No other metal in solu- 
 tion gives this reaction with mercuric compounds. 
 
 STANNIC COMPOUNDS. 
 
 SnCl<. 
 
 Swlphydric Acid and Ammonic Sulphide 
 
 precipitate stannic compounds as STANNIC SULPHIDE, SnS 2 , 
 light yellow. Stannic sulphide dissolves readily in ammonic 
 sulphide, and in sodic hydrate, and is precipitated completely 
 from the solution, when a dilute acid is added gradually until 
 the reaction becomes strongly acid. Stannic sulphide is nearly 
 insoluble in ammonic carbonate. 
 
 Sodic and Ammonic Hydrates and Carbon- 
 ates precipitate stannic compounds as STANNIC HYDRATE, 
 Sn(HO) 4 , white. Easily soluble in sodic hydrate. 
 
 Mercuric Chloride does not give a precipitate with 
 stannic chloride. 
 
 Metallic Zinc precipitates TIN as crystalline metallic 
 particles from stannic compounds in an acid solution, and the 
 tin can be easily recognized by dissolving the metallic parti- 
 cles, after they have been washed by decantation. They are 
 dissolved by heating them with a few drops of strong chlorhy- 
 dric acid, and the stannous chloride thus obtained gives the 
 precipitate, above described, with mercuric chloride. 
 
 ANTIMONY. 
 
 Sb. 
 
 Metallic antimony is scarcely attacked by chlorhydric acid. 
 It is oxidized when heated with moderately strong nitric acid 
 
ANTIMONY. 57 
 
 to ANTIMONIC OXIDE, Sb 2 O 5 , which is almost completely inso- 
 luble in nitric acid, but is readily soluble in hot concentrated 
 chlorhydric acid. Aqua regia dissolves metallic antimony as 
 
 ANTIMONIC CHLORIDE, SbCl 5 . 
 
 Water precipitates solutions of antimony containing chlor- 
 hydric acid, particularly those of antimonious compounds, as 
 a BASIC CHLORIDE. The precipitation can be prevented, or 
 the precipitate dissolved by the addition of a sufficient quan- 
 tity of acid ; for this purpose tartaric acid is the most suitable. 
 
 Metallic Zinc partially precipitates METALLIC ANTI- 
 MONY from its acid solutions, and when a piece of platinum 
 in contact with a piece of zinc is introduced into solutions of 
 antimony containing an excess of chlorhydric acid, metallic 
 antimony is deposited upon the platinum as a dark-brown stain. 
 No other metal produces the same stain under like circumstances. 
 
 Antimoniuretted Hydrogen. To obtain this body, 
 follow exactly the directions given for obtaining arseniuretted 
 hydrogen. (See page 59.) 
 
 The mirror obtained with antimoniuretted hydrogen consists 
 of a black, sooty, metallic deposit. It dissolves very slowly in 
 hypochlorite of sodium. If the deposit is moistened with 
 yellow ammonic sulphide, an orange-yellow stain appears on the 
 spot when it is dried. These reactions are unimportant as 
 tests for antimony, but a knowledge of them is necessary, in 
 order that they may not be mistaken for evidences of the pres- 
 ence of arsenic. 
 
 ^Blowpipe Reactions* Antimony compounds, mixed 
 with sodic carbonate and potassic cyanide, are quickly reduced, 
 at a comparatively low temperature, to METALLIC ANTIMONY, 
 brittle shining grains. The metal is completely volatilized by 
 a strong heat. Metallic antimony, when heated on charcoal, 
 burns with a white smoke, and gives an incrustation which is 
 deposited at some distance from the place heated, and is very 
 volatile before the blowpipe flame. 
 
 Oxidation by Sodic Nitrate. When a sulphide of 
 
5 8 PART II. 
 
 antimony is oxidized by a mixture of sodic carbonate and 
 nitrate, ANTIMONIATE OF SODIUM, Na 3 SbO 4 , insoluble in water, 
 is formed. 
 
 ANTIMONIOUS COMPOUNDS. 
 
 SbCl 3 ; KSbOC 4 H 4 O 6 , Tartar Emetic. 
 
 Sulphydric Acid and Ammonic Sulphide pre- 
 cipitate antimonious compounds as ANTIMONIOUS SULPHIDE, 
 Sb 2 S 8 , orange red. The precipitation should be made in a cold 
 solution containing tartaric acid and very little free chlorhy- 
 dric acid. Antimonious sulphide dissolves in sodic hydrate 
 and in ammonic sulphide, most readily in yellow ammonic sul- 
 phide, and is precipitated completely from these solutions when 
 a dilute acid is added gradually until the reaction becomes 
 acid. It is nearly insoluble in ammonic carbonate. Yellow 
 ammonic sulphide converts it into ANTIMONIC SULPHIDE, Sb 2 S 6 . 
 (See below.) 
 
 Sodic and Ammonic Hydrates and Carbon- 
 ates precipitate antimonious compounds as ANTIMONIOUS 
 OXIDE, Sb 2 O 3 , voluminous white precipitate. The precipitation 
 only takes place after the lapse of a considerable time in solu- 
 tions containing tartaric acid. Antimonious oxide is soluble in 
 sodic hydrate. 
 
 ANTIMONIC COMPOUNDS. 
 
 SbCl 5 ; Sb 2 6 . 
 
 Sulphydric Acid and Ammonic Sulphide pre- 
 cipitate antimonic compounds in acid solution as ANTIMONIC 
 SULPHIDE, Sb 2 S 6 , orange red. The precipitation does not take 
 place immediately ; but first an orange color appears ; and it is 
 only after passing sulphydric acid for a long time through the 
 solution that all the antimonic sulphide is precipitated. The 
 precipitation should be made in a cold solution containing tar- 
 
ARSENIC. 
 
 59 
 
 taric acid and very little free chlorhydric acid. Antimonic 
 sulphide has the same properties as antimonious sulphide. 
 
 Antimonic oxide, Sb 2 O 5 , plays the part of an acid with bases, 
 and forms insoluble compounds with SODIUM, and soluble 
 compounds with POTASSIUM. 
 
 Compounds of antimony can best be recognized by the stain, 
 which they form on platinum, when the metal is precipitated 
 by zinc from their solutions. (See page 57, Antimony.) 
 
 ARSENIC. 
 As. 
 
 Metallic arsenic is readily oxidized to arsenious or to ar- 
 senic compounds by nitric acid and is dissolved. It is not dis- 
 solved by chlorhydric acid. Metallic arsenic volatilizes at a 
 heat below redness in a tube or on charcoal, and produces an 
 odor like garlic, very characteristic of compounds of arsenic. 
 The same odor is produced by heating a dry compound of 
 arsenic on charcoal. 
 
 Arseniuretted Hydrogen. To obtain this body, 
 and to use Marsh's test, provide a 4 oz. flask with a tight- 
 fitting cork, into which a funnel-tube and a small tube drawn 
 off to a point and bent at a right angle * are introduced. Fill 
 it two-thirds full of dilute sulphuric acid, and add several 
 pieces of pure zinc. A brisk evolution of hydrogen only com- 
 mences after a few minutes. When that point is reached, 
 wait five minutes for the expulsion of the air contained in the 
 flask (without this precaution there is danger of an explosion), 
 and light the hydrogen issuing from the point of the bent tube. 
 The opening of the point and the quantity of gas evolved 
 should be such that the hydrogen burns with a blunt flame 
 
 * In order to dry the gas, it is better to fit to the cork a chloride of cal- 
 cium tube, bent downwards, so that it will not tip the flask over by its 
 weight, and to adapt to the chloride of calcium tube, by means of a cork, 
 a suitable jet directed upwards. 
 
60 PART II. 
 
 about J in. in length. Hold a bit of porcelain in the flame, in 
 order to be certain that the reagents employed are free from 
 arsenic. (See below.) If this is the case, pour a little of an 
 acid liquid containing arsenic into the funnel-tube, while the 
 hydrogen flame continues lighted. The flame after a few 
 moments becomes white, and leaves a black stain of metallic 
 arsenic upon a cold porcelain object held in it, in the same 
 way that the flame of a candle would leave a deposit of soot 
 on a cold surface. The arsenic stain or mirror is shining black 
 (not sooty black, like antimony). It dissolves quickly in hypo- 
 chlorite of sodium. When the stain is moistened with yellow 
 ammonic sulphide solution, and the spot is dried, it is bright 
 yellow. 
 
 Reaction with the Aid of Heat. All dry com- 
 pounds of arsenic, except some compounds with the higher 
 metals, when they are mixed with dry sodic carbonate and po- 
 tassic cyanide, and heated in a sealed tube, give a sublimate 
 of metallic arsenic, which can be best recognized by breaking 
 the sealed end of the tube after the formation of the sublimate, 
 and by heating the sublimate quickly until it begins to volatil- 
 ize, and by smelling of the upper end of the tube. A smell 
 of garlic is proof of the presence of arsenic. 
 
 ARSENIOUS COMPOUNDS. 
 
 As 2 O 8 . 
 
 Arsenious oxide is sparingly soluble in water It dissolves 
 more readily in sodic hydrate or carbonate, or in chlorhydric 
 acid. 
 
 Sulphydric Acid and Ammonic SulpJiide pre- 
 cipitate arsenious compounds from acid solutions as ARSENI- 
 OUS SULPHIDE, As 2 S 8 , yellow. Arsenious sulphide dissolves 
 very readily in sodic hydrate and ammonic sulphide, and it 
 also dissolves, although less readily, in ammonic carbonate. 
 
ARSENIC COMPOUNDS. 61 
 
 It is precipitated from these solutions, when a dilute acid is 
 slowly added until the reaction becomes acid. 
 
 Sodic and Ammonic Hydrates and Carbon- 
 ates produce no precipitate in arsenious compounds. 
 
 ARSENIC COMPOUNDS. 
 
 As 2 O 5 . 
 
 By heating any compound of arsenic with strong nitric acid, 
 or by fusing a dry compound with a mixture of dry sodic car- 
 bonate and nitrate in a porcelain crucible, arsenic acid, 
 H 3 AsO 4 , or arseniate of soda, Na^AsC)^ is formed. These com- 
 pounds are soluble in water. 
 
 Sulphydric Acid and Ammonic Sulphide do 
 not immediately precipitate acid solutions of arsenic acid. 
 
 The solution (which should contain free chlorhydric acid) 
 at first becomes yellow, when a current of sulphydric acid is 
 passed through it, and finally a yellow precipitate is formed ; 
 reduction to the state of an arsenious compound takes place 
 slowly, and at the end of several days all the arsenic is pre- 
 cipitated as ARSENIOUS SULPHIDE, As 2 S 3 , yellow. This opera- 
 tion may be very materially accelerated by boiling the solution. 
 
 See the properties of arsenious sulphide above. 
 
 Argentic ^Nitrate produces no precipitate in acid 
 solutions of arsenic compounds. {When chlorhydric acid is 
 present a white precipitate of argentic chloride forms, and 
 must usually, after the addition of an excess of argentic nitrate, 
 be separated by decantation and filtration.) If sodic hydrate 
 is added to a clear solution containing arsenic acid, and con- 
 taining also an excess of argentic nitrate, until a permanent 
 precipitate forms, and if then acetic acid is added until the 
 reaction becomes acid, ARGENTIC ARSENIATE, Ag 3 AsO 4 , brick 
 red, is precipitated. This precipitate is soluble in ammonic 
 hydrate and in dilute nitric acid. 
 
 This is the usual test for arsenic. 
 
62 PART II. 
 
 GOLD. 
 
 AuCl 8 . 
 
 Metallic gold is insoluble in any single acid, but it dissolves 
 readily in aqua regia. 
 
 Ferrous Sulphate solution in considerable quantity 
 precipitates METALLIC GOLD, Au, purplish-brown powder, from 
 acid solutions containing gold. 
 
 Stan/nous Chloride solution precipitates METALLIC 
 GOLD, Au, purple flakes, very finely divided, from acid solutions 
 containing gold. (A very dilute solution of stannous chloride 
 should be added, drop by drop, to the gold solution.) 
 
 Sulphyciric A.d(l gas slowly precipitates gold com- 
 pounds in acid solutions as SULPHIDE OF GOLD, Au 2 S 3 , brown 
 flakes. The precipitate is soluble after a long digestion with 
 yellow ammonic sulphide. 
 
 ^Blowpipe Reaction. Gold is easily separated from 
 its compounds with non-metallic elements by heating them on 
 charcoal. It can be recognized as bright yellow globules. 
 
 Gold can easily be separated from the other metals and recog- 
 nized by its precipitation with ferrous sulphate, or by its insolu- 
 bility in any single acid. 
 
 Mix together solutions containing tin, antimony, and arsenic, 
 dilute the solution and add a considerable quantity of dilute 
 chlorhydric acid. Pass sulphydric acid through the solution 
 for an hour, and set it over night in a warm place. Filter the 
 precipitate, wash and dry it completely. It contains all the 
 metals in the form of sulphides ; separate them according to 
 Part III. (83) and the following tests. 
 
TESTS FOR ACIDS. 
 
 GROUP I. 
 
 ARSENIOUS, ARSENIC, CHROMIC, SULPHURIC, 
 PHOSPHORIC, BORACIC, OXALIC, FLUORHY- 
 DRIC, CARBONIC, AND SILICIC ACIDS. 
 
 Acids which are precipitated from neutral or slightly alkaline 
 solutions by baric chloride. 
 
 SECTION I. 
 
 ARSENIOUS, ARSENIC, AND CHROMIC ACIDS. 
 
 Acids which are precipitated as sulphides or reduced to an oxide 
 by sulphydric acid. 
 
 ARSENIOUS and ARSENIC acids are precipitated as sulphides, 
 and must always be detected by sulphydric acid. 
 
 Arsenic with this reagent plays the part of a metal. (See 
 pages 60 and 61.) 
 
 CHROMIC ACID. 
 
 H 2 CrO 4 ; K 2 Cr 2 O 7 . 
 
 Solutions containing chromic acid have a yellow color. 
 Those which have an acid reaction are of a deeper color, and 
 
 63 
 
64 PART II. 
 
 very concentrated acid solutions are red. Solutions contain- 
 ing chromic acid can usually be recognized by these colors. 
 Chromic acid can easily be reduced to chromic oxide (see 
 page 38) by boiling its solution with chlorhydric acid and 
 alcohol, or by adding to the solution sulphydric acid or am- 
 monic sulphide. Therefore, in all solutions to which these 
 latter reagents have been added, chromic acid is changed to 
 chromic oxide, and must be looked for among the metals. 
 
 JBaric Chloride precipitates chromic acid in neutral 
 solutions, or in solutions of which the only free acid is acetic 
 acid (/. e., solutions which have been made alkaline by sodic 
 hydrate, and acid by acetic acid), as BARIC CHROMATE, 
 BaCrO 4 , light yellow. Chromium may be discovered in the pre- 
 cipitate of baric chromate thus obtained, by dissolving some 
 of it in the borax bead. The green color of chromic oxide 
 appears, even though very little chromate of barium was con- 
 tained in the precipitate. 
 
 No other acid which precipitates baric chloride gives the same 
 color in the borax bead. 
 
 SECTION II. 
 
 SULPHURIC AND SULPHUROUS ACIDS. 
 
 Acids which are precipitated by baric chloride in acid solutions, 
 either immediately or after oxidation. 
 
 ^Blowpipe Reactions. All compounds containing 
 sulphur, when they are pulverized, moistened, and mixed with 
 sodic carbonate, and when the mixture is heated on charcoal, 
 form sodic sulphide. The sodic carbonate must be heated 
 until it soaks into the charcoal, then if the portion of the char- 
 coal which has absorbed it is dug out with a knife, moistened 
 with water, and laid on a piece of bright silver, the presence 
 of sulphur may be detected by the appearance'of a black stain 
 
SULPHURIC AND SULPHURO US A CIDS. 65 
 
 of sulphide of silver. If no silver is at hand, the sodic sul- 
 phide can be extracted by soaking with water the portion of 
 the charcoal which has been heated. The solution, after it has 
 been filtered, and acidified with acetic acid, gives a precipitate 
 of SULPHIDE OF LEAD, PbS, black, with a solution of ACETATE 
 
 OF LEAD. 
 
 For the separation of sulphuric and sulphurous acids, see Part 
 III. (123) and (124). 
 
 SULPHURIC ACID. 
 
 H 2 SO 4 ; Na 2 SO 4 . 
 
 Baric Chloride precipitates compounds of sulphuric 
 acid in neutral or acid solutions as BARIC SULPHATE, BaSO 4 , 
 white powder. Baric sulphate is insoluble in chlorhydric, nitric, 
 sulphuric, and the weaker acids. No other acid gives a like 
 precipitate with BaCl 2 , in acid solutions. COMPOUNDS OF LEAD 
 AND STRONTIUM and STANNIC OXIDE are the only others precipi- 
 tated by sulphuric acid. 
 
 Calcic Chloride in concentrated solutions gives a pre- 
 cipitate of CALCIC SULPHATE, CaSO 4 , but calcic sulphate is 
 soluble in 380 parts of water. 
 
 SULPHUROUS ACID. 
 
 H 2 SO 8 ; Na 2 SO 8 . 
 
 Sulphurous acid, in acid solutions, can be recognized by the 
 smell of sulphurous oxide which is given off. Sulphurous 
 oxide is evolved from acid solutions of sulphites when they 
 are heated. 
 
 JZaric Chloride precipitates compounds of sulphurous 
 acid in neutral solutions as BARIC SULPHITE, BaSO 3 , white 
 powder. Baric sulphite is decomposed and dissolved by chlor- 
 hydric and nitric acids. POTASSIC DICHROMATE oxidizes sul- 
 phurous acid in a solution, made acid with chlorhydric acid, to 
 SULPHURIC ACID, 2(H 2 CrO 4 ) + 3(H 2 SO 8 ) = Cr 2 (SO*) 8 + 5H 2 O. 
 5 
 
66 PART II. 
 
 Sulphydric Acid decomposes sulphurous acid in acid 
 solutions with precipitation of sulphur. 
 
 2 H 2 S + H 2 SO 8 = S 3 + 3H 2 O. 
 
 The Iron Reduction Test. POTASSIC FERRICYAN- 
 IDE AND FERRIC CHLORIDE, in consequence of the reduction 
 of the ferric chloride, give a blue precipitate when these so- 
 lutions are mixed, and a drop of the mixture is held on the end 
 of a glass rod in an atmosphere containing sulphurous acid. 
 Sulphydric acid gives the same reaction, and when sulphydric 
 acid is present it is necessary to add to the solution sufficient 
 plumbic acetate to precipitate it, before applying this test for 
 sulphurous acid. This reaction is produced equally well in a 
 solution ; but there are many other reducing agents which act 
 in the same manner, but which do not take the gaseous form, 
 and for this reason it is best to add chlorhydric acid to the 
 solution, to warm it, and to test the gas evolved for sulphurous 
 acid with potassic ferricyanide and ferric chloride. 
 
 Sulphurous acid is usually recognized by its smell, or ty the 
 above test. 
 
 SECTION III. 
 PHOSPHORIC, BORACIC, OXALIC, AND FLUORHYDRIC ACIDS. 
 
 Acids which are precipitated by baric chloride in neutral solutions , 
 but which are not precipitated in acid solutions. 
 
 For the tests by which these acids can be most easily detected, see 
 Part III. (125-128). 
 
 PHOSPHORIC ACID. 
 
 H 3 PO 4 ; Na 2 HPO 4 . 
 
 All the Metals of Groups II., III., IV., V., pre- 
 cipitate compounds of phosphoric acid in neutral or in slightly 
 alkaline solutions as PHOSPHATES. The precipitate is soluble 
 
BORACIC ACID. 67 
 
 in acids. The phosphoric acid cannot usually be separated 
 from the metal by treating the precipitate with ammonic 
 hydrate, and from most metals it can only be partially separ- 
 ated by a treatment with sodic hydrate or carbonate. MAG- 
 NESIC SULPHATE solution, to which a solution of ammonic 
 chloride and of ammonic hydrate is added, precipitates com- 
 pounds of phosphoric acid as MAGNESIC AMMONIC PHOSPHATE, 
 MgNH 4 PO 4 , white crystalline powder, soluble in acids but inso- 
 luble in ammonic hydrate. When the quantity of phosphoric 
 acid is very small, the precipitate only forms after some time. 
 This is the best reagent for phosphoric acid in neutral or slightly 
 alkaline solutions. 
 
 Molybdate of Ammonium * precipitates compounds 
 of phosphoric acid in an acid solution as a yellow crystalline 
 phospho-molybdate of ammonium. No other acid gives a like 
 precipitate. The best reagent for phosphoric acid in acid solu- 
 tions is molybdate of ammonium, when there are bases present 
 which are precipitated by ammonic hydrate. 
 
 Sulphydric acid and ferrocyanhydric acid also precipitate 
 ammonic molybdate solution. (See Part III., 127.) 
 
 BORACIC ACID. 
 
 H 8 BO 3 ; Na 2 B 4 O 7 , Borax. 
 
 All the Metals of Groups II., III., IV., V. preci- 
 pitate compounds of boracic acid in neutral or slightly alkaline 
 solutions as BORATES. The precipitate is soluble in acids. 
 
 The precipitation of boracic compounds by metals of these 
 groups is mostly incomplete, and the acid can be separated 
 from the metal in almost all cases by treating the precipitate 
 with ammonia or with sodic hydrate or carbonate. 
 
 Flame Test. Boracic acid in dry compounds or in con- 
 
 * It is best to pour a few drops of the solution to be tested into the am- 
 monic molybdate solution, instead of adding the ammonic molybdate to the 
 solution to be tested. 
 
68 PART II. 
 
 centrated solutions imparts a beautiful green color to the flame 
 of burning alcohol. 
 
 In order to perform the test, mix the dry boracic acid com- 
 pound with a few drops of strong sulphuric acid in a small 
 evaporating dish, add alcohol, and set the alcohol on fire. Stir 
 the contents of the dish constantly during the combustion of 
 the alcohol, and observe the green color, when the alcohol has 
 mostly burned away. 
 
 Turmeric Paper Test.* Boracic acid compounds, 
 when their solution has been rendered acid by chlorhydric 
 acid, turn a piece of turmeric paper, which has been dipped in 
 the solution and completely dried, brownish red. This is the 
 usual test for boracic acid. 
 
 OXALIC ACID. 
 H 2 C 2 O 4 ; KHC 2 O 4 - 
 
 All the Metals of Groups II., III. 9 IV., F. pre- 
 cipitate salts of oxalic acid in neutral solutions as OXALATES, 
 except chromic oxide compounds. The precipitate is soluble 
 in acids, and in the case of many metals the oxalic acid is 
 removed from the precipitate by ammonic hydrate and by sodic 
 hydrate and carbonate. 
 
 Sulphate of Calcium precipitates oxalic acid and its 
 salts in a solution to which sufficient ammonic hydrate to 
 render the solution strongly alkaline, and then sufficient acetic 
 acid to render the reaction acid, have been added, as CALCIC 
 OXALATE, CaC 2 O 4 , white powder. No other acid, except fluor- 
 hydric acid, produces a precipitate under these circumstances, and 
 calcic fluoride cannot easily be mistaken for calcic oxalate. (See 
 fluorhydric acid, below.) 
 
 Sulphuric Acid (concentrated) decomposes dry com- 
 pounds, or highly concentrated solutions of compounds of 
 
 * If iron is present it is necessary to boil the solution with sodic carbon- 
 ate in excess, to filter, and to use the filtrate for the turmeric paper test. 
 
FLUORHYDRIC ACID. 
 
 69 
 
 oxalic acid, with evolution of CARBONOUS and CARBONIC OX- 
 IDES. Effervescence takes place. H 2 C 2 O 4 = CO +CO 2 + H 2 O. 
 
 FLUORHYDRIC ACID. 
 
 HF ; CaF 2 ; NH 4 F. 
 
 Fluorhydric acid cannot be present in acid solutions in glass 
 vessels. 
 
 BARIUM, STRONTIUM, and CALCIUM salts in neutral solu- 
 tions precipitate fluorhydric acid as BARIUM, STRONTIUM, or 
 CALCIUM Fluoride, BaF 2 , SrF 2 , CaF 2 . The two former are 
 voluminous white precipitates. Calcium fluoride is a gelatin- 
 ous transparent precipitate, whose formation it is very difficult 
 to observe. 
 
 // is usually unnecessary to test for fluorhydric acid except in 
 solid substances. 
 
 Sulphuric Acid (concentrated) sets fluorhydric acid 
 free from its solid compounds, and the acid may be recog- 
 nized by its property of etching glass. To perform the test, 
 mix the pulverized substance, containing fluorine, with strong 
 sulphuric acid in a lead cup,* or in a platinum crucible. Pre- 
 pare a piece of glass by melting wax on it and pouring off all 
 that does not adhere, leaving a thin coating of wax on the sur- 
 face ; scratch lines in the wax, laying the surface of the glass 
 bare ; cover the vessel containing the fluorine compound and 
 sulphuric acid with the glass, and warm gently. After fifteen 
 minutes warm the glass, and rub off the wax ; the surface ex- 
 posed by the scratches will be etched by the fluorhydric acid. 
 
 Fluorhydric acid must always be removed, if present, by heat- 
 ing with strong sulphuric acid, before the other acids of Group I., 
 except sulphuric acid and carbonic acid, are tested for in a sub- 
 stance. 
 
 * A lead cup may be made by hammering up the end of a piece of one- 
 inch lead pipe until it is entirely closed, and by sawing off the pipe i$ 
 inches from the end. Such a cup may be warmed on a sand-bath suffi- 
 ciently for making the test without danger of melting it. 
 
7 o PART II. 
 
 SECTION IV. 
 CARBONIC ACID AND SILICIC ACID. 
 
 Acids which are precipitated by baric chloride in neutral selu- 
 tion, but which are set free by acids, and which cannot be present 
 in a' solution that has been evaporated with an excess of an acid. 
 
 CARBONIC ACID. 
 
 Na 2 C0 3 ; K 2 CO 3 . 
 
 Carbonic acid can only be present in considerable quantity 
 in a solution which has an alkaline reaction. 
 
 All the Metals of Groups II., III., IV., V. pre- 
 cipitate alkaline carbonates. 
 
 Every Acid sets free CARBONIC DIOXIDE, CO 2 , from a 
 solution of a carbonate. An effervescence or the formation of 
 bubbles can be observed when an acid is added to a solution of 
 a salt of carbonic acid. Carbonates insoluble in water are also 
 decomposed by a free acid, carbonic dioxide being evolved. 
 
 Carbonic dioxide can easily be recognized by the white 
 precipitate of CARBONATE OF CALCIUM, CaCO 3 , which the gas 
 produces in a drop of lime-water,* held in it on the end of a 
 glass rod. 
 
 SILICIC ACID. 
 
 K 2 Si0 3 . 
 
 Silicic acid combined with bases in a solution is set free by 
 all acids, and being decomposed into water and silicic oxide 
 (H 2 SiO 3 = H 2 O + SiO 2 ), the latter often separates as a pre- 
 cipitate ; frequently, however, it remains for months in a solu- 
 tion after an acid has been added. 
 
 * The lime-water is soon destroyed by the absorption of carbonic di- 
 oxide from the air, and before using it to test for carbonic acid, the stu- 
 dent should assure himself that a drop of it gives a precipitate with the 
 gas evolved from sodic carbonate, to which chlorhydric has been added. 
 
SILICIC ACID. 7I 
 
 When silicic oxide, or silica, has been separated from its com- 
 bination with a base by the addition of an acid, and the solution 
 has been evaporated completely to dryness, the silica remains 
 perfectly insoluble when the dry mass is treated with water or 
 acids to dissolve the bases. This is the characteristic test for 
 silicic acid. 
 
 When silicic acid is present in a solution it is recognized 
 and separated by the above process before any other tests are 
 performed. 
 
 Many solid compounds of silicic acid are not acted upon by 
 acids, and can only be brought into solution by the process 
 described, Part III., XXIII. 
 
 The following properties of silicic acid must be considered, 
 in order to determine whether a body, which has been left in- 
 soluble after a treatment with an acid, is silica, or some other 
 compound, which is likewise insoluble : 
 
 Silicic acid is precipitated from its solutions by acids in the 
 gelatinous form, or in the form of amorphous white flakes. 
 Silica which has been dried always has the latter form. 
 
 Silica, after it has been fused with four parts of a mixture 
 of sodic carbonate and potassic carbonate, forms a glass, which 
 is entirely soluble in water. When the solution thus obtained 
 is evaporated with an excess of nitric acid, and the dry mass 
 is treated with water, no metal should go into solution which 
 gives a precipitate with sodic carbonate. 
 
PART 77. 
 
 GROUP II. 
 
 CHLORHYDRIC, BROMHYDRIC, IODHYDRIC, CY- 
 ANHYDRIC, FERROCYANHYDRIC, FERRICYAN- 
 HYDRIC, AND SULPHYDRIC ACIDS. 
 
 Acids which are not precipitated by baric chloride, but which 
 are precipitated by argentic nitrate in nitric acid solution. 
 
 SECTION I. 
 
 CHLORHYDRIC, BROMHYDRIC, IODHYDRIC, AND CYANHYDRIC 
 
 ACIDS. 
 
 Acids which give with argentic nitrate a white or light yellow 
 flocculent precipitate, insoluble in dilute nitric acid, and are not 
 precipitated by salts of iron in acid solution. 
 
 CHLORHYDRIC ACID. 
 
 HC1; NaCl. 
 
 Plumbic, Mercurous, and Argentic Salts are 
 
 the only compounds which give, with chlorhydric acid, pre- 
 cipitates insoluble in nitric acid. 
 
 Sulphuric Acid (concentrated) sets free chlorhydric 
 acid from its compounds. Chlorhydric acid gas precipitates a 
 drop of argentic nitrate, held on the end of a glass rod in an 
 atmosphere containing it, as ARGENTIC CHLORIDE, AgCl, white 
 flakes. Compounds containing cyanhydric, chloric, and hypo- 
 chlorous acids, produce the same reaction. Chlorhydric acid 
 gas does not bleach indigo solution. See, however, Part III. 
 (48). 
 
BROMHYDRIC AND IODHYDRIC ACIDS. 
 
 73 
 
 Argentic Nitrate precipitates solutions containing 
 chlorhydric acid as ARGENTIC CHLORIDE, AgCl, white flakes, 
 turning purple on exposure to light, settling quickly after they 
 have been shaken. Argentic chloride is soluble in ammonic 
 hydrate, and completely insoluble in boiling nitric acid (con- 
 centrated). No other compound of silver except the ferro- and 
 ferricyanide remains undissolved after this treatment. When 
 ferro- or ferricyanhydric acid is present in a solution containing 
 chlorhydric acid, follow the directions given, Part III. (130) 
 (a), before applying the argentic nitrate test. 
 
 This is the usual test for chlorhydric acid. 
 
 BROMHYDRIC ACID. 
 
 KBr. 
 
 Nearly all bromides are soluble in water ; bromide of lead, 
 however, dissolves very sparingly, and the mercurous and 
 argentic salts are quite insoluble. 
 
 Argentic Nitrate precipitates from solutions of bro- 
 mides the ARGENTIC BROMIDE, AgBr, white, closely resembling 
 the chloride, but more difficultly soluble in ammonia ; insolu- 
 ble in hot nitric acid. 
 
 Bromides are decomposed by chlorine, hypochlorites, strong 
 sulphuric and nitric acids, bromine being set free, which im- 
 parts a yellow or yellowish-red color to the liquid. On shak- 
 ing the tube with a little ether or carbonic disulphide the 
 bromine will be dissolved and impart to the solvent a yellow 
 or reddish-brown color. 
 
 IODHYDRIC ACID. 
 
 KI. 
 
 The iodides resemble very much the corresponding chlorides 
 and bromides. 
 Argentic Nitrate produces in solutions of iodides a 
 
74 PART II. 
 
 yellowish-white precipitate of ARGENTIC IODIDE, Agl, very 
 slightly soluble in ammonia, but soluble in boiling concen- 
 trated nitric acid. 
 
 JVLeTCUTOUS NitTate precipitates yellowish-green MER- 
 
 CUROUS IODIDE, Hg 2 I 2 . 
 
 jMLevcwT-ic Chloride precipitates brilliant scarlet MER- 
 CURIC IODIDE, HgI 2 . 
 
 Chlorine or bromine water liberates iodine from iodides ; on 
 shaking with carbonic disulphide the iodine is concentrated in 
 this liquid, forming a violet-colored solution. 
 
 Free iodine imparts a blue color to starch paste. For this 
 purpose dilute starch paste is added to a solution of an iodide, 
 and then chlorine water, nitric acid, or sulphuric acid carefully 
 added to liberate the iodine. If too much chlorine is added 
 chloride of iodine is formed, which prevents the formation of 
 the blue color. The chloride of iodine is also produced by 
 nitric acid in presence of considerable amount of chlorides. 
 
 CYANHYDRIC ACID. 
 
 HCy; KCy. 
 
 The reactions of different classes of cyanhydric acid com- 
 pounds must be considered separately. 
 
 Soluble Simple Cyanides. Cyanides of metals of 
 Groups I., II., and III. are soluble in water. Cyanhydric acid 
 is set free from their solutions by even the feeblest acids (acetic 
 and carbonic). (MERCURIC CYANIDE is soluble in water, but 
 is not decomposed by alkalies nor by acids, except by sul- 
 phydric acid ; and the tests described for cyanhydric acid 
 cannot be applied to it. Sulphydric acid precipitates mercu- 
 ric sulphide, and sets cyanhydric acid free. ) 
 
 Insoluble Simple Cyanides. Cyanides of metals 
 of Groups IV. and V., except mercuric cyanide, are insoluble 
 in water, and the cyanides of metals of Group V. are not de- 
 
CYANHYDRIC ACID. 
 
 75 
 
 composed, or are decomposed with great difficulty by acids. 
 The insoluble cyanides dissolve readily in potassic cyanide, and 
 the ordinary tests for metals cannot be used with such solutions. 
 The cyanides can be precipitated from these solutions by the 
 addition of an acid, with some exceptions, the two most re- 
 markable of which are described separately. (See Ferro- and 
 Ferricyanhydric Acids.) 
 
 Free Cyanhydric Acid can be recognized by its 
 smell, which is like that of bitter almonds. (The acid is very 
 poisonous, and the fumes arising from a solution containing a 
 considerable quantity of it should be inhaled with caution.) 
 
 Argentic Nitrate precipitates soluble compounds of 
 cyanhydric acid in an acid solution as ARGENTIC CYANIDE, 
 AgCy, white flakes, which do not settle so readily, when shaken, 
 as argentic chloride, and which do not turn purple quickly in the 
 light. Argentic cyanide is wholly decomposed and dissolved 
 by boiling a few minutes with strong nitric acid. It is also 
 decomposed, and cyanhydric acid goes into solution, when it is 
 digested with dilute chlorhydric acid in contact with metallic 
 zinc. 
 
 Prussian Blue Test. When a feebly acid solution 
 containing cyanhydric acid is mixed with several drops of 
 ferrous sulphate solution, and with a drop of ferric chloride 
 solution, and sodic hydrate is added until a precipitate forms, 
 and the mixture is warmed for a minute, and then acidified 
 with dilute chlorhydric acid, a blue precipitate, or more frequently 
 a blue coloration, appears, either immediately or after the addi- 
 tion of a drop of ferric chloride. 
 
 When ferro- or ferricyanhydric acid or both acids are present 
 (see the following section), they must be removed from the 
 solution before the Prussian-blue test can be applied. To this 
 end add to a small quantity of the solution an equal bulk of 
 dilute sulphuric acid, and dilute with a considerable quantity 
 of water ; add ferric chloride or ferrous sulphate, or both to- 
 gether, according as ferro- or ferricyanhydric acid or both 
 
76 PART II. 
 
 acids are present, and then add baric chloride until the blue 
 precipitate appears of a much lighter shade ; shake thoroughly, 
 and allow the precipitate to settle for a few minutes, and filter. 
 If only the first few drops run through the filter blue, they 
 should be thrown away and the remainder of the filtrate taken. 
 If no clear filtrate can be obtained, add to the filtrate a little 
 baric chloride and filter again. The filtrate is to be tested as 
 above by the addition of sodic hydrate, and afterwards of an 
 acid for cyanhydric acid. A sufficiently capacious flask must 
 be chosen for the operation. The only object in adding baric 
 chloride is to facilitate the filtration from the blue precipitate. 
 Cyanhydric acid is the only acid which gives this reaction under 
 these circumstances. 
 
 SECTION II. . 
 FERROCYANHYDRIC AND FERRICYANHYDRIC ACIDS. 
 
 Acids which give with argentic nitrate colored precipitates, 
 which are not wholly destroyed on boiling with strong nitric acid, 
 and which are precipitated by ferrous or ferric salts, and by cu- 
 pric salts in dilute acid solutions. 
 
 FERROCYANHYDRIC ACID. 
 
 H 4 (FeCy 6 ) ; K 4 (FeCy 6 ). 
 
 Ferrous Sulphate precipitates ferrocyanhydric acid 
 compounds in acid solutions as POTASSIC FERROUS FERROCY- 
 ANIDE, K 2 Fe(FeCy 6 ), bluish white precipitate, which quickly 
 turns dark blue through oxidation by the air. 
 
 Ferric Chloride precipitates ferrocyanhydric acid com- 
 pounds in acid solution as PRUSSIAN BLUE, Fe 4 (FeCy 6 ) 3 , deep 
 blue. 
 
 Ferrocyanhydric acid is the only acid which gives this reaction. 
 
 Cupric Sulphate precipitates ferrocyanhydric acid com- 
 pounds in acid solution as CUPRIC FERROCYANIDE (Cu 2 FeCy 6 ), 
 brownish-red powder. 
 
SULPH YDRIC A CID. j j 
 
 The metals are left as oxides, and the ferrocyanogen is dis- 
 solved as sodic ferrocyanide, when these precipitates are di- 
 gested with sodic hydrate. 
 
 FERRICYANHYDRIC ACID. 
 
 H 6 (Fe 2 Cy 12 ) ; K 6 (Fe 2 Cy 12 ). 
 
 Ferrous Sulphate precipitates compounds of ferricy- 
 anhydric acid in acid solution as TURNBULL'S BLUE, Fe 3 (Fe 2 
 Cy M ), deep blue. 
 
 Ferricyanhydric acid is the only acid which gives this reaction. 
 
 Ferric Chloride does not precipitate compounds of fer- 
 ricyanhydric acid in acid solution. The color of the solution 
 is deepened. 
 
 Cupric Sulphate precipitates compounds of ferricyan- 
 hydric acid in acid solution as CUPRIC FERRICYANIDE, yellow- 
 ish-green powder. 
 
 The metals are left as oxides, and the cyanogen is dissolved 
 as sodic ferricyanide when these precipitates are treated with 
 sodic hydrate. 
 
 SECTION III. 
 
 SULPHYDRIC ACID. 
 
 An acid which gives a black precipitate with salts of lead, silver, 
 copper, and many others in an acid solution. 
 
 No other acid gives a precipitate of the same color with these 
 metals. 
 
 SULPHYDRIC ACID. 
 
 H 2 S ; (NH 4 ) 2 S. 
 
 Sulphydric Acid is set free from its solutions by all 
 other acids except carbonic and cyanhydric acid, and it can 
 be recognized by its smell. The sulphydric acid gas is given 
 off with effervescence when the solution is concentrated. 
 
7 8 PART II. 
 
 The Metals of Groups I. and II. form with 
 sulphydric acid soluble sulphides, which have an alkaline re- 
 action. 
 
 The Metals of Group IV. 9 when the acid with which 
 they are combined is neutralized, form with sulphydric acid 
 INSOLUBLE SULPHIDES, which, with the exception of the sul- 
 phides of cobalt and nickel, are dissolved by cold dilute chlor- 
 hydric acid, with evolution of sulphydric acid. 
 
 Metals of Groups V. and VI. form with sulphydric 
 acid insoluble sulphides, which are not decomposed by dilute 
 acids. (See also Mercury, page 48 and page 50.) 
 
 Lead-Paper Test. A piece of paper moistened with 
 plumbic acetate, and held over a solution from which sulphy- 
 dric acid is set free by the addition of a stronger acid, is black- 
 ened. No other acid gives this reaction. 
 
NITRIC AND CHLORIC ACIDS. 
 
 79 
 
 GROUP III. 
 
 NITRIC, CHLORIC, AND ACETIC ACIDS. 
 
 Acids which are not precipitated by any metal. 
 
 SECTION I. 
 
 NITRIC AND CHLORIC ACIDS. 
 Acids which deflagrate when tested with the blowpipe on charcoal. 
 
 NITRIC ACID. 
 
 HNO 8 ; NaNOs. 
 
 Nitric Acid 9 when concentrated, is readily decomposed 
 when heated with copper turnings, and red fumes of NITRIC 
 PEROXIDE, NO 2 , are given off. The reaction can be obtained 
 with a moderately dilute solution by adding to it concentrated 
 sulphuric acid. No reaction is obtained with very dilute so- 
 lutions. 
 
 J?errous Sulphate Test. Add a few drops of a solu- 
 tion containing nitric acid to concentrated sulphuric acid in a 
 test-tube, and pour upon this solution a layer of cold ferrous 
 sulphate solution. A brown or red color appears at the line of 
 separation of the two solutions, arising from the absorption of 
 nitrous gases by the ferrous sulphate. 
 
 This is the characteristic test for nitric acid. 
 
 CHLORIC ACID. 
 KC10 8 . 
 
 Sulphuric Acid (concentrated). When a small quan- 
 
80 PART II. 
 
 tity of a solid chlorate, or a very concentrated solution con- 
 taining a chloric acid compound, is added to strong sulphuric 
 acid, and heat is applied, a peculiar yellow gas (oxides of 
 chlorine) is evolved, which has a characteristic suffocating 
 odor, which precipitates ARGENTIC CHLORIDE in a drop of an 
 argentic nitrate solution, and which bleaches a drop of an in- 
 digo solution when these reagents are held on the end of a 
 glass rod in an atmosphere containing the gas. 
 
 This is the characteristic test for chloric acid. 
 
 Hypochlorous Acid, gives the same reactions as chloric 
 acid, but that acid is easily set free and evolved from its solu- 
 tion by dilute sulphuric acid, while chloric acid is not, and 
 moreover it is usually present only in alkaline solutions. 
 
 Chlorhydric Acid in the presence of an oxidizing 
 agent gives a similar reaction, but the yellow gas evolved 
 (chlorine) is much less intense in color, and has a different 
 odor. It is, however, very difficult to distinguish between the 
 reaction given by chlorine in such a case and that given by 
 chloric acid compounds. 
 
 SECTION II. 
 
 ACETIC ACID. 
 
 An acid which does not deflagrate on charcoal. 
 
 ACETIC ACID. 
 
 HC 2 H 8 2 ; NaC 2 H 8 2 . 
 
 The Strong Mineral Acids set acetic acid free 
 from its combinations. 
 
 Acetic acid can be recognized by the odor of vinegar pecu- 
 liar to it. 
 
 Sulphuric Acid Test. When an equal bulk of alco- 
 hol is added to strong sulphuric acid, and a small quantity of 
 a solution containing a compound of acetic acid is added, and 
 
ACETIC ACID. 8 1 
 
 the mixture is heated, a characteristic odor of acetic ether is 
 given off. 
 
 In case gases are given off, which make it difficult to recog- 
 nize the odor of acetic ether, it is advisable to provide the test- 
 tube in which the reaction is performed with a tube for distil- 
 lation,* and to distil a small quantity of the alcohol into an- 
 other test-tube, to mix the distillate with water, to neutralize it 
 with sodic carbonate, and to warm it ; the odor of acetic ether 
 can then be recognized in the liquid which was distilled. 
 
 This is the characteristic test for acetic acid. 
 
 Argentic and Mercurous Nitrates precipitate 
 concentrated neutral solutions of acetic acid compounds as 
 
 ARGENTIC AND MERCUROUS ACETATES, AgC 2 H 3 O 2 and Hg 2 
 
 (C 2 H 3 O 2 )2, white crystalline scales. The precipitates are soluble 
 in dilute nitric acid, and also in a large quantity of water. 
 
 * Bend a 3-16 inch tube, of about one foot in length, at an angle of 
 about 80, so that one arm shall only be i^ inches long. Fit a cork to 
 the test-tube, and insert the bent tube in a hole bored through the cork 
 with a round file, 
 
 6 
 
PART III. 
 
 PRELIMINARY TESTS WITH NON-METALLIC 
 SOLIDS. 
 
 EXAMINATION IN A CLOSED TUBE. 
 
 USE a piece of hard glass tubing three-eighths of an inch 
 in diameter, closed at one end (see page 26) for this examina- 
 tion. Introduce the substance, pulverized or in small pieces, 
 into the tube, wipe the inside of the tube if necessary with a 
 bit of rolled filter-paper, and heat the substance, gently at first, 
 but eventually to the highest temperature attainable with the 
 flame of a Bunsen's lamp or with the blowpipe flame. Ob- 
 serve carefully the changes which occur. 
 No Change. The substance contains no organic matter, 
 
 (1) no readily fusible body, no readily volatile body, and 
 no water. 
 
 Pass to the Examination on Charcoal (page 85). 
 Wat er* Substances containing water (usually water of crys- 
 
 (2) tallization) deposit a film of moisture in the upper 
 part of the tube when they are heated. If the water 
 colors turmeric paper brown, AMMONIA is present. 
 
 Organic Matter* Substances containing organic matter 
 
 (3) blacken and give off gases when they are heated. 
 
 Should the substance contain organic matter it must 
 be burnt, until the organic matter is completely de- 
 stroyed,* by heating with the lamp or blowpipe, on 
 
 * In some special cases, as in* examinations for mercury and arsenic, 
 other processes of analysis must be employed, for which larger works must 
 be consulted. 
 
 82 
 
EXAMINA TION OF SOLIDS. 83 
 
 platinum foil or on a bit of porcelain, or in a porcelain 
 dish or crucible, before further analysis, commencing 
 with the examination on charcoal (page 85), is pro- 
 ceeded with. 
 
 A GAS IS GIVEN OFF. 
 
 May be recognized by its property of rekindling 
 (4) a glimmering match held in the tube. 
 
 PEROXIDES, NITRATES, and CHLORATES evolve oxy- 
 gen. 
 
 Nitrates and chlorates also deflagrate on charcoal. 
 See (14). 
 
 Sulphurous Oxide, SO 2 , can be recognized by its smell. 
 (3) Some SULPHATES of higher metals, and many 
 SULPHITES, evolve sulphurous oxide when they are 
 heated.* 
 
 SulpTiydric Acid, H 2 S, can be recognized by its smell 
 (0) and by its property of blackening lead paper. See 
 (45). 
 
 Some alkaline SULPHIDES, containing water, evolve 
 sulphydric acid when they are heated. 
 Carbonic Dioxide, CO 2 , can be recognized by its prop- 
 
 (7) erty of extinguishing a lighted or glimmering match 
 held in the tube. See also (41). 
 
 Some CARBONATES lose carbonic dioxide when they 
 are heated. 
 Hyponitric Oxide, NO 2 , appears as red fumes. 
 
 (8) NITRATES of the higher metals evolve hyponitric 
 oxide when they are heated. 
 
 Ammonia, NH 3 , can be recognized by its smell, and by its 
 
 * Many sulphides of higher metals give off sulphur in the form of sul- 
 phurous oxide when they are roasted with . access of air. The sulphides, 
 finely pulverized, may be heated red-hot in a tube, open at both ends, and 
 held in an inclined position to favor the draught, and the sulphurous oxide 
 may be detected by its smell at the upper end of the tube. 
 
84 PART III. 
 
 (9) property of turning moist turmeric paper brmvn. Salts 
 of ammonia, in the presence of alkalies, and some 
 organic substances, evolve ammonia when they are 
 heated. 
 
 [Better tests for these bodies, with the exception of oxygen, 
 are given in the following pages, since it is often difficult to 
 observe the formation of a gas in a small tube ; the phenomena 
 described above should, however, be looked for when sub- 
 stances are heated in a closed tube.] 
 
 A SUBLIMATE FORMS. 
 
 An opinion may be formed of the volatility of the sublimate, 
 according to the distance from the heated part of the tube at 
 which it is deposited. 
 Sulphur sublimes easily and solidifies in reddish-brown 
 
 (10) drops, which become yellow or yellmrish-brown on 
 cooling. 
 
 Some METALLIC SULPHIDES give off a portion of 
 their sulphur when they are heated. 
 Ammonic Salts form white sublimates. Touch the sub- 
 
 (11) limate with a drop of sodic hydrate, or with a bit of 
 paper moistened with sodic hydrate, and if the smell 
 of ammonia is given off it consists of an ammonic salt. 
 
 Mercury. Metallic mercury sublimes as & grey film, which 
 
 (12) augments to form globules when the quantity of mer- 
 cury is large. 
 
 MERCURIC SULPHIDE, HgS, gives a black sublimate, 
 which becomes ra/when it is rubbed. 
 
 MERCUROUS CHLORIDE, Hg 2 Cl 2 , and MERCURIC 
 CHLORIDE, HgCl 2 , give a white sublimate, which turns 
 black when it is moistened with ammonic sulphide so- 
 lution. 
 Arsenic. Metallic arsenic sublimes and deposits itself as a 
 
 (13) brilliant black metallic ring in the tube. 
 
EXAMINA TION ON CHARCOAL. 85 
 
 ARSENIOUS OXIDE, As 2 O 8 , forms a white crystalline 
 sublimate, which turns yellow when it is moistened 
 with sulphydric acid solution. 
 
 ARSENIOUS SULPHIDE, As 2 S 3 , forms a sublimate, 
 which is reddish yellow when hot, and yelloiv when 
 cold. It is somewhat less volatile than sulphur. 
 
 RECAPITULATION (10) TO (13) SUBLIMATES. 
 
 The substance is heated in a closed tube. 
 
 WHITE SUBLIMATE ammonic salts (11); mercurous chloride, 
 
 Hg 2 Cl 2 , and mercuric chloride, HgCl 2 (12) ; and arsenious 
 
 oxide, As 2 O 3 (13). 
 YELLOW SUBLIMATE sulphur (10) ; and arsenious sulphide, 
 
 As 2 S 3 (13). 
 
 BROWN SUBLIMATE (while hot) sulphur (10). 
 REDDISH-YELLOW SUBLIMATE (while hot) arsenious sulphide, 
 
 As 2 S 3 (13). 
 
 GRAY METALLIC SUBLIMATE mercury (12). 
 BLACK SUBLIMATE arsenic (13) ; and mercuric sulphide, 
 
 HgS, red when rubbed (12). 
 
 EXAMINATION ON CHARCOAL. 
 
 Hollow out a small cavity in a piece of charcoal (see page 
 25), and heat a portion of the solid substance with the blow- 
 pipe flame. 
 titrates and Chlorates enter into a vivid combus- 
 
 tion, called deflagration, when they are heated on 
 
 charcoal. 
 
86 PART III. 
 
 Potassium and Sodium Salts melt, and some of 
 (15) them are imbibed by the pores of the charcoal when 
 
 they are heated. 
 
 Compounds of the Metals of Groups II. and 
 (16) III., also Zinc Compounds and Silicic 
 Oxide, remain as a white infusible mass on the char- 
 coal after heating. Frequently, when heat is first ap- 
 plied, they melt in their water of crystallization, and 
 afterwards become solid. 
 
 ALUMINIC OXIDE becomes blue, and ZINC OXIDE be- 
 comes green when they are moistened with cobaltic 
 nitrate and heated in the oxidizing flame. 
 Salts of the Metals of Groups IV. and V. leave a 
 
 (17) dark-colored residue when they are heated on char- 
 coal. The oxides of these metals generally assume a 
 darker color when they are heated. Exceptions : Zinc 
 and mercury. 
 
 Salts of Ammonia and Mercury, also Com- 
 
 (18) pounds of Arsenic and Antimony, which 
 do not contain another metal, volatilize completely 
 when they are heated on charcoal. 
 
 Gold and Silver Compounds, also Oxides of 
 
 (19) Lead and Bismuth, give bright metallic globules 
 when they are heated on charcoal. 
 
 FUSION WITH SODIC CARBONATE. 
 
 (20) When metals of Groups IV., V., and VI. appear to 
 be present (see 17), mix a small quantity of the pul- 
 verized substance with two or three times its bulk of 
 sodic carbonate in the palm of the hand, moisten with 
 water, and form the mixture by working it with a knife- 
 blade into a ball the size of a pea. Place the ball in 
 a cavity scooped out of a piece of charcoal, and heat 
 
EX AM IN A TION ON CHARCOAL. gy 
 
 with the inner blowpipe flame until almost all of the 
 carbonate of soda has been imbibed by the charcoal. 
 Many metals are reduced and appear as metallic 
 globules in the cavity of the charcoal, and those which 
 are volatile deposit an incrustation of their oxides on 
 the charcoal. This incrustation is to be looked for at 
 a greater or less distance from the cavity, according 
 to the volatility of the metal, and always in the direc- 
 tion in which the metallic vapors are blown by the 
 flame. 
 
 The physical and chemical properties of the glob- 
 ules and the color of the incrustations afford means of 
 recognizing several metals, usually, however, only 
 when they are not associated with others. 
 Tron, Cobalt, Nickel, and Manganese Com- 
 
 (21) pounds give neither globule nor incrustation. 
 Gold, Silver, and Copper Compounds give mallea- 
 
 (22) ble globules, which can be distinguished by the re- 
 spective colors of the metals. They give no incrusta- 
 tion. 
 
 Zinc Compounds give no globules, but a white incrusta- 
 
 (23) tion, ZnO, near the spot heated. The incrustation is 
 yellow while hot. It is not volatile in the oxidizing 
 flame. It becomes green when it is moistened with 
 nitrate of cobalt, and heated in the oxidizing flame. 
 
 Tin Compounds give very ductile white globules. 
 
 (24) The incrustation, SnO 2 , produced by tin compounds 
 is dirty yellow when hot, and lighter when cold. It is 
 deposited in the immediate vicinity of the cavity, and 
 it is very difficult to distinguish it from the ash of the 
 charcoal. 
 
 Lead Compounds give very ductile globules. 
 
 (25) The incrustation, PbO, is bright yellow when hot, 
 and pale yellow when cold. It is deposited at a greater 
 distance than SnO 2 from the cavity. When the blow- 
 
88 PART III. 
 
 pipe flame is directed upon the incrustation of PbO 
 it vanishes, and the flame is colored blue. 
 Bismuth Compounds give brittle globules. The incrus- 
 
 (26) tation, Bi 2 O 3 , is orange yellow when hot, and bright 
 yellow when cold. It vanishes when the blowpipe 
 flame is directed upon it, but it does not impart a blue 
 color to the flame. 
 
 Cadmium, Compounds give no globules, but a yellow to 
 
 (27) reddish-brown incrustation, very different in color from 
 that of any other metal. 
 
 Antimony Compounds give brittle globules, but metallic 
 
 (28) antimony is so volatile that frequently these are driven 
 off by the heat required for their reduction. Some- 
 times fumes, arising from* the vapor of antimony, are 
 visible. The incrustation, Sb 2 O 3 , is white. It is de- 
 posited at a greater distance than PbO from the cavity, 
 and it can easily be driven from one place to another 
 on the charcoal by the heat of the blowpipe flame. 
 
 Arsenic Compounds give no globules, but a character- 
 (20) istic garlic odor. The incrustation, As 2 O 3 , is white, 
 
 and it is still more volatile than Sb 2 O 3 . 
 When the compound contains several metals that can be re- 
 duced, they alloy with each other, and it is usually impossible 
 to recognize the metals in the presence of each other by their 
 physical properties ; also, the incrustation given by one metal 
 frequently obscures that given by another. 
 
 If a sufficient quantity of the metal can be easily reduced, 
 
 it is always advisable to treat it with solvents in the manner to 
 
 be described under metals. (See page 98.) 
 
 Sulphur. The following modification of the fusion with 
 
 (30) carbonate of sodium on charcoal is a valuable test to 
 
 discover sulphur in baric sulphate and in sulphides. 
 
 If a piece of the charcoal which has imbibed the 
 
 soda is moistened and laid on a silver coin, a black 
 
 stain appears, if the coin is washed after a few minutes, 
 
EXAMINA TION ON CHARCOAL. 89 
 
 in case sulphur is present. The charcoal may also be 
 pulverized and treated with water, and if the solution, 
 after being filtered, gives a black precipitate with 
 plumbic acetate solution, sulphur is present. For this 
 test the flame of a candle, oil lamp, or alcohol lamp 
 must be used, else the sulphur in the burning gas will 
 vitiate the results. 
 
 RECAPITULATION OF THE EXAMINATION ON 
 CHARCOAL. 
 
 The substance is heated on charcoal. 
 
 DEFLAGRATION. Nitrates and chlorates (14). 
 
 FUSION. Potassium and sodium salts (15). 
 
 WHITE INFUSIBLE RESIDUE. Compounds of metals of Groups 
 II. and III., zinc salts and silicic acid (10). 
 
 DARK-COLORED RESIDUE. Compounds of metals of Groups 
 IV. and V., except zinc and mercury (17). 
 
 COMPLETE VOLATILIZATION. Ammonic and mercuric com- 
 pounds, and compounds of arsenic and antimony, which 
 contain no other metal (18). 
 
 BRIGHT METALLIC GLOBULES. Silver and gold compounds, 
 and the oxides of lead and bismuth (19). 
 
 The substance is mixed with sodic carbonate and heated on 
 charcoal. 
 
 METALLIC GLOBULES WITHOUT INCRUSTATION. Gold, silver, 
 and copper (22) ; tin (24). 
 
 METALLIC GLOBULES AND INCRUSTATION. Lead (25) ; bis- 
 muth (26) ; antimony (28). 
 
9 o PART III. 
 
 INCRUSTATION WITHOUT GLOBULES. Zinc (23) ; Cadmium 
 
 (27). GARLIC ODOR. Arsenic (29). 
 FORMATION OF SODIC SULPHIDE. All compounds containing 
 
 sulphur (30). 
 
 PRELIMINARY TESTS WITH METALLIC BODIES. 
 
 EXAMINATION IN A CLOSED TUBE. 
 See page 26. 
 
 Mercury. Amalgams containing mercury give a subli- 
 
 (31) mate of METALLIC MERCURY when they are heated. 
 At first a gray film forms in the upper part of the tube, 
 and, when the amount of mercury is considerable, fine 
 globules of metallic mercury are formed, which ag- 
 glomerate and become more distinctly visible when 
 they are rubbed with a copper wire. 
 
 Arsenic* Some metallic compounds, containing ARSENIC, 
 
 (32) give a metallic mirror or ring in the upper part of the 
 tube when they are heated. 
 
 EXAMINATION ON CHARCOAL. 
 
 Heat a piece, one-fourth as large as a pea, of the 
 (33) metallic substance in a cavity on a piece of charcoal. 
 See Part I., page 25. The phenomena to be ob- 
 served are the formation of the incrustations de- 
 scribed (pages 87 and 88), the smell of ARSENIC, and 
 the vapors of MERCURY and ANTIMONY. 
 
 COPPER colors the blowpipe flame green, or in the 
 presence of chlorine blue. 
 
SULPHURIC ACID TEST. 9 ! 
 
 PRELIMINARY TESTS WITH NON-METALLIC 
 SOLIDS (continued}. 
 
 TESTS WITH THE BORAX BEAD. 
 
 If the substance to be tested appears to be the oxide, or an 
 oxygen-salt of a higher metal (see 17), dissolve some of it in 
 the borax bead. 
 Cobalt colors the bead blue in the oxidizing and in the re- 
 
 (34) ducing flame. 
 
 Copper colors the bead green when hot, and blue when cold, 
 
 (35) in the oxidizing flame. It colors the bead red, when 
 cold, in the reducing flame. 
 
 Chromium colors the bead green in both flames. 
 
 (36) 
 Iron* colors the bead brownish red when hot, and yellow when 
 
 (37) cold, in the oxidizing flame. 
 
 Nickel colors the bead violet when hot, and/0/<? brown when 
 
 (38) cold, in the oxidizing flame. The color disappears in 
 a good reducing flame. (See page 44.) 
 
 IVLaifigd'nese colors the bead amethyst in the oxidizing flame. 
 
 (39) The color disappears in a good reducing flame. 
 
 (40) THE OXYGEN COMBINATIONS OF THE REMAINING 
 METALS color the bead very slightly or not at all. 
 
 CONCENTRATED SULPHURIC ACID TEST. 
 
 CONCENTRATED SULPHURIC ACID, with the aid of heat, sets 
 free other acids from most of their combinations with metals, 
 and frequently in such a form that they can be recognized by 
 simple tests. This reaction is not, of course, a method of 
 separation ; one acid may obscure the test for another, and 
 the possible cases are so complicated that it would be useless 
 to attempt to describe them all ; therefore, if the result of the 
 
9 2 
 
 PART III. 
 
 sulphuric acid test appears doubtful, it is best to reserve judg- 
 ment of its value until after the tests for acids in solution have 
 been applied. 
 
 The reactions of acids or their salts, when a small quantity 
 of a solid substance or of a very concentrated solution is added 
 to a few cubic centimetres of strong sulphuric acid in a test- 
 tube are described below. Heat should be applied after the 
 reaction, which takes place at the ordinary temperature, has 
 been observed. 
 Carbonic Acid. Effervescence. Carbonic dioxide gas 
 
 (41) renders turbid a drop of lime-water held on the end 
 of a glass rod in the test-tube (see page 70, foot-note). 
 Carbonic acid is also detected by the chlorhydric acid 
 test, and in many cases that test is preferable to the 
 one with sulphuric acid, since oxalic acid does not 
 give the same reaction with chlorhydric acid. See 
 
 (73). 
 
 Oxalic Acid. When a dry compound of oxalic acid is 
 
 (42) added to strong sulphuric acid, and the mixture 
 is heated, carbonous oxide and carbonic oxide are 
 evolved. 
 
 The sulphate of calcium test for oxalic acid (125) 
 is more accurate than that with sulphuric acid. 
 Cyarihydric, Ferro- and Ferricyanhydric 
 
 (43) Acids. Compounds of these acids evolve, when 
 perfectly dry, carbonic oxide with effervescence when 
 they are heated with strong sulphuric acid. Usually, 
 however, a faint odor of cyanhydric acid can be de- 
 tected. The special tests are more valuable. See 
 (75), and (133), (134), and (135). 
 
 Fluorhydric Acid. When a solid substance or a con- 
 
 (44) centrated solution containing fluorine is heated with 
 strong sulphuric acid, fluorhydric acid is evolved, which 
 etches glass (see page 69) ; when silicic acid or a sili- 
 cate is present, fumes of fluoride of silicon, which give 
 
SULPHURIC ACID TEST. 93 
 
 a precipitate of silica in a drop of water held over 
 them on the end of a glass rod, are evolved. 
 
 If fluorhydric acid is discovered, the substance 
 which is to be used for page 104, III., and the follow- 
 ing tests, must be heated with sulphuric acid in a pla- 
 tinum vessel until the fluorhydric acid is driven off 
 completely. 
 
 A separate portion can be used for testing for sul- 
 phuric acid. See (30). 
 
 Sulphydric Acid. Compounds containing this acid 
 (4#) evolve it (often with effervescence) when strong sul- 
 phuric acid is added to them. Sulphydric acid may be 
 recognized by its smell and by its property of black- 
 ening paper dipped in plumbic acetate solution. See 
 (74) and (129). 
 
 Sulphurous Acid. Compounds containing this acid 
 
 (40) evolve sulphurous oxide, SO 2 , with effervescence when 
 
 strong sulphuric acid is added to them. Sulphurous 
 
 oxide, when free from sulphydric acid, and from some 
 
 others, can be recognized by its smell. See (7#). 
 
 Chloric and Hypochlorous Acids. (See page 80.) 
 
 (47) Compounds containing these acids evolve a. yellow gas 
 on the addition of strong sulphuric acid, even when 
 the mixture is not heated. The gas can best be recog- 
 nized by its color, its odor, and its strong bleaching 
 action on a drop of indigo solution held in the tube 
 on the end of a glass rod. This gas precipitates nitrate 
 of silver. 
 
 Chlorhydric Acid. Compounds containing chlorhydric 
 
 (48) acid evolve the gas, HC1, frequently with efferves- 
 cence, on the addition of sulphuric acid. The gas 
 does not bleach indigo solution, but precipitates ni- 
 trate of silver held on the end of a glass rod in the 
 tube. Chlorhydric acid, in the presence of an oxidiz- 
 ing agent, evolves chlorine under the same circum- 
 
94 PART in. 
 
 stances. The gas produces the same reactions as are 
 produced by the gas evolved by chloric and hypo- 
 chlorous acids, but it can be distinguished from them 
 by its smell and by its color, which is a less intense 
 yellow. For special test, see (130). 
 Hromhydric Acid. Bromides are decomposed by strong 
 
 (49) sulphuric acid with evolution of bromhydric acid, 
 which, if the sulphuric acid is concentrated and in ex- 
 cess, is partly decomposed, with separation of bromine 
 and formation of sulphurous oxide. (See also page 73.) 
 
 lodhydric Acid. All the iodides are decomposed by 
 
 (50) strong sulphuric acid on the application of heat. 
 Iodine is set free, which escapes in violet vapors and 
 imparts a blue color to paper moistened with starch. 
 (See also page 73.) 
 
 Nitric Acid. Compounds containing nitric acid in con- 
 
 (51) siderable quantity produce reddish fumes when heated 
 with sulphuric acid in the presence of copper turnings 
 or of any other reducing agent. The following test 
 for nitric acid is more delicate : Mix a little of the 
 powdered substance or solution with strong sulphuric 
 acid, and pour cautiously upon the acid a solution of 
 ferrous sulphate. A brown or red color at the line of 
 separation of the two solutions indicates the presence 
 of NITRIC ACID, HNO 8 . See (136) (a) for this test 
 in the presence of ferro- or ferricyanhydric acid. 
 
 Acetic Acid gives with sulphuric acid an odor of vine- 
 gar. The following test is more delicate : Add an 
 equal volume of alcohol to strong sulphuric acid, and 
 then add the solid substance or concentrated solution 
 supposed to contain ACETIC ACID, C 2 H 4 O 2 . If this 
 acid is present, the odor of acetic ether will be per- 
 ceptible on heating the mixture. It is well to add 
 pure acetic acid at the same time to a mixture of sul- 
 phuric acid and alcohol, in order to compare the odor 
 
SOL UTION OF NON-ME TALLIC BODIES. 95 
 
 produced with that observed in the test. If other 
 gases render the odor of acetic ether difficult to per- 
 ceive, the precautions described on page 81 must be 
 observed. 
 
 RECAPITULATION OF SULPHURIC ACID TEST. 
 
 A colorless gas is given off. 
 
 THE GAS is WITHOUT ODOR. Carbonic acid (4:1) ; oxalic 
 
 acid (42). 
 A PUNGENT SUFFOCATING ODOR. Fluorhydric acid (44) ; 
 
 chlorhydric acid (48) ; bromhydric acid (40). 
 AN ODOR OF BITTER ALMONDS. Cyanhydric, ferrocyanhydric, 
 
 and ferricyanhydric acids (43). 
 AN ODOR OF ROTTEN EGGS. Sulphydric acid (43). 
 AN ODOR OF BURNING SULPHUR. Sulphurous acid (4:6). 
 AN ODOR OF VINEGAR. Acetic acid (52). 
 
 A colored gas is given off. 
 * The gas has also a peculiar suffocating odor. 
 
 YELLOW GAS. Chloric and hypochlorous acids (4:7). 
 FAINT YELLOW GAS. Chlorine (48). 
 RED FUMES. Nitric acid with copper (SI). 
 VIOLET FUMES. lodhydric acid (50). 
 
 METHODS OF DISSOLVING NON-METALLIC 
 BODIES. 
 
 Substances to be tested may be divided into three classes. 
 
 1st Class : Bodies soluble in water. 
 
 To ascertain whether a body is entirely soluble, 
 take a few grains of the substance, which, when it dis- 
 solves with difficulty, must be in as finely divided a 
 
96 PART III* 
 
 condition as possible, and digest them with a con- 
 siderable quantity of water in a test-tube. If the sub- 
 stance does not dissolve completely, boil it for a few 
 minutes with water. If it still does not dissolve com- 
 pletely, filter and evaporate a few drops of the filtrate 
 on platinum foil, in order to see whether anything has 
 dissolved. 
 
 If the substance is partly soluble in water, treat a 
 considerable quantity in the manner directed above, 
 repeat the boiling with water once or twice, wash 
 thoroughly with water, and filter, taking care that as 
 little as possible of the solid substance goes upon the fil- 
 ter, and use the insoluble portion for the following test. 
 
 A separate analysis should usually be made of the 
 part of a substance which is soluble, and of that which 
 is insoluble in water. 
 
 If the substance is wholly insoluble in water, it may 
 be used immediately for the following test. It should 
 usually be finely pulverized. 
 
 2d Class : Bodies insoluble in water, but soluble in chlor- 
 (54:) hydric acid or in nitric acid or in aqua regia. Digest 
 a few grains, or a very small quantity of the finely 
 powdered substance with dilute chlorhydric acid. If 
 it does not entirely dissolve, boil it with the acid. If 
 it still does not entirely dissolve, pour off the liquid, 
 and boil the insoluble substance with strong chlorhy- 
 dric acid. If the substance is not wholly soluble in 
 chlorhydric acid, repeat the trial with nitric acid in 
 the same way, using a fresh portion of the substance. 
 If the substance is still insoluble, use a mixture of both 
 .acids (aqua regia). 
 
 If sulphur, which may be recognized by its color and 
 its low specific gravity, or silicic acid, which may be re- 
 cognized by its peculiar gelatinous aspect, separate out on 
 the addition of chlorhydric or of nitric acids, the sub- 
 
SOLUTION OF NON-METALLIC BODIES. 97 
 
 stance must be considered as soluble, and after filtration 
 (in case of silicic acid, see 64) the solution must be tested 
 according to page 104 (///.), or pag'e 107 (VI r .), accord- 
 ing as nitric or chlorhydric acid has been the solvent. 
 
 If the substance is only partially soluble after treat- 
 ment as above, the insoluble portion must be washed 
 carefully and separated by filtration from the soluble 
 portion, and it must be subjected to the tests for the 
 3d Class. See page 137. 
 
 In all cases before proceeding to test the solution 
 obtained by treatment with acids, it must be so pre- 
 pared that it will be moderately acid with chlorhydric 
 acid, and will not contain any free nitric acid, or great 
 excess of strong acid of any kind. It the solution was 
 made by the use of dilute chlorhydric acid it is already 
 in this condition. If nitric acid was used for the solu- 
 tion the excess must be replaced by chlorhydric acid. 
 In this case add a quantity of the latter about half 
 as great as that of the nitric acid used ; evaporate the 
 liquid almost to dryness in the hood, and add from 25 
 to 50 c.c. of water to the residue. If the substance 
 was dissolved in concentrated chlorhydric acid, the 
 solution should be evaporated down in the same 
 manner to expel the excess of that acid, and the resi- 
 due diluted considerably with water. If the addition 
 of chlorhydric acid should cause a precipitate in the 
 clear acid solution, or if the residue, after evaporation, 
 does not dissolve entirely in water, it can only be occa- 
 sioned by the presence of argentic, plumbic, or mercu- 
 rous chloride. In this case, after the addition of water, 
 filter off the insoluble powder and examine for those 
 metals as directed (OS), (67), (69). Use the fil- 
 trate for the sulphydric acid test (76). 
 
 3(1 Class ; Bodies insoluble in water, and in chlorhydric, 
 nitric, and nitro-chlorhydric acids (aqua regia). 
 7 
 
98 PART III. 
 
 The tests to be applied to bodies of this class follow 
 those in the scheme for testing bodies in solution. See 
 Page 137. 
 
 METHODS OF DISSOLVING METALLIC BODIES. 
 Metals are divided into three classes. 
 
 1st Class : Metals which are not attacked by NITRIC ACID. 
 (50) If the metal does not appear to be entirely soluble 
 in dilute nitric acid, even after boiling, use strong ni- 
 tric acid. 
 
 GOLD is insoluble, also alloys containing a very large 
 proportion of gold. 
 
 2d Class : Metals which are attacked by nitric acid, and 
 (57) converted into oxides (white powder], which are inso- 
 luble in the acid. 
 
 TIN and ANTIMONY belong to Class II. 
 
 3d Class : Metals which dissolve entirely in nitric acid. 
 (58) All the remaining metals dissolve entirely in nitric 
 acid. With metals of this class the solution should be 
 effected by heating with nitric acid in an evaporating 
 dish until no more red fumes are given off. Chlorhy- 
 dric acid should then be added, the greater part of the 
 acid evaporated, water added, and the solution should 
 be tested according to (page 104, III.). If the metal 
 remains partly insoluble, a small portion of it should 
 be carefully tested by boiling with strong nitric acid, 
 to see whether it will not dissolve by using a sufficient 
 quantity of the acid. 
 
 An alloy may contain metals of each class, therefore, 
 after the treatment with nitric acid described in the 
 last paragraph, if an insoluble residue is found it should 
 be dissolved in aqua regia, and the solution boiled 
 until the smell of chlorine ceases to be given off. 
 Then if the residue had a metallic appearance the solu- 
 
SOLUTION OF METALLIC BODIES. 99 
 
 tion must be tested for gold. See (87). If it was a 
 white powder it must be tested for tin and antimony. 
 See (86) and (85). 
 
 SOLUTION IN CHLORHYDRIC ACID. 
 
 (00) Sometimes a test shows that the metallic body can 
 be readily dissolved in chlorjiydric acid (see ZINC, 
 page 39, "and IRON, page 41), and in this case the chlor- 
 hydric acid solution should be preferred, and it should 
 be tested according to page 107, VI. 
 
TESTS FOR METALS AND 
 ACIDS. 
 
 BODIES IN SOLUTION. 
 CLASSES I. AND II. (See page 95.) 
 Sttbstances dissolved in water or in acids. 
 
 (The tests I. and II., which follow, are only to be used for 
 substances dissolved in water, or where the solvent is 
 unknown.) 
 
 I. REACTION WITH TEST-PAPER. 
 
 (01) Observe the reaction with litmus paper. 
 
 (02) If the reaction is acid, add to a small portion of the 
 solution sodic carbonate, drop by drop, until the effer- 
 vescence ceases, and then heat to boiling. 
 
 If metals of the 2d, 3d,. 4th, and 5th groups are 
 present, a precipitate will be formed, except in a few 
 special cases. 
 
 Some idea of the amount of free acid present in the 
 solution can be formed by observing the amount of 
 sodic carbonate required to neutralize it. 
 
 II. EVAPORATION. 
 
 (03) Test one or two drops of the solution by evaporation 
 on platinum foil (or on a bit of glass or porcelain, if 
 there is danger of injury to the platinum), to see 
 whether it leaves a residue. 
 
 100 
 
PRELIMINARY EXAMINATION. 
 
 If this test shows the presence of solid matter in the 
 solution, a portion of the latter may be evaporated to 
 dryness in a small porcelain dish, and some of the tests 
 for solids, particularly the EXAMINATION ON CHARCOAL 
 (page 85), and the CONCENTRATED SULPHURIC ACID 
 TEST (page 91), may be applied to the dry substance 
 thus obtained. It is best only to use the evaporation 
 and tests applied to the dry substance, to settle any 
 doubts that remain in regard to the constitution of the 
 solution, after the tests usually applied to solutions 
 have been performed. (See page 104, etc.) 
 
 When a solution of unknown origin is presented for 
 analysis it should always be heated, in order to see 
 whether a gas is given off that can be recognized by 
 the tests (page 92, and page 93). 
 
 Silicic Acid is usually not present, except in alkaline solu- 
 tions ; it occurs, however, in small quantities in spring 
 and river waters, and it may also exist in acid solu- 
 tions. 
 
 Unless silicic acid is known to "be absent it should 
 be tested for, and removed before making the remain- 
 ing tests. To this end render the solution acid with 
 chlorhydric acid, and evaporate carefully to dryness 
 in a small porcelain dish. Care must be taken not to 
 heat the dish over the lamp after its contents have be- 
 come dry. Treat the dry substance with a few drops 
 of acid, and then boil with water. If there is an inso- 
 luble residue, it consists of silica, SiO 2 . (See page 71.) 
 
 Subject the chlorhydric acid solution to the remain- 
 ing tests, beginning with page 107, VI. ; and if there 
 is reason to suspect that the precipitate, left after the 
 treatment with an acid, contains other substances be- 
 sides silica, examine it according to page 105, IV. 
 
PART III. 
 
 TESTS FOR METALS. 
 
 THE following scheme of testing for metals is founded upon 
 the successive precipitation of a number of groups, which in- 
 clude all the metals as far as Group II. After the metals of 
 the higher groups have been removed by precipitation, or have 
 been found to be absent, those of Group II. are precipitated 
 successively. Sodium' and potassium are detected by the 
 colors of their flames in a solution which has been freed from 
 all the higher metals except magnesium. Ammonia can be 
 detected in a solution without having reference to its other 
 constituents. 
 
 After the separation into groups has taken place by precipi- 
 tation with the general reagents, each precipitate, which may 
 contain one or all the metals belonging to its group, is usually 
 dissolved, and the further analysis is performed by testing in 
 the several solutions for all the metals which they may contain. 
 These special tests sometimes require the separation of the dif- 
 ferent metals, one after another, in a particular order, while 
 sometimes a test for a metal may be applied to the solution 
 without regard to the presence of other metals. In all cases 
 the conditions requisite for applying the tests will be described. 
 
 The general tests must be applied in the following order : 
 ist, Chlorhydric acid to effect the precipitation of silver and 
 mercurous compounds, and of lead if it is present in large 
 quantity. 2d, Sulphydric acid in an acid solution to precipi- 
 tate small quantities of lead and the metals of Group V., Sec- 
 tion II., and of Group VI. (From this precipitate the metals 
 of Group VI. are separated by dissolving them in ammonic sul- 
 
TESTS FOR METALS. IO3 
 
 phide.) 36, Ammonic hydrate until the reaction becomes 
 alkaline, ammonic chloride and ammonic sulphide to precipi- 
 tate the metals of Groups III. and IV. 
 
 Whenever a single metal or a group of metals is precipitated 
 some of the liquid containing the precipitate must be poured 
 on a filter, and the first drops of the solution which run x 
 through must be tested with some of the reagent which was 
 used to produce the precipitation, in order to ascertain whether 
 it has been completely effected. Should a fresh precipitate 
 make its appearance, everything must be poured back from 
 the filter into the test-tube or flask and more of the reagent 
 must be added. This operation must be repeated until no 
 precipitation is produced by the same reagent in the liquid 
 which runs through the filter. After a little practice it is easy 
 to estimate how much of a reagent is required to effect a com- 
 plete precipitation. A surplus over this quantity is called an 
 excess of the reagent. 
 
 When it has been found that a slight excess of the reagent 
 has been added, the whole of the liquid and the precipitate 
 together must be poured upon a fresh filter and the liquid 
 must be allowed to drain off. The liquid is to be tested for 
 metals of the succeeding groups, and the precipitate must 
 usually be completely freed from it ; otherwise the separation 
 has no value. To this end water must be blown on the preci- 
 pitate from the wash-bottle and allowed to drain off. Care 
 must be taken not to let the water overflow the edge of the 
 filter. The washing must be continued until the water which 
 flows through the filter is proved, either by evaporation on 
 platinum foil or by the application of tests for the succeeding 
 groups, not to contain any metal in solution. The last portions 
 of the wash-water which pass through the filter may be 
 thrown away, as they contain very little of the substances to 
 be tested for. 
 
 The value of analyses depends upon the care with which 
 the separation of precipitates from the liquid in which they 
 
104 
 
 PART III. 
 
 have been formed is executed. For special directions for 
 filtering, see pages 23 and 24. 
 
 When the reaction of a liquid in a test-tube is to be tested, 
 always close the tube and shake it thoroughly before dipping 
 the test-paper in. 
 
 METALS OF GROUPS VI. AND V. 
 
 III. CHLORHYDRIC ACID TEST. 
 
 Metals in acid solution. 
 
 GROUP V., SECTION I. 
 
 In case the solution is known to contain chlorhydric acid, pass to 
 page 107 (VI.). 
 
 If the solution has an alkaline reaction, pass to page 105 (IV.). 
 
 (63) Add to a very small quantity of the solution to be 
 tested a few drops of dilute chlorhydric acid. If a 
 precipitate forms, continue to add chlorhydric acid, 
 drop by drop, as long as it seems to increase in quan- 
 tity, then add a quantity of chlorhydric acid about 
 equal to that already added. 
 
 If no precipitate is formed, or if it is dissolved on 
 further addition of dilute chlorhydric acid, LEAD IN 
 
 LARGE QUANTITY, SILVER, AND MERCUROUS SALTS 
 
 ARE ABSENT. Pass to page 107 (VI.). 
 (06) If a permanent precipitate is formed, it may consist 
 
 Of PLUMBIC, ARGENTIC, AND MERCUROUS CHLORIDES. 
 
 Most of the succeeding tests for the detection of bases 
 must be performed in a solution freed from these 
 metals. Therefore if a precipitate has been observed, 
 treat a considerable quantity of the solution in the 
 same way that the small portion was treated, and after 
 an excess of chlorhydric acid has been added shake 
 
CHLORHYDRIC ACID TEST. 
 
 105 
 
 the liquid for a minute or two. The precipitate will 
 then settle in a short time, leaving the solution nearly 
 clear. The solution should be decanted through a 
 filter, and the precipitate washed twice by decantation 
 through the filter with water acidulated with chlor- 
 hydric acid. 
 
 Test the filtrate according to page 107 (VI.). 
 
 Lead. Add a small quantity of water to the precipitate^ 
 
 (67) and boil, then let the precipitate settle, and decant the 
 clear liquid through the same filter which was used in 
 (66) into another vessel. Add to the filtrate an equal 
 bulk of alcohol and a small quantity of dilute sul- 
 phuric acid. If a white precipitate forms it consists 
 
 Of SULPHATE OF LEAD, PbSO 4 . 
 
 If lead is found, the precipitate must be washed as 
 before, with boiling water by decantation, until the 
 filtrate gives no black precipitate with ammonic sul- 
 phide. If no precipitate remains after the washing, NO 
 ARGENTIC OR MERCUROUS SALTS are present. Pass to 
 page 107 (VI.). 
 Silver. If a precipitate remains after washing with boiling 
 
 (68) water, add to it ammonia, pour the solution through 
 a filter, and acidify with nitric acid. If a precipitate 
 forms, it consists of ARGENTIC CHLORIDE, AgCl.. 
 
 ftfercurous Salts. If a precipitate remains after ammo- 
 
 (69) nia has been added, it will have a gray or black color. 
 The precipitate consists of A MERCUROUS COMPOUND 
 
 OF AMMONIA. 
 
 IV CHLORHYDRIC ACID TEST. 
 
 Metals in alkaline solutions. 
 
 (70) Add chlorhydric acid until the reaction becomes 
 distinctly acid, and if a precipitate forms, wash it 
 thoroughly with cold water upon a filter until the fil- 
 
I0 6 PART III. 
 
 trate is no longer acid to test-paper. Observe whether 
 sulphydric acid is given off. 
 
 If a precipitate is formed with chlorhydric acid, filter 
 and test the filtrate according &>page 107 (VI.). 
 (71) If no precipitate is formed, pass to page 107 (VI.). 
 
 (a) If the precipitate is white, it may consist of : 
 Plumbic Chloride. Boil a little of it with water, and 
 
 test one portion of the solution for LEAD according to 
 (07)) and test another portion for CHLORINE by add- 
 ing nitric acid and argentic nitrate. 
 
 Plumbic Sulphate. Test according to (140). 
 
 Argentic Chloride. Test according to (141). 
 
 (b) If the precipitate is colored it may contain 
 
 The Sulphides of Arsenic, Antimony, and Tin. 
 Test according to page 109 (VIII.). 
 
 Sulphur may be precipitated, accompanied by a disengage- 
 ment of sulphydric acid. The precipitated sulphur 
 can be recognized by its appearance, and its insolu- 
 bility in aqua regia. 
 
 V. CHLORHYDRIC ACID TEST FOR ACIDS. 
 See also Silicic Acid (64). 
 
 If a solution is acid, many of the following tests, 
 particularly (73), (74), and (7S), can be applied 
 by simply heating the solution. 
 
 Carbonic Acid is only present in alkaline solutions. CO 2 
 is evolved with effervescence, when an acid is added, 
 until the solution has an acid reaction. Hold a drop 
 of lime-water on the end of a glass rod in the tube ; if 
 CARBONIC DIOXIDE, CO 2 , is present, a white precipitate 
 forms. (See page 70, foot-note.) The same test can 
 be applied to solid carbonates. See also the sulphuric 
 acid test (41). 
 
 The following acids need only be looked for when an 
 
S ULPH YDRIC A CID TEST. ! o 7 
 
 odor can be perceived after heating the solution, or after 
 adding chlorhydric acid and heating. 
 Cyanhydric Add, in its soluble combinations with 
 
 (73) most metals, is set free by chlorhydric acid. It can 
 be recognized by its smell. See also the prussian blue 
 test (133), 
 
 Sulphydric Acid is evolved from alkaline solutions 
 
 (74) (often with effervescence), on the addition of chlorhy- 
 dric acid, when the solution is heated. It can be re- 
 cognized by its smell and by the lead paper test. 
 (See the sulphuric acid test (4#) and the argentic 
 nitrate test (129). 
 
 SulpJlllTOUS Oxide is evolved from alkaline, neutral, or 
 (7>) slightly acid solutions of sulphites, on the addition of 
 chlorhydric acid. Mix a little potassic ferricyanide 
 and ferric chloride, and hold a drop of the mixture on 
 the end of a glass rod in the tube after chlorhydric 
 acid has been added and the tube heated. If a blue 
 color appears, SULPHUROUS ACID is present. If sul- 
 phydric acid is present add sufficient plumbic acetate 
 to precipitate it before performing the test. See the 
 sulphuric acid test (40). 
 
 VI. SULPHYDRIC ACID TEST. 
 
 Metals in acid solutions. 
 
 (76) Add sulphydric acid solution to a small quantity of 
 the solution to be tested, and warm gently. In case 
 metals of Group VI. are to be tested for, it is better 
 to pass sulphydric acid gas into the dilute solution, 
 made acid with chlorhydric acid. The total precipi- 
 tation of the metals of Group VI. is frequently only 
 effected after one or two days. 
 
 If no precipitate forms, no metals of Grozips V. and 
 VI. are present. Pass to page 115 (X.). 
 
I0 8 PART III. 
 
 (77) If a precipitate forms observe the color. It may con- 
 sist Of the SULPHIDES OF LEAD, PbS J BISMUTH, Bi 2 S 8 J 
 
 COPPER, CuS ; MERCURY, HgS ; and GOLD, Au 2 S 3 , 
 when it is black ; ARSENIC, As 2 S 3 ; TIN (BISULPHIDE), 
 SnS 2 ; CADMIUM, Cd$>, yellow j TIN (MONOSULPHIDE), 
 SnS, brown ; ANTIMONY, Sb 2 S 3 or Sb 2 S 5 , orange. The 
 presence of a black sulphide hides the color of the 
 other sulphides, so that all may be present when the 
 precipitate is black. 
 
 If only a light, fine, white precipitate, which is not de- 
 stroyed by acids, is formed, it consists of sulphur, and is 
 frequently due to the presence of a ferric salt or a chro- 
 mate in the solution. In case only sulphur is precipi- 
 tated, pass to page 115 (X.). 
 
 (78) If a precipitate forms in a small portion of the solu- 
 tion, a sufficient quantity for use in all the succeeding 
 tests for metals must be treated with sulphydric acid 
 until the metals of Groups V. and VI. are completely 
 precipitated as sulphides ; and the precipitate thus 
 obtained must be washed on a filter quickly, with 
 warm water containing sulphydric acid, until the ad- 
 dition of ammonic hydrate to the filtrate ceases to 
 produce a precipitate, and it must then be treated ac- 
 cording to (79). 
 
 Test the filtrate for metals of Groups IV., III., II., 
 and I. (See page 115, X., etc.) 
 
 VII. SOLUBILITY OF THE SULPHYDRIC ACID 
 PRECIPITATE IN AMMONIC SULPHIDE. 
 
 Add ammonic sulphide to a small quantity of the 
 precipitated sulphides (76), and warm gently. 
 
 If the precipitate dissolves entirely, it consists of the 
 sulphides of metals of Group VI. Those of Group V. 
 
METALS OF GROUP VI. 
 
 109 
 
 are absent. Test the remainder of the precipitate accord- 
 ing to (82) (a). 
 
 (80) If a part of the precipitate does not dissolve, add 
 four or five parts of water, and separate the solution 
 by nitration from the undissolved precipitate. 
 
 The part of the precipitate which is insoluble in ammo- 
 nic sulphide, after being carefully washed, must be tested 
 for sulphides of metals of Group V. (See page 112, 
 IX.) 
 
 (81) The ammonic sulphide solution obtained in (80) 
 may contain metals of Group VI. Add to it gradually 
 dilute chlorhydric acid until the solution becomes acid, 
 and observe the color and general appearance of the 
 precipitate which is produced. It is well to boil the 
 liquid after the formation of a precipitate. 
 
 If only a fine white precipitate forms, which remains a 
 long time in suspension in the liquid, even after boiling, it 
 consists of sulphur, and metats of Group VI. are absent, 
 and the tests described in VIII. can be omitted. 
 
 A flocculent precipitate, or one that becomes so on boil- 
 ing, indicates the presence of metals of Group VI., and 
 the color of the precipitate shows what metals predomi- 
 nate. Pass to the following tests : 
 
 VIII SEPARATION OF METALS OF GROUP VI. 
 
 (82) If the test (page 108, VII.) has shown the presence 
 of metals of Group VI., and if the precipitate with 
 sulphydric acid (77) was not entirely soluble in am- 
 monic sulphide, the whole of that precipitate must be 
 treated two or three times with ammonic sulphide, as 
 directed (79), (80), and (81) ; and the sulphides of 
 Group VI. must be precipitated from the solution, and, 
 after careful washing, treated as described in (83). 
 
 (a) Were the sulphides, precipitated by sulphydric acid, 
 
HO PART III. 
 
 wholly soluble in ammonic sulphide (see 79), it is 
 sufficient to wash and dry the portion of the precipi- 
 tate (77) which was not treated with ammonic sul- 
 phide, and to use it for (83)* 
 
 (83) Free the precipitate (82) as completely as possible 
 from water by pressing the filter and its contents be- 
 tween several thicknesses of filter-paper, remove the 
 precipitate from the filter and heat it with concen- 
 trated chlorhydric acid. The sulphides of antimony 
 and tin will be dissolved, while the sulphide of arsenic 
 remains undissolved. Collect the residue in a filter, 
 wash, and dry it at 100, and test for arsenic as fol- 
 lows : 
 
 Arsenic. Place a small portion of the dried residue in a 
 (84:) small tube closed at one end, and put over it about six 
 times its bulk of a mixture of equal parts of sodic car- 
 bonate and potassic cyanide. Heat the portion of the 
 tube above this mixture, and afterwards the mixture 
 itself, gently ; if as the result of this any moisture is 
 deposited in the upper part of the tube, wipe it out 
 carefully with a rolled-up strip of filter-paper. When 
 the whole is thoroughly dry, heat the lower part of the 
 tube with its contents to a red heat. ARSENIC, if 
 present, is sublimed and deposited on a black or 
 brownish ring in the upper and cooler part of the 
 tube. 
 
 Antimony. Concentrate the solution of the other two sul- 
 (85) phides that was filtered from the arsenious sulphide 
 
 * By heating the dry precipitate in a glass tube, or with less accuracy, 
 by heating it before it is dry, on a bit of glass or porcelain, an approximate 
 test may be made (see 13) ; and in case arsenic sulphide alone is indicated 
 by the complete volatility of the precipitate, this test is conclusive, and the 
 remaining tests in the separation of metals of Group V. may be omitted. 
 The test has little value except when the pure yellow color of the precipi- 
 tate gives rise to the suspicion that only arsenic sulphide is present. 
 
METALS OF GROUP VI. m 
 
 (S3) in an evaporator, put a small piece of zinc in the 
 concentrated solution, and bring the edge of a piece 
 of platinum foil in contact with the zinc for a minute 
 or two. If ANTIMONY is present the portion of the 
 platinum immersed in the liquid will be stained black 
 by a thin deposit of that metal. 
 Tin. Put the rest of the concentrated solution of the sul- 
 
 (86) phides in an evaporator with more zinc, collect the 
 precipitated black flakes that may appear after a time 
 on a filter, wash carefully, and pour a little concen- 
 trated chlorhydric acid on the filter. TIN, if present, 
 will be dissolved, and the solution that passes through 
 will give a white precipitate, or perhaps a gray one if 
 much metal is present, with mercuric chloride. (See 
 page 56.) 
 
 Gold. (The analytical chemist usually knows whether it is 
 
 (87) necessary to test for gold or not. In ordinary analyses 
 its presence would be improbable.) When gold is 
 present, it will be found accompanying the sulphide of 
 arsenic remaining after treatment with hot chlorhydric 
 acid (83). Heat a portion of this residue in a por- 
 celain crucible or on a piece of a broken evaporator 
 until the arsenic and excess of sulphur have been vol- 
 atilized. Dissolve the portion which remains in a 
 mixture of chlorhydric and nitric acids. Evaporate 
 the solution nearly to dryness, dilute with water, and 
 add ferrous sulphate solution. The formation of a 
 brown or purple precipitate of METALLIC GOLD, either 
 immediately or after heating, indicates the presence 
 of the metal. 
 
IT2 PART III. 
 
 IX. SEPARATION OF METALS OF GROUP V. 
 (SECTION II.) 
 
 (88) If a portion or the whole of the precipitate obtained 
 with sulphydric acid is insoluble in ammonic sulphide 
 see (80), free it by careful washing from the liquid in 
 which it was formed, or from the ammonic sulphide 
 which was used to dissolve the soluble portion, place 
 it in a porcelain dish, pour upon it pure concentrated 
 nitric acid, and heat it gently, if red fumes are given 
 off, until they cease. In any case, complete the oper- 
 ation by adding a little water, and boiling the contents 
 of the dish for a few minutes. 
 
 If no part of the precipitate ', or if only yellow particles 
 of sulphur remain insoluble, mercury is absent. Pass to 
 
 (90). 
 
 In this test, when the liquid holding sulphur in sus- 
 pension is boiled, the sulphur melts, and may enclose 
 particles of black sulphide, which, then become very 
 difficult to dissolve, and the appearance of the sulphur 
 may, in such a case, lead to an erroneous conclusion 
 that mercuric sulphide is present. It is for this reason 
 that the precipitate is oxiflized with strong nitric acid, 
 as far as possible at 'a- temperature below its boiling 
 point, before the sulphides are finally boiled with a 
 somewhat weaker acid. The same cause makes the 
 confirmatory test for mercury with stannous chloride 
 (89) necessary, when there appears to be an insolu- 
 ble black sulphide. 
 Mercury. If a black sulphide, HgS, remains insoluble 
 
 (89) after the above treatment (88), a MERCURIC SALT is 
 probably present. Confirmatory test : Boil the black 
 insoluble sulphide with chlorhydric acid and a little 
 potassic chlorate in a porcelain dish, and evaporate 
 until the greater part of the acid is volatilized ; dilute 
 
METALS OF GROUP V. (SECTION II.) 113 
 
 with water (it is not necessary to filter), and add stan- 
 nous chloride. A white precipitate of MERCUROUS 
 CHLORIDE, Hg 2 Cl 2 , is formed, if mercury is present. 
 
 If mercury is present, dilute a few drops of the 
 nitric acid solution of the sulphides (S3) with water, 
 and add sulphydric acid. If a black or brown preci- 
 pitate is formed the solution must be tested for LEAD, 
 BISMUTH, COPPER, and CADMIUM. Pass to (90). If 
 the precipitate is yellow, cadmium alone is present 
 (93). If no precipitate is obtained all of these four 
 metals are absent. Pass to page 115, X. 
 LectcL. Add a few drops of strong sulphuric acid to a small 
 
 (90) portion of the nitric acid solution of the sulphides 
 (88) and evaporate until dense, white fumes of sul- 
 phuric acid appear, and dilute with a considerable 
 quantity of water. If a white precipitate forms, it 
 
 Consists Of SULPHATE OF LEAD, PbSO 4 . 
 
 The test can be made more delicate by adding an 
 equal bulk of alcohol to the solution after it has been 
 diluted with water. If lead is discovered, treat the 
 whole of the nitric acid solution of the sulphides in 
 the same way ; filter and use the filtrate for (91). 
 ^Bismuth. Add ammonic hydrate to alkaline reaction, to 
 
 (91) the nitric acid solution of the sulphides, or to the fil- 
 trate from the lead precipitate, if lead was present. 
 If bismuth is present it is precipitated as the HYDRATE 
 OF BISMUTH, Bi(HO) 8 , white. If bismuth is present, 
 filter and use the filtrate for (92) and (93). 
 
 Copper. If copper is present it is dissolved by the ammo- 
 
 (92) nic hydrate, and imparts a blue color to the solution. 
 Cadmium. If the ammoniacal solution obtained in (91) 
 
 (93) was colorless, copper is absent, but cadmium may be 
 present. In this case add sulphydric acid. A yellow 
 precipitate, CdS, indicates CADMIUM. 
 
 If copper is present, neutralize the blue ammoniacal 
 8 
 
114 PART III. 
 
 solution with chlorhydric acid and add sulphydric 
 acid. This will precipitate both CdS and CuS ; the 
 former is soluble in hot dilute sulphuric acid. After 
 washing the mixed sulphides treat the mass with hot 
 dilute sulphuric acid and filter. Cool the filtrate and 
 pass sulphydric acid again through the liquid. If 
 cadmium is present there will be a yellow precipitate of 
 cadmic sulphide, CdS. 
 
METALS OF GROUPS IV. AND III. 
 
 METALS OF GROUPS IV. AND III. 
 
 X. AMMONIC SULPHIDE TEST. 
 
 (94i) To a small portion of the nitrate from, the precipi- 
 tate produced by sulphydric acid, or to a portion of 
 the original solution, if no precipitate is produced in 
 it by sulphydric acid, add sufficient ammonic hydrate 
 (free from carbonate) to make the reaction alkaline, 
 and then, whether a precipitate is formed or not, add 
 ammonic sulphide. If the solution contains no chlor- 
 hydric acid, it is necessary to add a small quantity 
 before neutralizing with ammonic hydrate. 
 
 If a black precipitate is formed, it may contain the 
 
 SULPHIDES OF NICKEL, NiS J COBALT, CoS ; IRON, FeS ; 
 
 MANGANESE, MnS ; and ZINC, ZnS ; and the HYDRATES 
 OF ALUMINIUM, A1 2 (HO) 6 , and CHROMIUM, Cr 2 (HO) 6 . 
 If the precipitate is white, flesh-colored, or light green, 
 it can only consist of the SULPHIDES OF MANGANESE 
 AND ZINC, and the HYDRATES OF ALUMINIUM AND 
 CHROMIUM. In this latter case omit (97), (98), and 
 (101). 
 
 If no precipitate is formed no members of Groups III. 
 and IV. are present. Pass to page 123, XII. 
 (,9t>) If a precipitate was formed in the above test (94), 
 treat a considerable quantity of the solution in the 
 same way, heat the liquid, filter as quickly as possible, 
 and wash immediately with boiling water, until the 
 nitrate has no longer an alkaline reaction. 
 
 The filtrate must be tested according to page 123, XII. 
 
IT 6 PART III. 
 
 SULPHIDES INSOLUBLE IN DILUTE CHLOR- 
 HYDRIC ACID. 
 
 (96) Add to the precipitate (95) cold, dilute chlorhy- 
 dric acid ; if a black residue is insoluble, it consists of 
 the SULPHIDE OF NICKEL or COBALT. Filter and 
 examine the residue on the filter according to (#7) 
 and (98). 
 
 The filtrate must be tested according to (99). 
 If the precipitate dissolves entirely, or if only a white 
 residue is left, no NICKEL or COBALT are present. Pass 
 to (99). 
 Cobalt* Dissolve a portion of the precipitate, which proved 
 
 (97) to be insoluble in cold, dilute chlorhydric acid, in the 
 borax bead, and expose the bead to the action of the 
 outer blowpipe flame. If the bead is blue, COBALT is 
 present. If the bead is brown, NICKEL is present in 
 large quantity. The bead is blue even when more nickel 
 than cobalt is present. To test for traces of cobalt in 
 a nickel bead, detach the hot bead from the wire, heat 
 it two.or three minutes on charcoal in a good reduc- 
 ing flame, remove it from the charcoal, and melt it on 
 the platinum wire in the reducing flame. Even if 
 only traces of cobalt are present, the bead will be 
 colored blue. 
 
 Nickel. When nickel is present, and when it is nearly free 
 
 (98) from cobalt, it can be discovered by the brown color 
 which it imparts to the borax bead, and the test is 
 conclusive. 
 
 To discover small quantities of nickel in the pres- 
 ence of considerable quantities of cobalt, dissolve the 
 precipitate, insoluble in cold dilute chlorhydric acid 
 (see 90), in concentrated nitric acid, and neutralize 
 with sodic carbonate. Potassic cyanide is added till 
 the resulting precipitate has dissolved, and then sodic 
 
METALS OF GROUPS IV. AND HI. II7 
 
 hypochlorite till the liquid smells strongly of it, even 
 after being shaken. It is then boiled. If nickel is 
 present, a black precipitate of the sesquioxide (Ni 2 O 3 ) 
 is obtained. 
 
 SULPHIDES SOLUBLE IN DILUTE CHLORHYDRIC 
 
 ACID. 
 
 (00) Boil in an evaporating dish the solution which was ob- 
 tained by the treatment of the ammonic sulphide pre- 
 cipitate with cold, dilute chlorhydric acid (see 06), 
 until the smell of sulphydric acid has entirely disap- 
 peared ; add a few drops of strong nitric acid, and boil 
 a minute longer, and filter if a precipitate of sulphur 
 has formed. 
 
 If oxalic, phosphoric, and boracic acids are not known 
 to be absent, pass to page 119, XI. 
 
 (100) Add sodic hydrate to the solution obtained in (00) 
 until the reaction becomes very strongly alkaline, dilute 
 with water and boil for a few minutes. 
 
 If a precipitate forms immediately, or after boiling, 
 test it according to (101) and (102) ; it may con- 
 tain the HYDRATES of IRON, Fe 2 (HO) 6 , MANGANESE, 
 
 Mn(HO) 2 , and CHROMIUM, Cr 2 (HO) 6 . In this case 
 filter and test the filtrate, which may contain ZINC and 
 ALUMINIUM, according to (103) and (104). 
 
 If no precipitate forms, pass to (103). The solu- 
 tion may contain ZINC and ALUMINIUM. 
 Iron. Put a small quantity of the precipitate (100) on a 
 
 (101) watch-glass, dissolve it in a single drop of dilute chlor- 
 hydric acid, dilute with water, and add potassic sul- 
 phocyanate. A red color indicates the presence of 
 IRON. 
 
 Manganese and Chromium. Fuse a portion of the 
 
 (102) precipitate obtained in (100) on platinum foil with 
 
n8 PART III. 
 
 sodic carbonate and sodic or potassic nitrate. A green 
 mass indicates MANGANESE, owing to the formation of 
 sodic manganate ; a yellow mass, CHROMIUM, from the 
 formation of the sodic chromate. A very small quan- 
 tity of manganese can readily be detected in the pres- 
 ence of a considerable quantity of chromium by the 
 green color imparted to the mass. In case of doubt, 
 however, place the platinum foil in an evaporator, cov- 
 er it with water, add a few drops of alcohol, and boil. 
 The sodic chromate will be entirely dissolved, along 
 with the excess of sodic carbonate, while the sodic man- 
 ganate will be decomposed, yielding brown flakes of 
 manganese sesquioxide. These can be filtered through 
 a small filter, washed thoroughly with hot water, and 
 fused again upon platinum foil, as before, when the 
 green color will appear very distinctly. To detect a 
 small quantity of chromium in presence of excess of 
 manganese, treat the mass as above with boiling water 
 and filter. The appearance of a yellow color in the 
 solution is proof of the presence of CHROMIUM. 
 
 Acetate of lead may also be added to the solution, 
 acidified with acetic acid, and it then takes a deeper 
 yellow color and becomes turbid, or a yellow precipi- 
 tate is formed if chromium is present. The formation 
 of a white precipitate under these circumstances only 
 indicates that the carbonate of sodium employed con- 
 tained sulphate of sodium as an impurity. 
 Zinc. Add sulphydric acid to a portion of the solution in 
 (103) sodic hydrate, obtained in (100), after it has been 
 filtered, if a precipitate was formed. If a white, floccu- 
 lent precipitate forms, it consists of SULPHIDE OF ZINC, 
 ZnS. 
 
 If chromium was discovered in (102), zinc may 
 also be present in the precipitate obtained in (100). 
 Therefore, in that case, dissolve a portion of the precip- 
 
METALS OF GROUPS IV. AND III. n 9 
 
 itate by boiling with a very little dilute chlorhydric acid, 
 add sodic hydrate until the reaction is alkaline, acidify 
 with acetic acid and add sulphydric acid. A white, 
 flocculent precipitate consists of SULPHIDE OF ZINC, 
 ZnS. 
 
 Aluminium. Add to another portion of the sodic hy- 
 (104) drate solution, obtained in (100), chlorhydric acid 
 until the reaction becomes acid, and then ammonic 
 hydrate until it becomes alkaline, and boil. If a 
 white, flocculent precipitate forms, it consists of ALU- 
 MINIC HYDRATE, A1 2 (HO) 6 . This precipitate is at first 
 gelatinous, and it may easily escape notice ; it is there- 
 fore best to set the test-tube aside, and to wait a quar- 
 ter of an hour for the precipitate to settle. 
 
 XL AMMONIC SULPHIDE TEST IN CASE PHOS- 
 PHORIC, OXALIC, AND BORACIC ACIDS ARE 
 PRESENT. 
 
 If PHOSPHORIC, OXALIC, or BORACIC ACID was pres- 
 ent in the original solution, these acids, together with 
 the metals of Group II., may be contained, wholly or 
 in part, in the solution obtained in (99), for the acids 
 would be precipitated with any of the metals of Groups 
 II., III., and IV., during the treatment with ammonic 
 hydrate and ammonic sulphide in (94), and the pre- 
 cipitates would be dissolved during the treatment with 
 dilute chlorhydric acid in (96), and consequently the 
 metals and acids might be contained in the solution 
 (99). The solution (99) must be tested for phos- 
 phoric, oxalic, and boracic acids, and freed from them 
 before the ordinary course of analysis can be pro- 
 ceeded with. 
 
 PllOSpJioric Acid. Use the test (127) with a small 
 portion of the solution obtained in (99). 
 
120 PART III. 
 
 If phosphoric acid is present, use (107) and the suc- 
 ceeding tests. 
 
 If phosphoric acid is absent, use (100) and the suc- 
 ceeding tests. 
 
 In either case first perform the operation described in 
 the next paragraph. 
 
 (103) Before testing for OXALIC AND BORACIC ACIDS it 
 is necessary to set them free from their combinations 
 with the metals of Groups II., III., and IV. To effect 
 this end add sodic carbonate to a small quantity of the 
 solution obtained in (99) until the reaction becomes 
 strongly alkaline, and boil for a few minutes and filter. 
 The metals are precipitated, with the exception of a 
 portion of the aluminium, and the acids remain in the 
 solution. 
 
 [If no precipitate is formed with sodic carbonate it 
 * is unnecessary to test further for these metals or acids, 
 as in that case they cannot be present in the solu- 
 tion (99).] 
 
 Oxalic Acid. Test a portion of the filtrate obtained after 
 boiling with sodic carbonate for oxalic acid according 
 to (125). 
 
 Itoracic Acid. Test another portion of the same filtrate 
 for boracic acid according to (128). 
 
 If one or both these acids are found \ the whole of the 
 remainder of the solution obtained in (99) must be treated 
 with sodic carbonate as in (103), and the precipitate thus 
 obtained must be dissolved in dilute chlorhydric acid. The 
 solution must be used for (107) and the succeeding tests, 
 if phosphoric acid was discovered. 
 
 If phosphoric acid is absent, the solution must be used 
 for (100) and the succeeding tests. 
 
 Aluminium. If it was found necessary to treat the solu- 
 
 (106) tion obtained in (99) with sodic carbonate according 
 to (105), the filtrate from the precipitate produced 
 
METALS OF GROUPS IV. AND III. I2I 
 
 by sodic carbonate may contain a portion of the alumin- 
 ium. Add to the nitrate in (103) chlorhydric acid 
 until the reaction becomes acid, and then ammonic 
 hydrate until the reaction becomes alkaline. If a 
 white, flocculent precipitate forms immediately, or after 
 long standing, it contains aluminium. 
 Iron. Add a few drops of potassic sulphocyanate to the 
 
 (107) solution obtained in (99) ; a red color indicates the 
 presence of IRON. 
 
 (108) Add to the remainder of the solution obtained in 
 (99), if phosphoric' acid alone is present, or to the 
 solution obtained in (105), if oxalic or boracic acid 
 is likewise present, ferric chloride, until a few drops, 
 treated with ammonic hydrate on a watch-glass, give 
 a yellow and not & white precipitate ; dilute largely with 
 water, render the solution alkaline with ammonic 
 hydrate, and add a considerable excess ; then add 
 acetic acid until the solution has a slight acid reaction, 
 and boil for a few minutes in a flask. 
 
 The precipitate contains all the IRON, ALUMINIUM, 
 CHROMIUM, and PHOSPHORIC ACID present in the solu- 
 tion so treated. Test according to (109) and 
 (110). 
 
 The filtrate contains the MANGANESE, ZINC, and 
 probably part of the BARIUM, CALCIUM, and MAGNE- 
 SIUM which were present in the original solution. 
 
 The operations described, page 115, X., #;z^page 123, 
 
 XII., must be repeated with this filtrate, omitting those 
 
 which relate to the separation of NICKEL and COBALT, 
 
 and the detection of IRON, CHROMIUM, and ALUMINIUM. 
 
 Chromium. Test a portion of the precipitate obtained 
 
 (109) in (108) for chromium according to (102). 
 Aluminium. Boil the remainder of the precipitate ob- 
 
 (110) tained in (108) with sodic hydrate, and test the 
 solution for aluminium according to (104). 
 
I22 PART III. 
 
 [The method of precipitation by boiling the acetic 
 acid solution used in (10S) can be used in all cases 
 for the separation of aluminium, chromium, and iron 
 (ferric salts) from the metals of all other groups, except 
 the Groups V. and VI., and it is preferable to any other, 
 but it demands more skill in manipulation.] 
 
METALS OF GROUP II. 
 
 METALS OF GROUP II. 
 
 XII. DETECTION OF BARIUM, STRONTIUM, CAL- 
 CIUM, AND MAGNESIUM. 
 
 Should the filtrate, after removal of the metals of Groups 
 III. and IV., have a brown color, it can only come from the 
 presence of nickel, a small quantity of the sulphide of that 
 metal having been dissolved in the excess of ammonic hy- 
 drate and ammonic sulphide used. Before proceeding to the 
 detection of metals of Group II., the nickel must be entirely 
 removed, and this can be readily accomplished by boiling for 
 a few minutes and filtering again. 
 
 (Ill) To a small portion of the solution, to which the 
 previous tests have been applied, or to a solution which 
 has been found to contain no metals of the higher 
 groups, add ammonic hydrate until the reaction be- 
 comes alkaline, and then ammonic carbonate, and boil. 
 (If the solution does not already contain ammonic 
 chloride, this also must be added, to prevent the precip- 
 itation of magnesium.) 
 
 If a white precipitate forms, it can only consist of 
 
 BARIC CARBONATE, BaCO 3 , STRONTIUM CARBONATE, 
 
 SrCO 8 , and CALCIC CARBONATE, CaCO 3 . 
 
 If no precipitate forms, BARIUM, STRONTIUM, and 
 CALCIUM are absent ; pass to (116). 
 
 If a precipitate was formed with ammonic carbonate, 
 the whole of the solution must be treated as described 
 
I24 PART In ' 
 
 above. Filter, wash the precipitate, and test the fil- 
 trate according to (116). Dissolve the precipitate by 
 pouring a very little dilute chlorhydric acid on the 
 filter, and use the solution thus obtained for (112), 
 (113), (114), and (115). 
 Barium. To a small portion of the solution in chlorhydric 
 
 (112) acid add a considerable quantity of calcic sulphate. 
 If a precipitate forms immediately, it consists of BARIC 
 SULPHATE, BaSO 4 . 
 
 Strontium. If on the addition of calcic sulphate, as 
 
 (113) in (112), a precipitate appears only after some little 
 time, it consists of STRONTIUM SULPHATE, SrSO 4 . 
 
 ( 1 14) If BARIUM or STRONTIUM is discovered by means of 
 calcic sulphate, add dilute sulphuric acid to another 
 portion of the chlorhydric acid solution. Boil, filter, 
 and test the filtrate for calcium (11&). Examine 
 the precipitate on platinum wire moistened with chlor- 
 hydric acid in the flame. (See pages 26 and 34.) 
 Barium and strontium can both be detected, even when 
 a small quantity of one is present with a large quantity 
 of the other. After placing a small particle of the 
 precipitate on the loop of the platinum wire the parti- 
 cle should be repeatedly moistened in chlorhydric 
 acid and subjected again to the action of the flame. 
 Where strontium is present in very small proportion 
 the barium color will, after repeated moistening with 
 chlorhydric acid, finally give place to the crimson of 
 strontium. Where, however, the reverse proportion is 
 found, the detection of barium is not so easy by this 
 method. In this case add to a portion of the chlor- 
 hydric acid solution obtained in (111) a solution of 
 strontium sulphate. A faint white cloudiness appear- 
 ing after some time indicates the presence of barium. 
 
 Calcium. If barium, or strontium, or both, were present in 
 
 (115) the precipitate obtained on the addition of ammonic 
 
METALS OF GROUP II. 125 
 
 carbonate, add ammonic hydrate to alkaline reac- 
 tion to the filtrate obtained after precipitation with 
 dilute sulphuric acid (114), and then ammonic 
 oxalate. If a white precipitate forms, it consists of 
 
 CALCIC OXALATE, CaC 2 O 4 . 
 
 If neither barium nor strontium was found, the whole 
 of the precipitated carbonate (111) must have con- 
 sisted of calcic carbonate. Confirm by flame reaction, 
 Page 35. 
 
 It is quite possible that owing to the presence of a 
 large excess of ammonic chloride in the solution, the 
 calcium may have escaped precipitation by ammonic 
 carbonate. In this case it would come down as a floc- 
 culent precipitate in the test for magnesium (116). 
 To detect calcium under these circumstances filter 
 the precipitate obtained in (110) and dissolve it in 
 acetic acid, dilute considerably and add ammonic 
 oxalate. If calcium was contained in the precipitate 
 - obtained by hydric disodic phosphate, it will now be 
 precipitated as the oxalate, which is insoluble in acetic 
 acid. Filter the calcic oxalate, and to the filtrate add 
 ammonic hydrate to alkaline reaction, when the ammo- 
 nio-magnesic phosphate will be re-precipitated. 
 Magnesium. To the filtrate from the precipitate pro- 
 (116) duced by ammonic carbonate (111), or to the solu- 
 tion in which no precipitate was obtained on addition 
 of that reagent, add sodic phosphate. If a white pre- 
 cipitate forms (frequently only after the lapse of some 
 minutes), it consists of the AMMONIO-MAGNESIC PHOS- 
 PHATE, MgNH 4 PO 4 . 
 
126 PART III. 
 
 METALS OF GROUP I. 
 
 XIII. DETECTION OF SODIUM, POTASSIUM, AND 
 AMMONIUM. 
 
 Before testing for sodium and potassium, precipitate the 
 metals of Groups V. and VI. which are present with sulphy- 
 dric acid, and precipitate those of Groups II., III., and IV. 
 with a mixture of ammonic carbonate and sulphide, and use 
 the solution, freed from those metals, for the following tests 
 (117) and (118): 
 Sodium. Evaporate the solution to dryness, and drive off 
 
 (117) ammonia salts, if any are present, by heat. Sodium, 
 if present, can be distinguished by the yellow color 
 which a small quantity of the solid residue held in 
 
 the flame of a gas or alcohol lamp imparts to it. See 
 page 26. 
 
 Frequently sodium can be detected in a solution 
 without evaporation, by dipping a platinum wire in the 
 solution, and then holding it in the flame. 
 Potassium can be recognized by the violet color which it 
 
 (118) imparts to the flame. The solid residue obtained in 
 (117) can be tested for potassium in the same way 
 that it is tested for sodium. 
 
 If sodium is also present, its greater coloring power 
 will obscure the potassium flame, but by looking 
 through a piece of blue glass at the flame, the violet 
 color can be distinguished even when sodium is pres- 
 ent. (The color of the potassium flame is almost the 
 same as that of the heated wire, while the sodium 
 
METALS OF GROUP I. 127 
 
 flame is much more blue, if it is not excluded entirely 
 by the glass.) 
 
 Add to a portion of the original solution a 
 (HO) few drops of sodic hydrate, and heat. 
 
 AMMONIA can be recognized by its smell, or by hold- 
 ing a piece of moistened turmeric paper or red litmus 
 paper at the mouth of the test-tube, taking care not to 
 let it touch the sides, which may be moistened with 
 sodic hydrate. The turmeric paper will be turned 
 brown, and the litmus paper blue, if ammonia is 
 present. 
 
I2 8 PART III. 
 
 TESTS FOR ACIDS. 
 
 IT is usual to take a fresh quantity of the solution to test 
 for acids, and generally the tests for metals precede those for 
 acids, in order that information gained by the first series of ex- 
 periments may point out the most convenient way of detecting 
 the acids. Silicic acid, however, is always first precipitated 
 from a solution as directed, page 101 (64), and the acids no- 
 ticed under the chlorhydric acid test (page 106, V.) are to be 
 looked for while that test is applied to the detection of the 
 metals. Phosphoric, oxalic, and boracic acids interfere with 
 the ordinary methods of testing for the metals of Groups II., 
 III., and IV., and consequently these acids must be tested for 
 as directed, page 119, XL, during the application of the tests 
 for the metals of those groups. 
 
 In all other cases, when the presence of any metals or acids 
 in the solution interferes with the performance of the test for 
 an acid, directions are given under the head of each acid for 
 removing them. 
 
 It is obvious that metals and acids which precipitate each 
 other cannot be present together in a solution, and that when 
 certain metals have been found, the number of acids which it 
 becomes necessary to look for is restricted within limits de- 
 termined by this consideration. Therefore the knowledge 
 already acquired of the composition of a solution must be 
 brought to bear upon the problem of testing for acids. 
 
 First, the reaction which the solution gives with test-paper 
 must be considered, and then the tables IV. and V. must be 
 
ACIDS OF GROUP I. 
 
 129 
 
 consulted to ascertain what acids can exist in a solution pos- 
 sessing the observed reaction, together with the metals which 
 have been discovered. 
 
 For instance, if a solution contains lead, and is neutral or 
 nearly neutral, the only acids which can be present in it are 
 acetic, chlorhydric, chloric, and nitric. If the reaction is 
 strongly acid the solution may contain all the acids except 
 sulphuric, sulphydric, ferro- and ferricyanhydric. Moreover, 
 the solution cannot contain a large quantity of lead and chlor- 
 hydric acid at the same time, because the chloride of lead is 
 only soluble in 135 parts of water. 
 
 ACIDS OF GROUP L 
 
 Arsenious and Arsenic Acids are always detected 
 
 (120) by the sulphydric acid test in searching for the metals. 
 When these acids are discovered, they must always be 
 precipitated by sulphydric acid before testing further. 
 
 Chromic Acid cannot be present in a solution to which 
 
 (121) SULPHYDRIC ACID or AMMONic SULPHIDE has been 
 added (see page 64). If chromic acid is present in a 
 solution, it must be contained in the precipitate ob- 
 tained with baric chloride (122), and can be detected 
 by heating a small portion of the precipitate in the 
 borax bead. If CHROMIC ACID, H 2 CrO 4 , is present, 
 the bead will be colored green. Chromic acid can 
 often be recognized by the yellow color which it im- 
 parts to solutions which contain it, and by the yellow 
 precipitate, PbCrO 4 , which is obtained by adding 
 plumbic acetate to the neutral or slightly acid solu- 
 tion. 
 
 9 
 
I3 o PART III. 
 
 XIV. BARIC CHLORIDE TEST. 
 
 Baric nitrate and nitric acid should be used instead of baric 
 chloride and chlorhydric acid, when lead, silver, or mercurous 
 salts have been discovered in the solution. 
 
 (122) Put a piece of litmus paper in the solution, and if 
 the reaction is acid, add ammonic hydrate, drop by 
 drop, until it becomes slightly alkaline. If a precipi- 
 tate is formed in consequence, add dilute chlorhydric 
 acid only in sufficient quantity to dissolve it. Add 
 baric chloride, and if a precipitate forms, it indicates 
 the presence of ARSENIOUS, ARSENIC, CHROMIC, SUL- 
 PHURIC, SULPHUROUS, OXALIC, FLUORHYDRIC, PHOS- 
 PHORIC, BORACIC, CARBONIC, OR SILICIC ACIDS. (Sili- 
 
 cic acid cannot be present after the operation (04) 
 has been performed). 
 
 If these acids are absent, pass to the argentic nitrate 
 test (page 132, XVIIL). 
 
 (The baric chloride test is of little value except 
 when the substance is soluble in water with a neutral 
 or slightly alkaline reaction, or in case the reaction is 
 acid, when the metals of Groups II., III., IV., and V. 
 are absent. The following special tests are more ac- 
 curate :) 
 Sulphuric Acid. Acidify (if the solution is not already 
 
 (123) acid) with dilute chlorhydric acid, in considerable ex- 
 cess, and add baric chloride as long as a precipitate 
 continues to form. (Use nitric acid and baric nitrate 
 if chlorhydric acid produces a precipitate.) If sul- 
 phuric acid is present, it is precipitated as BARIC SUL- 
 PHATE, BaSO 4 , fine white powder. The solution must 
 not be heated when it is intended to use the filtrate 
 for the next test. 
 
 Sulphurous Acid. To the nitrate from the precipitate 
 
 (124) produced by baric chloride, or to the acid solution to 
 
CALCIC SULPHATE TEST. 131 
 
 which baric chloride has been added without produc- 
 ing a precipitate, add potassic dichromate, and boil. 
 If a precipitate forms, it consists of BARIC SULPHATE, 
 BaSO 4 , produced by the oxidation of SULPHUROUS 
 ACID, H 2 SO 3 , contained in the solution. Usually the 
 CHLORHYDRIC ACID TEST (7<>) is more convenient 
 and sufficiently accurate. 
 
 XV. CALCIC SULPHATE TEST. 
 
 If the solution contains metals which are precipitated by 
 sulphydric or sulphuric acids, they must first be removed by 
 adding a slight excess of those precipitants and filtering. 
 
 Add, if the solution is alkaline, add acetic acid 
 until the reaction becomes acid ; if it is acid, add sodic 
 hydrate until the reaction becomes alkaline, and then 
 add acetic acid until it becomes acid, and test as below 
 for oxalic acid. If a precipitate forms and does not 
 dissolve in the acetic acid, add to the original solution 
 a considerable excess of sodic carbonate, boil, fil- 
 ter, add to the filtrate acetic acid until its reaction 
 becomes acid, and test as follows for oxalic acid : 
 Add calcic sulphate in considerable quantity. If a 
 precipitate forms, it consists of CALCIC OXALATE, 
 CaCsO*, white powder. Fluorhydric acid is the only 
 other acid which precipitates calcic sulphate under 
 these circumstances, and as the precipitate is almost 
 transparent and gelatinous, it cannot easily be mistaken 
 for that produced by oxalic acid. 
 
 Acid can only be present in alkaline solu- 
 (126) tions in glass vessels. If there is reason to suspect 
 the presence of this acid, add calcic chloride and am- 
 monic hydrate to the solution, and if a precipitate 
 forms, collect it on a filter, and examine it for fluorine 
 according to (44). 
 
132 PART III. 
 
 XVI. AMMONIC MOLYBDATE TEST. 
 
 Phosphoric Acid. Make the solution strongly acid (if 
 
 (127) it is not so already) with nitric acid, and add a small 
 portion of it to a considerable quantity of ammonic 
 molybdate solution.* If phosphoric acid is present, 
 PHOSPHO-MOLYBDATE OF AMMONIUM, yellow crystalline 
 powder, is precipitated. If the quantity of phosphoric 
 acid in the solution is very small, the precipitate does 
 not form until after several hours. 
 
 If sulphydric acid is present, it is necessary to heat 
 the acid solution until it is expelled, before perform- 
 ing the test. See (136) (a) for this test in the pres- 
 ence of ferro- or ferricyanhydric acid. 
 
 XVII. TURMERIC PAPER TEST. 
 
 Boradc Add.\ Strongly acidify the solution (if it is not 
 
 (128) already acid) with dilute chlorhydric acid, dip a piece 
 of turmeric paper in it, and dry the paper by holding 
 it over the lamp-flame, without charring it. If a red 
 or brownish-red stain appears upon the paper when it is 
 dry, it is due to the presence of BORACIC ACID, H 8 BO 8 . 
 
 ACIDS OF GROUP II. 
 
 XVIII. ARGENTIC NITRATE TEST. 
 
 Acidify with nitric acid, if the solution is not already acid, 
 and add argentic nitrate. 
 
 Sulphydric Acid. If a black precipitate is formed, it 
 (129) must contain ARGENTIC SULPHIDE, Ag 2 S, showing 
 
 *See foot-note, page 67. f Ibid., page 68. 
 
ARGENTIC NITRATE TEST. I33 
 
 that sulphydric acid was present in the solution. The 
 precipitate may also contain ARGENTIC CHLORIDE, 
 
 BROMIDE, IODIDE, CYANIDE, FERRO- and FERRICY- 
 ANIDE. 
 
 If a precipitate is formed which is not black, it can 
 only be due to presence of CHLORHYDRIC, BROMHY- 
 
 DRIC, IODHYDRIC, CYANHYDRIC, FERROCYANHYDRIC, 
 and FERRICYANHYDRIC ACIDS. 
 
 If no precipitate is formed, none of the above acids are present. 
 Pass to (137). 
 
 Chlorhydric Acid. See also (48). Acidify strongly 
 
 (130) with concentrated nitric acid, add argentic nitrate in 
 excess, shake thoroughly if a precipitate forms, and 
 allow it to settle, decant the liquid, and pour on the 
 precipitate strong nitric acid, and boil for five minutes ; 
 if a precipitate remains undissolved, it consists of 
 
 ARGENTIC CHLORIDE, AgCl, if violet. If yellow, of 
 
 ARGENTIC BROMIDE. See (132) (a). 
 
 See (136) (a) for this test in the presence of ferro- 
 or ferricyanhydric acid. 
 IBromhydric Add gives with argentic nitrate a white 
 
 (131) curdy precipitate, which darkens on exposure to light, 
 is insoluble in boiling concentrated nitric acid, and is 
 not so readily soluble in ammonia as the argentic 
 chloride. See also (49). 
 
 lodliydric Acid gives with argentic nitrate a yellow pre- 
 
 (132) cipitate of the iodide which is nearly insoluble in am- 
 monia. See also (50). 
 
 (132) Detection ^/CHLORHYDRIC, BROMHYDRIC, and IODHY- 
 (d) DRIC acids in presence of each other. Mix the liquid to 
 be tested with a few drops of dilute sulphuric acid, 
 then with a little starch paste, and add a few drops of 
 fuming nitric acid, or a solution of hyponitric oxide 
 in sulphuric acid, when the blue color characteristic 
 
I34 PART III. 
 
 of iodine will appear. (See page 74.) Add now 
 chlorine water until that reaction has disappeared. 
 On continuing the gradual addition of chlorine water 
 the bromine will be set free, and will impart a yellow 
 or brownish color to the liquid. (See page 73.) 
 Chlorine in presence of bromine can best be detected 
 by the following method : Evaporate the solution to 
 dryness and mix the residue with a little potassic di- 
 chromate. Place the mixture at the bottom of a clean 
 test-tube and pour on it a few drops of concentrated 
 sulphuric acid. On the application of heat dark red 
 
 drops Of CHROMYL BICHLORIDE, Or CHROMIC OXY- 
 
 CHLORIDE, CrO 2 Cl 2 , will condense in the upper por- 
 tion of the test-tube. Bromides, under the same 
 treatment, give a similar result ; but in the latter case 
 the distillate consists of bromine, which is instantly 
 . decolorized by a drop of ammonic hydrate. In the 
 case of chlorides, when the CrO 2 Cl 2 is treated with 
 ammonic hydrate, it gives a yellow solution, owing to 
 the formation of ammonic chromate [(NH 4 ) 2 CrOj. 
 
 Where the substance to be tested for bromine and 
 iodine was not soluble in water it should be fused on 
 platinum foil with sodic carbonate. The mass is then 
 treated with water and the aqueous solution used for 
 the foregoing method of separation. 
 
 XIX. PRUSSIAN BLUE TEST FOR CYANHYDRIC 
 
 ACID. 
 
 Cyanhydric Acid. Add to the solution ferrous sul- 
 (133} phate and a few drops of ferric chloride ; add sodic 
 hydrate until a precipitate forms (unless the solution 
 is alkaline and a precipitate forms without the addition 
 of sodic hydrate), warm for a minute, and add dilute 
 chlorhydric acid until the reaction becomes acid. 
 
FERRIC CHLORIDE AND FERRO US SULPHA TE TESTS. 
 
 135 
 
 The appearance of a blue precipitate or a blue color in 
 the solution is evidence of the presence' of cyanhydric 
 acid. 
 
 See ( 136) (b) for this test in the presence of ferro- 
 or ferricyanhydric acid. 
 
 XX. FERRIC CHLORIDE TEST. 
 
 FerrocyanTiydric Acid. Add a little ferric chloride 
 (134=) to the acid solution. If ferrocyanhydric acid is pres- 
 ent, a precipitate of PRUSSIAN BLUE, Fe 4 (FeCy 6 )3, deep 
 blue, is formed. 
 
 XXI. FERROUS SULPHATE TEST. 
 
 Ferricyanhydric Acid. Add a little ferrous sulphate 
 (135) to the acid solution. If ferricyanhydric acid is pres- 
 ent, a precipitate of TURNBULL'S BLUE, Fe 3 (Fe 2 Cyi 2 ), 
 deep blue, is formed. 
 
 (130) If ferro- or ferricyanhydric acid is present, before 
 (a) performing the tests for PHOSPHORIC ACID (127), and 
 CHLORHYDRIC ACID (130), the following steps must 
 be taken : Add dilute sulphuric acid, dilute with 
 water if the solution is not dilute, add cupric sulphate, 
 and finally add enough baric nitrate * to render the 
 precipitate of a decidedly lighter color ; heat almost 
 to boiling, allow the precipitate to settle for a few 
 minutes, filter, and use the filtrate for the tests (127) 
 and (130). In case the test (51) for nitric acid is 
 to be used take the same preliminary steps ; using 
 baric chloride in place of baric nitrate. 
 
 * Sulphuric acid and baric nitrate or chloride are only added in order to 
 produce a heavy precipitate of baric sulphate, which carries down with it 
 the lighter particles of the other precipitates, and renders the filtration 
 easier. 
 
I3 6 PART III. 
 
 (136) If ferro- or ferricyanhydric acid is present, the test 
 (b) for CYANHYDRIC ACID (133) is to be modified in the 
 following manner : Dilute with water if the solution is 
 not dilute, add dilute sulphuric acid, then add, accord- 
 ing as ferro- or ferricyanhydric acid is present, ferric 
 chloride or ferrous sulphate, or both, in sufficient 
 quantity to precipitate the ferro- or ferricyanhydric 
 acid or both acids ; finally, add baric chloride * until 
 the color of the precipitate has become decidedly 
 lighter ; shake thoroughly, allow the precipitate to 
 settle for a few minutes, and filter. Add to a portion 
 of the filtrate sodic hydrate until a precipitate forms, 
 warm gently, and add dilute chlorhydric acid until 
 the solution becomes acid. The appearance of a blue 
 precipitate, or of a blue color, is evidence of the pres- 
 ence Of CYANHYDRIC ACID. 
 
 ACIDS OF GROUP III. 
 
 ACIDS WHICH ARE NOT PRECIPITATED BY ANY 
 METALS. 
 
 Chloric Acid. See sulphuric acid test (47). 
 
 (13V) 
 Nitric Acid. See sulphuric acid test (SI). 
 
 (138) 
 Acetic Acid. See sulphuric acid test (52). 
 
 (139) 
 
 *See note, page 135. 
 
INSOLUBLE SUBSTANCES. 
 
 CLASS III. 
 
 SUBSTANCES WHICH ARE INSOLUBLE IN WATER 
 AND IN ACIDS. 
 
 (See page 97.) 
 
 The only substances which are insoluble after the treat- 
 ment described on page 97 are the following : 
 
 Plumbic Sulphate (not absolutely insoluble in acids). 
 
 Argentic Chloride (slightly soluble in chlorhydric 
 acid). 
 
 Sulphur. 
 
 Carbon. 
 
 Baric Sulphate, Silica, and many Silicates, 
 and some Oxides. 
 
 XXII. SOLUTION IN AMMONIC ACETATE AND 
 POTASSIC CYANIDE. 
 
 Plumbic Sulphate. Boil a portion of the substance 
 (14=0) with ammonic acetate, and test the solution (after fil- 
 tration, if necessary) with ammonic sulphide. If 
 LEAD is present, a black color, or a black precipitate is 
 formed. 
 
 Test the solution also for SULPHURIC ACID, accord- 
 ing to (123). 
 
 If lead is discovered, repeat the treatment with am- 
 monic acetate until no more lead is dissolved. 
 Argentic Chloride. Digest a portion of the substance, 
 free from plumbic sulphate with potassic cyanide, 
 
138 PART III. 
 
 warm (unless it blackens by warming), and test the so- 
 lution (after nitration, if necessary) with ammonic 
 sulphide. A black precipitate indicates the presence 
 of SILVER. If a black precipitate forms, wash it, dis- 
 solve it in strong nitric acid, and test with chlorhydric 
 acid according to (68) in order to confirm the pres- 
 ence of SILVER. 
 
 If silver is present, repeat the treatment with potas- 
 sic cyanide until no more silver is dissolved. 
 Sulphur. Test the substance, free from plumbic sulphate 
 
 (142) and argentic chloride, for sulphur according to (10). 
 If the substance is moist, it must be carefully dried by 
 heating it in a porcelain dish over a water-bath before 
 applying the test. 
 
 If sulphur is present, heat the substance in a cov- 
 ered porcelain crucible until the sulphur is completely 
 volatilized. 
 Carbon. If the substance has a black or gray color, which it 
 
 (143) loses when it is heated with the blowpipe on platinum 
 foil, carbon in some form is probably present. If 
 carbon is present, the substance, free from plumbic 
 sulphate, argentic chloride, and sulphur, should be 
 burnt, until as much as possible of the carbon is de- 
 stroyed, by heating it red-hot on platinum foil or in a 
 porcelain crucible. 
 
 XXIII. FUSION WITH POTASSIC AND SODIC 
 CARBONATES AND SODIC NITRATE. 
 
 Baric Sulphate, Silicic Acid, and many Sili- 
 
 (144) cates, and some Oxides. Mix the finely pow- 
 
 (a) dered substance, free from plumbic sulphate, argentic 
 
 chloride, and sulphur, and as nearly free from carbon 
 
 as possible, with two parts of potassic carbonate, two 
 
INSOLUBLE SUBSTANCES. 139 
 
 parts of sodic carbonate, and one part of sodic nitrate ;* 
 bring as much of the mixture as can be heated at 
 once on the platinum foil, and heat the under side of 
 the foil with a blast-lamp until the whole mass is in a 
 state of quiet fusion. Repeat this operation two or 
 three times, if much substance is required for the ana- 
 lysis. 
 
 (b) Detach the fused mass from the platinum foil each 
 time by plunging the foil, while it is hot, in distilled 
 water. Boil the product of fusion with water, and if 
 it does not dissolve completely, filter, and wash the 
 precipitate on the filter with distilled water, rejecting 
 the washings. Continue the washing until baric 
 chloride ceases to produce a precipitate in the water 
 which runs through the filter. 
 
 (143) The first filtrate may contain ARSENIC ACID, see 
 (120), (its occurrence is rare) ; CHROMIC ACID, see 
 (121) ; SULPHURIC ACID (123), (the tests referred 
 to above may be applied successively to a single por- 
 tion of the filtrate) ; FLUORHYDRIC ACID (its occur- 
 rence is rare), see (126) and (44), and PHOSPHORIC 
 ACID (127). The two last tests may be applied suc- 
 cessively to another portion of the filtrate. No com- 
 pound of these acids, except BARIC SULPHATE, is by 
 itself insoluble, but insoluble substances sometimes 
 contain small quantities of the acids. CALCIC FLU- 
 ORIDE is only decomposed completely by the treat- 
 ment with sulphuric acid described in (44). 
 Silicic Add* The principal portion of the filtrate should 
 (140) be tested according to (04) for silicic acid. After 
 
 * The sodic nitrate is added in order to destroy carbon or other reduc- 
 ing substances. If the substance to be analyzed appears to contain much 
 carbon, increase the quantity of sodic nitrate. If the substance contains 
 no carbon, the use of sodic nitrate is usually unnecessary. 
 
140 
 
 PART III. 
 
 separation of silica the only metals * that can be pres- 
 ent in the chlorhydric acid solution are LEAD, ALU- 
 MINIUM, and ZINC. Test for lead by adding an excess 
 of dilute sulphuric acid and alcohol to the solution. 
 If a precipitate of PLUMBIC SULPHATE forms, filter. 
 Test for ALUMINIUM, in the solution, free from lead, 
 by adding ammonic hydrate in excess. If a precipi- 
 tate of ALUMINIC HYDRATE forms, filter. Test for 
 ZINC in the solution, free from lead and aluminium by 
 adding to the solution containing ammonic hydrate in 
 excess ammonic sulphide. A flocculent, white precip- 
 itate Consists Of SULPHIDE OF ZINC. 
 
 (14f) If a portion remains insoluble after boiling the fused 
 mass with water (1.44) (&), dissolve it in chlorhydric 
 acid. If much silica was discovered (see 146), it 
 is best to evaporate the chlorhydric acid solution to 
 dryness, and to proceed as directed in (64). 
 
 Test for metals in the chlorhydric acid solution accord- 
 ing to page 107, VI., etc. 
 
 * It is evident that sodium and potassium in insoluble silicates cannot 
 be detected by this process. All reliable methods for their detection re- 
 quire the use of platinum vessels and great care in manipulation. Larger 
 works on analysis must be consulted for such methods, 
 
EXPLANATION OF TABLES. 
 
 THE Tables I., II., and III. contain a synopsis of the course 
 of analysis of bodies in solution given in Part III., and they 
 are intended as an index to the methods which are there de- 
 scribed in detail. They may also serve as guides in analytical 
 work to students who have made themselves acquainted with 
 the detailed descriptions of Part III. 
 
 A skeleton form, similar to that of the tables, should be filled 
 out with the results of an analysis, and the reactions which oc- 
 cur on the application of each test should be noted. 
 
 The sign ^ placed under the formula of a compound 
 indicates that it is formed as a precipitate during a reaction. 
 This sign is used in the following tables, and it will also be 
 found convenient in noting the results of analyses. 
 
 The Tables IV. and V. are intended to indicate the degree 
 of solubility in water, and in many cases in alcohol, acids, and 
 alkalies, of the combinations of the metals and acids men- 
 tioned in Part II. 
 
 The properties of a salt are described in the square formed 
 by the intersection of the column devoted to an acid with that 
 devoted to a metal. 
 
 The Roman numerals standing after the symbols of the 
 metals indicate their quantivalence, and the formula of a salt is 
 made by putting the symbol of a metal in the place of the 
 symbol of an equivalent number of atoms of hydrogen in an 
 acid. When an acid contains more than one atom of hydro- 
 gen, several classes of salts may be formed, according as one 
 or more atoms of hydrogen are replaced by a metal. The 
 normal or regular salts are those which are formed by the re- 
 
 141 
 
142 
 
 PART III. 
 
 placement of the greatest passible number of atoms of hydro- 
 gen by a metal. 
 
 The descriptions of the tables refer to normal salts, but the 
 following cases are exceptions, because the salts specified are 
 more commonly met with in analysis ; and in using the tables 
 the formulas below must be substituted for those of the normal 
 salts : 
 
 ARSENATES. MgNH 4 AsO 4 ; MnNH 4 AsO 4 ; (Hg 2 )HAsO 4 ; 
 
 HgHAs0 4 . 
 PHOSPHATES. (NH 4 ) 2 HP0 4 ; BaHPO 4 ; CaHPO 4 ; MgNH 4 
 
 P0 4 ; MnNH 4 P0 4 ; HgHPO 4 ; Na 2 HPO 4 . 
 The ARSENATE OF ALUMINIUM probably contains more acid 
 
 than the normal salt. 
 The CHROMATES OF ALUMINIUM and of IRON (ferric chro- 
 
 mate) contain a larger proportion of metal than the 
 
 normal salts. 
 ARSENITES. The arsenites referred to in the table have only 
 
 two atoms of hydrogen replaced by a metal, except Mg 3 
 
 (AsO 3 ) 2 and Ag 3 AsO 3 . 
 The ARSENITES OF COBALT and MANGANESE contain less than 
 
 two atoms of hydrogen replaced by the metal. 
 BORATES (NH 4 ) 2 B 4 O 7 ; BaB 2 O 4 ; CuB 4 O 7 ; FeB 4 O 7 ; (Fe 2 ) 
 
 B 3 O 6 ; PbB 4 O 7 ; CaB 2 O 4 ; MnB 4 O 7 ; NiB 4 O 7 ; K 2 B 2 O 4 ; 
 ; Na 2 B 4 O T ; ZnB 4 O 7 . 
 
EXPLANATION OF SIGNS IN TABLES IV. AND V. 
 
 Numbers = number of parts of water required to dissolve one part of the 
 
 anhydrous salt * at the ordinary temperature, 
 oo . insolubility.! The sign of infinity indicates that an infinite 
 
 quantity of water is required to dissolve the salt, 
 s. = soluble to a considerable extent in water, 
 s.s. = slightly soluble, 
 del. = deliquescent, or capable of dissolving by attracting moisture 
 
 from the air. 
 dec. = decomposed. Examples : dec. = decomposed by water. 
 
 dec. = decomposed by acids. 
 = acids. Example : s. = soluble in acids. 
 + = sodic or potassic hydrate. Example : + s. = soluble in so- 
 
 dic or potassic hydrate, 
 am. = ammonic hydrate, 
 am. cl. = ammonic chloride. 
 
 al. = alcohol. Example : al. 00.= insoluble in alcohol. 
 When no solvent, such as , am., al., etc., is indicated, the signs co., 
 s., s.s., and dec. refer to the action of water on the salt. 
 
 * The salt without water of crystallization is referred to. 
 
 t Most of the salts marked insoluble in the table are not really more insoluble than 
 salts like baric sulphate, but they are described as insoluble because they are known to 
 be nearly so, and because the quantity of water required to dissolve them has not been 
 determined. 
 
 143 
 
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 SULPHYD 
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 TABLE V. EXPLAN 
 
 
 Acetic Acid, HC a H 3 O 2 
 
 Arsenic Acid, H 3 AsO 4 
 
 Arsenious Acid, ) 
 H,As0 8 . S~~ 
 
 Boracic Acid, H 3 BO 3 . 
 
 Bromhydric Acid.HBr 
 
 V^arbonic Acid, H a CO 3 
 
 ^XfhlorhydricAcid.HCl. 
 
 Chloric Acid, HC1O S . 
 
 Chromic Acid,H 2 CrO 4 
 
 Cyanhydric Acid, HCy 
 
 -U 
 
 || 
 
 ^ 
 
 Fluorhydric Acid, HF. 
 
 X 
 
 .g 
 
 O 
 
 K 
 
 y 
 
 'B 
 
 Oxalic Acid, H 2 C 2 O 4 . 
 
 y 
 
 If 
 
 N/Sulphuric Acid, H 2 SO 4 
 
 !2 
 
 P" 
 
 ,/Sulphydric Acid, H 2 S. 
 
~1VO 'OOSIONVMd NVS 
 
 " '!/! .SH3AVSSV 
 
 ^!AiI^S 
 ( 2H: ilJLSflf 
 
 .Vki Ci'lOW 
 
14 DAY USE 
 
 RETURN TO DESK FROM WHICH BORROWED 
 
 LOAN DEPT. 
 
 This book is due on the last date stamped below, or 
 
 on the date to which renewed. 
 Renewed books are subject to immediate recall. 
 
 8 May'GSPSf 
 REC'D LP 
 
 MAY 8 1963 
 
 200ef65CD 
 
 REC'D LD 
 
 OC1 6'65-9pM 
 
 4 - 1968 
 
 LD 
 
 8EC1JLQ. FEB 23 71 -2PM 5 7 
 
 LD 21A-50m-8,'61 
 (Cl795slO)476B 
 
 General Library 
 
 University of California 
 
 Berkeley 
 
M81896 
 
 c.7 
 
 THE UNIVERSITY OF CALIFORNIA LIBRARY