A METHOD FOB THE DIRECT VOLDHBHUC DETSQCDUTICW OF SOIFATE
DISSBRTinCE Presented In Partial Fulfil Tjseait of the Requirements for the Degree Doctor of RzllosopfcQr In the Obradaate School of The Ohio State University
hr
ROBHtT LUMKUiTH 3TCFFER, A«B* The Ohio State University 19SU
Approved by:
Adviser & AcnroiaEDcaan To &r« Bari* R. Caley the author expresses his sincere appreciatlcu for the suggestion u d g n U u e i of tfals Investigation sad for Inspiration which will servo throughout his career in cheats try. 11
TABLE CF COBTBTTS
Page Introduction • •••••••.••«•...... 1 Apparatus and Ch—lcals . • * 1* Preliminary Experiments *•••••.•...... $ Systematic Experiments ...... 10 Discussion of Data ...... • . . 19 Procedure •*.••••••••...... UO Notes on the Procedure ...... la R e s u l t s ...... 1*2 Interferences 1*2 Applications ...... 1*8 Sonmary •••.••••••••••.•••• 1*9 Bibliography ...... $1 Autobiography 5k 1 A HBTBGD FOR THE D3HSCT TOUOSSRIC DBTffiHIHATSCV GT SUIAXE
The difficulty of a given deteml nation may often be judged by the amount af literature on the ntjeot, for there seons to be a direct relation between the tec* That the Tdlawtrle determination of aolfate boa been a difficult problem eeene evident fTon the large mfcer of papers that have been published on this subject, and fron the nidaly varying conclusions reached by the different Investigators as to the procedure to employ* The best procedure to use In the gravlnstrlc deterulnatlon of sulfate in itself a highly controversial subject* Kolthoff (15) has pointed out that the difficulties of coprecipitation and solubility mfcdoh, beset gravinstric Methods are likely to cause even greater disturbances In volumetric work* To these are added difficulties in the observation of the end point* The vclanetrlo m e t h o d s recorded in the literature for the datermination of sulfate nay be roughly divided into tuo classess first, indirect nethods} and second, direct Methods* In the category of Indirect nethods are placed all procedures uhlch involve f i l t r a tion* or back titrations* Of the Indirect nethods uhlch Involve nitrations, Mohr* a (20) method appeared first* It consisted of adding an excess of standard barium solution to the sulfate sample, con verting the excess of barium to carbonate, filtering off the nixed precipitate, and titrating the barium carbonate in this precipitate uith acid* Clomi (7) modified thin procedure by titrating the excess of carbonate in the filtrate * 2 Chromate « 8 first used by Sehsers (51) la a procedure involTii* filtration* Be added an «xcooa of lead nitrate to hie enlfate samples, filtered off the lead sulfate, and dstexwdnad the eaae-eee of lead In the filtrate by titrating It with a potassium chroan te eolation using Oli ver nitrate as an external indicator* Andrews (1) determined the excess of chr caste fay precipitating It as barium chromate, fllterli^ off the precipitate, and titrating the chromate In the precipitate lodcnetrically after dissolving it In add* The ben si dine method Is probably the best known of the filtration methods* Ben si dine was first need fay Muller (22) • After filtering off the ben si dine sulfate precipitate, he titrated the excess of ben- sldlno hjrdrochlaride In the filtrate with standard base* Soon after, Ranchig (26) introduced M s method of titrating the banrldlnn sulfate precipitate Itself; and there aroee a keen rivalry between theee two men as to whose method was the best* Phosphate (17) and oxalate (28) salts have also been used in somewhat similar filtration procedures* The main disadvantage of all the procedures mentioned thus far is the fact that they require a filtration, which necessarily makes the methods rather long* Indirect methods which do not require a filtration are usually based upon back titration* An excess of a standard reagent is added to precipitate the sulfate, and then the excess of this reagent is titrated with a second standard reagent* WLldenstein (hX) was first to introtksce a method involving back titration* »ia method depended upon the addition of an excess of barium chloride; the excess of barium was determined by titrating with 3 chronste solution until the first f e w drops of an areas of chrosnte colored the supernatant liquid yellow* Other Investigators have used carbonate (12), stearate (6), palxdtate (3), sulfate (3$, 38), sal ethylseedljMdae tetraacetate (23) to titrate the exoeea of barlusn*
Lead salts were first ased 1a a back titration Method by Oddo sad Beretta (25) . They precipitated the sulfate by adding an excess of lead nitrate and then back titrated the excess of lead with ehronate airing silver cfcrcsnte as an indicator* Potaasinw fsrricyanide (21) has also been used to titrate the excess of lead. In general, however, lead sulfate is too soluble to sake these Methods readily applicable
1a dilate solutions. The naln disadvantage of procedures that require a back titration is the fbct that at least tm> standard solutions are needed* Further- acre those Methods which depend upon bydroilyvU - those that use carbonate, stearate, palsdtate, and in scan cases chronate - are subject to Interference frcei the presence of salts that exert a buffering action. Direct titration Methods usually have a considerable advantage over indirect Methods in speed, and often in simplicity of the opera tion. The first direct vduMStrlc proceAiro for sulfate was worked out by Oay-Lussac (13) • Be detemlned the end point of the titration of a sulfate solution with barium* chloride by withdrawing a swell portion of the supernatant liquid and adding a anall anount of the barlAM chloride solution to see if any turbidity was produced. Pro cedures based upon this Method are very tedious and tine consualng. Level (16) was the first to use load nitrate as the titrant in a h direct determination, Fotutlni Iodide nee meed as the Indicator with yellow lead Iodide being famed at the end point. Here again, aa with all methods using lead salts ms titraxits, the solubility of lead sulfate is the main disadvantage. Adsorption Indicators (U, UO) have been tried to determine the point of a direct titration of sulfate, they have not proved to be very successful because of interferences and the difficulty in observ ing the end point. la 1933 Schroeder (30) introduced tmtrmhydr ***** ms mn internal indicator. This indicator has had rather m wide acceptance although It has been criticised by some who have had trouble in observing the color change. Various instrumental methods based on canhactonetrlc (10), refrac tome trie (5), t h e m urns trie (8), potentlometrlc (27, 33), or amperometric (18) measurements have been proposed for the direct determination of sulfate* Host of these methods require special laboratory apparatus, and m a n y of them have various limitations, Mazy modifications of the procedures mentioned in this brief survey have appeared in the literature, and comprehensive lists of references are given by Kuril sale (29), Schroeder (30), Thomson (37), and Siegfrledt et ml. (32). The purpose of this investigation was to develop a direct volu metric procedure with a readily detectable end point for the deter mination of sulfate.
ATOKAICS AMD GHBCDCAI8 All weighings were made on an Ainsworth magnetically damped chalnoswtlc balance using calibrated stainless steel weights* Dryings and Ignitions sere accomplished in a De Khotinsky type constant temperature oren and a Heri Duty type woffle furnace, re spectively* A Beckman Modal O pH Meter equipped with, a General Purpose glass electrode and a ground glass sleeve type ral own! reference electrode was used* The stirrer used vas a Cenco variable speed stirrer. Buffer solutions vere prepared from potassluu a d d phthalate and borax as described by Dole (9)* However, Dole’s use of a saturated solution of sodium broad.de to maintain borax at a constant weight In a hygrostat has been proved unsatisfactory by Mensel (19), who recost- mended a solution saturated with respect to sucrose and sodium chlo ride* The author found this solution also to be unsatisfactory and used instead a saturated solution of sodium sulfate deeahydrate which successfully leapt the borax at constant weight and in proper propor tions with respect to sodium borate and water* The percent of sodium borate in the borax was checked with standard hydrochloric acid. All chemicals used were reagent grade except where otherwise indicated, and where necessary were dried according to reccumended conditions*
PREL1MXHAKT EXPERIMENTS Prom a consideration of the various possibilities, it seemed likely that a titration based on a pH change at the end point of a titration of sulfate with standard barium solution offered the best possibility for the development of a new and satisfactory method. 6
It h m first hoped that a method could be devised that would wake possible the wee of ordinary acid-base indicators in soch a titration* However, this hope was not realised, as It was soon seen that the pH change in all the nethods tried was neither sharp enough nor extensive enough to permit the use of such indicators* Therefore, all changes of pH at end points ware observed by weans of a glass electrode as the indicator. Tbs possibility of determining sulfate by the use of carbonate as an Indicating agent was first studied, a swall amount of carbonate being added to the solution in which was sulfate to be determined. Because carbonate ions hydrolyse and tie up hydrogen Ians, this solution is initially alkaline. If carbonate ions are then withdrawn from the solution by barlun ions at the end point, a drop in pH occurs as hydrogen ions are released. Due to the great difference in solu bility between barlow sulfate and barium carbonate, virtually no barium oarbonate precipitates from a sulfate solution containing a snail amount of carbonate until practically all of the sulfate has been pre cipitated by the addition of barium chlorideto the solution, and the amount of sulfate remaining In solution is insignificant. This precipitation of barium carbonate should then cause a narked change in pH. The theoretical pH of a solution at various stages daring the process just described can be calculated frcn the solubility products of barium sulfate and barium carbonate, and the ionisation constants of carbonic acid. The pH values calculated for the titration of 50 ml. sajqilee of 0.075 M potassium sulfate containing amounts of 0.1 M 7 sodium carbonate varying from 0.0J> ml. to 0.5 al* with 0*1 M barium chloride are shown In Table I and illustrated In Figure 1, When these titrations sera performed using the same amounts and concentrations as specified In the calculations, results sere obtained d d c b corresponded vary closely to those Which mare calculated. The pB change in these titrations did not occur as rapidly as with an acid-base titration, and the pH of the solution would drift with time after an addition of barium chloride had been made. Tbs addition of ethyl alcohol to the sulfate solution prior to the titration Increased the extent of the change in pH to a ma r l m of approximately three pH units. However, this pH change took place relatively slowly, and drifting of the pH of the solution still occurred. Chloroform, agar, gelatin, starch, and nltrobensene were added to the sulfate solution with the hope of improving the extent of the pH change, but none of these substances proved helpful. The sub stitution of dibasic potassium phosphate or potassium chromate for sodium carbonate as the indicating agent gave a slightly sharper drop in the pH. It is not surprising that the addition of alcohol improves the pH change, m an alcoholic medium the ionisation constant of the acid corresponding to the anion of the indicating agent is snch smaller than in water, and the solubility product of barium sulfate Is de creased. Both of these factors tend to make the change in pH more pronounced at the equivalence point. However, the presence of alcohol produces an effect, which will be described later, that probably greatly overshadows either cf these two factors. CABLE I
Theoretical pH Change Daring the Titration of 50 XL* of 0*075 X with 0*1 M Bad2 0*1 M BagOO^ as the Indicating Agent
______a ______Ba& 2 Added 0*05 ML* of BagOOj 0*1 Ml. of BajOO, 0*2 XL. of B a ^ 0.5 XL. of B a ^ Ml. (Approx* 1 Drop) (Approx* 2 Drope) (Approx. It Trope; (Approx* 10 Drops)
0.0 9.85 10.06 10*28 10*53 36.5 9.66 9.90 10*12 10*39 37*0 9.66 9.90 10*12 10.38 37.5 9.66 9.90 10*12 10.38 37.6 9.66 9*82 10.01 10*33 37.7 9*51 9.61 9.82 10.25 37.8 9.36 9*1*3 9.61 10.16 37.9 9.25 9.30 9JU3 10.00. 38.0 9.15 9.20 9.30 9.82 38.1 9.08 9.11 9.20 9.61 38.2 - - 9.11 9Ji3 38.3 8.95 8*98 9.0b 9.30 38.lt --- 9.20 38.5 8.86 8.87 8.93 9 0 1
00 C 60r
9801
pH
<$4 0 — Theoretical pH change during the titration of 50ml of 0 0 7 5 M K2S04 with O' M BaCl2 using Oi M Na2C0 as the indicating ogent
9 .0 0 0 05 mL of No2C0
0 2 ml of Na 2C03 0 5 ml of Na?C0*
8 201— ------l ______|______|______1 3^ 370 375 380 38.5
Ml of Oi M BoGi2 Added
FlGJRE l 10 Slnoa the drifting of the pH of the eolation eaggeeted that equilibrium m e not being established In the eolation, it m e found n e c e s s a r y to keep the earn rate of addition of the barlna chloride titrant for all titrations in order to obtain reprodaolble titration curves* aismunc mwuxBtis A wyet i— tli set of eaperlaaats using 25 si* of 0.15 M potassium sulfate, to ^rich use added various aaouots of 9$% ethyl alcohol and mter to produce the desired nedlnm, ae the sanpies and 0*1 >1 barium chlarlda ae the titrant m s then carried out to study the effect of different medians and of the different Indicating agents an the pH change when the following procedure was used* With a moderately high constant rate of stirring, the barlna chloride titrant m e added at a rate of 5 al> per minute to within several ml* of the equivalence point* This rate corresponds to a rather fast drop rate* For each reading of the pH, the stopcock was closed, and one art mb s was allowed to elapse before the reading m s mde* The titration was then con- timed as before until the next reading m e mde* This stoppage of the titrant allowed the solution to cone closer to a state of eqoi- llbrlna* When the equivalence point was only several ml* away, the rate of addition was slowed to U drops per minute with a reading made after each alnrte* Topical titration data obtained using this procedure are shown In Tables n and H I and are illustrated graphically In Figures 2 and 3* Figures U - 6 show sons of the other typical titration curves obtained* Table 17 sumarlses the results of the experiments to determine the 11
TABU n Titration of 25 KL. of 0.15 H with 0.1 H BaCLg
25 Ml. of Wrter Added 25 ML* of Sthaaol Added to the Hanoi e to the Senle BaCl2 Added BaCl^ Added Ml* PH Ml. pH
0.00 6.63 0.00 7.10 10.00 8.73 10.00 9.53 20.00 9.23 20*00 9.80 25*00 9 J*H 25.00 9.87 30.00 9*60 30.00 9.91* 3l*.00 9.67 31* .00 10.00 35.00 9.68 35.00 10.01 3509 9.69 35.18 10.05 35*39 9.69 35*33 10.07 35.59 9.69 35*50 10.06 35.79 9*66 35.67 10.07 35.99 9*66 35.83 10.02 3 6 0 9 9*63 35.99 9.98 3 6 0 8 9*62 36.16 9.93 36.58 9*61 36.32 9*86 36.78 9*59 36.50 9.81 36.98 9*58 36.68 9.77 3 7 0 7 9.57 36.87 9.69 37.37 9.56 37.01* 9.62 37*57 9.51* 37.22 9.55 37.76 9.53 37*39 9*1*7 37*95 9.52 37.57 9.36 38.1U 9.50 37.73 9.27 38.31* 9 .1*8 37.90 9.13 38.5U 9.H7 38.09 9.02 38.73 9J*5 38.27 8.87 38.93 9*>*3 38*1*1* 8.71 39.12 9.1*1 38.61 8.52 39.33 9.39 38.79 8.13 39.53 9.37 38.99 7.67 39.72 9 0 6 39.16 7.20 39.92 9.33 39.31* 6.88 1*0.12 9.33 39.53 6.71* 39.70 6.67 39*88 6.60 1*0.06 6*53 12
TABLE ZZX Titration at 25 ML. of 0*15 M X^SO^ vlth 0.1 M B a d 2 Oclng 2 Stop* of 0*1 If 1 2 ^ ^ a* the TiwH^frtwg Agent
25 Ml* of Utter Added 25 KL. of BUaanol Added to tlM W— >1l to tbi flnplt B a d 2 Added B a d 2 Ml. pH Ml. PH
0.00 7.85 0.00 8.37 10.00 8.63 10.00 y.53 20.00 9.12 20.00 9.82 25.00 9.25 25.00 9.88 30.00 9 J a 30.00 9.96 3b .00 9.52 3b .00 10.01 35.00 9 £ 2 35.00 10.02 35*19 9 £3 35.19 10.06 35*39 9.52 35*37 10.06 35.58 35.5b 10.07 35.77 IS 35*70 10.02 35*97 9 M 35*86 9*98 36*16 9*b0 36.03 9.92 36.36 9*37 36.20 9.83 36*55 9.31 36.38 9.76 36.71 9 .29 36.55 9.69 36.J3 9.26 36.71 9.62 37.12 9*22 36.89 9.52 37.32 9.18 37.05 9.3b 37.51 9.16 37.21 9.22 37.70 9*13 37*39 8.73 37.90 9.07 37*56 7.97 38.09 9.0b 37.71 7*35 36.29 8.99 37*89 6.9b 38J.8 8.96 38.07 6.77 38.67 8.93 38.23 6.63 38.86 8.87 38 JiO 6*53 o — — -- 0- pH change daring the titration of 25 ml. pH 25 ml, of water added to the sample 25 mi of etnanoi added to the sample 0 10 20 30 40 Ml of 01 M BaCl2 Added fig u re 2 pH change during the titration of 25 ml of 015 M K2504 wth 01 M 0aCl2 usmg 2 drops of 0 I M K2HP04 as the indicating agent ------25 ml of *ater added to the sample ------25ml of ethanoladded to the :ample ______L_ fO 20 30 Mi of 0 i M BaC 12 added FIGURE 3 o — -o- pH change during the titration of 25 ml of 0 15 M K2504 with 0 1 M BaClz using 5 drops of 0.1 M K2H P04 as the pH mdicoting agent 25 ml of water added to the sample 25 ml of ethanol added to the sample 30 Ml of 0 1 M BaCu added FIGURE 4 & - o — o pH change during the titration of 25 ml of 0 15 M K2S04 with 0 I M BaCI2 using 2 drops of 0 1 M NOgCOj PH as the indicating agent 25 ml of water added to the sample 25 ml of ethanol added to the sample 20 40 Ml of Oi M Boa, odded FIGURE 5 /y. — — —O -0^0 O— — -0 pH pH change during the titration of 25 mI of of 01 M K2O 0 4 as the indicating ogent 25 nnl of water added to the sample 2 5 mi of ethanol added to the somple 6 - 0 10 20 30 40 Ml of 0 1 M BaCl2 Added FIGURE 6 TABUS 17 Stadjr of Indicating Aganto and tha Effaet of Alcohol 4 pH M a x i m - In Varlooa Martin— d V 50 Ml* fetor 1(0 Ml. fetor 25 Ml* fetor 25 Ml. fetor Indicating Agont 0.0 Ml* 955 Bthanol 10 Ml* 955 Sthanol 25 Ml. 955 Ethanol 50 Ml* 955 Kfea— 1 Uood (0.05 Alcohol) (205 Alcohol) (505 Alcohol) (675 Alcohol) Mona 0.1 0.3 2.2 2.5 Two Dropo of 0.1 M fe^OOj 0.1 0.3 2.6 2.3 Plva Eropa of 0.1 M 803003 o a 1.3 1.9 2.2 T— Dropo of 0*1 H I^HPO^ 0*2 2.3 3.1* 2.2 Fire Dropo of 0.1 M £2® ^ 0.6 1.7 2*6 1.9 Two Dropo of 0,1 M KjCrO^ 0*1 2.3 2.9 2j| FIt o Dropo of 0*1 M KjCrO^ 0.3 2.0 2.9 2*6 19 affbct of the various alcoholic mediums and of tha different 4w*««»^ne agents* Fran this study It ues found that tha eptlnan amount of alcohol to ba aaad to give tha boat pH change aaa an amount equal to tha folaaa of tha original aqueous a dint Ion and that tha agant which gave tha boat pH ohanga naa 2 dropo of 0*1 H dibasic potaaalan phosphate* D330P3810H pH change during the titration of 25 ml of 25 ml of ethanol added to the sample 0 10 20 3 0 Ml of O i M BaCl2 Added FIGURE 7 pH change during the titration of 25 ml. of 0 0 2 M k2S04 #ith 0 0 2 M BaCl2 25 ml of ethanol added to the sample C 10 20 30 Mi. of 0 02 M BoCl2 Added FIGURE 8 pH change during the titration of 25 ml of 002 M 25 ml of ethanol added to the sample 6 - Mi of O.i M BoCU Added FIGURE 9 pH change during the formation of BaS04 25 ml of O.l M BaCl2 titrated with 01 M K2SO4 titrated *ith 01 M BaCl 25 ml. of ethanol added to each sample 6 - 20 Mi of Tit rant Added FK3JRE 10 pH change during the titration of 25 ml. of 0.2 M NalO, with 0.1 M BaCL 25 ml. of water added to the sample 25 ml. of ethanol added to the sample 0 10 20 30 Ml. of O.l M BaCI2 Added FIGURE II pH change during the titration of 25 mi of 0 I M BaC^ with 02 M NaI03 25 mi of water added to the sample 25 ml of ethanol added to the sample ± 0 10 20 30 Ml of 02 M NaI03 Added FIGURE 12 pH change during the titration of 25 ml. of 0.1 M K2S04 with 0.1 M P b(N 03)2 25 mL of water added to the sample 25 mL of ethanol added to the sample o- N x -o 0 10 20 30 Ml of 01 M Pb(N0.)3 2 Added FIGURE 13 pH change djnnq the titration of 25 ml of 01 M Pb(N03)2 with 0 1 M K2S04 — 25 ml of water added to the sample — 25 ml of ethanol added to the sample / 4 J i ______1__ L C 20 30 V; of C i M * 2S04 Added F'GJRE '4 pH change during the htration of 25 ml. of 0 2 M NaI03 with 0 1 M Pb (N03)2 25 ml. of woter added to tne sample 25 mi. of ethanol added to the sample >C 20 30 V . of C ■ M Pb;N0i.,2 Added F G jr E 5 c pH change during the titration of 25 rnl of 0. I V Pb(N03)2 with 0.2 M No 10 3 25 rn‘ of water odded to the sample 25 ml. of ethanol added to the sample 0 10 20 30 Ml of 0 2 M N al03 Added FIGURE 16 —o pH 25 ml of water added to the sample 25 ml of ethanol added to the sample 0 10 20 30 Ml. of 0.1 M AqN03 Added FIGURE 17 pH chonge during the titration of 25 ml of O.l M AgN03 with 0.1 M N a I 0 3 25 ml. of water added to the sample 25 ml. of ethanol added to the sample - o pH — O e, 0 10 20 30 Ml of O.l M NoI03 Added F IGlJRF 18 pH chonge during the titration of 25 ml. of 0 1 M KCI with 0.1 M A gN 03 25 ml of water added to the sample 8 25 ml. of ethanol added to the sample ------a - 'O-O-O- __ -o- ± 1 0 10 20 30 Ml of 0 I M AgNOj Added FIGURE 19 £r pH change during the titration of 25 mL of 0.1 M AgN03 with 0.1 M KCI 25 ml. of water added to the sample 25 ml. of ethanol added to the sample - o 10 20 30 Ml. of 0.1 M KCI Added FIGURE 20 36 Adsorption depends upon the specific nsture of the precipitate as well as vpon the mediae and the rate o f id Ting* Silver precipitates particularly do not seen to be affected by hydrolytic adsorption. The titration curve of the pH change when iodate Is precipitated free 50% ethanol by the addition of lead nitrate as shown by figure 15 seests to be unusual, lbs drop in pH during the first part of the titration nay be caused by the addition of the acidic lead nitrate and by the formation of basic lead salts. In an aqueous solution lead hydroxide or basic lead salts begin to precipitate at about a pH of 7, and In an ethanol solution they probably begin to precipitate at a lower pH, The abrupt rise in pH just before the equivalence point occurs at the sane tine that a slower rate of addition of the lead is begun, A slower rate of addition allows the solution to cone closer to a state of equilibria*, end if the basic lead salts are unstable in this solution, their dissolution causes the pH to rise. When this titration was carried out, there was a gradual drifting of the pH upward with tine after a snail portion of lead nitrate had been added. This upward drift shews that eqoillbrlun had not been established and that scene process was occurring that diminished the hydrogen Ion concentration. After the equivalence point the pH drop is no dcxibt caused by the addition of the acidic titrant. In the development of a procedure for the deteruinatlon of sulfate based upon the pH change doe to hydrolytic adsorption, the use of acid-base indicators proved unsatisfactory because the pH drops too slowly near the equivalence point. The slope of the titration curve near the equivalence point is almost constant. This fact prevents the 37 end point from being determined an the point of maximum slop#, f oar the slope nay bo at a warlnmi value over a relatively largo range of the titration c u m . When a uniform rate of addition of reagent la used, titration curves of the pfl change during the preolpltatlon of barium sulfate can be reproduced ulth some degree of accuracy* This reproducibility suggested that a definite pH value on the titration curve near tbs equivalence point night be selected as the end point* What pH value to select Is rather arbitrary although a value on the curve where the slope Is at the marl mi provides the sharpest end point* Although the results su— xiaed in Table 17 indicate that 2 drops of 0_1 H dibasic potassium phosphate gave the best pH change, the precision obtained when this Indicating agent mas used to standardise barium chloride against potassium sulfate was not as good as with some of the other Indicating agents* Table V shows the precision obtained when various indicating agents were used* The standardisation of the barium chloride was carried out according to the following procedure* The potassium sulfate was dissolved in 2f> ml* of water, 2 $ ml* of alcohol was added, and then the designated amount of Indicating agent was added* Using a moderately high rate of stirring, the barium chloride was added at a rate of f> ml* per minute* Hear the equivalence point, the pH begins to drop rather suddenly* At this point the rate of addition was slowed to U drops per minute. The pH values chosen for the end points for the various indicating agents were 7*70, 7*80, 7*80, 7*80, and 7*90 for 2 drops of 0*1 M dibasic potassium phosphate, $ drops of 0*1 H dibasic potassium phosphate, 2 drops of 0*1 M potassium 36 o I o• o • o • o • o* o• o • o • o • o• o • o • o • I ► ► 5 a £-co >o r— m i 8 u p • • i • ♦ 't'l' * • • * • o o o o o o o o o o o o o o ♦ . 0.0971 ► ► t A o vQOJO Vf\ r— w f cf— O N - N VO vO H t- r-vo vO'O'Op B S N o• o • o • o • o o o o o• o • o * o • f f m MCM X X # O 0 \5 * §• §• §• h a £ e I fi 1 TABUS 7 (CO01BDED) Potaeain* Sulfate Indicating Agent Taken Ml, of Ba&2 Molarity of Deflation Uaed g. Soln, Repaired B4C12 Soln, frc® the Mean Five ETope of 0 J M K2Cr0U 0.6712 39JiO 0.0978 0.0000. 0*6?2lt 39 j e 0.0980 0.0003 O.6767 39.87 0.0971 0.0003 0.6793 h0.01 0.097li 0.0003 At . 0.0977 At . 0.0003 Two Drope of 0,1 M 5a2CO3 O.670I1 15.58 O.O6U1 0.0006 0.6733 h S M 0.0660 0.0002 O M r t Ii5.26 0.0866 0.000b 0,6775 1*5.31. 0.0668 0.0006 At . 0.0662 At . 0.0006 $ Uo chromate, 5 drops of 0*1 M potassium chromate, and 2 drops of 0*1 M sodium carbonate, respectively* From this tsblo It esn bo soon that the use of 5 drops of 0*1 H dibasic potassium phosphate as the Indicating agent gars the best results* If 0*1 M barium chloride is used as a tltrant, rather large sulfate samples must be used In order to use a convenient volume of tltrant* A titrant whose concentration is 0*02 H with respect to barium chloride is tery suitable far samples containing approximately 75 mg* of sulfate, and the procedure below is based on a tltrant of this concentration* Since the pH value chosen for the end point may not occur at the equivalence point, it is advantageous to standardise the barium chloride empirically using the same procedure as that used to deter mine sulfate in unknown samples* bpirlcal standardisation nay also eliminate other inherent errors in the procedure* The barium chloride is standardised against dried, reagent grade potassium sulfate* PROCKDPRB From the results obtained in the foregoing experiments the following prooedure is recommended* A sample containing about 75 milligrams of sulfate is dissolved in 25 ml* of water in a 250 ml* beaker* If the amount of potassium in the saaple Is not equivalent to the sulfate present, 200 mg* of potassium chloride is added and dissolved* To this solution are added 25 ml* of 9$% ethyl alcohol and 5 drops of 0*1 M dibasic potas sium phosphate* This solution is then placed in position so that the Ul tips of the electrodes sod the stirrer ere 1— eised. The pH of the solution Is adjusted if necessary to a value between 8*90 m H 9*00 with 0*02 V potassium hydroxide or 0*02 X hydrochloric acid* With a moderately high rate of stirring* the 0.02 M barium chloride tltrant Is added at a rate of 5 ul* per ulxaite until the pH begins to drop rapidly* This rapid drop In the pH usually begins at a point several si. before the end point* At this point the rate of addition is slowed to one drop per 15 seconds until the pH of the end point* 7.70, Is reached* Standardisation of the barltm chloride tltrant is carried out in the sane way* H a a s OH THE HtOCKDCKE The volume of water used to dissolve the sample and the volume of alcohol added need be measured no more accurately than Is possible with a graduated cylinder* If the pH of the solution is quite distant trcm the required range of 8*90 to 9*00 a more concentrated solution of potassium tydroxide or hydrochloric acid than that recoomendad in the procedure should be used to bring the pH of the solution within this approximate range* The final pH adjustment can then be made with the more dilate solution* The use of concentrated solutions is necessary in some cases in order to keep the final volume as close to 50 ml* as possible. From the beginning of a titration the pH gradually falls until the titration is well over half completed* Before the pH begins to drop rapidly near the end point* there is quite some tine that the pH ream Ins constant* or it may even rise slightly during this time in U2 sens euec, of this baha-rior, the beginning of the rapid fall In pH ean t u U y bo detected, and consequently the point at which to slow the addition of the barium chloride In fairly clearly defined* The tine neoeeaary to c deplete a single titration is usually 15 to 20 minutes* RBSPU8 Results obtained by using the above procedure are susnarised in Tables VI and VU* High results obtained with the snail ar sulfate sanplee nay be ex plained by the fact that with snail or amounts of sulfate the total amount of barton sulfate present is lees, and therefore the anount of adsorption 1s decreased. As the anoent of adsorption is decreased, the change in pH bee ones slower, and a longer tine is needed before the pH of the end point is reached* Since the anount of barlun chloride added is directly related to the tine, nore barlun chloride is used than should be* jgEHrouncgs A systenatic study *as conducted to detemine the interference of sense of the nore ecnann ions that night be present nlth sulfate* To £>*00 ml. sanplee of a known sulfate solution were added increasing anoonts of the substance whose interference was to be studied. The determination of the sulfate was then carried out in the usual way* Frctn the results listed In Table vill it can readily be seen that except for potassium, sodium, and chloride the other ions tested produce errors of one mg* of sulfate or nore when present in only 1*3 TABLE VI Detcndjmtlon of Sulfate in Potaoaiun Sulfate Saagd.ee Sulfate Sulfate Error Taken Pound ■g. Mg. Mg. of Sulfate Pte,/1000 91.7 9U.5 0.2 -2 91.7 9U.7 0.0 to 85.7 86*0 0.3 +1* 82* Ji 8UJi 0.0 to 69.2 69 Ji 0.2 +3 66.6 66 JU 0.2 -3 5U.8 55.5 0.7 +13 51.9 52.7 0.8 +15 36 Ji 36.9 0.5 +U* 35.7 36 0.7 + 20 TABLE V U Determination of Sulfate in Sodlua Sulfate Samples* Sulfate Sulfate Taken Found Error ng. Mg. Mg. of Sulfate Pte./I000 89.6 90.0 Oj* + i* 89 .1* 90.3 0.7 + 8 7U.0 71* .5 0.5 7 73.8 71* .3 0.5 7 55.8 56.6 0.8 + li* 51* .8 55.7 0.9 <16 ♦200 eg. of KCL added h h u b i £ v m Study of Interference* In the Petered nation of Sulfate in Potwi— Sulfate 8— plot Sulfate Sulfate Error Compound Takan Found Pfiwnt ng. ng. Mg. of Sulfate Pte^lOOO 10 ng. Ed 72.2 72.1 0.1 -1 10 ng. KCL 72*2 72.2 0.0 i-0 20 ng. KCI 72.2 72.3 0.1 +i 20 ng. SCI 72.2 72.3 0.1 +i 50 ag. K d 72.2 72.6 0.U + 6 50 Mg. KCL 72.1 72.2 oa +1 5o ng. K d 7 2 A 72.3 0.2 + 3 ioo m * K d 72*1 72.5 0.U +6 loo m * K d 72a 72.6 0.5 + 7 200 ng. K d 7 2 a 72.5 OJi 6 200 ms* K d 72.1 72.1 0.0 -0 200 ng. Kd 72.1 72.ii 0.3 k 200 ng. Kd 72.1 72.6 0.5 7 500 K d 72.1 72.fi 0.7 10 500 Mg. K d 72.1 73.0 0.9 12 10 H a d 72.1 72.H 0.3 U 10 ng. Had 72.1 72.5 0.U 6 20 ms* H a d 72J. 72.7 0.6 8 20 ms* H a d 7 2 ^ 72.ii 0.3 U 50 ng. B a d 72.1 72.2 oa 1 50 ms* H a d 72.1 72 a 0.0 -0 100 ws* H a d 72 A 72 *k 0.3 h 100 ms* H a d 72.1 72.2 oa 1 200 ms* H a d 7 2 A 72.2 oa 1 200 ng. Ha d 72.1 72.6 0.5 7 200 ng. Had 72.6 72.9 oa U 500 H a d 72.6 72.8 0.2 3 500 ng. Had 72.6 72.7 oa 1 1000 ng. H a d 73.2 7 3 ^ 0.1 -1 1000 w** H a d 73.2 73 *U 0.2 3 U5 TABLE T U I (CONTIIUED) Sulfat* Sulfate & *ro r uoafioaDa T u n i rgona Present ag* a g . Mg. o f Solfata pt«.Aooo 5 a g . MO 3 7 2 .6 7 3 .2 0 .6 + 8 5 a g . MUK)3 7 2 .6 7 3 .1 0 .5 + 7 10 Mg. BMBO3 7 2 .6 73 J t 0 .8 +11 10 Mg. MO 3 7 2 .6 7 3 .6 1.0 +U i 20 MO 3 7 2 .6 7 3 .6 1 .0 + Iii 20 Mg. RaK03 7 2 .6 7 3 .3 0 .7 + 10 50 a g . MaH0 3 7 2 .6 7 U a 1.5 + 21 50 n g . HaB03 7 2 .6 714.0 lJl + 19 100 n g . n*xoi 7 2 .6 7U .1 1 .5 + 21 100 Mg. *a*o3 7 2 .6 7 3 .9 1 .3 + 18 10 K2HPOli 72*8 7 7 .2 UJ* 60 10 ng. K2HP0li 72.8 77.3 u o 59 10 a g . » 3 P 0 U 7 2 .8 7 8 .9 6 .1 + 8h 10 Mg. B^FO^ 7 2.8 7 8 .7 5 .9 + 81 1*6 relatively mall nuutte. With 10 eg. of either amwrsilum chloride, magnesium chloride, or pilclw chloride present In the sample, the titration carve wee changed eo n e b that no quantitative studies could be made on their Interferences* It is not surprising that ma^y Ions do Interfere with this proce dure because the procedure itself is based upon a rather snail pH change* Any ions that exhibit a buffering action or are thane elves coprecipitated produce the greatest anount of interference* Various Methods were used in an attempt to remove calcium and magnesium from sulfate samples prior to determining sulfate by the proposed procedure* The method that proved to be most successful removed the calcium and magnesium as the carbonates* The separation was carried out In the following way* A potassium sulfate sample containing approximately 600 mg* of sulfate and UOO mg* each of mag nesium and calcium as the chloride is dissolved In $0 ml* of water* To this solution is added li grams of sodium carbonate, and the solu tion is thoroughly stirred* The solution is then heated to just short of boiling and digested for one hour on a hot plate regulated to main tain a t saiperature of 90° C* After the digestion is completed, the supernatant liquid is filtered thrcmgh a Schleicher and School! white ribbon filter paper, and the precipitate 1s washed four times by de- camtation using a hot 0.2% solution of sodium carbonate as the wash liquid. The precipitate is then transferred to the filter paper, and the washing is continued until l£0 ml* of filtrate has been collected* The filtrate is then neutralised with concentrated hydrochloric acid using methyl red as an indicator* To drive off the eatress carbon U7 dioxide, the eolation is boiled for two minutes, sad nore hydrochloric sold is added if necessary. After the solution is cooled to room turn perature, it is treasferred to s 200 ml. volumetric flask sad diluted to the nark. TwenfcjMfive ml. aliquot samples sre then withdrawn, sad the sulfate deterwined in the proposed wanner. The results of this procedure la rsnoring the interference of calcium sad magnesium sre shown in Table IX* TABUS IX Removal of Calcium sad Magnesium Interference Mg Present Ca Present Sulfate Sulfate as HgCl as CaCl Taken Found Brror mg* mg. mg. mg. mg. 50 SO 72.7 73.7 1.0 50 50 72.7 73.6 1.1 50 50 72.8 7U.1 1.3 50 50 72.8 7U.1 1.3 50 50 72.8 73.8 1.0 50 50 72.8 73 .U 0.6 Ion exchange was tried as a method to remove calcium and magne sium* Dowex 50, Amberlite IR-100, and Amberlite IRCJ-50 were investi gated as possibilities for truus removal. None of these resins gave a successful separation under the experimental c udltlons used* Since the final volume of effluent had to be restricted in order to obtain a convenient concentration of sulfate, the amount of water used to wash all of the sulfate from the column bad to be limited. This limitation on the amount of wash water also limited the else of the column. A modification of a method used by Noyes (2U) in a qualitative scheme for the earth metals was also tried to remove calcium ae and magnesium* This method consisted of precipitating the carbonates of calcium and In a very ibbctiIb u s I solution* CasqxLete removal of calc Ion and w g w l m ana accomplished, bat attempts to destroy the n— ntil nm Iona whose buffering notion would subsequently interafere proved to bo t.1 we consuming and unsuccessful* Attempts to remove calcium and magnesium froai anlfata samples aa oxalates In an acetic acid nedixni according to a Modification of a procedure recommended by Siring and Calay (U) for the deteralnatlon of magnesium resulted In aolvtlona which did not give the characterlatlc pH change when the sulfate waa titrated In the recoamended manner . A F P U C A T i a B The proposed procedure waa applied to the determination of the aulfate content in mixtures of aulftrric and hydrochloric acids * Samples were made by adding increasing anonnto of hydrochloric acid to a known amount of a standard sulfuric acid solution* The sulfuric acid solution had been standardised gravimetrically as barium sulfate, and tbe standardisation checked volume trioally against sodium carbon* ate* These synthetic samples were neutralised with 1 H potassium hydroxide using methyl red as an Indicator* Enough water was added to each sample to bring the final volume up to 2$ ml*, and the sulfate was then determined in the usual way* Results of this study are shown In Table Z where the aulfate content is expressed as sulfuric acid. It can be seen from Table Z that even without the presence of any hydro* chloric acid, there la a constant positive error of 0*8 mg* of sulfuric acid* If this constant error Is Incorporated into a titration blank h 9 which la subtracted from the Mount of sulfuric acid found in the mixtures, the results are somewhat Improved as Is shown by Table XI* s t m k a k t There are numerous examples in chemistry where a worker has wade use of a seemingly unfavorable occurrence* Coprecipitation has been utilized In several ways* For Instance, radiant has been extracted from pitchblende by eoprecipitatiag it on barium sulfate* Herein has been proposed a sulfate determination which also depends upon the coprecl- pitation on barium sulfate* The determination is relatively fast, and the and point is easily detected by means of a pH meter* Studies on Interferences Indicate that the method is best suited for solutions containing besides the sulfate ion only potassium, sodium, or chloride ions* 5o TABLE I Deteralnatlon of Sulfuric Acid la Mixture* of Sulfuric and ^fdrochloric Aclde B d HZSOU Hasoh Present I t l n Pound Erra ■g> a s* n g . Voae 78 Jt (0.16 H) 79.3 0.9 Hons 78 J t 79.2 0 .8 Hone 78 J t 79.2 0 .8 10 (0*023 H) 78 J t 79.3 0.9 10 78 .k 79.6 1.2 20 (0.056 H) 78 J t 80 wO 1 .6 20 78 J t 79.7 1.3 50 (0.11* H) 78 J t 7 9 .8 1 Ji 50 78 Jt 80*1 1.7 100 (0.26 H) 78 J t 81.5 3 a 100 78 J i 80 Ji 2.0 100 78 J t 8 0 .6 2.2 100 78 Jt 81.8 3 J i 200 (0.56 H) 78 J t End point never reached 200 78 J t Dad point never reached TABLE XI Detunwinatloo of Sulfuric Acid in Mixtures of Sulfuric and Efcndtrochloric Acids B d H2S0U H230r IVesent Taken Found Drrar n g . *g. eg. ■g. 10 (0.028 N) 78.)4 (0.16 H) 78.5 0.1 10 78. h 78.8 OJi 20 (0.056 H) 78 J4 79.2 0 .8 20 78 Ji 78.9 0.5 50 (O.lli H) 78 Jk 79.0 0 .6 50 78.U 79.3 0 .9 100 (0.28 H) 78 J i 80.7 2 .3 100 78 J* 79.6 1 .2 100 78 J4 79 J i lJ * 100 78.U 81.0 2 JS 200 (0.56 H) 78 J* Did point never reached 200 78.U End point never reached 51 BXKLXttmPHT 1* Aadrews, L. V,f "On * Vblwtrlc Hethod of Oeaeral Applicability for the Drternlnation of Coabined Sulphuric Acid," An, Che*. 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" ------5U n r r o B B X B A m I, Robert Ilewellym Staffer, n a born In Berth Georgetown, Ohio, Septeafcer 16, 1927* 1 received nr secondary school ehcttloe In the public schools of the city of Alliance, Ohio* lQr undergraduate train ing w e obtained at Meant Union College and Ashland College, and I received the degree Bachelor of Arts firm the latter school In 1950* Military service in the U* S* Arwy Intermpted mj undergraduate train ing for seventeen nonths during the years 19lt6-19li8 • In 1950 I entered the (feradnate School of The Ohio State University share I specialised In the Departnent. of Clwlntrj » Dazing the year 1950- 1951 I held the position of University Scholar; and daring the year 1952-1953, the position of University Fellow* In 1953 I received a Rational Science Foundation fellowship, and I held this position for one year while eoeqpleting the reqairweents for the degree Doctor of Philosophy*