QD 41 53 EXCHANGE BULLETIN or THE UNIVERSITY OFTEXAS NO. 210 FOUR TIMES A MONTH OFFICIAL SERIES No. Chemistry in High Schools BY E, P. SCHOCH, PH. D., Professor of Physical Chemistry The University of Texas PUBLISHED BY THE UNIVERSITY OF TEXAS AUSTIN, TEXAS Entered as iccond-class mail matter at the postoffice at Austin, Texas BULLETIN OF THE UNIVERSITY OF TEXAS NO. 210 FOUR TIMES A MONTH OFFICIAL SERIES No. 64 DECEMBER 8, Chemistry in High Schools BY E. P. ^CHOCH, PH. D., Professor of Physical Chemistry The University of Texas PUBLISHED BY THE UNIVERSITY OP TEXAS AUSTIN, TEXAS Entered as second-class mail matter at the postoffice at Austin, Texas EXCHANGE Acknowledgment. Many teachers of science in Texas schools have kindly fa- vored the writer with expressions of their experience on several topics considered in this paper, and their collective advice has been carefully incorporated. Indebtedness is hereby grate- fully acknowledged. WHAT SOME OF OUR SMALLER TOWNS HAVE DONE FOR CHEMISTRY. Chemical Laboratory, Winnsboro (Texas) High School. Chemical Laboratory, Corpus Christi (Texas) High School. WHAT SOME OF OUR SMALLER TOWNS HAVE DONE FOR CHEMISTRY. Chemical Laboratory, San Marcos (Texas) High School. Chemical Laboratory (on left), Eagle Lake (Texas) High School. TABLE OF CONTENTS. PART I. EQUIPMENT. Page. The Teacher Qualifications . , 9 When and Where to Order Laboratory Supplies 9 Selection of a Room for the Laboratory. .' 10 Draft Hoods 11 Arrangement of Desks, Shelves, etc., in the Laboratory. ... 13 Details of Desks 14 Plumbing 17 Cost of Desks and Plumbing 18 Other Furniture 18 Other Desk Designs and Ready-Made Desks 18 Laboratory Burners and Fuel Supply 19 A Home-Made Gasoline Gas Machine 21 Source of Current for Electrolytic Work 23 A Modern Current Rectifier 24 A Small Cheap Current Rectifier 25 Supplies 25 Individual Outfit 26 Notes on Student's Outfit 27 General Laboratory Supplies for a Class of Twelve Students 28 Special Apparatus for Quantitative Experiments 32 Apparatus for the Section of Electrolysis, lonisation, Bat- tery Cells, etc 32 Apparatus for the Demonstration of Combining Volumes of Gases 33 Chemicals 33 Cost of Supplies 36 First Cost of Laboratory 37 Cost of Maintenance 37 PART II. PLAN AND CONDUCT OF THE COURSE. Main Object of the Course. 38 Which First Chemistry or Physics 38 Time Allowance for the Course 38 Time Allowance for Preparation of Laboratory and Lecture Table Experiments 39 Manner of Conducting Laboratory Work 40 Note on Quantitative Experiments 41 The Note Book 42 How to Begin the Course 44 TABLE OF CONTENTS. vii The Introduction of Symbols and of the Notions of Atoms and Molecules 44 The Introduction of Valence 46 The Fundamental Facts of Electrolysis and Its Introduction in the Course 46 General Discussion of an Outline for the First Part of an Introductory Course 47 Other Subjects That May Be Studied 51 Experiments Which Should Accompany, the Study of Avo- gadro 's Hypothesis 53 Chemical Information of Direct Economic Value in Texas . . 56 The University of Texas Requirements for One Unit Entrance Credit 60 Some Suitable Text-Books in Chemistry 60 PART III. OUTLINE OF AN INTRODUCTION TO THE FIRST PRINCIPLES. Oxygen 62 Hydrogen 63 Chlorine 64 Hydrogen Chloride 65 Acids. Bases and Salts 65 Hydration of Oxides 67 Solubility of Salts 68 Table of Solubilities of Salts 68 Acid Salts 69 Carbon 71 Sulphur 72. Ammonia 73 Other Optional Topics 73 lonisation, and the General Relation Between Dissolved Substances which Results in Metathetical Reactions 73 Exercise 81 Electrolysis 81 The Electro-Motive Force of Galvanic Cells 87 Action of General Reagents Upon Solutions of Salts 91 Sodium (or Potassium) Hydroxide as a Reagent 91 Ammonia as a Reagent 93 Soluble Sulphides as Reagents 95 Reactivities of Sulphides 96 Colors of Sulphides 97 List of Useful Special Properties and Reactions of Metals . . 98 Chemical Problems 99 Text-Book Reading 100 Chemical Changes Involving Oxidation and Reduction. . 101 viii TABLE OF CONTENTS. Exercise . 102 Exercise 104 Nitric Acid and Its Reduction Products 105 Note on the Oxidizing Action of Sulphuric Acid 108 Table of Electromotive Forces of Battery Poles 109 APPENDIX: THE DETAILS OF CONSTRUCTION AND ACTION OF ALTER- NATING CURRENT RECTIFIERS. The Action of the Electrolytic Cell 113 Construction of Large Electrolytic Jars for Rectifier Set No. 1 114 Design and Action of Special Transformer 115 Preparation of Solution for Electrolytic Jars 117 Manipulation of the Rectifier Set No. 1 117 A Small Cheap Rectifier 118 How to Rectify Both Alternations Without a Transformer. . 119 CHEMISTRY IN HIGH SCHOOLS. PART I : EQUIPMENT. The Teacher First and most important is the teacher. He should be above all a science man, both by inclination and training, a person whose inclinations and desires naturally make him fond of the subject. He should have had some sound training in chemistry and he should have received this training at a place where that subject is taught extensively rather than at a place where only introductory work is given, since much necessary information is absorbed from the proper "atmosphere." The extent of the person 's training in chemistry should not be less than two thor- ough courses in a first-class college or university, and in addi- tion some college training (at least one course) in either physics or in a biological subject (botany or zoology). This minimum requirement is not excessive, and frequently persons may be obtained who have had much more training than this. When and Where to Order Laboratory Supplies. The teacher who is to use the laboratory should design its equipment and order the materials. The supplies should be ordered early in June. It is quite profitable to submit large orders to several of the large import- ing houses for quotations for duty free importation and deliv- ered at your railway station. If the ordering is delayed beyond July 1st it will scarcely be possible to secure a duty free im- portation by the time the session opens in the fall. All of the large chemical dealers mentioned in the list below will import apparatus duty free if requested to do so. C. H. Stoelting Co., successors to Chicago Laboratory Supply & Scale Co., 31-45 W. Randolph St Chicago, 111. 1C ' B /** University of Texas Bulletin Central Scientific Co., 20-28 Michigan St Chicago, 111. Edward P. Martin Co., 144-146 E. Erie St Chicago, 111. L. E. Knott Apparatus Co Boston, Mass. Bausch & Lomb Optical Co Rochester, N Y. Woldenberg & Schaar, 1025 S. State St Chicago, Til. Eimer & Amend, 205-211 3rd Ave New York. The Kny-Scheerer Co., 404 West 27th St New York. E. H. Sargent & Co Chicago, 111. Selection of a Room for the Laboratory. In the selection of a room for the chemical laboratory the fol- lowing requirements must be met: The room must have such means of ventilation that all fumes may be removed rapidly and not blown into other parts of the building. Communicating with the laboratory there must be a small room where the supplies may be kept under lock, and in which the teacher may prepare experiments and solutions. The store and preparation room in turn must communicate with the class room so that the setting up and removal of lec- ture apparatus may be facilitated. If the teacher also teaches physics, then the store and prepa- ration room, as well as the class room, may be used for both subjects; but in that case the physical laboratory should be located conveniently near so that apparatus may be transferred to and from the store room without unnecessary trouble. The simplest way to meet the requirement of an adequate draft is to select a room with outer walls and windows on at least two sides. If an inner room with only one exposure (south!) must be used, then there must be provided at the far- ther end of the room from the windows one or more ventilating shafts so located that all parts of the room may be swept by the air current coming in at the windows and passing out through the shafts. Such a shaft should be one to two feet in diam- eter, and it should extend straight to the top of the building. The cover or cap on the roof over the top of this shaft should be placed very high (say twelve inches) above the top of the shaft, so that it will not impede the draft. Cornice men frequently Chemistry in High Schools 11 place this cap entirely too low. It is desirable to aid the draft by putting a gas flame or an electric fan at the bottom inlet. The form of ventilators frequently employed for ventilating ordinary buildings (that is, narrow shafts extending up in the wall, and opening into the rooms by means of a series of holes near the ceiling of the room) provides insufficient ventilation for a chemical laboratory. Draft Hoods. Obnoxious gases should not be allowed to escape freely into the atmosphere of the laboratory, but should be confined to special draft hoods. Such draft hoods must be carefully designed, other- wise they are worse than useless. The following rules should be observed in their design : 1. The hood should be as small as practicable. More and smaller hoods should be the rule. 2. The stack should be vertical throughout its entire length. 3. The hood proper, which is the part below the stack in which the apparatus is placed, should be as small as possible, so that the velocity with which the air moves through it may not be less than one-third of that with which it moves in the stack. The velocity of the draft should be great enough to lift the heaviest vapors, as, for instance, those of sulphuric acid. The inlet for the draft should be placed so that the current of air from that point to the chimney may sweep through all parts of the hood. For this purpose some stops are usually placed so as to prevent closing entirely the front door of the hood. This carefully adjusted .air inlet is essential to the successful opera- tion of the hood. The hood proper should be made as low as convenience will permit in order that the length of the slowly moving air column in the hood may be as short as possible. 4. The stack or chimney should be proportionate to the size of the hood and its cross sections should not be essentially less than one-third of the cross section of the hood proper. 5. A large gas burner should be placed at the base of the stack at a point at which it is possible to light the gas by reach- ing up into the hood with a burning candle on a stick. This gas flame need not always be lit when the hood is used, but in 12 University of Texas Bulletin la. Fiq 1 b, Examples of Faulty Hood Design. .2 a. front Examples of Proper Hood Design. From the Chemical Engineer. Chemistry in High Schools 13 quiet weather it must be lit in order to establish a sharp draft in the chimney. 6. A cap covering the stack for the purpose of keeping out rain is frequently not necessary and naturally the hood will operate best without the cap. Should a cap be found necessary, then care must be taken to place it very high above the top edge of the stack in order that the draft may not be impeded by it. Hoods are frequently constructed without any consideration of the fundamental principles to be observed in order to get a good draft. Fig. 1-a shows a hood in which the stack is en- tirely too small. In Fig. 1-b, the stack suffers from the fur- ther impediment of not being straight. These mistakes in de- sign are frequently found in actual hood construction. Fig. 2-a shows the ideal hood construction. In this hood the stack runs straight up from the closet and is large enough to provide for a sharp draft. Fig. 2-b shows a mode of construction in which the closet is built to connect with a large flue in the wall. If the diameter of this flue is large enough and a flame is placed as shown, then the hood will operate properly. All these fig- ures show the form in which the hood proper, or closet, is gen- erally built. The bottom of the closet is usually at the height of the ordinary laboratory desk and the height of the closet above the table to the contracting portion is 30 to 36 inches. The sliding front door, as well as the two sides, should be mnde of glass. Arrangement of Desks, Shelves, etc., in the Laboratory. In the placing of the laboratory desks and the supply shelves care must be taken to leave enough room for the students and instructor to pass each other and secure their supplies readily. The distance between the desks should be 4 ft. 6 in., certainly not less than 4 ft. Desks should not be placed adjacent to the wall anywhere, and a space at least 6 ft. wide should be left between the wall and the desks all around the room. However, this space is large enough to place shelves along the wall wherever desired, as well as shelves for blast lamps, balances, and other special apparatus. It is advisable to place a platform with a suitable table and a black-board in the room so situated that 14 University of Texas Bulletin the instructor may make special demonstrations before the class or give special explanations while laboratory work is in progress. In placing the desks, the manner of running the gas, water and sewer mains should be considered: It is most desirable to run these underneath the ceiling of the room below the labora- tory hence open to view and accessible for repairs at any time ; this permits the placing of the desks in any way desired. If, however, the pipe mains must be placed in the floor, then they should run so as not to cross the floor joists, and the desks must be placed accordingly. The floors over the pipe mains should not be nailed down again, but made "trap-door" fashion. There should be no high shelves or other super-structure on top of the laboratory desks which would prevent the instructor from looking across the room and seeing what is on top of any desk. The sets of reagent bottles are to be placed on small, low, movable shelves, 36 to 40 inches in length, designed to hold one row of bottles on each side, with a partition at the centre not over six inches high. The bottoms of the shelves should be six inches "clear" above the table top, but the total height of the shelf and bottles should not exceed 12 to 14 inches. Details of Desks. In most schools the laboratory is to be used by two sets of students on different days. This is easily arranged, because the amount of desk top space required by a student while working is large enough to construct two independent lockers under- neath it. In some cases instructors have even arranged for three independent lockers under this space so that the laboratory may be used by three distinct groups of students at different periods. They place four deep drawers in the space of the two lockers used in the former arrangement : three of these are locked sep- arately and in them are placed the perishable apparatus dealt out to the students individually. The fourth is not locked and in it are placed the iron ring stand, the wooden funnel stand, the burners and other large and usually non-breakable appa- ratus to be used by three students in common. This arrange- ment is not as convenient as the two student arrangement. The table top space allowed per student, while at work, Chemistry in Hiy the product of their numbers per cc., say aXb (if a is the num- ber of N per cc. and b is the number of P). When a is increased by adding a little of the saturated solution of naphthalene, then aXb is increased and the equilibrium disturbed. In order to regain equilibrium, some of the free N and free P combine, thus reducing a and b until aXb reaches its former value again; and this newly formed NP separates from the solution because the latter is saturated with NP. The corresponding change takes place when picric acid is added. Chemistry in High Schools 77 Inches Inside- D.C Main. Rheostat Conductivity Trough and Connections 78 University of Texas Bulletin In order to demonstrate the determination of the per cent of dissociation in electrolytes it is better to use the conductivity method rather than others such as a freezing point method, be- cause the connection between the experiment and the notion to be demonstrated is more direct. The following experimental arrangement makes the procedure extremely simple. This ex- periment also serves to demonstrate that the degree of ionisation of a dissolved substance increases with dilution. Have a carpenter make a wooden vessel of the shape and di- mensions given in Fig. 8. Stout cypress or soft pine boards, at least ly^ inch thick and perfectly smooth on both sides, are to be used. The vessel should be as nearly water tight as the car- penter can make it. It should then be thoroughly covered on the inside wilth melted paraffin to make it absolutely water tight, and in order that the boards may not take up any of the solutions poured in it. Now secure from a plumber or cornice maker a piece of fairly stiff sheet copper, at least 16 inches square. Cut it from one corner diagonally across to the oppo- site corner into two triangular pieces. Trim each piece if nec- essary, so that each sheet of copper may cover exactly one of the triangular ends of the trough. Bend over the excess of each plate at the top, arid clamp or place each copper plate so that it may stick closely to the wooden end of the trough it covers. Provide any suitable means for connecting the copper sheets to the wires of an electric circuit. Shake up about 200 grams of crystals of copper nitrate with about 200 cc. of water until a saturated solution is obtained. Take about 100 cc. of this solution and pour it into the trough. (1) Secure an am- meter with a total capacity of 1 or only a few amperes, (2) a volt-meter with a capacity of 3 volts or only a little more, (3) a source of direct electric current with a voltage of 2 to 10 volts, and (4) a rheostat of such capacity and construction that the current used in this experiment may be controlled to within 0.01 ampere. Connect up all this apparatus and the trough as shown in Fig. 9. Turn on the current and adjust it with a rheostat until it is 1-5 or 1-6 of the total that the ammeter can carry, but not exceeding % ampere. Note both ammeter and voltmeter readings that are then shown. In the following op- erations adjust the current so that the first voltage between the Chemistry in High Schools 79 poles (which the voltmeter indicated) is kept constant, and record in parallel columns the amount of current that flows after each dilution of the solution. Dilute the solution by add- ing measured amounts of distilled water: at first in portions of 100 cc. at a time, subsequently in larger portions of sev- eral 100 cc. At the beginning, the increase in current will be relatively large with each addition of water; then it will be- come less until finally no further increase in current is obtained. Divide the final (maximum) value into the first value (ob- tained with the original solution) : this gives the fraction of the salt present as ions in the original solution. The larger current obtained after dilution shows that the num- ber of ions arriving at the poles each second is larger after dilu- tion than before. Since the attractive force (voltage) is kept con- stant and hence the ions move at the same speed, and since on dilution they are moved parallel to the poles but remain, as a whole, at the same perpendicular distances from them, there remains no other way to account for the greater number of ions arriving at the poles after dilution except the conclusion that the number of ions is increased with dilution until all the ions pos- sible have been formed. Point out next the general conditions under which metathet- ical reactions take place, and apply this to the neutralization of an acid with a base and to the experiments given below. The following outline may be helpful here. (a) Whenever two solutions (or a solution and a slightly sol- uble solid) are mixed, the (accidental) meeting of the ca- tions from one solution with the anions from the other solution will produce at least small amounts of all the new compounds possible. (b) Note the amount (in a general way, that is, whether large or small) of each free ion in the mixture at the beginning. (c) Note the amount (i. e. large or small) of free ions that can be produced by each of the resulting substances. This is the amount of its ions with which each one would be in equilib- rium, this is very small in the case of slightly soluble and slightly dissociated substances.* *To be exact, we should consider here the product aXb, of the con- centrations of each pair of ions in place of their amounts (see page 75). 80 University of Texas Bulletin (d) If an insoluble or slightly dissociated substance is among the resulting ones, and the amount of its ions in the original mixture is much larger than the small amount with which it would be in equilibrium,* then this pair of free ions will com- bine and thus they will reduce their amounts until these are small enough for equilibrium. As thus the free ions disappear, any undissociated portions of their original compounds will ionize and be used up in turn. If one of the original substances is only slightly soluble (as in d below), then at first it will dissolve only in the small amount which saturates the solution. Then as the amount in solution is changed by reaction, more solid will dissolve, and so on. All metathetical reactions are due to such extensive reduc- tion of the original number of a certain pair of ions. A slight formation of a new compound, by ions combining to a slight extent, is not considered to be a reaction. The student should make test-tube trials with the following mixtures, and point out in each case the particular pair of ions which by combining extensively serve in a sense as the primary cause of the reaction: (a) Mix any one of several barium salt solutions with any one of several sulphates, producing in this way barium sul- phate from at least nine different mixtures. (b) Produce silver chloride from several different mixtures. (c) Produce ferric hydroxide from several different mix- tures. (d) Dissolve calcium phosphate in dilute hydrochloric acid. Here the least ionized combination is a combination of H ions with P0 4 ions, for instance, (H 2 P0 4 ) ', the calcium salt of which acid radical is soluble. (e) To the mixture obtained in (d), add ammonia to neu- tralize the acid. The ammonium phosphate thus produced will introduce many P0 4 ions, and calcium phosphate will be ob- tained as a precipitate. (f) Add solution of sodium carbonate to an aqueous solu- *Or, to be exact, the original amounts of its ions form a larger ion-product than that with which the insoluble or slightly-dissociated substance can be in equilibrium. Chemistry in High Schools 81 tion of the salt of any metal that forms insoluble carbonates. Filter and dissolve the precipitate by means of the addition of any acid that forms a soluble salt. Repeat with the salts of five other such metals. These experiments enable the student to see that whenever he knows the general relations between the ionized substances in a given mixture, then he may predict whether or not reaction will take place. Exercise. Which of the substances given below when mixed will react? give a reason for your answer. 1. Lead oxide and water. 2. Barium oxide and water. 3. Calcium carbonate and hydrochloric acid. 4. Sodium acetate solution and hydrochloric acid. 5. Copper sulphate and sodium phosphate solution. 6. Since dry sodium chloride and concentrated phosphoric acid react to form HC1 by metathesis, what are likely to be the ionisation relations in this liquid medium (concentrated phos- phoric acid) to bring about this change? 7. If the liquid medium in (6) is changed by the addition of much water, will any reaction take place extensively? What are the ionisation relations in this latter case. Other questions of similar character may ~be added. Electrolysis. The first experiment to be given is the electrolysis of hydro- chloric acid, and this should be carried out with the following apparatus: Secure two small porous cups, 3 to 4 inches high and V/2 to 2 inches in diameter. Dip the upper edges to a depth of one inch into melted paraffin in order to close up the pores in that portion of the cup. Fit rubber stoppers to the 82 University of Texas Bulletin cups. Through each stopper cut two holes, one to fit a piece of retort carbon, the other to fit the glass tubing with which the apparatus is to be connected.* Secure also a porcelain jar 4 or 5 inches in depth and 5 or 6 inches in diameter in which the porous cups are to be placed: This jar is to be filled with a saturated salt solution. Secure a tall, slender bottle, of at least one quart capacity, fit it with a rubber stop- per and two glass tubes one of which extends to the bottom of the bottle while the other terminates just below the stopper. In place of this bottle, the cylinder shown in Fig. 10 may be used. This bottle or cylinder serves to retain the chlorine, and allows an equal volume of air to be discharged in place of the chlorine. Air may be collected over water without ap- preciable loss, while chlorine cannot be collected over water because it is too soluble. Glass tubes for connections should now be bent as shown in Fig. 10. The tube fitted to the cup on the left should have no rubber joints in it, because hydrogen is to be evolved at this pole, and this gas would dif- fuse through the rubber tubing joints, and thus vitiate the ex- periment. The ends of the delivery tubes which are to be placed under the burettes should be drawn out to a small open- ing. Secure two burettes and two dishes or beakers full of water. Place the burettes in position with the lower ends ex- tending into the water, and fill them by drawing up water by means of a piece of rubber tubing. By this means they may be quickly refilled when the experiment is to be repeated. For electric connection, twist some bare copper wire around two arc light carbon rods, insert these through the rubber stoppers into the porous cups, and connect the copper wire with the ter- minals of an electric circuit which supplies direct current at a voltage of 10 to 25 volts. The electrode on the right should be connected to the positive terminal. Insert a switch in the circuit with which the current may be conveniently turned on or off. Fill the porous cups three-fourths full of a mixture of equal *For cutting holes in rubber stoppers, sharpen the edge of a cork borer by means of a file, thus producing a rough saw-like edge which is very effective, dip the end of the borer into caustic soda solution or into alcohol and proceed as with cork stoppers. Chemistry in High Schools 83 \ , ! \ + ! v o 1 4-1 O 2 s i *\ tf %l - 84 University of Texas Bulletin parts of concentrated hydrochloric acid and .water. To the cup on the right add a few crystals of potassium permangan- ate. By this means the liquid is immediately saturated with chlorine. Now join up the apparatus, and turn on the cur- rent, but allow the gases discharged by the delivery tubes to escape into the air and not collect in the burettes. After elec- trolysis has been in progress for a minute or two interrupt the current, place the burettes over the openings of the de- livery tubes and turn the current on again. It will be found that the two gases are evolved at equal rates by volume. Replace the hydrochloric acid in the cathode (negative pole) cup by dilute sulphuric acid, and turn the curent on again, i.e., repeat the experiment with the two cups thus filed with different solutions. Take the apparatus apart and place the porous cups in dis- tilled water to leach out the solution contained in the pores. Salts crystallizing in the pores would crack the jars. The last experiment indicates that the fact that hydrogen and chlorine are obtained in equal volumes cannot be due to any ' ' de- composing" of hydrochloric acid because this substance is present at the chlorine pole only, and the hydrogen liberated is certainly not obtained from the hydrochloric acid because practically no hydrogen passes through the intervening salt solution a fact that could be easily demonstrated. Before taking up the theory of the changes in this electro- lytic cell, it is necessary to present the views held at present concerning the nature of electricity. "Electrical matter'* is considered to be made up of definite unit charges, or particles of definite size. Electrically neutral substances contain an equal number of positive and negative unit charges, while ions contain an excess of positive or negative charges corre- sponding to their valence. Only negative charges are mobile, and may move from one material particle to another, or from one kind of matter to another. These mobile negative charges are called electrons. The positive charges are fixed upon each particular particle of mater, and are never transferred from one particle to an- other. Chemistry in High Schools 85 In accordance with this theory, the change at the cathode takes place as follows: the electromotive force of the battery forces electrons to pass from the cathode to those hydrogen ions which are next to the surface of the cathode (negative pole) thus changing these ions to neutral hydrogen, i. e. or- dinary, gaseous hydrogen. During the same period of time, an equal number of chlorine ions which are next to the sur- face of the anode (positive pole) give up their negative charges or electrons to the pole, thus changing to neutral chlorine or ordinary gaseous chlorine. Since a drop of any solution is electrically neutral to the out- side it follows that it must contain as many ions with postive charges as it contains ions with negative charges. When ions are discharged at the cathode, the drops of liquid which have lost the cations are momentarily left with an excess of negative charge while the opposite state of affairs exists simultaneously at the anode. These drops at the extreme end act upon their neighbors so as to obtain from them the kinds of ions that are necessary to re-establish their electric neutrality. This action communicates itself from drop to drop, from one pole to the other, and results in a slight shift of all the positive ions to- ward the cathode and all the negative ions toward the anode, as a result of which all the drops of solution regain .their elec- tric neutrality. With the discharge of the next ions at the poles, this action again takes place and so on. These are the essentials that have to be brought out in con- nection with electrolysis. The writer realizes that this treat- ment is very dogmatic, and that there has been no attempt made to show why scientists have arrived at the fundamental notions which are involved her, but he believes that the young mind is scarcely ready to follow the experimental evidence. However, it must be emphasized that the fundamental no- tions presented above are very important, and should be taught in the manner in which they are presented here. The following further experimental demonstrations of elec- trolysis should now be brought before the students. Make a 10 or 15 per cent solution of zinc chloride, and fill two porous cups with this solution. As before) place the two cups in a jar filled with salt water, and into one place 86 University of Texas Bulletin a carbon rod for an anode, and into the other a strip of sheet copper for a cathode, but leave the cups open i. e. unstop- pered. Make the necesary electrical connections and pass a current through this cell. Zinc will be deposited at the cathode, and will show itself by its gray color; while chlorine will be evolved at the anode and will reveal itself by its odor. Replace the zinc chloride in the cathode (negative pole) cup by- a copper sulphate solution, insert a clean sheet copper pole and turn on the current again. Copper will be deposited upon the cathode, while the reaction at the anode is the same as before. Show next the electrolysis of dilute sulphuric acid, using either the apparatus used above for the electrolysis of hydro- chloric acid, or a Hoffmann apparatus (see page 53 with ref- erence to the purchase of the latter.) Next fill the same apparatus with a solution of sodium sul- phate in place of sulphuric acid and show that when this solu- tion is electrolyzed, hydrogen and oxygen are liberated in the same proportion by volume as with sulphuric acid. Do not at- tempt to force a large current through this solution because it is a poor conductor, and the heat effect of the current is liable to crack the apparatus at the electrodes. Then secure two clean porous cups, the pores of which contain neither an acid nor an alkali. Fill them with a solution of sodium sulphate, insert platinum electrodes, and pfece the cups in a salt water jar as before. Before the current is turned on, show with the aid of litmus paper that the solution in the cups con- tains neither an acid nor an alkali. Then turn on the current, and show that the solution in the cathode cup becomes alkaline and the solution in the anode cup becomes acid as the result of the passing of the current. Replace the liquid in the cathode cup by a dilute solution of copper sulphate, and show that copper is deposited when the current is passed through this apparatus. These experiments show that the nature of a chemical change at a pole during electrolysis is determined by the nature of the substances present right at the pole, and that it is not depend- ent upon the nature of the substances present at the other pole or anywhere else. The electrolytic discharge of the ions zinc, copper, hydrogen, Chemistry in High Schools 87 and chlorine, requires no further explanation ; but the discharge of hydrogen out of the solution of a sodium salt and of oxygen out of a solution of a sulphate require further consideration. For an understanding of these phenomena, the following funda- mental or general facts must be recognized: 1. Water and all aqueous solutions contain the ions of water (i. e. hydrogen and oxygen ions practically), and although the amounts of these ions present are small, yet if they are ever used up, more will be formed as fast as the others are used up. 2. If two or more different ions which may be discharged are present at a pole, that particular one will be discharged which will set up the least opposing electromotive force after its discharge. In this connection see page 91, and also page 112. In other words, the ion discharged is the one requiring the least voltage for its discharge. With these two fundamental facts the discharge of hydrogen from a solution of a sodium salt is readily understood. Since the voltage required for the discharge of the sodium ions is much greater than that required to discharge the hydrogen ions which are also present, the latter are discharged; and the hydroxyl ions which are formed from the water as the hydrogen ions are discharged impart the alkalinity to the solution, or in other words, they together with the sodium ions form sodium hydroxide. Since oxygen is obtained during electrolysis at a platinum anode surrounded by a solution of a sulphate, it follows that the oxygen ions of the water are more easily discharged than the sulphate ions themselves; and the hydrogen ions which are formed from the water as the oxygen ions are discharged impart the acidity to the solution, or in other words, they together wtih the sulphate ions form sulphuric acid. The Electro-Motive Force of Battery Cells. The subject of electrolysis has not been presented fully unless the electro-motive force of the product formed at the poles has been pointed out, and since this electro-motive force when con- sidered by itself presents the cell as a galvanic battery, it is ad- visable to consider the latter subject in this connection. In the treatment of this topic as in the treatment of elee- 88 University of Texas Bulletin trolysis, it must be realized that the fundamental fact is the utter independence of the chemical actions which take place or tend to take place at the two poles. Without this a logical presenta- tion of the facts involved becomes practically impossible and it is on this account that the methods of presentation now ordi- narily employed are comparatively valueless, if not actually harmful. For the following experimental demonstration, secure four small porous cups, a jar with strong sodium chloride solution, and a sensitive voltmeter with a total range of three volts or very little more. Fit the cups with cork or rubber stoppers, which are to be perforated to fit the electrode rods mentioned below. With one porous cup make a chlorine pole by using a carbon electrode rod, filling the cup with a mixture of equal parts of concentrated hydrochloric acid and water, and then saturating the solution with chlorine (preferably by electrolysis, otherwise by adding a little potassium permanganate). For the second pole, fill the cup with zinc sulphate solution, and use a rod of zinc for the electrode rod. For the third pole, use a plat- inum pole (see page 33) or, less suitably, a carbon rod, and fill the cup with ferric chloride solution to which has been added a little of a ferrous salt (FeSOJ. For the fourth pole use a copper rod and copper sulphate solution. Put the copper sulphate pole into the salt solution jar and then put with it in turn each one of the other three poles and measure the voltage between each' of them and the copper sulphate pole. Now plot the results on a line as follows: mark a point on the line to denote the voltage of the copper- copper sulphate pole and refer to it arbitraily as a "zero voltage." Since the zinc sulphate pole is the negative pole when combined into a cell with this copper-copper sulphate pole,, and since the combination has a voltage of (about) 1.1 volts, then if we consider the force of the copper pole to be zero (arbitrarily), it follows that the voltage of the zinc pole is 1.1 volts. To represent this on our plot we measure from the copper pole point eleven spaces of any arbitrary length to the left arid mark this point (see Fig. 11). Then lay off the voltage of the chlorine pole (which is about +1.0 volt) and the voltage of the ferrous- ferric salt pole (the lat- Chemistry in High Schools 89 ter will be about +.4 to +.5 volts.) Measure next the voltage of all other combinations of any two of these poles and com- pare the values obtained with the number of unit lengths be- tween them shown in the plot. It will be found that, allow- ing for slight inaccuracies due to this rough procedure, the numbers will agree. This shows that the voltage of a cell is the sum of the two independent voltages of the poles. The primary cause of the action of a pole is the natural tendency to give up or take up electrons which some sub- stance present at the pole possesses; e. g., metals and hydrogen have a tendency to give up electrons and become positively charged ions; while non-metal (oxygen, chlorine, bromine, iodine) possess a tendency to take up electrons and become negatively charged ions. Such tendencies to give up or take up electrons depend also on the particular resulting substances formed: many resulting substances may be changed back simply by reversing the direc- tion of the electric change (i. e., reversing the direction of the current), and such substances actually exert tendencies which oppose the tendencies of the original substances. Hence the force of a pole depends on the nature of the orig- inal substance and on the particular kind of resulting substance formed during its action. Furthermore, the force- depends on the concentrations of the original and of the resulting substances : both the forward and the reverse tendencies to change increase somewhat with the amounts, per cc., of the original and of the re- sulting substances respectively. Hence, in recording measure- ments of the electromotive forces of poles, we must state the ex- act concentrations of the original and of the resulting substances in the poles when the measurements are made. Note how this is done in the table on page 110. These pole-changes may be expressed by equations. Thus for copper and zinc, which are used in the demonstrations above, the equations are : Cu 2( )=Cu** Zn 2 =Zn**. The Cu and Zn on the left side of the equation are marked with a zero to denote the fact that they are elements, i. e. they have 90 University of Texas Bulletin no ionic charges. An electron by itself is represented by ( ), and the expression 2 ( ) denotes that two electrons are given up by one atom of copper or of zinc. For the other two substances, the expressions of the pole changes are Fe***+l( )=Fe** by which is meant that these substances take up one electron per atom and become chlorine ion and ferrous ion respectively. But although a particular "pole" may have a tendency to give up electrons, it cannot actually do so unless it is connected with another pole which will take up electrons. When the "copper- copper sulphate" cup is placed in the connecting salt trough, together with the chlorine cup, and the poles are connected by a wire, then the copper actually changes to Cu** ions because the electrons given up by the copper are "transferred" in a sense through the wire connection to the chlorine pole and are there taken up by neutral chlorine, which becomes Cl' ions. In this particular combination both elements change to ions when the current flows. But when a "copper-copper sulphate" and a "zinc-zinc sulphate" pole are combined into a cell, then only one of these can change from metal to ions, i. e., only one metal can give up electrons, and the other takes up electrons and changes in the reverse manner, as represented by this ex- pression Cu**+2( )=Cu The direction in which a pole actually changes when it is com- bined with another pole to form a cell depends upon the relation of the tendencies of the two poles. This relation has been ascer- tained for all poles by ascertaining the relations of all other poles to a common * ' reference " or " zero ' ' pole : since the tendency of a single pole is independent of the other pole with which it is combined, the relation between this reference pole and all other poles gives also the mutual relations between these other poles. This procedure was illustrated in the demonstration above. All the data and further information needed in this connection is Note. An asterisk designates positive electric charges; an apos- trophe, negative charges. Chemistry in High Schools 91 given in the "Table of Electromotive Forces of Battery Poles," pag'e 109, which see. Further general consideration of the galvanic cell need not be given here, but will be given in connection with the subject of oxidation and reduction (which see). It remains only to point out the connection between battery cells and electrolytic cells. The substances produced by electrolysis (hydrogen from acids or water, copper from copper salts, chlorine from chlo- rides, etc.) naturally exert a tendency to regain the ionic state, thus each pole acts as the pole of a battery with an electromo- tive force opposite to the applied force. This is the electromo- tive force of polarization. It is absent before electrolysis be- gins and hence at the beginning any applied force, however small, will produce a current; but as soon as products of the electrode discharges are present, their electromotive force is exerted, and in order that electrolysis may continue, the ap- plied force must be larger than this opposing force. The Action of General Reagents Upon Solutions of Salts* The action of reagants, the results of which depend only on our general knowledge of solubility and ionic dissociation, has already been illustrated above. Many other reagents, however, present special conditions of ionic dissociation which require separate consideration. Of the reagents which present special conditions, only those commonly used will be considered. It will be found that the special properties of the latter are also essential parts of chemical information. 1. Sodium (or Potassium) Hydroxide as a Reagent. Experiment. Secure solutions of water soluble salts of the following cations: Cu, Ag, Zn, Cd, Hg (ous), Hg (ic), Pb, Fe (ous), Fe (ic), Ni, Al, Mg, Ca. To a few cubic centimeters of each of these solutions in a test-tube add a few drops of sodium hydroxide solution, shake in order to mix thoroughly, observe the effect, add a few drops more of the reagent, and so continue until the reagent has been *Read pages 49 and 50 in this connection. 92 University of Texas Bulletin added in a relatively large amount, say about twice as much reagent by volume as of salt solution, if the two solutions are of approximately equivalent concentrations. Compare the re- sults with the statements in the table below. General Facts. By strict metathetical reaction the hydrox- ides should be obtained in the mixtures below. However, in many cases the substances finally obtained are not the hydrox- ides of the metals but substances derived from them through one or both of the two following changes: (a) Dehydration, complete (Hg(OH) 2 to HgO) or partial (Cn(OH) 2 toCu 8 2 (OH) 2 ). (b) Dissolution of the precipitated hydroxide by excess of the reagent, thus showing that the precipitated hydroxide func- tionates as an acid, e. g. Zn(OH) 2 dissolves in excess of NaOH solution as per equation: H 2 Zn0 2 +NaOH= Na 2 ZnO,+2H 2 0. The formula for zinc hydroxide is thus written to suggest its functionating as an acid. Table showing Results of Action of NaOH (or KOH) Solu- tions upon ' * Water Soluble ' ' Salts of the Following Metals : Ba Ba(OH) 2 is precipitated only from concentrated solu- tions because it is soluble in 20 parts of water white. Sr Ppt. Sr(OH) 2 sol. in 60 parts of H 2 white not pre- cipitated from dilute solutions. Ca Ppt. Ca(OH) 2 sol. in 700 parts H 2 white not pre- cipitated from very dilute solutions. Mg Ppt. Mg(OH) 2 sol. in 6000 parts H 2 white. Al Ppt. A1(OH) 3 white, sol. in excess of reagent, giving NaA10 2 . Zn Ppt. Zn(OH) 2 white, sol. in excess Na 2 Zn0 2 . Pb Ppt. Pb(OH) 2 white, sol. in excess Na 2 Pb0 2 . Ferrous Fe(OH) 2 white ppt., darkens on exposure to air (oxidizes). Ferric Fe(OII) 3 reddish brown, flocculent ppt. Mn Ppt. Mn ( OH ) 2 "flesh" colored, darkens on exposure to air. Ni Ppt. Ni(OH) 2 pale green. Chemistry in High Schools 93 Cu in cold solution, Cu(OH) 2 bluish white ppt., soluble in large excess of reagent; in hot solution, CuO black ppt. Cd Cd(OH) 2 , white ppt. Bi Bi(OH) 3 , white ppt. Ag Ppt. Ag 2 greyish brown. Hg( ous) Ppt. Hg 2 black. Hg (ic) Ppt. HgO yellow in order to avoid the forma- tion of different colored basic salts, the Hg salt solu- tion must be poured into the reagent. Exercise (a). In what way can it be readily ascertained whether or not a substance has been completely precipitated, say Fe(OH) 3 by NaOH? Hence would sodium hydroxide be employed as a precipitant for aluminium if it be required that precipitation be complete? From the metals given in the table above select those for which sodium hydroxide may thus be used as a precipitant. (b) Since it is easy to convert a hydroxide completely to any salt by treatment with the necessary acid, and there is no troublesome ( ?) by-product formed, it is plain why the pre- cipitation of metals in the form of hydroxides is very desirable. By means of such an intermediate step many metals may be readily changed from one salt to another a change that may not be possible in one step otherwise. Give directions for changing the magnesium in magnesium sulphate to magnesium chloride; also for changing copper in copper nitrate to copper chloride. (c) Given a solution containing aluminium and ferric salts. State how these metals may be obtained separately in the form of any compound. Similarly, how may zinc and cadmium be separated ? 2. Ammonia as a Reagent. Preliminary Considerations. Since a solution of ammonia contains OH ions we expect to find that it reacts just as sodium hydroxide does in the preceding exercise; that is, as a reagent to precipitate metal hydroxides. However, it differs from so- dium hydroxide in being a weak base, and hence it has no effect 94 University of Texas Bulletin on solutions of salts of the strong bases (which are these?). With reference to the other bases, the following holds. Am- monia precipitates the hydroxides of the trivalent metals com- pletely under all conditions, but it does not precipitate these of the other weak-basic cations if enough of an ammonium salt (i. e. } many NH* ions) is present (exceptions lead and mer- cury, which with ammonia always form insoluble, complex am- monium compounds). This influence of the NH* ion is ex- erted as follows: since ammonia is very slightly ionized, the product (NH 4 *)X(OH'), which holds an equilibrium with the undissociated NH 4 OH, must remain practically constant; hence when (NH 4 *) is increased by the addition of an ammonium salt (e. g. HN 4 C1) which is largely dissociated, (OH') must de- crease, or in other words, ammonia is less dissociated in the presence of an ammonium salt than in the, absence of the latter. The concentration of the OH ion in such a mixture is so small that the product obtained by multiplying this OH ' concentration with the concentration of any bivalent metal ion that may be added ( [M*] X [OH] ') will be less than the "ion-product" that a resulting precipitate itself could produce. The formation of a precipitate does not take place because thereby the ion-product would be increased, and reactions take place only when an ion- product is decreased. But the ion-products of the tri-valent- metal-hydroxides are smaller than those of the bivalent-metal- hydroxides, and whenever a tri-valent-metal-ion is introduced into a solution containing any OH' ions, the product [M] X [OH] is always greater than the ion-product of the precipitate, hence the precipitate is formed because thereby the product of these ions in this solution is reduced. To demonstrate this effect, add some ammonia to one portion of magnesium chloride solution, and to another portion add some ammonium chloride solution and then ammonia. Exercise. By the use of ammonia, separate bismuth from a solution containing bismuth and copper salts; Al from a solu- tion containing Al and Ni salts; Fe (ic) from a solution con- taining ferric and Mg Salts. Give directions for changing the aluminium in potash alum completely to aluminium chloride. Chemistry in High Schools 95 3. Soluble Sulphides as Reagents. Note. Hydrogen sulphide is poisonous, hence the generators of this gas should be used either outside of the building or in well drawing hoods. Preliminary Considerations. Since all sulphides except those of the alkali and alkaline earth metals (Na, K, "NH 4 ," Ca, Sr, Ba, Mg) are insoluble, it is to be expected that a precipi- tate of the corresponding sulphide will be obtained whenever a soluble sulphide is added to a solution of a salt of any of the other metals. And such is the case except when hydrogen sulphide is used. Judging from its properties, this substance appears to be an exceedingly weak acid very slightly disso- ciated and its ion product (H*)X(S") is very small. A great increase of H* ions, through the addition of a strong acid, lessens its S" concentration greatly for the same reason that an increase of NH 4 ion concentration decreases the OH' con- centration in ammonia. Hence the following general fact con- cerning the effect of hydrogen sulphide as a regeant:- Even in the presence of a moderate concentration of hydro- gen ions (a strong acid), hydrogen sulphide precipitates com- pletely the sulphides of some metals, among them Ag, Hg, Pb, Bi. and Cd, which for convenience may be called the hydrogen sulphide group; and only in the absence of hydrogen ions (ab- sence of acid) does it precipitate completely the sulphides of other metals, among which may be mentioned Fe, Zn, Ni, Mn, the ammonium sulphide group. Demonstration. Put 10 to 20 cc. of dilute copper sulphate solution and an equal amount of zinc sulphate solu- tion in a small flask, add 10 to 20 drops of dilute HC1, heat to boiling, treat with H 2 S in excess and filter. What is the sub- stance on the filter paper? The reaction that has taken place is metathetical write the equation. What is the by-product? Evidently the concentration of the hydrogen ion has been in- creased during the progress of the reaction. Since with exces- sive concentration of hydrogen ion, hydrogen sulphide is un- able to precipitate sulphides of this group even (?), and since it is possible that in this experiment the concentration of the acid has become excessive, the last portion of copper may not 96 University of Texas Bulletin be precipitable under the conditions. Hence a small part of the solution must be diluted largely by adding an equal volume or more of water, or the acid may be partially neutralized by the cautious addition of a base. Then it should be heated again and treated with hydrogen sulphide. If necessary, treat the main portion of the solution in this way and repeat this pro- cedure until precipitation is complete. Before proceeding to precipitate the zinc sulphide, it is de- sirable to remove the remnant of hydrogen sulphide which re- mains dissolved in the liquid, since otherwise it starts precip- itation :as soon as the acid is neutralized by ammonia or any other base. To remove this gas, boil the liquid well for two or three minutes. Then add ammonia to the clear (and odorless) solution until it is plainly alkaline. If enough ammonium salt is present (from what source?), then no precipitate will be formed. Next, treat the solution with hydrogen sulphide again. A white pre- cipitate (ZnS) will be formed (a trace of iron compounds in- troduced by the hydrogen sulphide may impart a "dirty" color to the precipitate). Filter and wash the precipitate. The ammonia added above performs two functions: first, it neutralizes the free acid present, and second, it forms ammo- nium sulphide with the hydrogen sulphide. Write the meta- thetical reaction of the action of ammonium sulphide with the zinc salt. Transfer the CuS to a dish, add a little dilute HN0 3 , and warm the mixture. A solution of Cu (N0 3 ) 2 will be obtained by complex reaction. Note the free sulphur formed. Drip dilute HC1 slowly upon the ZnS on its filter paper until all has been dissolved and collected in a beaker placed below. The reaction is metathetical ( ? ) . Note the evolution of H 2 S. What is the object of this whole experiment? Reactivities of Sulphides. The reagents under (b) and (c) below will also react upon the sulphides in the groups above their own; but the reagents of (a) and (b) do not react upon the sulphides in the groups Chemistry in High Schools 97 below them. In all cases the metals dissolve because they are changed to water soluble salts. (a) Sulphides which react metathetically with dilute HC1 or dilute H 2 S0 4 : ZnS, FeS, MnS, CdS. CdS requires a stronger "dilute" HC1 than the others. (b) Sulphides which react with dilute HN0 3 action com- plex, resulting in free sulphur and nitrates of metals: Ag 2 S, PbS, CuS, Bi 2 S 3 . (c) Sulphides which react with aqua regia only (aqua regia is essentially a concentrated solution of free chlorine which acts to produce free sulphur and chlorides of the metals) : NiS, CoS, HgS. Colors of Sulphides. The colors of sulphides are a valuable means of identifying the metals. By proper means (?) prepare small portions of the colored sulphides. They have the following colors: FeS, NiS, Ag 2 S, HgS, PbS, Bi 2 S 3 dull black. CuS brownish black. ZnS White (when pure!). CdS yellow. MnS pink or flesh color. Note: during precipitation HgS frequently exhibits several other colors black, yellow, red but with excess of H 2 S it finally becomes black. Exercise. For how many metals may soluble sulphides be used as precipitation reagents? Into how many classes may metals be separated by means of soluble sulphides? The an- swers to these two questions will indicate why soluble sulphides are the most valuble general reagents known. Precipitate separately as sulphides the metals from a solution containing any one of the following pairs of metals copper and zinc, cadmium and nickel, bismuth and manganese. Continue the treatment for the precipitation of the first metal until the second sulphide can be obtained pure as shown by its color. Wash the copper, cadmium, or bismuth sulphide with distilled water while it is on the filter paper (to remove remnants of the salt of 98 University of Texas Bulletin the second metal), and convert it to the nitrate. Convert the zinc, nickel, or manganese to the chloride. List of Useful Special Properties and Reactions of Metals. These properties and reactions are only briefly indicated here. Consult a text book for further information. These properties and reactions need not be definitely remembered, since they may be looked up when needed. The student should make simple trials to acquaint himself experimentally with these facts. 1. AgCl is soluble in ammonia, from which solution AgCl may be reobtained by neutralizing the ammonia with an acid (HNO,-). 2. PbCl 2 is very soluble in hot water, though not in cold water. 3. HgCl forms a black compound with ammonia, for which the formula and reaction need not be learned. 4. HgCl 2 is reduced to HgCl or Hg by SnC'l 2 solution, the latter becoming SnCl 4 . 5. Bi salt solutions hydrolyze when the concentration of free acid (H* ion!) in the solution falls below :a certain limit. "Hy- drolysis "is a metathetical reaction between water and a salt which produces the free acid and the free base (or a basic salt), e. g. BiCl 3 +H 2 0=BiOCl+2HCl. A1,S 3 +6H 2 0=2A1(OH) 3 +3H 2 S. All salts of weak acids or of weak bases are hydrolyzod, though some only so slightly that the effect can not be noticed except by some delicate means such as the effect upon litmus. For demonstration, test solutions of the following normal salts with litmus : Sodium carbonate. Sodium acetate. Copper sulphate. Zinc chloride. Chemistry in High Schools 99 6. Note that the following salts and their solutions are col- ored : Copper salts blue or green. Nickel salts blue or green. Ferrous salts pale green. Ferric salts reddish yellow. Manganous salts pale amethyst. 7. Compounds of the following metals will color the Bunsen flame. To try this use a clean platinum wire, dip it into solu- tions of these metals and hold the drop of the solution in the flame : Bright red Sr. Brick red Ca. Yellow Na. Yellowish green Ba. Green to blue Cu. Blue, pale Pb. ' Violet K. 8. Ammonia in compounds is revealed by treating them with a strong base (e. g. NaOH solution), warming the mixture, and noting the odor. 9. Sodium hydroxide cannot be used as a precipitating rea- gent in the presence of ammonium salts because it reacts with the latter. Ammonium salts may be removed by evaporating the solution to dryness and then heating the dish and contents to low redness until fumes cease to be given oft*. Chemical Problems. The following problems are to be solved by means of the fore- going facts: 1. Students should be given solutions of water soluble salts, each solution should contain only one metal, and the student should ascertain what the metal is. Salts of the following inetals may be given copper, silver, bismuth, lead, mercury (both valences), cadmium, iron (ferric only) manganese, nickel, zinc, calcium, strontium, barium, magnesium, sodium, potassium, am- monium. Twelve to fifteen solutions should thus be worked out by every student. 100 University of Texas Bulletin In trying to find the metal, the student should note the color of the solution; he should ascertain if it would color the flame of the Bunsen burner; and on very small portions of the solu- tion he should try the effect of reagents in the order given be- low. (a) Add dilute hydrochloric acid. If a precipitate is ob- tained, then special tests 1 to 3 should be tried on the ppt. after it has been filtered off and washed. (b) Add dilute sulphuric acid. (c) Acidify moderately with either HOI or H 2 S0 4 (which- ever produces no precipitate), and then treat wtih hydrogen sulphide (see "g" below). (d) Add ammonium chloride and ammonia. Ammonium chloride will produce a ppt. if HC1 did, but such a ppt. should be removed by filtration before ammonia is added. (e) Irrespective of the presence or absence of any ppt. pro- duced by ammonia, treat the resulting mixture from (d) with hydrogen sulphide (see "g" below). (f ) Add sodium hydroxide. (g) If a black sulphide has been obtained, then to decide which metal sulphide it is, try special tests 4 and 5 above, and also ascertain the reactivity of the sulphide with acids after it has been filtered off and washed. 2. Prepare KNO 3 by the commercial method, from NaN0 3 and KC1. (See Smith, Art. 127.) 3. Prepare NaOH by the commercial method, from Na 2 C0 3 and Ca(OH) 2 . (See Smith, Art. 126.) 4. Prepare pure sodium chloride by precipitation with HC1. (See Smith, Art. 132.) Text-Book Reading. Parallel with this laboratory work on the metals some of the usual text-book reading found under the various metals should be given. However, it should be more or less confined to the commerically imporant facts and compounds, such as the com- mercial sources, and the methods of preparation, from these sources, of the commercially important products. The choice of what is commercially important varies some- Chemistry in High Schools 101 what with the locality. In an agricultural state such as Texas the source of potassium salts and the composition of fertilizers is of vastly greater importance than the source of copper, and the methods of its extraction from its ores; while in a mining state such as Arizona just the reverse holds good. The empha- sis in the informational reading should be laid accordingly. Chemical Changes Involving Oxidation and Reduction. lonisation and Valence. In contradistinction to the changes studied thus far (metathetical changes and hydration) in which the valence of all ions (or constituent parts of compounds) re- mained constant, the changes now to be studied specially . are those involving changes of valences. As has been shown, valence means the number of electric charges on an ion (or that would be on any particular part of an acid, base, or salt after it had been changed to an ion). Thus the valence of the N0 3 ion is 1 because in the ionization of any one of its com- pounds e. g. NaN0 3 or HNO 3 , the Na or H gives one negative ionic charge (electron) to the N0 3 . So far the attention has not been directed to the complete ionisation of such complex compounds as HN0 3 , NH 4 C1, CuS0 4 , etc., which would yield each element of a compound in the form of a separate ion, and this must now be taken up. Thus the pri- mary ionisation of NH 4 C1 yields NH 4 * and Cl', but complete ionisation into the elemental parts requires the further ionisa- tion of NH 4 *. Since 4H* result from this ionization, and thus 3 new (-]-) charges appear, it follows that 3 ( ) have been given up by the H's and left upon the N ion (written N 3 ' or N'"). The whole compound in the form of ions may be ex- pressed thus (4H*, N 3 ', Cl'). In the same way the complete ionisation of H.,SO 4 requires, after the primary ionisation into 2H* and S0 4 ", the further ionisation of S0 4 ". Since 40" would result from this secondary ionisation, and hence 6 new ( ) charges appear, then the 6 ( + ) produced simultaneously Note. An asterisk designates positive electric charges; an apos- trophe, negative charges. 102 University of Texas Bulletin must reside on the S ion (written S 6 * or S******).' The whole may be written 2H*, S 6 *, 40". It is a general fact that in aqueous solutions hydrogen and metals become cations while (or OH), the halogens (F, Cl, Br, I) and the cyanogen radical (ON) all become anions. It has been shown above that in electrolytic and battery cells, during action, the substances at one pole undergo one kind of a valence change while the substances present at the other pole undergo the opposite kind of a valence change. The proportion of the substances changing valences simul- taneously at the two poles depends merely upon the fact that the number of electrons simultaneously transferred "at" the two. poles are equal (but they are transferred in opposite senses at the two poles). Thus if the two "pole changes" are then the proportion in' which the two substances change simul- taneously is Zn :2C1. When the results produced upon a substance by these changes in valence are compared with the results produced upon it by oxidizing (or reducing) agents, it is found that an increase of positive ionic charges (or decrease of negative) produces the same result as oxidation, and the reverse electric change pro- duces the same result as reduction. This has been found to be always true. For illustration, consider the production of chlo- rine from HC1. Hence the following simple and perfectly gen- eral definition: oxidation is the increase of positive charges or valences of an element or radical (or the decrease of negative charges). Reduction is the reverse change. Exercise. Ascertain whether the elements italicized when changed from the first to the second form undergo oxidation or reduction. Express the amount of change in numbers of ionic charges per atom of the changed element. Chemistry in High Schools 103 First Form. Second Form. (a) KCl HC7 (b) Chlorine HCl (e) ZnS0 4 ZnCl 2 (d). Zinc ZnC\ 2 (e) K 2 S0 4 KCl (f) UNO, #H 4 OH (g) KMnO, MnS0 4 (h) \FeS0 4 FeCl 3 (i) K 2 CrO 4 CrCl, In the light of this definition of oxidation and reduction, it is seen that in every electrolytic cell and every battery cell, oxida- tion takes place at one pole, and reduction at the other. Study next the "Table of Electromotive Forces of Battery Poles," page 109. If the "valence-changing" substances of the two poles of a battery cell are put in direct contact, (i e. mixed together) they will undergo the same changes without the use of the metallic poles and their connecting wires. Here evidently the transfer of the electrons is direct from one valence-changing substance to the other. This is easily shown by the following test-tube trials with some of the same substances used in the battery cells shown before : (a) Put granulated zinc into some copper sulphate solution, and shake the mixture at intervals until the solution is colorless. Then test solution .with H 2 S for zinc ion. Con- clusion ? (b) In place of chlorine in a chloride solution it is better to use iodine dissolved in KI solution (or bromine water). Treat this with zinc until the free iodine (or bromine) has disappeared. Then test the solution for zinc ion. Conclusion ? (c) Treat a ferric chloride solution with zinc until the red- dish brown color of the ferric ion has disappeared. Test with NH 4 C1 and NH 4 OH to ascertain if the change is complete. 104 University of Texas Bulletin k The three equations for these reactions are: (a) Zn +Cu 2 'S0 4 :=ZnS0 4 -f Cu to 2 to (b) Zn +21 =ZnI 2 . to 2* to 1' (c) Zn -f-2Fe 3 *Cl 3 =ZnCl 2 +2FeCl 2 . to 2* to 2* The numbers of electric charges on the valence changing sub- stances are written as shown in order to show at a glance which substances are oxidized and which are reduced. Point out which are oxidized, which reduced. The proportion between the valence-changing substances in a reaction is determined by making the number of electrons given up by one substance equal to the number of electrons taken up by the other substance. Thus in (b) above, one Zn gives up 2( ) while one I takes up only 1( ) ; hence their ratio in the reaction is IZn to 21. Another way of stating the fact is to say that the amount of oxidation in a reaction is equal to the amount of reduction. Exercise : Make test tube trials of the following substances : 1. Treat ferric chloride solution with excess of hydrogen sul- phide, and test with ammonium hydroxide. 2. Add a slight excess of bromine water to a ferrous salt solu- tion and test with NH 4 OH. . 3. Treat dilute bromine water with hydrogen sulphide. 4 Put a rod of zinc into a concentrated solution of stannous chloride. 5. Put a piece of copper into some mercuric chloride or nitrate solution. Write the equations for these reactions as in the illustration above. The foregoing reactions are evidently quite simple, yet they present all the principles involved, and other reactions which appear to be more complex really involve no new facts. As a Chemistry in High Schools 105 rule they differ only by involving some additional metathetical actions. This is well illustrated by the oxidizing actions of HNO, which will now be discussed. Nitric acid and its Reduction Products. The following should be among the experiments shown : 1. The preparation of HN0 3 by distillation from NaN0 3 and concentrated H 2 S0 4 . 2. The ease with which HN0 3 gives up oxygen is strikingly illustrated by dropping a glowing piece of charcoal into some hot fuming nitric acid in a test tube. 3. The preparation of NO by the reduction of dilute HN0 3 with copper. To be performed as usual. Show that NO combines with oxygen to form N0 2 and that this dis- solves in cold water forming HN0 3 +HN0 2 . 4. The extreme reduction of the nitrogen in the nitrate ion by means of a metal very high in the electromotive force table (see page 110), e. g. zinc or aluminium: take concentrated sodium hydroxide solution, add a little of a nitrate, then add either one of these metals in the form of shavings or powder. Warm the mixture. After violent reaction has set in, the odor of ammonia can be easily noticed. 5. Treat some cold, freshly prepared, ferrous sulphate solu- tion with dilute HN0 3 . A black compound (of FeS0 4 , NO) will be formed. Heat the mixture to boiling. The NO escapes, and the reddish brown color of the solution shows the presence of ferric iron. 6. Pass H 2 S through warm dilute HN0 3 . Sulphur and NO will appear. In order to consider properly the oxidizing actions of HN0 3 the following facts must be noted: The nitrogen in nitric acid and in nitrates is in the highest state of oxidation in which it ever occurs; and the nitrogen in ammonia or ammonium compounds is in the lowest state of oxidation, it cannot be reduced further. Between these two limits there are a number of oxidation stages, the order and re- lation of which is shown in the following table. Note that the free element occupies an intermediate position. 106 University of Texas Bulletin Nitrogen Name. Ion. Nitric acid, HN0 3 and nitrates N 5 * Nitric peroxide or nitrogen textroxide, N0 2 or N 2 4 .... N 4 * Nitrous acid, HN0 2 and nitrites N 3 * Nitric acid or nitrogen dioxide, NO N 2 * Hyponitrous acid, NHO, and nitrous oxide N 2 N* Nitrogen, N 2 N Ammonia, and its salts, (NHJ * N 3 ' The equations for the experiments above may now be de- rived. Experiment (3). When completely ionized, HN0 3 is H*, N 5 *, 30")- Since the nitrogen finally appears as NO, i. e. (N 2 *, O") it has received 3( ), three electrons. Since these electrons are obtained from Cu becoming Cu 2 *, it follows that these two substances must change in the ratio 3Cu to 2HN0 3 in order that the number of electrons given up by the Cu shall be completely taken up by the changing HN0 3 . But when 2HN0 3 change (to 2NO), then 2H* and 40" will become free ions, and since 0" ions cannot remain free in the presence of H* ions, they will form 4H 2 0. This requires 8H*, for which the 2HN0 3 just considered supply only 2H*, and 6 extra HN0 3 must give up their H* ions. The 6NO 3 ' from these 6HN0 3 balance up with the 3Cu** formed in this action. Considering all this we write 3Cu +2HN 5 *0 3 +6HNO 3 =3Cu(NO 3 ) 2 +2NO+4H 2 O to 2* to 2* Which substance is oxidized? which reduced? Is the amount of oxidation equal to the amount of reduction? Experiment (4). In the change from nitrate ion to ammonia, each "N" receives 8 electrons, while each "Al" gives up 3 elec- trons ; hence the ratio between the number of Al and the number of N0 3 ' which change simultaneously is 8A1:3NO 8 ' In the presence of sodium hydroxide, the Al*** ion combines with INa* and 20" to form NaA10 2 . The N 3 ' ion forms NH 3 , Note. All aqueous solutions contain H* and O" ions, and when these free ions are used up more will be formed by further ionisation. Chemistry in High Schools 107 hence it combines with 3H*. Considering all this, we arrive at the following equation: 8A1 -f 3NaN 5 *O 3 +5NaOH+2H 2 O==8NaAlO 2 + 3NH 3 . to 3 to 3' Which element or substance is oxidized? which reduced? Is the amount of oxidation equal to the amount of reduction? These questions should be asked in connection with all equa- tions of oxidation and reduction. Exp. (5). One Fe** gives up one electron to become Fe***, and since the nitrogen is reduced from HN0 3 to NO (i. e., 3 elec- trons are taken up by 1 N), then the ratio between the valence- changing substances is 3 FeS0 4 :HN0 3 . The other considerations are the same as in (3) above. The equation arrived at is: 3Fe**S0 4 +HN 5 *0 3 +3HN0 3 = to 3 to 2* Fe 2 (SOJ 3+Fe (N0 8 ) 8 +NO+2H 2 0. Exp. (6). S" ions change to S. Since NO is formed from NO 3 ', the ratio between the valence-changing substances is 3H 2 S :2HNO 3 . Hence the equation is: 3H 2 S 2 ' +2HN 5 *0 3 ==2NO+3S-f4H 2 O. to to 2* One more equation should be added here: the reaction of ''aqua regia. " Considering that NO is the reduction product and free chlorine is the oxidation product, the ratio between the valence changing substances must be 3HC1:1HN0 8 and the equation is 3HC1 '-fHN 5 *0 3 =2H 2 0+3Cl+NO. to to 2* The foregoing arithmetic considerations are second in impor- tance to the real chemical knowledge to be gained from the ex- periments above. The latter consists of knowing what particular products of oxidation or reduction are obtained under particular circumstances. This must be remembered. The "chemical" as- pect of the experiments with nitric acid involves the following general facts (see also the Table of Electromotive Forces of Battery Poles) : The extent of reduction of nitrogen in compounds depends 108 University of Texas Bulletin upon the strength of the reducing agent and upon the dilu- tion of the nitrogen compound. For example, with dilute nitric acid, copper produces NO, while zinc produces nitrogen. Again, with concentrated nitric acid, copper produces N0 2 ; and with very dilute nitric acid, zinc reduces the nitrogen from nitric acid to ammonia. These general relations should be impressed by drill with spe- cial examples as illustrated in the following exercise. Exercise. 1. Figure out the equation of the reaction that takes place when copper sulphide is dissolved in dilute nitric acid if at the end the sulphur is present as free sulphur and the nitrogen as nitric oxide. 2. Do the same with bismuth sulphide, and again with silver sulphide. 3. If aqua regia, after its preparation by mixing the acids, is essentially a solution of chlorine (see equation above), what will be the equation for its action upon mercuric sulphide which re- sults in a solution of mercuric chloride? 4. Note the position of silver, relative to copper, in the Elec- tromotive Force Table, page 110, and hence infer what will be the probable reduction product of nitric acid if the concentrated acid reacts with silver. Write the equation for this action. 5. Do the same for lead and dilute nitric acid ; and again for cadmium and very dilute nitric acid. Note on the Oxidizing Action of Sulphuric Acid. Sulphuric acid is occasionally used as an oxidizing agent. The student may have occasion to meet it in connection with the dissolution of metals in it, or in its action upon potassium bromide and potas- sium iodide. Since the reduction products of sulphuric acid are all known to the student, it is advisable to add here the few re- marks necessary to inform him fully concerning these particular reactions. The order of the reduction products of sulphuric acid and of their ions is as follows: Chemistry in High Schools 109 Order Substance Ion. 1 Highest state of oxidation H 2 S0 4 S * 2. Next lower H 2 S0 3 or S0 2 S 4 3. Next lower S S 4. Lowest H 2 S S" With weak reducing agents such as metallic silver and metallic copper concentrated sulphuric acid will be reduced to S0 2 only. Of course the sulphate of the metal is produced in the reaction. With the strong reducing agents such as zinc, the reduction product is H 2 S. Concentrated sulphuric acid does not oxidize chlorine from HC1; but it oxidizes bromine from bromides forming free bro- mine, and the sulphuric acid itself is reduced the least amount only, namely to S0 2 . Iodine from iodides is so easily oxidized that the sulphuric acid is reduced to H 2 S even. For a numerical exercise, the first five equations for the five reactions just mentioned may be figured out. Table of Electromotive Forces of Battery Poles. In the table below, of electromotive forces of some battery poles, the voltages given are those of the cells formed by combin- ing each of these poles with a hydrogen pole, which latter serves as a zero or reference pole. The prefixed algebric signs indicate the polarity of the " meas- ured" pole in this combination. The accompanying scale of these forces shows their mutual relation. When a pole-component changes so as to be deprived of or give up electrons, then it is oxidized, or acts as the negative pole of a battery cell. Any pole in the table combined with any other pole will form a battery cell the force of which is shown in the ' ' Scale showing the relation of EMF's": it is represented by the distance be- tween the poles. The pole uppermost in the table acts as the negative pole in this combination, here oxidation takes place. When the substance that would be used up at one pole during the action of a cell is mixed or placed in direct contact with the substance that would be used up at the other pole, then 110 University of Texas Bulletin p 5-9 c a -2 1= .? . . 1 . ^ .D.S 2 g o 3 -- PM ftZ5 o o c o o o o o o o 2 2 .2 .2 .2 .2 .2 .2 .2 5 .2 o B wquCJjoJVVwww 5}3!)J)3$2>Q)q;q2qj?Qjq;'t^ "t* $ Q? $ t^ $ $ Tj''''''''' oooo XliXlX5 --> o ~ X2 a " ^ t5 I |*6 a* III S 80 *** * O * 03 * m * * " - O ^^ A S*J* "**no'* OQ ** a* *** ^tr! 'oaix-soQ*'a' c * 8 os MJL rt 3 a>nb,aO.;a a* o S -*a> &0&J.S .S ^ S(^ M^OS^N^^O^^(^tW'-gO g^ p p o o o r- r- IH i-! i-I i-< i-J i-4 IH ,-i i i TTT TT IT T +++++++++ ++ ++++ + + Scale showlnr relation of EMFs. -K 3.0- Na Oa 2.0 -JAf Al -1.0- Zn S Cd PblO -Ni Pbl8. 8n 0.0-H -Aa. Sb. Bl Ou I Fel8 +1.0- Br HNOs Or23 Ol Mn27. O F +2.0- Chemistry in High Schools 111 these two substances react in the same way as in the battery cell. Direct contact makes the connecting wires, etc., for the transfer of the electrons unnecessary. All mixtures undergoing oxidation and reduction reactions may be considered to be made up in this way. The substances in the higher states of oxidation (in the wide column on the left) are all oxidizing agents : they appear in the ascending order of strength as oxidizers. The reducing agents are in the column on the right: they ap- pear in the descending order of strength as reducers. Any reducing agent in the table will react with any oxidizer below it in the table. The following two common examples of this fact should be noticed : (a) Any metal in the table will displace (reduce) any metal below it from a solution of its simple salts. Thus, zinc or iron will displace copper from copper sulphate, silver from silver nitrate,' or hydrogen from hydrochloric acid. (b) Anyone of the non-metals (Cl, Br, I, S), will be displaced (oxidized) from its simple-ion compounds (chlorides, bromides, iodides, sulphides), by any one of the non-metals below it in the table. Thus, chlorine liberates bromine from bromides, sulphur from sulphides, etc. The metals above zinc in- the table below, and fluorine, cannot be used as poles for practical batteries, or as reducing (respect- ively, oxidizing) agents in ordinary chemical (aqueous) mix- tures because these elements react with water. Note on the potential of chlorine: its potential is less noble ("noble" means the direction in the table from hydrogen to- wards gold) if the concentration of the free chlorine is less or if the concentration of the chloride compound is greater than as given in the table. Hence the formation of the first portions of free chlodine in a concentrated solution of hydrochloric acid will take place with a potential between +1.35 and +1.0 volts. This explains why nitric acid oxidizes hydrochloric acid. Note on the "Back" Electromotive Force of the Poles During Electrolysis. The products formed at the poles during electrolysis exert an "opposing" electromotive force, i. e. "opposed" to the ap- 112 University of Texas Bulletin plied force. They exert about the same force as the battery poles composed of the same materials. Of different chemical changes possible at the cathode of any particular cell that one will take place which will exert the least "zincic" potential (zincic means the direction in the table from hydrogen toward zinc). At the anode that particular change will take place which re- quires the least " noble " potential. For the liberation of hydrogen on metals other than platinum, a more zinzic voltage is required than for platinum. This addi- tional voltage amounts to as much as 0.5 0.7 volts for the metals lead, zinc, and mercury. Chemistry in High Schools 113 APPENDIX. The Details of Construction, Action and Operation of Alter- nating Current Rectifyers. The Action of the Electrolytic Cell. The electrolytic cell used to "rectify" the alternating cur- rent, as ordinarily constructed, has one pole of aluminium, the other of lead, and between them a solution of sodium phos- phate. When prepared for operation, the cell allows only that alternation of the current to pass which makes the aluminium pole the cathode, or negative pole, and it prevents the passage of the reverse alternation, which would make the aluminium the anode or positive pole. This property of the cell is due to the fact that the aluminium is covered with a non-conduct- ing film of aluminium hydroxide or basic salt with a gas (oxygen) held in its "meshes" or pores. Such a film over a metal does not prevent the passage of the current when this metal serves as the cathode; but it hinders the passage of the current when this metal serves as an anode. It is as though the cations need not actually touch the metal of the pole in order to be discharged, while the anions cannot be discharged unless they come in direct contact with the metal surface. Hence a cell in which one pole is covered with such an insulating film and the other not covered, or with a "conduct- ing" surface, is " uni-directional " in its action. The materials used for this purpose, aluminium and lead, are particularly well suited for this purpose on account of the fol- lowing advantageous properties possessed by them : (a) when a bare surface of aluminium is exposed to anodic discharge (in solutions of phosphates, sulphates, borates, etc.) a film (of aluminium hydroxide) is almost immediately formed; and this film is not disturbed by any cathodic dis- charge such as takes place in sodium salt solutions. Hence when a bare aluminium pole is connected with an alternating current, that alternation which makes it the anode will quickly cover it with such an insulating film, and the formation of this 114 University of Texas Bulletin film will not be hindered by the reverse alternation of the cur- rent. (b) a lead pole, which in the rectifier acts mostly as the anode, has its surface converted to lead peroxide by the anodic discharge, and this material, in contradistinction to the alumin- ium film is a good conductor. If the lead acts as a cathode, this , peroxide is reduced to metallic lead. Thus neither anodic nor cathodic action covers the lead with an insulating layer. With high voltages, or hot solutions, the aluminium film breaks down frequently in its weaker spots, and considerable leakage of current occurs. Free acid in the electrolyte also increases the leakage current. But with cold, neutral solutions and a voltage below 70 volts, the leakage current is not exces- sive. Construction of Large Electrolytic Jars for Rectifier Set No. 1. Secure 2 jars of 2 to 4 quarts capacity, with a depth of from 7 to 11 inches. Have a carpenter cut 4 discs of one-inch pine lumber, two discs of a size to fit inside of the upper edge of the jar, and the other two to fit over the top of the jar. By means of screws fasten one of the smaller and one of the larger discs together with the grains of the two pieces at right angles to each other; cut a rectangular hole through the center of these discs of such a size that the aluminium plate which is to be used will just pass through this hole. One side of the hole should be cut practically vertical, the other should be cut slantingly so that a suitable thin wedge inserted on the slant- ing side of the hole may serve to hold the aluminium plate rigidly in a vertical position when this wooden cover is placed on top of the jar. Secure, by mail, from the U. S. Aluminium Co. (Pittsburgh, Pa, or Old Colony Bldg., Chicago, 111.), speci- fying the softest grade for which their stock number is 2SO, four aluminium' plates 3 inches wide, 12 inches long, 1-8 inch thick, 13-5 lbs., - which will cost 80 cents. Drill a hole near one end of each of two plates to screw on a binding post. (It is almost impossible to fasten the connecting wires on the alu- minium by soldering.) Another way to fasten the connecting wire is to drill a hole about the size of the wire through the Chemistry in High Schools 115 upper edge of the plate, put in the end of the wire, and clamp it by hammering the plate together at that point. Secure from a plumber four strips of sheet lead about 3 to 4 inches wide,, 6 to 7 inches long and thick enough to be stiff. Flatten each out thoroughly, and at one end bend a strip % inch wide, sharply at right angles to the rest of the plate; on this strip solder the end of a 6 to 10 inch piece of copper wire (No. 14). By means of brass screws driven through this strip into the wooden cover, fasten two lead plates to the under side of each cover, so that they will be parallel to the aluminium plate at a distance of % to % inch, one on each side. Then dip the wooden cover into hot paraffin, so that all the wood and the lead plate fastenings will be covered with it. Insert the alu- minium plates, fill the jars three-fourths full with the phos- phate solution (prepared as directed below), and make the con- nections shown in Fig. 6. The two jars may be reduced to one by placing both alumin- ium poles in the same jar with one lead plate between them and one on the outside of each aluminium plate. Thus three lead plates are used, which are all connected to the same ter- minal. The connections are the same as with two jars. , i ,--...} Design and Action of Special Transformer. The transformer to be used in this- connection is made of only one coil. One connection for the current to be drawn from it is at the center of the coil (see Figure 6). The " positive" direct current flows outward from this, or in other words, this is the direct current + pole connection. In addition there are other pairs of connections to the coil, the two connections of each pair being " balanced" with reference to the center, i. e., between the center and one connection there are as many ' ' turns of wire" as between the center and other connection of a pair. One connection of each pair may be connected with the alumin- ium pole of one cell, and the other to the aluminium pole of the other cell. The lead poles of the two cells are connected to a common wire from which the negative direct current flows, in other words, it is the direct current pole connection. 116 University of Texas Bulletin The voltage between any two points of the primary coil of a transformer is practically such a fraction of the voltage of the primary circuit as the ratio that the number of "turns of wire" intercepted on the coil by the two points bears to the number of turns in the whole coil. Hence between the center and either extremity one-half of the primary voltage is ob- tained, and between the center and any other points, lesser volt- ages proportionate to the number of turns between it and the center are obtained. The direction of the electric pressure (or sign of the voltage) on one side of the centre is the opposite of that exerted simul- taneously at the other side. If at any particular moment the voltage on the left of the centre is such as to make the left con- nection negative, then this would tend to send a current into the (left) cell with which it is connected, thus making the aluminium pole of this cell the cathode, and the lead pole the anode. Under these conditions the aluminium pole opposes no hindrance to the passage of the current. However, exactly the reverse condition exists simultaneously at the connection to the right of the center of the coil and the aluminium pole in this second cell opposes (practically prevents) the passage of the current from this side. With the next alternation of the cur- rent all the conditions are reversed. This time the current is free to flow through the cell on the right, but it cannot flow through the cell on the left. Thus the current which flows out of the coil at the centre is always positive, but it comes from the right side of the coil at one moment, and from the left side at the next moment. In ordering such a coil suitable for "Rectifying Sets No. 1" above, the following should be specified: the transformer is to be of the single coil or "auto" type, and is to be designed for a "110 volts, 60 cycle" circuit; capacity, 5 amperes at 110 volts. It is to be without case, but equipped with frame cast- ings in order that it may be fastened to the table or wall. The following connections are to be made nine leads to be con- nected as follows: Two, at the ends of the coil for the main line. One, at the centre. Chemistry in High Schools 117 Three pairs for connections to be " balanced" with reference to the central tap and connected to the coil so as to secure the following voltages (approximately) : 13% volts between the center and either one of the first pair. 27% volts between the center and either one of the second pair. 41*4 volts between the center and either one of the third pair. In other words, the transformer is designed for eight circuits of 13% volts each. The Moloney Electric Co. of St. Louis, Mo., will furnish such a transformer for $13.00 or slightly less, the general appearance of which is shown on page 43 of their catalogue. Preparation of Solution for Electrolytic Jars. For each quart of solution use one-half pound of sodium phos- phate crystals, and add phosphoric acid in small portions until the mixture does not turn litmus paper a deep Hue, but a lighter shade not pink, however. Operation of the Rectifier Set No. 1. When the current is to be turned on for the first time after the rectifier has been assembled (or for the first time after a long period of rest) the lowest alternating current voltage should be connected to the electrolytic cells, and this should remain con- nected at least several minutes before any higher voltage may be connected. After that no further precaution need be taken. In the absence of an insulating layer over the surface of the alumin- ium the current will pass direct through both cells, but the liquid resistance of the cells will be large enough to prevent this current from being excessively large if the voltage is small. This current in passing forms the films over the surface of the alu- minium plates. During action aluminium phosphate is slowly formed and appears as a white flocculent sediment. Thus the aluminium pole and the phosphate ion (the P0 4 radical) are used up very slowly. Hydrogen also is discharged at the cathode. To re- 118 University of Texas Bulletin place this loss, small amounts of phosphoric acid must be added to the solution from time to time, with the aid of litmus as in preparing the solution. Naturally, any water lost by evapora- tion must be replaced. Under these conditions the cells will operate practically indefinitely. Depending upon the rate at which the current is drawn, the voltage of the direct current obtained will be from 10 to 25 volts less than the alternating current voltage applied to the electrolytic jars. When a current greater than 5 amperes is drawn continuously for more than an hour, the liquid in the electrolytic jars is heated so rapidly that it may finally become too hot for safe or economical operation. Hence when large currents are needed for longer periods of time, the electrolytic jars must be cooled in some way. This can be done by replacing the glass jars by heavy tin or sheet iron cans of the same size (which can be made by any tinner), and placing these electrolytic vessels in a large vessel filled with cold water. In extreme cases, ice water may be used. A Small Cheap Rectifier. Secure a 1 quart fruit- jar or a similar vessel. Secure a. strip of sheet lead about 3 inches wide and 10 to 12 inches long. Hang this lengthwise into the jar, along the wall, and clamp it in position by bending the upper end over the rim, and solder a connecting wire to this. Fasten a wooden cover (6x6 inches) on the top of the jar by means of pieces of wood nailed as cleats to the under side of the cover. Secure by mail, from the U. S. Aluminium Co. (Pittsburgh, Pa., or Old Colony Bldg., Chicago) rods of aluminium % to % inches in diameter, and 10 to 12 inches long. Specify the softest grade (Cat. No. 280). Four rods V2xl2 inches, will cost 50 cents. To fasten a copper connecting wire to the aluminium rods, drill, near one end of each, a hole large enough to admit the wire. When making connection insert the wire and clamp it by a few blows of a hammer on the side of the rod. Near the center of the cover, bore one (or two) holes of the exact size of the aluminium rods, so that the rods extending Chemistry in High Schools 119 through these holes will be held in vertical position. Con- nect this cell in series with either a lamp-bank or a suit- able rheostat. The lamp-bank should have 8 to 12 incan- descent lights of 32 candle-power, all connected in parallel. In place of this lamp-bank it will be more convenient to use a 11 theatre dimmer" for 12 lights (Ward-Leonard Electric Co., BronxviUe, N. Y. (Cat. No. SD 16), price about $2.50). After filling the jar with sodium phosphate solution, the rectifier is ready for use and can be used without any particular precau- tions. How to Rectify Both Alternations Without a Transformer. Without the aid of the special transformer, the only way to rectify both alterations is to use four simple cells and connect them as shown in Fig. 7. It is absolutely necessary to insert a large resistance in the alternating current circuit until the sur- faces of the aluminium plates are formed. For this purpose cells of the size first described above require two 32 c. p. lamps. As the surface "forms" the lamps become dim. Then the lamps may be taken out of the circuit because the film is able to stand 110 volts. A single-pole, double-throw switch is usually used to facilitate this change. The direct current obtained from this rectifier will have a voltage of 80 to 95 volts. If, as is usually the case in laboratories, the current is to be used at a low volt- age, a lamp-bank or some "high resistance" rheostat must be inserted in the circuit, either in the alternating or in the direct current side. This form of rectifier, together with its necessary accessories, costs fully as much as the first rectifier described, and is less convenient and less suitable. '"if L " (df\ filvm (QJ A Four- jar Rectifier (without transformer) 14 DAY USE RETURN TO DESK FROM WHICH BORROWED LOAN DEPT. This book is due on the last date stamped below, or on the date to which renewed. Renewed books are subject to immediate recall. 5 LD 21A-50m-8,'61 (Cl795slO)476B General Library University of California Berkeley Manufactured by GAYLORD BROS. Inc. Syracuse, N. Y. Stockton, Calif. SS261418 THE UNIVERSITY OF CALIFORNIA LIBRARY