STUDIES O S T H E HTDROLYSIS O F COMPOUSDS K H I C H l L 4 Y
OCCUR IK PORTLAND CEMENT' BY W M . L E R C H A N D R. H. BOGUE
The reactions occurring when portland cement is gauged with water are complicated because of the presence of a number of constituent compounds, each of which undergoes chemical change by interaction with water. These reactions are involved in the processes of setting and hardening. Hydrolysis and hydration both assume important roles in these processes. A given compound of portland cement may interact with water by the process of hydrolysis or by the process of hydration, or both, but the rate with which these reactions proceed in the case of any compound appears to be dependent on several factors. Inasmuch as the value of a cement in a mortar or concrete seeins to be concerned with the rate with which these reactions proceed, and the nature and amount of the end products formed, it becomes imperative that we understand these processes and conditions of equilibrium if the greatest advantage is to be taken of the possible means for controlling them. In the development of an orderly program for the study of these reactions, it is necessary to differentiate between hydrolysis and hydration and to study each separately. Otherwise the underlying causes of the several reactions are not entirely clear. It is also necessary to investigate first the pure constituent compounds of cement and to determine their behavior under conditions that are made arbitrarily ideal. There is no immediate relation between these conditions and those of cement practice but the former must be determined before the latter can be sufficiently explained. I t should be borne in mind however that it is not permissible to assume that, because e . g. tricalcium silicate hydrolyzes to monocalcium silicate h>drate in a given time in a certain large amount of water, it also will proceed to the same point in the same time when present in the concentrated solutions from freshly mixed concretes and in the presence of the many soluble substances of the wet mixture. This report is a preliminary survey concerned only with pure cement compounds and not at all with commercial portland cement. ' I here are considered : The equilibria conditions in the reactions of hydrolysis. (I) (2) the rate with which these reactions proceed under various concentrations of hydroxyl ion, and (3) the nature of the end products resulting under the several conditions imposed. Published hy permission of the Director of the Kational Bureau of Standards, V. S. Department of Commerce. Paper No. I I of the Portland Cement Association Fellowship a t the Bureau of Standards.
1628
WM. LERCH AND R. n. BOGCE
These have been studied independently with each of several compounds which have been reported to exist in portland cement clinker.
Preliminary Discussion Since the terms “hydrolysis” and “hydration’] are sometimes used interchangeably, it is desirable to define each. By hydrolysis is meant such reactions of a compound (salt type) with water that a part or all of the compound becomes decomposed either completely into the corresponding acid and base or incompletely, forming an acid or basic salt or a derivative of the same. Thus the complete hydrolysis of 3Ca0.A120dwould give: 3Ca0.A1203 X H Z O ~ ~ C ~ ( OA1203.aq. H ) ~ The partial hydrolysis of 3CaO.SiOz could give: jCa0.Si02 x H 2 0 ~ C a 0 . S i 0 2 . a q . aCa(OH)2. By hydration is meant the direct addition of the elements of water as: CaO H 2 0 e C a ( O H ) 2 or 3Ca0.A1203 x H ~ O ~ 3 C a O . A l 2 0 3 . a q . The reactions of hydrolysis are not to be considered as unique or different from those of the electrolytic dissociation of neutral salts except in the degree to which the dissociated ions of the salt combine with the H’ and the OH’ ions of water to form undissociated acid or base. Indeed it has been suggested that the term hydrolysis might be omitted altogether since no break is required in the thermodynamic considerations of the equilibria, expressed in the Mass Law, which apply to the development of an excess of H’ or OH’ ions as a result of the dissociation. The term however has found a distinct place in the literature and will be employed primarily to differentiate this type of reaction with the ions of water from the reactions of hydration. The major constituents of cement are apparently compounds formed from weak acidic oxides,-silica, alumina, and ferric oxide,-and a relatively strong base, calcium oxide. Under the influence of water, such compounds would be expected to undergo hydrolysis with the formation of salts of lower basicity (or even, in the limiting case, of the free acidic components as hydrous silica, hydrous alumina or hydrous ferric oxide) and the free base, Ca(OH)2. Aqueous solutions of cement compounds are found to be alkaline in reaction. Since such reaction is determined by the ratio of the ionization constants and the concentrations of the respective H’ and OH’ ions, it follows that Ca(OH)z possesses either a higher ionization constant than the acidic components] or that its concentration in solution is much greater, or both. I n either case, hydrolysis has occurred, and in accordance with the mass law, will proceed until an equilibrium is established or until the available solvent is removed. In order to make clear the ionic relations involved in the hydrolysis of cement compounds, the following hypothetical reactions are presented. These represent possible stages in the hydrolysis of jCa0.Si02, and are offered only in an attempt to analyze previously obtained experimental data, and to furnish the premises for the present study.
+
+ +
+
+ +
HYDROLYSIS OF COMPOUNDS I S P O R T L A S D C E N E S T
I
3Ca0.Si02a-gCa" 6H20-60H'
t4
1629
+ SiOj""" + 6"
tl
3Ca(OH)? HsSi05 (Si02.aq.) 2.
+
Ca2SiOj" 3CaO.SiO2-Ca" 2 H 2 0 a 2 0 H ' zH'
tl
+
Ca(OH)? zCa0.Si02+-Ca" zH20+z0H
ti
Ca(OH)* 4.
ti
+
CazSiO, H 2 0 (zCa0.Si02.aq.)
+ CaSiO," ' + 2H'
11
+
CaSiOa HzO (CaO.Si02.aq.)
+
CaO.SiO1=Ca" Sios" n H 2 O ~ z o H ' zH'
tl
+
tl
Ca(OH)2 H2Si03 (SiOs.aq.) Equation No. I represents the several equilibria which would be expected to obtain under conditions where the gCaO.SiO2 were caused to proceed towards complete hydrolysis. I t is indicated in the report which follows. however, that intermediate products are formed, and that the hydrolysis will not go to completion unless the soluble Ca(OH)* is continuously removed. Equations 2 , 3) and 4 represent possible stages in the hydrolysis wherein one, two and three mols, respectively, of CaO are removed. Reaction 3 would be expected to begin before the completion of reaction 2 and reaction 4 before the completion of reaction 3. Still other intermediate reactions are of course conceivable. Assuming, for the present, that the reactions as given represent the sum total of the hydrolysis reactions, a consideration of the mass law would require that each proceeded until its respective equlibriuni condition were realized. Thus, in reaction z 1 nCa0.Si02 aq. is one of the products, and in reaction 3, this product is dissociated into Ca0.Si02 aq. If zCaO.SiO2.aq. were the least soluble in any given set of conditions, or produced the lowest ion concentration (solubility product) of any of the products of dissociation, it would form in largest amount when the components represented were brought together. Rut if CaO.SiO2.aq. mere less soluble, or produced a lower ion concentration under the conditions of the experiment, this mould be formed in larger amount. Again, by reaction 4, the CaO.SiO2.aq. would be dissociated to Si02.aq if ~
1630
K M . L E R C H AND R.
n.
BOGUE
the latter were the more insoluble or the least ionized under the given conditions. In each of these cases, however, the concentration of the Ca and OH ions determines the position of the equilibrium. This may be seen better by writing one of the reactions in a different form. Equation J may be written: Ca0.Si02.aq. xH20+Ca(OH)2 SiOn.aq.
+
+
any increase in the C'a(OH)* in solution must result in a decreas8 in the Si02.aq.in solution and a consequent forcing of the reaction to the left, until the constant is restored to the original value, 12. From the above analysis of the hydrolysis reactions of jCaO.SiOn,it is apparent that the reaction product may conceivably consist of a combination of the several intermediate and end products mentioned (or perhaps others) and that the position of this equilibrium condition depends, among other factors, on the concentration of Ca(OH)zpresent. X corollary of this conclusion will indicate that if a sufficiently high concentration of Ca(OH)2is introduced at the start, it may serve almost completely to prevent the progress of the reaction, or even, in some cases, to reverse the hydrolysis. For example, aCaO.SiOz might conceivably be caused to proceed to the production of the more basic 3CaO.SiOz. zCa0.Si02 Ca(OH)2=3CaO.SiO* H20 I n this case the zCaO.SiO2would function as an acidic component. It should be possible, in view of the above, to find a concentration of OH ion at which each of the cement compounds will be in equilibrium with the solution, provided this theoretical concentration is within the range of available values. If a basic solution is added to a solid portion of one of the compounds, and the pH of the solution is thereby raised, hydrolysisistakingplace, the compound acting basic with respect to the solution. If, on adding a more basic solution, the pH is lowered as a result of the reaction, combination is taking place, the compound acting acidic. At some intermediate OH ion concentration, there will be neither hydrolysis nor combination. This is the equilibrium concentration. The problem is further advanced by a study of the speed of the reactions under various concentrations of hydroxyl ion, and by a study of the nature of the intermediate and end products formed under these various conditions. These several problems constitute the subject of this report.
+
+
Preparation of the Cement Compounds The compounds studied in this investigation are as follows. j-3 calcium aluminate, jCa0.3hl2O3, Tricalcium aluminate, 3CaO.Al2O3, Gamma dicalcium silicate, yzCa0.Si02, Beta dicalcium silicate, P2Ca0.Si02, Tricalcium silicate, 3Ca0.SiOz, Dicalcium ferrite, aCaO.FezO3
163I
HYDROLYSIS O F COMPOUNDS IW PORTLASD CEMENT
Rankin2and Bates3 state that these include themajorconstituentsof awellburned portland cement clinker, making up over 90 per cent of the product. These compounds were prepared from calcium Carbonate, alumina, silica, and ferric oxide of the composition given in Table I. TABLE 1 Compositions of Haw Materials4
Calcium Carbonate
Alumina
Si02 Ah03 Fe203
nil 0.08
98.48
FeO CaO MgO
55.73
0.01
0.01
Alkalies SO3
c1
co2 Loss on ignition
0.01
0,002
nil -
Silica
Ferric Oxide
98.75
0.25
0.61
-
0.31
99.02
-
-
0.30 -
0.20
-
-
trace trace
-
44.10
-
-
-
1.42
0 . I2
0.59
0.04
-
trace -
Calczuin Alummates.- The calcium aluminates were made by mixing calcium carbonate and alumina in the proper proportions with water and molding into sticks. These were then heated in an updraft, gas-fired kiln. The temperature was slowly raised to 1350’~and maintained at that temperature 20’ for four hours. In each case the product was ground, molded, and heated a second time to insure homogeneity. On examination under the microscope6 each of the calcium aluminates was found to be practically pure. They were found to contain no free lime by White’s test,’ or by the ammonium acetate titration method.*
*
Gamma Dzcalczum Stlzcate.-Gamma dicalcium silicate was prepared by mixing calcium carbonate and silica in the proper proportions and molding with water into sticks. These were heated to 1500’ and held a t that temperature for about z hours. The gas was then turned off and the sticks removed from the furnace. These sticks, which were very hard at high temperatures, crumbled completely on cooling. This crumbling or “dusting” is due to the J.Ind.Eng.Chem.,3,zr1(1911). Concrete-Cement Age (C. M . S.) 2,3 (1913). hnalyses by H.C. Stecker. Temperatures thruout the paper are expressed in degrees Centigrade. 6 All microscopic examinations were made by F. W. Ashton. A H. White: J. Ind.Eng.Chem., 1, 5 (1909). Cf.Lerch and Bogue: Ind. Eng. Chem., 18,739(1926).
1632
WM. LERCH AND R. H. BOGUE
inversion of beta dicalcium silicate to the gamma form, with an accompanying increase in volume of about I O per cent.s The product was of such fineness that 95 per cent passed a IOO sieve, and about 85 per cent passed a zoo sieve. Beta Dicalcium Silicate.-In the preparation of beta dicalcium silicate, it was necessary to prevent the inversion of this compound to the gamma form, Bates and Klein'O had found that the addition of small quantities of boric oxide or chromic oxide would entirely prevent this inversion of the dicalcium silicate to the gamma modification. I n the preparation of beta dicalcium silicate, 0.5 per cent of boric oxide was added to the mixture. The sticks, in this case, were heated to about 1500' and maintained a t that temperature for about three hours. They were removed from the furnace while hot, and quenched in air, since rapid cooling is a further aid in the prevention of dusting. Examined under the microscope, the product was found to be practically pure beta dicalcium silicate. Tricalcium Silicate.-A special procedure was necessary in the preparation of tricalcium silicate. It was found that when all of the calcium carbonate, required to give the proper composition, was added before the first heating, it was impossible to attain complete combination at the temperature obtainable (approximately 1550'). This may be due to the formation of crystals of high-burned lime before the latter enter into combination with the silica. These crystals, when once formed, are very inactive. In the method finally adopted,11 the calcium carbonate was added in increments, and the mixture heated after each addition. The first composition was approximately 2 . 3 CaO:ISi02, the second ~.6Ca0:1SiO2,the third 2.9Ca0:1SiO~. This last product wm then analyzed and the amount of calcium Carbonate, necessary to bring the composition to that of tricalcium silicate, was added. Each mixture was heated to I ~ O O ~ - Iand ~ Zmaintained ~ ~ at that temperature for three hours. The final product examined under the microscope was found to be practically pure tricalcium silicate, with a trace of beta dicalcium silicate but no free lime. Dicalcium Ferrite.-In the preparation of dicalcium ferrite, the sticks, composed of the proper proportion of ferric oxide and calcium carbonate, were heated to and held at a temperature of 13jo~-1400'for three hours. I n this preparation it was necessary to keep below 1436' as the compound dissociates at that temperature.12 Each of these compounds was grGund to such fineness that I O O per cent passed through a 2 0 0 sieve. The Rate of Hydrolysis A . Hydrolysis when the Soluble Products are not removed.-The first experiment involved a determination of the rate of hydrolysis at 30' of the several cocstituent compounds when these were placed in a quantity of water Day and Shepherd: J. Am. Chem. SOC.,28,1089(1906). Bureauof StandardsTechnicalPaper No. 78 (1917). l1 First applied by W. C. Hansen in thia laboratory. I2 SosmanandMerwin: J. Wash. Acad. Sci., 6,15 (1916). Io
HYDROLYSIS O F COMPOUNDS IN PORTLAND CEMENT
I633
sufficientlylarge that all of the Ca(OH)*contained in them would, if liberated, &ill fail to saturate the solution. The OH ion concentrations resulting in each case were determined at definite time intervals until no further change in pH with time was observed. These determinations served, t,herefore, both as a measure of the rat’e of hydrolysis, and as a measure of the total hydrolysis which had taken place. I n these experiments, 2 0 0 cc. of water, free from carbon dioxide, were used, and such a quantity of each compound as would result in the formation of 0.29 gr. Ca(OH)2 if the lime which each contained were converted entirely into the hydroxide. Pure Ca(OH), was used in the same amount in 2 0 0 cc. water as a standard of comparison. The reactions were carried out in tightly stoppered flasks to prevent contamination with the carbon dioxide of the air. The flasks were shaken from time to time to prevent the compounds from setting. The OH ion concentrations of the resulting solutions were determined by the electrometric method with a saturated calomel half-cell, platinum-platinum black electrodes, and a Type K potentiometer. Results, reported in terms of pH, are given in Table I1 and plotted graphically in Fig. I .
TABLE I1 Hydrolysis of the compounds when the soluble products are not removed Time OH ion concentration of resulting solutions expressed as pH in Cs(0H) 3CaO.cCiO. @Ca0.8i02y2Ca0.6iO2 3Ca0.A120, gCa0.3Ab03 zCaO.Fe.Os h urs I rz.1011.j6 11.08 10.91 11.43 10.25 11.70 3 12.28 11.62 11.16 10.98 11.44 10.42 11.89 11.92 11.46 10.82 11.33 11.07 24 12.2911.86 11.42 96 12.29 11.95 11.00 11.93 11.46 11.19 11.45 11.24 11. I O 192 12.28 1 2 . 0 1 288 1z.281z.00 11.47 11.27 11.93 11.45 11.12 11. I 2 11.47 11.28 384
From these results, it is evident that, under the conditions of this experiment, the amount of lime liberated by hydrolysis of tricalcium silicate and tricalcium aluminate is greater than that liberated by the other compounds, even though the total amount of lime present in the system is the same in each case. The calcium aluminates react with water very rapidly and reach an apparent equilibrium in about one day, while the calcium silicates and dicalcium ferrite require 8 to 1 2 dags to reach this condition. The “lime liberated by hydrolysis” is not necessarily the total lime brought into sohtion, but only that part of the total lime as may be regarded as free calcium hydroxide. Lime in the form of calcium salts is not included by this term. B. Hydrolysis when the Soluble Products are remoced.-It has been found in the experiments just described that the compounds examined when placed in water, undergo hydrolysis, some of the lime being split off and going into solution as calcium hydroxide. When the solution is allowed to remain in contact with the solids, as above, equilibrium is eventually established be-
1634
WM. LERCH AND R . H. BOGUE
tween the solid phase and the calcium hydroxide in solution. Further experiments were undertaken to determine the nature of the hydrolysis which would result on the removal of the calcium hydroxide as it was formed in the solution.
nm
I"
Hour
FIG I Hydrolysis of cement compounds as measured by pH at varying time intervals, the total Ca(OH)* content in each case being o 145 per cent.
Molar equivalent weights of the three silicates were used, ( 5 grams of jCaO.SiO,, and 3.67 grams of pzCaO.SiO, and of yzCaO.SiO2). These amounts were employed because of the very rapid liberation of one mol of
T , ~ .I n
nap
FIG. 2 Per cent of CaO removed from 3 CaO. SiOz,0 2 CaO SiOz and y z CaO. Si02 at various time periods by intermittent extraction with water removed every 2 or 3 days.
.
CaO from the 3Ca0.SiOz. The subsequent hydrolysis of the three residual silicates will then be somewhat more directly comparable as they will contain more nearly equal percentages of CaO. These were placed individually
HYDROLYSIS O F COMPOUNDS I N PORTLAND CEMENT
163 5
in one liter quantities of distilled water and the flasks shaken several times a day to stir the contents and to prevent setting. Every second or third day 900 cc. of the supernatant solutions were siphoned off, and replaced by 900 cc. of distilled water. The amounts of lime removed in this manner were determined by titrating with standard hydrochloric acid aliquot portions of the solutions removed. The solutions contained a small amount of silica or silicates in each case. Table I1 gives the amount and the percentages of the total lime removed from each compound at various periods of time up to 1 1 2 days, and Fig. 2 shows the percentages of the total CaO removed in solution plotted against the time. The hydrolysis of tricalcium silicate proceeded so rapidly a t first that in four days the amomt of lime extracted waa more than a third of the total lime present in the original compound. This is equivalent to about one mol. After this period, the extraction proceeded much more slowly. The rate of extraction of lime from beta dicalcium silicate slightly exceeded in this exTABLE
111
The Extraction of CaO from the Lime Silicates, at Room Temperatures 3Ca0.Si02
Vsed 5 gr. 3.68 gr. CaO Days
Grams % of total CaO CaO removed removed
8zCa0.Si02 #.3a0.SiOz* Used 3.67 gr. 2.39 gr. CaO Used 3.67 gr. 2.39 gr. CaO Grams CaO Days removed
0.8170 I ,3620 I . 5636 1.7139
22.2
2
37.0 42 ' 5 46.5
4 6 9
1.8444 ' 9695 2.0874 2.1882
50.0
I1
1
53.5 j6.8 59.5
I3 16 18
2.2818 2.3529 2.4186 2.4807
62.0 64.0 65.7 67.4
23
29 31 33 35
2.5293 2.5797 2.6202 2.6508
68.8
37 39 41 44
2.6913 2.7300 2.7606 2 ' 7983
73.1 74.2 75.0
2
4 6 8 IO
13 15
I7 20 22
24 27
j0,o 71.2 72.0
j6.0
Days same as B2Cas.Si02.
20
25
27
% of total CaO removed
Grams % of total CaO CaO removed removed
0.1197 0.2106 0.3015 0.3924
5.0
0 .I 107
8.8 12.6 16.4
0.1863 0.2592 0.3312
4.6 7.8 10.8 13.8
0,4833 0,5688 0.6543 0.7389
20.2
23.8 27.4 30.9
0.4041 0,4752 0.5481 0.6201
16.9 19.9 22.9 25.9
0.8199 0.9063 0.9828 1.0557
34.2 37.9 41.1 44.2
0.6912 0.7659 0.8370 0.9081
28.9 32.0 35.0 38.0
29 3' 33 35
I.
I259 1673 I . 2366 1,3077
47.0 48.8 51.6 54.7
0.9792 I ,0040 1.1124 I . 1844
40.9 43 ' 7 46.5 49.5
37 40 42 45
I ,3680 1,4391 I . 4922 I . 5624
57 ' 2 60.2 62.5 65.3
I.
2501 1.3122 I , 3689 I . 4283
52.3 54.8 57.3 59.7
I.
1636
WM. LERCH AND R. H. BOGUE
TABLE I11 (continued) ,6128
67.5
I . 4823
I . 6740
70.0
1.7352 I . 7856
72.6 74.8
5426 1 ' 5984 I . 6542
56 59 61 63
1.8378 1.8981 1.9431 I . 9962
76.9 79.3 81.3 83.5
84.9 85.7 86.7 87.5
66 68 70 73
2,0574 2.1078 2 . I510 2.1987
86.1 88. I 90. I 91.9
I . 9818 2.0187
79.7 81.5 82.9 84.5
3.2537 3.2843 3.3149 3.3455
88.3 89.3 90. I 90.9
75 77 80
2.2293
82
2.2653 2.2806
93.3 94.1 94.8 95.4
2,0493 2.0790 2.1078 2.1330
85.7 86.9 88.2 89.3
84 86 88 91
3.37'6 3.3959 3.4184 3 4409
91.6 92.3 92 ' 9 93 ' 5
84 87 89 91
2.2959 2.3049 2.3103 2.3157
96.0 96.4 96.7 96.9
I537 I753 2.1888 2.2014
90.2 91 .o 91.5 92. I
93 95 98
3 ' 4643 3.4841 3.5066 3.5264
94.2 94.7 95.3 95.8
94 96 98
2.3193 2.3247 2.3301 2.3337
97.1 97.3 97.5 97.7
2.2131 2.2248 2.2465 2.2573
92 ' 7 93 ' 1 94.0 94.4
3.5426 3.5579 3.5732 3 5876 3.6011
96.3 96.7 97.1 97.5 97.9
2.3382 2.3427 2.3463 2.3493 2.3523
97.9 98.0 98.2 98.4 98.5
2.2636 2.2699 2.2798 2.2870 2.2930
94.7 94.9 95.3 95.6 95.9
46 49
76.9 78.0 78.8 80.0
47 49 52 54
I
53
2.8343 2.8703 2.9063 2.9441
56 58 60 63
2.9828 3.0134 3 ' 0503 3.0881
81.1 81.9 82.9 83.9
65 67 70 72
3.1223 3 ' 1565 3 . I925 3.2231
74 77 79 81
51
IO0 I02 105
107 I09 112
'
IO1
103 105 I 08 IIO
I12
2.2500
I.
1.7073 7613
I .
1.8117 1.8612 1
'
9044
I . 9458
2. 2.
62 . o 64.6 66.9 69.2 71.4 73.7 75.8 77.9
periment that from the gamma form, but both of these were far below that for tricalcium silicate at the early periods. After I I 2 days the titration curves indicated that from 95 to 98 per cent of lime had been extracted from the several calcium silicates. At this point the remaining residue was filtered off, ignited and analyzed. The results are given in Table IT'. TABLE IV Compound
3 C a O . SiOn P2CaO. SiOz
72CaO. SiOz
Total Residue o 6466 gr. 0.5040 0.6010
SiO,
0.6202 gr. 0,4955 0,5882
La0
0.0240 gr. 0.0085 0.0154
HYDROLYSIS O F COMPOUNDS IS PORTLAND CEMENT
I637
Thus the remaining residue was found to be almost pure hydrous silica with only a small amount of lime. There were originally present about 1.3 grams of SiOn;the difference between this value and the amount found in the residues is to be accounted for in part by loss through solution and in part by mechanical loss on removing the solutions. A most interesting observation obtained by an inspection of the curves is that the rate of extraction of the dicalcium silicates is nearly a straight line function of the time for about two months, but that after about 30 days (See also Table 111) this rate is greater than that of the tricalcium silicate. After months there is very little difference in the three curves. I t appears to be impracticable to measure the rate of extraction of the calcium aluminates and the dicalcium ferrite in the above manner because of the solubility of the alumina and the ferric oxide in the lime solution and because of the presence of colloidal suspensions. Some of the alumina and the ferric oxide would doubtless be in molecular combination with the lime in one or more of the several possible intermediate hydrolytic products, some would be ionized in the solution, and some perhaps would be uncombined. There seems to be no reliable manner, in this case, for determining the partition of the compounds among these several products. It certainly would not be permissible to assume the total CaO content of the solution to be free Ca (0H)z resulting from hydrolytic cleavage. For these reasons a method involving the determination of the total line in solution would not be a measure of the rate of hydrolysis. The end products of the hydrolysis, however, have been determined in another manner. Small quantities of 5Ca0.3.&03, 3Ca0.;11203 and aCaO. Fe203were placed separately in Soxhlet extraction thimbles. These were suspended in separate flasks of distilled water and the solutions replaced by fresh quantities of water at frequent intervals. At the end of five months all of the lime had been extracted from the 5-3 calcium aluminate, but it required eight months to complete the removal of all of the lime from the tricalcium aluminate. The residue remaining in the thimbles a t the end of these periods consisted, in each case, only of hydrous alumina, A1203.aq. I t required but three months to complete the hydrolysis of the dicaleium ferrite, whereby a residue was obtained consisting entirely of hydrous ferric oxide, Fe203.aq. Hydroxyl Ion Concentration of Hydrolytic Equilibria A s stated in the preliminary discussion, one of the objects of this inves-
tigation has been the determination of the hydroxyl ion concentration at which each of the several compounds neither combines with water in the reactions of hydrolysis, nor combines with base in the formation of compounds of higher basicity. This OH ion concentration would be expected to be different for each compound. The experimental method consisted in noting the change in pH resulting upon adding the compounds to solutions of calcium hydroxide and sodium hydroxide. These will be treated separately.
1638
WM. LERCH AND R. H. BOGUE
A . Equilibria in Solutions of Calcium Hydroxide.-The pH of equilibrium of the several compounds was determined by placing 4 grams of the compounds in 2 0 0 cc. of calcium hydroxide solutions of varying concentration from zero to saturated. The use of 2 0 0 cc. of solution gave an ample quantity for pH determinations, and 4 g. of the compounds would be sufficient to make a measurable change in hydroxyl ion concentration of the solution if the solution and the compound were not in equilibrium. The samples were placed in an air bath maintained at 30' by means of a Beaver regulator.13 The pH determinations were made a t the same temperature. It was not necessary in in each case that sufficient time be allowed to attain complete equilibrium, but only that time be allowed for a partial reaction to take place which would indicate the direction of equilibrium. At the end of a week, pH measurements were made on the original solutions and the solutions in contact with the cement compounds. The results are given in Table V. When the three calcium silicates were placed in unsaturated solutions of calcium hydroxide, an increased pH always resulted, indicating that these compounds hydrolyzed with the formation and ionization of calcium hydroxide. When these compounds were placed in saturated calcium hydroxide, there was still a slight increase in pH. I t seems necessary to conclude that these silicates continue to hydrolyze even in a saturated solution of calcium hydroxide. This is further demonstrated by the formation of crystals of calcium hydroxide on the sides of the flask. In the case of tricalcium silicate these crystals appear within 2 4 hours, both when placed in distilled water and in saturated lime water. When the dicalcium silicates are placed in saturated calcium hydroxide, the crystals appear after about I O weeks, and when these compounds are placed in distilled water, 4 grams in 2 0 0 cc., the crystals appear in about 3 months. The formation of the calcium hydroxide crystals apparently is due to the continued hydrolysis of the calcium silicates, even in saturated lime water, with the resultant formation of Ca(OH), which crystallizes from the saturated solution. The increase in pH which is noted in these samples over and above that obtained in a saturated solution of precipitated Ca(OH), at the same temperature is of interest. Since crystals of the latter are present in many cases, being deposited on the sides of the flask, there cannot exist a permanent supersaturation, but a temporary supersaturation may account for the difference noted. Control tests have demonstrated that there is no appreciable reaction of the calcium hydroxide on the glass (Pyrex) and that the high pH values are not due to such reaction. I t might be expected that an apparent increase in the OH' ion concentration could result through the effect of the calcium silicate on the activity concentration of the calcium hydroxide. The degree of the observed increase however seems to be greater thah would be expected by the very small conl3
J I n d Eng. Chem , 1 5 , 3 5 9 (1923)
H’iDROLYSIS OF COMPOUNDS I S PORTLASD CEMENT
Ibjl)
1640
WM. LERCH A S D R. H. BOGUE
centrations of the salt in solution. A further indication in favor of the view of supersaturation was furnished by analysis. A portion of the solution was separated from the tricalcium silicate after seven days, evaporated to dryness and analyzed for lime and silica. It contained 0.00~5gram Si02 and 1.286 grams CaO per liter. h saturated solution of calcium hydroxide a t the same temperature contains I . I 6 grams CaO per liter. The other products of hydrolysis of these calcium silicates have not yet been definitely determined. Le ChatelierI4 believed that the hydrolysis continues until a hydrated monocalcium silicate (Ca0.Si02.z+H20)is formed, that this is the principal constituent of hardened cements. On the other hand, Kewberry and SmithL5found that one sample of tricalcium silicate, after shaking with water, gave a varying residue of 1.5 to z equivalents of CaO to I of SiO2, and a fused sample gave a hydrated silicate of the composition 2Ca0.Si02.7H20. Rankin l6 suggested that gelatinous hydrated silica might be released when the finely powdered tricalcium silicate is mixed with water, and that it is this silica which serves as the binding agent in concrete. I n studying the hydration of beta dicalcium silicate, Bates and KIein'O found that this compound undergoes hydrolysis, resulting in the formation of crystals of calcium hydroxide large enough to be visible to the naked eye a t the end of six months. We have also found this to be true of both the beta and gamma dicalcium silicate. Thus it seems that under these conditions a hydrated dicalcium silicate cannot be the final product of hydrolysis of the calcium silicates. I n our studies of the hydrolysis of the calcium silicates it was found that when the soluble products are not removed no one of the calcium silicates will give up all of the lime in combination. In each case an equilibrium was established between the calcium hydroxide in solution and some less basic calcium silicate. On the other hand, a sample of freshly precipitated hydrous silica which had been placed in saturated calcium hydroxide solution for twelve months was found to have been changed, so that 97 per cent of the Si02 was acid soluble. The soluble residue had a composition 1.07 CaO:ISiOz. This is evidence that a calcium silicate is formed by the interaction of calcium hydroxide in solution upon hydrous silica. It would not be expected, therefore, that the hydrolysis of a calcium silicate would proceed to the formation of any large amount of hydrous silica under conditions where an excess of calcium hydroxide is present. The 5-3 calcium aluminate was found to undergo hydrolysis in all solutions of pH I I .6 or less but in solutions of p E I I .8 or more this aluminate enters into a reaction with the solution which causes the pH to decrease. This results from the combination of lime with the 5-3 calcium aluminate. This reaction also has been demonstrated in another manner. A quantity of the 5-3 calcium aluminate was placed in a flask and covered with a satBull., 42,82 (191x3). J . SOC.Chem. Ind., 22,94 (1903). '6Trans.FsrsdaySoc., 14,23 (1918-19). l4
'5
1641
HYDROLYSIS OF COMPOUNDS I N PORTLAND CEMENT
urated solution of calcium hydroxide. A thimble containing Ca(0H)z was suspended in the solution, the flask carefully sealed and set aside for I O months. The thimble with the remaining insoluble material waa then removed and the residue in the flask filtered off and analysed. The analysis of this residue was as follows: Grams
Mols
A1203
. '37'
I
CaO
,2177
2.90
TheRe experiments show that the gCa0.3A1203combines with Ca(0H)Z. The first experiment shows this to take place when the calcium hydroxide is present in such amounts that the solution is brought to a pH of 11.8or higher. An equation may be written: gCa0.3A1203 4Ca(OH)* aq.S3(3CaO.Al2O3.aq). From the molar ratio of 2.90 CaO to I Alz03obtained in the last experiment, it appears that the compound 3CaO.AlZO3.aq.constitutes the major product in the de-hydrolysis of the 5-3 aluminate, when this reaction is carried out in the presence of a saturated solution of calcium hydroxide. I t should be emphasized however, that the nature of the hydrolysis precludes the probability of any one compound being formed as an end product, to the complete exclusion of other compounds of different basicity. This will be apparent upon a consideration of the reactions suggested for the silcates given earlier in the paper. Tricalcium aluminate undergoes hydrolysis when placed in calcium hydroxide solutions of pH 12.2 or less, and the resulting solution has a higher pH. But in no instance did this hydrolysis proceed far enough to produce an OH ion concentration equivalent to that of a saturated solution of calcium hydroxide. On the other hand when this compound is placed in saturated calcium hydroxide solutions, the pH of the resulting solution is always lowered slightly. Kuhl and Thuring," and LafumaI8 who have observed a similar phenomenon, believed the change to be due to the formation of a tetracalcium aluminate, but RadeffIg attributes the change to adsorption of calcium hydroxide by the tricalcium aluminate. That there is no appreciable quantity of tetracalcium aluminate formed under these conditions has been shown by a different means. Four grams of tricalcium aluminate were placed in a flask containing a saturated solution of calcium hydroxide. A Soxhlet thimble containing 0.83 gr. of CaO was then suspended in the solution and the flask sealed. After 15 months the flask was opened and the thimble with the remaining lime removed. A portion of the material on the bottom of the flask was found on analysis to contain: 0.1260 gr. A1203 0.2093 gr. CaO
+
17 18 '9
Zement, 13,243(1924). Le Ciment, 1925,175. Zement, 14,177(1925).
+
1642
W M . LERCH AND R. H. BOGUE
This is equivalent to a composition 3Ca0.A1203and gives no indicatioi. of a reaction of lime with hydrated 3Ca0.A1203to form qCaO.Al2O3.aq. The dicalcium ferrite is found to hydrolyze in all concentrations of calcium hydroxide, which are less than saturated, but appears to be unchanged in saturated Ca(OH)2. B. Equilibria in Solutions of Sodium Hydroxide.-The hydroxyl ion concentrations obtainable with calcium hydroxide are limited by the low solubility of lime. I t seemed desirable to measure the equilibria of these compounds in more basic solutions, and accordingly sodium hydroxide was used. All samples were treated in the manner previously described for calcium hydroxide solutions. The results of this study are given in Table VI. The calcium silicates were found to react much the same in sodium hydroxide solutions as in those of calcium hydroxide. They all were found to undergo hydrolysis in the less alkaline solutions, producing an increase in pH and, eventually, a formation of calcium hydroxide crystals on the sides of the flask. However, as the concentration of the sodium hydroxide solutions increased, the pH changes decreased. This would be expected on the basis of the hydrolytic equilibria, for as the OH ion concentration is increased, the hydrolytic reaction potential will be decreased, and, at some definite value, will reach zero. At a pH of 13.7 there appeared to be no hydrolysis whatsoever. This was further borne out in this experiment by the failure of crystals of Ca(OH)2to appear on the sides of the flask even after eight months. The 5-3 calcium aluminate reacts as with calcium hydroxide solutions. At pH values below 11.6 hydrolysis takes place, but at pH 11.8 and above the pH of the sodium hydroxide solutions is decreased. This indicates that combination is being effected between the compound and the base resulting in the formation of a more basic aluminate. In this case, sodium appears to substitute for calcium. TABLE
Change in pH of constituent compounds of cement in solutions of sodium hydroxide pH of solution after 4 grams of cement compounds PH had been allowed to react with solution for one week original sodium hydroxide solution 3Ca0.Si02 paCa0.Si02 ynCaO.SiO2 ,?CaO.Al2O3 sCaO.lA1 On ~ C a 0 F e . 0 ~ 13.72 13.72 13.74 13.71 13.72 13.55 13.72 13.01 12.95 12.52 13.00 13.03 13.01 12.94 12.67 12.63 12.33 12.65 12.58 12.64 12.59 12.03 12.40 12.37 12.61 12.28 12.20 12.25
11.93 11.53 I O . 19
12.58
1z.15
12.07
12.20
11.79 11.63 11.07
The tricalcium aluminate also reacts in a similar manner as with calcium hydroxide solutions at pH of I 2 .zand below. I n these ranges hydrolysis takes place.
HYDROLYSIS OF COMPOUNDS I N PORTLASD CEMENT
I 643
At pH values higher than 12.6, however, no change is observed. This indicates that hydrolysis can occur only below pH 1 2 . 2 and that above this value there is no combination with sodium hydroxide t o form a more basic aluminate. The dicalcium ferrite seems to hydrolyze in sodium hydroxide solutions which are below pH 13.0 but above this no change is observed.
Discussion and Summary The experiments which have been described have been concerned primarily with the reactions of hydrolysis under several conditions, but the reactions of hydration proceed at the same time and cannot well be separated completely from the former. I n all of the experiments studied, the hydrolyses have been brought about in the presence of quantities of water or solution much greater than can be present in a properly tempered mortar or concrete. The reason for this lies in the difficulty or uncertainty of following the course of the hydrolytic reactions except by examination of the resultant solutions. The latter procedure offers, at least, the most obvious method for such a study. With the data herein obtained, however, it is hoped to proceed to a more direct examination of the limits reached in the hydrolytic cleavage occurring during the processes of setting and hardening in the cements of commercial mortars and concretes. In the experiment on the rate of hydrolysis when the soluble products were not removed, and the water was present in excess of that required to dissolve all the CaO, several points are of interest. I. All of the six compounds studied undergo hydrolysis and give up Ca (OH), to the solution but none of them give up all of their CaO to the solution, regardless of the time period. A condition of equilibrium appears to be approached in the case of each compound, whereby the lime in solution becomes sufficient greatly to retard or even to prevent further hydrolysis. 2. The 3Ca0.Si02 and the 3CaO.ALO3 give up by hydrolysis more lime as calcium hydroxide to the solution than the other compounds, and of these two, the 3Ca0.Si02gives the greater amount. Xext in order are p nCaO.SiO2, jCa0.3A1~03,yzCaO.Si0, and lastly nCa0.Fez03. 3. The calcium aluminates reached a condition of apparent equilibrium in one day, tricalcium silicate in eight days, all of the others attained this condition by the twelfth day. A point of interest is that the dicalcium silicates required only a little more time to reach this condition than wasrequired of the tricalcium silicate under the conditions of this experiment. I n the experiments on the rate of hydrolysis when the soluble products m e removed, it is brought out that : 4. The 3Ca0.Si02 loses a much greater percentage of its total content of CaO in periods up to about eighteen days than do the dicalcium silicates. The loss in the former case is very rapid at first, and becomes slight at later periods, while with both of the dicalcium silicates the loss is almost a straight line function of the time for about two months. Thus after eighteen days the rate of loss of CaO from these latter exceeds that from the 3Ca0.Si02. The
1644
WM. LERCH A N D
R. H. BOGUE
tricalcium silicate loses one mol (one third of its CaO) in three to four days, and z mols in twenty-four to twenty-six days. The beta dicalcium silicate loses one mol (one half of its CaO) in thirty-two days and the gamma form in thirty-six days. I t may be pointed out that here again the dicalcium silicates show a lower rate of hydrolysis than the tricalcium silicate only in the early periods. 5 . The end product of the above reactions (the soluble products being intermittently removed) appears to be hydrous silica and not a silicate. Thus at the end of about a month all of the silcates have been reduced to the composition CaO.SiOz.aq. but the hydrolysis proceeds continuously beyond this point. 6 . The rate of hydrolysis of the aluminates and the ferrite could not advantageously be studied by this method, but the end products bydialysis, with the dialysate intermittently removed, have been found to consist of hydrous alumina in the case of the two aluminates and hydrous ferric oxide in the case of the dicalciuni ferrite. 7 . The time periods required to effect complete hydrolysis under the conditions of this experiment were approximately : gCa0.3A1203, 5 months 3Ca0.A1203, 8 months zCa0.Fez03, 3 months The experiments on the hydroxyl ion concentration of hydrolytic equilibria in solutions of calcium hydroxide and of sodium hydroxide showed several points of interest. 8. None of the three calcium silicates are in a condition of initial equilibrium in solutions of calcium hydroxide of any concentration. This is indicated not only by the increase in H ion concentration of calcium hydroxide solutions upon the addition of these silicates, but also by the continued deposition of crystals of calcium hydroxide in their solutions. I n solutions of sodium hydroxide the calcium silicates reached a pH of hydrolytic equilibrium at 13.7, which value was further confirmed by the failure of the silicates in such a solution to deposit crystals of Ca(0H)Z. 9. The end products in the hydrolysis of the silicates when the products are not removed, but in the presence of an excess of water probably consist essentially of CaO.SiOs.aq. and Ca(0H)Z. This conclusion is arrived at by deduction rather than by direct experiment. Thus both 3Crt0.Si02 and zCaO. SiOz (both forms) deposit crystals of Ca(OH)2 from solution, showing that they hydrolyze, even in saturated lime water, to a composition of lower basicity than zCa0.Si02. However, hydrous silica is found to combine with Ca(OH)2 when these are brought together. Hence some intermediate product between nCaO.SiOz and SiOz.aq. must be the most stable compound of silica in the presence of saturated Ca(OH)zand an excess of water. It seems probable that this most stable compound is CaO.SiOz.aq., but it will be recalled from the preliminary discussion that no one product can be expected to result
HYDROLYSIS OF COMPOUNDS IS PORTLAND CEMEXT
164j
:it equilibrium to the complete exclusion of all other hydrolytic products. Hence there probably results in saturated Ca(OH), solution an equilibrium of silicate hydrates of low basicity. IO. The pH of hydrolytic equilibrium of jCa0.3A1203 is about 11.7 as determined bot,h in cnlcium hydroxide and in sodium hydroxide solutions. The jCa0.3A1203 hydrolyzes in Ca(OH)2solutions only when the pH is 11.6 or less. At alkalinities of pH 11.8 and higher this aluminate cornbines with some of the CaO to form a more basic aluminate. Additiona1,evidence obtained by allowing SCa0.3A1?03 to remain in a solution of Ca.(OH)2containing an excess of the base, indicates that the major product of the reaction of this aluminate in saturated Ca(OH)2solution is 3Ca0.A1203.aq. The pH of hydrolytic equilibrium for the 3Ca0.A1203is about 12.3 11. as determined both in calcium hydroxide and in sodium hydroxide solutions. At OH ion concentrations of pH 1 2 . 2 or less hydrolysis takes place. At OH ion concentrations of saturated Ca(OH)2solution, about pH 12.4, the introduction of the 3CaO.&o3 lowers this value slightly. This may be due to the formation of a more basic aluminate (4CaO&0,) or perhaps could be explained on the basis of a decreased solubility of Ca(OH), in the presence of the tribasic aluminate. The zCa0.Fe203hydrolyzes in all solutions of Ca(OH)2that are less 12. than saturated. In saturated solutions t,here is no change in pH but this does not determine the position of the equilibrium. In solutions of sodium hydroxide, the pH of hydrolytic equilibrium is found to be about 13.0.
Conclusions The hydrolyses of several compounds which may occur in portland cement clinker have been investigated under various conditions. These compounds include gCaO.SiOz, PzCa0.Si02, yzCa0.Si02, jCa0.3A1203, 3CaO.h1203 and aCao.Fe~O3. The conditions under which these were investigated were made arbitrarily ideal in this preliminary study in order that the fundaniental reactions of hydrolysis might be examined with the least possible interference due t o uncontrolled physical factors or foreign salts. It is believed that: I. the direction and relative rate of these reactions have been established in each cwe, 2. the OH ion concentrations required to prevent hydrolysis have been ascei-taincd, 3. the nature of the probable end products has been determined throughout the possible range of concentrations. All of the above experiments have been performed on pure materials and no attempt has been made in this study to learn the effect on the reactions due to t'he presence of more than one compound or to foreign materials. ?io attempt can be made at this time to correlate these data with the conditions met in cement mortars and concretes. The essential conclusions may be set down as follows: I. Each of the compounds studied was found t o undergo hydrolysis as determined by the liberation of hydroxyl ions in aqueous solutions.
1646
WM. L E R C H AND R. H. BOGVE
2. The relative rate of hydrolysis under the conditions studied, as measured by the pH of the solutions, indicates that the aluminates come to equilibriiim more rapidly than the silcates and the ferrite. 3. The OH io11 concentration necessary to prevent hydrolysis is different for the several compounds, as follows:
PH gCaG.3A120a 3Ca0.A1203 2CaO Fe20a 3CaO.SiOq) /32Ca0.SiC2) yzCa0.Si02)
II .
;
12.3
13 . o
1 ~ 37.
The pH of saturated Ca(OH)? at 30' is 1 2 . 4 . Hence the aluminates do not hydrolyze in such solutions, the ferrite hydrolyzee slightly, and the silicates hydrolyze to a greater estent. I.. If the soluble products are intermittently removed, the silcates will hydrolyze eventually to hydrous silica, the aluminates to hydrous alumina, and the ferrite to hydrous ferric oxide. In each case Ca(OH)? is the other product. 5 . If the soluble products are not removed, hydrolysis will proceed to the formation of Ca(OH), and a less basic silicate, aluminate or ferrite. ThP composition of the end products in this case will be determined by the OH ion concent,ration and the amount of available water. In the presence of a large amount of saturated lime water, the silicates seem to go largely to a composition which may be represented by CaO.SiO?.aq., and the aluminates to a composition 3CaO.Al1O3.aq. Bureau of Ctandards, January 21, 1927.
