Q D 457 B7 MAIN B M 57M LABORATORY OUTLINES IN PHYSICAL CHEMISTRY BY T. R. BRIGGS ITHACA, NEW YORK 1920 LABORATORY OUTLINES IN PHYSICAL CHEMISTRY BY T. R. BRIGGS ITHACA, NEW YORK 1920 LABORATORY OUTLINES IN PHYSICAL CHEMISTRY ' INTRODUCTION It is the purpose of this course to acquaint the student with some of the factors governing physical and chemical change and to enable him to recognize these factors and to measure their intensity by their effects. Painstaking accuracy is not required in most of the experi- ments, which have been designed primarily to illustrate principles and to encourage intelligent thinking. It is believed that work of this kind proves more interesting and stimulating to the average student than do the more tedious and exact measurements carried out commonly in laboratories of physical chemistry. Completion of the work of this course entitles the student to three hours of University credit per term, two of which are given for experiments performed satisfactorily in the laboratory and one for written reports based upon these experiments. The following Laboratory Outlines describe the work to be accomplished during the year, though certain of the experiments may be omitted at the discretion of the Professor in charge. No work of a similar nature done elsewhere at another college or university is required to be repeated, provided the work be submitted to the Professor in charge for his approval. LABORATORY In performing the majority of these experiments, students are to work in groups of two. Partners are to be chosen at the first labora- tory period and this partnership is to be maintained throughout the year so far as possible. It is absolutely essential, however, that both partners work in cooperation on the same experiment. Independent work on different experiments in a given group will not be permitted. Since this course is introductory in nature, the student is not given the most delicate instruments or the purest materials. The apparatus supplied will nevertheless be found quite sufficient for the require- ments of these experiments. When the student has determined how closely his calculations must be made, he can readily ascertain the allowable error, how carefully his measurements must be made and what degree of delicacy he must look for in his measuring instruments. All burettes and pipettes should be calibrated according to the methods of Experimental Group I and should be cleaned in chromic- 1 sulf)lno-40 nyi * EXPERIMENTAL GROUP VIII DISSOCIATION The following experiments are designed to illustrate qualitatively the dissociation of ehemical compounds, either as the result of an increase in temperature or as the result of dissolving the substance in a solvent. Dissociation of the first type is called thermal; dissocia- tion of the second type is called electrolytic when ions are formed. We have already studied some of the phenomena due to electrolytic dissociation, especially in Experimental Groups IV and V. Other instances of electrolytic dissociation and its effects will be studied in the Experimental Groups which follow. References. Solutions and Electrolytic Dissociation. Abegg: Die Theorie der elektrolytischen Dissociation, Ahren's Sammlung 8 (1903). Arrhenius: Theories of Solution (1912). Findlay: Osmotic Pressure (2nd Ed. 1919). Jacques: Complex Ions (1914). Jones : The Nature of Solution (1917) . Ostwald (Muir) : Solutions (1891). Rothmund: Die Loslichkeit (1907). Scxidder : Electrical Conductivity and lonization Constants (1914) . Seidell: Solubilities of Inorganic and Organic Compounds (1919). Stieglitz: Qualitative Analysis, Vol. I (1917). EXPERIMENT 1 Thermal Dissociation of Nitrogen Tetroxide In a test tube heat a small quantity of Pb(NO 3 ) 2 and pass the resulting gas through a delivery tube into a test tube which is surrounded by a freezing mixture of ice and salt. The NO 2 will condense, under these conditions, as a bluish green liquid, N 2 O4. On removing from the cooling bath the colorless gas N 2 O4 will be formed first and on further heating this will dissociate into NO 2 . Note color changes. References. Nernst, 453 (1911). Ostwald: Principles of Inor- ganic Chemistry, 329 (1908). EXPERIMENT 2 Thermal Dissociation of Limestone Heat some powdered marble in a hard glass tube. Show that dis- sociation takes place. References. Bigelow 4 etc. For study of the reaction used in lime burning read Kremann- Potts: 107; LBR, 398. 27 Define "dissociation pressure" and draw a curve showing how dissociation pressure changes with the temperature for the following reaction: 2NaHCO 3 = Na 2 CO 3 + H 2 O + CO 2 . Reference. LBR, 398. * EXPERIMENT 3 Electrolytic Dissociation and Color Part 1. Compare the colors of concentrated 'solutions of the fol- lowing salts: CuSO 4 , CuCl 2 , CuBr 2 . Dilute until they have the same blue color. Start with about one cc. of solution. Explain. Part 2. Add concentrated hydrochloric acid to a greenish-blue solution of CuCl 2 . Note color change. Also heat some of the same solution. Explain. References. Ostwald: Prin. Inorg. Chem., 642 (1908); also Mellor: Inorganic Chemistry, 468 (1914). Part 3. Color changes with CoCl 2 solutions. Dissolve a little cobalt chloride in absolute alcohol. Add two or three drops of water to the solution. Add ether to the solution. Add water again. Also to the pink solution in water add (1) solid magnesium chloride ; (2) concentrated hydrochloric acid. References. Donnan and Bassett: Jour. Chem. Soc., 102, 939 (1902). Ostwald: Prin. Inorg. Chem., 623 (1908); Nernst, 389 (1911). EXPERIMENT 4 Reactions depending upon Degree of Dissociation Part 1. Pass chlorine gas into AgNO 3 solution. Does a precipi- tate form at once? Part 2. Add carbon tetrachloride to AgNO 3 solution. Explain. Part 3. Add chloroform to AgNO 3 solution. Does a precipitate form at first? Let -the mixture stand in the light until the next period. Does a precipitate form on standing? Explain. EXPERIMENT 5 Complex Ions Part 1. To AgNO 3 solution add KCN in excess. Test for silver with NaCl. Do not use HCU Beware of HCN and remember that KCN is poisonous. Use hoods. Part 2. To AgNO 3 solution add sodium thiosulphate'in excess. Test for silver. Explain results. Cf. Walker, 343. 28 EXPERIMENT 6 Relative Stability of Complex Ions Part 1. Add KCN in excess to dilute CdSO 4 . Test for cadmium with H 2 S. Part 2. Add KCN in excess to dilute CuSO 4 . Test for copper with H 2 S. Explain. Cf. Walker, 343. Caution. Cyanogen is formed in Part 2. Use hoods. EXPERIMENT 7 Hydrolysis Part 1. Test KCN and Na 2 CO 3 solutions with litmus paper. Explain. Part 2. Test CuSO 4 and A1 2 (SO 4 ) 3 solutions with litmus paper. Explain. Part 3. Precipitate PbSO 4 completely from lead acetate solution by adding A1 2 (SO 4 ) 3 . Then add water and boil. Filter and test the filtrate for lead and aluminum. Precaution. It is essential to use very little Al2(SO 4 ) 3 in excess. At any rate, add plenty of water and boil thoroughly for several minutes. Explain. EXPERIMENT 8 Conductivity and Electrolytic Dissociation Discussion. The conductivity is the reciprocal of the resistance. From the resistance of a solution, its conductivity may be calculated. In this experiment the relative resistance of N/10 HC1 and N/10 CH 3 COOH is measured by reading the current and voltage across graphite elec- -p trodes which dip into the solution. From Ohm's law, I = ' the R resistance may be computed. By maintaining the temperature constant, keeping the electrodes the same distance apart, and having them immersed to the same extent, a rough approximation of the conductivity of these two equivalent acid solutions may be obtained. The conductivity of a solution depends, among other things, upon its dissociation. If two solutions are of equivalent concentration and at the same temperature and if both are placed in the same vessel for measuring the conductivity, the better conducting solution is either more completely ionized or else contains the more mobile (the more rapidly moving) ions. If the difference in conductivity is very great, as in the present case, the poorly conducting solution is almost certainly the less strongly dissociated. Since both solutions have the hydrogen ion in common and since the chlorine and acetate ions are about 29 equally mobile, the relative conductivity is here a very nearly exact measure of the relative ionization. Procedure. Measure the relative resistance of N/10 HC1 and N/10 CH,COOH solutions. Follow the procedure used in the experiment which showed the distinction between an electrolyte and a non-electrolyte Use alternating current and a-c meters. Look up the per cent ionization of N/10 HC1 and N/10 CH 3 COOH. Are your conductivity values proportional? 30 EXPERIMENTAL GROUP IX SOLUTION AND SOLUBILITY The experiments of the following group are designed to illustrate the process of solution, the properties of saturated solutions, the cor- rosion or solution of metals and the determination of solubility. References. See under Group VIII Dissociation. EXPERIMENT 1 Quantitative Determination of Solubility References. F, 302; OW, 176; G, 234, etc. The solubility of a salt in water depends chiefly upon the nature of the salt and the temperature. The rate at which the salt dissolves depends upon the same factors plus several others besides, such as size of particles, rate of stirring, presence of catalysts, and so forth. Solubility may be determined directly, provided the salt is not too slightly soluble, by saturating a solution with an excess of salt at a desired temperature, and analyzing a definite weight or volume of the solution. Determine the solubility of an assigned salt at 25 C. Place in a bottle an excess of finely powdered salt, add water and shake in a thermostat until equilibrium is reached, or until there is no change of density between successive tests, when measured with a delicate hydrometer. In a second bottle place finely divided salt and add, not water, but a solution of the salt saturated at some temperature (usually a higher one) at which the salt is more soluble than it is at 25 C. Shake as before and determine the density of the saturated solution. The final densities should be the same in both bottles. Withdraw samples for analysis using a dry pipette and a small filtering tube to prevent the entry of solids. Determine the concen- tration of the saturated solution either by chemical analysis, or by evaporating a weighed sample to dry ness in an oven or desiccator. Check results. Determine the density of the solution at 25 C. and calculate the solubility of the salt in grams per 100 grams of solution; also in terms of the "molar fraction" of the solute. EXPERIMENT 2 Cryolite and Water Add a little finely powdered cryolite to water in a test tube. Does it dissolve? Explain. 31 EXPERIMENT 3 Solution and Catalysis Chromic chloride appears in two forms, as the hexahydrate (CrCl 3 . 6H 2 O) green in color, and as the anhydrous salt (CrCl 3 ) which is violet. The anhydrous form appears to be nearly insoluble in water while the hydrate dissolves readily. According to Moissan the violet form dissolves slowly at high temperatures to a green solution, and Ostwald believes that the apparent insolubility at ordinary temperatures is due to the extreme slowness with which solution occurs; in other words, that the violet form is not really in equilibrium with water. Drucker under Ostwald's direction showed that the violet modification dissolves readily in the presence of chromous chloride (CrCl 2 ) in solution, the latter acting as a catalyst. With these facts in mind perform the following tests: (1) Try to dissolve violet CrCl 3 in water. (2) Dissolve some green hexahydrate in water. (3) To a small quantity of the violet salt add water plus a crystal of the green hexahydrate. Add a bit of zinc and acidify with HC1. See whether the violet salt dissolves in time. Explain. (4) To th'e violet salt add a bit of metallic chromium, then add dilute HC1. Does the salt dissolve? (5) Prepare chromous chloride by dissolving metallic chromium in dilute HC1. Add this solution to a few particles of the violet salt. Do they dissolve? (6) Repeat (d) adding zinc instead of chromium. Explain. References. Ostwald: Prin. Inorg. Chem. 615; Mellor: Inorg. Chem. 258. Drucker: Zeit. phys. Chem., 36, 173 (1901). EXPERIMENT 4 Relative Solubility Part 1. Precipitate PbSC>4, let it settle, wash once or twice by decantation, then add KI solution to the residue. Note the color change. Then warm it. What color change occurs? Part 2. Precipitate AgCl, repeat procedure in (a) using KBr solution. What change occurs in the precipitate? Explain. Part 3. Repeat (2) using KI solution. Walker, 356 (1913). Part 4. Precipitate AgBr, add KC1 solution. Is there any visible change? Explain. Part 5. Prove by simple experiments, which is the more soluble, CaSO 4 or CaCO 3 . EXPERIMENT 5 Compound Solvents Part 1. Add about 20 cc. of impure commercial sulphuric acid to an equal volume of water. What is the precipitate? Explain. 32 Part 2. Add about 5 cc. of 95 per cent ethyl alcohol to (1) a saturated solution of Na 2 SO4 (2) a saturated solution of Na 2 CO 3 . cf. Group V, Experiment 3. Part 3. Determine by experiment qualitatively the effect of sodium chloride on the solubility of phenol in water. Repeat with sodium acetate instead of sodium chloride. EXPERIMENT 6 Solubility Product Discussion. When a salt, dissociating into univalent cations and anions, is in equilibrium with its saturated solution, the Law of Mass Action leads to the conclusion that the product of the concen- trations of cation and anion is a constant for a given temperature, provided the nature of the solvent undergoes no change. The product of the ion concentrations when the solution is saturated is called the solubility product. Thus : [cation] [anion] = Ks, (1) where the symbols "[cation]" etc., represent the concentrations. Reference. Stieglitz, I 141 (read page 142 for criticism of theory) ; Washburn, 298. When the salt dissociates into polyvalent ions or into ions of mixed valence, the relation is more complex. Cf. Washburn, 301. It is possible to distinguish between two cases, as follows : (1) When to a solution saturated with a given solid electrolyte there is added a soluble salt containing a common ion, the product of the concentrations of cation and anion momentarily becomes greater than the solubility product. The solution is no longer in equilibrium with the saturating solid salt and the latter is precipitated, until new conditions of equilibrium are established. These new conditions correspond to diminished solubility. (2) When the concentration of one or both of the ions produced by the saturating solid is decreased by any kind of physical or chemical reaction, the product of the concentrations of cation and anion momentarily becomes less than the solubility product. The solution is no longer in equilibrium with the solid and fresh solid dissolves until new conditions of equilibrium are established, the latter corresponding to increased solubility. Procedure. Part 1. To a BaCl 2 solution in a test tube add concentrated HC1, then add water. Explain. Part 2. Repeat, using a CaCl 2 solution. Part 3. To a saturated solution of NaCl, add concentrated HC1. Part 4. To a saturated solution of HgCl 2 add a saturated solution of NaCl. Account for what happened in (1) and (3), by applying the theory of the solubility product. Explain the very different results of (2) and (4). 33 How might all these experiments be explained in the light of Experiment 4? Reference. Ostwald: Prin. Inorg. Chem., 675 (1908). Part 5. Treat some freshly precipitated and washed AgCl with (a) Na 2 S 2 O 3 solution; (b) with KCN solution (poison); (c) with NH 4 OH. Part 6. Treat some freshly precipitated calcium oxalate with HC1. Part 7. Shake a little HgO with a solution of KI. Note any color change. Filter and test the nitrate with red litmus. Part 8. Prepare some Cd(OH) 2 and wash thoroughly with water. Shake with water and test the supernatant liquid with red litmus. The solution should be neutral. To one-half of the Cd(OH) 2 add a small amount of KNO 3 and shake again. Test the supernatant liquid with red litmus. To the second half of the Cd (OH) 2 add a little KI, shake and test the supernatant liquid* with red litmus. Explain. Reference. Ostwald: Prin. Inorg. Chem., 637 (1908). EXPERIMENT 7 Solubility of Glass Part 1. Phenolphthalein Test. Boil some clean, finely-powdered glass with water in a beaker, then add a drop of phenolphthalein. Explain. Part 2. Eosin Test. "If a glass surface is brought into contact with watery ether, it draws water from the solution and gives up alkali to it. On the other hand, the orange-yellow solution of iod- eosin in ether is changed by the alkali into red. Mylius, who had previously used the color reaction for another purpose, has applied it to the practical testing of glasses. Commercial ether is shaken up with water at ordinary temperature until it is saturated with water. It is then poured from the rest of the water and eosin is added in the proportion of 0.1 g. to 100 cc. of the liquid. The solution is filtered "Glass vessels are tested by pouring in the solution. The first step is to clean the surface from any products of weathering which may adhere to it, by carefully rinsing with water, with alcohol, and lastly with ether. Immediately after the cleaning with ether, the eosin solution is poured in, the vessel is carefully closed and the solu- tion is allowed some twenty-four hours to do its work. It is then emptied out and the glass rinsed with pure ether. The surface of the glass is now seen to be colored red; and the strength of the color furnishes an indication of the susceptibility of the glass to attack by cold water." Reference. Hovestadt (Everett) : Jena Glass and its Scientific and Industrial Applications. 34 Following these directions make up 100 cc. of eosin solution. Then test the surface of a new 50 cc. beaker and a new test tube as described above. Place a small sample of powdered glass in an 8-dram vial and add some eosin solution. Note the color the powdered glass assumes on standing twenty-four hours. Note also the color of the walls of the vial. If the powdered glass becomes colored, filter it and wash thoroughly with water. Does the water remove the color? Pour off the water and add alcohol. Does the alcohol remove the color? Eosin as Indicator. Take a few cubic centimeters of the eosin solution and add a few drops of dilute NaOH. Part 3. Tetrachlorgallein Test. Add to a beaker of boiling dis- tilled water a few drops of alcoholic tetrachlorgallein. Continue the boiling and observe the color change. Make a blank test with fresh distilled water. EXPERIMENT 8 Corrosion of Metals Discussion. Many metals dissolve more or less readily in aqueous solutions, appearing in the solution in the form of cations for at least a limited time and displacing during this process an equivalent weight of some other cation, usually hydrogen, from the solution. Thus zinc and sulphuric acid give zinc sulphate and hydrogen; zinc and copper sulphate give zinc sulphate and metallic copper, the salts and acids being in solution. Under these circumstances the zinc is said to corrode. It is generally believed that the process of corrosion is electro- chemical in nature. For-example, when zinc corrodes, two so-called "electrochemical" reactions take place as follows: (1) Metallic zinc gives zinc ions plus negative charges, the latter being retained by the metal. (2) Hydrogen ions in solution give hydrogen gas plus positive charges, the latter neutralizing the negative charges on the metal. Represented by symbols, these reactions may be written: +26 (1) +20 (2) If one adds reactions (1) and (2), the total reaction becomes Zn + 2H+ ^Zn + + + H 2 (3) It is interesting to note that the anions appear to play no part whatsoever. Applying the Law of Mass Action to the two reactions given above, it is possible to draw the following conclusions regarding the rate of corrosion : (1) A metal tends to corrode more readily in an aqueous solution the greater its "electrolytic solution pressure," i. e., the greater the driving force of reaction (1) or the greater the ion-forming tendency of the metal. 35 (2) The smaller the concentration of the dissolving metal as ion in the solution, the faster is the corrosion. The' ion concentration may be kept low by the formation of complex ions, by hydrolysis, etc. (3) The greater the hydrogen ion concentration in the solution the faster the corrosion. Other things being equal, metals tend to cor- rode more readily in acids than they do in alkaline solutions. (4) Anything that reacts with and removes the discharged hydro- gen tends to aid corrosion. Oxidizing agents may do this, in which case they are called "hydrogen depolarizers." Note the part played by air in the experiments; also the formation of nitrites in Part 2b. (5) The absence of stable, difficulty soluble protecting films (pas- sivity) favors corrosion. (6) Miscellaneous. Metal should have irregularities, etc., in surface to aid in setting up local "galvanic" couples. Also the "overvoltage" for hydrogen should be low. These points belong properly under electrochemistry and cannot be discussed here. All the conditions favoring corrosion do not have to be fulfilled simultaneously. Copper for example corrodes in aqueous ammonium hydroxide in the presence of air. The electrolytic solution pressure of copper is very small and the hydrogen ion- concentration in ammo- nium hydroxide solution is very slight, but these conditions which tend to prevent corrosion are more than offset by the fact that the copper ion concentration in the solution is practically zero (complex Cu(NH 3 ) 2 cations) and air oxidizes the discharged hydrogen under the conditions of the experiment. The reaction as a whole may be written : Cu + 2NH 4 OH + O >Cu (NH 3 ) 2 (OH) 2 + H 2 O Iron corrodes readily in moist air. Moisture is essential inasmuch as it furnishes the hydrogen ions which are displaced by the iron, the latter entering the solution in the form of ferrous ions. These are almost immediately oxidized by air to ferric ions which combine with the hydroxyl ions of the water to form hydrous ferric oxide. The iron thus passes from solution and corrosion is thereby accelerated. Carbon dioxide stimulates corrosion by dissolving in the film of mois- ture and thus increasing the hydrogen ion concentration by the forma- tion of H 2 CO 3 . Air increases corrosion by removing the dissolved iron as explained above and by serving as the hydrogen depolarizer. Procedure. Part 1. Solubility of Metals in Acids and Alkalies, (a) Place a small strip of copper foil in aqueous NH^OH in a test tube. Shake thoroughly from time to time. Note the color change and explain. (b) Experiments with concentrated H 2 SO4. In a few cc. of concentrated H 2 SO4 test the solubility of cast iron, iron wire, nickel wire, and copper wire. Set aside for an hour. Dilute the acid five fold with water and repeat, using the same test pieces. Dilute the acid until the rate of solution is rapid. Caution. Dilute the acid properly. 36 (c) Experiments with concentrated HNO 3 . In a few cubic centimeters of concentrated HNO 3 , test the solu- bility of iron wire and nickel wire. Set aside for an hour. Repeat with acid diluted twice. Why are metals often more readily attacked by HNO 3 than they are by HC1? Test the solubility of aluminum in caustic soda solutions. Explain. Aluminum forms complex anions in NaOH. Part 2. Solubility of Metals in Salt Solutions. Clean the metal thoroughly, and, after weighing, set aside for ten days in a test tube with 10 cc. of the salt solution. Cover up loosely with filter paper. Shake from time to time. Clean the test piece and weigh again. Record the time and note any change in the metal. (a) Copper in 10 per cent NaCl, test alkalinity of filtered solution at end. (b) Cadmium in 10 per cent NH4NO 3 , test alkalinity of filtered solution at end, and also test for nitrites. (c) Iron in 10 per cent sodium tartrate. Test as before. Reference. Chem. News, 90, 142 (1904). Part 3. Passive Iron. Discussion. The passivity of iron is probably due to an adsorbed and stabilized film of a higher oxide, the formula of which is possibly FeO 2 . The oxide, which is very difficultly soluble in HNO 3 , is formed by certain oxidizing agents such as HNO 3 , NO 2 , etc., or when iron is made anode in an electrolytic cell through which a sufficiently high current passes. Passivity is removed and activity is restored by destruc- tion of the oxide film. Reducing agents may destroy the film or the same thing may be done by making a passive rod cathode with a sufficiently high current. Consult the Instructor. Procedure. (a) Make an iron rod passive in concentrated nitric acid (sp. gr. = 1.4). Wash in water carefully and dip in dilute HNO 3 (sp. gr. 1.2). What happens? (b) Having immersed the rod in the dilute acid, touch the rod with a fresh (active) iron rod. Explain. Repeat, touching passive rod with zinc. (c) Immerse an active and a passive rod in dilute (1.2) HNO 3f taking care to dip the active rod deeply and the passive rod only slightly beneath the surface of the liquid. Connect the two rods out- side of the cell with a copper wire. What happens? (d) Repeat experiment (c), having a large surface of the passive rod and only a small surface of the active one dipping into the acid. To understand (c) and (d) see Bennett's paper, p. 220. (Schonbein's experiments) . (e) Immerse a passive rod in dilute acid and scratch the surface. Does the rod become active? Reference. Bennett and Burnham: Trans. Am. Electrochem. Soc., 29, 217 (1916). 37 EXPERIMENTAL GROUP X REACTION VELOCITY AND CATALYSIS This group of experiments is designed to illustrate in a semi- quantitative manner the Law of Mass Action and its bearing on the velocity of chemical change. Simple experiments illustrating cata- lysis are also included. Standard References. Bancroft: Papers in Jour. Phys. Chem. (1917 ). Henderson: Catalysis and Its Industrial Applications (1918). Herz: Ahren's Sammlung, 11, 103-145 (1906). Jobling: Catalysis and Its Industrial Applications (1916). Mellor: Chemical Statics and Dynamics (1609). Ostwald: Uber Katalyse (2nd Ed. 1911). Rideal and Taylor: Catalysis in Theory and Practice (1920). van't Hoff: Lectures; Vol. 1, Chemical Dynamics (1898). van't Hoff (Evan) : Studies in Chemical Dynamics (1896). Woker: Die Katalyse (1915-16). Law of Mass Action. The rate at which chemical change occurs is a function of the concentration of each of the substances taking part in the reaction. The rate is also a function of the temperature and pressure and it is affected by catalysts and by various other influences, such as light, electrical and surface forces. The law is illustrated by the reaction between bromic and iodic acids 6 HI + HBrO, -^ HBr + 3 H 2 O + 3 I 2 , in which the course of the reaction can be followed color imetrically, using starch as an indicator. The rate at which iodine is set free is directly proportional to the ion concentrations of iodine and bromate and to the square of the concentration of hydrogen as ion. Clark: Jour. Phys. Chem., 10, 700 (1906). If one keeps the concentration of hydrogen ions con- stant and does not allow the volume of the solution to vary, the velocity with which iodine is liberated at any moment is expressed in terms of the mass law by the equation _^ = k(a x)(b x) (1) dt in which a and b refer respectively to the amount of iodine and bro- mate present as ions at the beginning of the experiment and are there- fore proportional to the initial quantity of HI and HBrO 3 , while x refers to the amount of iodine or bromate ions used up and is accord- ingly proportional to the quantity of free iodine liberated. 38 If the reaction is allowed to proceed for a relatively short time only and in such a way that x is small by comparison with a and b, the velocity equation takes the form ^ X =kab (2)' whence, on integrating between the limits x = o and x = Xi ; t = o, and t = t, the following expression results : t = constant ^L (3) ab General Procedure. In the experiments which follow iodide and bromate are mixed in acid solution and the reaction is allowed to proceed until a definite constant quantity of iodine is liberated, as determined by the forma- tion of a definite "standard" blue color with starch as indicator. The initial quantities of iodide and bromate are varied and the time required to reach the standard blue is determined by means of a stop- watch. Under these experimental conditions, it is evident from equation (3) that the time required to reach a standard blue at constant temperature and volume varies inversely as the product of the initial quantities of iodide and bromate, as long as the amount of iodine set free is small. It is also obvious that this statement becomes less exact as the depth of the standard blue becomes greater. For comparison times, the relative values of a and b may be sub- stituted for absolute values. EXPERIMENT 1 Mass Action Acid Mixture. 800 cc. distilled water 26 cc. N/2 HC1 (shelf) 20 cc. starch solution To prepare the starch solution rub one gram of starch with 5 cc. of cold water in a mortar; pour 150 cc. of boiling water over it, allow the undissolved part to settle, and decant the supernatant liquid. Standard Blue. Take two 100 cc. bottles (glass stoppered) and in one make a standard blue solution as follows: 80 cc. distilled water 2 cc. starch solution (described above) 3-6 drops "iodine mixture" (shelf) Procedure. Part 1 . In the test bottle place 80 cc. acid mixture 1 cc. N/2 KBrO 3 (shelf) 1 cc. N/2 KI (shelf) in the order named. Add the KI quickly and take the time from the moment it is added. Shake at the moment of adding KI and note the time required for the solution to assume the same blue as the standard. Run a parallel. 39 Notes. Place the standard and the test bottle against a white background. Avoid using a standard with too deep a blue. The time taken in Part 1 should not exceed two minutes nor be less than one minute. Be careful to work throughout at constant temperature (20 C.). Record. Part 2. 80 cc. acid mixture 2 cc. bromate 1 cc. iodide Shake. Note time as before. Run a parallel. Part 3. 80 cc. mixture Part 4. 80 cc. mixture 1 cc. bromate 2 cc. bromate 2 cc. iodide 2 cc. iodide Shake. Note time. Run a Shake. Note time. Run a parallel. parallel. EXPERIMENT 2 Catalytic Effect of Acids The effect of acids in accelerating certain chemical reactions is roughly proportional to their electrical conductivity. The effect is dependent primarily on the hydrogen ions. Prepare a mixture as follows : 400 cc. water 10 cc. bromate (shelf) 10 cc. iodide (shelf) 10 cc. starch solution Part 1. Take 80 cc. of the above mixture in a 100 cc. bottle, add 2 cc. N/2 HC1. Shake. Note time required. Run a parallel. Part 2. Take 80 cc. of mixture and 2 cc. of N/2 H 2 SO 4 . Shake. Note time required. Run a parallel. Part 3. Take 80 cc. of mixture and 2 cc. of N/2 CH 3 COOH. Note time required. Shake. Run a parallel. Explain. EXPERIMENT 3 Catalytic Effect of Ferrous Sulphate Mixture of 160 cc. H 2 O. 8 cc. KI (shelf) 8 cc. KBrO 3 (shelf) 4 cc. starch solution Part 1. Take 80 cc. of the above mixture and 10 cc. of N/2 acetic acid. Shake. Note time required to become blue. Part 2. Take 80 cc. of the mixture and 10 cc. of N/2 acetic acid to which is added one drop of neutral saturated FeSO4. Proceed as before. Explain. 40 Part 3. To 25 cc. of an extremely dilute solution of chromic acid (CrO 3 ) add a little starch solution. (a) To 5 cc. of this solution add 2 to 3 drops of KI solution. Note time as before. (b) To 5 cc. of the solution add KI as before and a little iron dust. Note time. (c) To another 5 cc. portion add KI and a few drops of a ferrous sulphate solution. Note time. (d) To another 5 cc. portion add KI and a few drops of ferric sulphate solution. Note time. Part 4. (a) Mix in the following order: Dilute CrO 3 solution, ferrous sulphate solution and starch; shake and wait ten minutes; then add KI. Note time to reach standard blue after adding KI. (b) Mix in the following order: CrO 3 solution, KI solution and starch; wait ten minutes ; then add ferrous sulphate. Note time after adding FeSO 4 . Compare (a) and (b) and explain. EXPERIMENT 4 Hydrolysis of an Ester Catalysis Place 50 cc. of distilled water and 5 cc. of ethyl acetate in a clean, glass stoppered bottle. Shake thoroughly and titrate duplicate samples (2 cc.) with N/10 NaOH, phenolphthalein as indicator. In a second bottle place 50 cc. N/2 HC1 plus 5 cc. ethyl acetate. Shake and titrate as before. Set both bottles aside for 24 hours (shaking occasionally) and again titrate duplicate samples (2 cc.). . Note differences and explain. How is this phenomenon used to measure the strength of acids? EXPERIMENT 5 Reactions in Heterogeneous Systems Part 1. Size of Particles. Whenever one of the reacting sub- stances is a solid, the speed of the reaction is a function of the surface area of the solid, or more accurately, of the surface per unit weight of solid (specific surface). The specific surface, in turn, is a function of the size of the particles and increases rapidly as the particles become smaller. Read W9 Ostwald: Grundriss der Kolloidchemie, 30 (1912). Prepare about 2 grams of finely divided copper by placing some granulated zinc in a concentrated solution of CuSO 4 . Shake from time to time to remove the finely divided copper from the zinc. After most of the copper has been precipitated, remove the zinc, wash the precipitate with water and dry in an air bath. Mix the finely divided metal with powdered sulphur and ignite cautiously with a match. What is formed? 41 Dissolve sulphur in CS 2 and into this solution dip a clean strip of copper. What is the substance formed on the copper? Show how this experiment illustrates the principle discussed. Part 2. Protecting Films, (a) Clean a strip of aluminum foil by immersing it in 10 per cent NaOH. Rinse and plunge the wet metal quickly into clean mercury. Hold it there until amalgamated. Remove and rub off the excess of mercury adhering to the aluminum, then expose to the air. What happens? Explain. (b) Place freshly amalgamated aluminum in contact with warm water. What happens? Compare with sodium. (c) Dip a rod of metallic magnesium into warm water. What happens? (d) Dip a rod of metallic magnesium into warm NH 4 C1 solution. What happens? Explain. The passivating films might be regarded as negative catalysts. References. Ostwald: Prin. Inorg. Chem., 560 (1900); Wis- licenus: Jour. Praktische Chemie, (2), 54, 41 (1896). Note. The amalgamated aluminum may be prepared by cleaning the metal in 10 per cent NaOH, rinsing carefully and then dipping the wet metal into dilute mercuric chloride. Part 3. Reactions between Solids. Incompatible Hydrates. Use small quantities in proportions approximately equivalent. Weigh out roughly, except in (a), where a few crystals are enough. (a) Grind together in a mortar HgCl 2 + KI. (b) " " "" " Na 2 SO 4 -10H 2 O + NH 4 NO 3 . (c) " " " " " (NH 4 ) 2 SO 4 + NaNO 3 . 00 g o O I-H 'S 1 ^ M w T3 i_) o ~ pS "o o OJ ^o 'o O 1 l-a 5=1 > I OH K^S TJ o 1 o 'S - b'A M ^ac' - a'c/ B Procedure. Prepare a solution of 50 g ; sodium sulphate decahydrate (Glauber's salt) and 10 g. sodium chloride in 100 cc. of distilled water. Filter the hot solution. Cool to 45 C. and analyze the solution for sodium chloride and sodium sulphate. See below for procedure. Keep the solution in a stoppered flask or Erlenmeyer. Run in duplicate. Cool the solution until solid crystallizes out in considerable amount, then carefully pipette two samples of the solution for analysis. It may be found advisable to fit to the end of the pipette a bit of glass tubing containing glass wool or cotton to serve as a filter. Separate some of the solid and wash with a very little water. Has any sodium chloride been precipitated? Analysis. Determine NaCl in one sample (1 g.) with standard silver nitrate (shelf) using K 2 CrO 4 as indicator. Evaporate a second sample to dryness (being careful to avoid spattering) and determine total chloride and sulphate. Determine water by difference. Using equation (6) determine the chemical formula of the solid phase, assuming that no solid solutions are formed in this experiment. How else might one determine approximately the composition of the solid in the above experiment, using, of course, an indirect method? Outline the procedure in case component C also separates out in the solid phase. See references (Triangular Diagrams). How could one distinguish between compound and solid solution? Why must the two salts have an ion in common? How can one tell whether the number of solid phases precipitated from the solution is one or two? EXPERIMENTAL GROUP XVI COLLOID CHEMISTRY This comprehensive group of experiments serves to illustrate some of the more important and interesting properties of colloidal systems. Typical colloids are prepared and studied, particularly from the point of view of Bancroft: Jour. Phys. Chem., 18, 549 (1914). Read the article before beginning experimental work in this group. General Texts in Colloid Chemistry. Alexander: Colloid Chemistry (1919). Bancroft: Applied Colloid Chemistry (1920). Burton: Physical Properties of Colloidal Solutions (1916). Cassuto: Der Kolloide Zustand der Materie (1911). Freundlich: Kapillarchemie (1909). Hatschek: An Introd. to the Physics and Chemistry of Colloids (1919). Miiller: Chemie der Kolloide (1907). Ostwald (w) : Grundriss der Kolloidchemie (1911-12). Ostwald (w) (Fischer): Theoretical and Applied Colloid Chemistry (1915). Ostwald (w) (Fischer) : Handbook of Colloid Chemistry (1915) Svedberg: Die Methoden zur Herstellung kolloider Losungen usw. (1909). Taylor: The Chemistry of Colloids (1915). Willows and Hatschek: Surface Tension (1915). Zsigmondy (Alexander) : Colloids and the Ultramicroscope (1909) . Zsigmondy: Kolloidchemie (1912). Zsigmondy (Spear): Colloid chemistry (1917). Journals Journal of Physical Chemistry, (special articles). Kolloidchemische Beihefte (special articles). Kolloid-Zeitschrift. (1906). SUB-GROUP 1 DIFFUSION, DIALYSIS AND MEMBRANES EXPERIMENT 1 Diffusion of Solutions. Obtain six test tubes, fitting each with a rubber stopper (one hole), and prepare six 15 cm. lengths of narrow-bore (2. 5-3 mm. internal diam.) glass tubing. Seal one end of each length of tubing and fill 70 completely with distilled water. Place 10 cc. of solution to be tested in each test tube, insert a water-filled diffusion tube in the stopper and place it in the test tube, immersing open end of the diffusion tube just below the surface of the solution. Work carefully. Set aside the test tubes in a safe place and make observations at regular inter- vals, recording the time. Test the following solutions: KMnO 4 solution N/50. KMnO 4 solution N/5. Congo red 1/5 of one per cent. Methyl violet or safranine 1/5 of one per cent. Arsenious sulphide sol. (See Part 3 below). Ferric oxide sol. (Sse Part 3 below). Optional Method. The following experiments are similar to those of Graham. A small, two-dram vial is fastened to the bottom of a tall, narrow beaker (250 cc. capacity) by means of paraffin. Fill the vial carefully with the solution containing the solute whose rate of diffusion is to be measured and cover it securely with a small cover-glass (20 millimeters). Be sure that no solution is spilled from the vial during the process of filling and covering. Pour dis- tilled water into the beaker until it is nearly full and the vial is well covered, taking care to have the water level at the same height in each beaker. Finally, slide the cover glass carefully off the mouth of the vial by means of a clean glass rod. A two cc. test-sample is then pipetted from the liquid in the beaker at a point about three centimeters above the open mouth of the vial. Mark this position by means of a label placed on the wall of the beaker. Be careful not to stir the liquid. Test for chlorine as ion with silver nitrate making a rough nephelometric estimation of the relative amounts of silver chloride formed in each sample. Test for organic matter by evaporating a test sample to dryness in a clean porcelain dish and carbonizing the residue. It is essential that the water levels be the same in each beaker, that the sample be pipetted from equal distances above the mouth of the vial and that the beakers and solution remain absolutely undisturbed. Withdraw test samples at the beginning and after 1, 2, 4 and 7 days, noting the exact time. The following solutions are to be tested: (1) One per cent solution of gelatine. (2) Five per cent solution of sodium chloride. (3) Twenty-five per cent solution of sodium chloride. Note. Prepare a 5 per cent solution of gelatine for this and subse- quent work as follows: Soak 2 g. of gelatine in cold water until soft, pour off. the water and to the softened gelatine add enough warm water to make about 40 cc. of solution. On cooling, a jelly will form which readily melts when the beaker with the jelly is warmed on the steam bath. Do not warm over a flame as the beaker will almost certainly crack. Dilute the gelatine solution as required. 71 EXPERIMENT 2 Diffusion Through a Jelly Obtain eight small test tubes and fill each half full of liquid 5 per cent gelatine and allow this to solidify. Pour into the tubes, on top of the gelatine, the solutions or sols specified below, being careful that the latter are cold so that they do not liquefy the jelly. If they diffuse, the substances in solution will tend to pass from the upper aqueous layer into the lower portion occupied by the gelatine and the process may be observed by means of the coloration produced in the jelly. If the colored substance forms a true solution, the diffusion of the solute through a jelly occurs almost as rapidly as through pure water itself. On the other hand, colloidal solutions show practically no evidence of diffusion. We may, therefore, dis- tinguish between the two classes of solution by means of this method, provided the jelly is not "semi-permeable" to the dissolved solute. Observe the condition of each tube after twenty-four hours and again after a week. Keep the tubes in a cool place. Use the follow- ing solutions (shelf) : (1) Eosine (2) Congo red (3) Safranine (4) Picric acid (5) Methylene blue (6) Arsenious sulphide sol 1/5 of one per cent. 1/5 of one per cent. 1 /5 of one per cent. 1/5 of one per cent. 1/5 of one per cent, (see below). -(7) Ferric 'oxide sol (see- below). (8) Mixture Congo red and picric acid, picric acid in excess. From the data obtained in these experiments what conclusion do you draw regarding the nature of the above solutions? EXPERIMENT 3 Dialysis v/ith Collodion. Instead of using parchment, prepare collodion dialyzing tubes as follows: Take one of the inner test tubes of heavy glass used in the free2ing point determinations and wet the inner walls completely with a fairly thick film of collodion solution (soluble cotton in a mixture of ether and alcohol) . Do this quickly while spinning the tube to make the collodion film uniform. As soon as the collodion "sets" blow air into the tube to remove the ether. This process should take about five minutes. Then pour water into the test tube and gradually loosen the collodion from the glass. With moderately careful manipulation, a transparent, tough dialyzing tube can be obtained which is more convenient and less expensive than the parchment dialyzers ordinarily used. Having prepared the tube, test for leaks by filling with water and if intact, immerse completely in a large beaker of water to remove the alcohol. Soak until the next period, changing the water from tirn.e to time. Make three dialyzing tubes. Fill one nearly full with a mixsd solution containing 1 per cent gelatine plus 5 per cent of sodium chloride. Place this in a beaker of distilled water and test the water at stated interval for NaCl and .gelatine. 72 Fill the second tube with a solution of safranine. Place this in a second beaker of water and observe diffusion. In the third tube place a solution of Congo red. Does this diffuse? EXPERIMENT 4 Semipermeable Membranes Into a small bottle pour, very carefully and in the order given, the following liquids: Chloroform, water, and ether. Three layers should be present. Note the thickness in mm. of each layer. Let the bottle stand undisturbed for a week and again measure the thickness of the layers. Continue the experiment until one of the three original layers disappears. Explain. Reference. Kahlenberg: Jour. Phys. Chem., 10, 146 (1906). EXPERIMENT 5 Osmosis and Semipermeable Membranes Part 1. Fill a test tube with a M/2 CuSO 4 , then, by means of a pipette placed in this solution add slowly and carefully a small amount of M/2 potassium ferrocyanide. A globule should form, consisting of the solution of ferrocyanide surrounded by a gelatinous membrane of brown copper ferrocyanide. Carefully detach the globule from the end of the pipette and it will sink, owing to the greater density of the ferrocyanide solution. Observe carefully any changes that may occur in the copper sulphate solution surrounding the globule. Set aside the test tube and keep it constantly under observation. What happens? Explain. Part 2. Plant-like Growths. Fill a small beaker with dilute sodium silicate (water glass) solution and drop into the liquid one or two crystals each of CuSO 4 , MnSO 4 , CoSO 4 , etc. What happens? Explain. SUB-GROUP 2 ADSORPTION The following experiments are designed to illustrate adsorption phenomena. Adsorption is the basis of colloid chemistry. All the experiments of Sub-groups 3 and 4 illustrate this point. EXPERIMENT 1 Adsorption by Bone Black Part 1. Boil a dilute solution of litmus with bone black. Filter. Part 2. Repeat, using dilute solution of indigo. Are the colors removed? Explain. Part 3. Prepare a dilute solution of silver nitrate. Divide this into two portions. To one portion add about one-tenth its volume of bone black and shake vigorously for at least three minutes. Then 73 filter and add NaCl to both portions. Compare the amounts of precipitated silver chloride. Bone black or animal charcoal contains 85 per cent. of calcium phosphate and about 15 per cent of carbon. EXPERIMENT 2 Selective Adsorption Part 1. Prepare about 250 cc. of indicator solution as follows: To 250 cc. of distilled water add a little phenolphthalein and a trace of NaOH, just enough to color the liquid pink. Part 2. Ina test tube shake fuller's earth with distilled water and add some of this muddy suspension to one of the test tubes containing the indicator. Is the color removed? Part 3. Allow this muddy suspension to settle and then add the supernatant clear liquid to a second test tube colored with indicator. Filter the supernatant liquid to remove all the fuller's earth. Is this filtered liquid acid? Part 4. Moisten a little fuller's earth with boiled water and test with blue litmus by pressing the latter down on the earth. Reference. Cameron: Jour. Phys. Chem., 14, 400 (1910). Part 5. Add blue litmus solution to fuller's earth suspended in water. Notice the change. Part 6. Add some fuller's earth to a dilute solution of methyl violet and shake. Filter, noting color of filtrate and of earth. Is the color removed from the earth by water or alcohol? Part 7. Repeat the last experiment, using eosin instead of methyl violet. Note any differences in behavior. Part 8. Moisten some absorbent cotton with freshly boiled water (free from CO 2 ) and wrap it around a strip of blue litmus paper. For comparison of the original and the final color, let about half an inch of the paper protrude beyond the cotton. Explain your results. Compare Part 4, above. EXPERIMENT 3 Adsorption by Iron Oxide. The Antidote for Arsenic Poisoning Hydrous ferric oxide is precipitated from a solution of ferric sulphate or chloride by adding an excess of magnesia. Shake vigorously. Then prepare a dilute solution of As 2 O 3 and filter, and test the filtrate for arsenic with H 2 S. Be sure that the As 2 O 3 solution is very dilute. Test half the original solution with H 2 S for arsenic. Only a slight test should be obtained, if the experiment is to work well. Then test the second half of the As 2 O 3 solution after treatment with the ferric hydroxide mixture. Has the arsenic been adsorbed? Should the arsenic be completely adsorbed? Explain. 74 EXPERIMENT 4 Adsorption Compounds. Carey Lea's "Photohalides" Reference. Carey Lea: Am. Jour. Science, (3) 34, 349, 480, 489 (1887). Method suggested by Luppo-Cramer : Kolloid-Zeit., 2, 360 (1908). To 3.5 cc. of 10 per cent KBr add 5.5 cc. of 10 per cent AgNO 3 . To this mixture containing AgBr plus AgNO 3 in excess add the following solution: 7.5 cc. Rochelle salts (1:3) plus 2.5 cc. of ferrous sulphate (1:3). Do not add the Rochelle salts and ferrous sulphate solutions sepa- rately. Wash the dark colored precipitate several times by decantation and finally with a mixture of equal parts concentrated HNO 3 (1.4 sp.gr.) and water. An intense blue-violet color should develop. The photohalides of silver are adsorption compounds of silver with silver chloride and are similar to the "subsalts" of silver composing the "latent image" in an exposed photographic plate. EXPERIMENT 5 Selective Adsorption and Capillary Diffusion Part 1. Place several drops of a mixed solution of CuSO 4 and CdSO 4 (shelf) on the center of a square of blotting paper (6 by 6 in.). Allow the drops to diffuse until a large round spot has formed, then hold the paper in a stream of H 2 S gas. Which "diffuses" farthest, water, CuSO 4 , or CdSO 4 ? Cf. Gordon: Jour. Phys. Chem., 18, 337 (1914). Part 2. Suspend strips of blotting paper (1 cm. broad and 20 cm. long) in water solutions of the following substances: Congo red; picric acid; cosin; methylene blue; methylene blue plus cosin. Note the height to which the water and dye rise. Reference. Goppelsroeder: Kapillaranalyse (1906). EXPERIMENT 6 Adsorbed Air in Charcoal Fit a cylinder (100 cc.) with a three-hole rubber stopper. Into one hole introduce the delivery tube of a burette filled with water. In the second place a thermometer. In the third place a glass tube lead- ing to a water- filled graduated cylinder (capacity 250 cc.) inverted over water in a trough. Place a volume of 50 apparent cc. of granular cocoanut charcoal in the cylinder. Then add water slowly from the burette, recording the volume added. Continue to add water until its level rises to the surface of the charcoal. Measure the volume of air displaced. Take the temperature before and after adding the water. Have the water in the burette and the charcoal at the same temperature in the beginning. 75 SUB-GROUP 3 PEPTIZATION EXPERIMENT 1 Peptization by Adsorbed Ions Lottermoser: Jour. Praktische Chemie, [2] 72, 39 (1905); 73, 374(1906); Zsigmpndy (Spear) 179; Ostwald (Fischer) : Theoretical and Applied Colloidchemistry, 115. Part 1. Prepare a small quantity of silver bromide and wash the precipitate thoroughly by decantation. Place approximately equal amounts of the freshly prepared silver bromide in each of five stoppered test tubes. In the first test tube place distilled water (10 cc.); in the second, N/100 KBr; in the third, N/30 KBr; in the fourth, N/10 KBr, and in the fifth, N/5 KBr. Shake thoroughly and after allowing the test tubes to remain standing several minutes, describe the appearance of each tube. In which is the supernatant liquid most turbid? The process constitutes a dispersion method of preparing colloidal silver bromide. Part 2. Fill two burettes with N/20 AgNO 3 and N/20 NH 4 CNS (shelf). Fit a small Erlenmeyer flask with a solid rubber st.opper. Perform the following experiments: (a) To 10 cc. AgNO 3 in flask add quickly 10 cc. NH 4 CNS, stopper and shake. (b) To 10 cc. AgNO 3 in flask add quickly 12 cc. NH 4 CNS, stopper and shake. (c) To 10 cc. NH 4 CNS in flask add quickly 10 cc. AgNO 3 , stopper and shake. (d) To 10 cc. NH 4 CNS in flask add quickly 12 cc. AgNO 3 , stopper and shake. What striking differences do you observe and how do you account for them? (e) Refill the burettes and, placing 10 cc. AgNO 3 in a flask run in NH 4 CNS from a burette (not too rapidly) until floccula- tion occurs. Shake and note the volume of NH 4 CNS added. Repeat adding NH 4 CNS more slowly as end- point is reached. The end-point represents the isoelectric point (define). (f) Place 10 cc. NH 4 CNS in a flask and add AgNO 3 following the procedure of (b) 5 preceding. Explain. EXPERIMENT 2 Peptization by Adsorbed Colloid Prepare 5 per cent solutions of chromic and ferric chlorides. Mix in the proportions specified below. Then add 10 per cent NaOH in excess. Note the color and appearance of the precipitate (if any) and of the supernatant liquid. Use test-tubes and shake. 76 Ferric Chloride Chromic Chloride Remarks (cc). (cc.) 10 8 2 5 5 3 7 2 8 1 9 10 Cf. Nagel: Jour. Phys. Chem., 19, 331, 569 (1915). EXPERIMENT 3 Peptization by Adsorbed Colloid (Protective Colloids) Solution A: 5 cc. N/2 AgNO 3 + 5 cc. of 5 per cent gelatine. Solution B: 5 cc. N/2 KBr + 5 cc. of 5 per cent gelatine. Part 1. After thoroughly mixing each solution, pour B into A, shake and note any changes. Place the mixture in the sunlight and note results. Repeat the above experiment, replacing the gelatine solution by an equal volume of pure water. Was AgBr formed in the first experiment with gelatine. How might one prove this? Part 2. Prepare some silver bromide, wash by decantation and remove to a filter paper. Divide into two portions. Place one por- tion in an air bath and dry for an hour at 120, being careful not to exceed this temperature. To the freshly prepared moist silver bromide add a few cubic centi- meters of hot 5 per cent gelatine and shake vigorously. Is a suspen- sion formed? Do the same thing with the dried silver bromide and note any differences in its behavior compared with that of the freshly prepared substance. What is the effect of "ageing?" Part 3. Grind a little roll-sulphur with a 5 per cent gelatine solu- tion in a mortar until a milky suspension is formed. Pour some of this suspension into water and note the color. SUB-GROUP 4 PREPARATION AND FLOCCULATION OF SUSPENSIONS EXPERIMENT 1 Colloidal Arsenious Sulphide (Condensation Method) In a clean beaker, boil about 6 grams of As 2 O 3 with 100 cc. of distilled water for fifteen minutes. Cool, filter and dilute to 100 cc. Pass clean hydrogen sulphide gas into the solution of arsenious acid until no further action takes place. Remove excess of H 2 S by blow- ing a slow stream of air through the suspension and then filter. Describe the appearance of the suspension as to color, turbidity, etc., and perform the following tests. (See also diffusion experi- ments) . 77 (a) To 10 cc. colloidal As 2 S 3 add 2 cc. M/20 HC1. (b) To 10 cc. colloidal As a S 3 add 2 cc. M/20 NaCl. (c) To 10 cc. colloidal As 2 S 3 add 2 cc. M/20 MgCl 2 . (d) To 10 cc. colloidal As 2 S 3 add 2 cc. M/20 Al (NO 3 ) 3 . (e) To 10 cc. colloidal As 2 S 3 add 2 cc. M/20 Na 2 SO 4 . Which produces flocculation most quickly? Explain. Colloidal As 2 S 3 thus prepared is a negative suspension. That is, the particles of the disperse phase carry a negative charge due to preferential adsorption of anions from H 2 S present in solution. Place a test tube containing 10 cc. of As 2 S 3 suspension in an ice salt freezing mixture until frozen solid. Warm the test tube gently until the ice is melted. What effect upon the suspension is noticed? EXPERIMENT 2 Colloidal Ferric Oxide (Condensation Method) Add about 0.5 gram of crystallized ferric chloride to 100 cc. of boiling distilled water. Then boil the solution gently for about ten minutes, replacing the water boiled away. Note the color and appear- ance of the hot solution, and compare with the color of a solution made by adding FeCl 3 to cold water. Explain the change. What is this process called? (a) To 10 cc. of the iron oxide suspension add 2 cc. M/20 NaCl. (b) To 10 cc. of the iron oxide suspension add 2 cc. M/20 MgCl 2 . (c) To 10 cc. of the iron oxide suspension add 2 cc. M/20 Na 2 SO 4 . (d) To 10 cc. of the iron oxide suspension add 2 cc. M/20 citric acid (e) To 10 cc. of the iron oxide suspension add trace of H 2 SO 4 . Which causes the most rapid flocculation? Explain. What is the precipitate? The ferric oxide suspension as prepared above is positive. Optional Experiment. Colloidal Ferric Oxide (Dispersion Method) Reference. Kratz: Jour. Phys. Chem., 16, 126 (1912). Prepare Fe 2 O 3 suspension by the method of washing out the coagu- lating salt, following Kratz's procedure. % EXPERIMENT 3 Mutual Flocculation of Two Suspensions Study the mutual flocculation of colloidal As 2 S 2 and Fe 2 O 3/ two oppositely charged suspensions. Plan your own experiments. EXPERIMENT 4 Colloidal Silica. (Condensation Method) To 10 cc. of syrupy sodium silicate solution add 30 cc. of water and pour the resulting solution into a mixture of 25 cc. of concentrated hydrochloric acid previously diluted with an equal volume of water. A limpid mixture will result, consisting of a suspension of hydrated silica. 78 Warm some of this solution nearly to boiling and allow it to stand undisturbed for a few minutes. What has occurred? Can the sus- pension be restored? Study the jelly obtained. How does it differ from gelatine or agar agar? EXPERIMENT 5 Colloidal Metals (Condensation Methods) Part 1. Colloidal Silver. Gelatine as Protecting Colloid. To 5 cc. of water in a test tube add about 1 cc. of M/10 AgNO 3 solution, mix well and treat with NaOH in slight excess. What is formed? To 5 cc. of a 5 per cent gelatine solution in a test tube add about 1 cc. M/10 AgNO 3 , mix well and treat with NaOH in slight excess. Note any unusual action. Then heat the test tube until contents are about to boil. What color changes occur? Dilute some of the silver sol so formed with water and describe its color. What reduces the silver oxide? Repeat the above experiment, using a drop or two of hydrazine hydrate as the reducing agent, besides gelatine. If unsatisfactory results are obtained, repeat the experiment, using smaller amounts of AgNO 3 solution and varying other conditions until successful. Part 2. Colloidal Silver. Method of Carey Lea. Prepare two solutions as follows : Solution A. Mix: 10 per cent silver nitrate solution 20 cc. 20 per cent Rochelle salts solution 20 cc. distilled water 80 cc. Solution B. Mix: 30 per cent ferrous sulphate solution . . 10.7cc. 20 per cent Rochelle salts solution 20 cc. distilled water 80 cc. Pour B slowly into A, stirring rapidly. The solutions must be freshly prepared and the work should be done in light as weak as possible. Throw out the precipitated silver by means of a centrifuge, wash with 2 per cent Rochelle salts solution and again separate in a cen- trifuge. Obtain a camels-hair brush and paint some of the silver on a watch glass. Dry slowly (without heating above 50 C.) and note the color of film obtained. Place a crystal of iodine in the center of the yellow silver film. Record all that happens. Explain. References. Carey Lea: Am. Jour. Science, (3) 37, 476 (1889); 38, 47, 129, 237(1889); 41,179,259,482(1891); Blake: Zeit, anorg. chem., 37,243(1903); also Svedberg: Herstellung (1909). Part 3. Colloidal Copper (Gelatine as Protecting Colloid.) Mix equal volumes (5 cc.) of 10 per cent gelatine solution (freshly prepared and warm) and 5 per cent copper acetate. To this solution 79 add, with shaking, a very slight excess of sodium hydroxide (20 per cent). A purplish-blue, clear solution should result. If a persistent precipitate remains, repeat the experiment, using a more concentrated gelatine solution. Perform the same experiment, using 5 cc. of water in place of the gelatine. What is the precipitate? Does it dissolve in an excess of sodium hydroxide? Heat some of the purplish-blue copper oxide-gelatine solution to boiling and add a few drops of hydrazine hydrate. The latter is a very powerful reducing agent and will reduce the oxide to metallic copper in alkaline solution. Continue gently to heat the reaction mixture until a dark, blood-red liquid is produced. The red color is due to finely divided copper. Pour some of the liquid into water, noting its beautiful color. In this connection cf. Paal: Ber. 35, 2206,2219 (1902). EXPERIMENT 6 Colloidal Sulphur (Condensation Method) Reference. Raff 6: Kolloid-Zeit., 2, 358 (1908); 8, 58 (1911). Place a cylinder containing 70 grams of concentrated sulphuric acid (sp. gr. 1.84) in ice water or in a freezing mixture and into it pour, drop by drop and with constant stirring a cold solution of 50 grams of pure crystallized sodium thiosulphate in 30 cc. of distilled water. Work at the hoods, as H 2 S and SO 2 are given off. When the reaction is complete, transfer the mixture to a beaker, add 30 cc. of distilled water and warm to 80 on a water bath until SO 2 and H 2 S cease to be given off. Then prepare a Buchner funnel and filter, connect with the suction and pour in hot water until the funnel and filter-flask are warm. Pour out this wash water and filter the hot sulphur hydrosol. Cool the warm filtrate in ice water and decant the supernatant acid liquid. To some of the precipitated sulphur add water. Is it peptized? To 10 cc. of this suspension add a little saturated K 2 SO 4 . What happens? To 10 cc. add some Na 2 SO 4 solution. Is flocculation so easy? Note difference between K 2 SO 4 and Na 2 SO 4 . Flocculate some of the sulphur suspension by adding a soluble salt of potassium and allow the sulphur to settle. Decant the super- natant liquid and wash once by decantation. Then add water to the precipitate of sulphur and shake until a coarse yellow suspension of sulphur is formed. To this add a tiny crystal of Na 2 SO 4 . Con- tinue to add salt cautiously until a clear yellow suspension of sulphur is formed. What is this process called? When a large excess of sodium sulphate is added, what happens? SUB-GROUP 5 EMULSIONS References. Bancroft: Jour. Phys. Chem., (1912-1918); Briggs: Ibid., 19, 210, 478 (1915); 24, 147 (1920). 80 EXPERIMENT 1 Oil-in-Water Emulsions Part 1. In a 150 cc. glass stoppered bottle place 45 cc. of benzene plus 5 cc. of 1 per cent sodium oleate solution. Then shake the bottle and contents steadily and without interruption until the ben- zene is completely reduced to a milk-white emulsion and no free benzene remains floating at the surface. Note the time required and the approximate number of shakes. Part 2. Discard the emulsion by emptying into the bottle marked "benzene residues" and repeat the experiment making a change, however, in the method of shaking. Give the bottle two violent up and down shakes and then let it stand on the desk for a "rest interval" of about thirty seconds. Continue the intermittent shak- ing until emulsion is completed. Note the time and approximate number of shakes. Compare with (1). Explain. Part 3. Again discard and make the emulsion in still another way, as follows : In glass stoppered bottle, place 2 cc. of sodium oleate solution and to this add 1 cc. of benzene from a burette. Shake thoroughly until all the benzene is emulsified. Then add another cc. of benzene and again shake. Repeat this process until about 100 cc. of benzene have been emulsified. An emulsion having the consistency and appearance of blanc-mange should result. As the volume of emulsion increases, more benzene may be added each time before shaking, but if too much is added the emulsion may "break" and a fresh start become necessary. Add a drop of HC1 to some of this emulsion. What happens? Explain. In this emulsion the oil (benzene) exists in drops (disperse phase) and the soap solution is the dispersion medium. EXPERIMENT 2 Water-in-Oil Emulsions In a 200 cc. bottle, as in the previous experiment, place 10 cc. of a benzene solution of magnesium oleate. Add water from a burette slowly and with shaking, following a procedure similar to that of the preceding experiment, until 40 cc. of water have been added. How does this emulsion compare with the benzene-in-water one? In this case the water forms the drops (disperse phase) and the soap solution is the dispersion medium. This may be proved as follows: Proof. On a glass plate place a drop of water and with a glass rod stir in some of the emulsion formed in Experiment 1 . Does it mix easily? On another portion of the plate place a drop of benzene and stir in some of the emulsion. Does it mix easily? Do the same thing with some of the emulsion obtained in Ex'-eri- ment 2, that is, stir it into water and into benzene. If the aqueous liquid is the outside phase the emulsion will mix easily with water, but not with benzene. The reverse is true when benzene forms the outside phase. Newman: Jour. Phys. Chem. 13,35(1914). 81 EXPERIMENTAL GROUP XVII THERMOCHEMISTRY It is the purpose of the following group of experiments to study the thermal effects accompanying chemical action, change of state and similar phenomena. Many instances of such thermal effects have been met with in previous experiments. References. Thomsen (Burke): Thermochemistry (1908). Thomsen: Thermochemische Untersuchungen (1882-1886). Sackur (Gibson) : Thermochemistry and Thermodynamics (1917) . Journal articles. Mathews and Germann: Jour. Phys. Chem., 15, 73 (1911); Richards and Rowe: Proc. Amer. Acad., 43, 475 (1908) ; Richards: Jour. Am. Chem. Soc., 31, 1275 (1909). Procedure in Laboratory. F, 273-293 (1917); OW, 119-138; T, 132-152. General Directions. For this work a simple, home-made calorimeter may be obtained from the Instructor. Two special thermometers are also supplied. These must be com- pared with each other in the usual way by heating in a well-stirred water-bath between 10 and 30 C. Number each thermometer and reduce all subsequent readings of temperature to readings on one of these thermometers. Having assembled the calorimeter, determine the water equivalent by experiment several times. How does this compare with the calculated water equivalent? Note. Mix weighed and approximately equal amounts of cold and warm water so that the final temperature of the mixture is about equal to that of the room. Weigh out water to grams only on the large balance. Report the water equivalent before proceeding with the experiments which follow. EXPERIMENT 1 Heat of Solution Part 1. Qualitative. Half fill a test tube with finely powdered dry NH 4 NO 3 and close tube with a rubber stopper. Then add quickly an equal volume of cold water and mix violently to produce instantaneous solution. Then observe the temperature of the solu- tion. Explain the extraordinary drop in temperature. How does 82 this method of making a freezing mixture compare with the usual one (ice-salt) ? Explain. Read the quaint old paper on this subject by Robert Boyle, re- printed in the Philosophical Transactions of the Royal Society (Lon- don), 1, 86 (1666). Part 2. Quantitative. Procedure. T, 137. The weighed solute is introduced into a known amount of water contained in the calorimeter. A convenient method is to make a thin walled glass bulb, fill it with the solute and place it in the calorimeter. When bulb and water are at the same temperature, break the glass and allow the solute to dissolve as quickly as possible. See that {he solute is very finely pulverized. Take the substance assigned from the following: (1) NH 4 NO 3 in 200 gram molecules of water. (2) KNO 3 in 200 gram molecules of water. (3) NH 4 C1 in 200 gram molecules of water. (4) KC1 in 200 gram molecules of water. Measure the heat of the solution and derive equation (1) below, Computations. S = p(a + w) (t a ti) (1) S = heat of solution in small calories; t t = initial temperature of water and bulb in calorimeter; t 2 = final temperature when solution is complete; a= grams of water;, w = water equivalent; 1/p = fraction of required molecular quantities actually used experimen- tally. For further explanation refer to Experiment 3 following. EXPERIMENT 2 Heat of Dilution Procedure. T, 139. In this experiment the solution to be diluted is placed in the upper vessel and the water is placed in the calorimeter. The solution and water are then mixed and the thermal affect measured. Determine the heat of dilution when a solution represented by NaCl -f 10H 2 O is diluted with 40 gram molecules of water. Derive equation (2) below. Computations. D = p { (tf tb) [(a + b) c + w] (t a tb) (a + w) } (2) D = heat of dilution in small calories ; ta = initial temperature of water; tb = initial temperature of solution; tf = corrected final temperature of mixture whose specific heat = c ; w = water equiva- lent; a = grams of water; b = grams of solution to be diluted; 1/p = fraction of required molecular quantities actually used experimentally. 83 EXPERIMENT 3 Heat of Neutralization of 'Acids and Bases Procedure. T, 133. Place the acid in the calorimeter and" the base in the upper vessel. Mix and measure the heat change. Computations. N = p [b (tf tb) + (a + w) (tf t a )] (3) N = heat of neutralization in small calories; t a = temperature of acid; tb = temperature of base ; tf = temperature of mixture ; a = grams of water contained in solution of acid; b = grams of water contained in solution of base; w = water equivalent. . 1/p = frac- tion of required molecular quantities used experimentally. Here the specific heat of the mixture is assumed to be unity. Derive equation (3). Part 1. Sulphuric Acid and Sodium Hydroxide. Measure the heat of neutralization for each of the following cases: (a) (2 NaOH + 200 H 2 O) + (1/2 H 2 SO 4 + 200 H 2 O). (b) (2 NaOH + 200 H 2 O) + (H 2 SO 4 + 200 H 2 O). (c) (2 NaOH + 200 H 2 O) + (2 H 2 SO 4 + 200 H 2 O). Part 2. Phosphoric Acid and Sodium Hydroxide. (a) (H 3 P0 4 + 200 H 2 0) + (NaOH + 200 H 2 O). (b) (H 3 PO 4 + 200 H 2 O) + (2 NaOH + 200 H 2 O). (c) (H 3 PO 4 + 200 H 2 0) + (6 NaOH + 200 H 2 O). In this work one is dealing with molecular quantities of the sub- stances involved. For instance (2NaOH + 200H 2 O) means 80 grams of NaOH dissolved in 3600 grams of H 2 O. Again, (1/2H 2 SO 4 + 200 grams H 2 O) means 49 grams of H 2 SO 4 in 3600 grams of H 2 O. Obviously such volumes of acid and base cannot be handled con- veniently, so one chooses some convenient fractional part of the acid and base solution, for example, 1/16 whence 1/p = 1/16. The quan- tity of the solutions to use in the case of H 2 SO 4 and NaOH (Part 1) would be found thus: 1/16 (80 + 3600) = 230 grams of the NaOH solution. 1/16 (49 + 3600) = 228 grams of the H 2 SO 4 solution. To make up this acid solution mix 3.06 grams of H 2 SO 4 with .225 grams of H 2 O. H 2 SO 4 and H 3 PO 4 tables may be found in the Kalendar, Vol. 1, and elsewhere. EXPERIMENT 4 Thermoneutrality of Salt Solutions Measure the heat change when solutions of the following pairs are mixed: 1 . NaCl + 200H 2 O and KNO 3 + 200 H 2 O. 2. NH 4 C1 + 200H 2 O and KNO 3 + 200 H 2 O. Take the pair assigned, placing one solution in the upper vessel and the other in the calorimeter. 84 EXPERIMENTAL GROUP XVIII PHOTOCHEMISTRY The purpose of the following experiments is to study qualitatively the action of light in producing and accelerating chemical change. References. Bancroft: Electrochemistry of Light, Jour. Phys. Chem. (1908-1912). Bancroft: Orig. Comm. 8th Int. Cong. App. Chem., 20, 31 (1912). Sheppard: Photochemistry (1914). EXPERIMENT 1 Soluble and Insoluble Sulphur Saturate 10 cc. of CS 2 with roll sulphur. Work in the hood. Then divide into three portions and place in loosely stoppered test tubes. (a) Expose one test tube to direct sunlight. After precipitation of amorphous sulphur has occurred, set aside in a dark place. The amorphous sulphur will dissolve. It may be necessary to wrap the test tube in dark paper to protect it from the light. (b) Place another portion in a test tube which is immersed in a solution of CuSO 4 . (c) Place the third portion in a test tube which is immersed in a solution of K 2 Cr 2 O7. Reference. Rankin: Jour.*Phys. Chem , 11, 1 (1907). Note. Do not stopper the test tubes too tightly when exposing to the sunlight. Discussion. This experiment shows in a very satisfactory way how light dis- places the equilibrium: Q " C soluble insoluble. It also shows that the activity of light differs for different wave lengths. In a certain sense "light is a mixture of reagents." Light of a particular wave length is active if it is absorbed, and absorbed light tends to shift the equilibrium in such a way as to favor the production of the substance which absorbs the particular light less readily. 85