Analysis of Highly Hydrated Dolomitic Lime Critical Comparison of Chemical and Instrumental Methods GEORGE L. CLARK AND ROBERT S. SPRAGI. E
A'oyes Chemical Laboratory, Cnirersity of Illinois, L-rbana, IlZ., and Coolidge Chemicul Laborutory. Harrard C-nirersity, Cambridge, M a s s .
The general13 accepted method of chemical analy sis of highly hydrated dolomitic lime, an important industrial material, containing oxides, h j droxides, carbonates, sulfates, etc., of calcium and magnesiirm iniolves t h e assumption of complete hjdration of calcium oxide to calcium hydroxide and conipntation of the amount of magnesium h5droxide from the residual amount of combined water dri\en off by heating. The purpose of this work is to test rritically the reliability and accuracg of this method through comparati\e anal) sis of mixtures of pure compounds and five commercial limes bl other available techniques-heats of solution, x-ray diffraction, differential thermal analyzer, Richardson thermal technique, thermal balance, gravimetric siicrose. and iodonietric titration procedures for calcium hydroxide. When uncombined (adsorbed) water is determined from loss in weight a t 280" instead of 120' C. now recommended, and other experinlentall) established precautions are observed, the results of these independent chemical and instrumental analjses support the validity of computing the per-
I
centages of calciuni and niagnesiuni hl-droxides from the content of combined water. The quantitatiie reliability and limitations of each method of analysis of these mixtures are established. I n spite of anomalous thermal behavior of magnesium hydroxide toward combined and uncombined water, as compared with calcium hydroxide, a satisfactorymethod correct to better t h a n 27" is derived. Excellent agreement for calcium hydroxide is found by the several techniques for four out of five of the commercial hydrates, impurities such as silica and sulfur trioxide accounting for deviations in the fifth sample. A s a result of these critical comparative experimental studies the long-standing difficulties and controversies attending the analysis of dolomitic limes and plasters should be largely resolved. Of prime importance, to the industry and to builders, is knowledge of how much magnesia remains unhydrated in highly hydrated, dolomitic limes used in finishing plasters; and whether it hydrates i n such plasters, with attendant possible volume change and plaster failure, or persists unchanged.
NTHE course of years of research on limes at the University of
Illinois and elscLvhere it has become increasingly evident that Table I. Chemical tnalJsis of Test Hydrates 0 1 1 %s-Kecei\ed Basis the analysis of t,he very familiar oxides, hydroxides, carbonates, 1 B C D E and sulfates of calcium and magnesium, especially in mixt,ures. -tlrates arc shon-n in Table 11.
Table 11. Composition of Test Hydrates (on ks-Heceited Basis) as Coniputed from Chemical 4nal) sis 4
oc C'lt(0H)~ h\Ig(OH)z CaCOa CaSOr
.5l 4
32.6 .I. J
MgO Unconihined H?O
0 5 3.8
0.2
B
C
D
E
c1
70
co
"*E
.52 1 36.4 5.2 0.5 3.2 0.4
49.6 38.4 7.5 0.7 3.3 0.3
51.4 40.2 3.6 2.7 1.3 0.5
51 2 39.1
5 0
0.8 3.3 0.3
If thr i ~ ~ n i ~ ~ o actually u n d s assumed present in the hydrate are in fact 1)resent a n d thc,rr : i i ~no others, the degree of unrertaint:, in t h c x c:ilculatc:d pewentag? of the various compounds is entirely tlepc~ridenton the d(:gi,ce of uncertainty of the various oside constituents :is oh1 :lined from t h r rhemical analysis. C'onsider :i hypothc,ticxl hj-tlratcs hitving the follon.i~igchemical iimlysis Rith the iiidiibatcil uncertainties (in absolute tcrnis) for t h r various v:ilucxs: c-lto, ?< 42.A z 0 1 MgO. c ; c 0 ? ,'I Total 1 0 ~ son ignition Uncotiihined H2O. 5%
ci
3 0 . 6 =t0 . 1 2.2 h O . 1 26.9 = 0 1 0 . 5 =t0 0.5
Sulfur ti,ioxitir is iiot )icing cwiisitlr~i~rd in this hypothetical c it is usually present to 1t.s than 0.5'%. The uncertaii ercentagrs art' based on the aut,hors' t of many samples of hjdrated doloniiti include not only the small uncvrtainty inherent in the chemical analysis itself but also the uncei,t:hty due to the fact t h a t the ua~i~ple actually :inalyzed is iiot iq)resentative of the entirr, s:inipIe liecausc. of ~ionhomogc~ieitj-, Tvhich is principally due t o reaction of calc.iuni hj-droside with atmospheric carbon dioxide, 1,rsulting iri a higher conc8r)ritr:ition of calcium cqi,bonnte a t the surface of the main m i i j than in t h e interior. The calculated 1ierwnt:igi. composition 1 d on the chemical amilysis ii-ith assoc-iatetl uncertainty val for the pcrcentages of the vayious t*ompountls is as i'ollon.~: as
5.0 32.5 37.0 5 0
3
0 2
I0 . 3
+ 0.9
+ 0.7
h s the cmnk)inetl water is customarily obtained bj- subtracting the sum of the unconit)ineti n-ater and carhon dioside from the total loss on ignition, a significant error in the determination of the uncombined water has a profound effect on the calculated perventages of niagneuiuni hj-tlroside and magnesium oxide. In th(, hypothetical case under consideration, to show \\-hat a systematic error of 0.2% absolute-i.e, if the uncombined water were actually 0.7% instead of 0.5%-\vill do to the computed composition, the per cent niagnesiuni hydroside becomes 36,4y0 and the per cent magnesium oxide becomes 5.4%, these figures deviating from the former values 1,s 2 and 8% (relative), respectively. It' is shown sulmquently in t,his paper that the present method of unronihined n-ater gives results as much ute) low: hence, the calculated values for in hydroxide anti per cent magnesium oxide in particular cannot be regarded as very reliable. The per cent calcium hydroxide (on the as-received basis) is, however, inde-
pendent of the water coiiteiit of thc sample aiid depends onlj- on the validity of the assumption that the only calcium compounds present are calcium sulfate, calcium carbonate, and calcium hytiroside, and that all the carbon diositle arid sulfur triosicle in the sample are combined with calcium. I n this investigation, great care was taken to minimizc the ,c,ffects of sample nonhomogenrity anti cmhon dioside contamination. The test hydrates \yere niised thoroughly and stored in 1-quart. screw-cap jars. Smaller hamples vere removed from these jars from time to time (after thor.ough mixing of the contents) for the analytical work antl these smaller portions (about 25 to 30 grams) were stored in small screw-cap jars which were in turn kept in a desiccator containing fused pot'assium hydroside. Despite all precautions, values for the various oxide constitursnts determined chemically were reproducible only to the esteiit 1)reviously indicated. This lacli of high preclixion in the chciiiical tleterminations must be bornc, in mind r\-hrxn sutisequrnt coniparisons are made between t h e 1)ert~~nt:ige of certain compounds determined direct]>-and thc percc~nt:tgc~as c,:ilculated froni the t~hemiralanalysis. I t has tieen proposed ( 1 6 , 17 j to c1ic~l.rthe idculated coii1l)os.ilion of hydrated dolomitic lime :E computed from the chemical analysis by measurement of thc hi,:it of iolutiori of the hydrate in I~iIi,:+td tlolomitic. linirs. So cariticism of the hrat of solution n i r ~ t h w las such is iniplic~tl,:LS it may >-ield significant iriforniation on other tvpes of s p e i h e i i s :itid has heen used effectively 113- l\7c~lls:ind associates at the Xational Bureau of Standartlu. € I o n - c ~ c rthe . lack of high pri+ cision to make this mcthotl i w l l y tlc~trriiiincitiveis indicatvtl by the data ( 1 6 ) on 88 s:implrs of whitr-coat plaster. Diecrep:incies b e t w e n observed and c:ilculatetl h w t s oi .solutions range from +13 to -13 calorie?, with the fortuitouq rmult that the 011served arid calculated averagcls for :ill wniples agree cloiely, d i e r e a s for most of the specimens tlic, difference is of the order of 3 to 6%. Preparation of Calcium Hydroxide and Magnesium Hydroxide of Known Composition. In older adcquatelr to test various methods of analysis the conipoundi present in dolomitic t o havr availal)l(~calcium carlionate, hydrates, it is iiecws calcium hydroxide, m raiuni h>-di~ozi(li~. anti magnesium oxide of known composition. ('alcium cw1)oriatr is reddily obtainable fi,oni crystalline calcite or reagent grade calcium carbonate,. hut, in the authors' esperienccl, iirither ould lit, relied upon unless their purity is confirmed by :+iial> , .Uternatively, pure ridciuin carbonate may be ohtaincd from wagent grade c:iIriuni chloride hexahydrate and ammonium carbonate. For this xork after esperimerital test of many variations in niaterials and techniques, calcium hydroside of known reproducible composition was prepared as follo\vP :
Tn-enty-five-gram portions of pure pon-dereii calcium carbonate rwre placed in porcelain ciuciblrs and calcined for 3 hours a t 950" C. The calcinate was hydrated hy tieing refluxed with 500 nil. of freshly boiled distilled water for each 100 grams of cnlciuiii oxide
690
ANALYTICAL CHEMlSTRY
( , a O . 'lo Total loss on ignition.
cos, %
rc
Uncombined HzO. c; CaO moles per 100 g. CO2,'mole per 100 g. Combined Hz0, iiioles ~ , r ' r100 g,
Ca(OH)s, % CaC03.
%
i4.91 23.00 1.30 0.15 1.336 0 029.5 1.307 96.8 2.96
JIacnesium IIvdrox-ide MgO, 68.0l.j Total loss on ignition (7 32.00
coz, 9c
Loss a t 125' C., 5c Loss a t 280' C., % MgO, moles per 100 g COS, mole per 100 g Total H20, moles p ~ 100 r l\k(OH)z, %
J k O ! 70
0.12 0 34 0 63 1.688
0,003 1.760 98 3
75 37 24.60 0.64 0.05 1 344 0.01-1 1 329 98 5 1 46
69.82 30.44 0.13 0.14 0.24 1 732 0.003 1.68'2 97.4 2.42
H20 in excess of that c o l i t o Mg(OH)z, 70 1.4: 0.24" a 70 loss in weight a t 280' c'. is c~on.~idt~red t o be water in cxrc-5 oi t l i n t associated with Mg(0H)s in tlii- preparation.
for 3 hours and then the suspension (protected from atmospheric carbon dioxide) was allowed to stand for several days to ensure complete hydration. The solid was then filtered, m-ashed with water, and dried a t 110' C. for 3 hours to constant weight. The dried material was pulverized in an agate mortar to a fine powder, mixed thoroughly, then placed in a vacuum desiccator containing fused potassium hydroxide to complete the drying of the solid. The dried material %-asmixed thoroughly a second t,inie and was stored in a glass-sioppered bottle, which was in turn kept in a desiccator contaking fused potassium hydroxide. The analysis of txvo such 11re1iarationsof calcium liydroside is shown in Table 111. In both cases, qualitative tests disclosed that no detectable amounts of heavy metals, magnesium, or sulfate were present. The usual chemical methods were used t o obtain the per cent calcium oxide, per cent carbon dioxide, and per cent total loss on ignition. The per cent uncombined water was determined by measuring the loss in n-eight of a sample heated for 2 hours a t 125' C. in a carbon dioxide-free atmosphere. The comparison of t,he moles of combined water per 100 grains is excellent in bot,h cases, demonstrating that the per cent calcium hydroxide as obtained from the chemical analysis is correct to 0.2% relative or better. The calcium hydroxide so prepared and S O stored was found to remain essentially unchanged in composition for 1 year. Furthermore, the material could be weighed and handled in air with no particular precautions against carbon dioxide contamination escept speed in weighing. Thus 1 gram of the calcium hydroxide was found to gain 0.5 mg. in 4 minutes, while on the balance pan, a negligible change. The material is stable toa-ard carbon dioxide principally because it is dry, as the reaction between calcium hydroxide and carbon dioxide requires the presence of n-ater in order to proceed a t a reasonable rate. Magnesium hydroxide and magnesium oside for this work mere prepared following tests of many experimental variables, as follows: Reagent grade magnesium sulfate heptahydrate (246 grams, mole) was dissolved in 1 liter of water, and 250 grams of reagent grade ammonium carbonate were dissolved in 1 liter of water and 300 ml. of 15 M ammonia. The ammonium carbonate solution was then added slowly and with good stirring to the magnesiuni solution. The solution and precipitate were digested for 24 hours on the sbeam bath, then the magnesium ammonium carbonate precipitate was filtered with suction and washed with water until the washings gave a negative sulfate test. The solid so obtained was dried a t 110' C. for 24 hours, then the dried material was ground to a fine powder in an agate mortar. Then 25-gram portions of this dried material were calcined at 950" C . for 3 hours in porcelain crucibles for the magnesium oxide of known composition. The magnesium oxide so prepared was stored, as was the calcium hydroside and was freshlj. cnlcincd whenever pure magnesium
oxide was desired. Other portions of the dried magnesium ammonium carbonate were calcined at. 600" C. for 3 hours (a temperature high enough to drive off all the carbon dioxide present and yet not overburn the oside so that it n-ould be resistant, to hydration) to produce the magnesium oside for hydration to magnesium hydroxide. This magnesium oside was hydrated in a manner similar to that used for the preparation of calcium hydroxide, except that the suspension was boiled for 24 hours and the solid was allowed to remain in contact with water for 1 week. The solid nas then filtered with suction, washed with water, and dried at 110" C. for 3 hours. The dried material was then powdered in an agate mortar, miyed thoroughly, and dried for 1 w e k in an evacuated desiccator over anhydrous magnesium perchlorate, This material was again well mixed and was stored, as was the calcium hydroside. The analysis of two such magnesium hydroxide preparations i-; shox-n 'in Table 111, the results being obtained by the usual methods. Qualitative tests on these preparations disclosed no det'ectable amounts of heavy metals, calcium, or sulfate. -4sthcs summation of moles per 100 grams of carbon dioside and total water compared n-ith the moles per 100 grams of magnesium oxide shoiri. that in the first case, excess mater is present, and in the second, there is a deficiency of water, and as the formula oi the magnesium carbonate species present is in doubt, the percentage compo9ition based on the chemical analysis is much less certain than in the case of calcium hydroxide. The most probable magnesium carbonate species in the hydroxide before drying i* magnesium carbonate t,rihydrate, based on the aork of Davis (6 J. However, heating at 110" C . and drying under low pressure over magnesium perchlorate would convert the trihydrate t o some extent t o the nionohydrate. In any case, t8hepresence of magnesium carbonate species ~nalcesit impossible to determine accurately the uncomhined water by thermal methods. In the case of preparation 1, the percentage composition was calculated hJ- assuming 1 mole of magnesum oxide per mole of carbon dioxide and the remaining magnesium oxide was calculated as magnesium hydroside, since an excess of water is present and it is assumed that there is no free magnesium oxide. For preparation 2, the water equivalmt to the magnesium hydroxide was obtained by subtracting thP per cent loss in neight a t 280" C. from t,he total water. The free mag.nesi1.nn osicle was then ohtained by subtracting t,he sum of the magnesium oxide equivalent to the carbon dioside and thp magnesium oside equivalent to the magnesiuni hy(1rosidP from the total magnesium oxide. The justification for consitlering the loss at 280" C. to be the nonessential witcxr is given subsequently. In view of these uncertainties, the per cent magnesiu:ii hydroxide in these preparations is prohably reliable to only i O . 5 % relative. Evidently, magnesium hydroside prepared h!. hydration of the oxide (analogous to the mode of formation of magnesium hydroxide in hydrated dolomitic lime) is not so n.t~ll defined a chemical species as is calcium hydrazide. The significant difference, het>weenthe per cent loss a t 125' C. (run in accordancr nith the S S T M procedure for uncombined water) and the loss at 280" C. for these magnesium hydroside preparations is noteworthy. Evidently, magnesium hydroxide can retain unconihinrd water at temperatures well above those normally employed for thermal drying. Calcium hydroxide, on the othef hand, appears to give up its uncombined v-ater at 125' c'. l N F O R ~ l . ~ T I O FROM N X-RAY POU DER DIFFRACTION
The newer technique of recording s-ray polvder diffraction patterns on a strip chart by a spectrometer utilizing a GeigerPliiller tube to convert the diffracted w a y s to pulses of electrical energy v a s found highly satisfactory for recording the patterns of dolomitic hydrates and similar calcium-magnesium-containing mixtures, as compared to the older film method. Using a Norrlco instrument of the original 90" type, a diffraction pattern could be obtained in 30 minutes, whereas the film method required 4 hours' exposure or longer. Furthermore, the spectrometer tracing is more easily int.erpreted than is the film pattern, and the spec-
691
V O L U M E 2 4 , NO. 4, A P R I L 1 9 5 2
st o ~ i e ill most cases) and nould be coiivcxted to anhydrous hwdburned calcium sulfate which remains unhydrated during the suhsequent treatment because of inactivation due to overhurriing The sample was ground in a n agate mort'ar t'o pass 200 mesh appears to he a reasonable one, but confirmation of this would he and was mounted in a sample holder of Lucite or brass having a desirable. If a magnesium carbonate species is present, thc most iiiilled rectangular depression inch deep. The brass sample likely possibility would be magnesium carbonate ti'ihydrate. holder was found to be superior to the Lucite holder because a large halo is ohtained from the Lucite at low diffraction angles, Bernays ( 4 ) and the authors have both found, using x-ray which tends to mask weak diffraction masima in this region. methods, that t,his is the principal product of the reaction brtween Tlie 1- or 2-r.p.ni. scanning motors were found to give equally magnesium hydroxide, carbon dioxide, and ,water vapor, and that satisfactory results, the 2-r.p.m. motor having the advanthis compound can be formed in the presence of calcium hytage that i t yielded a pattern in one-half the time of the 1r.p.m. motor with less danger of significant carbon dioxide condroxide. Evidently, if magnesium rarbonate trihydrate is tamination of the sample during the exposure. Cu K a radiation present in the test hydrates, its conrrntration is below the detec(n-ave length = 1.5-4A . ) vias used and a diffraction pattern covertion limit (1 to 2%). ing the region of diffraction angles from 70' to 10" was obtained An effort wxs made to obtain quarititative information from the in 30 minutes using thr 2-r.p.111. scanning motor. diffraction patterns by preparing mixtures of known composition from the previously preparrtl f calcium hydroside, magnesium hydro\itie. magI nesium oxide, and reagent grade calcium carn ~honate. and comnarina the heights of thetliffracI -CALCIUM HYDROXIDE 1 2- MAGNESIUM HYDROXIDE tion peaks of the known mixtures t o t h o v of the 3- C A L C I U M CARBONATE 4- MAGNESIUM OXIDE test hydrates. However, the reproducibilit) of the instrument wa5 of the order of 10% a t I>eqt; hence, no reliablr ronclusions could hi. drr i w d from this study. Work on the known mixtures did permit the approwmate lorn-er limit of drtrction of the cornpounds present in low concentration to be ascertained. I n the cases of ralcium carbonate, calcium hydroxide, or magntwuin hv28 32 36 40 44 48 52 56 60 64 drouide, 3% vias readily detected, but 1 7% rould DIFFRACTION ANGLE-COPPER RADIATION not be detected with certainty; 1.7% magnr\ium Figlire 1. X-Ray I'owder Diffraction Pattern (Spectrometer Type) of Oxide but O.*% not'. limit is far helox the 10% claimed for magneHighly Hydrated Dolomitic Lime sium oxide detection in hardened plaster. (16). These observations were made on cornpounds The spectrometer pattern of a highly hydrated dolomitic lime whose particle size was in the optimum range; where the is shown in Figure 1. Diffraction peaks characteristic of calcium particle size is below the optimum range. the smallest concmcarbonate, calcium hydroxide, magnesium hydroxide, and magtration that. ran be detected is larger. nesiuni oxide are apparent. Powder diffraction patterns of the The data on diffraction peak height from the diffraction patfive test hydrates were all similar to Figure 1, in that no comterns of the five test hydrates are tabulated in Table I17. Runs pounds other than these four were detected. Hence, inforniation A-1 and ;1-2 are successive runs made on the same portion of as to the sulfate species present or the possibility of some type of sample, while runs R-1 and 13-2 arc. successive runs on a new portion of the same sample. Instrurnerltal settings were the same magnesium carbonate being present was not obtained. The assumption that the sulfur is mainly introduced during the burning for all patterns and the 1-r.p,in. scanning motor was u;ed. The pl'ocess of the limestone (since there is so little in the original strongest diffraction maximR nirasurrd for each conil~ounds r e
trometer method was fully as sensitive as the film method for the detection of compounds present in low percentage (1 to 570).
c
.
/I
Tahle IV.
X-Ray Powder Diffraction Data for Test Hydrates
Diffraction N a x i m a Measured ~ _ d spacing, A. 20 Cu radiation. 2. 63 34.1 2.35 38.6 3.04 24.3 2.10 43.0
Run
.4
-4-1 -4-2 B-1 B-2
R
C
D
E
~
I
___
Hydrate
Y
A-1 A-2 B-1 n-2 .%-I '4-2 B-1 B-2 4-1 A-2 B-1 B-2 .4-1 A-2 n-1
n-z
R e j e c t e d f r o m ;average.
Diffraction Naxinia Heights, Cm. C a ( 0 H ) z Alg(OH/; CaCOa XgO 9.20 8.55 7.30 6.50 10 00 9.50 10.15 8.90 10.65 9.0j 9.20 8 90 9.70 R 05 8 10 7.15 7 30 7 50 7.50 6883
2.85 2.85 2 55 2.85 3.80 4.00 3 80 :3.50 4.15 4.00
3 80 3 80 4.45 4 00 3.65 3.20 3 00 X 20 8.00
2,j.j
1.05 1.10 0 80 1.20 1.45 1.90 1.10 0.95 1.90 1.30 1.10 1.30 0.80 0.80 0 80 0 65 0.95 0 80 0.80 0.95
Ratio. of Heirlit&of ?lg(OH)r, CaCOa. a n d J I g O Diffraction Maxima t o Height of Ca(OH). Diffraction Maxima .\Ig(OH)s CaCOa MgO
_
1.45 1,75 1.60 1.45 1.45 1.45 1.60 1.60 0 95 0.80 0 95 0.80 0.50 0.65 0.50 0.50 0.80 0 95 0.95 0.95
Run .%-I .4-2 R-1 R-?
.I\..
Ca(0H);
n xx
0.33 0 . 3.5 0,440 0 33
Ca(0H);
CalOH);
0.11 0 13 0 11 0.18" 0 12
0.20 0.22 0 . '72 0.16'I 0.21
0.12 0.18" 0 14 0.12 0.14 0.13 0 08 0 09 0 10 0 09 0 09 0 11 0.13 0: 1 3 0.14 0 13
0 . 1.5 OR 0 OR 0 10 0 09 0 09 0 03 0 Oi 0 06 0 Oi 0 06 0.11 0.13 0.13 0.14
A-1
-4-2 B-1 E-2 .4r. -4 - 1 -4-2
R-1 B-2
Ir. -4-1 -4-2 B-1 B-2
A!.. .4-1 -4-2 B-1 B-2 -4v.
n . 38 0 ?!I,$ 0 44 0 41
0.43 0.43 0 46
0 1.1 0 45 0 44 0.46 0.41 0 43 0.40 0.37a 0 41
n
n 13
ANALYTICAL CHEMISTRY
692 shown in Table IV. The onl5- possible interference occurs in the case of the magnesium oxide peak at 28 = -13.0' ( d = 2.10 -4.) where there is a lesser calcium carbonate peak a t 43.2" ( d = 2.09 A ) . Hov-ever, this peak does not appear unless the calcium carbonate concentration is greater than l5%, and furthermore, these tw.0 peakb are resolvable when the 1-r.p.ni. mot,or is used. The peak height's above the estimated background have been measured to the nearest 0.05 cni. A study of the table reveals that the variation in peak height foi :i particular compound in a particular sample is too great to permit any justifiable conclusion as to the quantitative amounts prtwnt. I s the cheiiiic%l anal!-sis indicates that the per cent calciuni hydroside is betn-ren 49.6 arid 52.l%,aspreadof2.5%, the assumption has been niade that the variation in calcium hydroside among the various samples is too small to cause significant tiiff erences in the calvium hy(1roxide peak heights. IVork on known mixtures indicated that this assumption was correct. Therefore, the calrium hydroxide has been used as an "internal standnrd" and ratios of magnesiuni Ii\droside, calcium carbonat,e, and magnesium oxide peak heights t o the calcium 11)-dioxide peak height for c w h I'un have been tal)ulated in Table IV. Xn inspection of these data aho\vs that, if a fen- highly devi:tting result. are rejected, the ratios :igwe far better than do the indivitlual peali heights, as they trind to c.ompensate for iristrumental variation from run to run. -1c~orup:irison of the t,fJlative amounts of niagneqiuni hydroside, caalcium cxrhonate, and magnesium oside as tit~terniiiiedfrom the s-ray patterns Lvith the relative amounts as c~;ilcuI:itedfrom the cheniiral anal! X-Ray Patterns
Clieni. .Inal>siz
The agreement is, in general, good, except for hydrate C whic~h, from the s-ray patterns, alipears to have more magnesium hydroxide anti less magnesium oxide than calculated from chmical analysis. Or, comparing A v i t h the other samples for magnesiuni hydroxide and magnesium oxide ratios, the values in Table 1-are obtained for x-ray anti chemical analyses. (Since these results were ohtained with the Sorelco spectrometer, new trials on synthetic samples containing the constituents of these hydrates have been made on a nem- Ohio Spectron, a Geiger spectrometer for 165" angle scanning with a 1-ictorern logarithmic scaler. Although not in final adjustment necessitating estensive further tests, there is definit,e evidence of decided improvement in reproducibility arid reliability of peak intensities in evaluation of concentrations.) The agreement, while far from perfect, is sufficiently close to estend the hopes for the s-ra>-method n-hich potentially is so well adapted for material5 of this kind.
teniberature difference between the sample under test and a reference material (commonly aluminum oxide). I n the case of hydrated dolomitic lime, the characteristic endothermic breaks for magnesium hydi oside, calciuni hydroxide, and calcium carbonate in the differential thermal analyqis curve permit identification oi theye compounds in the sample.
T E M P E R A T U R E -'C
Figure 2.
Pigui,c, 2 shon-s the differenti:\l thernial anal!-& of a highly hydrated dolomitic lime. The first endot,herni, beginning a t 320" C. and cshibiting a niasimum heat absorption a t 415' C. is that for the decomposition of magnesium hydroside. The nest endotherm, lleginning a t 465" C. and eshihiting a niasinium absorption at 5TO" C., is dur to thr decomposition of calcium hydroside. The third endotherm, beginning a t 73.5" C . and exhibiting a masimum ahsorption at 865" (',> is due to calciuni carbonate. Diffewntial thermal analysis curves iverr pi,epared for the five test hydrcites, and were all similar to Figure 2 in that no compounds othrr than magnesium hytli,oside, cal(,iuiiihydroside, and calcium (,arbonate were detected.
p
-02-
I
-
>
-04r
2
-06-
0
~ and Magnesium Talde \-. Ratios of IIagnesium 1 1droxide Oxide in H>drates Referred to Sample A B/.\ C .I D/A E/A 31n(OH)z
X-ray Chemical
1.09 1.22
1 23 1.08
1.29
1.13
Differential Thermal inalysis Curie for Highly Hydrated Dolomitic Lime
B
-
c
-08,-
1.17 1.10
3IgO X-ray
Clieinical
.
, 4.-
'
,bo.
Figure 3. ISFORMATION FRO11 DIFFEREYTIAL T H E R M A L ANALYSIS CURVES
The techniques of differential thernial analysis, as described by Grim and Rowland ( I O ) , so successfullj- applied to the qualitative analysis of clay minerals, have been used to detect the t'hermally decomposable compounds present in doloniitic hydrated linir. The method detects esothermic or endothermic changes exhibited by the test. sample as it is heated a t a constant rate (usually 10" C. per minute) from room temperature to 1000° C. This is accomplished by plotting on an aut,oniatic recorder the
'
I
'
'
300°
'
'
'
'
5009
V I
,
/ ;
700'
,
,
,
,
9001
T E M P E R ATUR E - *C Differential Thermal Analysis Curve for Calcium Hydroxide
Some work on mistures of known composition demonst,rated t.he fact that the differential thermal analysis method, as ordinarily carried out on hydrated dolomitic lime, did not yield results capable of accurate quantitative interpretation on the basis of the evaluation of concent,rations fi,om peak intensities. An iniprovenient was not,ed in calculating ratios of arem under peaks, measured with a planimeter. But variations considerably greater than with s-ray patterns persisted espwially for magne-
V O L U M E 24, NO. 4, A F R I L 1 9 5 2
693
siuin hydroxide, and it. was deemed inadvisable to extend the techniques beyond qualitative observations. This apparent lack of quantitative reliability arises from the fact that the method is a dynamic, nonequilibrium one and a number of factors, some of which are not readily controlled, influence the magnitude of the thermal breaks. Some of these are the packing of the samples, the placement of the differential thermocouple with respect to the test sample and reference material, the heating rate, the particle size of the test sample, and finally, the fact that the endothermic reactions involved have associated with them very large heats of reaction compared to the one"C to which t,he method is normally applied. The sensitivity of the method is about equal t o the x-ray method-i.e., a compound niust he present to the extent of 1 to 2% to be detected. A comparkon of the magnitude of the endothermic breaks for calcium c:irt)onate in the thermal analysis curves for the five test hydrates Sh0\\.9 the following relative order of per cent calcium carbonnte:
C
> -1,B, 12 > D
This is in agreement n-ith the x-ray data and the chemical analysis. Hence these breaks for calcium carbonate are a more reliable indication of composition than the ratios in Tables I V and r-. THERMOCHEMICIL \IET€IOI) ( R I C I 1 R D S O N TECHKIQUE)
Inspection of Figures 2 and 3, the differential thermal analysis curves for a highly hydrated dolomitic lime and for calcium hydroxide, suggests a method for the direct, determination of magnesium hydroxide and calcium hydroxide in dolomitic hydrates. A measurement of the loss in \\eight (or the aeight of water evolved) at controlled temperatures should represent the amount of water equivalent t,o the magnesium hydroxide and the water equivalent t o the calcium hydroxide, provided that the proper critical temperatures c:in be determined. The differential therni:il analj curves of Figures 2 and 3 give an approximate, idea of these temperatures. The deconiposition of magnesium hJ-droxide appears to become measurably niaterial temperature of 325" C. rapid n t a "furnace" 01' i~t~ference and reaches a maximum at, 415' C. Deroniposition of calcium hydroxide appears t,o hegin at 400" C. and the rate of decomposition reaches a maximum a t 550" to 5i0° C. A knowledge of the equilibrium water vapor pressures from the solid hydroxide a t various teniperntures ivould serve to locate these critical temperatures more exactly. I:vidently, the decomposition of the hydroxide hy ill hegin n-hen the pressure of the water vapor from the solid exceeds the pressure of the water vapor in the atmosphere surrounding the solid. If this atmosphere is a normal or uncontrolled one, the prevailing xater vapor pressure will be variable and the initial decomposition temperature ill also be variable. As t,he water vapor pressure in the atmosphere is usually between 10 and 20 nun. of mercury, the initial decomposition of the hydroxide would he expected to occur a t temperatures corresponding to this water vapor pressure. The decomposition of the hydroxide will proceed a t a masinium rate when the water vapor pressure from the solid exceeds the prevailing atmospheric pressure. Inspection of the literature reveals some discrepancies in the data for the equilibrium water vapor pressures of the two hydroxides as a function of teniperat,ure, particularly for niagnesium hydroside. A summary of some of the values, with their sources, is as follows: Fur the system Ca(OH), (sj = CaO (s) H?O (g)
+
Temp.,
c'
PreyiurP
H20. Mm.
301 348 397 547
2.7 13.4 24.3 760
3'20 513
7 60
10
Reference
For the system Mg(OH)n (s) = AIgO (s) Temp.,
35 79
C.
+ H,O (g)
Pressure
U?O, 311~1. 7.6
171
113 63ii
300
10
Reference
(10)
Johnston's values ( 11 ) are quoted in the International Critical Tables and those for calcium hydroxide appear to he in line with the differential thermal analysis of Figure 3, hut the values for niagnesium hydroxide are in serious disagreement with Figure 2. Inasmuch as Johnston's values were obtained by free energy calculations from heat of hydration data, the data used for magnesium hydroxide are probably in error, part,icularly in view of the fact that the rate of hydration of magnesium oxide varies so markedly LT-ith the calcination temperature used in its preparation. Furthermore, the speed with w.hich magnesium hydroxide attains equilibrium with magnesium oxide and water vapor appears t o be slow, xhich means that in practice a higher temperature than that predicted must be used to complete the reaction in a reasonable period of time. There are several reports in the literature on the temperature range for the thermal decomposition of calcium hydroxide and magnesium hydroxide as determined by a thermobalance in which the sample is suspended from the balance arm and held directly in the heating unit, so that, t h r loss in iwight as a function of temperature may be directly determined. Using such an arrangement, Gill (9) found that caloiurn hydroxide decomposed most rapidly at 520" to 525' C., but the decomposition was not complete until 675" C. Jlagnesium hydroxide as found to decompose most rapidly at 375' to 380" C., but was not complete until 475" C. Duval and Duval (Y),using an automatically recording therniohalance, found that the thermal decomposition of magnesium hydroxide prepared by precipitation from aqueous s o h tion began at 250" C., v a s most rapid at 350' C., but was not complete until 819" C. These Fame ivorkers found that a sample of hrucite (naturally occurring crystalline magnesium hydroxide) showed an initial thermal decomposition at 288" C. and thereaft,er a continuous loss in weight to 815' C. The fact that the thermal decomposition of magnesium hydroxide is not completed until very high temperatures is in confirmation of the findings of Fricke and Hiittig ( 8 ) ,who state that the last traces of water are not removed from magnesium oxide until 1150" C. It is also significant that. the Duvals ( 7 ) found that the precipitated magnesium hydroxide retained uncombined water up to 224" C. The success of the therniochemical method will evidently depend on an accurate differentiation of uncomhined and combined water, as well as differentiation of the combined water belonging to the t n o hydroxides. In view of t,he preceding discussion, it appears that the loss in weight at 250' to 300" C. should represent the uncombined water, and the increased loss in weight a t 375" to 390" C. should represent the water belonging t o the magnesium hydroxide (although the decomposition of magnesium hydroxide may not be completed in a reasonable period of time a t this temperature); then the increased loss a t 525" to 550" C. should represent the water belonging t o the calcium hydroxide, as it has been shown (18) that this is the upper temperahre limit for heating calcium carbonate without significant decomposition. Some means must also he provided t o exclude carbon dioxide from the sample during t,he heating period, since calcium hydroxide (or calcium oxide) will react with carbon dioxide up to the threshold decomposition temperature of calcium carbonate (525' to 550" C.). The thermochemical method was originally applied to hydrated dolomitic lime by Richardson ( 1 2 ) in 1926. The weighed sample was placed in a crucible, onto which a cover equipped with glass stopcocks was sealed so that air free from carbon dioxide and water could be passed over the s a m d e during the heating period. The crucible and sample were then
ANALYTICAL CHEMISTRY
694 Table VI. Thermobalance Data for Calcium Hydroxide, Magnesium Hydroxide, Calcium Carbonate, and Test Hydrate A Calcium Hydroxide (Prepn. 1) Total elapsed Accumulated time, Temp., weight loss, c. % hours 0.5 1 1.5 2 2.26 2.5 2.75 3 3.25 4 25 4.;. 4. , a
6 4.5 D 75 6 7 7.5 8.7 Const. w t .
110 150 175 190 260 295 326 350 390 400 420 445
465 490
520 525 526 600 600 900
300 450 b50 600 bOO
600 700 900
0.02 0.06 0.08 0.09 0.09 0.12 0.13 0.14 0.14 0.14 0.14 0.70 2.44 11.25 18.47 23.09 23.09 23.52 23.58 25.02
0.00 0.00 0.03 0.05 0.09 0.13 0.65 43.91
Magneaium Hydroxide ( P w p n . I ) Total elapsed Accumiilated time,Temp., \wight loss, hours c. 40 0.6 1 1.; 2 2.5 3
3.5 4 4.5
5 6 8 8.5 9 23.5 25 26 27 28 28.5
Const. wt. 0.5 1
;.
2.5 3 3.5 4 4.5 5 5.5 5.76
6
6.6
7
7.5
R
s.6 9 9.5
Const. wt.
heated in a muffle furnace for a 20-minute period at a specific temperature, the loss in weight was determined, and the 20minute heating period was repeated a t a new higher temperature. A graph of loss in weight versus temperature myas then constructed and the break which occurred in the region 400" to 450" C. was taken to be the water equivalent to the magnesium hydroxide, while the nest break, occurring in the region 500' to 550' C., was taken to lie the water equivalent to the calcium hydroxide. Working with calcium hydroxide and magnesium hydrosidc of only approximately known composition, Richardson found that magnesium hydroxide dissociated most rapidly a t 400"to 450' C., being virtually complete at 460" C., while calcium hydroside decomposition began at 430" C. and was completed a t 575" C. Hence, there was some overlapping of the t,wo decomposition regions, although the error due to incomplete magnesium hgdroxide decomposition is somewhat compensated for by some calcium hydroxide decomposition. Evaluation of the accuracy of the method was difficult, as reliable standards were not awtilable. It therefore seemed desirable to reinvestigate the therniochemical method, as it seemed to be the most logical method for the direct determination of magnesium hydroxide in the presence of uncombined water, magnesium oxide, calcium hydroside, calcium carbonate, and calcium sulfate, and calcium hydroxide and magnesium hydroxide of known composition were available. Preliminary work on these preparations using essentially Richardson's technique was carried out. The sample m-as contained in a platinum crucible, which n-as in turn placed (on a piece of asbestos paper) in a cast iron Skidmoretype crucible, and this in turn uas heated in a sand bath. A stream of carbon dioxide- and water-free air was passed into the entrance tube of the Skidmore during the heating period and the temperature was measured by a calibrated iron-constantan thermocouple which was run through the exit opening of the Skidmore and was in contact with the inside of the platinum crucible. For comparison, runs viere made in an ordinary muffle furnace a t the same temperatures and for the same lengths of time. When
the sample is heated in contact with a stream oi carbon dioxide- and water-free air, the water vapor pressure in the immediate vicinity of the sample is virtually zero (depending on the desiccant used to dry the air stream-anhydrous magnwiuni perchlorate was used in this case): hence, the decomposition of the hydroside becomes appreciable at a much lower temperature than if the samplc is heated in contact with a normal atmosphere.
This preliminary ~ o r ghoxeti k that significant decomposition of magnesium hydroside occurred on heating for 1 hour a t 280" C'. i n a carbon diosideand water-free atmosphew, nhile significant decomposition of calcium hydroxide occurred on heating for 1 hour at 300" C. under the same conditions. On the other hand, heating nugnesium hydroside for 1 hour at 250" C. in a normal atmosphere produced no decomposit,ion, while calcium hydroside could be heated for 1 hour at 300" C. under the same condit,ions viith no decomposition. Carbon dioside analyses of thc residues heated in a normal atmosphere shon-ed that no car1)on dioxide was gained by t,he magnesium h ~ d r o s i ~ lduring e the 1ie;lting period, d i i l e calcium hydroxide tended t o gain carbon dioxide. t h c net gain in weight from this effect being about 0.6 mg. per gram of calcium hydroxide per hour. IToxever. this effect could not mask any signifirnnt ;tnlount of calciuni h>-droside decomposition. It becaanie apparent that the 01)served loss in weight a t a given temperature using Richardson's technique was critically dependent' on time and is evidently the rewon why Richardson arbitrarily chose 20-minute heating intervals. For example, a sample of calcium hydroxide (preparation l), heated a t 300" C. for 0.5 hour in a carbon dioxide- and water-free atmosphere showed a loss in weight of 0.33% (uncombined wat.er in sample O.lJo/o), but another sample, heated for 1 hour under thc same conditions, showcrl :t loss i n weight of 7.08%. ANALYSIS WITH A SI3IPLE A S D SUC
S S € U L THERMOBALASCE
Inasmuch as the time factor cannot he precisely controlled when the sample must, bc geriodicsll!- placed in and removed from t,he heating unit and some time iy rrquired for the sample to mich the prevailing tenll)erat,ure of the heating unit, a thernioImlance type of arrangement :lppe:tred t o he the logical method for the determinat,ion of t,he loss i n \\-right as a function of both teniperature and time wit,h a n~asimumcontrol of conditions. h very siniple thermobalance was therefore set up and its essential features are depicted in Figure 4. The observed loss in weight of a sample as determined by this thermobalance is considered t o be correct t o f1 mg. or bebter. ;LS a 5-gram sample was customarily used, this corresponds to a n uncert,ainty in the loss in weight of &0.02% absolute. Test runs on samples of known loss on ignition (as barium chloride dihydrate and calcium carbonate) gave reaukq'for the per cent loss in weight correct to 0.1 to 0.2% relative. The operation of the thermobalance for a dolomitic hydrate is as follows: a 5.000-gram sample is trarzqferred to the sample crucible, which is imniediate13- connccted to the suspension in the muffle furnace, a stream of air free from carbon dioxide and water being passed into the muffle during this operation and for the remainder of the heating period. A few minutes are allowed for the system to equilibrate, then the initial weight is determined (since the empty weight of the system is knoa-11, a check on the sample weight is thereby obtained). The niufflr is then turned on, this lieing considered zero time, and the elapsed time is measured from this point. The muffle temperaturc is then allowed to rise to the previously determined critical temperature, weight readings being taken a t 0.5-hour or 15-minute intervals. When constant weight is reached, the muffle temperature is a l l o ~ ~ etod rise to the nest critical temperature and is held there until the weight loss is constant, and PO on. In general, constant weight vas considered to have been rcached n-hen the observed change in weight
V O L U M E 2 4 , NO. 4, A P R I L 1 9 5 2
695
did not exceed 2 mg. (for a 5-gram sample) for a 15-minute heating period a t a particular temperature. Therrnobalance runs on the previously prepared calcium hydroxide and magnesium hydroxide and 011 reagent grade calcium carhonate n-ere made t o determine the various critical temperatures under the thermol)alanre conditions. Typical results are tabulated in Table VI. T h e behavior of magnesium hydroxide is especially notexvorthy. Constant weight is apparently reached in the region 280" t o 300" C.; nevert'heless, less than half the calculated excess of water shown in Table I11 has been lost. Some of this excess water is proh;tk)iy combined with the magnesium carbonate species and would not be released until temperatures above 300" C. have been reached (4);however, the per cent carbon dioxide is too small t o correspond t o enough water t o account for the large observed difference. Above 300' C., magnesium hydroxide begins decomposing (rapidly above 350" C.) and an appxreiit constant weight is reached in the region 390" t o 400" C., but the observed loss in weight, in this region, after subtracting the c:tlculated excess water and the carbon dioxide, most of which should he driven off at, this point, corresponds t o about 89% of the calculated amount of magnesium hydroxide present. Further work on the magnesium hydroxide preparations demonstrated that when an apparent constant weight was reached in the region 375" t o 400" C., 90. 4~ 1% of the magnesium hydroxide present was decomposed. T h e time required for this "pseudo-equilibrium" t o he reached was of the order of 4 t o 5 hours at 375" t o 400' C. for 5 grams of magnesium hydroxide. In cwnfirmation of previously reported work (4))temperatures in escess of 900" C. were necessary to achieve complete decomposition of the magnesium hydroxide. Carbon dioxide determinations on residues from simples of niagnwiuni hydroxide heated to const:uit weight a t 280" t o 300" C. shoved that t,here \vas no carbon dioxide gained by the sample during the heating period. The data in Table 1-1 indicate that calcium hydroxide can be heated for 1 hour at 390" to 400" C. without significant deconiposition, rdpid decomposition commencing above 420" C.; however, with heating periods of several hours in the region 390' t o 400" C,> some decomposition of calcium hydroxide occurred. Hence, t o decompose magnesium hydroxide (st least t o the ex-
I
1 *BALANCE
T
PYREX
1 1" H O L E
SUPPORT LEDGE FOR BALANCE
7
-r-
,,2"
I MUFFLE
HOLE
PT'
28"
T CRUCIBLE
Figure 4.
--
MUFFLE FURNACE
Diagram of Thermobalance Arrangement
tent of 90%), but not deconipose significant amounts of calcium hydroxide, it is necessary t o keep the sample between 375" and 390" C., since below 375" C. t h e decomposition of magnesium hydroxide is excessively slow. At 525" C. an equilibrium weight loss for calcium hydroxide is reached corresponding t o 97.4% of the calcium hydroxide priginally present. T h e equilibrium weight loss a t 600" C. corresponds to 99.6% of the calcium hydroxide originally in the sample. The data for calcium carbonate in Table 1-1indicate that decomposition begins at ca. 550" C., but even after heating for 2 hours a t 600" C., only 0.3% of the calciuni carbonate taken has been decomposed. Carbon dioxide analyses of residues of samples of calcium hydroxide heated t o constant weight a t 525" and 600" C. showed that ca. 10 mg. of carbon dioxide were gained by a 5-gram sample during the heating period. ilfter correcting for this small carbon dioxide effect, the decomposition of calcium hydroxide is found to be 98% complete a t 525" C. and virtually 1 0 0 ~complete o a t 600" C. Table VII. Thermochemical Data and Thermochemically Computed Per Cent Magnesium and Calcium Hydroxides for Test Hydrates Hydrate
I B
c
D
E
Accuniulated Weight Loss, % 280' C. 380' C. 0.46 10.1 10.6 0.43 0.47
11.4
0.59 0.49
11.3
Calcd.
Calcd.
% hIg(OH)x
Ca(OH)?
34.6 36. < 39.3 41.1 38.9
47.6 50.2 50.5
12.0
%
51.7 51.7
Table VIII. Magnesium Hydroxide Computed from Chemical Analysis on Basis of Loss at 280" and 120" C. Hydrate A B
c
D E
Calcd. from Chem. Analysis by Considering Loss at 280' C'. t o be Uncombined W a t e r , yo 34.8 36.4 37.7 39.8 38.3
Table I1 Value (1200 Loss), % 35.6 36.4 38.4 40.2 39.1
I n view of the fact that the decomposition of magnesium hydroxide is only 90% complete a t 375" t o 390" C. andstill incomplete a t 600' C., the thermobalance method is useful only for the direct determination of magnesium hydroxide by application of the empirical correction factor t o the ohserved loss at 375" t o 390" c. The thermobalance data for hydrate -1in Table VI are typical of the thermochemical behavior of all the test hydrates. These thermobalance runs revealed the fact that a t 300' C., the loss in weight continuously increased by 0.04 t o 0.05% per 0.5 hour (after the large initial loss in weight had occurred) and no better constancy of weight could be reached. Although the data of Table TI for magnesium hydroxide (preparation 1) indicated t h a t not all t h e excess water was driven off when apparent constant weight was reached a t 300" C., it appears t h a t t h e magnesium hydroxide of the test hydrates is decomposing a t a slow rate a t 300" C. As a n apparent constant weight is reached a t 280" C., the loss in weight a t this temperature is taken t o be the uncombined water in the sample. The data for the per cent loss in weight at 280' C. for the test hydrates may be found in Table VII. With the exception of hydrate R , these losses are 0.1 t o 0.3% absolute higher than t h e per cent loss in weight as determined by the ASTM method (1) for uncombined water. T h e lossp at 280" C. were reproducible to &0.05% absolute. It is possible, in view of the behavior of magnesium hydroxide (preparation l), that the loss a t 280' C. still does not represent all the free water in the sample; nevertheless, it is believed that the loss in weight a t this temperature is a better approximation of the true value than the loss in weight a i 125' C. Two possible sources of error for this determination of uncombined water a t 280" C. have been found t o be negligible. One of thew, carbon dioxide pickup by the calcium hydroxide of
696
ANALYTICAL CHEMISTRY
the sample, resulted in a net gain in weight of only 1.0 m g . for a 5-gram sample after reaching constant w i g h t a t 280" C. (2 hours t,otal time in the thermobalance). It has also been suggested (1;) that the heating of the sample t o drive off the uncombined water vi11 cause some of this lvater t o react with the magnesium oxide present t o form magnesium hydroxide, since the rate of hydration of mag'nesium oxide increases xith increasing temperature, thus giving low resulk for the determination of unconiliined water. To check this point, k n o m amounts of water were added t o xeighed samples of hydrate .4 (the hydrate containing the largest percentage of free niagnraium oxide) arid these samples were then immediately heated in :t normal atniopphere a t 110' C. for 1 hour or a t 250" C. for 0.5 hour. The loss in xveight \vits determined and the difference t~et\veen the loss in weight of a sample without additional water and the loss in weight of a sample t o n-hich water had been added \vas conipared with the known ;rmourit of water added. The results w r e a s follo\vs: Temp.,
c.
Heating Tiiile, IIOlll
110
1
110
1 0 5
250
Water Added, LIg,
Water R ~ c o v e r e t l , Mg.
37 0 86 6 35,6
37 0 86 8 35.3
the amount+ of wstei, added were 10 to 20 tinies the estiniated amount of uncombined water in the 1-gram samples of hydrate used, it, is concluded that fixation of uncombined water to magnesium hydroxide during the heating does not occur: :it least in highly hydrated dolomitic lime. Experiments conducted on the test hydrates showed that the same results for the loss a t 280" C. as in t'he thermobalance (within 3Z0.05% absolute) were obtained by heating the samples in an oven or muffle in a normal atmosphere. I n either case the gain of carbon dioxide by the sample caused a negligible error. Heating a 5-gram sample of hydrate a t 280" C. for 1 hour \vas sufficient to produce constancy in weight. It is therefore proposed that uncombined water in hydrated dolomitic lime be determined in this manner. The temperature should be controlled to 3Z 10" C. or better. JVater from any calcium sulfate dihydrate whose presencgis probably unlikely in comparison wit,h anhydrous calcium sulfate ( 1 , 2, 1 7 ) and some of the water of hydration of magnesium carbonate species (approximately two-thirds of the water of hydration of magnesium carbonate trihydrate that might be present) will be included in this determination of the unc-ombined kvater. As previous work on calcium hydroxide indicated that prolonged heating at 390" to 400" C. resulted in some deconiposition of the hydroxide, the test hydrates Tvere heated t,o 370" to 380" C. to decompose the magnesium hydroxide. -1true constant' weight could not be reached a t this trniperature. By recording the weight loss a t 15-minute intervals in this temperature range, after 1.5 hours' heating, a minimum change in the weight loss was reached, followed by an increasing change in the weight loss as the heat,ing was continued a t 380" C. This behavior is due in part to slight, continued decomposition of magnesium hydroxide, but because the change in weight loss per change in t,ime actually increases with time after passing through a minimum, it seems likely that calcium hydroxide is now slowly decomposing. The decomposition of calcium hydroxide is largely suppressed while the magnesium hydroxide is decomposing, because the dcconiposition of the magnesium hydroxide maintains the water vapor pressure in t,he vicinit,y of the sa!iiple a t a sufficiently high value. The minimum change in the per cent weight loss with time was of the order of 0.05% in 15 minut,es. It' is necessary to use a thermobalance to locate this minimum. Efforts to reproduce these results by Richardson's technique failed because the time factor could not be precisely regulated. The values for the per cent loss in weight a t 380" C. corresponding to the minimum change in weight loss with time for the test hydrates have been recorded in Table IT and were found to be reproducible to +O. 1%
absolute. Carbon dioxide analyses of residues heat,ed for 1.5 hours a t 380" C. showed that the gain in weight due to carbon dioxide pickup was ca. 4 mg., demonstrating that the magnitude of the "carbon dioxide error" increases with increasing temperature. However, the carbon dioxide error at 380" C. is somewhat less than the observed precision, whicnh is about 1% relative. The difference between the acwniulated weight loss a t 380" C. and the accumulat,ed loss a t 280" C. is taken to represent 90% of the water belonging to the magnesium hydroxide, and the per writ magnesium hydroxide in the test hydrates has been so calcul::ted, the results appearing in Table VII. I n vieu. of the observed precision for the determination of th(1 losses at 380" and 280" C., the uncertainty in the empiricallj. determined correction factor, and the carbon dioxide error, the prob:rt)le error in this determination of the per cent magnesium hydroside is considered to be about 2% relative. The per cent citlcium hydroxide has been computed by subtracting the ivater equivalent to the magnesium hydroxide as determined thelmocheniicall!-, the loss a t 280" C. representing the uncombined JvaTer, and the per cent carbon dioxide (all carbon dioxide being assumed to be present as calcium carbonate) from the total loss on ignition, the difference being taken as water equivalent to the calcium hydroxide. These results appear in Table 1-11. As it has been previouslj- shown that the loss a t 280' C. is a }letter approximation of the ur;romt)ined water than is the loss a t 120' to 125' C., the per cent magnesium hydroxide as obtained thermochemically should be compared with the per cent miignesiuni hydroxide calculated from the c*heniicalanalysis by taking the loss at 280" C. to be the uncombined water, rather than with the values for the per cent iiiagnesiuni hydroxide in Table 11. These recalculat,ed values appear in Ta1,lc VI11 along with the v:tlues repeated from T a b k 11. \\.it11 the exception of 1.1) Ir:ites C: and D, the values agrct' with the thermochemical value to 1.6% (wlative) or better. This :rloncl, however, cannot be cwnsiclri,rd justification for considering thcx loss a t 280" C. to be the uncombined water, as the per cent magnesium hydroxide for hydrates .A, B,and E as calculated in Table VI1 are in about as good agreement with the thermoc (tal values as the recalcu1:iteti d u e s from the chemical anal) Tablr YII. Hydrates C and D a p p e : ~to be out of line with the c*h(micdcalculations. The x-ray data of Table I V showed that the per cent magnesium hydrmicle of hydrate C was a little larger than that of hydrate E, which is in ag1,eement with the thermochemical calculations. Howrvrr, thcx per cent calcium hydrouide for hydrate C, as calculated from the thermochemical dat,a, does not agree with the value as calculated froin the chemical analysis. The therniochemically calculated pw cent calcium hydroxide for hJdrate D is also out of line \vith the c4iemical analysis value, while in the case of hydrates -1,B,and I.:, the thermochemical values for the per cent calcium hytirouitlr :igree to 1.4% (relative) or better with the chemical analysis values. I t is shown suhnrquently that the per cent calcium hJ-droside of hydrates A, H, D, :rnd E as obt,ained by direct cheniiral methods agrees to I .5y0relative or bett,er xith that, calculated thermochemically. Howcvcr, t>het'hermochemical value for the per cent calcium hydroxide of hydrate C could not be chrcked by direct chemical methods. The fact that the thermochemically determined per cent calcium hydroxide for four of the five tcst hj-drates agrees within the probable experimental error with that determined by direct chemical methods is strong confirmation of the validity of the t>hermochemicalvalues for the per cent magnesium hydroxide, and, to a lesser extent, the validity of the assumption that the per cent loss a t 280" C. represents the unconihined water. In the case of hydrate C, direct chemical methods for the per cent, calcium hydroxide agree sonieivhat better with the calculated value from the chemical analysis than with the thermochemical value. The only evidence supporting this apparently high thermochemical value for the per cent magnesium hydroside of hydrate C is the x-ray data. Aside from the difficulties inherent in the thermochemical method, which are primarily due t,o the reluctance of magnesium hydroxide t o give up its uncombined water as well as its conibined water a t temperatures below the initial decomposition
V O L U M E 24, NO. 4, A P R I L 1 9 5 2 temperature of calcium hydroside, the thermochemical method for the per cent magnesium hydroxide is comparatively free of interferences. A possible interference would come from t,he presence of magnesium carbonate species in the sample. As stated previously, the most probable compound of this type present in the sample would be magnesium carbonate trihydrate. This compound teiids to lose 2 molecules of water below 300" C. to form the nionohydrate ( 6 ) . Such water would be counted in with the uncombined water from the loss at 280" C. in the thermochemical method. Above 300" C., magnesium carbonate monohydrate tends to lose both water and carbon dioside, forming magnesium oxide. This bvater would be counted as v-ater belonging to magnesium hydroxide in the thermochemical method, while the carbon dioxide tends to react with calcium hydroside present, forming calcium carbonate and releasing an equivalent amount of water, which is also counted as water belonging to magnesium hydroxide, This decomposition of magnesium carbonate monohydrate occurs slowly :it temperatures from 300' to 400" C., then more rapidly from 400' to 500" C., so that the error introduced in the per cent magnesium hydroxide is not equivalent to all the niagnesium carbonatemonohydrate formed from the decomposition of t8he magnesium carbonate t,rihydrate originally present. Although no evidence has been obtained for the presence of significant amounts of magnesium carbonate species in the test hydrates, its possible presence in dolomitic hydrates and consequent interference in the thermochemical mcthod for magnesium hydroxide niust be borne in mind.
697 Ascarite and anhydrous magnesium perchlorate is placed in the top of the funnel. \Vhen filtration has been completed, the carbon dioxide trap is removed and the extraction flask and filter are washed with 25 ml. of sucrose solution in three portions, then 30 ml. of water in three portions. The filtrate is then transferred to a 400-ml. beaker, and the filtering flask is rinsed thoroughly with water, then with 25 ml. of 12 JI hydrochloric acid, which is added to the solution in the beaker. This solution is diluted to 200 ml., 3 grams of reagent grade ammonium oxalate dihydrate and 3 drops of 0.1% methyl orange are added, and the solution is heated to 80" to 90" C. all solid material has dissolved, the solution is removed from the source of heat and the calcium oxalate is precipitated by the dropwise addition of filtered 15 M ammonia with good stirring, the addition of the animonia cont,inuing until the indicator just assumes a yellow color. The precipitate and solution are allowed to stand without further heating for several hours or until the calcium oxalate has conipletely settled out. The precipitate is then filtered, washed with 0.1% ammonium oxalate solution, ignited at 950" C., and weighed as calcium oxide. The per cent calcium hydroxide in the original sample is then computed from the n-eight of calcium oxide obtained,
Table IX. Results on Known Amounts of Calcium Hydroxide by Gravimetric Sucrose Method Eupt.
la
Sufficient sample to represent 0.5 gram of calcium hydroxide (1.000-gram samples of highly hydrated dolomitic lime viere used) is weighed and transferred to a 150-ml. estraction flask. Fifty ml. of a 30% sucrose solution (prepared by dissolving 428 grams of commercial cane sugar in 1 liter of carbon dioside-free distilled water) and a small amount of paper pulp are added, the flask is tightly stoppered, and the suspension is vigorously stirred by a magnetic stirrer for 1 hour. Extraction is a t room temperature, since t,he solubility of calcium hydroxide decreases with increasing temperature and higher t,emperatures favor the precipitation of insoluble tricalcium sucrate ( 3 ) . The suspension is then immediately filtered through a special filtration arrangement which is designed to minimize carbon dioside contamination. This device is prepared by placing a 1-inch W t t e plate in the apex of an ordinary Gooch filtering funnel (1.5 inches in diameter), then placing a 2-mm. layer of asbestos over the IF-itte plate. The asbestos layer is washed thoroughly with 1 liter of water, then a second \Vitt,e plate is placed over it. The suspension in the extraction flask is quickly transferred to t'he prepared Gooch filtering funnel, the filtrate being received in a 125-ml. filtering flask. When the susuension has been transferred to the filtering funnel, the extractibn flask is quickly rinsed out two or threetimes with water, then a calcium chloride-type drying tube containing I
I
CaCOi
2b 3
4 5 6 7 8 9 10 11 12 13 14
C.4 LCIUM HYDROXIDE BY GRAVIMETRIC SUCROSE hlETIIOD
The inci,eased solubility of calcium hydroxide (and ot,her calcium coinpounds) in sucrose solutions, as well as other organic solvents containing adjacent C-OH groups such as ethylene glycol or glycerol, has long been known (14, l j ) , and has been made the basis for a rapid volumetric method for the deterniination of calcium hydroxide (as \Tell as calcium oxide) in the presence of calcium carbonate in lime. Shead and Heinrich ( I S ) have demonstrated that calcium hydroxide could be quantitatively extracted from magnesium hydroxide (ormagnesiumoxide) by a 30% (by weight) sucrose solution \T-ith no solubilization of magnesium compounds. They utilized this behavior as the basis for a rapid method for the determination of calcium and magnesium in limestones. This principle appeared readily extendable t,o dolomitic hydrates to give a direct method for the determination of calcium hydroside. A series of experiments was therefore carried out using known amounts of calcium hydroxide from preparations 1 and 2 described in Table IIT, to test this method, with essentially the technique of Shead and Heinrich ( I S ) . The esperimental technique worked out for t,his determination is as follows:
Other Compoundz Piesent
18 19
CaSOa (anhrd.) Same a s 17 Same as 17
a b
MgO .\Ig(OH)z
70
Gram
0.4939 0.4944
0.491.:
0.5188
-5.6
0.4915
0.4979
+1.3
0 4915
0.4917
0.0
0 4015 0.4915
0.4947 0 4940
4-0 6
0 5
17
16
+0.6
Errol
0.4925 0.4840 0.4840 0.4840 0.4915 0.4915 0.4915 0.4915 0.4915 0.4915 0.4915 None Sone 0.4915
0 5
MgO
(Rel.)
Giam
0 05 0.1 0 1
.\lg(OH)z CaCOa CaSOa. 2H20 CaCOs CaSOi (anhyd.) CaCOa,
15
Ca(OH)? Found Gram
0.05 0 05 0.0%5 0 05 0 05 0 05
CaCOa
Ca(0H)z Taken
0 5000 0.5000 0 05 0.05
n
4 ~ 2 6
+0.3
+n 2
n. F,.
0.05 0 .O b 0.03 0.0.5 0.05 0.05 0.35 0.01
fO
5
Extraction time 20 minutes. Extraction time 30 minutes.
I n all other experiments, extraction time was 1 hour.
The data in Table I S indicate that the method in general yields high results, but in the presence of the amounts of magnesium hydroside, magnesium oxide, calcium carbonate, and calcium sulfate (anhydrous) usually found in highly dolomitic lime, the results are correct to 0.6% (relative) or better. iis there was no blank on the reagents, the high results are principally due to the extraction of some calcium carbonate and calcium sulfate as shown by experiments 12, 13, 15, and 16. The interference of calcium sulfate dihydrate is far greater than that of anhydrous hard-burned calcium sulfate (prepared by calcining reagent grade calcium sulfate dihydrate at 950" C.). Because calcium sulfate is more soluble than calcium carbonate, t.he former would be espected to interfere to the great,er extent, as the data show. Experiments 6 to 11 show that the extent of the calcium carbonate interference is independent of the amount of calcium carbonate present, while experiments 15, 16, 17, 18, and 19 show that the extent of the calcium sulfate interference is dependent on the amount of calcium sulfate present,. For this reason, this method would be impractical for determining calcium hydroside in plasters and white coats. It would seem desirable to increase the specificity of the determination and shorten the time required t o perform it by titrating the hydroxyl from the Eolubilized calcium hydroxide with st'and-
ANALYTICAL CHEMISTRY
698 ard acid, any solubilized carbonate being removed by the addition of barium ion. I n such an event, it would still be necessary to filter off the unextracted solid, because the titration could not be precisely performed in the presence of magnesium hydroxide, whose natural pH is 10.4. Shead and Heinrich (13) state that the calcium-sucrose complex apparently ties up small amounts of hydroxyl which does not readily react with acid. This point has been confirmed by the authors. Direct titration of extracted calcium hydroxide solutions in the presence of barium by standard hydrochloric acid t o a thymophthalein end point (pH ea. 9) gave results that were 1.5% (relative) low, on the average. Carbon dioxide was excluded from the solution during the titration. Some information was obtained which seems to indicate that the calcium-sucrose complex actually involves the binding of hydroxyl (or calcium hydroxide a such), rather than involving the dissociation of some hydrogen ion on formation of the complex-which neutralizes an equivalent amount of hydroxyl. The titration of known amounts of hydroxyl from a standard sodium hydroxide solution (the carbonate in the base being removed by barium) by standard hydrochloric acid in the presence of 30% sucrose to a thymophthalein end point gave results nithin 0.1% (relative) of theory. The addition of calcium in the form of a soluble calcium salt to the sucrose solution prior to the addition of the standard sodium hydroxide and titration with standard acid gave similarly correct results. It appears that the solubilization of calcium hydroxide by sucrose solutions involves the tying up of hydroxyl in the complex, so that it reacts only very slowly or not a t all with hydrogen ion.
Table X.
Calcium Hydroxide in Test Hydrates as Determined by Gravimetric Sucrose Method A % 51.7
Av.
51.6 51.4 51.6
B
C
% 51.8 51.8 51 7 51.8
% 48 8 48 8 48 7 48.8
D % 50.7 50 3 50 2 50.4
E
% 51 2 51 2 50 8
51.1
The gravimetric sucrose method was then applied to the test hydrates and the results of individual determinations appear in Table X. Except in the case of hydrates C and D, reasonable agreement between these results and the per cent calcium hydroxide as calculated from the chemical analysis (Table 11) and the per cent calcium hydroxide from the thermochemical data (Table VII) has been obtained. The gravimetric sucrose result for the per cent calcium hydroxide of hydrate C lies almost exactly between the value as calculated from the chemical analysis and the value from the thermochemical value. Hydrate D appears t o be much less uniform than the other samples and tended to give less precise results, a fact which must be borne in mind in making such comparisons. Nevel-theless, in the case of hydrate D, the per cent calcium hydroxide as determined directly appears to be significantly different from the value obtained by calculation from the chemical analysis. The case of hydrate C is less clear-cut and requires further data to arrive a t a conclusion. The gravimetric sucrose method for calcium hydroxide (as well &s calcium oxide) is generally applicable to all types of lime materials and can be expected to give results that are correct to 0.6% (relative) or better, provided that the calcium sulfate content is not more than 1%. A more subtle interference is due to the possible presence of a magnesium carbonate species. Assuming that this species is magnesium carbonate trihydrate, the reaction: bSgC08.3Hz0
+ Ca(0H)z
= hfg(OH)2
+ CaC03 + 3Hz0
tends t o proceed toward the right in aqueous solution, because the products areless soluble than the reactants. This removes calcium hydroxide from reaction with the sucrose and would give low results. The extent of this interference would depend on the
amount of magnesium carbonate trihydrate present and the relative rates of the calcium hydroxide-magnesium carbonate trihydrate and the calcium hydroxide-sucrose reactions. As the calcium hydroxide-sucrose reaction is undoubtedly much faster, this interference would be negligible unless large amounts of magnesium carbonate trihydrate are present. The principa1,drawback of this method is that it is extremely time-consuming. Because of this fact and because it was desired to obtain further information on the correctness of the per cent calcium hydroxide as obtained by direct methods, a volumetric method was tested. IOWMETRIC METHOD FOR CALCIUiM HYDROXIDE
A large number of methods have been proposed for the determination of calcium hydroxide (and calcium oxide) in the presence of calcium carbonate by taking advantage of the greater basicity of calcium hydroxide. However, the extension of these methods to dolomitic hydrates is not always practical because of the magnesium hydroxide (or magnesium oxide) present. Although magnesium hydroxide is much less soluble than calcium hydroxide, it can maintain the pH of the solution a t a sufficiently high value (10.4 is the natural pH of a saturated magnesium hydroxide solution) so that a sharp end point cannot be observed in the acidimetric titration of hydroxyl or other basic ions equivalent to the calcium hydroxide. I n some cases, it may be possible to filter the solution before the final titration, but this detracts from the rapidity and possibly t h e accuracy of the method. Bailar (3) has described a method for the determination of calcium hydroxide and calcium oxide in the presence of calcium carbonate which makes use of t h e following reaction: 313-
+ 3Ca(OH)z = 3 C a + + + 81- +
+ 31120
(1) ,4n excess of standard iodine is added, and the excess is backtitrated with standard thiosulfate. Approximate equilibrium calculations demonstrate that this method would give fairly accurate results for calcium hydroxide in the presence of magnesium hydroxide, provided that the reaction between magnesium hydroxide and iodine is slow. The reaction between iodate (99.8% of the iodine in the standard solution is present as iodate) and hydroxyl occurs in two steps as follows: 13OH- = H I 0 21(2) 3HI0 3 0 H - = 211033Hr0 (3)
+
+
+
the over-all reaction being: 318-
+ 60H-
=
81-
+
108-
+ +
103-
+ 3H20
(4)
This is essentially a reversal of the well-known reaction. 103-
+ 81- + 6 H +
313-
+ 3Hr0
(5)
The equilibrium constant for Reaction 6 from standard potential data is 3.7 x l O S 4 a t 25" C. Combining this value with the ion-product constant of water, the equilibrium constant for Reaction 4 is found to be 2.7 X 1019 a t 26" C. This treatment assumes that the disproportionation of hypoiodite is quantitatively complete and therefore only the over-all equilibrium represented by Reaction 4 need be considered. Calculations based on the standard free energy of hypoiodite and hypoiodous acid show that this is a reasonable assumption. In carrying out the reaction between iodate and calcium hydroxide, a standard iodine solution is used that is 0.2 M in iodate and 0.6 M in iodide. To calculate the completeness of the reaction, let it be assumed that 0.5 gram of calcium hydroxide is taken a t the beginning and that the excess volume of standard iodine solution a t the completion of the reaction is 10 ml. in a total volume of 100 ml. As the calcium hydroxide is completely d k o l v e d a t the completion of the reaction, its solubility product constant does not enter into the equilibrium calculations. Assuming concentrations equal to activities, the hydroxyl ion concentration is found to be 8 X 10-4 M a t equilibrium and the reaction is therefore 99.5% complete. The p H a t equilibrium for this reaction has been measured and found to be 10.3, indicating that the reaction is somewhat more complete than the calcula-
V O L U M E 24, NO. 4, A P R I L 1 9 5 2 tions indicate. Further calculation indicates that any base capable of maintaining the hydroxyl ion concentration a t a value greater than 8 X lo-&M would consume iodine. As the hydroxyl ion concentration of a saturated magnesium hydroxide solution is 2.5 X 10-4, it would be expected t h a t magnesium hydroxide would interfere. However, this interference is largely suppressed until most of the calcium hydroxide has reacted because the hydroxyl in solution from the unreacted calcium hydroxide represses the solubility of the magnesium hydrouide. As the rate of reaction between the calcium hydroxide and the iodine is slow, the accuracy of the method will evidently depend on the determination of the minimum reaction time for the calcium hydroxide reaction, so that the excess of standard iodine can be back-titrated before appreciable reaction Tvith the magnesium hydroxide takes place. The tiiPpropoi-tionat,ioriof hypoiodite to iodide and iodate is slow at room temperature but proceeds rapidly on warming to 80' to 90" C. In Bailar's procedure (3),boiling xvater was added to the sample, the measured volume of standard iodine was added while the solution m-as still hot, and the suspension was shaken and then allowed to stand for 15 minut.es before being backtitrated with standard thiosulfate. As the solubility of calcium hydroxide decreases with increasing temperature while that of magnesium hydroxide increases with increasing temperature, it seemed desirable to carry out the solubilization of the calcium hydroxide a t room temperature, then warm briefly t o complete the disproportionation of the hyp0iodit.e. Furthermore, there is less danger of loss of iodine on warming by t'his procedure, as the small excess of iodine present after completion of t8he calcium hydroxide reaction finds itself in t'he presence of a very large excess of iodide. Experiments demonstrated that this procedure gave better results in the presence of magnesium hydroxide bhan did the method where the iodine is added to a hot suspension of sample. If the warming step is omittctl, t'he results are 1 to 1.5% lorn. The amount of iodine added is somewhat critical. If the excess iodine at equilibrium is less than 5 nil. (of 0.2 M iodate), the reaction tends t o be slow and incomplete. On the other hand, if the excess is of the order of 20 nil. or greut'er, t'here is greater likelihood of significant consumption of iodine by the nlagnesium hydroxide and loss of some iodine on rvarming. An excess of 10 nil. of standard iodine a t equilibrium was found to be optimum. The procedure whereby the excess standard iodine is backtitrated with thiosulfate was found to be inapplicable in t h e presence of magnesium hydroxide, because a t this p H (CR. IO), the t,hiosulfate is partially oxidized to sulfate and the stoichiometry of the back-titration is therefore indefinite. The back-titration was consequently carried out with a standard arsenite solut>ionin the presence of borax (sodium t,etraborate) as a buffer. A large excess of borax was found to be deeirable to provide adequate buffering action and to remove the calcium ion by precipitation as a calcium borate to prevent possible precipitation of calciuni arsenite during the back-titration. A disodium hydrogen phosphate buffer would serve equally well. Sodium hydrocarbonate is unsatisfactory because the hydrogen ion formed by the arsenite-iodine reaction generates carbon dioxide which tends to sweep iodine out of solution. The arsenite solution may be prepared determinately and used to standardize the iodine solution. A 0.2 M iodate solution was used, as it requires a convenient sample size (1 gram) for highly hydrated dolomitic lime. Twice as much iodide was used in preparing this solution as is customary to decrease the volatility of the iodine and eliminate difficulties from this source during the determination. The solution concentration did not change by more than 0.1% per month, so that restandardization a t more frequent intervals is unnecessary. The procedure applicable to highly hydrated dolomitic limes and similar samples in which the percentage of uncombined magnesium oxide is less than 10% is aa follows: The standard iodine solution (0.2 M in iodate and 0.6 M in iodide) is prepared by dissolving 45 grams of reagent grade
699 iodine and 132 grams of reagent grade potassium iodide in 150 ml. of carbon dioxide-free nater, then diluting to 1 liter with carbon dioxide-free water. The solution is stored in a dark bottle. It may be standardized against weighed quantities of reagent grade arsenious oxide by the usual procedure or against the standard arsenite solution. I n any case, the volume ratio of iodine to arsenite should be determined by titration of the iodine by t h e arsenite in the presence of 5 grams of reagent grade borax. The standard arsenite solution is prepared deterniinately by weighing 19.78 grams of primary standard grade arsenious oxide (dried a t 110" C.) and dissolving in 200 ml. of water in which 20 grams of reagent grade sodium hydroxide have been dissolved. Solution is completed by gentle warming, then the solution is neutralized to a phenolphthalein end point with 12 M hydrochloric acid. The solution is cooled and diluted to 1.000 liter in a volumetric flask and mived Tvell. This solution is 0.2000 M in arsenic( 111).
Table XI. Results on Known Amounts of Calcium Hydroxide by Iodometric Method
IO-min. stirring 15-min. stirring 20-min. stirring 5-min. warming
CaCOi CaCOa CaCOs CaCOs hIg(0H)z
Ca (OH); Error Found (Rel.) Gram Gram 00 0.477'7 - 1 . 2 0 . 0 5 0.483: 0 . 0 5 0.4925 0 , 4 9 1 5 -0.2 0 . 0 5 0 , 4 8 3 5 0.4824 - 0 . 2 0 . 0 5 0.4835 0.4790 - 0 . 9 0.5
10 niin. warming None Sone None None Xone
Same a s CaCOs CaCOs CaCOj CaCOj CaCOs hIg(0H)z
4 0.03 0 05 0.3 0.5 0.95 0.3
0.4836 0.4925 0,3546 0.3546 0.3546 0.4925
0.4809 0.4927 0.3532 0.3539 0.3838 0,4927
-0.5 0.0 -0.4 -0.2 -0.2 0.0
11 12= 13a 14
1-hour stirring None Sone Sone
9.4835 0.4992 0.4991 0.4913
0.4856 0.4967 0.4995 0.4926
+0.4 -0.5 +o. 1 10.2
15 16
Xone Sone
Same a s 10 None None 0.05 CaCOI, 0.35 Mg(OH)z 0.05 MgO. Cas04 0.01 Same as 14 CaCOI 0.05 0.3
0 4915
0.4934 0.3671
+O. 4 i3.5
17
Sample boiled m t l i water under re0ux for 30 min.. t h e n iodine. p i c . . added Same a s 17 Sample stood 30 min. with water. then iodine, eic., added 50 ml. of boiling water added i o sample, then stand 30 min.. then iodine, etc., added
Expt.
1 2 3 4 5 6
b
9 10
18a 190
20
Variation in Other Compounds Standard Procedure Present Gram
Ca (OH12 Taken
0 3546
Same as 16 0.3
0.3546 0.3379 0 . 4 9 7 2 0 4630
-4.7 -6.9
Same as 18
0 4988 0.5049
+1.2
MgO
0 . 0 5 0 4915 0 . 4 9 2 9 +0.3 0.5 21 Sanie a s 20 CaCOs 0.05 0.4915 0.4891 -0.5 0 Ca(0H)Z for these experiments obtained f r o m weighed quantities of freshly ignited CaO obtained from pure CaCOa. CaCOs
hlgO
To determine the calcium hydroxide conknt of a highly hydrated dolomitic lime (or similar material in which the amount of uncombined magnesium oxide is less than IO%), sufficient sample to contain 0.5 gram calcium hydroxide (1.000-gram samples of highly hydrated dolomitic lime were used), is weighed and transferred to a 250-nil. glass-stoppered Erlenmeyer flask. The ensuing operations are done as quickly as possible to minimize carbon dioxide contamination. Fifty milliliters of carbon dioxide-free s-ater and 5 nil. of a 1% gelatin solution are added, then sufficient standard iodine solution to provide an excess of 5 to 15 ml. (40 nil. of standard iodine solution were u . d for the 1-gram samples of highly hydrated dolomitic lime) after reaction with the calcium hydroxide is measured into the flask, most conveniently from a buret because of the very slow drainage time of the iodine solution. The flask is then stoppered and the contents are stirred vigorously for 30 rt 5 minutes. The flask is then unstoppered slightly and placed on the steam bath for 10 to 15 minubes. The solution should not be heated above 90" C. during this warming period. The flask is then cooled quickly under running water, 5 grams of rea ent grade sodium tetraborate are added, and the solution is swirfed briefly. The excess iodine is then titrated nit,h the standard arsenite s o h tion until the brown color is faint;. 5 ml. of 0.2% potato starch are added and the titration is contmued to the discharge of the
A N A L Y T I C A L CHEMISTRY
700 blue color. As the reaction is slon- near the end point, care must, be taken to allow each drop of arsenite to react before the next is added. From the volume of arsenite used in the back titration, the total volume of iodine added, the volume ratio of iodine to arsenite, and the concentration of the iodine solution, the per cent calcium hydroxide in the sample is computed. The procedure as described \vas applied to known amounts of calcium hydroxide in the presence of various potential interferences and with several variations of the standard procedure, and the results are shown in Table XI. Zxperiments 1, 2, and 3 indicate that the reaction b e h e e n the calcium hydroxide and the iodine is complete n-ith 15 minutes' stirring; nevertheless, a 30-minut'e stirring time was standardized on to ensure complete reaction in all cases. Experiment 10 shom-s that correct results are obtained in the presence of an equivalent, amount of magnesium hydroxide with a 30-minute stirring time, but if the time is increased to 1 hour, somewhat higher results are obtained, as experiment 11 shows. Esperiment,s 4 and 5 indicat,e that the warming time on the steam bath should be a t least 10 minutes, but should not, be prolonged beyond 15 minutes, as loss of iodine by volatilization or side reaction with magnesium hydroside may occur. Experiments 6, 7, 8, and 9 show that calcium carbonate in amounts ranging from the usual quantities found in dolomit,ic hydrates to quantities equivalent to the calcium hydroxide are without significant effect on the accuracy of the determination. The most serious interference is that of magnesium oxide, as sh0n.n by experiments 16 and 19. This interference \vas found to be of a twofold nature. The small amounts of freshly formed magnesium hydroxide from the hydration of the magnesium oxide (or perhaps the magnesium oxide itself) have a very specific ability to adsorb iodine. Such adsorbed iodine does not give a blue color with starch and it reacts only very slowly with arsenite. l l a g nesium hydroxide prepared hj- hydration of the oxide by boiling and long standing did not sho\y this adsorptive tendency toward iodine. The dfect is evidently one of particle size, the freshly formed niagiiesium hydroxide having a small particle size and a high surface activity. The addition of 5 ml. of a 1% gelatin solution as a protective colloid was found to minimize this adsorption (which actually rrpiwented very little iodine but made the visual detection of the end point extremely difficult). Even in the presence of gelatin, samples containing magnesium oxide in amounts equal to one-tenth the weight of calcium hydroxide did not give a completely colorless solution a t the iodin?-arsenite end point, but retained a faint brovinish tinge. It is for this reason that the starch end point is utilized. The second type of magnesium oxide interference is due to the fact that the magnesium oxide reacts far more rapidly with iodine than does aged magnesium hydroxide, again undoubtedly a particle size and surface activity effect. The magnesium oxide used in these esperiments \vas prepared by calcination of niagnesium hydroxide at 950" to 1000" C,. Lvhich approximates the thermal history of magnesium oxide in dolomitic hydrates. Presumably, the oxide prepared by calcination at, lower temperatures ~ o u l ddisplay even greater activity toTT-ard iodine. Because of this magnesium oxide interference, the iodometric method for calcium hydroxide as described is suitahle only for highly hydrated dolomitic lime and other materials in which the free magnesium oxide content is 10% or less. Experiments 14 and 15 demonstrate that in the presence of the usual constituents and usual amounts in highly hydrated dolonlitic lime, the method gives slightly high results, but the results are correct to 0.4% (relative) or better. I n the absence of magnesium oxide, the method gives results that are usually correct to 0.3% or better, but in any case where the magnesium oxide content is less than lo%, the method can be espected t o give results that are correct to 0.570 or bett'rr. A logical procedure to minimize the interference of large amounts of magnesium oxide would be to hydrate the magnesium oside forcibly and so inactivate it. Boiling the sample with water under reflux for 30 minutes was found to inactivate the magnesium oxide (as shoivn by the lack of adsorption of iodine
after the end point), but apparently also inactivated the calcium hydroxide, as very low results were obtained (esperiments 17 and 18). The sample was protected from carbon dioside contaminat,ion during this treatment. This inactivation occurred even when calciuni oxide was used as the source of the calcium hydroxide, rather than calcium hydroxide itself. Allowing the sample t o stand with water for 30 minutes improved t,he situation somewhat ( a good visual end point was obtained), but the result was significantly high (experiment 19). A set of conditions intermediate between these two extremes was found to give the best results. The sample was treated wit,h 50 nil. of boiling water and allo~vedto stand for 30 minutes prior to addition of the iodine, then \vas treated in the regular manner. Experiment 20 shows that the accuracy is about the same as that obtained wit,h samples in which the initial amount of free niagnesium oxide is one-t,enth the calcium hydroxide (experiments 14 and 15). In this case, an excellent end point was obtained. Esperiment 21 shows that the boiling water treatment may have inactivared the calcium hydroxide slightly, but the result is not significant,ly low when compared with experiments 7 and 12. There is probably some compensation of errors, the small incompleteness of t'he iodine-calcium hydroxide reaction balancing out the side reaction between magnesium oxide and iodine. It is therefore indicated that the procedure should be modified for samples containing large amounts of free magnesium oxide (as normal dolomitic hydrates) by t'reating the sample with 50 ml. of boiling water and allowing it to stand for 30 minutes without further heating, then adding the gelatin and iodine, and proceeding in the prescribed manner. .kny soluble or sparingly soluble h>.droside or carbonate will interfere in this method. Magnesium carbonat,e trihydrate will int,erfere by the metathetical reaction previously described. Calcium sulfate is, of course, n-ithout interference in any amount. On considering possible interferences present in burnt lime, hydrated lime, and similar materials, the method is seen to be specific. I t does not differentiate between calcium oxide and calcium hydroxide. The behavior of niagnesium oxide and magnesium hydroxide in this met,hod is further evidence of the enormous variation in reactivit). that these tn.0 species can exhibit and points out the anomaly t,hat fairly hard-burned magnesium oxide, which is resistant to hydration under normal conditions, can be active enough to interfere where aged magnesium hydroxide does not interfere.
Table X11. Calcium Hydroxide in Test IIydrates as Determined by lodometric Method
7c
D 70
54
51.9
49.1 48.8 48.5
50.0 49.6 49.5
51.3 51.1 51.1
o
48 8
49 7
51 2
B
51.4 51.3 51.3 51.3
52 0 52.0
/c
A,..
C
6
1,'
52
E
The iodometric method for calcium hydroxide, as described, was nest applied to the test hydrates and the results of individual det,erminations appear in Table XII. Hydrates A , B, and E: all showed good precision, with hydrates C and D somewhat poorer but not excessively so. The average of the iodometric results agrees with the average of the gravimetric sucrose results (Table X) to 0.6y0or better for hydrates A (0.370 lower), B (0.2% higher), C (identical), and E (0.1% higher), while the iodometric results for hydrate D (0,7yolower) agree with the gravimetric sucrose results to 1.4%. The excellent agreement obtained between t,he two methods for four of the t,est hydrates is strong evidence that both methods give correct results for the per cent calcium hydroxide in highly hydrated dolomitic lime m-ithin the probable experimental errors indicated, since the principal interferences with the two methods are different.
V O L U M E 24, NO. 4, A P R I L 1 9 5 2 Table XIII.
701
Summary of Analytical Data B
A
C
D
E
Ca(0H)z Computed from chemical analysis X-ray diffraction (order) Thermal analyzer (order) Thermal Sucrose Iodometric
52.1(1) 1 2 1 51.7(1) 57.7(1) 51 6 (2) 5 7 . 8 (1) ,57.3(2) 52.0(1)
51.9(2) 4
49.6(5) 2
51.9(2) 57.2(4) 3 5 4 3 47.6(5) 50.2(4) 50.5(3) 48.8 ( 5 ) 5 0 . 4 (4) 51.1 (3) 48.8(5) 49.7(4) 57.2(3) 5
for magnesium hydroxide as desrribed. Considering the overall agreement between the values for the per cent calcium hydroxide and per cent magnesium hydroxide as calculated from the chemical analysis and as determined by direct methods for the test hydrates, the validity of computing the percentage composition of highly hydrated dolomitic lime of the northn-estern Ohio type from its chemical analysis hy making the1 usual assumptions is established. ACKNOWLEDGMENT
hIg(OH)z Computed from chemical analysis (free water 1200)
Same (free water Z50°) X-ray diffraction Thermal analyzer Thermal
33.6(5) 3.1 8 ( 5 ) J
3 34.6(5)
36.4(4) 38.4(3) 40.2(1) 3 6 . 4 (4) 3 7 . 7 (3) 3 9 . 8 (1) 2 1 4 1 4 2 3 6 . 7 ( 4 ) 3 9 . 3 ( 2 ) 41.1(1)
39.1(2) 38.3( 2 ) 3 3 38.9(3)
The significmtly higher sucrose rekult for hydrate D compared to the iodometric values is probably due to the fact that the per cent calcium sulfate in hydrate D is three to five times that of the other test hydrates. COMPARISON O F R E S U L T S BY VARIOUS M E T H O D S
A summary of analytical results for calcium hydroxide and magnesium hydroxide in the five test hydrates includes t'he qualitative observation of order of intensities in s-ray patterns and thermal analyses curves givrn in Table S I I I . Comparison of the calcium hydroxide as determined the sucrose, iodometric, and thermochemical methods n-ith the per cent calcium hydroxide as calculated from the chemical analysis s h o w that the agreement among the four values for hydrates A, B, and E is satisfactory, the niasimum deviation between rstrcme values for a p:trticular hydrate being 1.4'%. In the case of hydrate C, the sucrose and iodometric results are in agrceinent and are 1.670 lower than the chemical analysis value. The thermochemical value for hydrate C, which is probably in error, is 2.4% l o v x ~than the iodometric arid sucrose values and 4y0lon-er t.han the chemical analysis value. The sucrose, iodometric, and thermochemical values for hydrate D are in fair agreement, the deviation for extreme values being 1.4%, hut these values are significantly lower than that calculated from the c~hemicnlanalysis. The chemical analysis of hydrate D was c:ircfully rechecked and a search \\-as made for significant aniounts of undetermined constituents as a mcnns of establishing the reason for this large deviation, but no explanation was forthcoming from this. Both hydrates C and D displayed poorer Iircvision in their chemical analyses than did the other indicating less homogeneity; nevertheless, the difthe calcium hydroxide results in the case of hydrate D is too large to be accounted for on this basis, and no exp1an:ition for this difference has been established. The fact that the values for the per cent calcium hydroxide by the direct mrthods are generally l o w r than the value calculated from the chemical analysis can be explained by the, fac.t that' calcium may tie associated with the silica and alumina (as dicalcium silicate or trind this has not been taken into account in calculation. .lis convincingly proved in extrnded researches b ~ -Brocard (j), these compounds tend to hydrolj-ze in water, the calcium then reacting as calcium hydroxide, but the hydrolytic reactions are slow a t room temperature, so that not all the calcium tied up in this way would react. In support of this idea, it may be not,ed that the hydrat,es having the smallest percentage of silica and R,Oi (B and E ) shoived the least deviation between the chemical analysis value and the values by the direct methods. The agreement between the t'hermochemical value for the per cent calciurn hydroxide and the values obtained by the iodometric and sucrose methods for hydrates A, B, D, and E is evidence supporting the validity of the thermochemical method
The authors wish to express their thanks to R. E. Grim and associates at the Illinois State Grological Survey and to J. A. LIurray and associates at t.he Massachusetts Institute of Technology for the preparation of thr differential thermal analysis curves used in this research. LITERATURE CITED
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