IRLF 277 fifi7 Summer Sohool 0* ahemistry* LIBRARY " . OF THE UNIVERSITY OF CALIFORNIA. D>(f T *ea!h &CO ; . 551 Sansome St;S;F;Calr - Received (J^tfl^s ...., ifyte- Accession No. Class No. ^ O 1 r 1 Q r THE ELEMENTS CHEMICAL ARITHMETIC A SHORT SYSTEM OF ELEMENTARY QUALITATIVE ANALYSIS. BY J. MILNOR COIT, PH.D., MA8TBB IN ST. PAUL'S SCHOOL, CONCORD, N.H. Qt TH3 ' BOSTON: D. C. HEATH & COMPANY. 1895. 1 COPYRIGHT, MAR. 22, 1886, BY J. MILNOR COIT. J. S. Gushing & Co. Berwick & Smith. Boston, Mass., U.S.A. PEEFACE. THIS little manual is intended to supplement the teaching of the text-books of descriptive chemistry, and to be used as a companion to them, by those who desire to make the whole subject more practical. It is the result of the author's experience after several years of elementary science teaching. Part I. contains some of the more important rules and principles of chemical arithmetic, followed by a series of problems, which will not be found to be above the comprehension of the average student in the schools. The matter relating to chemical theory, and the rules, have been collected from the best authorities. Part II. is devoted to an elementary system of quali- tative analysis, and the best methods have been adopted. This part of the book can be used separately, and can be taught together with any good work in descriptive chemistry, such as Eliot & Storer's, Shepard's, RemseSPi, or Avery's Chemistry. An intelligent student can, with the occasional supervision of his instructor, work out by himself the reactions and the separations as given in iv PREFACE. the tables. The tables are those generally in use. Tests are given for the more common metals and acids only, and the reagents indicated are those which almost any school laboratory will afford. J. M. C. ST. PAUL'S SCHOOL, Concord, N.H. 1711 PAKT I. CHEMICAL ARITHMETIC. CHAPTER I. INTRODUCTION. 1. Matter is anything that occupies space. 2. Divisions of Matter. Three divisions of matter are recognized in science, masses, molecules, and atoms. A mass of matter is any portion of matter appreciable by the senses. A molecule is the smallest particle of matter into which a body can be divided without losing its identity ; or it is the smallest portion of matter which can exist by itself. An atom is a still smaller particle produced by the division of a molecule; or it is the smallest portion of matter that can go into combination. EXAMPLES. The sun and a grain of sand are masses of matter. The smallest particle of salt which can exist and which exhibits the properties of salt is a molecule. The minute particles of chlorine and sodium which com- pose the molecule of salt are atoms. A mass is made up of molecules, and a molecule is com- posed of atoms. 3. Attractions of Matter. The three forms of attrac- tion admitted in science are : 2 CHEMICAL ARITHMETIC. First. Gravitation, or the attraction between masses. Second. Cohesion, or the attraction between like mole- cules; adhesion between unlike molecules. Third. Chemical attraction, or the attraction between unlike atoms. EXAMPLES. The attraction between the sun and the planets, or between the earth and all bodies upon it, is gravitation. The attraction between the molecules of a piece of marble is cohesion. The attraction between a liquid and solid, as, for instance, when you dip your hand into water it becomes wet, or between two different solids at the surface, as shown by the action of cements, is adhesion. The attraction between the unlike atoms of chlorine and sodium by means of which we have an entirely different substance, salt, is chemical attraction. 4. Province of Physics. Physics is that department of physical science which studies the results which come from the molar and molecular conditions of matter. 5. Province of Chemistry. Chemistry studies matter in its atomic condition. It investigates the laws and con- ditions of chemical changes, and seeks to account for some of the phenomena connected therewith. 6. Physical Changes. Physical changes are those which take place outside the molecule; they have no effect upon the molecule itself nor alter the identity of the matter operated on. The study of physics is a study of physical changes. 7. Chemical Changes. Chemical changes take place through the atoms and within the molecule. They alter the character of the molecule, and hence destroy the INTRODUCTION. 3 identity of the matter itself. The study of chemistry is a study of chemical changes. EXAMPLES. The change of water into ice and steam, or the change of any solid into a liquid, or of any liquid into a vapor, are physical changes. But when water is subjected to the influence of the electric current, it undergoes a more radical change ; the water disappears, and in its place appear two gaseous substances, oxygen and hydrogen, entirely different from the water from which they were derived. This is a chemical change. 8. Physical Properties. Physical properties are those properties which bodies possess in virtue of their molar or molecular condition. 9. Chemical Properties. Chemical properties are those which result from the atomic composition of the molecule. 1C. Chemistry defined. Chemistry is that branch of physical science which treats of the atomic composition of bodies, and of those changes in matter which result from an alteration in the kind, the number, or the relative position of the atoms which compose the molecule. 11. Analysis and Synthesis. The two processes by which the chemist seeks to find out the composition of matter are analysis and synthesis. Analysis consists in separating the molecule into its constituent atoms. Synthesis consists in putting together constituent atoms to form the molecule. CHEMICAL ARITHMETIC. CHAPTER II. MOLECULES AND ATOMS. 12. Chemical Definition of the Molecule. A mole- cule is the smallest particle of any substance which can exist in a free state in nature. Molecules classified. Molecules are of two classes : First. Elementary molecules, or those whose atoms are alike. Second. Compound molecules, or those whose atoms are unlike. 13. Simple Substances are those whose molecules con- tain like atoms. 14. Compound Substances are those whose molecules contain unlike atoms. 15. Number of Simple Substances. There are sixty- eight elementary substances, as far as has been investi- gated by chemical science | that is, sixty-eight substances whose molecules contain like atoms. Therefore it is obvious that there are sixty-eight different kinds of atoms. From combinations of these sixty-eight kinds of atoms all the different varieties of matter result. We cannot resolve a simple substance into any other substances or atoms. 16. Ampere's Law. "Equal volumes of all gases, simple as well as compound, under like conditions of tem- perature and pressure, contain the same number of mole- cules." MOLECULES AND ATOMS. <" 5 From this law, which is the most -important law of modern chemistry, it results, First. That the molecules of all bodies in the gaseous state are of the same size. Second. That the weight of any molecule, compared with that of hydrogen, is proportional to the weight of any given volume, also compared with the same volume of hydrogen. EXAMPLES. If 1 liter of nitrogen, which weighs 14 times as much as a liter of hydrogen, contains the same number of molecules, then it is obvious that each molecule of nitrogen must be 14 times as heavy as a molecule of hydrogen. 1 7. Number of Atoms in the Molecule of Hydrogen. Assuming that 1 volume of hydrogen contains 1000 molecules, then, according to the law of Ampere, 1 vol- ume of chlorine will contain 1000 molecules also. Suppose these volumes (that is, 1 volume of hydrogen containing 1000 molecules and 1 volume of chlorine con- taining 1000 molecules) be mixed together and exposed to the action of the sunlight, they combine, forming 2 volumes of the new substance, hydrochloric acid gas, which 2 volumes, by the same law, will contain 2000 molecules. Upon analysis, each molecule of hydrochloric acid gas will be found to contain 1 atom of hydrogen and 1 atom of chlorine. That is, the 2000 molecules will contain 2000 atoms of hydrogen and 2000 atoms of chlorine. The 2000 molecules will contain, therefore, 4000 atoms ; or, each molecule will contain 2 atoms. Hence each molecule of hydrogen is made up of 2 atoms. 6 CHEMICAL ARITHMETIC. 18. Molecular Weights/ If the weight of the hydro- gen atom be taken as 1, then, since its molecule contains 2 atoms, its molecular weight will be 2. Since the molecule of a compound gas or vapor occupies a volume twice as large as that occupied by the atom of hydrogen, it is obvious that the specific gravity of the gas or vapor may be found from the molecular weight by dividing the latter by 2. The specific gravity of a com- pound gas or .vapor is, therefore, one-half its molecular weight. The molecular weight of any substance may be obtained by multiplying its density in the state of gas by the molecular weight of hydrogen ; that is, by 2. EXAMPLES. The density of oxygen gas, for example, is 16 ; that is, any volume as 1 liter weighs 16 times as much as 1 liter of hydrogen. Its molecule must be, therefore, 16 times as heavy. The molecular weight of hydrogen is 2 ; therefore the molecular weight of oxygen will be 16 x 2 = 32. The weight of 1 liter of hydrogen is called 1 crith, and the weight of the hydrogen atom 1 microcrith. 19. Number of Atoms in the Molecule. The number of atoms in a molecule is obtained by dividing the molecu- lar weight by the atomic weight. EXAMPLE. The molecular weight of oxygen is 32, and its atomic weight 16. The number of atoms in the mole- cule is 32 divided by 16 = 2. The molecular weight of phosphorus is 124, and its atomic weight 31; its mole- cule, therefore, contains 4 atoms. MOLECULES AND ATOMS. 7 PROPERTIES OF ATOMS. 20. Definition. An atom is the smallest particle of simple matter which can enter into the composition of a molecule. 21. Atomic Weight. The relative weight of any atom referred to hydrogen as unity is its atomic weight. It is the smallest weight of any simple substance which can take part in the formation of a chemical compound. The molecular weight of any substance is the sum of the weights of its constituent atoms. 22. Quaiitivalence. The quantivalence of an atom is the quality of its combining power, expressed in hydrogen units. It expresses the number of hydrogen atoms with which it can combine or for which it can be exchanged. EXAMPLES. The quantivalence of zinc is 2, because 1 atom replaces 2 of hydrogen. The quantivalence of carbon is 4, because 1 atom of carbon requires 4 of hydrogen to satisfy it in combination. Atoms are called monads, dyads, triads, tetrads, pentads, hexads, and heptads, according to their quantivalence. The Latin nu- merals are used for the adjective terms. These atoms are univalent, bivalent, trivalent, quadrivalent, quinquivalent, sexivalent, and septivalent. Atoms whose quantivalence is even are called artiads ; those whose quantivalence is odd are called perissads. An atom may form several compounds with the same substance. Therefore its quantivalence may vary. It always increases or diminishes by 2, so that it may have quantivalence of 1, 3, 5, or 7, or of 2, 4, or 6. A perissad atom can never become an artiad by such a change, nor can an artiad become a perissad. 8 CHEMICAL ARITHMETIC. EXAMPLES. 'Iron in iron sulphate is a dyad, in pyrites it is a tetrad, and in ferric acid a hexad. Chlorine forms a series of compounds with oxygen in which its quantivalence is 1, 8, 5, and 7. Atoms are divided into two classes, according to the quality of their combining power. First. Positive atoms are those which are attracted to the negative pole in electrolysis, and whose hydrates are bases. Second. Negative atoms are those which go to the posi- tive electrode, and whose hydrates are acids. 23. Atomic Symbols. Berzelius, in 1815, proposed an abbreviated form of chemical language. In this system each atom has for its symbol the first letter of its Latin name. When the names of two different atoms begin with the same letter, a second letter suggestive of the name is added. EXAMPLES. Ag stands for an atom of silver ; Fe, for an atom of iron ; Sn, for one of tin, etc. (on page 89 will be found the table of the symbols of the elements). Each atomic symbol stands not only for the atom, but represents its atomic weight. EXAMPLES. Fe (ferrum) represents 56 weight-units of iron; Hg (hydrargyrum), 200 weight-units of mercury; O, 16 weight-units of oxygen. 24. The quantivalence of an atom is indicated by plac- ing Roman numerals above or a little to the right of the symbol. Sometimes minute-marks are used. EXAMPLES. 1. H or H' stands for the monad hydrogen atom; 2. S or S" stands for the bivalent sulphur atom; 3. P or P" ; for the trivalent phosphorus atom; 4. C or C' r " for the quadrivalent carbon atom. MOLECULES AND ATOMS. 9 Sometimes graphic symbols are used to represent the quantivalence atoms, the graphic symbols being a circle with lines called bonds radiating from it ; as, for example, Monad. Dyad. Triad. Tetrad. ^ Pentad. Hexad. 6 -o- A -9- ')6; The circles are usually omitted, the bonds radiating from the symbol. The number of bonds and not their direc- tion is significant, as, for example, -Q- Q- O~ stands I equally for 1 atom of dyad oxygen. N=, N=, or N= equally represent the atom of trivalent nitrogen. 25. Multiplication of Atoms. Atoms are multiplied by placing an Arabic numeral below and to the right of the symbol. EXAMPLES. C 2 represents 2 atoms of carbon. N*, 4 atoms of nitrogen. C1 3 , 3 atoms of chlorine. Molecules are multiplied by enclosing their symbols in brackets and placing the numeral outside, below, and to the right. EXAMPLES. (H 2 ) 6 represents 6 molecules of free hydrogen. (Br 2 ) 2 stands for 2 molecules of bromine. 10 CHEMICAL ARITHMETIC. CHAPTER III. COMPOUND MOLECULES AND VOLUME RELATIONS. 26. Compound Molecule. A compound molecule is one whose constituent atoms are unlike. Compound mole- cules are formed by the union of atoms according to the law of quantivalence. 27. Molecular Weight. The molecular weight of a compound molecule is the sum of the atomic weights of its constituents. It is always equal to twice the density of the substance in the state of gas. 28. Classification of Compound Molecules. Com- pound molecules are divided into two classes : first, those whose atoms are directly united, called Binaries ; second, those whose atoms are indirectly united, called Ternaries. A binary compound is formed by the union of two simple substances, the termination IDE being the characteristic : as, for example, sodium and chlorine yield sodium chloride; silver and sulphur yield silver sulphide ; calcium and iodine yield calcium iodide. In some cases the number of atoms of each constituent is to be indicated. This is done by prefixing Greek numerals to each of the names given ; as, for example, 1 atom of C and 2 of O form carbon dioxide, 1 atom of P and 5 of Br form phosphorus pentebromide. 29. Definition of an Acid. An acid molecule is one which consists of one or more negative atoms united by COMPOUND MOLECULES AND VOLUME RELATIONS. 11 hydrogen and Coxygsib It is a compound of hydrogen and oxygen with some non-metallic element, and possesses the property of turning blue litmus paper or solution red. 30. Definition of a Base. A basic molecule is one which contains one or more positive atoms united by hydrogen and oxygen. It is a compound of hydrogen, oxygen, and some metallic element, and possesses the property of restoring the color to vegetable blues, which have been reddened by an acid. 31. Definition of a Salt. A saline molecule is one which contains a positive atom or group of atoms, united by oxygen to a negative atom or group of atoms. It is formed by the action of an acid upon a base, and since it contains no hydrogen, has no action upon vegetable colors. 32. Compound Radical. A compound radical is a group of atoms, which goes into combination like a single atom. It may be composed of two or more elements; as, for example, (NH 4 ) ammonium, (C 2 H 5 ) ethyl. 33. Normal, Acid, Basic, and Double Salts. A salt is formed by substituting a metal for the hydrogen of an acid, each bond of the metal displacing one atom of hydro- gen. A normal salt is formed by displacing all the hydro- gen of the acid with an equivalent metal. An acid salt is formed by exchanging a part of the hydrogen of an acid for an equivalent of metal. A basic salt is formed by the substitution of a metal in part for the hydrogen of an acid, and in part for the half or the whole of the hydrogen of water (H 2 O). Double salts are those containing two or more different positive or metal atoms. 12 CHEMICAL ARITHMETIC. EXAMPLES. K' with HNO 3 forms KNO 3 , displacing H. KNO 3 is a normal salt, formed also by acting upon HNO 3 by KHO, as : - HNO 3 -f KHO = KNO 3 + H 2 O. K' with H 2 SO 4 may form KHSO 4 , an acid sulphate formed also by KHO + H 2 SO 4 = HKSO 4 + H 2 O. Bi'" with g a forms Bi | Q 3 , usually written BiONO 3 , a basic nitrate. NaCa"SbO 4 , sodio-calcium antimonate is an example of a double salt, or Ba"Zn"SiO 4 , baro-zincic silicate. Monobasic acids can form only normal salts. Polybasic acids can form normal, acid, and double salts. 34. Chemical Equations. A chemical equation is the expression in symbols of a chemical reaction, or change. The sign plus (+) indicates added to, and the sign minus ( ), taken from, and the sign of equality ( = ), equals to. The equation must be a true equation ; that is, the sum of the weights of the atoms on one side must equal the sum of the weights of the atoms on the other side. 35. The substances entering into the reaction are called factors; these constitute the first member. The substances issuing from the reaction are called products; these constitute the second members. The equation, representing the reaction of two molecules upon each other may be written by the following rule : - Place the formulas of the factors, connected by the sign plus, as the first member of the equation, and the formulas of the products, also connected by the sign plus, as the second. COMPOUND MOLECULES AND VOLUME RELATIONS. 13 EXAMPLES. Let AB and EF be two molecules. The reaction between them would be represented by the equation AB + EF= AF + BE. 36. Weight of the Factors and Products. The quantities of matter taking part in a chemical change are definite in weight, since each formula represents a definite weight, viz., the molecular weight. For the same reason no loss of weight can be the result of any chemical reac- tion. 37. There are three kinds of chemical reactions : First. Analytical reactions ; that is, the separation of a complex molecule into simpler ones. Second. Synthetical reactions, or the union of two or more simple molecules to form a more complex one. Third. Metathetical reactions, or the transposition or exchange of atoms between molecules. EXAMPLES. An analytical reaction may be represented by the general equation or, taking an actual example, HgO = Hg + O; Mercuric oxide. Mercury. Oxygen. that is, one molecule of mercuric oxide will yield one molecule of mercury and one molecule of oxygen. Synthetical reactions may be represented by the general equation E+F=EF\ or, taking an actual example, Fe + S = FeS; Iron. Sulphur. Iron sulphide. 14 CHEMICAL ARITHMETIC. that is, one molecule of iron and one molecule of sulphur yield one molecule of iron sulphide. Metathetical reactions may be represented by the gen- eral formula AB -f EF= AF+ BE ; or, practically, Zn + H 2 SO 4 = ZnSO 4 + 2H. Zinc + Sulphuric acid = Zinc sulphate -*- TT^ -ogen. 38. The conditions which form chemical change depend upon the facility with which the atoms of any molecule may be rearranged. When substances are in the gaseous or liquid state, these changes between atoms take place most readily. Hence, fusion or solution or vaporization facilitate chemical action. Heat is therefore the great aid to the chemist. ST01CHIOMETRY. 15 CHAPTER IV. ( STOICHIQMETBY. ] CHEMICAL ARITHMETIC. 39. Definition." By Stoichiometry we mean that de- partment of chemistry which treats of the numerical rela- tions of atoms. The calculations of these numerical relations, whether of volume or weight, depend upon the fact that every atom has its own weight, called the atomic weight. The atomic weight is the smallest portion by weight of any simple or elementary substance referred to the atom of hydrogen as unity which can take part in a chemical change. 40. Unit of Weight. The weight of the hydrogen atom is called a microcrith. (The weight of one liter of hydrogen under general conditions of temperature and pressure is one crith.) We adopt the term microcrith for convenience' sake. 41. All chemical changes take place between definite quantities of matter, as represented by a chemical equa- tion. An equation expresses not only the fact of chemical reaction between two bodies, but also indicates the quan- tities by weight concerned in it. RULES. 42. From the Formula of a Substance to find its Molecular Weight. The molecular weight of a com- pound is the sum of the atomic weights of all the atoms of the elements which compose it. 16 CHEMICAL AEITHMETIC. i The name of each element present being written in a column, and opposite to each the multiple of its atomic weight which is present in the compound, on adding these numbers together the molecular weight of the compound is obtained. Thus the molecular weight of sulphuric acid, H 2 SO 4 , is H= 1x2=2 S~=32x 1 = 32 O = 16x4 = 64 98 43. To find the Percentage Composition of any Sub- stance in a Molecule. Rule. Multiply the atomic weight by the number of atoms, and this product by 100. Divide the final product by the molecular weight, and the quo- tient will be the percentage amount of that constituent. EXAMPLE. What is the percentage composition of carbon dioxide, Co 2 ? Carbon =12 Oxygen, 16 x 2 = 3? Molecular weight, 44 Carbon = 12 x -^ = 27.27 per cent. Oxygen = 32 x -^ = 72.73 per cent. This rule can be expressed by a general formula. Let m represent the molecular weight, a the atomic weight of any constituent, n the number of atoms, and x its percen- tage amount ; then we have the proportion : tnfian: : 100 :#, from whence the formula m STOICHIOMET&Y. 17 In the above formula, when any three of the quanti- ties a, w, w, and x are known, the fourth can be found. Whence, to find the number of atoms of any constituent in a molecule, "a?," "a," and "w" being known, we have, by transposing formula (1) : - n = -^L, (2) alS a= m * (3) am * m = an X 1QQ (4) 44. To calculate from an Equation a Mass. Mule. Find the multiples of the atomic or molecular weights of the substances given and asked in the equation, and work the proportion. The molecular weight of substance given : the molecu- lar weight of substance asked : : the real mass of sub- stance given : the real mass of substance asked. Thus, to find now many grams of sodium sulphate can be ob- tained from 100 grams of sodium hydrate : EXAMPLE 1. 2 NaHO + H 2 SO 4 = 2 H 2 O + Na 2 SO 4 . 2 x 40 142. 2 X 40 : 142 : : 100 : x x = J-4JMLO = 177.5 grams. EXAMPLE 2. In the equation KNO 3 + H 2 SO 4 = HNO 3 + HKSO 4 101 + 98 = 63 + 136. 125 grams of KNO 3 yield 77.97 grams of HNO 3 , whose molecular weight is 63. What is the molecular weight of KN(X? 18 CHEMICAL ARITHMETIC. This rule can be simply expressed by the general pro- portion : M : tn : : W: w, where M represents the molecular weight of the substance given, m the molecular weight of the substance asked, W the real mass of the substance given, and w the real mass of the substance asked ; whence 45. The Relations of Weight to Volume. 1. To find the volume occupied by a given weight of any gas. Rule. Divide the weight of the gas given by the weight of 1 liter ; the quotient is the number of liters. 2. To find the weight of any given volume of gas. Rule. Multiply the number of liters of gas by the weight of 1 liter; the product is. the weight of the given volume. EXAMPLES. 1. What volume is occupied by 6.08 grams of oxygen, the weight of 1 liter of oxygen being 1.43 grams ? 6.08 -=- 1.43 = 4.25 liters. Am. 2. What is the weight of 25 liters of nitrogen gas, 1 liter weighing 1.26 grams? 1.26 x 25 = 31.5 grams. Ana. 46. Density of Gases. The density of any gas ex- presses how many times the gas is heavier than hydrogen. Knowing the density, the weight of 1 liter may readily be obtained by multiplying it by the weight of 1 liter of STO1CHIOMETBY. 19 hydrogen, 0.0896 grams, or 1 crith. The molecular weight of any substance being the weight of 2 volumes of the substance in the state of gas, it is evident that its density in the state of gas may be obtained by dividing its molec- ular weight by 2. With few exceptions, the density of any elementary gas is expressed by the same number as its atomic weight, and that of any compound gas is expressed by the same number as half its molecular weight. Thus, oxygen, O = 16 ; density, 16 ; or 1 liter weighs 16 criths. Ammonia, NH 3 = 17 ; density, 8.5 ; or 1 liter weighs 8.5 criths. 47. Relation of Gaseous Volume to Pressure. To calculate the change in volume of a mass of gas produced by a change in pressure. Boyle's Law. The volume of a mass of gas varies in- versely as the pressure upon it; or the volume of a mass of gas, multiplied by the pressure at any one time, is equal to the volume of the same mass of gas multiplied by the pressure upon it at any other time. Thus, let V equal the volume of a gas under the pressure P, and let V equal the volume under the pressure P' ; then ' , or F= If the pressure upon 1000 cc. of gas be increased from 400 mm. to 800 mm., what is the new volume ? 48. Relation of Gaseous Volume to Temperature. Gruy Lussac's Law. When 273 volumes of gas at C. are heated, they increase by one volume for every 1 C. through which thev are heated. Thus : 20 CHEMICAL ARITHMETIC. 273 volumes of gas at C. become at 1 C. 273 + 1 volumes, 273 " . " " " " 2 C. 273 + 2 " 273 " " " " " 3 C. 273 + 3 " 273 " " " " u t C. 273 + " where t expresses any number of degrees on the centigrade scale. The coefficient of the expansion of a gas is ^ of the volume of the gas at for every degree centigrade. Hence v volumes at t C. become at T C. .. 273 + T volumes ; which, if V stands for the volume of the gas after change of temperature t C. to T C., is usually written: 273 +* EXAMPLE 1. Find the new volume, if 1000 cc. of gas are heated from 17 C. to 27 C. The formula is : - 1000 (273 + 27) = 1000 X 300 = 1(m g c 273 + 17 290 EXAMPLE 2. If 1000 cc. of gas at - 23 C. are heated to 27 C., find the new volume. F= 1000 (273 + 27) = 1000 X 300 = 12Q() cc 273 - 23 250 49. If the pressure on the gas, as well as its tempera- ture, be changed, the above formula must be combined with the one given in (46). 273 + * P' J STOICHIOMETBY. 21 EXAMPLE. If -500 cc. of gas are cooled from 39 C. to 13 C., the pressure being decreased from 800 mm. to 300 mm., find the new volume. 500(273 + 13) 800 = 1222<2cc> 273 + 39 300 5O. Density of Oases. When the temperature of and the pressure on a gas are not mentioned, it is supposed to be at 760 mm. and C. A gas under these conditions is said to be normal. The formulae given in this and the succeeding section only apply to normal gases ; hence, when necessary, the gas under consideration must be rendered normal by using the formula : F== 0X273 P' 273+* 760' and, conversely, the volume found by these formulas is nor- mal, and must be reduced to the required temperature and pressure by T 760 273 m In the case of gases, the liter, = 1000 cc., is taken as the unit of volume, and the mass of one liter of normal hydro- gen, called a crith, = .0896 gram, is taken as the unit of mass. The density of a gas, then, is the number of criths con- tained in one liter of it, measured at C. and 760 mm. ; or the number of times it is heavier than an equal volume of hydrogen. Hence the mass in grams of a liter of any normal gas can be found by multiplying its density by .0896. 22 CHEMICAL ARITHMETIC. EXAMPLE. The density of carbon monoxide is 14 ; re- quired the weight of one liter. 14 x .0896 = 1.2544 grams. The density of a gas referred to air may be obtained by multiplying its density referred to hydrogen by .06926, the density of hydrogen referred to air. EXAMPLE. Nitric oxide is 15 times as heavy as hydro- gen ; how many times is it heavier than air ? 15 X .06926 = 1.0389. If the density of a gas referred to air be given, its den- sity referred to hydrogen can be obtained by multiplying its density referred to air by 14.436. EXAMPLE. Sulphur dioxide is 2.22 times as heavy as air ; find its density and molecular weight. 14.436 x 2.22 = 32.042 is the density referred to hydrogen, and 32.042 x 2 = 64.084 is the molecular weight. 51. Volume and Mass of Gases. It is found by ex- periment that 22.32 liters of any normal gas weigh a num- ber of grams equal to the number expressing the molecular weight of the gas. Thus : 22.32 liters of hydrogen (H 2 = 2) weigh 2 grams. 22.32 " " oxygen (O 2 = 32) " 32 " 22.32 " " nitrogen (N 2 = 28) " 28 " 22.32 " "chlorine (C1 2 = 71) " 71 "' This volume, 22.32 liters, is commonly spoken of as "two volumes " and expressed by the symbol CD. STOIC HIOMETBY. f^ y 23 Since 22.32 liters (or, if great accuracy be not required, 22.4 liters) of any gas weigh its molecular weight in grams, a liter of any gas weighs its molecular weight in grams divided by 22.32 (or 22.4); and one gram of any gas occupies 22.32 (or 22.4) liters divided by its molecular weight. Hence the mass in grams of any volume of a gas can be found by multiplying the number of liters of it by its molecular weight, and dividing by 22.32. - EXAMPLE 1. Find the mass of 250 liters of chlorine (C1 2 = 71). Conversely, the volume in liters of any gas can be found by multiplying the number of grams of it by 22.32, and dividing by the molecular weight. EXAMPLE 2. Find the volume of 225 grams of hydro- gen sulphide (H 2 S = 34). 225 X 22.32 = 147 ^ rs> 34 EXAMPLE 3. Find the mass of 80 liters of oxygen (O 2 = 32) measured at 52 C. and 7.40 mm. The gas must be reduced to C. and 760 mm. 80X273 X : x **- = = 93.47 273 + 52 760 22.4 95 52. Equation and Volumes of Gases. When the volume of one gas is given, and that of another gas is asked, since each molecular weight expresses two volumes of the gas, the result may often be obtained directly. 24 CHEMICAL ARITHMETIC. EXAMPLE 1. What volume of hydrogen chloride is formed when 10 liters of chlorine combine with hydrogen? C1 2 + H 2 =2HC1. j-ri+rTj = 2 m Two volumes of chlorine form twice two volumes of hydrogen chloride ; hence 10 liters of chlorine form 2 x 10 = 20 liters of hydrogen chloride. EXAMPLE 2. If 10 liters of hydrogen at 15 C. are burned, what volume of steam at 300 C. is formed ? H 2 -}-0 = H 2 0. mam The volume of the steam would be equal to that of the hydrogen, if the temperatures were the same, making the correction for the change of temperature. 10 (273 + 300) = 5730 = 273 + 15 288 When the mass of a solid or liquid is given or asked, and the volume of a gas is asked or given, the equation can only be solved in terms of the mass of the gas. EXAMPLE 3. How much lead sulphide can be precipi- tated by 17 liters of hydrogen sulphide? .^ HNO 3 + PbS. 34 239 17 liters of H 2 S weigh 17 * 34 grams. 34 grams of H 2 S precipitate 239 grams of lead sulphide. 239 1 gram of H 2 S precipitates - grams of PbS. 17x 84 grams of H,S precipitate 17X84 x = 181.3 grams of lead sulphide. STOIC HIOMETKY. 25 53. Gaseous Diffusion. Graham's Law. "The veloc- ity of the diffusion of any gas is inversely proportional to the square root of its density." This law applies of course to volumes. That is, when two gases diffuse through the same apparatus for equal times under similar conditions, the volume of the one gas diffused multiplied by the square root of its density is equal to the volume of the other gas diffused multiplied by the square root of its density. EXAMPLE. 4 liters of hydrogen diffuse through an apparatus in 10 minutes, and 1 liter of oxygen in an equal time under similar conditions ; find the density of oxygen. Z>=16. For method of determining the empirical formula of a substance from its percentage composition, and for methods of determining the relative density of solids, liquids, and gases, see Appendix, pp. 93 and 95. 26 CHEMICAL ARITHMETIC. EXAMPLES. MOLECULAR WEIGHTS. 1. Find the molecular weight of (a) carbon monoxide, CO; (b) magnesia, MgO; (c) lime, CaO; (d) alumina, A1 2 3 . 2. Find the percentage of oxygen in each of the above- mentioned bodies. 3. Find the molecular weight of (a) nitric oxide, NO; (>) sodium hydrate, NaHO ; (c) ferric oxide, Fe 2 O 3 . 4. Find the molecular weight of (a) zinc sulphate, ZnSO 4 .7H 2 O; (5) copper sulphate, CuSO 4 .5H 2 O; ammonium salts, mercury, arsenic, antimony. Metalic mirror indicates arsenic. If a gas is evolved : oxygen indicates chlorates, nitrates, peroxides; carbon monoxide indicates oxalates ; nitrogen tetroxide indicates nitrates ; ammonia indicates ammonium salts ; carbon dioxide indicates carbonates. If the substance alters in color : black indicates organic matter ; yellow while hot indicates zinc oxide. * Mercury, sulphur, ammonia, though not acids, arc included. 36 ELEMENTARY QUALITATIVE ANALYSIS. Take another portion of substance under analysis, and add HC1. Notice whether a gas is evolved with efferves- cence. If it smell. like burning sulphur, it indicates sulphites, or hyposulphites. If it has the odor of rotten eggs, sulphides. If it has the odor of bitter almonds, cyanides. If it has the odor of chlorine on heating, peroxides, chromates, or hypochlorites. If it renders lime-water turbid, carbonates. Take another portion of substance, and try if it is soluble in water ; if so, add BaCl 2 solution to a portion of the solution and notice whether a precipitate form. A white precipitate insoluble in HC1 indicates sulphates. White and soluble in HC1 indicates phosphates, silicates, oxalates, borates, and fluorides, also carbonates and sulphites. Yellow indicates chromates. If BaCl 2 gives no precipitate, add AgNO 3 to another portion of the solution and notice if a precipitate form. White precipitate indicates chlorides, also cyanides. Yellowish-white indicates bromides and iodides. Black indicates sulphides. In case neither water nor HC1 has dissolved the sub- stance, try HNO 3 . If this does not dissolve it, try aqua regia ; and if that fails, try method described in Table II. If the substance is dissolved in HNO 3 or aqua regia, it must be evaporated to dry ness with HC1 before proceed- ing to the examination for base. EXAMINATION FOR BASE. 37 SECTION II. Examination for Base. 55. Having obtained a solution, add HC1. If it pro- duces a precipitate, it indicates silver, had, or mercurous salts. Add HC1-|-H 2 S. If it produces a precipitate, Black indicates mercuric salts, lead, bismuth, or copper ; Yellow indicates arsenic, stannic salts, or cadmium; Orange indicates antimony ; Brown indicates stannous salts. If (NH 4 )HO + (NH 4 )C1+(NH 4 ) 2 S produce a precipi- tate, it indicates Black, iron, nickel, cobalt; White, zinc or aluminum ; Flesh-colored, manganese; Green, chromium. If (NH 4 )HO + (NH 4 )C1+(NH 4 ) 2 CO 3 produce a precipi- tate, it indicates \ Barium (tinges flame green), Strontium (tinges flame crimson), Calcium (tinges flame dull red). If the solution is not precipitated by any of the above reagents, it indicates magnesium, potassium, sodium, am- monium, of which the following are the individual tests: Magnesium is precipitated by Na 2 HPO 4 h (NH 4 )HO, white. Potassium is precipitated (except in very^jlilute solu- '' I17SESIT7 38 ELEMENTARY QUALITATIVE ANALYSIS. tions) by PtCl 4 , precipitate insoluble in alcohol ; also tinges the flame violet. Sodium is precipitated by H 2 SiF 6 ; also tinges flame in- tense yellow, not visible through blue glass. Ammonium salts heated with NaHO give smell of NH a . TABLE II. Examination of Insoluble Substances. The follow- ing substances are, under certain circumstances, insoluble in acids, and must be examined specially: Silica, /Silicates. Alumina, Aluminates. Oxides of Antimony, Chromium, and Tin. Chrome Iron Ore. Sulphates of Barium, Strontium, and Lead. Certain Fluorides (e.g. of Calcium). Certain Sulphides (e.g. of Lead). Chloride, Bromide, and Iodide of Silver. Carbon. Sulphur. Heat the substance in a dry tube as before, and notice if it fuses and volatilizes completely. If it smells of SO 2 , it indicates sulphur. If it fuses, but does not volatilize, indicates chloride, bromide, or iodide of silver (also will yield metallic silver on fusing on charcoal with Na 2 CO 3 ). If it is infusible, but disappears on heating, carbon (de- flagrates when heated with KN0 3 ). If it is infusible, but darkened in color while hot, regain- ing its color on cooling, tin dioxide and antimony pentoxide ' EXAMINATION OF INSOLUBLE SUBSTANCES. 39 (confirmed by blow-pipe test tin bead malleable ; anti- mony bead, brittle). Notice whether. It yields a green bead with borax or microcosmic salt ; it indicates chromium oxide, or chrome iron ore. It swims undissolved in a bead of microcosmic salt, silica and silicates (fuse with four times its weight of a mixture of K 2 CO 3 and Na 2 CO 3 . Allow to cool, dissolve in water, add HC1, and evaporate to dryness. Silica will separate out as a gelatinous mass). It yields a colorless bead,' with microcosmic salt, alu- mina. (Heated on charcoal, and moistened with CO(NO 3 ) 2 and reheated, it yields a blue, infusible mass.) It is white and infusible, but quite unaltered by heat- ing. Lead sulphate yields, when heated with Na 2 CO 3 in blow- pipe reducing flame, malleable metallic bead. Barium sulphate fused with Na 2 CO 3 yields BaCO 3 . Boil the fused mass with water, filter and wash; the residue dissolved in HC1 yields BaCl 2 (flame color green), precipi- tated by SrSO 4 solution. Strontium sulphate fused with Na 2 CO 3 yields SrCO 3 . Boil the fused mass with water, filter and wash; the residue dissolved in HC1 yields SrCl 2 (flame color crimson), preci- pitated by CaSO 4 solution. Calcium fluoride heated with H 2 SO 4 yields HF, which etches glass. It is black and infusible, and yields a malleable metallic bead when fused with Na 2 CO 3 in the blow-pipe flame. Lead sulphide (bead leaves mark on paper), and when dissolved in HNO 3 gives a white precipitate on addition of H 2 SO 4 . The action of strong H 2 SO 4 often affords a valuable indi- 40 ELEMENTARY QUALITATIVE ANALYSIS. cation of the nature of a salt, whether soluble or in- soluble. ' Thus evolution of Sulphur dioxide indicates sulphites or hyposulphites. Sulphuretted hydrogen indicates sulphides. Hydrocyanic acid indicates cyanides. Oxygen indicates peroxides, chromates, permanganates. Carbon dioxide indicates carbonates. Carbon monoxide indicates oxalates, formates, ferro* cyanides. Chlorine indicates hypochlorites. Hydrochloric acid indicates chlorides. Hydrofluoric acid indicates fluorides. Nitric acid indicates nitrates. Acetic acid indicates acetates. Chlorine tetroxide indicates chlorates. REACTIONS OF THE COMMONLY OCCURRING METALS WITH THE METHODS OF SEPARATION. GROUPING OF THE METALS. 56. The metals are divided into five groups, according to their behavior with certain substances termed group reagents. Group I. (Silver Group.) Group reagent, HC1. Metals whose chlorides are insolu- ble in water. They are precipitated from the solutions of their salts by the first group reagent, hydrochloric acid. Silver, mercury (rnercurous salts), lead. GROUPING OF THE METALS. 41 Group II. (Copper Group.) Group reagent H 2 S in presence of HC1. Metals which in acid solutions form insoluble sulphides, are precipitated from their acidulated solutions by the second group rea- gent H 2 S (hydrosulphuric acid). Arsenic, antimony, tin, lead, bismuth, copper, cadmium, mercury (mercuric salts). The three metals, arsenic, antimony, and tin, form a sub-group, as their sulphides are soluble in (NH 4 ) 2 S 2 , whilst the sulphides of the remaining metals are insoluble in that reagent. Group III. (Iron Group.) Group reagent (NH 4 ) 2 S in presence of (NII 4 )C1 and (NH 4 )HO. Metals whose sulphides and hydroxides are insoluble in water, but decomposed by dilute acids, are precipitated from neutral solutions by the third group reagent, ammo- nium sulphide. Aluminium, and chromium are precipi- tated as hydrates ; the others as sulphides. Iron, nickel, cobalt, zinc, manganese, as sulphides. Group IV. (Barium Group.) Group reagent (NH 4 ) 2 CO 3 in presence of (NH 4 )HO and (NH 4 )C1. Metals whose carbonates are insoluble in water, and are precipitated from their solutions by the fourth group rea- gent, ammonium carbonate ; barium, strontium, calcium. Group V. (Potassium Group.) Metals not precipitated by any of the above group rea- gents, as their chlorides, sulphides and carbonates are 42 ELEMENTARY QUALITATIVE ANALYSIS. soluble in water. They are, therefore, distinguished by individual tests : magnesium, potassium, sodium, ammo- nium. 57. Each group reagent will precipitate the metals of preceding groups. The metals distinguished by being insoluble as chlorides (Group I.) are also insoluble as sulphides (with Groups II. and III.) and as carbonates (with Group IV.). The second group sulphides are pre- cipitated both from acid and from neutral solutions, though the third group sulphides are precipitated from neutral, but not from acid solutions, and second arid third group metals form insoluble carbonates, as well as those of Group IV. In the work of analysis, the first group metals may be worked with the second, but thereafter the metals found in each group must be completely removed before testing for the next group. After filtering out a group precipitate, the reagent which produced it should be again carefully applied, with the proper conditions, to the filtrate before testing it for the next group. The student should at first have several metallic salts given to him, and be asked merely to determine to which of the above groups each salt belongs. He ought next to make himself familiar with the indi- vidual tests for each metal which follows, and then proceed to the separations of the different metals. It will also be well for him to attempt to frame a table of separations for each group before consulting those given in the book. REACTIONS OF THE METALS OF THE SILVER GROUP. 43 r 58. Reactions of the Metals of the Silver Group (Group I.). Silver, Ag', 108. Solution for Reactions, AgNO 3 . 1. HC1 produces a white, curdy .precipitate of AgCl, insoluble in hot water, soluble in Rff 4 fi 6. Zn also precipitates copper solutions. 7. K 2 CrO 4 precipitates a brownish-red basic chromate, soluble in HNO 3 and in (NH 4 )H0. 8. Compounds of Cu, when heated in Bunsen flame, im- part a green color, especially after addition of AgCl. 9. Heated on charcoal with NaHCO 3 in reducing flame, yields brittle metallic globules of bright-red color, soluble in HNO 3 or concentrated H 2 SO 4 . Characteristic reactions, 3, 5, 6. Cadmium, Cd", 112. Solution for Reactions, CdN 2 O 6 . 1. H 2 S precipitates yellow CdS, soluble in HNO 3 , in- soluble in KHS, KCN, and (NH 4 ) 2 S. CdS is dissolved by hot dilute H 2 SO 4 . 2. KHO precipitates Cd(HO) 2 , insoluble in excess of reagent. REACTIONS OF THE METALS OF THE COPPER GROUP. 49 3. (NH 4 )HO precipitates Cd(HO) 2 ; soluble in excess of reagent. 4. Zii precipitates Cd in brilliant scales. 5. Heated on charcoal with NaHCO 3 in the reducing flame, yields a brown incrustation of CdO. Cd dissolves readily in HNO 3 . Characteristic reactions, 1, 5. 50 ELEMENTARY QUALITATIVE ANALYSIS. 1 ... I c - S v | 1 i m ( ^wft- Group A ) . nearly to dryness, ter. FILTRATE. Groups III., IV., i (NH 4 ) 2 S 2 forabou1 FILTRATE. !^ 1 111 ^ H | |*| . fi ^rH 3 ^^' p i ^ ^ 1 *g * 6 i: ^ O p g 1 M e II 'OB , ft 5 , tfl 8 1 & W 00 "1 g aj ,ej bjo i g 6^ o a; T3 5 o< I 1 S" 1 ^ ^ ,0 fcc ! 5 1 w P A ^ s 08 e ""O Q CD 2 ^ 0) , tf || W EJ 5 H 4 |rf 1 H "o M 8 f | 5 en 1 Q OD 1 1 1 K * 02" H 0> ? fH 2 r= -- t 2 ^ 2 * s s I ^ CO ft ^ P H GO 1 1 HH o" ^' i i 1 d "i" . QB & . K SI ^ 3 1 ! . W) 4^ r^ 03 ^ QQ a- ^cc 2. ^,~ M ^ ,_j ,_4 ^ ,_- H0 tg ^ ^ ^ C7 1 - o^ cr o' CT cr cr CT cr 1 cr a 1 CT "* . ! , . * L^ ^ ^ [ - ; - > IT era era era CfQ s> era J9 dT era ' > o CJ CfQ W >-. W W era c? era cp 2? cr; H 5 DP? H JQ ^^ GO G^ GO GO GO GO GC -V K w o 4 hH ^ oF' - tf C^ _ CC OQ crq ^v W J3 S W -*F ^ * &i * era * Mi 1 o| era" 1 GO Q '* O o P 059 g W td w c?P * P tt hH o o o "* jS ?" N N N N N N N cs ., 3 -_ 3 _ 3 _ 3 3 o ^ "*> t> o j> <-t "* *i * o ;> 6 b> ^ !> Q j- 1 J ^^ t> *^> ^Z" K>-^" ^^ ^ hj - ~^ ii * ^~" ' C^vT^ ^- ^ ^ ^T* "** 1 1 (D ^ n> "* W p * ' fD Jt. -P" 1 hj g g ^ K^ f? fe! ^ g taj g HH J3 _p 3 r^~* f"v 3 d M O * ^ g O o - 3 - B j-s bd F w SB F W td ja Q w ^ v: GO GO GO GO ? o d o ^ ^ Q o SB fa O M p <^ SB . Q w g g g g g pH era era era 72 ELEMENTARY QUALITATIVE ANALYSIS. REACTIONS OF THE ACIDS. s 83. Grouping of the Acids. The acids can be ap- proximately classified by means of certain group reagents. They are divided into two great classes : inorganic and organic acids. These are easily distinguished by the action of heat. Salts of inorganic acids, when heated to redness, are not charred ; salts of organic acids are at once charred, owing to decomposition and separation of carbon (with the ex- ception of acetic and formic acids). 84. Grouping of the Inorganic Acids. Group I. {Sulphuric Acid Group.) Group reagent, BaCl 2 in presence of HC1. Sulphuric acid, hydrofluo-silicic acid. The acids of this group are precipitated by BaCl 2 , and the precipitate is not dissolved on addition of HC1. Group II. {Phosphoric Acid Group.) Group reagent, BaCl 2 . Phosphoric, boric, hydrofluoric, carbonic, silicic, sul- phurous, arsenious, arsenic, iodic, chromic acids. The acids of this group are precipitated in neutral solutions by BaCl 2 . Group III. {Hydrochloric Acid Group.) Group reagent, AgNO 3 . Hydrochloric, hydrobromic, hydroiodic, hydrocyanic, and hydro sulphuric acids. The acids of this group are precipitated by AgNO 3 , and not by BaCl 2 . REACTIONS OF THE ACIDS. 73 Group IV. (Nitric Acid Group.) Nitric, chloric, and perchloric acids. These acids are not precipitated by any reagent, as all their salts are soluble in water. Reactions of the Inorganic Acids belonging to Group I. 85. Acids precipitated by BaCl 2 in presence of HC1. SULPHURIC ACID, HYDROFLUO-SILICIC ACID. Sulphuric Acid, H 2 SO 4 , 98. 1. BaCl 2 precipitates a white BaSO 4 , insoluble in HC1 or HNO 3 . In very dilute solutions the precipitation is not immediate, but on standing, the solution becomes clouded, and ultimately the precipitate subsides. 2. Pb(NO 3 ) 2 precipitates a heavy white PbSO 4 , soluble in NaHO, and in boiling HC1 (on allowing this solution to cool, PbCl 2 crystallizes out). 3. fused on charcoal with Na 2 CO s in the reducing flame of the blow-pipe, a sulphide is produced. If the fused mass be moistened with HC1, the odor of H 2 S is at once perceptible ; or if it be placed on a bright piece of silver and moistened with water, a black stain of Ag 2 S is produced. 86. Hydrofluo-silicic Acid, H 2 SiF 6 , 144. 1. BaCl 2 precipitates a crystalline BaSiF 6 , insoluble in HCL 2. KC1 precipitates a gelatinous K 2 SiF 6 . 3. Heated with H 2 SO 4 in a leaden cri cible covered with 74 ELEMENTARY QUALITATIVE ANALYSIS. a piece of glass, the latter will be etched by the evolved HF. .Reactions of the Acids belonging to Group II, . 87. Acids precipitated by BaCl 2 in neutral solutions. PHOSPHORIC, BOEIC, HYDROFLUORIC, CARBONIC, SILICIC, SULPHUROUS, ARSENIOUS, ARSENIC, IODIC, AND CHROMIC ACIDS. Phosphoric Acid, HsPO^, 98. 1. BaCl 2 precipitates a white BaHPO 4 , readily soluble in HNO 3 or HC1, but "with difficulty in NH 4 C1. 2. Mg 2 SO 4 , along with (NH 4 )HO and NH 4 C1, precipi- tates a white crystalline Mg(NH 4 )PO 4 + 6 H 2 O, insoluble in (NH 4 )HO, but soluble in HC1, HNO 3 , and acetic acid. In dilute solutions the precipitation does not take place till after the lapse of some time, but is promoted by stirring and gentle warming. 3. AgNO 3 precipitates a yellow Ag 8 PO 4 , soluble in HNO 3 , and also in (NH 4 )HO. 4. Lead acetate precipitates a white Pb 8 (PO 4 ) 2 , soluble in HNO 3 , but almost insoluble in acetic acid. 5. Fe 2 Cl 6 , in presence of excess of sodium acetate, pre- cipitates a yellowish FePO 4 , soluble in HC1, and in excess of Fe 2 Cl 6 , which must.be added drop by drop. 6. Ammonium molybdate produces in solutions acidi- fied by HNO 3 a yellow color, and then a precipitate ; this reaction is hastened by warming. 88. Boric Acid, B(HO) 8 , 62. 1. BaCl 2 precipitates a white Pa(BO 2 ) 2 , soluble in acids. REACTIONS OF THE ACIDS. 75 2. AgNOg produces in strong solutions a yellowish-white precipitate. In dilute solutions Ag 2 O is precipitated. 3. H 2 SO 4 or HC1, added to hot concentrated solutions of alkaline borates, on cooling, precipitates a crystalline B(HO),. 4. If alcohol containing free boric acid be kindled, it burns with a green flame, best seen on stirring the mix- ture. Borates may be examined in this way by first adding strong H 2 SO 4 , to liberate the B(HO) 3 . 5. If the solution of a borate be made distinctly acid with HC1, and turmeric paper dipped into it, the latter, on gentle warming, acquires a brown tint, which is turned blue by caustic soda. 89. Hydrofluoric Acid, HF, 20. 1. BaCl 2 precipitates a white BaF 2 , soluble in HC1, and sparingly in NH 4 C1. 2. Ca01 2 produces a gelatinous and almost transparent precipitate of CaF 2 , made more apparent on addition of (NH 4 )HO. The precipitate is very difficultly soluble in II Cl, even on boiling, and is nearly insoluble in acetic acid. 3. Heated with H 2 SO 4 , all fluorides are decomposed with evolution of HF, which is recognized by its power of etching glass. 4. Heated with a mixture of borax and HKSO 4 , on a loop of platinum wire in the non-luminous flame, BF 3 is produced, coloring the flame green. 90. Carbonic Acid, H 2 CO 3 , H 2 4 CO 2 . 1. BaCl 2 in neutral solutions precipitates a white BaCO 3 , soluble in acids with effervescence. 76 ELEMENTARY QUALITATIVE ANALYSIS. 2i Treated with dilute HC1, all carbonates at once evolve COz with effervescence, which turns lime water a milky white from the formation of CaCO 3 . 91. Silicic Acid, Si(HO) 4 , 96. 1. BaCl 2 precipitates a white SiBa 2 O 4 , which is decom- posed on addition of HC1, and Si(HO) 4 separates out as a gelatinous precipitate. 2. HC1, added drop by drop to a strong solution of a silicate, precipitates a gelatinous Si(HO) 4 ; but if added to a dilute solution or in large excess, no precipitate is obtained until the mixture has been evaporated to dryness and ignited, when SiO 2 separates out, and this is not re- dissolved on addition of HC1. 3. Fused with Na 2 CO 3 in a loop of platinum wire in the non-luminous gas-flame, effervescence occurs from the dis- engagement of CO 2 , and the bead is transparent on cool- ing, unless the Na 2 CO 3 be in excess. 4. Fused with microcosmic salt on a loop of platinum wire in the non-luminous gas-flame, solution does not take place, but the silica floats about on the bead undissolved. 92. Sulphurous Acid, H 2 SO 3 , 82. 1. BaCl 2 precipitates a white BaSO 3 , soluble in HC1. On addition of chlorine water, gives a white precipitate of BaSO 4 , the sulphite being oxidized to the sulphate. 2. AgNO 3 gives a white precipitate of AgSO 3 , dark- ened on heating. 3. Added to a mixture of Zn and HC1, H 2 S is pro- duced, and recognized by its smell and by its action on paper moistened with a solution of a lead salt, blacken- ing it. REACTIONS OF THE ACIDS. . 77 4. H 2 S decomposes pure H 2 SO 3 , with separation of sulphur. 5. H 2 SO 3 is decomposed by HC1, with evolution of SO 2 . Y 93. Arsenious Acid, H 3 AsO 3 , 126. 1. AgNO 3 gives in neutral solutions a yellow precipi- tate of Ag 3 AsO 3 , soluble in (NH 4 )HO. 2. MgSO 4 + (NH 4 )C1 + (NH 4 )HO give no precipitate. *A 3. H 2 S precipitates As 2 S 3 , yellow. 94. Arsenic Acid, H 3 AsO 4 , 142. 1. AgNO 3 gives in neutral solutions a light-brown pre- cipitate of AgAsO 4 . 2. MgSO 4 +(NH 4 )Cl + (NH 4 )HO give a white pre- cipitate of MgNH 4 AsO 4 . 3. H 2 S precipitates AsgSa, yellow. 95. lodic Acid, HIO 8 , 176. 1. BaCl 2 gives a white precipitate of BaI 2 O 6 , soluble in HN0 3 . 2. AgNO 3 precipitates white crystalline AgIO 3 , easily soluble in (NH 4 )HO, but sparingly so in HNO 3 . 3. SO 2 gives at first a precipitate of I, which is con- verted into HI on addition of excess of reagent. X 4. HIO 3 is decomposed by H 2 S, with formation of an iodide and separation of S. 5. lodate salts, on heating, are decomposed, oxygen being evolved. In some cases iodine is given off in violet vapors. 96. Chromic Acid, H 2 CrO 4 , 118.2. 1. BaCl 2 precipitates a yellow BaCrO 4 , soluble in HC1 and HNO 3 , but insoluble in acetic acid. 78 ELEMENT AKY QUALITATIVE ANALYSIS. 2. H 2 S in presence of HC1 reduces the solution to Cr 2 Cl 6 (green), with separation of S. In neutral solutions, Cr 2 (HO) is precipitated along with S. 3. SO 2 reduces solutions of chromates to the chromic salt, the color of which is green. Chromates are likewise reduced by zinc and a dilute acid, by oxalic acid and dilute sulphuric acid, by strong H 2 SO 4 , by strong HC1, "and by boiling the solution, acidified with HC1 or H 2 SO 4 , along with alcohol. /x 4. AgNO 3 precipitates a dark-red Ag 2 CrO 4 , soluble in HN0 3 and in (NH 4 )HO. 5. Lead acetate produces a bright-yellow precipitate of PbCrO 4 , soluble in NaHO, but soluble with difficulty in dilute HNO 3 . 6. H 2 CrO 4 is precipitated by (NH 4 ) 2 S as Cr 2 (HO) 6 . .Reactions of the Acids belonging to- Group III. 97. Acids precipitated by AftN^, an^ " n * ^y ^n ni - HYDROCHLORIC, HYDROBROMIC, HYDRIODIC, HYDRO- CYANIC, AND HYDROSULPHURIC ACIDS. Hydrochloric Acid, HC1, 36.5. 1. AgNO 3 precipitates a white curdy AgCl, which be- comes violet on exposure to light. The precipitate is in- soluble in HNO 3 , but soluble in (NH 4 )HO, in KCN, in Na 2 S 2 O 3 , and also to some extent in NaCl. 2. Heated with H 2 SO 4 and MnO 2 , chlorides yield chlorine gas, recognized by its smell, bleaching action, and green color. 3. Dry chlorides, when heated in a retort with H 2 SO 4 and K 2 Cr 2 O 7 , yield CrO 2 Cl 2 (chromium oxychloride), REACTIONS OF THE ACIDS. 79 which distils over into the receiver as a dark-red liquid, decomposed by addition of water or (NH 4 )HO, yielding a yellow solution, which, on addition of a lead salt, gives a yellow precipitate of PbCrO 4 . 98. Hydrolromic Acid, HBr, 81. 1. AgNO 3 precipitates a pale-yellow AgBr, insoluble in dilute HNO 3 , soluble in strong (NH 4 )HO, and readily in KCN and Na 2 S 2 O 3 . 2. TTpaf.prl wit.h_ T-TpS^X and MnQ 9 , bromides yield red vapors of Br, recognized by its powerful odor. 3. Heated in a retort with K 2 Cr 2 O 7 and H 2 SO 4 , dry bromides yield dark-red vapors, which condense in the receiver to a liquid of the same color, which consists of pure bromine, and is decolorized on adding excess of (NH 4 )HO. 99. Hydriodic Acid, HI, 128. AgNO 3 precipitates a pale-yellow Agl, insoluble in dilute HNO 3 , and very difficultly soluble in (NH 4 )HO, but readily in KCN and Na 2 S 2 O 3 . 2. Cuprous sulphate precipitates a dirty-white Cu 2 I 2 , which separates most completely if the solution be made slightly alkaline with Na 2 CO 3 . 3. KNO 2 produces no reaction in solutions of iodides until a few drops of HC1 or H 2 SO 4 are added, when iodine is at once liberated and colors the solution yellow. If a little starch solution be now added, a deep-blue coloration results from the formation of starch iodide. 4. Chlorine water (or the gas) liberates iodine from iodides, but excess of Cl causes the formation of IC1 8 , 80 ELEMENTARY QUALITATIVE ANALYSIS. which is colorless, and gives no blue coloration with starch solution. 5. Heated with MnO 2 and dilute H 2 SO 4 , violet vapors of iodine are obtained, which color starch paper blue. 1OO. Hydrocyanic Acid, HCN, 27. * 1. AgNO 3 precipitates a white AgCN, insoluble in HNO 8 , with difficulty in (NH 4 )HO, but readily in KCN and Na 2 S 2 O 8 . AgCN is decomposed on ignition, and metallic Ag remains; this serves to distinguish it from AgCl, which is not decomposed on ignition. 2. If a solution of FeSO 4 which has become oxidized by exposure to the air, be added to the solution of a cyanide made alkaline with NaHO, a bluish-green pre- cipitate is formed, which is a mixture of Prussian blue with the hydrated oxides of iron. On adding HC1, these last are dissolved, and the blue precipitate remains. 3. HC1 decomposes nearly all cyanides with evolution of HCN, recognized by its odor, resembling bitter almonds. 4. Hg(CN) 2 cannot be detected by the above methods. The dry substance is detected by igniting in a small tube, when cyanogen gas is evolved, or the solution is decom- posed by H 2 S and filtered from the HgS : the filtrate contains HCN. 1O1. Hydrosulphuric Acid {Sulphuretted Hydrogen}, H 2 S, 34. 1. AgNOg precipitates a black Ag 2 S, insoluble in dilute acids. 2. Lead acetate, even when highly dilute, precipitates black PbS. REACTIONS OF THE ACIDS. 3. Sodium nitro-pmsside, in presence of NaHO, produces a reddish-violet coloration, even in very dilute solutions. The color disappears in a short time. 4. HC1 or H 2 SO 4 decomposes most sulphides with evo- lution of H 2 S, recognized by its disagreeable odor and by its blackening paper moistened with solution of lead. Reactions of the Acids of Group IV. 1O2. Acids not precipitated by any reagent. Nitric Acid, HNO 3 , 63. 1. Nitrates when heated evolve oxygen, and in some cases nitrous vapors also. On fusing a nitrate and adding a fragment of charcoal, vivid deflagration occurs. 2. Free HNO 3 heated with Cu gives red fumes ; boiled with pieces of silk or wool turns them yellow. 3. If to a solution of a nitrate FeSO 4 and concentrated H 2 SO 4 be poured carefully into the test-tube, a dark ring will appear on top of the H 2 SO 4 , which will be violet, red, or dark-brown according to the quantity of HNO 3 present. The ring disappears on warming. 103. Chloric Acid, HC1O 3 , 84.5. 1. H 2 SO 4 decomposes chlorates with evolution of C1 2 O 4 , a greenish-yellow gas having a powerful odor. If heated, violent explosions occur ; the mixture ought therefore to be kept cold, and only very small quantities should be used. 2. When chlorates are heated, oxygen is evolved, and a metallic chloride remains, which may be dissolved in water, and precipitated as AgCl by AgNO 8 . 82 ELEMENTARY QUALITATIVE ANALYSIS. 3. Chlorates are reduced by SO 2 with liberation of chlorine or its oxides ; hence if the solution of a chlorate be colored blue with indigo solution, it is decolorized on adding H 2 SO 4 and solution of Na 2 SO 8 . (Distinction from perchlorates.) 4. HC1 decomposes chlorates with evolution of Cl and C1 2 O 4 , a mixture called euchlorine. 5. Heated with charcoal, chlorates deflagrate violently. 1O4. Perchloric Acid, HC1O 4 , 100.5. 1. H 2 SO 4 does not act upon perchlorates in the cold, and on heating, white fumes of HC1O 4 are given off, but no explosions occur. 2. KC1 in strong solutions precipitates a white KC1O 4 . 3. Indigo solution is not decolorized when added to perchlorates warmed with HC1, as euchlorine is not evolved. 4. Dry perchlorates evolve oxygen on heating. 5. Perchlorates are not reduced by SO 2 . REACTIONS OF THE ACIDS. 83 TABLE VIII. Detection of Inorganic Acids in Mixtures. (The following acids are found in examination for bases, which ought always to precede examination for acids : H 2 SO 3 , H 2 S 2 O 3 , H 2 C<> 3 , H 2 S. Si(HO) 4 , H 2 Cr0 4 , H 3 As0 3 , H 3 As0 4 .) I. Acids in Soluble Bodies. 1. Neutralize a portion of the solution with (NH 4 )HO, and add BaCl 2 (or Ba(NO 3 ) 2 if Ag, Hg 2 , or Pb be present) : precipitate indicates H 2 S0 4 , H 3 PO 4 , H 3 AsO 3 , H 3 As0 4 , Si(HO) 4 , H 2 CrO 4 , and large quantities of B(HO) 3 and HF.* To precipitate, add H 2 0, and then HC1: if a precipitate remain, H 2 S0 4 was present. 2. To another portion of the neutralized solution add AgN0 3 : a precip- itate indicates one or more of these acids ; i.e., (a) HC1, HBr, HI, HCn, H 4 Fe(Cn) 6 , H 3 Fe(Cn) 6 , H 2 S. (6) H 3 PO 4 , H 3 As0 4 , H 3 As0 3 , H 2 CrO 4 , Si(HO) 4 , B(HO) 3 .* To the precipitate add cold dilute HN0 3 . Acids under (a) are insol- uble; those under (6), soluble. Detection of Acids under (a). To a portion of the solution add starch paste and one drop of a solution of N./) 3 in H 2 SO 4 . Blue coloration indicates HI. Add chlorine water till the blue color disappears, and shake with chloroform. Reddish brown * Oxalic, citric, and tartaric acids will also be shown, if present. 84 ELEMENTARY QUALITIVE ANALYSIS. color indicates the presence of HBr. HC1 is detected in presence of the others by boiling down the solution to dryness and distilling residue with K 2 Cr 2 7 and H 2 SO 4 (see 97, 3). Detection of Acids under (6). Test separately for each acid by the methods already given. Separation of H 3 As0 3 , H 3 As0 4 , and H 3 P0 4 . Acidify solution with HC1, add Na 2 S0 3 , and heat until no smell of S0 2 is given off. Pass H 2 S through the hot solution, filter, and test for H 3 P0 4 with am- monium molybdate : yellow precipitate indicates H 3 P0 4 . Precipitate another portion with magnesia mixture, and test both pre- cipitate and filtrate for arsenic. Test for the other acids by the following reactions, given under each acid : For HCn, by test 3, 100. For H 2 S, by test 4, 101. For HN0 3 , by tests 2 and 2, 102. For HC1O 3 , by tests 1 and 2, 103. For B(HO) 3 , by tests 4 and 5, 88. For Si(HO) 4 , by tests 2 and 4, 96. For H 2 S0 3 , by test 3, 92, and smell of SO 2 on adding HC1. For CO 2 , by test 2, 90. APPENDIX. To determine the empirical formula of a substance from its percentage composition. RULE. Divide the percentage amount of each constituent by its corre- sponding atomic weight; then divide each quotient so found by the lowest number } and reduce them to their simplest ratios. EXAMPLES. 1. A body on analysis yielded the following percentage composition: Carbon, 27.273 % Oxygen, 72.727 100.000 Calculate its formula. The atomic weight of carbon is 12. The atomic weight of oxygen is 16. Then, C = = 2.2727 ; 1^ = 1^21 = 4.5454. 16 Simplest ratio between the carbon and oxygen is as 1 : 2 ; for 2.2727 : 4.5454 : : 1 : 2. Hence the formula is C0 2 . 94 APPENDIX. 2. A compound was found to have the following percentage composition : Nitrogen, 82.353 Hydrogen, 17.647 100.000 Calculate its formula. The atomic weight of nitrogen = 14, and of hydrogen = 1. v 82.353 JN = = 14 H = Hl= The simplest ratio between nitrogen and hydrogen is as 1 : 3 ; for 5.882 : 17.647 : : 1 : 3. The formula of the body therefore is NH 3 . 3. A compound of iron and oxygen has the following per- centage composition : Iron, 70.01 Oxygen, 29.99 100.00 Calculate its formula. d Atomic weight of iron 56.0, and of oxygen 16.0. 4. Deduce the formulae of the following substances : Nitrogen, 30.43 Oxygen, 69.57 100.00 APPENDIX. 95 5. Potassium, 28.73 Hydrogen, 0.73 Sulphur, 23.52 Oxygen, 47.02 100.00 6. Carbon, 20.00 Oxygen, 26.67 Sulphur, 53.33 100.00 Relative Density of Solids, Liquids, and Gases; Vapor Density. The specific gravity (sp. gr.), or relative density of a solid or liquid substance, is the ratio of its mass to the mass of an equal volume of some liquid taken as unity. The standard universally adopted is pure water at its maxi- mum density. The number expressing the relative density of a solid or liquid expresses how much heavier or lighter the substance -is than an equal volume of water at 4 C. The relative density of a solid is generally ascertained by the following formula, the body being first weighed in air and then in water at 4 C., and the weights carefully taken. Eel. dens. = - Weight of substance (W) Weight of equal vol. of water at 4 C. W where W = Weight of substance in water at 4 C. If the solid be lighter, volume for volume, than water, a sinker is attached, whose weight in water = x, and rel. dens. = cL 96 APPENDIX. The relative density of the substance lighter than water is then expressed by the formula where W" = weight of sinker and solid in water. If the relative density be required at t C., then Eel. dens. = ^ x rel. dens, of water at t C. The relative density of a liquid is commonly found by 1st. The specific gravity flask method. Let x = weight of flask empty, W= weight of flask filled with water at t C., W'= weight of flask filled with liquid under examination; then -nri _ x Eel. dens. = - x rel. dens, of water at t C. W x 2d. By weighing a solid of constant volume in water and then in the liquid. Let cc = weight of solid in air, W= weight in water at C., W'= weight in liquid; then g. _ TTTf Eel. dens. = - x rel. dens, of water at t C. x-W Density of Gases and Vapors. The specific gravity, or relative density of a gas or vapor, as has been shown on pp. 18 and 19, is the ratio of its mass to the mass of an equal volume of hydrogen, measured at the same temperature and pressure. By the density of gas or vapor, we mean the relative density. APPENDIX. 97 One liter of hydrogen gas at C., and 760 mm. barometric pressure at the sea-level weighs .0896 gram. The relative density of a gas is determined by weighing a known volume of the gas, and comparing it with the weight of an equal volume of hydrogen under like conditions of temper- ature and pressure (p. 21). The effusion method may also be used. "The rel. dens, varies directly as the square of the time of effusion of equal volumes." Graham's Law (see p. 25) may also be applied. EXAMPLES. 1. Calculate the relative density of a solid from the follow- ing data : Weight of substance in air, 2.4554 grams. Weight of substance in water, 2.0778 grams. 2. Determine the relative density of wood from the follow- ing data : Weight of wood in air, 4 grams. Weight of silver sinker in air, 10 grains. Weight of wood and sinker under water, 8.5 grams. Relative density of silver = 10.5. 3. A solid weighs in vacuo 100 grams, in water 85 grams, and in another liquid 88 grams. 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