ESSENTIALS OF MEDICAL CHEMISTRY WOODY THE ESSENTIALS OF Medical Chemistry AND URINALYSIS BY SAM E. WOODY, A. M., M. D., PROFESSOR OF CHEMISTRY AND PUBLIC HYGIENE, AND CLINICAL LECTURER ON DISEASES OF CHILDREN, IN THE KENTUCKY SCHOOL OF MEDICINE. REVISED, ENLARGED AND ILLUSTRATED. THIRD EDITION, PHILADELPHIA: P. BLAKISTON, SON & CO., xoi 2 Walnut Street. 1890. Copyright, 1890, by P. Blakiston, Son & Co. PRESS OF WM. F. FELL & CO. 1220 24 SANSOM STREET, PHILADELPHIA, PREFACE. The Third Edition treading so closely on the heels of its prede- cessors, assures the writer that his little book has found use in the hands of many medical students, and that his labor has lessened theirs. As long as the effort is made to crowd the whole science of medicine into a five months’ course, the hurried student must have such a book as this to present the essential facts, so that he need not wade through the more exhaustive text-books, or be compelled to take voluminous notes, which are unavoidably inaccurate and unsatisfactory. The selection of material and the plan of presentation is the outgrowth of the author’s experience as a general practitioner and as a teacher of medical chemistry for the past twelve years. The subjects treated are so numerous that the descriptions are necessarily brief, but the principles of the science and the application of the facts to medicine have been stated more fully. 1506 West Walnut Street, Louisville, Ky. CONTENTS. INTRODUCTION, 9 Table of Elementary Bodies, with their Symbols and Atomic Weights, 12 PAGE PART I.—INORGANIC CHEMISTRY, 15 Classification of the Elements, 15 I. Hydrogen and Oxygen Group, 16 Hydrogen, 16. Oxygen, 17. Compounds of Hydrogen with Oxygen—Water, 20. Table of Valence, 27. 11. The Chlorine Group, 28 Fluorine, Chlorine, Bromine, lodine, 28. 111. The Sulphur Group, 34 Sulphur, 35. IV. The Nitrogen Group 40 Nitrogen, 40. The Atmosphere, 41. Phosphorus, 46. Arsenic, 49. Antimony, 53. Bismuth, 55. V. The Carbon Group, 56 Carbon, 56. Silicon, 61. Tin, 62. Lead, 63. VI. Metals of the Alkalies, 65 Lithium, 65. Ammonium, 66. Sodium, 67. Potassium, 68. Caesium and Rubidium, 71. VII. Metals of the Alkaline Earths, 71 Magnesium, 71. Calcium, 73. Strontium, 75, Barium, 75. VIII. Metals of the Earths, 75 Boron, 76. Aluminium, 76. Cerium, 77. IX. The Zinc Group, 77 Zinc, 77. Cadmium, 79. X. The Iron Group, 79 Chromium, 79. Manganese, 80. Iron, 81. Cobalt, 85. Nickel, 85. CONTENTS. PAGE XI. The Copper Group, 85 Copper, 85. Mercury, 87. XII. The Silver Group, 91 Silver, 91. Gold, 92. Platinum, 93. Table to Determine the Metallic Radical of a Salt, 94 Table to Determine the Acidulous Radical of a Salt, 95 Table—The Solubility or Insolubility of Salts in Water, 96 PART lI.—ORGANIC CHEMISTRY, 97 Hydrocarbons, 98 Alcohols, 101 Ethers, 105 Aldehydes, 108 Organic Acids, 109 The Carbohydrates, 112 Glucosides, 115 Nitrogenous Bodies, 116 Alkaloids, XiB PART HI.—THE URINE, 121 Physical Properties, 122 Chemical Constituents, 126 Tests for Albumin, etc., 133 THE ESSENTIALS OF MEDICAL CHEMISTRY INTRODUCTION. “ Chemistry is that branch of science which treats of the composition of substances, their changes in composition and the laws governing such changes.” ( Webster.) The distinctive characteristic of chemical action is change in com- position.* A bar of iron is the same in composition, whether hot or cold, luminous or non-luminous, magnetized or unmagnetized. But when it undergoes chemical action a new substance is formed which, though it contains iron, is entirely different from it in composition and properties.! Matter is that of which the sensible universe is composed. It is indestructible. Substances may undergo many changes, assume a great variety of forms, and even become invisible gases. Yet in none of these changes and combinations can a particle of matter be created or destroyed.! All matter has weight. By balances in the * Experiment.—Heat pieces of platinum and magnesium wire. Note that while the platinum is unaltered the magnesium burns and is converted into a white powder of magnesium oxide. t Experiment.—Weigh a small porcelain crucible containing powdered iron. Heat it, and it ignites; when combustion is complete, weigh again, and note the increase of weight and that a new substance is formed, which, though it con- tains iron, is not iron. J Experiment.—Burn a piece of charcoal (carbon) in a jar of oxygen gas. (Fig. i.) It disappears and, so far as we can judge by the senses of sight and touch, is lost, for it has combined with the oxygen to form an invisible gas. Add lime-water and shake. The gas combines with the lime and forms a white precipitate, which, if gathered and weighed, would exactly represent, besides the lime, the charcoal burned and the oxygen required to burn it. ESSENTIALS OF CHEMISTRY, open air we get the apparent weight of a body; but to obtain the abso- lute weight it must be weighed in a vacuum where there is no air to buoy it up. (For measures of weight, see table at back of book.) But of most importance to the student of chem- istry is the specific weight or specific gravity, by which we mean the weight of a substance as compared to the weight of an equal volume of some other substance specified as a standard. The standard for solids and liquids is water; for gases, air or hydrogen.* Fig.i. Matter exists in one of three states, solid, liquid, or gaseous. In the solid state the particles are held together * The specific gravity of solids is determined on the principle of Archimedes: A body immersed in a liquid displaces its own volume, and loses weight equal to the weight of the liquid displaced. Therefore, the weight a body loses when weighed in water, is the weight of its own volume of water and the standard with which the weight of that body must be compared, For example, suppose A piece of iron weighs 150 grains. Suspended in water it weighs 130 grains, Loss (or weight of its volume of water) ... 20 grains, Specific gravity of the iron (150-f-20) . . . 7.5 In case the body is lighter than water, a sinker is attached and the same method pursued, except that the loss of weight of the sinker is also obtained separately and subtracted from the total loss to ascertain the loss of weight of the lighter body. A body soluble in water may be weighed in some liquid of known specific gravity in which it is insoluble; e.g., suppose a lump of sugar weighs 100 grains, and in turpentine 45.62 grains. Loss 100 45.62 = 54.38 grains. 100 - 54.38 = 1.84 as the sp. gr. referred to turpentine. Multiply this by .87, the sp. gr. of the turpentine, and we get 1.6 as the true sp. gr. of sugar. For liquids we use the specific- gravity fiask (Fig. 2), which is made and marked to contain a certain weight of water. Fill the flask with the liquid to be in- vestigated and weigh it. Deduct the weight of the flask and divide this result by the weight of water the flask will hold. In practice the hydrometer is generally used. This is a hollow glass float with a graduated neck at its upper end, which indicates the specific, gravity by the depth to which it sinks in the liquid. The urinometer [the illustration of the urinometer is in Part 111, on the urine] is a hydrometer whose scale is constructed to measure the specific gravity of urine. For very accurate measurements of specific gravity the liquids must be at the standard temperature, which in this country is 6o° F. Fig. 2. INTRODUCTION. 11 so rigidly as to give the body a definite shape; while in the liquid state the attraction is so slight as to allow the particles to move freely upon each other and the body to take the shape of the vessel that contains it. In the gaseous state the mutual attraction of the parti- cles is entirely overcome, and their distance from each other depends upon the pressure to which the gas is subjected. The term fluid is applied to anything capable of flowing, whether liquid or gaseous. It is highly probable that all substances, which are not decomposed by heat or cold, are capable of existing in all three states. Heat is absorbed and cold produced wherever the particles are to be driven farther apart, as in the passage of a substance from the solid to the liquid or from the liquid to the gaseous state. When the two solids, ice and common salt are mixed, they form a liquid, and great cold is produced.* Perspiration in evaporating assumes the gaseous state, and absorbs in the change so much heat that the body is kept at its normal temperature in spite of the hottest weather.f On the other hand, when a substance passes from a rarer to a denser state it gives out again the heat absorbed in its passage in the opposite direction. If we examine the infinite variety of substances upon our earth we find most of them are compounds, i. e., they can be decomposed into two or more other substances, distinct in their properties from the sub- stance from which they were derived and from each other. There are some substances which have never been decomposed. These are called elements. Only seventy elements are at present known ; but, as our methods of investigation improve, this number may be increased by the discovery of other elements, or decreased by decomposing some of those now considered elements. Only about one-half of these enter into the materia medica, and will be noticed in this work. * Experiment.—Fold tin-foil into the shape of a little dish ; add powdered ice and salt. Spill water on the table and set the dish in it. Note how quickly it is frozen fast to the table. f Experiment.—Put a little water in a similar dish. Against the sides and bottom throw a spray of ether. Note that the evaporation of the ether is so rapid that the water is quickly frozen. 12 ESSENTIALS OF CHEMISTRY. TABLE OF ELEMENTARY BODIES, WITH THEIR SYMBOLS AND .ATOMIC -WEIGHTS. ( The more important are printed in capitals.) Name. Symbol Atomic Weight. Name. Symbol Atomic Weight. Aluminium, .... Antimony (Stibium), A1 27 Molybdenum, .... Mo 96 Sb 120 Nickel, Ni 58 Arsenic, As 75 Niobium, Nb 94 Barium, Ba 137 Nitrogen, .... N 14 Beryllium, Be 9 Norwegium, .... Ng 214 Bismuth, Bi 208 Osmium, Os 198 Boron, Bromine, B 11 Oxygen, 0 16 Br 80 Palladium, Pd 106 Cadmium Cd 112 Phosphorus, .... P 31 Caesium, Cs 133 Platinum, Pt 194.4 Calcium, Ca 40 Potassium (Kalium), K 39-1 Carbon, Cerium, C 12 Rhodium, Rh 104 Ce 141 Rubidium, Rb 85 Chlorine, Cl 35-5 Ruthenium, Ru 104 Chromium, .... Cr 52 Samarium, Sm 150 Cobalt, Co 59 Scandium, Sc 44 Copper (Cuprum), . . Cu 634 Selenium, Se 79 Didymium, D US Silicon, Si 28 Erbium, Fluorine, E 166 Silver (Argentum), . Ag 108 F 19 Sodium (Natrium), . Na 23 Gallium, Ga 70 Strontium, .... Sr 87.5 Germanium, .... Ge 163 Sulphur, S 32 Gold (Aurum), . . . Au 197 Tantalum, Ta 182 Hydrogen, H 1 Tellurium, Te 128 Indium, In ”3-4 Thallium, T1 204 Iodine, I 127 Thorinum, Th 231 Iridium, Ir 192 Tin (Stannum), . . . Sn 118 Iron (Ferrum), . , . Fe 56 Titanium, Ti SO Lanthanum, .... La *39 Tungsten, or Wolfram, W 184 Lead (Plumbum), . . Pb 207 Uranium, U 240 Lithium, Li 7 Vanadium, V 51.2 Magnesium, .... Mg 24 Ytterbium, Yb 173 Manganese, .... Mercury (Hydrargy- Mn 54 Yttrium, Zinc, Zirconium, Y Zn 90 65 rum), Hg 200 Zr 89.6 To explain the laws governing chemical phenomena we adopt the old atomic theory,* * Democritus, 460 b. c., said: “The atoms are invisible by reason of their smallness; indivisible by reason of their solidity; impenetrable and unalterable.” INTRODUCTION. 13 We will take up the theories and laws, not in the order of their enunciation, but of their natural sequence. It is assumed that matter is composed ultimately of infinitely small particles called atoms; that each element is composed of atoms, all of a certain size, weight, etc. Atoms do not exist alone, but in groups called molecules. In an element the molecule is composed of pairs of atoms of the same kind ; in compounds they consist of two or more atoms of different kinds. It has been determined that equal volumes of all substances in the gaseous state, and under like conditions, con- tain the same number of molecules. So a gallon of hydrogen gas and one of oxygen gas containing the same number of molecules, and those molecules consisting of pairs of atoms, must contain the same number of atoms. Furthermore, it is found that the gallon of oxygen is sixteen times as heavy as the gallon of hydrogen. So each oxygen atom must also be sixteen times as heavy as the hydrogen atom. Hydrogen being the lightest substance known, its atomic weight vs, taken as 1, and consequently the atomic weight of oxygen is 16. The atomic weights of other elements are determined in a similar way. By “atomic weight” is not meant the absolute weight of atoms (for that is not known), but the weight of the atom as com- pared to the hydrogen atom. The atomic weight of carbon is 12. If carbon combines with oxygen, atom for atom, the new substance (CO) resulting from that action will consist of molecules, in each of which the carbon will weigh 12 and the oxygen 16, and, as the whole mass is composed of these molecules, the same proportion obtains throughout the new compound. So 12 is found to be the combining weight of carbon, and 16 of oxygen. If, however, the combination should occur in the proportion of one atom of carbon to two atoms of oxygen, then each molecule must consist of 12 by weight of carbon to 32 of oxygen, and that must be the proportion throughout the entire substance. Between these two compounds no intermediate one can occur, for the carbon atom must take one or two, or more, oxygen atoms. It cannot take a fraction of one, for atoms are indivisible. Hence, we deduce the following Law : Substances combine in certain fixed pro- portions (atomic weights), or in multiples of these proportions. Symbols are abbreviations of the names of the elements. They consist of the initial letter of the Latin name; but if the names of several elements begin with the same letter, the single-letter symbol 14 ESSENTIALS OF CHEMISTRY. is reserved for the most common element, and for the others another letter is added. Thus, we have nine elements whose names begin with C; the most common is carbon, whose symbol is C ; the others add other letters, as chlorine, Cl; cobalt, Co; copper (cuprum), Cu, etc. The symbol indicates just one atom. When more than one atom is to be represented, the number is written just after and below the symbol, thus, C 4. Formula are to molecules what symbols are to elements. They indicate the kind and number of .atoms composing the molecule. When more than one molecule is to be indicated, the number is placed in front of the formula, thus, 5H20. A parenthesis inclos- ing several symbols or formulas should be treated as a single symbol, thus, 2(NH4)2C03 = N4H16QA. An equation is a combination of formulae and algebraic signs to indicate a chemical reaction and its results. As no matter is ever lost or created in a reaction, the number of each kind of atoms before the equality sign must be the same as after it. PART I.—INORGANIC CHEMISTRY. Classification of the Elements.—The elements are usually divided into two great classes: (a) Metals, about fifty-five in number, possessing a peculiar lustre, good conductors of heat and electricity, and whose oxides when combined with water form bases ; (b) Non- metals, about fifteen in number, possessing little lustre, poor conduct- ors of heat and electricity, and whose oxides combine with water to form acids. A better classification, and the one we shall adopt, is the following, based upon chemical properties : I. The Hydrogen and Oxygen Group. 11. The Chlorine Group : Fluorine, Chlorine, Bromine, lodine. 111. The Sulphur Group : (Oxygen) Sulphur, Selenium, Tellu- rium. IV. The Nitrogen Group; Nitrogen, Phosphorus, Arsenic, An- timony, Bismuth. V. The Carbon Group : Carbon, Silicon, Tin and Lead VI. The Alkaline Group : Lithium, Ammonium, Sodium, Potas- sium, Rubidium and Caesium. VII. The Alkaline Earths Group; Magnesium, Calcium Strontium, Barium. VIII. The Earths Group ; Boron, Aluminium, Lanthanum, Ce- rium, Didymium, etc. IX. The Zinc Group; Zinc, Cadmium. X. The Iron Group ; Chromium, Manganese, Iron, Cobalt, Nickel. XL The Copper Group: Copper, Mercury. XII. The Silver Group : Silver, Gold, Platinum. ESSENTIALS OF CHEMISTRY. I. Hydrogen and Oxygen Group. In a strict arrangement hydrogen would be placed in Group I of the metals and oxygen in the sulphur group. But we will consider them in a group to themselves, because (a) of all the elements hydrogen is taken as the standard for atomic weights, combining weights, valence, etc.; {!>) oxygen plays a most important role in chemistry, and its deportment with the other elements forms the basis of our classifi- cation ; (c) the chemistry of these two will best serve as an introduc- tion to the study of the other elements. Fig. 3. HYDROGEN (H —1) occurs free in volcanoes, gas wells, etc.; com- bined in water and all organized bodies. All acids are salts of hydrogen. Prepared in various ways from its compounds,* the most convenient being to displace it from sulphuric acid by zinc, thus— Physical Properties.—Transparent, colorless, odorless, tasteless gas ; the lightest substance known, fourteen and a half times as light as air ; hence used in balloons. Long suspected to be a metal, because it displaces metals in chemical compounds, forms alloys with certain H2S04 +Zn = ZnS04 + Ha. (Fig. 3.) * Experiment.—By means of a wire gauze spoon hold some sodium beneath the water and under a cylinder. The hydrogen gas liberated by the sodium from the water will rise in bubbles, fill the cylinder, and displace the water. PART I.—INORGANIC CHEMISTRY. 17 metals, and conducts electricity. This was proved in 1877, when Pictet condensed it under great cold and enormous pressure into a bluish metallic liquid. Chemical Properties.—Hydrogen does not support ordinary com- bustion or animal respiration. It burns in air with a pale but very hot flame.*f With pure oxygen it forms the oxyhydrogen flame. This is the hottest flame known, and a stick of lime held in it glows with dazzling brilliancy, forming the calcium or Drummond light. Mixed with air or oxygen, it explodes violently on contact with a spark. J Fig. 4. OXYGEN (O—16).—Sources. Most abundant of the elements, com- prising one-fifth of the air, eight-ninths of water, one-half of the crust of the earth, and three-fourths of all organized bodies. Prepared most easily by heating potassium chlorate (Fig. 4):— KCI03 = KCI + Oj. * Experiment.—lf an inverted jar of the gas is suddenly turned up, and a flame held a foot or two above, the gas escaping from the jar rises rapidly, and in coming in contact with the flame burns with a slight explosion. f Experiment.—If a jar of the gas be held mouth down and a candle be passed up into it, the gas ignites and burns quietly at the open end, while the candle passed up into the gas is extinguished, but may be relighted again by the burning gas as it is withdrawn. J Experiment.—Fill a bladder or rubber bag with two parts of hydrogen and one of oxygen or five of air. Attach a tube and blow up soap bubbles in a basin. Touched with a flame, they explode. ESSENTIALS OF CHEMISTRY. If manganese dioxide (Mno2) be mixed with the chlorate, the gas is liberated more quietly and at a lower temperature. The manga- nese dioxide is unaltered in the reaction. It seems to act by its mere presence, an influence called catalysis. Physical Properties.—Gas ; liquefied (Pictet, 1877) by great cold and intense pressure; colorless, odorless, tasteless; 1.10 times as heavy as air. Water dissolves only three volumes to the hundred, but enough to sustain aquatic life. Chemical Properties.—Intense affinities; combines with every element except fluorine. The product of its action is called an oxide, and the process oxidation. Oxidation so rapid as to produce heat and light is called combustion; if no light, slow combustion. Sub- stances that burn in air burn more brilliantly in oxygen,* and many substances that do not burn in air will burn in this gas.f By this property oxygen is usually recognized and distinguished from most other gases. Oxy- gen, especially in its diluted form (air), is the great supporter of combustion, for which its abundance and universal presence emi- nently fit it. Combustible and supporter of combustion are only relative terms. When a combustible substance burns in a sup- porter of combustion the union is mutual, one being as much a party to the action as the other. A jet of air Jor oxygen burns as readily in coal gas as a jet of coal gas burns in air or oxygen. The Fig. s. * Experiment.—A bit of phosphorus, dried by pressing between folds of blotting paper, is placed in a combustion spoon, ignited, and lowered into a jar of oxygen. The combustion is so intense that the phosphorus volatilizes, and its vapor burns throughout the jar with a brilliancy so dazzling that it is called the “ phosphorus sun.” f Experiment.—A watch-spring is wound into a spiral, tipped with a bit of tinder or a piece of yarn dipped in sulphur. This is lighted and lowered into a jar of oxygen. (Fig. 5.) The iron catches fire and burns with brilliant scintillations, globules of melted iron falling and melting into the glass, unless the bottom be covered with sand or water. ■{; Experiment.—Secure an ordinary lamp chimney (Fig. 6) and a wide cork to fit its lower end. Pass through the cork a narrow tube (a) connected by rubber hose with the house gas, and a wider one opening into the air. Turn FART I.—INORGANIC CHEMISTRY. 19 one in greatest abundance is usually called the supporter of combus- tion. Oxidizing agents are compounds in which oxygen is held so feebly it is readily given up to substances having greater affinity for it. Uses.—The process of respiration is a species of combustion, and, as oxygen is the best supporter of combustion, it is the best (and only) supporter of animal respiration. Admin- istered in capillary bronchitis, oedema glottidis, etc., when the patient cannot take in a volume of air sufficient to supply the requisite amount of oxygen. Fig. 6. Ozone.—If through a portion of air or oxygen electric sparks be passed,* a part of the oxygen will acquire a pungent odor and peculiar properties. The same change may be induced by various chemical pro- cesses, e.g. by mixing permanganate of potassium and sulphuric acid, or when phosphorus partially covered with water is exposed to the air. This modified oxygen is called ozone. It is one and a half times as heavy as ordinary oxygen, for its molecule contains three instead of two atoms. Very energetic, oxidizing substances unaffected by ordinary oxy- gen. Oxidizes potassium iodide with on the coal gas and light it as it issues from the tube. The cork with the flame (not too large) is then inserted into the chimney, where it continues to burn, sufficient air entering through the wide tube (r). Upon turning on more gas the air is crowded out and the chimney filled with coal gas. The gas flame disappears from the tube (a), and an air flame appears upon the tube (c) as the entering air burns in the atmosphere of coal gas. The excess of coal gas may also be lighted as it escapes, showing a gas flame above and an air flame within the chimney. On lessening the flow of gas the air will again be in excess, and the flame again appear on the narrow tube (a). In the gas flame above the lamp chimney (Fig. 6) heat some potassium chlorate in a combustion spoon until it melts and oxygen begins to bubble up. Then lower it into the atmos- phere of coal gas within the chimney. The escaping oxygen burns brilliantly, the coal gas being the supporter of the combustion. * Siemens’ apparatus for ozoning oxygen (Fig. 7) consists of two tubes, the inner surface of the inner and the outer surface of the outer tube being coated 20 ESSENTIALS OF CHEMISTRY. liberation of iodine, hence its test: paper dipped in a solution of potassium iodide and starch is colored blue in the presence of ozone.* Ozone is found in the air, especially after thunder-storms, and when present in considerable amount (as much as .00005 Per cent.) is apt to irritate the respiratory tract; but by oxidizing infecting germs, etc., it prevents the spread of infectious diseases. Compounds of Hydrogen with Oxygen.—Two are known— hydrogen oxide, or water, H2O; hydrogen peroxide* or oxygenated water, H202. Water (H2O) occurs widely distributed in nature; an important Fig. 7. constituent of all organized tissues ; forms seven-eighths of the human body. Physical Properties.—Transparent, colorless, odorless, tasteless— liquid. Below 320 F. (o° C.) it is a solid (ice), and above 2120 F. (ioo° C.) a vapor (steam or water gas). In solidifying, water expands ; so ice floats. The boiling-point is higher than 212° F. under increased pressure or when it contains solid matter in solution. with tin-foil, and each connected with the poles of an induction coil or Toepler- Holtz machine. A current of oxygen passing between these tubes may be ozonized to the extent of fifteen or twenty per cent. * Experiment.—Pour a little ether into a beaker, across the top of which is a glass rod supporting a strip of blue litmus paper and one of paper dipped PART I.—INORGANIC CHEMISTRY. 21 Water is the greatest of all solvents. The watery solution of a fixed substance is called a “ liquor,” and of a volatile substance an “ aqua." One body is said to dissolve m another when they coalesce and their particles intimately mingle. This is possible only in the liquid and gaseous states. When a substance dissolves it takes on the physical state of the solvent, e.g., a solid or gas dissolving in water becomes a liquid and then mixes with the water, the gas elevating the temperature and the solid lowering it. Heat assisting the liquefaction of a solid, and opposing that of a gas, hastens the solution of the one and retards that of the other. Most solid substances when separating from a solution take with them, as a necessary part of the crystal, a certain definite amount of water—water of crystallization. This water does not modify the chemical nature of the substance, but is necessary for maintaining the crystalline form. If the crystal loses Its water of crystallization by heat or exposure, it effloresces into an amorphous powder. Some substances when exposed absorb water from the air and deliquesce (melt down). Fig. 8. Chemical Properties.—The chemical composition of water may be proved by (synthesis) combining its constituents (H. -f O = H2O) *or by (analysis) passing the galvanic current through water until it is decomposed into its component gases (H«0 =H2 + O). f Neutral in reaction; combines with the oxides of the metals to form hydrates (bases), and with the oxides of the non-metals to form acids. Natural waters are never pure. The nature of the impurities in in potassium iodide and starch-water. Hold a hot glass rod in the jar (Fig. 8) ; the ether will undergo slow combustion, producing acid fumes which redden the litmus, and ozone which blues the other paper. * A mixture of two volumes of hydrogen and one of oxygen exploded in a eudiometer (Fig. 9), produces only water. f Fill the apparatus shown in Fig. 10 with water acidulated with sulphuric acid. Connect with a battery. The electricity passing through the water decomposes it into two volumes of hydrogen which collects in one tube and one volume of oxygen in the other. 22 ESSENTIALS OF CHEMISTRY. water depends on the condition of the atmosphere through which it has fallen as rain, and the nature of the geological strata through or over which it has passed, for water dissolves something from almost everything it touches. Good, potable (drinkable) water should be cool, clear and odorless. It should contain just enough dissolved gases and solids to give it an agreeable taste, neither flat, salty, nor sweetish; and should dissolve soap without forming a curd. Water impregnated Fig. 9. Fig. 10. with inorganic matters, especially salts of calcium, is called hard. A much more serious contamination is with organic (animal and vege- table) matters. Such water is a prolific source of diease. It is probable, in fact almost proven, that most infectious diseases are due to microorganisms, many of which find the most favorable conditions for their life and growth in water contaminated with organic, especially animal, matter. Though chemical analysis cannot detect the disease- producing elements, it can detect organic impurity, without which they PART I.—INORGANIC CHEMISTRY. 23 cannot exist. This is easily done thus: (1) Half fill a clean bottle with the water, warm, agitate, and critically smell it. A foul odor indicates organic impurity. (2) Fill a clean pint bottle three-fourths full, add a teaspoonful of the purest white sugar or gelatin ; set aside in a warm place for two days, when, if it becomes cloudy (bacteria), it is unfit to use. These rough-and-ready tests are those best suited to the practitioner, the more exact methods being practicable only to the chemist. To purify natural waters, they may be boiled to kill living organ- isms, and filtered to remove suspended matters; but for chemical Fig. ii. purposes, where great purity is desired, they are distilled* {aqua destillata, U. S. P.). Mineral waters are those possessing special therapeutic value They may be classed as follows: I, Carboiiated, those charged with carbonic acid. * When a liquid is rapidly vaporized, and the vapor, passing through a colder vessel, is recondensed, the process is called distillation (Fig. 11). If a solid be similarly treated it is called sublimation. When water containing solid matter in solution is distilled, the solids remain in the vessel, while the water passes over, enabling us to obtain perfectly pure water. When a mixture of two or more liquids is heated, the one having the lowest boiling-point distills first, leaving the others behind. This is called fractional distillation. 24 ESSENTIALS OF CHEMISTRY. 2. Sulphur, containing H2S or some soluble sulphide. 3. Alkaline, containing alkaline salts of potassium, sodium, or lithium. 4. Saline, containing neutral salts. 5. Chalybeate, containing iron. 6. Thermal, or hot waters. Hydrogen Dioxide—Oxygenated Water (ILCk).—Prepared most easily diluted by passing C02 through water holding barium dioxide in suspension. Ba02 + C02 + H2O = BaC03 + H202. The BaC03 may be allowed to subside and the clear solution poured off. Properties.—When concentrated it is a colorless, syrupy liquid, with a pungent odor and taste—prone to decompose into H2O -j- O. Used to bleach* the hair and skin, converting brunettes into blondes; as a disinfectant to ulcers, ozaena, and in diphtheria, espe- cially when the membrane has invaded the nose; also as a test for pus in urine, with which it causes an effervescence. The so-called “ ozonized ether ” used in the guaiacum test for blood is a mixture of hydrogen peroxide and ether. RADICALS.—Every molecule is composed of two parts, called radicals, held together by chemical affinity. Both radicals may be elements, as in H Cl, or one may be elementary and the other compound, as H N03, or both compound, as NH4 N03. Some compound radicals can be isolated, e.g. by heat, Hg CN = Hg -f- CN. Others decompose whenever set free. Whenever a gal- vanic current is passed through a compound, the chemical affinity is overcome by the electricity, and the molecule separates into its two radicals, one of which goes to the positive and the other to the negative pole.f Unlike electrical conditions attract, so the radical * Experiment.—Secure an old oil painting darkened with age, or take paper dipped in lead acetate and blackened by hydrogen sulphide : wash it with hydrogen dioxide, and the dark stain will be made white by the lead sul- phide being oxidized into sulphate. f Experiment.—lnto a jar put some water; add solutions of red litmus, potassium iodide, and boiled starch; connect with the galvanic battery. The electric current decomposes the potassium iodide into iodine, which gathers at the positive pole, producing a blue color, with the starch, and potassium at the negative, where it produces alkali, turning the red litmus blue. PART I.—INORGANIC CHEMISTRY. 25 going to the negative pole must be electro-positive, and the one going to the positive pole electro-negative. The metallic radicals are usually electro-positive and the non-metallic electro-negative. Some radicals are more intensely electro-negative or electro-positive than others. In the following list the more common elements are so arranged that each is usually positive to those following it and nega- tive to those preceding : Positive end: -j- K, Na, Mg, Zn, Fe, Al, Pb, Sn, Bi, Cu, Ag, Hg, Pt, An, H, Sb, As, C, P, S, N, I, Br, Cl, F, O—Negative end. A radical is electro-positive or electro-negative only in its relation to other radicals; for, while Cis positive to O, it is negative to K. In formulae the electro-positive radical is written first and the electro-negative next. The greater the difference between the electrical condition of two radicals, the greater the energy with which they unite and the more stable the product, and, vice versa ; e.g., O has a strong affinity for K, a weak one for Cl, and will not unite with F under any circum- stances. An idea once prevailed that the relations of affinities were fixed and constant between the same substances, and great pains were taken to construct tables similar to the above to show what was called the “ precedence of chemical affinities.” These tables showed the order of affinities for the circumstances under which the experi- ments were made, and nothing else. The circumstances attending chemical reactions are so complicated that in the greater number of cases it is impossible to predict the pre- cedence of affinities and the result of an untried experiment. Among these modifying causes may be mentioned : 1. Temperature, changes of which often reverse chemical affinities. Moderately heated, mercury and oxygen will readily combine, but when highly heated their affinity is destroyed, and they will refuse to unite, or, if already combined, will separate. Ordinarily free carbon has no affinity for oxygen, but at high tem- peratures it surpasses all other elements in its greediness for that substance, even taking it from a metal so extremely electro-positive as potassium. 2. Volatility.— Whenever i7i a mixture of two or more substances it is possible, by a rearrangement of the radicals, to form a compound 26 ESSENTIALS OF CHEMISTRY. volatile at the temperature of the experiment, such rearrangement will occur and the volatile compound be formed. For example : Fe S+ H2 SO, = Fe SO, +H2 S; or, 2 NH, Cl +Ca C03 = (NH,)2CO3 +Ca Cl2; or, H3803 -f 3NaCI = 3H Cl -f Na3 B03. 3. Insolubility.— Whenever, on mixing two or more substances in solution, it is possible, by rearrangement of the radicals, to form an insoluble compound, that rearrangement will occur and the insoluble compound be formed as a precipitate. For example:— Ca Cl2 + (NH4)2 C03 =Ca COs + 2 NH4 Cl. It is especially important to remember this law, for its application in tests, incompatibilities, and antidotes. 4. Nascent State.—Ordinarily the atoms of an element are grouped in pairs, and hence somewhat indifferent to the attractions of other atoms ; but just as they are being liberated (born) from a compound they are alone. Each atom, having no fellow, readily enters into combination with any atom it meets. This state is called nascent {nasci, to be born). 5. Catalysis.—This is the name given to the unexplained influence exerted by some substances of inducing chemical reactions between other substances without itself undergoing any change. The valence of a radical is its combining value, or its value in ex- change for other radicals * Here again hydrogen is taken as the standard. A radical that combines with or takes the place of one atom of hydrogen is said to be univalent (one valued); of two atoms, bivalent; three, trivalent; four, quadrivalent; five, quinquivalent; six, sexivalent. The valence is indicated by a Roman numeral just above and after the radical, thus: (NH,1), Ca11, (PO,)m, SiIV, Asv, SVI. The two radicals of every saturated compound must possess an equal number of unsatisfied valences. Hence, In HCI the radical Cl is equivalent to 1 atom of hydrogen; In H2O the radical Ois equivalent to 2 atoms of hydrogen In NHj the radical N is equivalent to 3 atoms of hydrogen ; In CH, the radical C is equivalent to 4 atoms of hydrogen. * The student should bear in mind that valence has nothing to do with the combining weight or the chemical activity of an element. PART I.—INORGANIC CHEMISTRY. 27 Therefore Cl is univalent, O bivalent, N trivalent, and C quadriva- lent. The same regard for valence is observed when radicals are made to displace each other, thus : H^SO*)11 requires two atoms of K1 or one of Zn11 to form KI2(S04)11 or Znn(SOi)n. Some elements exercise more than one valence : e.g., mercury may be univalent, as in Hgl, or bivalent, as in Hgl2; or iron may be biva- lent, as in FeCL, or the double atom (Fe2) sexivalent, as in Fe2Cl0. The termination “ous” is given to those compounds in which the positive element exercises its lower valence, and “ ic ” to those in which the higher valence is exercised, as, FeCl2, ferrous chloride; and Fe2ClG, ferric chloride. In the following table the most commonly occurring simple or ele- mentary radicals are arranged according to their valence ; I II Ill IV V VI F, Cl Ba, Sr A1 C, Si * Br, I Ca, Mg Au Pt H, Ag Cd, Zn Bo O. '. K, Na Pb, Sn Pb, Sn (NHJ, Li .. S, Se S, Se Fe2, Cr2 Mn2, Co2 Ni, Fe, Cr Mn, Co Ni •> N, P N, P Cu, Hg Cu, Hg Bi, Sb, As.. Bi, Sb, As.. Table of Valence. The next table shows the valences, together with the symbols and formulae, of the most common electro-negative (acidulous) radicals : Cl is the negative radical of all chlorides. Br is the negative radical of all bromides. I is the negative radical of all iodides. CN is the negative radical of all cyanides. HO is the negative radical of all hydrates. N03 is the negative radical of all nitrates. C 103 is the negative radical of all chlorates. C 2H30.2 is the negative radical of all acetates (Ac.). Univalent Radicals. 28 ESSENTIALS OF CHEMISTRY. 0 is the negative radical of all oxides. S is the negative radical of all sulphides. 503 is the negative radical of all sulphites. 504 is the negative radical of all sulphates. C03 is the negative radical of all carbonates. C 204 is the negative radical of all oxalates (Ox.). C 4H406 is the negative radical of all tartrates (T.). Bivalent Radical: Tri valent Radicals. C 6H507 is the negative radical of all citrates (Cit.). P04 is the negative radical of all phosphates. B03 is the negative radical of all borates. The student should learn these tables thoroughly, for with them he can easily know the formulae of all the principal inorganic and organic compounds. Fig. 12, 11. The Chlorine Group. Name. Derivation of Name. Symbol. At. Wt. Fluorine, Fluor spar, F, 19 Chlorine, %'aupog, green, Cl, 35.5 Bromine, Bpupog, stink, Br, . 80 lodine, Iviolet, I, 127 The members of this group are all univalent and much alike in their sources and physical and chemical properties. They differ in degree rather than in kind, forming a graded series. Hence we will consider them all together. PART I.—INORGANIC CHEMISTRY. 29 Sources.—Never free in nature. The principal source of fluorine is fluor spar (CaF2), while compounds of chlorine, bromine and iodine are abundant in sea and other salt waters. Preparation.—Free fluorine is obtained only with great difficulty; the others may be prepared by removing the hydrogen from their hydrogen salts (hydracids) by means of oxygen derived from manga- nese dioxide, thus:— 4HCI + Mn02 = MnCl2 + 2H20 + CI2 * 4HBr -j- Mn02 = Mnßr2 -f 2H20 + 4HI + Mn02 = Mnl2 + 2H20 + 12.I2. Physical Properties.—Fluorine is a colorless gas, with properties resembling chlorine, but more intense. Chlorine is a very irritating yellowish-green gas, two and a half times as heavy as air, slightly soluble in water (three volumes), forming “Aqua chlori, U. S. ” Bromine is a red liquid, giving off red vapors of a disagreeable, irritating odor; very slightly soluble in water. lodine is a solid, in bluish-gray scales, which, when warmed, give off violet vapors; insoluble in water except by the intervention-of an alkaline iodide; f soluble in alcohol; irritating, even caustic. Chemical Properties.—lntensely electro-negative; great affinity for the metals, J especially hydrogen. \ In negativeness, and conse- quently in affinity for the metals, F is greatest, Cl next, Br next, and I least. Therefore, in compounds with the metals, F will displace * Experiment.—lnto a flask standing in a cup of sand over a heater (Fig. 12) pour several ounces of hydrochloric acid and half as much black oxide of manganese, and agitate. The gas passes out; and, being heavier than air, collects in the bottle, where its yellowish-green color makes it visible. f Experiment.—To some water in a test-tube add a few scales of iodine; it does not dissolve. Now add a crystal of. potassium iodide; it dissolves easily. J Experiment.—Into a jar of chlorine introduce some copper or bronze foil, or sprinkle some powdered antimony. They inflame spontaneously. § Experiment.— (a) Into a jar of chlorine lower a lighted candle. The hydrogen of the tallow burns in the chlorine to form hydrochloric acid, and all the carbon is liberated as smoke. (I) Into a similar jar thrust a piece of paper dipped in warm turpentine. It inflames spontaneously and burns, evolv- ing dense clouds of smoke. 30 ESSENTIALS OF CHEMISTRY. Cl, and Cl will displace Br, and either F, Cl, or Br will displace I.* These elements destroy coloring matters and noxious effluvia in two ways: (i) by abstracting their hydrogen; (2) by abstracting the hydrogen of water, setting free nascent oxygen, which oxidizes the matters in question.! Medical.—Chlorine gas and bromine vapor are used for disinfec- tion. Inhaled they cause severe coryza and bronchitis. Taken into the stomach, bromine and iodine cause gastro-enteritis. The antidote is boiled starch. Locally bromine is used as an escharotic and iodine as a counter-irritant. Pharmaceutical.—The following preparations are officinal: Tinc- tura lodi (,SJ-0j) ; and Liquor lodi Compositus (Lugol’s Solution) (lodine 3yj, potassium iodide 5155, and water Oj.) The so-called colorless tincture of iodine is made by adding ammonia-water to the tincture until it is decolorized by converting the iodine into ammo- nium iodide. Tests.—ln the free state chlorine and bromine may be known by their bleaching, color, odor, etc. lodine is recognized by the blue color it strikes with starch. Note.—Acids.—All acids have, as their (positive) basylous radical, hydrogen, which may be replaced by metals to form salts. They may generally be recognized by a sour taste and the property of turning vegetable blues {e.g., litmus or purple cabbage) to reds. Acids whose acidulous (negative) radicals contain oxygen are called oxacids; those containing no oxygen, hydracids. The members of the chlo- rine group form both classes of acids. * Experiment.—Take two large test-tubes half full of water. Into one put a grain of potassium bromide, into the other potassium iodide; add chlorine- water to each. The chlorine will liberate the bromine in one and the iodine in the other. This may be shown (a) by their color; (b) by adding a few drops of carbon bisulphide or chloroform, which on agitation will gather all the free bromine and iodine, and be colored brown with the one and violet with the other; (r) add a few drops of starch-water, which will give brown with bromine and a deep blue with iodine. f Experiment.—(a) Into one bottle of chlorine gas introduce a piece of dry calico, into another a moist piece. The moist calico is rapidly bleached, while the dry is but slowly affected, [b) To a solution of indigo, cochineal, or some aniline color add chlorine-water. It is immediately decolorized. PART I.—INORGANIC CHEMISTRY. 31 The Hydracids of the chlorine group are as follows: H -f- F = HF—Hydrogen Fluoride*—Hydrofluoric acid. H -|- Cl = HCl—Hydrogen Chloride—Hydrochloric (muriatic) acid. H -)- Br HBr—Hydrogen Bromide—Hydrobromic acid. H I = Hl—Hydrogen lodide—Hydriodic acid. Fig. 13. Preparedhy treating the appropriate salt with H2S04, thus : CaF2 + H2S04 = CaS04 + 2HF. 2NaCI + H2S04 = Na2S04 + 2HCI.f 2KBr + H2S04 = K2S04 + 2HBr. 2KI + H2S04 = K2S04 + 2HI. * Binary compounds—i. e., those of only two elements—are named by call- ing first the name of the positive and then that of the negative radical, affixing to the latter the termination “ ide.” f Experiment.—To prepare hydrochloric acid gas, put several ounces of common salt and about twice as much sulphuric acid into a flask, and warm. The gas comes off in abundance and may be collected in a dry bottle (like chlorine, Fig. 12), or over mercury. The solution of the gas (the ordinary form) is obtained by passing the gas through a series of Wolff bottles contain- ing cold water and arranged as shown in Fig. 13. 32 ESSENTIALS OF CHEMISTRY. Physical Properties.—Colorless, irritating gases ; sharp, sour taste; very soluble, water dissolving several hundred times its own volume, forming aquce known by the simple name of the acid itself, thus : The officinal “hydrochloric acid” is a solution of the hydrochloric acid gas in water. Chemical Properties. Strong acids; true acids even without water. Uses.—HF attacks silicon energetically, hence is used to etch glass ; very poisonous, and burns made by it heal with difficulty. HCI is very useful in the arts. Aqua regia, or nitro-muriatic acid, is a mixture of nitric and hydrochloric acids. It is the only solvent of gold and platinum. The metals are attacked by the nascent chlorine evolved by the oxidation of the H of the HCI by the O of the HN03. In medicine HCI is often prescribed as a tonic. HBr, like all bromides, is a sedative. HI, like all iodides, is an alternative. Tests.—Fluoride -f- H2S04—etches glass.* Bromide -f- AgN03—yellowish-white precipitate, slightly soluble in ammonia. Chloride + AgN03—white precipitate, soluble in ammonia. lodide + AgN03—yellow precipitate, insoluble in ammonia. If to a bromide or iodide some chlorine-water and starch paste be added, the bromine and iodine will be liberated, the bromine striking a brown and the latter a blue color with the starch. Oxacids are formed by oxides of non-metals combining with water. The elements of the chlorine group, being very negative, have but little affinity for oxygen. lodine has most, bromine less, chlorine still less, and fluorine will not unite with oxygen at all. Chlorine, bromine, and iodine each forms a series of oxides per- fectly analogous, so we will notice only those of one—chlorine. The several oxides are distinguished by prefixes derived from the * Experiment.—On a plate of glass coated with wax or copper-plate var- nish (six parts of mastic, one of asphalt, and one of wax dissolved in turpen- tine) draw a design with a pointed instrument. Invert over a lead dish and warm gently. Hydrofluoric acid gas is evolved and attacks the glass wherever the wax has been scratched off. Upon removing the wax the design is found permanently etched on the glass. PART I.—INORGANIC CHEMISTRY. 33 Greek numerals indicating the number of oxygen atoms in the for- mula, thus:— C120—Chlorine Monoxide. Cl202 (?)—Chlorine Dioxide. C1203—Chlorine Trioxide. C1204—Chlorine Tetroxide. C120s—Chlorine Pentoxide. C1207—Chlorine Heptoxide. These oxides combining with water form the corresponding acids, thus:— C120 -f- H2O = 2HCIO —Hydrogen Hypochlorite—Hypochlorous acid. C1203 + H2O = 2HCI02—Hydrogen Chloride—Chlorous acid. C1205 -)- H2O = 2HCIO3—-Hydrogen Chlorate—Chloric acid. Cl2Oy H2O = 2HCI04—Hydrogen Perchlorate—Perchloric acid. Note.— The names of oxacids are derived from the negative element other than oxygen, and to this certain affixes and prefixes are added to indicate the degree of oxidation. The one containing more oxygen has the affix “■ic,” less oxygen, “ -ous.” If there is in the same series another acid with more oxygen than the “ -ic,” it is given the prefix 11 per-;” if less than the “ -ous,” the pre- fix “ hypo- ” (under). Acids ending in “-ic” form salts ending in “ -ate ;” those eliding in “-ous” form salts ending in “-ite.” The foregoing chlorine acids illustrate this. All these oxides, as well as their corresponding acids, are easily decomposed, sometimes with explosion ; hence much used as oxidiz- ing agents* and as explosive mixtures.! The most important of these salts is potassium chlorate, used in medicine and in the laboratory for the sake of its oxygen. * Experiment.—Their oxidizing action on combustibles may be shown by : (a) Mix together a drachm each of powdered potassium chlorate and sugar; place on a brick and touch off with a glass rod dipped in sulphuric acid. A vigorous combustion occurs. (I) Drop some crystals of potassium chlorate into a conical glass of water; add several bits of phosphorus; then by means of a pipette introduce sulphuric acid at the bottom of the glass. The phos- phorus takes fire and burns at the expense of the oxygen of the potassium chlorate. ! Experiment.—Mix on a sheet of paper a scruple of powdered potassium chlorate and five grains of some combustible powder, as sulphur, antimony sulphide, or tannin. Wrap it up in the paper, place upon an anvil, and strike with a hammer. It explodes violently. 34 ESSENTIALS OF CHEMISTRY. 111. Sulphur Group. Oxygen (already described), .0 16 Sulphur, . S 32 Selenium, Se 79.4 Tellurium, Te 128 The elements comprising this group are solid at ordinary tempera- tures ; bivalent and sexivalent; possess electro-negative affinities Fig. 14. which, as in other groups, decrease as the atomic weights increase; form hydracids as well as oxacids. The analogy between their compounds is shown in the following table;— Hydro-ic Acid. Dioxide. Trioxide. Hypo-ous Acid. -ous Acid. -ic Acid. H2S . . .so, . • S03 . . h2so2 . . H2S03 . . . H2S04. H2Se . . . Se02 . . Se03 . . H2Se03 . . . H2Se04. H2Te . . . Te02 . . TeOg . , . H2Te04. Selenium and Tellurium are of no medical interest, and will not be furthernoticed. PART I.—INORGANIC CHEMISTRY. 35 SULPHUR occurs free, especially in the neighborhood of volcanoes ; occurs combined as sulphides and sulphates in many valuable ores, and in small quantity in the animal and vegetable kingdoms. Preparation.—The native sulphur freed from stones is refined by distillation, as shown in Fig. 14. The crude sulphur is melted in the tank by the hot draft from the fire below, and then runs down through a pipe into the retort, where it is vaporized. This vapor, entering a large brick chamber, is condensed in fine, feathery crystals, called flowers of sulphur or sublimed sulphur. If the chamber be hot, it condenses into a liquid, which is drawn off and moulded into rolls, called roll brimstone. Sublimed sulphur is apt to contain more or less acid, and is washed (sulphur lotuin). Boiled with lime and precipi- tated with HCI, it forms sulphurprecipitatum, U. S. P. This mixed with water is milk of sulphur (lac sulphuris, U. S. P.). Physical Properties.—A brittle yellow solid; insoluble in water, hence, tasteless, etc. Chemical Properties.—lnflammable, hence called “brimstone” (burn-stone). Combines with metals,* forming sulphides.! Sulphur forms compounds remarkably analogous to those of oxygen, e.g. : H2O .... KHO .... C02 ... . H2C03 .... HCNO. H2S . . . . KHS .... CS2 ... . H2CS3 .... HCNS, Uses.—In the arts, to make gunpowder, matches, etc.; in medicine, as a laxative, parasiticide, and alterative. We have only theoretical explanations of the method of its absorption ; but that it is absorbed is certain, for persons taking it excrete enough to blacken silver car- ried on the person. Hydrogen Sulphide—H2S—Hydrosulphuric Acid or Sniphureted Hydrogeti—occurs in sewer gas and other effluvia from decomposing organic sulphurized matters, and in the water of sulphur springs. Prepared in laboratory by decomposing a sulphide, thus ; FeS + H2S04 = FeSO* + H2S. * Experiment,—ln a small glass flask, a little sulphur is heated to boiling. If now a bundle of fine copper wire or a piece of sodium, in a combustion spoon, be previously heated and then lowered into the vapor, it burns brilliantly. f Experiment.—Mix in a dish equal parts of iron filings and flowers of sulphur: moisten with water and set aside. Within a half hour it gets hot, vaporizes the water, and is converted into a black mass of FeS. 36 ESSENTIALS OF CHEMISTRY. Physical Properties.—Colorless gas, having the odor of rotten eggs or intestinal flatus; slightly soluble in water. Chemical Properties.—Very feeble acid ; burns with pale blue flame : H2S +O3 = S02 4- H2O * Forms characteristic precipitates with most metallic salts, hence a valuable test reagent.f Tests.—The presence of H2S even in minute quantities may be detected by its odor, and by its blackening paper moistened with a solution of lead acetate. Fig. 15. Preparation of H2S. Physiological.—When inhaled H2S is an active poison, combining with the hemoglobulin and destroying its oxygen-carrying power. Even when highly diluted, as in the atmosphere of city dwellings, * Experiment.—Burn the gas from a jet: [a) Hold near the flame a glass rod dipped in ammonia ; white crystals of ammonium sulphite are formed. (A) Hold a cold, dry bell glass over the flame; it is bedewed with water. f Experiment.—To show the action of Ii2S on metallic salts, connect several wash bottles with the generator A, as shown in Fig. 16. A dilute solu- tion of lead acetate is put in B, of tartar emetic (antimony) in C, of arsenic in D, of zinc sulphate in E. The gas in passing precipitates lead sulphide (black) in B, antimonious sulphide (orange) in C, arsenious sulphide (yellow) in D, zinc sulphide (white) in E. PART I.—INORGANIC CHEMISTRY. 37 clumsily “fitted with the modern conveniences,” it produces a low febrile condition. When concentrated, or even moderately diluted (one per cent, and over), the gas proves rapidly fatal. Carbon Disulphide—CS2.—Obtained by bringing S into contact with heated charcoal. A colorless, volatile liquid of a fetid odor, unless it is very pure. A valuable solvent for S, P, india-rubber, etc. Treatment.—Fresh air, artificial respiration, and stimulation. Fig. 16. Sulphur Oxides and Acids. Dioxide—S02 -|- II20 = H2S03—Sulphurous acid. Trioxide—S03 -f- H2O =• H2SO.4—Sulphuric acid. Sulphur Dioxide, S02, occurs whenever sulphur or any of its compounds are burned in air or oxygen. Prepared in laboratory by decomposing [sulphuric acid by copper or charcoal, thus:— 2H2S04 -f Cu = CuS04 + 2H20 + so2. 2H2S01 + C = 2S02 4- C02 + 2H20. Physical Properties. A colorless gas, with a suffocating odor 38 ESSENTIALS OF CHEMISTRY. (of burning matches); dissolves in water to form sulphurous acid (H2SO3). Chemical Properties.—Neither burns nor supports combustion; a strong deoxidizer; by removing O from coloring matters and infecting germs it bleaches* and disinfects. Uses.—Sulphur dioxide, sulphurous acid, and the sulphites possess the property of destroying microorganisms and arresting fermenta- tions. A sulphite digested with sulphur forms a hyposulphite, thus : Na2S03 + S = Na2S203. Sodium hyposulphite has the same uses as the sulphites, and is also a valuable solvent of the silver salts in photography. Sulphur Trioxide, S03.—Made by oxidizing S02 in the manu- facture of sulphuric acid. This is done upon a large scale by passing S02 from burning sulphur into a chamber kept filled with vapor of nitric acid, steam and air. f The nitric acid gives up a part of its oxygen to oxidize a portion of the S02 to S03. 2HN03 + 3502 = 3503 + H2O + N202. The S03 then combines with the water thus produced (SO3 -f- H2O = H2SO4), and more water is supplied by a jet of steam thrown con- stantly into the chamber. The N202 has the power of taking up oxygen from the air and be- coming N204, N202 +O2 = N204, which in turn parts with this oxygen to oxidize a new quantity of S02. N204 “h 2S02 N202 -f- 2S0s. Thus the process is kept up as long as the S02, air, steam, and N202 * Experiment.—Some sulphur is ignited beneath a tripod on which fresh flowers are placed, and the whole covered by a bell-glass. The flowers are bleached. The color may be restored by washing with some dilute alkali or acid that will combine with or displace the S02, or with very dilute nitric acid, which will restore the oxygen removed by the S02. t The manufacture of sulphuric acid may be illustrated on the lecture table by the apparatus shown in Fig. 17. The lead chamber is represented by a large flask. Into this are led (a) N202 from the flask on the right; {!>) S02 from a mixture of sulphur and manganese dioxide in the flask in the rear; (c) steam from the other flask, and (d) air or oxygen through the open tubes. PART I.—INORGANIC CHEMISTRY. 39 are supplied. The acid condenses with the water upon the floor of the chamber, and is drawn off, concentrated, and sold as Sulphuric Acid—H2S04—"Oil of Vitriol." Physical Properties.—A dense, colorless, oily-looking liquid, with- out odor. Chemical Properties.—Strong acid; very avid of water, not only dissolving in it, but combining with it, the act evolving considerable heat; chars organic matters by abstracting H Oto form water.* Fig. 17. Uses.—So important in the arts that the commercial prosperity of a country may be measured by the amount of H2S04 consumed. Prop- erly diluted, it is a refrigerant tonic, but concentrated it is a severe caustic. Tests.—(1) The concentrated acid, if placed on a piece of paper or * Experiment.—Pour strong sulphuric acid on an equal quantity of sugar or strong syrup; the sugar is dehydrated and a mass of carbon left. 40 ESSENTIALS OF CHEMISTRY. other organic material, will char it. If dilute, it will char the paper only after being warmed and concentrated by the evaporation of its water. (2) Sulphuric acid, or any other sulphate, will form with a solution of a barium salt a white precipitate (BaSO4) insoluble in nitric or hydrochloric acid. IV. Nitrogen Group. Nitrogen, ..* N 14 Phosphorus, P 31 Arsenic, As 75 Antimony (Stibium), Sb 122 Bismuth, Bi 210 Trivalent and Quinquivalent.—This group, as shown below, forms a graded series from nitrogen and the negative to bismuth at the positive end;— N P As Sb Bi 14 31 75 122 210 Sp.gr. 1.83, Sp. gr. 5.67. Sp. gr. 6-7- Sp. gr. 9.8. Gas, with full A soft solid. Solid. Hard solid. Very hard solid, negative ten- Non-volatil- dencies. Easily volatiliz- Volatilizable. Difficultly vola- izable. able. tllizable. Full metallic lus- Destitute of me- Some metallic Great metallic tre. tallic lustre. lustre. lustre. Full positive ten- Negative ten- Both negative More positive dencies. dencies. and positive tendencies, tendencies. The following will exhibit the relations of some of the most import- ant compounds;— Hydrides. Chlorides. Oxides. Sulphides. -ous. -ic. -ous. -ic. -ous. -ic. NH3 NCI,. . . N203, N2Os . . PH, PCI,. PC15 P,203, P205 P2S3, P2S5 AsPI3 AsC13, AsCl- As203, As205 As2S3, As2S3 SbH3 SbCl3, SbCl5 Sb203, Sb205 Sb2S3, Sb2S5 biC]3, . . Bi203, Bi203 • • • • NITROGEN occurs Uncombined in the atmosphere; combined in some mineral and in all vegetable and animal bodies, especially in the more highly organized tissues. Prepared most easily by burning phosphorus in a confined space PART I.—INORGANIC CHEMISTRY. 41 until the oxygen is removed from the air.* Prepared in this way it contains small quantities of other gases found in air. To prepare it pure, heat ammonium nitrite (NH4NO2 = 2H20 -f- N2). Physical Properties.—A colorless, tasteless, odorless gas, a little lighter than air. Chemical Properties.—Little tendency to combine with other ele ments, and its compounds, once formed, are very prone to decora- Fig. 18. pose, either with violent decomposition f or gradual putrefaction ; neither combustible nor a supporter of combustion ; negatively poison- The Atmosphere.—Air,J considered by the ancients one of the * Experiment.—A flat piece of cork floating on water supports a capsule containing a bit of phosphorus carefully dried. This is ignited and immedi- ately covered with a bell jar. The jar is filled with a dense white cloud from the combustion, which ceases only when the oxygen is all consumed. At first the air expands and some may be forced out. Upon cooling the water rises to take the place of the oxygen, and the white fumes gradually dissolve in the water, and the nitrogen is left clear and comparatively pure. f Experiment.—To tincture of iodine add excess of ammonia water. Filter to separate the precipitated iodide of nitrogen. Put portions of this on separate bits of paper and set aside. When dry they explode on the slightest touch. J Proofs that air is a mixture: (l) Its constituents are not in atomic pro- portions; (2) air can be made by mechanically mixing the gases; (3) solvents may remove one gas without affecting the others, each dissolving according to its own solubility. 42 ESSENTIALS OF CHEMISTRY. four elements, is neither an element nor a compound. It is a mixture mainly of nitrogen and oxygen, the function of the former being to dilute the latter. Miller gives the average composition of air as follows : Volumes. Nitrogen, 77-95 Oxygen, 20.61 Carbon dioxide, 04 Aqueous vapor, 1.40 Also traces of nitric acid, ammonia, sodium chloride, ozone, dust, bacteria, germs, etc. In the neighborhood of large cities various other substances are poured into the air from manufactories. Yet, owing to the*rapid diffusion of gases, the composition of the air is almost the same everywhere. Watery Vapor.—The higher the temperature the more water air will hold. A warm, dry air, when cooled, will appear damp, and the temperature at which it begins to deposit its water is its dew-point. A cold, damp air, when heated, becomes capable of holding more water, and appears dry; hence the necessity of supplying water to the heated air of our rooms in winter, especially in cases of bronchitis or catarrhal croup. Even in health a very dry air irritates the air-passages, pro- duces dryness of the skin and malaise; while a very moist atmos- phere retards evaporation from the skin and lungs, raises the body temperature, and becomes oppressive. Suspended matters in air are of a great variety of substances. The irritation of dust incident to certain trades may cause chronic bron- chitis, emphysema, and phthisis. Germs floating in the air are believed to be the cause of many contagious, infectious, and malarial diseases. The best disinfectants * are (a) free ventilation and consequent dilu- tion ; ([b) chlorine, bromine, iodine, and sulphur dioxide. Nitrogen Hydride—Ammonia, NH3. Occurs in the effluvia from decomposing nitrogenized organic bodies; for nitrogen and hydrogen unite only in the nascent state. (See page 26.) First obtained by distilling camel’s dung, near the temple of Jupiter Ammon in Libya ; hence called “ ammonia.” Obtained by heating clippings of hides, * Disinfectants destroy the power to infect, whether it be due to germs or other agent. Germicides destroy germs. Antiseptics prevent putrefaction. Antizymotics prevent fermentation. Deodorizers destroy offensive odors. PART I.—INORGANIC CHEMISTRY. 43 hoofs and horns,* especially of deer, in closed retorts (destructive distillation), it was called spirit of hartshorn. Coal contains about two per cent, of nitrogen, which in the manufacture of coal gas comes off as ammonia. In washing the coal gas the ammonia dissolves in the water. This aqua is its commercial source. Prepared in laboratory by driving the ammonia off from the aqua by means of heat. Fig. 19. Fig. 20. Physical Properties.—Transparent, colorless gas, of an irritating odor; condenses under a pressure of 6)4. atmospheres into a colorless liquid;f very soluble in water, which dissolves from 500 to 1000 times its own volume. J Administered by inhalation as a stimulant * Experiment.—Mix calcium, potassium, or sodium hydratfe with some nitrogenous substance, as albumin or clippings of horn, hoofs, flannel or lean meat. Heat in a test-tube. Ammonia gas is evolved, recognized by its odor, alkalinity, or by white fumes forming when a glass rod moistened with HCI is thrust into the tube. f Experiment.—Make ammonium silver chloride by passing ammonia gas over silver chloride. Enclose this in a bent tube (Fig. 20). The end contain- ing the compound is heated in a water bath, while the other is cooled in an ice mixture. Ammonia gas is driven off from the compound, and condenses into a colorless liquid in the cold end of the tube. % Experiment.—The absorption of ammonia gas by water may be illus- trated by filling a large bottle with the gas by upward displacement and closing the mouth with a rubber cork through which passes a glass tube sealed at the outer end. If this sealed end be plunged under water and then broken off, the water rushes in, forming a fountain (Fig. 19). If the water be colored with red litmus solution it will become blue as it enters the bottle, showing the alkalinity of the solution. 44 ESSENTIALS OF CHEMISTRY. in fainting fits, etc., but care must be taken, for its too liberal use may cause spasm of the glottis or induce a fatal bronchitis. Tests.—(i) Smell; (2) white fumes with HCI; (3) turns moistened red litmus paper blue. Nitrogen Oxides, Monoxide—N20 -j- H2O = 2HNO = Hyponitrous acid. Dioxide—N202. No corresponding acid. Trioxide—N203 -|- H2O = 2HN02 = Nitrous acid. Tetroxide—N204. No corresponding acid. Pentoxide—N205 4- H2O = 2HN03 = Nitric acid. Making N2O. Nitrogen Monoxide—N20 (Nitrous Oxide—Laughing Gas) Preparedly heating ammonium nitrate, as shown in Fig. 21. Physical Properties.—Colorless, odorless gas, of sweetish taste. Dentists keep it liquefied under pressure in iron cylinders. NH4N03 = N2O + 2H20. Chemical Properties.—By the ease with which it gives up its O it is a supporter of combustion and life, next to O itself. Medical.—lnhaled, diluted with air, it produces exhilaration of spirits, muscular activity, and then complete anaesthesia. Used in dental and other brief minor operations. PART I.—INORGANIC CHEMISTRY. 45 Nitrogen Dioxide—N202 (.Nitric Oxide).—Prepared by action of nitric acid on copper ;—* 3Cu + BHN03 = 3Cu(NO3)2 + 4H20 + N202. A colorless gas, which, when coming in contact with free O, forms red vapors of N203 and N204 ; hence a test for free O. Nitrogen Trioxide—N203 {Nitrous Acid—HN02).—Nitrous acid is known only in its salts, the nitrites. These are produced in nature by the oxidation of nitrogenous organic matter in the presence of cer- tain forms of microscopic life. This nitrification occurs in waters polluted with organic matter, and normally in the soil, where the acid so formed combines with bases. Hence nitrites in water is evidence of previous contamination with nitrogenous matter. Further oxidation forms nitrates. Nitrogen Tetroxide—N«04—occurs in company with N203 in the brown fumes given off whenever nitric acid is decomposed, as in cer- tain laboratory and manufacturing processes. The effect of breathing air thus contaminated is to produce chronic inflammation of the respira- tory tract. If the vapor be more concentrated the effects are more acute and serious. At first there is only a cough, in two or three hours a difficulty of breathing, and in about twelve hours, death. The remedy is ventilation. Nitrogen Pentoxide—N203—is of no medical interest. Nitric Acid—HN03 {Aqua Fortis)—occurs in traces in the atmos- phere and as nitrates in the soil. (See Nitrites.) Prepared by distilling a nitrate with sulphuric acid. 2KNO3 + H2S04 = K2S04 + 2HN03.t * Experiment.—Copper turnings, clippings, or wires are placed in a flask, and nitric acid diluted with half its volume of water is poured in, and the flask set in cold water. Red fumes soon fill the flask, but when these have escaped the gas appears colorless, turning red, however, on reaching the air. The colorless gas is collected over water. f Experiment.—ln the laboratory nitric acid may be prepared with the apparatus shown in Fig. 22. Equal parts of sodium nitrate and sulphuric acid are heated in the retort. The nitric acid produced is vaporized by the heat and recondensed in a receiver kept cool by a wet cloth, over which flows a stream of water from an elevated vessel. 46 ESSENTIALS OF CHEMISTRY, Physical Properties.—Heavy liquid, colorless, but if old and exposed to light it may be yellow or orange from presence of N203 and N204. Like all other nitrates, it is soluble in water. Chemical Properties.—Energetic oxidizer; * corrosive ; stains skin indelibly yellow. Medical Uses.—The strong acid is an escharotic, coagulating the albumin of the tissues; the dilute, a refrigerant tonic. Tests.—(i) Yellow stain. (2) Add H2S04, and then a crystal of FeS04 dropped in will be colored brown if nitric acid or any nitrate be present. Fig. 22. Making HN03. PHOSPHORUS (Light-bearer) occurs combined with O in the ancient unstratified rocks. These disintegrate and form soil, from which the P passes into the organisms of plants, and thence into the bodies of animals. First isolated by Brandt (1669) from urine ; now obtained from bones. Physical Properties.—A soft, yellowish solid, resembling unbleached * Experiment.—lnto a mixture of strong sulphuric and nitric acids pour from a beaker tied to a long stick some warm turpentine. The oxidation is so rapid that the turpentine is inflamed. PART I.—INORGANIC CHEMISTRY. 47 wax.* Insoluble in water, but soluble in carbon disulphide, ether, chloroform, oils, etc. Chemical Properties.—Very inflammable,! so kept under water; exposed to the air, it undergoes a slow combustion, emits the odor of ozone, and is luminous in the dark. Physiological.—Liable to inflame from careless handling, and burns by it are difficult to heal. In medicinal doses, a nerve tonic and aphrodisiac ; in larger quantities a virulent poison and gastro-irritant. Sometimes given with homicidal intent, but more frequently taken accidentally as rat poison, tips of matches, etc. Workmen in match factories suffer from irritation of stomach and bowels, caries of teeth, necrosis of bones, especially of lower jaw, and from fatty degenera- tion of various organs. This may be prevented by using the red allotropic variety, which is harmless. No good antidote. Evacuate the stomach; give copper sulphate! as emetic and antidote; give old turpentine, the ozone of which oxidizes the P. Avoid fats, for they dissolve it. Phosphorus Hydride. PH3 {Phosphoretted Hydrogen—Phos- phine)—occurs mixed with other hydrides of P in the gases arising from decomposing animal or vegetable matters, especially under water; hence seen as the ignis fatuus, or “Will-o’-the-wisp,” over marshes and graveyards. Tests.—(i) Shines in the dark; (2) emits garlicky odor. Prepared by boiling phosphorus in a solution of caustic potash. § * When heated to 500° F. in an atmosphere incapable of acting upon it, phos- phorus is converted into a reddish-brown powder, which, unlike ordinary phos- phorus, is not poisonous, not inflammable, and insoluble in the ordinary solvents. t Experiment.—Dissolve some phosphorus in carbon disulphide. Pour this on a sheet of filter paper hung on a retort stand. Soon the solvent evapo- rates and leaves the phosphorus in such a fine state of division that it inflames spontaneously. J Experiment.—Place a clean bit of phosphorus for a minute in a solution of copper sulphate. Remove, and note the coating of metallic copper. $ Experiment.—lnto a retort, whose delivery tube dips under water in a dish, add liquor potassae and a few bits of phosphorus. Expel the air by pass- ing hydrogen or illuminating gas through the retort, or by adding a few drops of ether, the vapor of which does the same thing. On applying heat the hydrogen or illuminating gas or ether vapor first escapes, then come bubbles of PH3, each of which, as it bursts into the air, ignites spontaneously, forming beautiful rings of white smoke rotating on their circular axes. These may ascend to the ceiling if the air be still. 48 ESSENTIALS OF CHEMISTRY. Properties.—Colorless gas, of a garlicky odor; inflames spontane- ously upon coming in contact with the air; very poisonous. Phosphorus Oxides. These are analogous to the oxides of hydrogen, and form, on the addition of water, analogous acids. Phosphorus Pentoxide (P206) is produced whenever P burns in Fig. 23, air* or O ; and forms three different phosphoric acids by combining with one, two or three molecules of water, thus:— P205 -f- 3H20 = H6P208 = 2H3P04 = Orthophosphoric acid. P205 -j- 2H20 = H4P207 Pyrophosphoric acid. P2Os -f- H2O = H2P206 = 2HPO3 = Metaphosphoric acid. Orthophosphoric Acid.—Never found free, but is widely dis- * Experiment.—A little stand in the middle of a dinner-plate supports a capsule, into which is put a bit of phosphorus freed from adhering water. This is ignited and covered with a bell jar. This jar is filled with clouds of P205, which, aggregating, fall into the plate like a miniature snow-storm. PART I.—INORGANIC CHEMISTRY. 49 seminated in the three kingdoms of nature in its salts, the phos- phates. Being the phosphoric acid most used in medicine (the other two are poisonous), it is usually called simply “ phosphoric acid.” Transparent, odorless, colorless, syrupy liquid. Being tri- basic, it forms three classes of phosphates by displacement of one, two, or three atoms of the basic hydrogen, thus: KH2P04, K2HP04, and K3POr In the diluted form (acidum phosphoricutn diluium) it is prescribed as a tonic, especially in dyspepsia. Tests.—Add a few drops of the magnesian fluid (MgS04, NH4CI, and NH4HO, each one part, water eight parts); a white precipitate indicates phosphoric acid or other phosphate. ARSENIC occurs mostly as sulphide, usually associated with other metals. The ore is roasted, and the resulting oxide heated with carbon (charcoal) gives the metal. This is a brittle, steel-gray, crys- talline solid, possessing a metallic lustre. Heated out of contact with air it sublimes; in air it burns with a bluish-white flame, emitting the odor of garlic and white clouds of As203. It combines with many elements; the metallic arsenides resemble alloys. Used in pyrotechny and in the manufacture of shot, pigments and fly-poison. All its compounds are poisonous. Arsenious Hydride—AsH:j—Arseniuretted Hydrogen—Arsine— is of great practical interest to the toxicologist, as its formation con- stitutes one of the most delicate tests for arsenic. Forms whenever hydrogen is generated in presence of an arsenical compound. Arsenious lodide—Asl3.—Prepared by fusing together atomic proportions of its constituent elements. It enters into Donovan s Solution, liq. arsenii et hydrargyri iodidi, U. S. P. Arsenious Sulphide—As2S3—occurs native as orpiment. Pre- pared by precipitating an arsenious compound with H2S. Bright yellow powder, insoluble in water or acid solutions, but soluble in alkaline. Another sulphide is realgar, AsS2. Both are used as pig- ments—the orpiment as a yellow, the realgar as a red. Oxides and Acids. As203 -f- 3H20 = 2H3As03, (ortho) Arsenious acid. As205 -f- 3H20 = 2H3As04, (ortho) Arsenic acid. ESSENTIALS OF CHEMISTRY. Arsenious Oxide—-As203. Arsenic, White Arsenic, Ratsbane, Arsenious Acid.—This is not only the most important compound of arsenic, but the most important of toxic agents, whether we consider the deadliness of its effects or the fatal frequency of its administration. When recently made it is in glassy lumps, which on exposure become crystalline and opaque. When sublimed it is deposited again in brilliant octahedral crystals. It is odorless, almost tasteless—slightly sweetish. When powdered arsenic is thrown upon water it does not all sink, notwithstanding its heaviness, but floats, showing a repulsion of the water. Very slightly soluble in water, even boiling water dis- solving less than two per cent. If the water be made acid or alkaline, it dissolves more readily. When arsenic dissolves in water it forms arsenious acid, H3As03. There are two officinal solutions, each containing one per cent, of arsenic: (i) Liq. acidi arseniosi, in which the water is acidulated with HCI; (2) Fowler s Solution, liq. potassii arsenitis, in which the water is made alkaline by K2C08. Arsenic Oxide.—Arsenic pentoxide is made when arsenious oxide (Asoo3) is treated with an oxidizing agent, as nitric acid. It is quite soluble in water, with which it forms a series of arsenic acids (ortho-, pyro- and meta-) analogous to the phosphoric acids. Toxicology of Arsenic.—The deadly effect of arsenical compounds has been known from remote antiquity, and they have probably been more used for homicidal purposes than all other toxic agents com- bined. Although chemistry has made its detection easy and certain, arsenic is so cheap, so readily administered to the unsuspecting victim, and so deadly, that it is still a favorite with the murderer. Owing to the extensive use of arsenical compounds as insect-powders (Paris green, etc.), and as pigments Tor wall-paper, toys, confectionery, etc., cases of accidental poisoning are quite common. Few physicians have the training and facilities to undertake an extended analysis, but they should all know the simpler tests, so as to promptly recognize the nature of the poison and combat it intelli- gently and successfully. Besides, the physician, being early in the case, can by wise precautions prevent breaks in the chain of evi- dence ; protecting the prisoner if innocent, and closing loop-holes of escape if guilty. If foul play is suspected, he should commit all his observations to writing, for notes to be admitted as evidence must be the original ones taken at the time. Having collected the urine, TART I.—INORGANIC CHEMISTRY. 51 faeces, vomit, and the suspected vehicle of the poison, and having tested some or all of them to verify his suspicion, he should place them under seal or lock and key. He should carefully reserve his opinion, lest he do injustice to the innocent or warn the guilty. In case of death, the coroner should be notified and an autopsy held, in presence of the chemist if possible. The stomach and entire intestinal canal, ligated at both ends, half of the liver, the whole brain, spleen, one kidney, and any urine remaining in the bladder should be saved. These, if possible, should be preserved in separate jars, to which a little pure chloroform may be added to prevent decomposition. These jars must be new and clean, closed with new corks or glass—not zinc caps. They are then to be labeled, and also sealed and stamped, so they cannot be opened without detection, and as soon as possible turned over to the chemist or prosecuting officer. The symptoms of arsenical poisoning are those common to all intense irritants, viz., nausea, vomiting, burning pain in the epigastrium, purging, cramps, thirst, fever, rapid pulse, etc., ending in collapse. Smallest fatal dose is two grains, and death usually occurs in twenty- four hours. Treatment.—Remove any unabsorbed poison from the stomach by emetics or stomach-pump. The best antidote *is freshly precipitated ferric hydrate, made by adding aqua ammoniae to a solution of a ferric salt. ‘‘Dialysed iron,” being a solution of ferric hydrate, may be used. It should be given at frequent intervals and in tablespoonful doses. Tests for Arsenic.—In the solid state: 1. Heated on a knife- blade over a lamp, it volatilizes with a white smoke, and leaves no residue. 2. Heated in a test-tube it sublimes, and is recondensed in the cooler portion of the tube (Fig. 24) as octahedral crystals (Fig. 25). 3. Heated in a small tube with powdered charcoal, the arsenic is reduced as it sublimes, and recondenses on the cooler portion of the tube in the metallic state. In the liquid state : 1. Through the solution, acidulated or rendered neutral, pass H,S ; a yellow precipitate (As2S3) falls. *An antidote is something harmless and capable of rendering the poison harmless. Since poisons are inert when insoluble, antidotes are usually such substances as are capable of combining with the poison to form an insoluble and therefore inert compound. 52 ESSENTIALS OF CHEMISTRY. 2. To an aqueous solution add a few drops of nitrate of silver, and then cautiously add ammonia, drop by drop, till a yellow precipitate, silver arsenite (Ag3As03) is obtained, showing the presence of arsenic. 3. Repeat the preceding, adding copper sulphate instead of silver nitrate, and the presence of arsenic is indicated by a green precipitate of copper arsenite (Scheele’s green or Paris green). The last two tests may be performed with greater ease and delicacy if the silver nitrate and copper sulphate each be previously treated with ammonia until the precipitate first formed is barely dissolved, forming solutions of a7nmonio-nitrate of silver and ammonio-sulphate of copper, which are filtered and set aside as test reagents. Fig. 24. Fig. 25. The Plating (.Reinsch's) Test.—Place a thin piece of pure copper in the solution acidulated with HCI, and boil. If arsenic be present, it will be deposited as a metallic film on the copper. If the solution be then poured off, and the piece of copper, carefully dried, be heated in a dry test-tube, the film will sublime and condense on the sides of the tube, and the preceding tests may be applied. The Hydrogen (Marsh's) Test depends on the fact that AsH3 is always formed whenever hydrogen is generated in the presence of any arsenical compound. Generate hydrogen (Fig. 26) in the usual way (Zn -f- H2S04), and if the chemicals are pure (free from arsenic), the gas burns with a pale yellowish flame, without odor, and does not stain a porcelain dish held in the flame. Then pour into the generator some of the suspected solution. If arsenic be present, there is an odor of garlic ; the flame becomes bluish-white, and a cold porcelain PART I.—INORGANIC CHEMISTRY. 53 dish held in the jet (Fig. 28) so chills the flame that only the H burns, and the As is deposited on the porcelain as a brilliant metallic film. Fig. 26. Fig. 27. If the delivery tube be heated (Fig. 27), the passing AsH3 is decomposed, and metallic ar- senic deposits farther out in the tube in a film of the same character as that on the porcelain. This may be distinguished from the him formed by antimony under similar circumstan- ces by (i) its greater metallic lustre, and (2) by its dissolving on the addition of chlorinated soda (Labarraque’s solution); (3) moisten the spot with nitric acid; evaporate the acid; a white stain is left, which is colored a red by AgN03 and yellow by H2S. The flame should now be extinguished and the delivery-tube made to dip into a solution of AgN03. This will be black- ened, and, if overlaid with aqua ammoniae, a yellow pre- cipitate will appear at the junction of the two fluids. Fig. 28. ANTIMONY (stibium) occurs native, but usually as a sulphide. Prepared by roasting the sulphide, and heating the oxide thus ob- tained with charcoal. Properties.—A bluish-white, brittle, crystalline solid, with a brilliant metallic lustre. Resembles metals and forms alloys. In chemical 54 ESSENTIALS OF CHEMISTRY. properties it plays the role of positive and negative radical with equal facility. Used in alloys, as type metal, Babbitt’s metal, Britannia, etc., to which it gives hardness and causes them to expand and fill the moulds on solidifying. The metal is not used in medicine, most of the com- pounds being obtained from the sulphide. Antimonious Hydride—SbH3 (Antimon lure tied Hydrogen—Sti- bifie), corresponding to AsH3,—This gas is formed wherever hydrogen is generated (nascent) in presence of a reducible antimony compound. Antimonious Chloride—SbCl3.—At ordinary temperatures a yel- low semi-solid; hence called butler of antimony. On addition of considerable water it decomposes, precipitating a white powder, the oxychloride (SbO.Cl),* formerly called powder of algaroth. Oxides and Acids of Antimony. Sb203 -f- H2O = 2HSb02—(meta) Antimonious acid. Sb205 -)- H2O = 2HSb03—(meta) Antimonic acid. Antimonious Oxide—Sb20t.—Prepared by treating the oxychloride with sodium carbonate to remove the chlorine, A whitish, insoluble, volatilizable powder. Antimony and Potassium Tartrate—(SbO)KTf (Tartar Emetic). Prepared by treating Sb203 with the bitartrate of potassium, thus : 2KHT -f Sb203 = 2(SbO)KT + H2O. Sweetish, metallic taste; soluble in water and slightly so in alcohol. Dissolved in Sherry wine it forms vinum antimonii, U. S. P. It enters also into unguentum antimonii and syrupus scillce compositus, U. S. P. Antimonious Sulphide—Sb2S3, the principal ore of antimony; occurs native in black, lustrous masses. It may be precipitated from any antimonial solution by H2S as an orange powder, iwhich is black when thoroughly dried. Poisoning by antimony occurs oftenest with tartar emetic, for that salt is used more than all the other compounds of antimony. The * SbO and BiO, called respectively antimonyl and bismutkyl, are univalent radicals, because two valences of the trivalent element being satisfied by the bivalent O, only one free valence is left. f (T) is used to represent the tartaric acidulous radical (C 411406),06), PART I.—INORGANIC CHEMISTRY. 55 symptoms are those referable to gastro-enteric irritation. Fortunately the salts of antimony are emetic, and cause spontaneous evacuation of the stomach. Encourage this, and give tannic acid or ferric hydrate, which will form an insoluble compound. The presence of antimony may be detected by (i) orange precipi- tate with H2S ; (2) by Marsh’s test (see page 52). BISMUTH occurs native and as sulphide. Prepared by roasting the sulphide in air, and reducing the resulting oxide with charcoal. Prope7'ties.—A brittle, white metal, with a bronze tint; volatilizes at a white heat. Forms compounds closely analogous to those of Sb, but is more positive, and plays the negative role with less facility. Used in alloys; e.g., pewter and stereotyping metal; the latter melts in boiling water. Bismuth Nitrate—Bi3N03.—Formed by treating bismuth with nitric acid. Dissolves in a little water, but if much water be added it decomposes, with precipitation of— Bismuth Subnitrate—BiONOs (.Bismuth Oxynitrate)—a white, tasteless powder, much used in medicine and as a cosmetic (pearl white). Bismuth Subcarbonate—(Bio)2C03.—Similar to the preceding in constitution, properties, and uses. Bismuth and Ammonium Citrate.—Obtained in pearly scales by dissolving the citrate in dilute ammonia-water, evaporating to a syrupy consistence and spreading on glass to dry. Being very soluble it is the preparation used in making the popular elixirs of bismuth. Physiological.—The bismuth salts are tonic, sedative, mildly astrin- gent and antifermentative. Used to allay gastro-intestinal irritation. Occasionally the irritation is increased from presence of arsenic which unscrupulous manufacturers often fail to remove as the Phar- macopoeia directs. When preparations of bismuth are taken, the stools are blackened by the sulphide formed with the H2S in the intestines. In severe cases of diarrhoea, with acid fermentation, this blackening does not occur, and its reappearance is a sign of improve- ment. Tests.—(1) H2S or NH4HS gives brownish-black precipitate; (2) the concentrated solution poured into water forms a white precipitate. 56 ESSENTIALS OF CHEMISTRY. V. Carbon Group. Carbon [carbo, a coal), C, 12 Silicon (silex, a flint), Si, 28 Tin [Stannum), Sn, 118 Lead [Plumbum), Pb, 207 Each element is bivalent and quadrivalent. The dioxide of each forms with water a dibasic acid;— C02 -j- II20 = H.2C03, Carbonic acid. Si02 ~f- H2O = II2Si03, Silicic acid. Sn02 -f- H2O = H2Sn03, Stannic acid. Pb02 -f- H2O H2Pb03, Plumbic acid. CARBON occurs free in its three allotropic forms, diamond, graphite, and coal; combined in carbonates and in all animal and vegetable substances. All its forms are probably traceable to organized life. Diamond.—Geological history unknown; transparent crystalline body of great brilliancy; hardest substance known. Used as a gem and for cutting glass, etc. Graphite (to write).—Owing to its resemblance to lead it has been called black lead or plumbago; almost pure carbon. Used for pen- cils, crucibles, stove polish, etc. Coal.—Mineral coal is a black substance, compact in texture, the remains of vegetable life of past ages. Charcoal is obtained by burn- ing heaps of wood with a limited supply of air. The volatile con- stituents pass off, leaving the carbon as a light, porous substance, retaining the form and structure of the wood. Animal charcoal is made by heating animal matters in closed iron retorts. Charcoal, especially animal, is a valuable absorbent of odorous gases* and coloring matters.f Properties.—Free carbon is solid at all temperatures, and insoluble in all menstrua. Ordinarily, free carbon is unaffected by chemical * Experiment.—Fill a test-tube with ammonia gas over mercury (Fig. 29). Introduce a piece of charcoal recently heated. The gas is absorbed as is shown by the rapid rise of the mercury. ■\ Experiment.—To a solution of indigo, cochineal, or potassium perman- ganate or beer in a flask, add some animal charcoal, shake up and filter. The filtrate is colorless, and in case beer is used it has also lost its bitter taste. PART I.—INORGANIC CHEMISTRY. 57 agents, but at high temperatures it surpasses all other elements in its avidity for O. Hence it is used to separate the metals from their oxides.* Carbon Monoxide CO—occurs whenever carbon is burned with an insufficient supply of air, as in anthra- cite stoves and furnaces, and in coal gas, but never occurs in nature. Fig. 29. Prepared in the laboratory by heat ing oxalic acid, or potassium ferrocya- nide, with sulphuric acid. Properties. Colorless, odorless, tasteless gas; burns with a pale blue flame ; very poisonous, combining with the coloring matter of the blood cor- puscles, and destroying their oxygen- carrying power. Artificial respiration is of little use. Transfusion of blood is the most promising treatment. After death the blood remains scarlet. The sources of danger are open charcoal fires, defective draught in stoves and chimneys, and illu- minating gas escaping into bed-rooms. Carbon Dioxide—C02. C02 -)- H2O = H2C03—Carbonic acid. Occurs sparingly (.0004) in the atmosphere, as a result of animal respiration, vegetable decay, and combustion. Plants absorb it, appropriating the carbon and returning the oxygen to the air. It often accumulates in cellars, beer-vats, wells, etc., where it is called choke-damp. Prepared by burning carbon, but most conveniently in the labora- tory by decomposing a carbonate with an acid. CaC03 + 2HCI = CaClj + H2O + C02. Physical Properties.—Transparent, colorless gas, of a pungent odor * Experiment. —lnto a slight depression in apiece of charcoal lay some metallic oxide—e.g., lead oxide—heat with a blow-pipe. The oxide is reduced by the heated charcoal around it, and globules of the metal appear which coalesce into a bright button. 58 ESSENTIALS OF CHEMISTRY, and sour taste. One and a half times as heavy as air,* Water dis- solves its own volume. Soda-water is only a solution of this gas under pressure.! Fig. 30. Making CO. Chemical Properties.—Neither burns nor supports ln water it exists as carbonic acid—H2C03. On attempting to concen- * Experiments.—To show the weight of carbon dioxide : (1) Pour it from one vessel to another. (2) Blow soap bubbles and allow them to fall into a wide vessel containing this gas. As soon as they reach the surface of the gas they stop and float upon it. (3) Pour a large beakerful of the gas into a light pasteboard box that has been balanced on a pair of scales. The box will at once descend. This gas accumulating in wells may be bailed in buckets, and tested by being poured upon a lighted candle. f That water will dissolve a greater quantity of carbon dioxide under pres- sure is shown by the rapid evolution of the gas whenever a bottle of soda or other carbonated water is opened and the pressure thereby removed. X Set a candlestick, holding several lighted tapers at different heights, in a large jar. Carbon dioxide is introduced at the bottom, and extinguishes the tapers one by one as the vessel fills up to their levels. PART I.—INORGANIC CHEMISTRY. 59 trate this dilute solution the acid decomposes again into water and C02; hence wet litmus reddened by it becomes blue again on dry- ing. Tests.—(i) The gas (15 per cent, and over) extinguishes a flame; (2) precipitates lime-water ;* (3) carbonates effervesce on adding a strong acid. Physiological.—lf the gas be undiluted, death is immediate from spasm of the glottis. If somewhat dilute (15 to 30 per cent.) there is loss of muscular power, anaesthesia, and death without a struggle. If quite dilute (5 to 10 per cent.) headache, giddiness, muscular weak- ness, and sometimes vomiting and con- vulsions occur. Fig. 31. The effects are more serious if the C02 comes from combustion or respiration, because of the removal of oxygen and the admixture of the deadly CO and ani- mal exhalations. Treatment.—Fresh air, artificial respira- tion, and stimulation. The preventive is ventilation. Ventilation.—More than 7 parts of C02 in 10,000 of air is oppressive. Taking this as the maximum impurity allow- able, 3000 cubic feet of fresh air per hour is needed by each person, and more in case of disease or when lamps are burn- ing. To secure this in a room containing 1000 cubic feet (10 XlO X IO) the air must be changed three times an hour. This would give a draught not uncomfortable or injurious. If the draught be properly distributed, a breathing space of 500 cubic feet changing six times an hour would be unobjectionable. Ventilation may be secured in two ways, by diffusion and by draught. * Experiment.—Two Wolff bottles are half filled with lime-water and arranged as in Fig. 31. Placing the rubber tube in his mouth, the operator can inspire through one bottle and expire through the other. The small amount of carbon dioxide in the inspired and the larger amount in the expired air is shown by a white precipitate, slight in the one and dense in the other bottle. ESSENTIALS OF CHEMISTRY. Diffusion.—Gases mingle more rapidly, liquids more slowly, to make a mixture of uniform density. Fig. 32. When two gases of different densities are sepa- rated by a porous partition, they mingle, the lighter passing through more rapidly than the heavier, the rapidity being in inverse ratio to the square roots of their densities.* This diffusion is more active in winter than in summer, because of the greater difference in den- sity of the warm air within the house and the cold air without. Damp walls are unhealthful mainly because being no longer porous they prevent this diffusion. Cyanogen—CN or Cy. Univalent because Nm can satisfy only three valences of CIV. A com- pound negative radical resembling in its chemical behavior the elements of the chlorine group. Prepared by strongly heating mercuric cyanide.f Hg(CN)2 = Hg + 2CN. A colorless gas, smelling like peach kernels. Burns with a peach-blossom flame; unites with metals to form cyanides, the most important being— Hydrocyanic Acid—H(CN), or HCy—{Prussic Acid, Hydrogen Cyanide').— Occurs in bitter al- monds, cherry-laurel water, etc. Properties.—Colorless liquid, having an odor like peach kernels. * Experiment.—Cement a porous earthenware battery cup at its open end to the top of a funnel tube, the end of which dips into a bottle of colored water. Support on a stand, as in Fig. 32. Bring down over the cup an inverted bell jar of hydrogen. The light H diffuses so much faster into the cup than the air diffuses out of it, that bubbles of gas escape rapidly through the water. Remove the bell jar and the conditions are reversed. The H now diffuses so rapidly out of the cup that the water is sucked up the tube. j- If mercuric cyanide cannot be obtained, a mixture of two parts of thor- oughly dried potassium ferrocyanide and three parts mercuric chloride may be used. PART I.—INORGANIC CHEMISTRY. For medical purposes only a dilute (2 per cent.) solution is used, and the dose is from two to five drops. Toxicology.—All the cyanides are very poisonous. One drop of the pure acid produces immediate death, and three grains of potassium cyanide kills in a few minutes. The respiratory centres are paralyzed, and the victim falls and dies in convulsions. Poisoning is liable to occur from handling the acid or the cyanides, which are largely used in the arts, or from eating vegetable products, e.g. peach and cherry seeds containing amygdalin, a substance easily decomposing into prussic acid and other products. Owing to the rapid action of the poison, antidotes are usually impracticable. Use artificial respiration and stimulate. If the patient survive an hour, the prognosis is good. Tests.—(l) Its odor; (2) silver nitrate—white precipitate soluble in boiling HNOs; (3) add ammonium sulphydrate, evaporate to dryness, and then add ferric chloride—a blood-red color. Cyanates.—Cyanic acid (HCyO) and ammonium cyanate (NH4- CyO) are the most interesting. The latter on being heated forms urea. Sulphocyanates are sulpho-salts corresponding to the cyanates (oxy-salts), and are good illustrations of the facility with which S forms series of compounds analogous to those of O. They, especially the potassium and sodium salts, are used as test reagents. Compound Cyanides.—Cyanogen shows a great tendency to form complex radicals, especially with iron; as, ferrocyanogen [FeII(CN)6I]IV or (FeCy6)IV, and ferricyanogen [Fe2VI(CN)i2rjVI or (Fe2Cyl2)vl. These two radicals contain ferrous and ferric iron re- spectively, and with hydrogen form acids known as hydro-ferrocyanic acid, H4FeCy6 (tetrabasic), and hydro-ferricyanic acid, HOFe2CNi2 or H6Fe2Cyi2 (hexabasic) ; the salts of these acids are termed ferrocyan- ides and ferricyanides. Potassium Ferrocyanide—K4FeCy6—commonly called yellow priissiate of potash, and potassium ferricyanide—K6FeCy12—red prussiate of potash, are important test reagents. The carbon compounds will be further considered under the head of Organic Chemistry. SILICON (also called silicium) resembles carbon, and occurs in three allotropic forms corresponding to coal, graphite, and diamond ; most abundant element after oxygen. It exists in only a few com- 62 ESSENTIALS OF CHEMISTRY. pounds, but they constitute the larger part of the earth’s crust. Its principal compound is— Silicic Oxide—Si02—occurring as flint, sand, rock-crystal, etc.; with water it forms silicic acid. Clay, soapstone, asbestos. Silicates of aluminium and magnesium are very abundant, as clay, soapstone, asbestos, etc. Glass is a mixture of several silicates, usu- ally of sodium, calcium and sometimes lead. It is made by melting sand (Sio2) with the carbonates or oxides of the metals. The addition of certain metallic oxides gives color; e.g., cobalt gives a blue, manga- nese an amethyst, and copper a ruby. If the glass consist of only an alkaline silicate {e.g-, sodium), it is soluble or water-glass, which is largely used in surgical dressings. TIN.—A bluish-white malleable metal,* not corroded by air or water; hence used to form a protective coating for iron and copper. Tin-ware is usually sheet-iron coated by being dipped into molten tin. Tin alloyed with lead is easily dissolved, and may cause lead- poisoning. Tin-foil (thin laminse of tin) is used in wrapping to exclude air and moisture. Tin enters into the composition of a great many alloys. Powdered tin is sometimes used as an anthelmintic. * The Metals.—Occurrence.—Some, as gold and copper, occur free, but most of them are found combined with non-metallic elements, especially sulphur and oxygen. Preparation.—lf combined with sulphur the ore is roasted until the sulphur is burned out, leaving the metal as an oxide, which is then heated with carbon to remove the oxygen, thus : Physical Properties.—Very opaque, with a “metallic lustre” (in fine powder, a dull black); bluish-gray, varying between the pure white of silver and the dull blue of lead. Yellow gold and red copper are exceptions. In weight, varying between lithium, specific gravity 0.58, and platinum, specific gravity 21.50. Most are solid, except mercury (liquid) and hydrogen (gaseous). All are absolutely insoluble. Chemical Properties.—Electro-positive, possessing great affinity for the non- metals and other electro-negative radicals. When two metals are fused together the product is an alloy. If one of the metals be mercury, it is called an a?nal- gam. Alloys are not chemical compounds, but mixtures, for the metals do not unite in definite proportions, and the alloy is not a new substance, but one with properties intermediate between those of its constituent metals. Used mostly in the arts. Of the fifty-five metals only about twenty-six, or rather compounds of these, enter the materia medica, and merit our notice. ZnS +O3 = ZnO + S02; then, ZnO +C= CO + Zn. PART I.—INORGANIC CHEMISTRY. 63 Tin forms two classes of compounds; the stannous, in which the atom is bivalent, and stannic, in which the atom is quadrivalent. These are of importance to the chemist, but of little interest to the physician. LEAD.—Its principal ore is its sulphide (PbS), called galena. It is a soft, heavy, blue metal, very slowly acted upon by most sub- stances ; hence used to make water-pipes and vessels that are exposed to corrosive liquids. Water containing nitrates or nitrites (from organic matter) dissolves lead slightly; but if it contains carbonates or sulphates, the lead is protected by an insoluble coating of lead carbonate or sulphate. Lead enters into the composition of many alloys: as pewter, solder, shot, type-metal, etc. The quadrivalent compounds of lead are of so little importance that the term phimbic is applied to the bivalent compounds. Lead Oxide—PbO—Litharge.—A yellow substance, found native; made artificially by heating lead in the air. It is by treating this with the appropriate acid that most of the lead salts are prepared. ' When rubbed with oil it decomposes the glyceryllic ethers and combines with the fatty acids to form lead soaps, one of which, the oleate, is lead plaster, emplastrum plumbi, U. S. P. Lead Dioxide, ox puce lead, is a dark-brown powder, forming one of the constituents of red lead (Pb3o4 or 2PbO.Pb02). Prepared by treating red lead with nitric acid to dissolve out the PbO. Lead Nitrate—Pb2N03. Made: PbO + 2HNO3 = Pb2N08 + H„0. Ledoyens disinfectant fluid is a solution of Pb2N08 (one drachm to the ounce). It corrects fetid odors by neutralizing H»S and NH4HS. Lead Acetate—Pb(C2H302)2, or PbAc2—Sugar of Lead. Used in medicine more than any other lead salt. Its solution will dissolve considerable quantities of PbO, forming the solution of the subacetate of lead, the liquor plumbi subacetaiis, U. S. P., Goulard's extract. It is astringent, and, like all the lead salts, sedative. Much used as a topical application in erysipelas, acute eczema, and other skin affections; and diluted {lead-water), it is used in conjunctivitis and other mucous inflammations. Made: PbO + 2HAc = PbAc2 + H2O. 64 ESSENTIALS OF CHEMISTRY. The following insoluble salts may be made by precipitation from solutions of the preceding soluble ones:— Lead Chloride—PbCl2.—Made: Soluble lead salt added to a soluble chloride ; e.g., PbAc2 + 2HCI = PbCl2 + 2HAc. Slightly soluble in warm water, but in cold it is always precipitated from solu- tions of moderate strength; hence classed with HgCl and AgCl as one of the three insoluble chlorides. Lead Sulphate—PbS04.—Forms as a white precipitate whenever a solution of a lead salt is added to a sulphate solution, thus : PbAc2 -|- ZnS04 = PbS04 -)- ZnAc2. Lead Carbonate—PbC03 White Lead. Made : PbAc2 -f- Na2C03 = PbC03 -f 2NaAc. Commercially, it is made by some modification of the old Dutch method, which consists in covering bars of lead with the refuse of the wine-press and barn manure. The acetic fumes from the grape husks attack the lead, forming lead acetate, which is decomposed by the carbonic acid from the manure. The acetic acid thus liberated com- bines with another portion of lead, which is again precipitated by the carbonic acid, and thus the process continues until all the lead is consumed. Used for painting, but blackens when air contains H2S. Lead Sulphide—PbS—is formed as a black precipitate whenever a lead solution is treated with a soluble sulphide, as H2S or NH4HS. Lead lodide—Pbl2.—A bright yellow precipitate on adding a soluble iodide to a lead solution ; as,— PbAc2 + 2KI = 2KAc + Pbl2. Lead Chromate—PbCro4. Made : PbAc2 + K2Cr04 = PbCr04 -f- 2KAc. Under the name of chrome yellow it is used in painting. Of late it has been used to color food products. Tests for lead consist in forming precipitates of the foregoing in- soluble compounds. Phys-iological.—All the lead compounds are poisonous. Acute poisoning sometimes occurs from the ingestion of a single large dose of a soluble lead salt. The symptoms are those of gastric irritation. Treatment. Give MgS04 to form the insoluble PbS04. PART I.—INORGANIC CHEMISTRY. 65 The chronic form of lead intoxication, painter s colic, is purely poisonous, and is produced by the continued absorption of minute quantities of lead by the skin of those handling it, and by the lungs and stomachs of those living in painted apartments, or using food and drink from leaden vessels. There is impairment of digestion, consti- pation, blue line along the edge of the gums, colic, and paralysis, especially of the extensor muscles. Lead once absorbed is eliminated very slowly, having combined with the albuminoids, a combination which is rendered soluble by the administration of iodide of potas- sium. The treatment for chronic lead-poisoning is to give MgS04, for the double purpose of overcoming the constipation and precipitating any lead remaining .unabsorbed in the alimentary canal; also KI to pro- mote the elimination of that which is combined with the albuminoids. Alum is a favorite treatment, seeming to perform all accomplished by both the MgS04 and KI. The paralyzed muscles must be treated with electricity, so that when the lead is eliminated and the nerve influence returns, it may not find them degenerated past redemption. VI. Metals of the Alkalies Hydrogen, ......... II i Lithium, Li 7 Ammonium, (NHJ 18 Sodium (Natrium), Na 23 Potassium (Kalium), K 39.1 Rubidium, Rb 85 Caesium, Cs 133. Univalent; very electro-positive (except H), so that when com- bined, unless it be with a strongly electro-negative (acidulous) radical, they form very alkaline compounds (hence the name). The positive affinities, as in the other groups, increase with the atomic weights. All their compounds are soluble. LITHIUM.—Sparingly but widely distributed in nature, especially in the waters of certain springs. Lightest of the solid elements. Its salts closely resemble those of sodium. Physiological.—Lithium urate being by far the most soluble com- pound of uric acid, salts of lithium, especially the carbonate, are given 66 ESSENTIALS OF CHEMISTRY. to gouty persons to promote the elimination of uric acid, which ac- cumulates in that disease. Test.—It colors the flame a beautiful carmine red. AMMONIUM.—When ammonia gas (NHj) combines with an acid, it appropriates the basic hydrogen and forms a salt in which NH4 is the positive radical; e.g. : NHj -f- HCI = NH4CI, corresponding to KCI or NaCl; NH3 + HHO = NH4HO, corresponding to KHO or NallO; NH3 -|- HN03 NH4N03, corresponding to KN03 or NaNOa; 2NH3 -j- H2S04 (NH4)2SO4, corresponding to K2S04 or Na2S04. This radical plays the role of a metal, like K and Na, and is called Ammonium. Does not exist uncombined, although Weyl claims to isolate it as a dark-blue liquid metal.* We can obtain it as an amal- gam by the reaction between sodium amalgam and ammonium chloride.f Ammonium Hydrate—NH4HO—Caustic Ammonia—is formed in solution whenever ammonia gas (NH3) dissolves in water, thus: NHS -(- H2O = NH4HO. It has been already stated that the watery solu- tion of a fixed substance is called a liquor ; of a volatile substance, an aqua. In like manner alcoholic solutions of fixed substances are called tinctures, and of volatile, spirits. There are four U. S. P. solu- tions of ammonia:— Aqua ammonia?, ..... • 10 per cent. Aqua ammonia fortior, 26 “ Spiritus ammonia 10 “ Spiritus ammonia aromaticus. * Experiment.— The supposed free ammonium. Sodio-ammonium is pre- pared by heating sodium in a sealed tube with ammonia gas. This is in turn heated with ammonium chloride in a sealed tube. A dark blue liquid, with metallic lustre, is obtained, but soon decomposes into ammonia gas and hydro- gen. f Experiment.—To some mercury in a test-tube add sodium, small bits at a time. On this sodium amalgam pour a strong solution of ammonium chlo- ride. Sodium chloride and ammonium amalgam are formed. The ammonium amalgam swells up and soon decomposes—(NH4 4~ Hg) = NH3 -f H -f Hg—the gaseous NH3 and hydrogen escape, and only the mercury remains. (Na + Hg) + NH4CI = NaCl -f (NH4 + Hg). PART I.—INORGANIC CHEMISTRY. 67 In all these solutions NH4HO exists, but has never been isolated, because, whenever we attempt to evaporate the water or alcohol, the NH4HO decomposes into NH3 -f H2O. Ammonium hydrate is very alkaline. Ammonium Hydrosulphide—NH4HS—occurs in decomposing nitrogenous, sulphurized organic bodies. Made by saturating a solu- tion of NH4HO with HoS. A yellowish solution; used as a test reagent. Ammonium Carbonate—(NH4)2C03.—Ammonii Carbonas, U. S. P. —Sal volatile—is prepared by heating a mixture of NH4CI and chalk (CaCO3) up to the temperature at which (NH4)2CO3 would be vola- tilized, when the following reaction will occur:— 2NH4CI + CaC03 = CaCl2 + (NH4)2CO3. (See Volatility, page 25.) Very prone to absorb C02 from the atmos- phere and become bicarbonate unless NH4HO be added. Other salts may be made by adding the appropriate acid to the carbonate or hydrate of ammonium. If we use the carbonate, we can tell when acid enough has been added by the cessation of -effer- vescence. If the hydrate be used there is no effervescence, and our only guide is the point at which the solution becomes neutral in reaction. This is determined by the use of test papers. These are made of white, unsized paper, steeped in a blue vegetable pigment called litmus, which is reddened by acids and restored to its blue by alkalies. Physiological.—The hydrate and carbonate are alkaline irritants, like the corresponding K and Na compounds, though in less degree. They, moreover, give off NH3, which, though irritating to the respira- tory tract, is a valuable stimulant in fainting fits, etc. Two drachms of aqua ammoniac have killed. The treatment, as for all alkalies, is to give a dilute acid or some oil. Tests.—(1) An ammonium salt warmed with liq. potassae gives off NH3, recognized (a) by its odor, (b) its forming a white cloud of NH4CI when a glass rod dipped in the HCI is held over the vessel, and (c) its changing moist red litmus to blue. (2) Heat the dry ammonium salt and it volatilizes. SODIUM occurs very abundantly as sodium chloride, or common salt, from which almost all the sodium compounds are now obtained 68 ESSENTIALS OF CHEMISTRY. instead of from ashes of seaweeds, as formerly. The preparation and properties of sodium and its compounds are so similar to those of potassium that we will omit their separate consideration. So much alike are the salts of the two metals that the choice between them is usually governed by considerations of economy, convenience, solu- bility, fashion, etc. On exposure to the atmosphere the sodium salts usually have a tendency to epfioresce, while the potassium salts tend to deliquesce. Tests.—No good precipitant; for all the compounds of sodium are soluble. However, the strong yellow color it gives a flame is a very delicate test; in fact, too delicate, for it shows traces of sodium in almost everything. Fig. 33. Fig. 34. POTASSIUM occtirs only in compounds. Prepared by heating one of its oxygen compounds with charcoal in an iron retort (K2CO3 + 2C = 3CO + K2). The metallic K distills over and is condensed in a flat receiver. Physical Properties.—Soft as wax; lighter than water; silvery lustre when freshly cut, but quickly tarnishes. Chemical Properties.—lntensely electro-positive ; hence it possesses great affinity for the non-metals;* takes O from H2O, even as ice, f setting fire to the escaping H, giving the flame the violet color char- acteristic of K (Fig. 34). * Experiment.—Potassium inflames spontaneously when lowered into a jar of chlorine (Fig. 33). Warmed with iodine or dropped into bromine it explodes violently. This should be done under a tubulated bell jar, because the potassium is scattered in every direction. -j-Experiment.—Load a strong toy cannon with gunpowder. On the fuse lay a small bit of potassium. Touching it with a piece of ice fires the cannon. PART I.—INORGANIC CHEMISTRY. 69 Potassium Carbonate—K2COs.—Obtained as an impure solu- tion (“lye”) by lixiviating the ashes of plants, especially forest trees. This, evaporated to dryness, forms “ concentrated lye.” This in turn, when purified, forms “pearl-ash,” which is further purified for medicinal use. Sometimes made by burning cream of tartar and lixiviating the residue, hence called salts of tartar. A white semi- crystalline or granular powder. C03 being a weakly acidulous radi- cal, K2C03 is very alkaline, even caustic. Acid Salts.—Salts are formed by a metallic radical displacing the basic Hof an acid. If all the Hbe displaced, the result is a normal salt, as, H2S04 +K2 = K2S04 + H2. But if part of the basic H of the acid remains, it is called an acid salt, as H2S04 + K = KHS04 + H. Sometimes acid salts are called “bi ” salts, because the proportion of the acidulous radical to the basylous is twice as great as in the normal; e.g., KHS04 is called potassium bisulphate, because the proportion of the acidulous radical S04 to the basylous radical K is twice as great as in the normal sulphate, K2S04. Potassium Bicarbonate—KHCOs.—Although an acid salt in constitution, it is alkaline in reaction, on account of the weakness of its acidulous radical. Made by passing C02 into a solution of K2COs. The reaction is as follows : K2C03 + H2O + C02 = 2KHCO3. Potassium Bitartrate—JvH (C4H406) or KHT—Cream of Tar- tar.—Prepared similarly to the above, by adding tartaric acid to a solution of the normal tartrate, thus: K2T -(- H2T = 2KHT. It exists naturally in grape juice, and, being insoluble in an alcoholic menstruum, is precipitated on the sides of the wine casks whenever vinous fermentation sets in. This is its commercial source. Other Salts.—Most salts of K are made by treating the car- bonate with the appropriate acid, e.g.:— The chloride—K2C03 -f 2HCI = 2KCI + H2O + C02. The sulphate—K2C03 -f- H2S04 = K2S04 -f- H2O -}- C02, etc. The decomposition is attended with an effervescence of C02. It is the formation of this volatile compound that determines the reaction. (See Volatility, page 25.) 70 ESSENTIALS OF CHEMISTRY. But the following salts are not made in this way : Potassium Hydrate—KHO—Caustic Potash—maybe made ex- perimentally by the reaction of metallic K on water, thus : II20 +K = KHO +H. But made in the shops by boiling K2C03 with slaked lime, thus:— K2COg + CaaHO = CaCOs + 2KHO. The insoluble CaC03 (chalk) sinks to the bottom, and the KHO dissolves in the supernatant liquid, which when clear is poured off (decanted). This watery solution, if of proper strength (SJ-Oj), forms “ Liquor potasses, U. S. P.” If this solution be evaporated to a syrupy consistence and poured into moulds, it forms the stick caustic potash. KHO is very alkaline, and a powerful cautery. Exposed to the air it absorbs C02 and forms carbonate : 2KHO + C02 = K2COs + H2O. Potassium lodide—KI:— 6KHO + 61 = SKI + KI03 + 3H20. The color disappears because the I goes to form colorless salts. Prepared thus, the KI is contaminated with KIOa (K-lodate).* But if the mixture be strongly heated the 03 is driven off and the KI alone remains. The addition of charcoal facilitates the removal of the 03. Potassium Bromide—Kßr—may be made similarly to the above. Sodio-Potassium Tartrate—NaKT—Rochelle Salt.—A neutral salt made by boiling acid potassium tartrate with sodium bicarbonate. KHT + NaHC03 = NaKT + H2O + C02. This is the reaction that occurs in baking with cream of tartar baking powders. Potassium Hypochlorite—KClO.—Made by passing chlorine into a cold solution of KHO. Yields free chlorine. The ordinary bleach- * Experiment.—The presence of KI03 in a commercial specimen of KI may be recognized by boiling a little starch in a test-tube, dissolving a crystal of the suspected salt, and then adding a few drops of a strong solution of tar- taric acid; if KI03 be present, I will be liberated, and a blue color struck with the starch. PART I.—INORGANIC CHEMISTRY. 71 ing solutions (Labarraque’s Solution or Javelle water) are solutions of impure sodium or potassium hypochlorite. Tests for Potassium.—(i) If the suspected solution be concentrated, add H2T and get a precipitate of KHT.* (2) Platinic chloride (PtCl4) gives a yellowish precipitate. But the PtCl4 is very costly, and all the potassium compounds so soluble that the above tests are but little used. The most convenient is the (3) flame test: dip the end of a clean platinum wire in the suspected solution, and hold in the color- less Bunsen or alcohol flame and notice the violet color. CAESIUM AND RUBIDIUM.—Rare metals, occurring in small quantities with potassium. Discovered in iB6O by means of the spec- troscope, and named from the colors of their lines in the spectrum (ccesius, sky blue, and rubidus, dark red). Of no medical interest as yet. Analytical.—To determine whether a salt be a compound of K, Na, NH4, or Li, heat samples of each ; the one that volatilizes is the salt of NH4. Confirm this by boiling with KHO and getting the odor of ammonia. To the other three salts apply the flame tests, getting the violet for K,f yellow for Na, and carmine for Li. VII. Metals of the Alkaline Earths. Magnesium, Mg 24 Calcium, Ca 40 Strontium, Sr 87.5 Barium, \ .Ba 137 Bivalent; their oxides and hydrates are very alkaline, but of an earthy character. Their positiveness or basicity, as in other groups, is in the order of the atomic weights. Their carbonates are decom- posable by heat and insoluble in water, unless it contains H2COs in solution. Their sulphates increase in solubility from the insoluble barium salt to the very soluble magnesium sulphate. MAGNESIUM.—Never free ; abundant in magnesian limestone (CaCO3.MgCO3). Asbestos, meerschaum, and soapstone are native *The addition of alcohol renders the test much more delicate, f The delicate violet of K may be masked by the intense yellow of Na, but can be seen if observed through a piece of blue glass, a medium that absorbs the yellow light. 72 ESSENTIALS OF CHEMISTRY. silicates. Most natural waters contain its salts. Silvery white metal; burns with a brilliant white light, rich in chemical rays, and used in photographing caves and other dark places. Magnesium Sulphate—MgS04—occurs in the waters of various springs, as those at Epsom; hence often called Epsom salts. Made artificially from the native carbonate, thus : MgC03 + H2S04 = MgSO, + (H20 + C02). White, crystalline, soluble salt of a nauseous bitter taste. It is a popular purgative. The nauseous taste and griping may be obviated by adding aromatics, acid, sulphate of iron (as in Crab Orchard salts) or by free dilution. Magnesium Citrate is the most pleasant of the saline purgatives. Usually given as the liquor magnesii citratis, which is prepared by adding a solution of citric acid to MgC03, and bottling immediately to retain the C03, Magnesium Carbonate—MgC03—occurs native. For medicinal purposes it is prepared by precipitation, thus ; MgS04 + Na2C03 = Na2S04 + MgC03. Similar to chalk in its physical and chemical properties. Magnesium Oxide—MgO—Magnesia. Made like CaO, by heat- ing the carbonate. MgCOs = MgO + C02. Insoluble and tasteless (earthy), but its alkalinity is shown by its turning moist red litmus paper blue when the solid MgO is dropped upon it. Magnesium Hydrate Mg(HO)2. Formed by precipitating a magnesium solution with potassium or sodium hydrate. Insoluble in water, but, like other salts of magnesium, soluble in the presence of ammonium compounds with which they form double salts. Suspended in water, it is called milk of magnesia. Magnesium Phosphates.—These resemble the calcium phosphates and are associated with them in the body, though in small quantity. The ammonio-magnesium phosphate (MgNH4P04) is precipitated whenever a soluble phosphate in neutral or alkaline solution finds itself in presence of an ammonium salt, as occurs in the alkaline fer- mentation of urine. PART I.—INORGANIC CHEMISTRY. 73 Physiological.—Magnesium oxide and hydrate being alkaline and tasteless, are popular antidotes for acids. These and the carbonate are given to correct acid conditions of the digestive tract, and com- bining with the acids they form soluble salts that are laxative. CALCIUM.—Never free, but its compounds are very abundant, as limestone, gypsum, etc. Calcium salts are necessary to animal life, the teeth and bones consisting mainly of calcium phosphate. Calcium Chloride—CaCl2. Made : CaC03 + 2HCI = CaCl2 + H2O + C02. A white salt; very avid of water and deliquescent; used to dry gases. Calcium Carbonate—CaC03.—Abundant as limestone, marble, corals, chalk, and shells of the Crustacea, mollusks, etc. Chalk con- sists of microscopic shells. Precipitated chalk is made by adding a soluble carbonate to a soluble calcium salt, as:— Na,CO3 + CaCl2 = 2NaCI + CaC03. The precipitate (CaCOs) may be separated from the CaCl2 in solu- tion, by— (a) Filtration.—Pouring the mixture into a cone of filter paper placed in a funnel, when the water with the dissolved salt will pass through, leaving the insoluble portion (the precipitate) on the filter, {b) Decan- tation.—Allowing the precipitate to settle to the bottom, and pouring off the clear fluid. In either case the precipitate may be washed from any remaining CaCl2 by adding pure water and repeating the process. Calcium Oxide—CaO—Lime, quicklime ; calx, U. S. P.—A white solid; made by heating limestone in rude furnaces called kilns. CaC03 = CaO -|- C02. When water is added to CaO there is a violent chemical union, great heat is evolved, and a hydrate is formed, thus : CaO + H2O = CaaHO. Calcium Hydrate Ca2HO—Slaked fane.—A white, odorless powder; very slightly soluble in water, less than one grain to the ounce, but enough to give “lime-water” (liquor calcis, U. S. P.) a 74 ESSENTIALS OF CHEMISTRY, decidedly alkaline taste and reaction. The presence of sugar greatly increases its solubility (liq. calcis saccharatus, Br.). Chlorinated Lime—Chloride of lime, bleaching powder, calx chlorata, U. S. P.—is a mixture of chloride of calcium (CaCl2) and calcium hypochlorite (CaaCIO). It is made by passing chlorine gas over slaked lime until it ceases to be absorbed. It is white, moistens on exposure to the air, absorbing C02 and giving off Cl. It is em- ployed as a source from which to get a gradual supply of chlorine for disinfecting and bleaching purposes. Calcium Sulphate—CaS04—occurs negative, as gypsum, which, when heated, loses its water of crystallization and forms a white amor- phous powder called plaster-of-Paris. If this plaster be mixed with water enough to form a creamy liquid, it will recrystallize or “set” into a hard compact mass. Much used in surgery to make casts to hold broken limbs in position. Very slightly soluble in water. Calcium Phosphate—Ca3(P04)2. It is the most abundant mineral ingredient of the body ; in every tissue and fluid, especially the teeth and bones, to which it gives hardness and rigidity. A white tasteless powder, soluble in dilute acids. Dissolved by lactic acid, it is given as syrupus calcii lactophosphatis, U. S. P., in scrofula, rickets, and other diseases of defective nutrition. Calcium Oxalate—CaC204, or CaOx—occurs in the juices of some plants and in the urine. Obtained as a fine white crystalline powder when a soluble oxalate is added to a calcium solution. Insoluble in water or acetic acid, but soluble in the mineral acids. Hard Waters are such as contain mineral matters, especially calcium (lime) compounds. Often water, in passing through the soil, becomes highly charged with carbonic acid, and dissolves con- siderable amounts of CaCO3, and is hard. This is called temporary hardness, because on exposure or boiling, the carbonic acid is driven off, the CaC03 is precipitated, and the water becomes soft. The solu- bility of CaS04 does not depend on the presence of carbonic acid, and boiling will not precipitate it. So water impregnated with CaS04 is said to be permanently hard. Drinking-water should contain a small quantity of lime; but very hard water impairs digestion. Hard water is unfit for washing, because the soluble alkali soap reacts with the lime salt to form an insoluble lime-soap. PART I.—INORGANIC CHEMISTRY. 75 STRONTIUM.—Of little importance, except that its nitrate is used in pyrotechny to make the red light.* BARIUM.—Of little interest to the medical student, except that its compounds are poisonous. Barium sulphate is very insoluble ; hence (i) the antidote of barium is some soluble sulphate, and (2) barium solutions (nitrate and chloride) are delicate tests for sulphates, and vice versa. (See Insolubility, page 26.) Barium gives the flame a green color; hence used (nitrate) in pyrotechny to make the green or Bengal light.f Analytical.—To determine whether a solution be one of barium, calcium or magnesium; Add potassium chromate; a precipitate indicates barium. If no precipitate, add ammonium chloride and then ammonium carbonate; a precipitate indicates calcium. If no precipitate, add sodium phosphate ; a precipitate indicates magnesium. VIII. Metals of the Barths. Boron B n Aluminium, ....... A 1 , . 27 Scandium, Sc 44 Yttrium, Y 92 Lanthanum, La 139 Cerium, Ce 141 Didymium, D 145 Samarium, Sm 150 Erbium E x6B Ytterbium, Yb '..... 173 Trivalent, though in compounds two atoms go together, forming a * Red Fire : Strontium nitrate, 800 grains; sulphur, 225 grains; potassium chlorate, 200 grains, and lampblack, 50 grains. | Green Fire : Barium nitrate, 450 grains; sulphur, 150 grains; potassium chlorate, 100 grains, and lampblack, 25 grains. For lecture-room experiments the following, without sulphur, are preferable: Green Fire: Two parts barium nitrate, two parts potassium chlorate, and one part ground shellac. Red Fire ; Two parts strontium nitrate, two parts potassium chlorate, and one part ground shellac. The ingredients should be dry, powdered separately, and mixed with as little friction as possible. 76 ESSENTIALS OF CHEMISTRY. sexivalent radical, as A12C16. Boron is so weakly positive that it is a non-metal. Aluminium is the most important member of this group, the others being rare metals associated with -it in various minerals. Their oxides and hydrates are of a neutral or earthy character. BORON occurs as a constituent of boracic acid and borax (sodium borate, U. S. P.). Has two allotropic forms, amorphous and crystal- line, corresponding to coal and diamond. Forms only one oxide (BoOa), which, combining with water, forms an acid:— B203 -|- 3H20 = H6B206 = 2H3B03—Boric acid. Boric or Boracic Acid occurs as pearly scales, soluble in water; feebly acid ; an unirritating antiseptic. Boiled with glycerine it was sold as boroglyceride, or mixed with borax as rex magnus to preserve foods, especially milk and meats. Test.—Compounds of boron, especially when moistened with sul- phuric acid, color the flame green. ALUMINIUM.—Never found free, but in the abundance and dis- tribution of its compounds it ranks next to oxygen and silicon—third among the elements and first among the metals. Isolated with difficulty, and therefore costly. Bluish-white metal, ductile and very light; does not tarnish in the air. With copper it forms a golden- yellow alloy, known as aluminium bronze. Aluminium Chloride—ALC1o.—Prepared industrially in the manu- facture of aluminium. A soluble, astringent salt. It absorbs and combines with H2S, PH3, and NH3. An impure solution is sold as a disinfectant under the name chloralum. Aluminium Sulphate—AL^SCU—Made by treating white clay with H2S04. It has properties similar to the above. Alum—Alumen.—An alum is a double sulphate of a trivalent and a univalent radical. Its constitution may be expressed thus :—• R2lIt3S04.R2IS04, or 2RmRi2S04. The trivalent radical (R111) may be Al, Fe, Cr, or Mn. The uni- valent radical (R1) may be K, Na, NH4, etc. So, by different com- binations of these radicals a variety of alums may be formed. The old potash alum (A123504.K2504) is giving place in the arts to the cheaper ammonium alum (A123504.(NH4)2504). The ammonio-ferric PART I.—INORGANIC CHEMISTRY. 77 ahan (Fe23504.(NH4)2504) is also much used in medicine. Burnt alum, alumen exsiccahim, is a white amorphous powder obtained by heating alum until its water of crystallization is driven off. Alum, like other salts in which the acidulous radical predominates, is astrin- gent ; burnt alum, on account of its avidity for water, is a mild escharotic. Aluminium Silicates.—Very abundant, as granite, clay, sand, etc. Clay is usually of a reddish or brown color from admixture of oxides of Fe, etc. Pure white clay (kaolin) is used in the arts to make porcelain, and in medicine as a vehicle for the external application of acids, etc. CERIUM is a rare metal. One of its salts, the oxalate, is used as a sedative to irritable stomachs, especially in the vomiting of preg- nancy. When pure it is a very efficient remedy ; but the commercial article is liable to contain salts of lanthanum, didymium, and other allied metals. The other members of this group possess little interest for the medical student. IX. The Zinc Group. Zinc, Zn 65.2 Cadmium, Cd 112 Bivalent; bluish-white metals, closely allied in sources and prop- erties. ZlNC.—When heated in air, zinc burns with an intense bluish- white light, forming clouds of oxide. It tarnishes quickly in air or water, but becomes coated with a film of oxide that protects it from further corrosion. Iron coated with zinc (“galvanized iron”) will withstand exposure to the weather an indefinite time. Alloyed with copper, zinc forms brass. Pure H2SQ4 is unaffected by pure zinc or zinc coated with mercury (amalgamated), unless it form a galvanic circuit* Commercial zinc is rapidly attacked by most acids. * Experiment.—lnto a large test-tube containing bits of zinc pour dilute sulphuric acid; there is a prompt effervescence of hydrogen. Add a little mercury, and agitate; the action ceases. Drop in a piece of copper; it begins again. 78 ESSENTIALS OF CHEMISTRY. Zinc Sulphate—ZnS04 White Vitriol—is made thus Zn + H2S04 = ZnS04 + H2. White, soluble salt, resembling MgS04 in appearance ; astringent and emetic. Zinc Chloride—ZnCl2.—Made: Zn -f 2HCI = ZnCl2 + H2. A white deliquescent salt; strongly astringent; severe caustic. Used as an injection to preserve anatomical subjects. Each of the following mixtures forms a hard, white mass, used for filling teeth;— (a) A strong solution of zinc chloride with zinc oxide. (.b) A strong solution of magnesium chloride with magnesium oxide. (c) Zinc oxide with phosphoric acid (zinc phosphate). Zinc Carbonate—ZnC03—is a white, insoluble powder made by precipitation ZnS04 + Na2C03 = Na2S04 + ZnC03. Used in the arts (zinc white') in place of lead carbonate, for it is not blackened by H2S; in medicine as a dusting powder for excori- ated surfaces, and in ointment. Zinc Oxide—ZnO—is prepared either by burning metallic zinc* or heating the carbonate, ZnC03 = ZnO -f- C02. It is a yellowish-white powder, used externally in ointment, inter- nally as a tonic and astringent, especially in the night-sweats of phthisis and diarrhoea of children. Zinc Sulphide—ZnS—is precipitated whenever a solution of a zinc salt is added to the solution of a soluble sulphide, unless the solution be acid in reaction. It is the only white sulphide, therefore a test for zinc. Poisoning.—All the salts of zinc that are soluble in the digestive fluids act as irritant poisons. Sodium chloride and organic acids dis- solve metallic zinc, therefore food kept in galvanized iron vessels is more or less poisonous, especially since commercial zinc usually con- * Experiment.—Place bits of zinc in a hessian crucible and heat strongly over a triple burner. The. metal is volatilized, and the vapor igniting bums with an intense bluish-white flame, yielding white flakes of zinc oxide, the lana philosophic a (philosopher’s wool) of the older chemists. PART I.—INORGANIC CHEMISTRY. 79 tains traces of arsenic. For this reason articles intended for toxico- logical analysis should never be kept in jars with zinc caps. CADMIUM resembles zinc in its properties and uses, except that its sulphide is yellow and insoluble in acid solutions. X. The Iron Group. Chromium, Cr 52.2 Manganese, Mn 55.0 Iron, Fe 56.0 Cobalt, Co 58.8 Nickel, Ni 58.8 These are hard metals and all more or less magnetic. By a variation in valence they form two classes of compounds : One in which the atom is bivalent, as in ferrous chloride (FeCk) ; the other in which the atom is irivalent, as in ferric chloride (Fe2Cl6). With oxygen they form acidulous radicals, which form the chromates, man- ganates, and ferrates, with the stronger bases. CHROMIUM.—So named because all its compounds are colored. The metal is of but little use. Its compounds are of great importance to the chemist and of considerable utility in the arts, but few are used in medicine. Chromium Trioxide—Cr03—is made by treating a strong solution of potassium bichromate with sulphuric acid, thus:— K2Cr207 + H2S04 = K2S04 -f H2Cr04 + Cr03. The Cr03 separates in crimson prisms. It is a powerful oxidant and a caustic. Sometimes improperly called chromic acid. Chromates.—The principal ones are potassium chromate, K2Cr04, a valuable test reagent, and lead chromate, PbCr04, a yellow pig- ment. Bichromates are not regular acid or bi-salts, but compounds of a ESSENTIALS OF CHEMISTRY. chromate and chromium trioxide. The most important of these is potassium bichromate, K2Cr207, or K2Cr04.Cr03. It forms large, red, soluble crystals. It is added to the sulphuric acid in batteries to oxidize* the nascent hydrogen. Chromates may be recognized by their color and by the yellow pre- cipitate on the addition of lead acetate. MANGANESE resembles iron in its properties. Used to alloy iron in the preparation of certain kinds of steel. Its most abundant ore is the Manganese Dioxide—Mn02—Black Oxide of Manganese—an insoluble steel-gray powder that readily gives up its extra atom of O. Used in large quantities in the preparation of chlorine and oxygen gas. Manganous Sulphate—MnS04. Mn02 + H2S04 = MnS04 + HaO +O. A soluble, rose-colored salt. Manganous Sulphide—MnS—is precipitated whenever a solution of a salt of manganese is treated with NH4HS. It is the only flesh- colored sulphide, hence its formation is a test of manganese. Manganates.—lf a mixture of KHO, KCI03, and Mn02 be heated together, there results a green mass of potassium manganate, K2Mn04. If this be dissolved in distilled water, it forms a green solution, which, on boiling, or even standing awhile, is changed to a purple, due to the formation of potassium permanganate, K2Mn208. The permanganate f gives up its oxygen so readily to organic matter, * Experiments.—The oxidizing action of the chromic salts can be shown in a number of reactions. (a) Any organic substance, as sugar, oxalic acid, or a chip of wood, boiled in the sulphuric acid and bichromate mixture, is oxid- ized, disappearing completely, with evolution of carbon dioxide, (b) Rinse out a beaker with strong alcohol, and then drop in a few crystals of chromic acid. The thin layer of alcohol is ignited with the odor of aldehyde. (c) Pour a few drops of absolute alcohol on the wick of a spirit lamp, and lay on several crystals of chromic acid. It ignites. f Experiment.—Powdered potassium permanganate treated with sulphuric acid gives off ozone. (See page 19.) So powerful an oxidizer is this mixture that alcohol, ether, benzol, carbon disulphide, flowers of sulphur, tannin, etc., are ignited on contact with it. PART I.—INORGANIC CHEMISTRY. at the same time losing its purple color, that it is used as a test for organic impurity in water and as a disinfectant. Physiological.—Associated with iron (i to 20), manganese is a normal constituent of the blood corpuscles ; hence its preparations, like those of iron, are blood tonics. Valuable in amenorrhoea. IRON occurs abundantly as oxide, carbonate, and sulphide ; occa- sionally free. Preparation.—The carbonate or sulphide is first roasted until con Fig. 35. Making Reduced Iron. verted into oxide. The oxide is heated in a blast furnace with coal and fluxes (limestone and silicates). The carbon of the coal removes the oxygen from the iron, which melts and sinks beneath the melted fluxes. The fused metal is then drawn off into furrows in the sand called pigs. This is cast iron, containing 4or 5 per cent, of carbon. Wrought iron contains little or no carbon, and steel an intermediate amount. Properties.—A bluish-gray metal, sp. gr. 7.5 ; rusts (oxidizes) when exposed to moist air, or water containing air. Reduced Iron—Ferrum Redactum, iron by hydrogen, Quevenne’s ESSENTIALS OF CHEMISTRY. iron.—lt is prepared by heating ferric oxide nearly to redness in a tube through which hydrogen is passed : Fe203 +H6 = Fe2 + 3H20 * It is a very fine, dark gray powder, which, if good and fresh, will ignite on contact with a lighted taper and burn with a red glow ; f pre- scribed in pill form. CHLORIDES. Ferrous Chloride—FeCl2,—Made by adding iron to hydrochloric acid until effervescence ceases, thus : Fe + 2HCI = FeCl2 + H2. Like most ferrous salts, it is green and prone to oxidize with the forma- tion of ferric compounds. Ferric Chloride—Fe2Cl6—is made by first forming the ferrous chloride as above, and then adding nitric and hydrochloric acids. The nascent chlorine evolved by the nitro-hydrochloric acid converts the ferrous into ferric chloride, thus:— liq. ferri chloridi, U. S. P., is the aqueous solution. This, when diluted with alcohol, forms the tinct. ferri chloridi, U. S. P. If citrate of potassium or sodium be added to this tincture, the solution loses its styptic taste, does not affect the teeth, and is not incompatible with solutions containing tannin. 6FeCl2 -j- 6HCI -f- 2HNO3 = 3Fe2CI6 -(- N202 -f- 4H20. Ferrous Sulphate—FeSQ,— Copperas, Green Vitriol.—Prepared: SULPHATES. * Experiment.—With the apparatus shown in Fig. 35, hydrogen is gen- erated from sulphuric acid and zinc in the Wolff bottle, and dried by passing through the U-shaped tube containing calcium chloride. It then passes through the porcelain tube containing ferric oxide (subcarbonate, U. S. P.) and is heated to redness in the furnace. After the reduction is completed, the iron should not be exposed to the air until cool, or it will ignite spontaneously. Experiment.—Faraday used to show that it is more inflammable than gunpowder, by pouring it mixed with gunpowder upon an alcohol flame burn- ing on a white dinner-plate. The iron burns with bright scintillations, while the gunpowder falls through the flame and is only ignited when the flame dies down and reaches the surface of the plate. One part of sulphur, two of reduced iron, and three of nitre make an iron gunpowder that burns as quickly and more brilliantly than ordinary gunpowder. PART I.—INORGANIC CHEMISTRY. 83 Fe -j- H2S04 = FeS04+ H2. Soluble, green crystals efflorescing upon exposure. A cheap and excellent disinfectant, destroying organic matters by abstracting their oxygen. When given in pill form it is first exsiccated. Ferric Sulphate—Fe23S04. Tersulphate is made by adding nitro- sulphuric acid (HNO3 -f- H2S04) to a solution of the ferrous sulphate, thus:— 6FeS04 + 3H2504 + 2HNO3 = 3Fe23504 + N202 + 4H20. Its officinal solution is the liq. ferri tersulphatis. Liq.ferri sub- sulphatis, U. S. P., Monsel’s Solution, is prepared similarly to the above, using only half the quantity of sulphuric acid. Fig. 36. A Dialyser. Ferrous Hydrate—Fe2HO—is precipitated on mixing solutions of a hydrate and a ferrous salt, as— HYDRATES. A green precipitate, which soon oxidizes and becomes brown. FeS04 + 2NaHO = Na2S04 + FeaHO. Ferric Hydrate—Fe26HO.—A brownish-red, gelatinous mass, precipitated by soluble hydrates from ferric solutions, e.g. : Fe2Cl6 + 6NH4HO = 6NH4CI + Fe26H0. This is the favorite antidote for arsenic, for which purpose it must be freshly prepared and given in large doses. Ferric hydrate dissolves freely in a solution of ferric chloride, forming a dark red liquid of a styptic taste. If this liquid be put in a dialyser (Fig. 36), a vessel with a bottom of parchment or animal membrane, and suspended in water, the 84 ESSENTIALS OF CHEMISTRY. chloride passes out through the membrane into the water. When barely enough ferric chloride remains within the dialyser to hold the fer- ric hydrate in solution and the styptic taste has disappeared, the liquid is removed and sold under the name of “ Dialysed Iron.” Ferric Nitrate—Fe26NO3. Made: Fe26H0 + 6HN03 = 6H20 + Fe26N03. Liq.ferri nitratis, U. S. P., is a reddish acid liquid. Used as an astringent, especially in dysentery. Ferrous lodide—FeL.—Prepared: Fe + I 2 = Fel2. Sometimes given in pill, but better with syrup, which acts as a pre- servative as well as a vehicle. Ferrous Carbonate—FeC03—is obtained by adding a soluble (alkaline) carbonate to a ferrous salt, thus:— FeS04 + K2C03 = K2S04 + FeCO,. It is insoluble in pure water, but slightly soluble in water containing carbonic acid, as in chalybeate springs. On exposure to the air it turns red from formation of ferric hydrate ; so it is preserved by mix- ing with sugar and honey, as in the ferricurbonas saccharatus, U. S. P. Ferrous Sulphide—FeS—occurs native, but may be made by heat- ing together iron filings and flowers of sulphur. Used in the prepara- tion of HoS. Scale Compounds of Iron.—These are ferric salts, mostly with organic acids. They do not crystallize readily, but are sold as thin scales. Made by evaporating their solutions to a syrupy consistence, poured upon plates, and when dry peeled off in scales. Often other bases, as potassium or ammonium, together with alkaloids, as quinine and strychnine, are incorporated in the compound. The following are officinal: Ferri citras, ferri et ammonii citras, ferri et quinice citras, ferri et strychnia citras, ferri et ammonii tar- tras, ferri et potassii tartras, and ferri pyrophosphas. Physiological.—lron is a normal constituent of the body, especially the blood corpuscles, where it performs an important function, as is shown by the great increase of blood corpuscles and of bodily vigor attending its administration. Many of its salts, especially the ferric salts of the mineral acids, are astringent and hemostatic. Iron is eliminated by various organs, but is mainly discharged by the bowels as sulphide blackening the faeces. PART I.—INORGANIC CHEMISTRY. 85 Tests for Iron.—Ferrous salts are usually green ; with NH4HS they give a black precipitate of FeS. Ferric Salts are usually red; they give a black precipitate with NH4HS ; a black precipitate with tannic acid ; and a blood-red with sulphocyanate of potassium. COBALT.—Its chief ore is a compound with arsenic, sold under the name of cobalt or flystone, for poisoning flies. Its salts are used in preparing sympathetic ink, for when free from moisture they are deep blue, but almost colorless when moist. Writing done with a dilute solution of chloride of cobalt is invisible until warmed, when it becomes blue, the color disappearing when the paper is cooled or moistened. Test for Cobalt.—lt imparts a deep blue color to a bead of glass or borax melted in the blow-pipe flame. NICKEL.—This is a hard, grayish-white metal that does not tarnish in the air. Used to electro-plate instruments made of metals more prone to corrode, and to make cheap coin. Mixed with brass, it forms German silver. XL The Copper Group. Copper (Cuprum), Cu 63.4 Mercury (Hydrargyrum), . . Hg 200 Each of these elements is univalent and bivalent, forming two classes of compounds, “ous" and "ic." At ordinary temperatures they are acted upon but slowly by the non-oxidizing acids, as H2S04 and HCI; but HN03 attacks them vigorously. COPPER is usually found combined with sulphur, etc., but often in the metallic state, especially on the southern shores of Lake Superior. Being found free, it was among the first metals wrought by man, so the bronze preceded the iron age. Copper is a red malleable metal; an excellent conductor of electricity. Cupric Sulphate—CuS04—Blue Vitriol, Blue Stone.—Obtained as an incidental product from silver refineries, copper mines, etc.; made experimentally by heating copper with strong H2S04. Forms beautiful blue crystals, soluble in water, but insoluble in alcohol. If the crystals be heated they lose their water of crystallization and form 86 ESSENTIALS OF CHEMISTRY. a white powder, which becomes blue again upon the addition of water. Hence, used as a test for water in alcohol. Like other salts in which the acidulous radical predominates, cupi'ic sulphate is astringent and coagulates albumen. A prompt emetic, but not used as much as ZnS04, because if, by chance, it be not all ejected from the stomach, a gastro-enteritis is liable to be set up. Cupric Hydrate—CuaHO—is formed as a bluish-white precipitate whenever a soluble copper salt is treated with a soluble hydrate, thus: When heated, even under water, it decomposes— CuS04 + 2KHO =K2S04 + CU2HO. Cu2HO = CuO + H2O. Cupric Oxide—CuO—Black Oxide.—Prepared by heating copper turnings in air. It gives up its oxygen easily, hence used as an oxid- izer in organic analysis. Cuprous Oxide—Cu20—Suboxide.—Made by boiling the cupric oxide with an oxidizable substance, as glucose (copper tests for glu- cose), which is oxidized at the expense of the oxygen of the cupric oxide. The precipitate is first yellow (hydrate), but soon becomes a bright red (oxide). Cupric Subacetate or Oxyacetate—-sometimes called verdigris (green-gray)—is made industrially by exposing plates of copper to the acetic fumes of grape husks, etc. It is apt to be formed whenever fruits containing acetic acid are placed in copper vessels. Tests.—l. Plating Test. Dip into the suspected solution a more electro-positive metal, as iron, and a plating of metallic copper will be deposited on the iron, an equivalent proportion of which takes the place of the copper in the solution. 2. Sulphur Test. Add H2S or NH4HS, and if copper be present a black precipitate (CuS) will be formed. 3. Ammonia Test. Add ammonia, and if copper be present a deep blue arnmonio-salt of copper will be formed. 4. Arsenic Test. To the ammonio-salt, described above, add an aqueous solution of As203, and a green precipitate of arsenite of cop- per (Paris green) will be thrown down. 5. Glucose Test. Add KHO, (CuSO4 + 2KHO K2S04 + Cu2HO) and boil (Cu2HO = CuO -(- H2O), with a little glucose, and a yellow- ish-red precipitate (Cu2o) indicates copper. PART I.—INORGANIC CHEMISTRY. 87 It will be seen from the last two reactions, above described, that a substance acted upon characteristically by a reagent is as good a test for the reagent as the reagent is for it—i. e., arsenic and glucose, being acted upon characteristically by copper, are as good tests for copper as copper is for them. Physiological.—Canned fruits, pickles, etc., that have been colored green with copper, and food, especially if acid, that has been cooked or kept in copper vessels, are apt to produce an acute gastro-enteritis. Chronic copper poisoning, so called, is perhaps always due to other substances, as lead or arsenic, and should be treated accord- ingly. Antidotes for acute copper poisoning: Encourage vomiting and give albumen (white of egg), which combines with the copper salt to form an insoluble albuminate; or iron filings, which will precipitate the copper in metallic state. MERCURY is the only metal liquid at ordinary temperatures, and resembles silver in appearance, hence the names hydrargyrum (water silver) and quicksilver (fluid silver). It is so heavy (specific gravity 13.56) that iron and stone float upon it as corks on water. (Fig. 37 represents a marble and a ball of iron floating on mercury.) -It does not tar- nish in the air unless contaminated with baser metals; dissolves all metals, except iron, to form amalgams. Fig. 37. Uses.—Metallic mercury is used extensively in the refining of silver and gold, in thermometers and other instruments, with tin in silvering mirrors, and in many other branches of the arts. Metallic mercury, rubbed up with various excipients until globules cease to be visible, forms several officinal preparations. Rubbed with chalk it forms “ gray powder,” hydrargyrum cum creta ; with confection of roses and licorice powder it forms ‘‘blue pill " pilula hydrargyri; and with lard and suet it forms “ mercurial ointment,” ungueiitum hydrargyri. The therapeutic activity of these preparations is not due to the metallic mercury they contain, but to small quantities of mercurous oxide formed by the oxidation of the finely divided metal. So their strength varies with the thoroughness of the rubbing, the extent of the expo- sure, and the age of the preparation. 88 ESSENTIALS OF CHEMISTRY. Mercurous lodide—Hgl—Proto-iodide, Green lodide, Hydrar- gyri lodidum Viride, U. S. P.—Made by rubbing together chemical equivalents of mercury (200) and iodine (127) until they combine and form a green mass. Mercuric lodide—Hgl2—Biniodide, Red lodide, Hydrargyri lodidiwt Ruhrum.—Made like the above, except that two equivalents of iodine (twice 127) are employed. Both the iodides, being insoluble, may be precipitated by adding a solution of KI to a solution of mercurous salt for the one and a mer- curic for the other, thus : HgNOg + KI = Hgl + KNOs and Hg2N03 + 2KI = Hgl2 + 2KNO3. The mercuric iodide is dissolved by excess of either the Hg2N03 or the KI. In precipitating, mercuric iodide is first yellow, but rapidly becomes red. If some of the dry red powder be placed on a sheet of paper and warmed over a lamp, it changes back to yellow, but on shaking or rubbing the red is restored. These changes in color are due to changes in crystalline structure. Mercurous Nitrate—HgN03—is formed when mercury is treated with cold dilute nitric acidi Mercuric Nitrate—Hg2N03.—Acid nitrate of mercury is formed if the mercury be boiled with strong nitric acid. Like all nitrates, both of the above are soluble. It enters into the liq. hydrargyri nitratis, U. S. P., and “citrine ointment,” ung. hydrargyri nitratis, U. S. P. Mercurous Sulphate— Hg2S04—is made by digesting sulphuric acid with excess of mercury. Mercuric Sulphate—HgS04—is made by heating mercury with excess of sulphuric acid. A white, crystalline salt, used in some forms of galvanic batteries. When diluted with water it decomposes into an acid salt, which remains in solution, and a yellow precipitate of oxysulphate, HgS04.2Hg0, called “ turpeth mineral,” hydrargyri subsulphas jiavus, U. S. P. Mercurous Chloride—HgCl—Calomel, mild chloride, Hydrar- gyri Chloridum Mite, U. S. P.—is made by heating mercurous sul- phate with sodium chloride (Hg2S04 + 2NaCI = Na2S04 + 2HgCI), when the mercurous chloride sublimes and is condensed in a cool receiver. TART I.—INORGANIC CHEMISTRY. 89 Calomel is a white, insoluble powder. Exposed to light it is slowly decomposed (aHgCI =Hg -f HgCL). With aqua regia, and more slowly with other soluble chlorides, it is converted into mercuric chlo- ride. Calomel probably passes through the stomach unaltered, but is converted into the mercurous oxide by the alkaline fluids in the small intestine. Mercuric Chloride—HgCI2—Bichloride of Mercury, Corrosive Subli7nate—is prepared by sublimation from a mixture of mercuric sulphate and sodium chloride, thus : HgSO, + 2NaCI == Na,SO4 + HgCl2. It is crystalline and soluble, with a disagreeable styptic taste, and is very poisonous; much used in antiseptic surgery. Mercuric Ammonium Chloride—Ammoniated Mercury, White Precipitate, U. S. P.—Formed by adding ammonia to a solution of mercuric chloride; mostly used in ointment. It is a double salt of mercury and NH.,, a derivative of ammonium. Its composition is that of NH4CI, in which two atoms of H are displaced by one of Hg, forming NH2HgCI. The ammonio-sulphate of copper previously described has an analogous composition. Mercurous Oxide—HgaO—Black Oxide of Mercury—is made by treating a mercurous salt with a soluble hydrate, as 2HgCI + 2KHO = Hg20 + 2KCI + II20. It is seldom used in medicine. Mercuric Oxide—HgO—Red or Yellow Oxide.—When prepared by decomposing mercuric nitrate by heat, it is crystalline and of a red color (hydrargyri oxidum rubrum, U. S. P.) ; but when made by pre- cipitating a mercuric solution with a hydrate, HgCl, + 2KHO = HgO + 2KCI + H2O, it is an amorphous yellow powder {hydrargyrioxidumflavurn, U. S. P.). The yellow variety, being amorphous and more finely divided, is less gritty and has greater therapeutic activity. Oleate of Mercury is made by warming the yellow oxide with oleic acid. A liquid or semi-solid. Applied to the skin it is rapidly absorbed. Mercurous Sulphide—Hg2S—is an unstable compound, which ESSENTIALS OF CHEMISTRY. falls as a black precipitate when a mercurous solution is treated with a soluble sulphide. Mercuric Sulphide—HgS—falls as a black precipitate when a mercuric solution is treated with a soluble sulphide. It is found in nature in crystalline masses called cinnabar. By certain processes it may be obtained as a deep-red crystalline powder, called ver- milion. Tests.—These consist in adding to the suspected liquid solutions of salts containing radicals capable of uniting with mercury and forming precipitates of the foregoing insoluble compounds. But the galvanic test is perhaps the best for clinical purposes. On a gold or copper coin put a drop of the suspected solution acidulated with HCI, and with a piece of baser metal, as a knife blade, touch the coin through the drop of fluid. Mercury, if present, will be deposited on the coin in a silvery film. Physiological.—Acute poisoning occurs from swallowing a single large dose of some of the mercuric compounds, especially corrosive sublimate. The minimum fatal dose of corrosive sublimate is three grains ; of white precipitate and turpeth mineral forty grains. Children tolerate mercury much better in proportion to their age than adults. The symptoms are those of severe gastro-enteric irritation. Give albumin, with which it forms an insoluble compound. Iron filings also act as a chemical antidote by decomposing the salt, taking the acidulous radical and depositing the mercury in the metallic state. Chronic poisoning is often called, from its most prominent symptom, salivation or ptyalism. It usually occurs from small, but often repeated doses of the mercurous preparations, as blue pill, calomel, etc. One of the first symptoms is a delicate red line along the margin of the gums, then comes a metallic taste, abdominal pains, nausea, vomiting, dysenteric diarrhoea, profuse flow of saliva, fetid breath, fever, emacia- tion, and paralysis. Sphacelation of the mouth and lips sometimes occurs. The treatment is to stop the ingestion of poison, and give some astringent, as tannin. PART I.—INORGANIC CHEMISTRY. 91 XII. The Silver Group. Silver (Argentum), Ag 108 Gold (Auruni), Au 197 Platinum, Pt 1944 These are heavy, bright metals, not easily corroded, rare and very valuable. Silver is univalent; gold, trivalent; and platinum, quadri- valent. SILVER occurs free, but often as a sulphide associated with lead in galena. A white, malleable, ductile metal, capable of a high polish ; best known conductor of electricity ; dissolved readily by nitric, but not by hydrochloric or sulphuric acid, except by the aid of heat; does not tarnish in air unless ozone or H2S be present. Used to plate mirrors and articles made of the more corrodible metals; alloyed with copper as coin; for tubes, sutures, etc., in sur- gery, for it does not corrode and irritate the tissues. Silver Nitrate—AgN03—ArgentiNiiras, U. S. P., Lunar Caustic. Made by the action of nitric acid on silver. If coin silver be used, the solution is blue, from the presence of copper. Silver nitrate is a crystalline salt, very, soluble. Its taste is acrid, and in large doses it acts as a corrosive poison, destroying the tissues by coagulating their albumin. For use as a cautery it is fused and moulded into sticks. Silver Oxide—Ag20—is precipitated as a brown powder on treat- ing a solution of silver nitrate with caustic potash or soda (2AgN03 + 2KHO = 2KN03 + Ag20 + H3O). Slightly soluble in water. The other salts of silver are insoluble, and made by precipitating a solu- tion of silver nitrate with a solution containing the appropriate radical. Silver Cyanide—AgCN. White precipitate, soluble in ammonium hydrate. AgNOg + KCN = AgCN + KNOs. Silver Chloride—AgCl.* AgN03 + NaCl = AgCl + NaN03. * There are three insoluble chlorides, viz., PbCl2, HgCl, and AgCl. They may be distinguished by ammonia, which dissolves AgCl; blackens HgCI, and has no effect on PbCl2. 92 ESSENTIALS OF CHEMISTRY. White precipitate ; insoluble in nitric acid, but freely soluble in ammonium hydrate. Silver Bromide—Agßr. AgN03 + KBr = Agßr + KN03. Yellowish-white precipitate ; slightly soluble in ammonium hydrate. Silver lodide—Agl. Yellow precipitate; insoluble in ammonium hydrate. AgN03 + KI = Agl + KN03. Effects of Light.—Light decomposes salts of silver, especially if organic matter be present, depositing metallic silver in a fine, black powder, hence their uses in photography, and in making indelible inks, hair dyes, etc. The black stain of silver on the hands or clothes may be removed by potassium cyanide or by applying tincture of iodine and washing in ammonia-water. When persons have taken silver salts for a long time, it sometimes occurs that the tissues, especially the skin, are permanently darkened. This is due to the decomposition of the silver salt under the influence of organic matter and light. Poisoning occtirs mostly from swallowing the which is the only soluble silver salt. It is a severe corrosive poison, destroying the tissues by coagulating their albumin. Its best antidote is a soluble chloride, as common salt, which forms the insoluble silver chloride. Albumin is also a good antidote. GOLD occurs widely, but sparingly distributed; always free, mixed with sand and quartz, from which it is separated by agitation with water or by dissolving it out with mercury. It is a soft, bright, yellow metal; so malleable that it may be beaten into sheets (gold leaf) less than one two-hundred-thousandth of an inch in thickness. These transmit green light. For coinage and general use gold is usually hardened by the addition of copper or silver, the amount of which is indicated by the term carat fine. Thus, pure gold is twenty-four carat, and eighteen, sixteen, and twelve carat signify so many twenty- fourths of pure gold. Gold does not tarnish in the air; is unaffected by any single acid, but nitro-muriatic acid (aqua regia) easily dissolves it, forming auric PART I.—INORGANIC CHEMISTRY. 93 chloride, AuC13, a caustic salt, which is sometimes given as a nerve tonic and aphrodisiac. Dose, one-twentieth to one-tenth of a grain. PLATINUM? occurs free, associated with the allied metals, palla- dium, rhodium, ruthinium and iridium. Owing to its scarcity it is almost as costly as gold. Resembles silver in appearance; can be melted only with very great difficulty, and very few substances cor- rode it; hence it is used to make vessels that are to be exposed to very high heat or to contain corrosive chemicals. Platinum wire is also used in flame testing. Platinum readily dissolves in nitro-muriatic acid, forming platinic chloride, PtCl4, a valuable reagent for potassium, ammonium and alkaloids. In the table of the elements are given the names, symbols and atomic weights of many substances that are as yet of little or no medical interest, and have therefore not received special description. They are rare elements, widely distributed but in minute quantities. Some of them are of considerable scientific? interest. Molybdenum, as ammonium molybdate, forms a valuable test-reagent for phos- phoric acid. Osmium, as osmic acid, Os04, is used in microscopy. 94 ESSENTIALS OF CHEMISTRY. TABLE.—TO DETERMINE THE METALLIC RADICAL OF A SALT IN AQUEOUS OR SLIGHTLY ACID SOLUTION BY SYSTEMATIC ANALYSIS. If HC1 gave no precipitate the metal is still in the liquid; pass H2S through the solution. Precipitate Hg (ous) Pb Ag. Collect, wash, and add If H2S gave no precipitate add NH4C1, NH4H0 and NH4HS. NH4HO. Hg ppt., blackened. Pb ppt., still white. Ag ppt., dissolved. Sb and Bi may also be precipitated by HC1, but are dissolved on adding more HC1. Precipitate Cd Cu Hg(ic) Pb Bi As Sb Sn Au Pt. Collect, wash, add NH.HS Precipitate If NH4HS gave no precipitate add (NH4)2C03. Insoluble. Cd, yellow Cu I . Hg(ic) l-ti Pb | Soluble. As (ous) & ic 1 | Sn (>c) | Zn Mn CO Ni A1 Fe Cr A\ }white- Cr, green. Precipitate Ba Sr Ca Collect, wash, dissolve in HC2H302, add K2Cr04. If (NH4)2C03 gave no precipitate add (NH4)2HAs04. Bi J-0 Sb, orange. Sn (ous) J Au U Pt J 3 cial tests for each nal solution. For e previous pages. Mn, skin-tint. Ni ) Co (-black. Ppt. If no precipi- tate, test origi- nal solution in flame on loop of Pt wire. Li, crimson. Na, yellow. K, violet. Apply spe to the origi these, see th Fe J Test specially for each in original solution. See previous pages. Ppt. Ba. Sol. Sr Ca. Add dil. h2so4. Mg. Ppt. Sol. Sr.—Ca. If neither, test orig. sol. for nh4. Add hydrochloric acid. PART I.—INORGANIC CHEMISTRY. 95 TABLE.—TO DETERMINE THE ACIDULOUS (NEGATIVE) RADICAL OF AN ORDINARY SALT IN AQUE- OUS SOLUTION, POUR SMALL PORTIONS INTO FIVE TEST-TUBES, THE SOLUTION RENDERED NEUTRAL, IF NECESSARY, BY AMMONIA. THEN ADD TO EACH RESPECTIVELY A FEW DROPS OF SULPHURIC.ACID, BARIUM CHLORIDE, CALCIUM CHLORIDE, SILVER NITRATE, AND FERRIC CHLORIDE. INTERPRET THEIR EFFECTS ACCORDING TO THE FOLLOWING TABLE: H2S04 decomposes BaCl2 precipitates CaCl2 precipitates AgNOa precipitates Fe2Cl6 precipitates Not precip- itated. f With effer- vescence ofH2Sand Sulphides | S02known Sulphites -j by smell. Carbonates and C02 having no noticeable odor. (with odor of Cyanides j HCN. j with odor of . acetic acid, Acetates j especially if [warmed. (All white.) Sulphates, insol. in HC1. Sulphites 1 Effervesce Carbonates) withacids. Citrates when Tartrates! heated on lartrates (platinumoil Borates. Oxalates. Phosphates. (All white.) Sulphates, sol. in much water. Borates ! soluble Carbonates )-in Citrates j NH4C1. Phosphates. Oxalates. Sulphites. Tartrates. All sol. in acetic acid, except oxalate and some sulphate and tartrate. All sol. in HC1, except much sulphate. Citrate and tartrate char when heated on platinum foil. Carbonate and sulphite effervesce with acids, evolving H2S and S02. Borates Carbonates Chlorates ssa, «*>»• Oxalates Sulphates | Tartrates J Bromid=.{v“- Phosphates } >’'1,ow' Sulphides, black. All soluble in dilute nitric acid, except chloride, bromide, iodide, cyanide, and sulphide. Phosphates {?“■ Borates, yellowish. Oxalates, yellow. Carbonates, reddish. Acetates, red if neutral. Sulphides, black. Nitrates. Chlorates. Apply special tests, (see previ- ous pages.) 96 ESSENTIALS OF CHEMISTRY. Acetate. co N{g Aniline (Phenylamine) is a colorless liquid, but its compounds (Cgh5 N -< H [H (the aniline dyes) are coloring matters of great brilliancy.* They are sometimes contaminated with arsenic used in their manufacture. Trimethylamine is sometimes confounded with propylamine. It is a colorless, volatile alkaloid, with an ammoniacal, fishy odor. It is found in many animal and vegetable sub- stances, but obtained from pickled herring. The hydro- chloride is the salt used. Dose, ten to fifteen grains. [ ch3 N CH3 [ch3 Antifebrin (acetanilide). This is a derivative of aniline in which the acetic radical is made to displace an atom of hydrogen. A crystalline, odorless, solid, slightly solu- ble in warm water, very soluble in alcohol. In doses of five to ten grains, repeated every two or three hours, it is an anti- pyretic and sedative. It is said not to affect the healthy temperature, but to rapidly lower a fever. [CgHS N -I H [c2h3o3 * Experiment.—Dissolve a few drops of aniline in water in two test-tubes. To one add solution of chlorinated lime—a purple color is produced; to the other add some sulphuric acid and potassium chromate mixture—a blue color appears. 118 ESSENTIALS OF CHEMISTRY. Phenacetine. The formula shows that this substance is closely ( c 6h4—O C 2H5 N \ H [c2h3o ( c 6h4—O C 2H5 N \ H [c2h3o allied to acetanilide. A white crystalline powder, but slightly soluble in water. In doses of fifteen grains it causes a fall of tem- perature and a profuse sweat. Its effect is more persistent, and per- haps more dangerous than antipyrin, producing symptoms of aniline poisoning with hsemoglobinuria and jaundice. Antipyrine, CXIH12N20, a derivative of the artificial alkaloid, chinoline, is a white crystalline powder, soluble in water and alcohol, of a slight tarry taste and odor. With nitrous acid it forms a poison- ous precipitate, and is therefore incompatible with spirits of nitrous ether. The hydrochloride is the salt used. In doses of ten to fifteen grains it is a valuable antipyretic and anodyne. Alkaloids (alkali-like).—These bodies are mostly of vegetable origin and bear a close analogy to the preceding, for they are ammonia substitution compounds, alkaline in reaction, and combine with acids to form salts. Of late years chemists have made substances very similar to, if not identical with, some of the natural alkaloids; and the time seems not far distant when our most costly alkaloids will be made cheaply by artificial means. In plants alkaloids are not found free, but combined with some vegetable acid forming a salt. Their salts (except tannates) are usually soluble and intensely bitter; the free alkaloids being much less soluble, are much less bitter. Those alkaloids (as conine and nicotine) that contain no oxygen are liquid; but the great majority of them are white powders. Alkaloids are so seldom prescribed in the free state that when the simple name of an alkaloid is written in a prescription the druggist puts up its most common salt. The names of alkaloids end in ine,” and are derived from the names of the plants in which they exist or from some characteristic property. The intense effect alkaloids exert on the animal organism makes them generally the active principles of the drugs in which they are found. But the active principle of a drug is not always an alkaloid. The alkaloids include the majority of our most potent remedies and powerful poisons. Tannin is a common antidote, but most important is the prompt evacuation of the stomach and the intelligent use of physiological antagonists. The alkaloids, even those of medical interest are so numerous that PART II.—ORGANIC CHEMISTRY. to give each separate consideration would cover a great portion of the materia medica. We can mention but a few of the most important. Name. Formula. Source. Remarks. r Crystalline; morphia gives a blue Morphine Codeine Narcotine Narceine c17h19no3 Ci8H21N03 c22h23no7 c23h29no9 1 Opium « with Fe2Cl6, and a red with HN03. These alkaloids and several others exist in opium in combination with'meconic acid, which gives with Fe2Cl6 a red color not discharged by HgCl2. Apomorphine c17h17no2 Morphine Made by heating morphine with HC1; a systemic emetic. Quinine c20h24n2o2 f Quinidine “ All crystalline except quinoidine, Quinicine “ which is a resinous mass. To Quinoidine (( bark test for quinine, add chlorine Cinchonine C20H24N2O water, shake, and then add aq. Cinchonidine it ammonia; a green color. Cinchonicine a Strychnine C21H22N202 Nux vomica Crystals; gives a purple with H2S04 and K2Cr207. Brucine U Crystals; gives a red with HNOs. ) Aconitine Aconite Colchicine Ci7H19NG5 Colchicum / V Crystals; very poisonous. Veratrine Yeratrum V i Atropine Hyoscyamine Ci7H23N03 c,5h23no3 Belladonna Hyoscyamus 1 Crystals; used to dilate the f pupils. Homatropine C’lellj.^NOs Atropine Caffeine Coffee Crystals; soluble in water; weakly basic. Theine Tea Crystals; soluble in water; weakly basic; anodyne. Cocaine Coco leaves Crystals; soluble in water; weakly basic; local anaesthetic. Physostigmine c15h21n3o2 Physostigma Crystals; contracts the pupils. (Eserine) (Calabar bean) Pilocarpine cuh16n2o2 Jaborandi Crystals; a powerful diaphoretic. Urea ch4n2o Urine Crystals; may be made artificially by heating NH4CNO. Nicotine QH7n Tobacco Liquid; powerful poison. Conine c8h15n Hemlock « U it Ptomaines (tr-u/xa, a corpse).—This name is given to certain alka- loids or amines formed in animal and some vegetable bodies during putrefaction, and in some pathological conditions during life. These 120 ESSENTIALS OF CHEMISTRY. are the products of bacteria, each species producing its own peculiar ptomaine. Thus, the typhoid-bacillus produces typhotoxine; the tetanus-bacillus, tetanine, etc. Many think the symptoms of the spe- cific fevers are only the effects of the ptomaines so produced, for the characteristic symptoms of the disease may be produced by the ad- ministration of its ptomaine. The poisoning that frequently results from eating spoiled meat, fish, etc., is due to ptomaines. The symp- toms resemble those of the vegetable alkaloids, except that there is usu- ally more gastro-intestinal irritation. The fact that certain ptomaines give physiological effects and chemical tests like such alkaloids as strychnine, morphine, conine, nicotine, atropine and veratrine, is apt to, and doubtless has often, caused'the condemnation of the innocent. Among the non-poisonous ptomaines may be mentioned putrescine, cadaverine and neuridine. Among the poisonous are: choline, which acts like curarine ; muscarine, from poisonous mushroons (mus- carius); tetanine, tetanotoxine and spasmotoxine, produced by the tetanus-bacillus and causing (?) the symptoms of tetanus ; typhotoxine, produced by the typhoid-bacillus and causing (?) the symptoms of typhoid fever, and tyrotoxine, found by Dr. Vaughan in poisonous cheese and milk and causing the symptoms of cholera infantum. Leucomaines are a class of alkaloidal substances produced in the living body as the result of fermentative changes or of the processes of retrograde metamorphosis, and are closely related to urea and uric acid, TJrey are eliminated in the various excreta, the urine, faeces, perspiration, etc. If retained, as in uraemia, or produced in abnor- mal amounts, as in dyspepsia, they act deleteriously on the nerve centres causing vertigo, lassitude, drowsiness, vomiting, purging and coma. Some elevate while others lower the temperature. Of the more important we may mention creatine, creatinine and xanthine. The chemistry of the ptomaines and leucomaines is new and still incomplete. PART 111. THE URINE. The urine is a fluid secreted continuously by the kidneys, and is the chief means by which the nitrogenous waste of the body is dis- charged.* A specimen, to be representative, should be a portion of the whole twenty-fours’ urine, for considerable variation in composition and properties may occur during the day. Especially is this true of traces of albumin and sugar. When this is impracticable, that passed before breakfast is generally preferable, because farthest from a meal. When significant variations during the day are suspected, several specimens may be taken at different hours. For microscopical ex- amination a few ounces of the urine in a stop- pered vial, or better still, in a covered conical glass (Fig. 40) are set aside for several hours until the sediment has settled to the bottom and can be examined. Fig. 40. G. T/£MAUN-CO. *The rationale of its secretion is one of transudation, osmosis, and cell elaboration. Owing to the resistance encountered by the blood in its exit through the efferent vessel, there is an increase of blood-pressure in the Mal- pighian tuft and a transudation of the water of the blood with some dissolved salts into the capsule. From loss of water the blood is very much thickened when it reaches the second capillary system surrounding the convoluted tubes, which contain the thin, watery transudation from the Malpighian bodies. Here are the essential elements of a complete osmometer—an animal mem- brane, composed of the thin wall of the capillary and the delicate basement membrane of the tube, with a dense fluid (the thickened blood) on one side and a thin saline solution on the other. An interchange now takes place of the water from the tube to the blood, and of the products of retrograde meta- morphosis (urea, etc.), and salts from the blood to the tubes, concentrating the fluid in the latter, making it urine, while the albuminous constituents of the blood, not being osmotic, are retained. An elaborative function has long been 122 ESSENTIALS OF CHEMISTRY, Physical Properties.—Normal urine is a transparent, aqueous fluid, of a pale yellow color, characteristic odor, acid reaction, and a specific gravity of 1020 when passed in the average quantity of about forty-five fluid ounces in the twenty four hours. This description is to be taken with much allowance, for very wide variations occur even in health. With these variations the student must become thoroughly familiar before he is capable of interpreting a specimen. Therefore the physical properties will be considered more particularly. Quantity.—ln health this depends upon (a) the amount of water ingested, and (b) its vicarious elimination by the skin, lungs, and bowels. Pathologically it is increased in diabetes, also in hysterical conditions associated with convulsions and high arterial pressure, and after the administration of diuretics. Transparency.—Normal urine is not always transparent, nor is transparent always normal. Some degree of opacity may be due to : (a) Miiczis, with some entangled epithelial cells, which may be observed in many specimens of healthy urine, especially of females, because of the larger area of mucous surface in that sex. (b) Urates (of Na, K, Ca, and Mg), which often form a precipitate in urine, especially when allowed to stand over night in a cold room. The test for this sediment is heat, which quickly dissipates it. (c) Earthy phos- phates (of Ca and Mg), which may give an opacity to normal urine, especially if it is alkaline or even weakly acid. The test for this sediment is the addition of a few drops of any acid which promptly clears it up, while heat would only increase it. (d) Fungi (bacteria, penicillia, sarcinas, etc.), especially in decomposing urine. A urine may be abnormally opaque from the above causes, or from the presence of blood or pus. When due to blood or pus the opacity is increased by heat or acids because of the precipitation of albumin always present in liquor sanguinis and liquor purls. Fhiidity.—Healthy urine is never otherwise than an aqueous fluid, attributed to the epithelial cells lining the convoluted tubes, for it was observed that whenever the tubes lost their epithelial lining (as in some forms of Bright’s disease), urea, etc., failed to be eliminated. This function of the cells may be demonstrated by injecting into the veins of a rabbit a solution of sulph-indigo- tate of sodium. If the animal be killed within a few minutes, none of the coloring matter will be found in the capsules, while the cells lining the tubes will be stained blue. If, however, an hour be allowed to elapse, even the cells will be found colorless and the coloring matter will be seen only in the urine. PART 111.—THE URINE. 123 flowing and dripping with ease; but in certain diseased conditions, abnormal quantities of mucus, or the presence of pus or fat, especially if the urine be allowed to decompose and become very alkaline, may give rise to viscidity. Color.—Healthy variations in color depend mainly upon the amount of water and the consequent degree of concentration or dilu- tion of the solid constituents. Aside from abnormal degrees of the above, pathological variations in color may be the result of (a) an increase or diminution of the normal coloring matters, as in fevers, etc.; (b) the presence of abnormal substances, as biliary and blood coloring matters. Moreover, the urine may be colored after the administration of certain drugs, as senna, santonin, rhubarb, prickly pear, etc. Odor.—When freshly passed, urine has, in addition to its charac- teristic odor, an aromatic fragrance due to certain volatile ethers. Alkaline urine has an ammoniacal odor, unless the alkalinity be due to fixed alkali, when the smell is fainty and sickening, like that of horses’ urine. Diabetic urine exhales a sweetish smell. In certain forms of dyspepsia and liver trouble, the odor of the urine is' almost pathognomonic. Medicines and certain articles of food often impart a peculiar odor, as turpentine the odor of violets, and asparagus and cauliflower a rank, disgusting smell. Reaction.—Normally the urine of the whole twenty-four hours will average an acid reaction ; but great variations occur during the day. Before meals it will have a high degree of acidity, but after eating becomes nearly neutral, or even alkaline. This is due to the ingestion of food which is largely alkaline and to the abstraction of acidulous principles from the blood to form acid gastric juice. It has also been observed that urine passed on rising in the morning is especially acid. This is probably due to the fact that during sleep less carbonic acid is exhaled from the lungs and less perspiration (acid) given off by the skin. The reaction of the urine is important to the physician, as it may favor or prevent the formation of sediments and concretions or irritation of the kidneys and bladder. The acidity of urine is due, not to free acid, but to acid sodium phosphate (NaH2PO4) occurring in con- sequence of carbonic, uric, and hippuric acids, seizing on to a portion of the sodium of the basic phosphate. An acid fermentation, attended with decomposition of mucus and coloring matters and a production of acetic and lactic acids, some- 124 ESSENTIALS OF CHEMISTRY. times occurs in urine that has stood for some time at a moderate tem- perature. After a while, more quickly in warm weather, the alkaline fermentation begins, caused by the development of the micrococcus Fig. 41. Acid Fermentation. Fig. 42. Alkaline Fermentation. urese (Pasteur). The urea is converted into ammonium carbonate, thus:— CH4N20 + 2H20 = (NHJ2CO3. PART 111.—THE URINE. 125 This gives the urine an ammoniacal odor and alkaline reaction, and it becomes opaque from the precipitation of urate of ammonium and the earthy phosphates and the development of bacteria. Pus and blood, or vessels tainted with urine previously fermented, greatly hasten this change. The reaction is determined by litmus paper. If acid, the blue litmus is turned red; if alkaline, the red litmus is turned blue; if neutral, there is no change in either. If alkalinity be due to ammonia (volatile alkali), the blued paper gets red again on drying. Specific Gravity.—Though the average specific gravity is 1020, it exhibits, even in health, great variations, the extremes being 1002 after copious use of water and diuretics, and 1040 after abstinence from fluid and the elimination of water through other means, as pro- fuse perspiration or copious diarrhoea. The amount of solids varying but little in health, fluctuations in specific gravity are due mainly to variations in the amount of water, and, as long as the inverse propor- tion between specific gravity and volume of urine is preserved, varia- tions need cause no alarm. Specific gravity is usually measured by an instrument called a hydrometer or urinometer (Fig. 43), which is a hollow glass float, weighted with mercury and having a long, graduated neck. The graduation begins above at 1000, because the heavier the urine the less deeply will the instrument sink and the further the neck will pro- trude from the surface. It is well to test a new urinometer by immer- sing it in water at 6o° F. (15.50 C.), into which it should sink to oor 1000 on the scale. Urinometers are usually provided with a cylinder or jar, as shown in the figure, but a large test-tube will answer. This is about three-fourths filled ; the urinometer is then introduced, and when still, the specific gravity is read off. The cylinder or test-tube should not be too narrow, lest the urinometer be attracted to and catch against the sides, and not rise as high or sink as low as it should. The fluid being attracted up around the stem, the reading should be made not along the line c d, as in the diagram, suggested by Dr. Leffmann, of Philadelphia, but a b, Fig. 43. ) sarcinae; (B5“ See pages 14 and ig for list of ? Quiz- Comp ends ? 12 STUDENTS’ TEXT-BOOKS AND MANUALS. PRACTICE. Taylor. Practice of Medicine. A Manual. By Frederick Taylor, m.d., Physician to, and Lecturer on Medicine at, Guy’s Hospital, London ; Physician to Evelina Hospital for Sick Chil- dren, and Examiner in Materia Medica and Pharmaceutical Chemistry, University of London. Cloth, 4.00 Roberts’ Practice. New Revised Edition. A Handbook of the Theory and Practice of Medicine. By Frederick T. Roberts, m.d. ; M.R.C.P., Professor of Clinical Medicine and Therapeutics in University College Hospital, London. Seventh Edition. Octavo. Cloth, 5.50; Sheep, 6.50 Hughes. Compend of the Practice of Medicine. 4th Edi- tion. Two parts, each, Cloth, 1.00; Interleaved for Notes, 1.25 Part i.—Continued, Eruptive and Periodical Fevers, Diseases of the Stomach, Intestines, Peritoneum, Biliary Passages, Liver, Kidneys, etc., and General Diseases, etc. Part ii.—Diseases of the Respiratory System, Circulatory System and Nervous System; Diseases of the Blood, etc. Physician’s Edition. Fourth Edition. Including a Section on Skin Diseases. With Index, x vol. Full Morocco, Gilt, 2.50 Tanner’s Index of Diseases, and Their Treatment. Cloth, 3.00 Wythe’s Dose and Symptom Book. Containing the Doses and Uses of all the principal Articles of the Materia Medica, etc. Seventeenth Edition. Completely Revised and Rewritten. Just Ready. 32m0. Cloth, 1.00; Pocket-book style, 1.25 PRESCRIPTION BOOKS. Pereira’s Physician’s Prescription Book. Containing Lists of Terms, Phrases, Contractions and Abbreviations used in Prescriptions Explanatory Notes, Grammatical Construction of Prescriptions, etc., etc. By Professor Jonathan Pereira, m.d. Sixteenth Edition. 32m0. Cloth, 1.00 ; Pocket-book style, 1.25 Stewart’s Compend of Pharmacy. Based upon Remington’s Text-Book of Pharmacy. Second Edition, Revised. Cloth, 1.00 ; Interleaved for Notes, 1.25 PHARMACY. SKIN DISEASES. Anderson, (McCall) Skin Diseases. A complete Text-Book, with Colored Plates and numerous Wood Engravings. Bvo. Just Ready. Cloth, 4,50 ; Leather, 5.50 Van Harlingen on Skin Diseases. A Handbook of the Dis- eases of the Skin, their Diagnosis and Treatment (arranged alpha- betically). By Arthur Van Harlingen, m.d., Clinical Lecturer on Dermatology, Jefferson Medical College; Prof, of Diseases of the Skin in the Philadelphia Polyclinic. 2d Edition. Enlarged. With colored and other plates and illustrations, xamo. Cloth, 2.50 Bulkley. The Skin in Health and Disease. By L. Duncan Bulkley, Physician to the N. Y. Hospital. Ulus. Cloth, .50 4®* See pages 2 to 3 for list of New Manuals. SURGERY. Caird and Cathcart. Surgical Handbook for the use of Practitioners and Students. By F. Mitchell Caird, m b., f.r.c.s., and C. Walker Cathcart, m.b., f.r.c.s.. Asst. Sur- geons Royal Infirmary. With over 200 Illustrations. 400 pages. Pocket size. Leather covers, 2.50 STUDENTS’ TEXT-BOOKS AND MANUALS. 13 Jacobson. Operations in Surgery. A Systematic Handbook for Physicians, Students and Hospital Surgeons. By W. H. A. Jacobson, b.a., Oxon. f.r.c.s. Eng.; Ass’t Surgeon Guy’s Hos- pital ; Surgeon at Royal Hospital for Children and Women, etc. 199 Illustrations. 1006 pages. Bvo. Cloth. 5.00; Leather, 6.00 Heath’s Minor Surgery, and Bandaging. Ninth Edition. 142 Illustrations. 60 Formulae and Diet Lists. Cloth, 2.00 Horwitz’s Compend of Surgery, including Minor Surgery. Amputations, Fractures, Dislocations, Surgical Diseases, and the Latest Antiseptic Rules, etc., with Differential Diagnosis and Treatment. By Orville Horwitz, b.s., m.d., Demonstrator of Surgery, Jefferson Medical College. 3d edition. Enlarged and Rearranged. 91 Illustrations and 77 Formulae, izrao. Cloth, 1.00; Interleaved for the addition of Notes, 1.25 Walsham. Manual of Practical Surgery. For Students and Physicians. By Wm. J. Walsham, m.d., f.r c.s.. Asst. Surg. to, and Dem. of Practical Surg. in, St. Bartholomew’s Hospital, Surgeon to Metropolitan Free Hospital, London. With 236 Engravings. See Page 2. Cloth, 3.00; Leather, 3.50 Acton. The Reproductive Organs. In Childhood, Youth, Adult Life and Old Age. Seventh Edition. Cloth, 2.00 URINE, URINARY ORGANS, ETC. Beale. Urinary and Renal Diseases and Calculous Disorders. Hints on Diagnosis and Treatment, ismo. Cloth, 1.75 Holland. The Urine, and Common Poisons and The Milk. Chemical and Microscopical, for Laboratory Use. Illus- trated. Third Edition. i2mo. Interleaved. Cloth, x.oo Ralfe. Kidney Diseases and Urinary Derangements. 42 Illus- trations. i2mo. 572 pages. Cloth, 2.75 Legg. On the Urine. A Practical Guide. 6th Ed. Cloth, .75 Marshall and Smith. On the Urine. The Chemical Analysis of the Urine. By John Marshall, m.d.. Chemical Laboratory, Univ. of Penna; and Prof. E. F. Smith, ph.d. Col. Plates. Cloth, 1.00 Thompson. Diseases of the Urinary Organs. Eighth London Edition. Illustrated. Cloth, 3.50 Tyson. On the Urine. A Practical Guide to the Examination of Urine. With Colored Plates and Wood Engravings. 6th Ed. Enlarged, nmo. Cloth, x.50 Van Niiys, Urine Analysis. Ulus. Cloth, 2.00 Bright’s Disease and Diabetes. Ulus. Cloth, 3.50 Hill and Cooper. Student’s Manual of Venereal Diseases, with Formulae. Fourth Edition, xamo. Cloth, 1.00 VENEREAL DISEASES. Durkee. On Gonorrhoea and Syphilis. Illus. Cloth, 3.50 See pages 14 and jj for list 0/ ? Quiz-Comp ends ? NEW AND REVISED EDITIONS. PQUIZ-COMPENDS? The Best Compends for Students’ Use in the Quiz Class, and when Pre- paring for Examinations. Compiled in accordance with the latest teachings of promi- nent lecturers and the most popular Text-books. They form a most complete, practical and exhaustive set of manuals, containing information nowhere else col- lected in such a condensed, practical shape. Thoroughly np to the times in every respect, containing many new prescriptions and formulse, and over two hundred and fifty illustrations, many of which have been drawn and engraved specially for this series. The authors have had large experience as quiz-masters and attaches of colleges, with exceptional opportunities for noting the most recent advances and methods. Cloth, each $l.OO. Interleaved for Notes, $1.25. NO.l. HUMAN ANATOMY, “ Based upon Gray.” Fourth Edition, including Visceral Anatomy, formerly published separately. Over 100 Illustrations. By Samuel O. L. Potter, m.a., m.d., late A. A. Surgeon U. S. Army. Professor of Practice, Cooper Medical College, San Francisco. Nos. 2 and 3. PRACTICE OF MEDICINE. Fourth Edi- tion. By Daniel E. Hughes, m.d., Demonstrator of Clinical Medicine in Jefferson Medical College, Philadelphia. In two parts. Part I.—Continued, Eruptive and Periodical Fevers, Diseases of the Stomach, Intestines, Peritoneum, Biliary Passages, Liver, Kidneys, etc. (including Tests for Urine), General Diseases, etc. Part ll.—Diseases of the Respiratory System (including Phy- sical Diagnosis), Circulatory System and Nervous System; Dis- eases of the Blood, etc. *** These little hooks can be regarded as a full set of notes upon the Practice of Medicine, containing the Synonyms, Definitions, Causes, Symptoms, Prognosis, Diagnosis, Treatment, etc., of each disease, and including a number of prescriptions hitherto unpub- lished. No. 4. PHYSIOLOGY, including Embryology. Fifth Edition. By Albert P. Brubaker, m.d., Prof, of Physiology, Penn’a College of Dental Surgery ; Demonstrator of Physiology in Jefferson Medical College, Philadelphia. Revised, Enlarged and Illustrated. No. 5. OBSTETRICS. Illustrated. Fourth Edition. By Henry G. Landis, m.d., Prof, of Obstetrics and Diseases of Women, in Starling Medical College, Columbus, O. Revised Edition. New Illustrations. BLAKISTON’S ? QUIZ-COMPENDS ? Continued. Bound in Cloth, $l.OO. Interleaved, for Notes, $1.25 No. 6. MATERIA MEDICA, THERAPEUTICS AND PRESCRIPTION WRITING. Fifth Revised Edition. With especial Reference to the Physiological Action of Drugs, and a complete article on Prescription Writing. Based on the Last Revision of the U. S. Pharmacopoeia, and including many unofficinal remedies. By Samuel O. L. Potter, m.a., m.d., late A. A. Surg. U. S. Army; Prof, of Practice, Cooper Medical College, San Francisco. Improved and Enlarged, with Ipdex. No. 7. GYNECOLOGY. A Compend of Diseases of Women. By Henry Morris, m.d.. Demonstrator of Obstetrics, Jefferson Medical College, Philadelphia. No. 8. DISEASES OF THE EYE AND REFRACTION, including Treatment and Surgery. By L. Webster Fox, m.d.. Chief Clinical Assistant Ophthalmological Dept., Jefferson Med- ical College, etc., and Geo. M. Gould, m.d. 71 Illustrations, 39 Formula;. Second Enlarged and Improved Edition. Index. No. 9. SURGERY. Illustrated. Third Edition. Including Fractures, Wounds, Dislocations, Sprains, Amputations and other operations; Inflammation, Suppuration, Ulcers, Syphilis, Tumors, Shock, etc. Diseases of the Spine, Ear, Bladder, Tes- ticles, Anus, and other Surgical Diseases. By Orville Horwitz, a.m., m.d.. Demonstrator of Surgery, Jefferson Medical Col- lege. Revised and Enlarged. 77 Formulae and 91 Illustrations. No. 10. CHEMISTRY. Inorganic and Organic. For Medical and Dental Students. Including Urinary Analysis and Medical Chemistry. By Henry Leffmann, m.d., Prof, of Chemistry in Penn’a College of Dental Surgery, Phila. Third Edition, Revised and Rewritten, with Index. No. ix. PHARMACY. Based upon “ Remington’s Text-book of Pharmacy.” By F. E. Stewart, m.d., ph.c., Quiz-Master at Philadelphia College of Pharmacy. Second Edition, Revised. No. 12. VETERINARY ANATOMY AND PHYSIOL- OGY. 29 Illustrations. By Wm. R. Ballou, m.d., Prof, of Equine Anatomy at N. Y. College of Veterinary Surgeons. No. 13. DISEASES OF CHILDREN. By Dr. Marcus P. Hatfield, Prof, of Diseases of Children, Chicago Medical College. Bound in Cloth, $l. Interleaved, for the Addition of Notes, $1.25. These books are constantly revised to keep up with the latest teachings and discoveries, so that they contain all the new methods and principles. No series of books are so complete in detail, concise in language, or so well printed and bound. Each one forms a complete set of notes upon the subject under consideration. Descriptive Circular Free. NOW READY. A NEW MEDICAL DICTIONARY GEORGE M. GOULD, Ophthalmic Surgeon, Philadelphia Hospital, etc. AN ENTIRELY NEW BOOK. Based On Recent Medical Literature. Small Octavo. g2O Pages. Handsomely Printed. Bound in Half Dark Leather, $3.25. Half Morocco, Thumb Index, $4.25. SEND FOR SPECIAL CIRCULAR. |UST PUBLISHED. A NEW Medical Dictionary, From PROF. J. M. DaCOSTA. . “ / find it an excellent work, doing credit to the learning and discrimination of the author. OPHTHALMIC SURGEON TO THE PHILADELPHIA HOSPITAL, CLINICAL CHIEF OPHTHALMOLOGICAL DEPT. GERMAN HOSPITAL, PHILADELPHIA. GEORGE M. GOULD, A.8., M.D., A compact, concise Vocabulary, including all the Words and Phrases used in medicine, with their proper Pronunciation and Definitions, BASED ON RECENT MEDICAL LITERATURE. It is not a mere compilation from other dictionaries. The definitions have been made by the aid of the most recent standard text-books in the various branches of medicine, and it will therefore meet the wants of every physician and student. It includes SEVERAL THOUSAND WORDS NOT CON- TAINED IN ANY SIMILAR WORK. It is printed on handsome paper, made especially for the purpose; from a new type selected on account of its clear, distinct face, and is bound so that it will lie open at any page. Small Octavo, 520 Pages, Half-Dark Leather, $3.25. With Thumb Index, Half Morocco, Marbled Edges, $4.25. It may be obtained through Booksellers, Wholesale Druggists and Dental Depots everywhere. [over] Gould’s New Medical Dictionary. The compact size of this Dictionary, its clear type, and its accuracy are unfailing pointers to its coming popularity/’ —-John B. Hamilton, Sr,rgeon-General U. S. Marine Hospital Service. There has been no dictionary accessible to the physician and student that has kept pace with the coinage of new words and terms during the past ten years. The growth of specialism in itself has increased the vocabulary by some thousands of words; and yet the busy practitioner or student has been offered no com- pact, thorough dictionary to which he could turn for a definition absolutely necessary to the proper understanding of the article he might be reading. This expressed want has led to the preparation of this work. The aim has been to prepare a handbook of sufficient scope to include everything of use to the general practitioner and student, and at the same time to be a compact, handy volume, giving the exact information desired at a quick reference. The wants of the specialist have also been taken into consideration, and the seeker after more extended knowledge will find much precise information relating to his special branch, to the etymology and meaning of words, etc. i. Of the ABBREVIATIONS used in Medicine, Prefixes and Suffixes of Medical Words, etc. IT CONTAINS TABLES 2. Of the ARTERIES, with the Name, Origin, Distribution and Branches of each. , 3. Of the BACILLI, giving the Name, Habitat, Characteristics of the Cultures (upon slides, gelatin, gelose, potato and bouillon). De- scription of the Cellules, the Influence of Oxygen and Heat, the Physiological Action, and Sundry Observations, 4. Of GANGLIA, with the Name, Location, Roots and Distribution of each. 5 Of LEUCOMAINES, giving the Name, Formula, Discoverer, Source and Physiological Action. 6 Of MICROCOCCI, giving the same information as in the case of the Bacilli. 7 Of MUSCLES, with the Name, Origin, Insertion, Innervation and Function. 8 Of NERVES, with the Name, Function, Origin, Distribution and Branches. 9. Of PLEXUSES,with the Name, Location, Derivation and Distribution. 10. Of PTOMAINES, with the Name, Formula, Discoverer, Source and Physiological Action. 11. Of COMPARISON OF THERMOMETERS ; of all the most used WEIGHTS AND MEASURES of the world; oftheMINERAL SPRINGS OF THE U. S., VITAL STATISTICS, etc., etc. Some of the material thus classified is not obtainable by Eng- lish readers in any other work. [over]