FURTHER PHASE STUDIES OF SULFAMATES IN AQUEOUS MEDIUM

Citation preview

INFORMATION TO USERS

This material was produced from a microfilm copy of the original document. While the most advanced technological means to photograph and reproduce this document have been used, the quality is heavily dependent upon the quality of the original submitted. The following explanation of techniques is provided to help you understand markings or patterns which may appear on this reproduction. 1. The sign or "target" for pages apparently lacking from the document photographed is "Missing Page(s)". If it was possible to obtain the missing ' page(s) or section, they are spliced into the film along with adjacent pages. This may have necessitated cutting thru an image and duplicating adjacent pages to insure you complete continuity. 2. When an image on the film is obliterated with a large round black mark, it is an indication that the photographer suspected that the copy may have moved during exposure and thus cause a blurred image. You will find a good image of the page in the adjacent frame. 3. When a map, drawing or chart, etc., was part of the material being photographed the photographer followed a definite method in "sectioning" the material. It is customary to begin photoing at the upper left hand corner of a large sheet and to continue photoing from left to right in equal sections with a small overlap. If necessary, sectioning is continued again — beginning below the first row and continuing on until complete. 4. The majority of users indicate that the textual content is of greatest value, however, a somewhat higher quality reproduction could be made from "photographs" if essential to the understanding of the dissertation. Silver prints of "photographs" may be ordered at additional charge by writing the Order Department, giving the catalog number, title, author and specific pages you wish reproduced. 5. PLEASE NOTE: Some pages may have indistinct print. Filmed as received.

Xerox University Microfilms 300 North Zeeb Road Ann Arbor, Michigan 48106

LD3907 oG7 195 ° •S*+

Seliksorn Bernard, 1922 Purther phase studies of sulfamates in aqueous medium* New Yorkc 1950 1 3 b typewritten loaves* diagrs*, tables* 29cm* Thesis (Ph°D.) - New York Univer­ sity, Graduate School, 1950 ® Bibliography: p*l33 -i3o*

C57606

c Shelf List

Xerox University Microfilms, Ann Arbor, Michigan 48106

THIS DISSERTATION HAS BEEN MICROFILMED EXACTLY AS RECEIVED.

X1BBABJ OF HEW TORI DNIV1R8IW DHIV1RSITY HEIGHTS’

Further Phase Studies of Sulfamates In Aqueous Medium

A dissertation in the department of chemistry submitted to the faculty of the Graduate School of Arts and Science of New York University in partial fulfillment of the requirements for the degree of Doctor of Philosophy

by Bernard Selikson 1950

The author wishes to express his gratitude toward the person whose direction and inspiration were ever present during the conduction of this work - Professor John E. Ricci

To my wife and to my parents without whose encouragement and assistance this work would never have been completed.

Table of Contents Introduction

5

Summary of Literature

6

General Discussion

15

Experimental Preparation of Salts

39

Methods of Analysis

48

Solubility Measurements

61

The System NaS - NaCl - H 20 at 25 °

65

The System NH 43 - NH4C1 - H 20 at 25 °

70

The System NH 43 - (NH4)2S0 4 - H 20 at 25 °

73

The System KS - KC1 - H 20 at 25 °

77

The System KS - KI - H 20 at 25 °

80

The System KS - K2S0 4 - H20 at 25 °

83

The System NH43 - KS - H20 at 9 °, 25 ° and 45 °

8b

The System (NH4)2S0 4 - K2S0 4 - H20 at 25 ° .

103

The System NH45 - K2S0 4 - H 20 at 25 °

106

Summary

132

Bibliography

133

Introduction Although sulfamic acid and its salts are finding increasing use both commercially and in the laboratory, relatively little information is published concerning this family of compounds.

The purpose of this research

is to extend the information, particularly with regard to the aqueous solubility of the salts in the presence of other salts, and also in the presence of their hydrolysis product (namely sulfate) in order to obtain a clearer picture of the purification process and in order to classify better the behavior of the salts, thereby obtain­ ing by comparison with other similarly behaving systems, more information concerning the nature of the sulfamates. There is presented here a report on the aqueous solu­ bility relations at 25°C. of the ammonium, potassium and podium sulfamates in the presence of the corresponding chlorides; the system potassium sulfamate - potassium iodide and water; and as part of the investigation of the reciprocal salt pair ammonium sulfamate - potassium sul­ fate in water at 25° (which is also reported here), the ternary systems ammonium and potassium sulfamates with corresponding sulfates and water (all at 25 °), and at 9 °, 25° and 45° the system ammonium sulfamate - potassium sulfamate - water.

Summary of Literature Sulfamic acid is reported^-»^ ^ and was found4- to be a colorless, odorless, non-volatile, non-hygroscopic solid, forming rhombic plates5 , melting at 205° with decomposition (Oberhauser and Urbina^ claim that it re­ quires 209 ° to decompose the melt). gravity in toluene of 2.126 at 25°

It has a specific and in ether of

2.03 at 12° 3 ,5 ,8 ^ yields ammonium acid sulfate by hydrolysis, reacts vigorously with nitrites to liberate nitrogen and sulfuric acid, and also reacts with con­ centrated nitric acid to yield nitrous oxide and sulfuric acid^.

It is reported as being isomorphous with its

potassium salt, as being less soluble than its salts, and as crystallizing better from aqueous solution than most of its salts8 .

It is suggested as a primary standard^

because of its high equivalent weight, (its strength approaching that of hydrochloric and sulfuric acids-*-), and because its hydrolysis would not change the titer of a standard solution.

However, the present writer finds this

suggestion unwise owing to the decomposition of wet sul­ famic acid by heating, the requirement of obtaining sul­ famic acid by desiccation rather than by heating, and the necessity of storing the acid in a dry atmosphere to avoid subsequent hydrolysis.

If a really pure grade of

sulfamic acid were available, the suggestion would be appropriate; as yet, it is not. The solubility of the acid at 25 ° is listed by

7

Cupery0 in various solvents in grams per 100 grams of solvent:

methanol, 4.3 g.; ethanol (containing 2% ben­

zene) 1.7 g«; acetone, 0.4 g.; ether .009 g.; formamide 2 0 .0

g..

It is insoluble in hydrocarbons, chlorinated

hydrocarbons, carbon disulfide, and sulfur dioxide.

The

aqueous solubility is decreased by sulfuric acid and by sodium sulfate to insolubility at 70% sulfuric acid con­ centration® »®, but it then rises again with further in­ crease in acid concentration0 .

The solubility as a

function of temperature is listed per 100 grams of water: 14.08 g. at 0°; 1 8 .

g. at 10°; 21.32 g. at 20°; 2 0 .0 9 g.

at 30 °; 2 9 .4 9 g. at 400; 3 2 .8 2 g. at 50 °; 3 7 .1 0 g. at 00 °; 4 1 .9 1 g. at 70 °; 47»08 g. at 80 °.

are given^:

Two additional values

1 4 .9 4 g. acid per 100 cc of aqueous solution

at 12° and 40.00 g. at 78°; with the statement that the increase in solubility with temperature is practi­ cally linear.

It is further noted that the pH of a 10 $

solution is 1 .8 as determined colorimetrically.

Heating

the solid is inadvisable^ because at 135 ° the water is "fixed", i.e.. hydrolysis takes place (this writer observed this phenomenon at 110°) and at 105° the solid does not become any drier than can be obtained by air drying or by storing over Anhydrone.

The present writer observed

that a saturated solution of sulfamic acid (approx. 2 .5 Molar)

underwent measurable hydrolysis at 25° in several days. This observation is in sharp variance with the findings of Oberhauser and Urbina^-® that a Normal solution of the acid is not hydrolyzed at room temperature over eight months but is hydrolyzed at 97° to tfre extent of 24$ in fifteen minutes and 30 $ in one hour, and with their statement that the degree of hydrolysis is higher in more dilute solutions (and therefore probably lower in more concentrated solutions). They also report^ the preparation of the following sulfamates by reaction with the corresponding hydroxides: aluminum sulfamate, crystallizing with 18 moles of water m.p. 230°; calcium sulfamate tetrahydrate - m.p. 200°; copper sulfamate dihydrate - m.p. 184°; potassium sul­ famate - m.p. 211°; sodium sulfamate - m.p. 250 °; zinc sulfamate tetrahydrate - m.p. 195°; ammonium sulfamate 128 °; ferrous sulfamate with n molecules of water - m.p. 135 °; nickel sulfamate tetrahydrate - m.p. 125°; by re­

action with the corresponding chlorides;

barium sulfamate

(does not melt under 300 °); magnesium sulfamate tetra­ hydrate - m.p. 91°; and by reaction with lead acetate: lead sulfamate dihydrate (does not melt under 300°). They further claim that the lead and barium salts are insoluble in water, a statement which is inconsistent with the properties of barium salt observed by this writer and with the r eported^>^

properties of both salts.

All the inorganic sulfamates reported in the literature

are water soluble (with the exception of a basic mercury salt) and insoluble in alcohol**»H.

BerglundH notes the

existence of the following sulfamates:

ammonium sul­

famate (m.p. s 125 ° without decomposition); sodium sul­ famate (very soluble needles); lithium sulfamate; thallium sulfamate; silver sulfamate (dissolving in 15 parts of water at 19°); barium sulfamate (soluble in 3 parts of water at ordinary temperature); strontium sul­ famate tetrahydrate (more soluble than the barium salt); lead sulfamate monohydrate (most soluble salt); nickel sulfamate trihydrate (very soluble); cobalt sulfamate trihydrate; manganese sulfamate trihydrate (very soluble); cadium sulfamate pentahydrate (very soluble); and copper sulfamate dihydrate (less soluble than others).

He also

notes that the sulfamates of magnesium, aluminum and uranium are very soluble.

Cupery0 lists the solubilities

at 25° of the following sulfamates per 100 grams of water: Ammonium sulfamate - 193 g. (s 65.6$ in weight per cent); sodium sulfamate - 10b g. (* 51 .5$); magnesium sulfamate 119 g.; calcium sulfamate - 67 g.; barium sulfamate 34.2 g.; zinc sulfamate - 115 g.; and lead sulfamate 218 g..

The increasing solubilities of the barium, calci­

um, and magnesium salts expressed in grams of anhydrous salt per 100 grams of water as a function of temperature are listed by King and Hooperl3> but their (interpolated) values differ markedly from those reported by Cupery0 : calcium sulfamate - 78.0 g.; barium sulfamate - 2 9 .7 g.;

10

magnesium sulfamate - 110 g..

The relative viscosities

and densities of aqueous solutions of sulfamic acid and of the ammonium, barium, calcium and magnesium salts over ranges of undersaturation are reported^-4-.

The

partial molal volume of (blue) copper sulfamate (degree of hydration undetermined) in aqueous solution at 25°, is 56.08 ml. per gram molecular weight at zero concentra­ tion and the solubility determined as 74.73 g. per 100 g. water.

Gordon and Cupery^0 note the existence of addi­

tional sulfamates:

potassium sulfamate (m.p. = 192°);

ammonium sulfamate (m.p. = 131 °); dihydrates of sul­ famates of lead and copper; tetrahydrates of sulfamates of calcium, magnesium and zinc; and aluminum sulfamate, containing 18 molecules of water. Butler and Audrieth^ describe the preparation of the ammonium salt by addition of sulfamic acid to liquid ammonia, permitting the excess solvent to evaporate, and recrystallizing the salt from hot ethanol.

Ammonium sul­

famate is also mentioned^ as one of the products of the reaction of ammonia and sulfur trioxide at from 80 ° to 250 °.

Its crystallographic properties are purported to be

reported.

Cupery^ lists the solubilities of this salt in

grams per 100 grams of water as a function of temperature: 134.8 g. (= 57.41 g. per 100 g. soln.) at 0.2°; 166.6 g. (= 6 2 .99 %) at 1 0 0 ; 2 0 0 .2 g. (= 66.69%) at 20 °; 2 3 2 .4 g. (* 6 9 .92 %) at 3 0 0 ; 2 7 9 .5 g. (= 7 3 .65 %) at 40 °; 357 g.

(*» 78.12%) at 500.

Brown and Cox? report crystal axial ratios for sul­ famic acid which are in poor agreement with the values given by MellorS.

Their values for the axial ratios of

the potassium salt are in good agreement with those in Mellorl® but they claim, from X-ray examination, that there is no evidence for hydrogen bonding from nitrogen to neighboring oxygens; whereas Yost and Russell^ remark, from X-fay data20 yielding the same tetrahedral spacing of oxygens and nitrogen around the central sulfur and the same interatomic distances as reported by Brown and Cox, that the ions "undoubtedly form hydrogen bonds from the nitrogen to the oxygens of neighboring molecules". Ricci and fc>elikson4 report the solubilities at 20 °, 23° and 330 of the acid, and the ammonium and potassium salts {mentioned later in this paper; plus the solubility of the monohydrated sodium salt reported there for the first time.

The sodium salt reported in the earlier

literature may have been the hydrate, for its reported solubility® (= 51*5 g. per 100 g. soln. at 25°) is some­ what closer to the value reported for the monohydrate4 (a 55.3 g. per 100 g. soln. at 25°) than the value cal­ culated by interpolation of the anhydrous solubility curve^l down to 25 ° (s 60 g. per 100 g. soln.).

However,

no mention is made in the literature of the striking properties of the hydrate and so it is doubted by the present writer that the hydrate was ever discovered until recently reported^.

The degree of hydration was demon­

la

strated both by direct drying experiments and also by using the indirect Schreinemakers method22 Qf algebraic23 extrapolation of tie-lines, using sodium iodide as the third component.

The system sodium sulfamate - water,

investigated by Laning and van der Meulen21, is presented at eight temperatures, demonstrating the occurrence of an incongruently melting solid undergoing transition at 38.3° and b2.3 g. anhydrous salt per 100 g. solution: 7 9 .9 g- per 100 g. water at 0° (= 4 4 .4 g* per 100 g. soln);

113.0 g.

at 20° (5 53.0%); (55*4# by interpolation at 25°);

1 3 5 .9 g.

at 30 ° (« 57.6ft); 1 5 1 .4 g. at 35 ° (= 6 0 .2ft);

108.6 g.

at 40 ° (* 6 2 .8 ft); 175.1 g. at 45 ° (= 63.6ft);

1 8 2 .2 g.

at 50 ° (* 6 4 .5%); 1 9 1 .3 g. at 55 ° (« 65.6ft).

Although the authors proved that the saturating phase above 38.3° is the anhydrous salt (by following the solu­ bility curve of the salt in the three component system up to 7.5ft added sodium chloride and employing extrapolations of the type mentioned above), they merely proved that at 20° the monohydrate is the stable phase under an aqueous vapor pressure in the neighborhood of 12 mm. and that it deliquesces in an atmosphere of 1 3 *9 mm..

To prove con­

clusively the water content of the saturating phase of solutions of sodium sulfamate below the transition point, the obvious procedure calls for an investigation of a ternary system below the transition temperature in the same manner as was done above it.

The extreme difficulty^

in effecting equilibration of the hydrate with its

13

saturated solution owing to its spongy nature, pronounced supersaturation tendency and sensitivity to slight changes in temperature, probably prevented pursuance of this approach.

Laning and van der Meulen report the long

transparent anhydrous needles to show parallel extinction, and to have a density at 20 ° of 2 .2 1 0 .

The salt is

further reported24 to have a molar volume of 53 *7 .

The

system ammonium sulfamate (m.p. = 132.9°) - ammonium nitrate is reported25 as simple but the data do not correspond at all to the curve shown (because of an error in plotting, according to a private communication).

The

binary system ammonium sulfamate - sodium sulfamate (m.p. * 250.5°) is reported26 with the occurrence of the congruently melting compound 2 ammonium sulfamate • 5 sodium sulfamate melting at 213 - 1 °»

Noting the de­

composition of the ammonium salt above 170° (and that of the sodium salt in the neighborhood of 200°), the authors admit that "an accurate determination of the upper eutec­ tic point was not possible, but it lies near a tempera­ ture of 212° and a composition of 73»0% sodium sulfamate" (7 2 .3% * theoretical percentage in compound).

The

closeness of the "eutectic" to the "melting point", the admitted decomposition of the ammonium sulfamate at this temperature, and the disregard of certain pertinent points just above the "eutectic" leave the congruency of the com­ pound with regard to melting in doubt.

The best that may

be said is that the compound is on or very close to the

14

border of incongruency.

The decomposition of the

ammonium sulfamate yields products which probably lower all the fusion curves and so it becomes impossible to distinguish further.

It is important to note the "ex­

tensive supercooling and exceedingly slow crystalliza­ tion even when the melt was seeded". The binary system sodium sulfamate - sodium nitrate is also reported in this paper2&f over an even narrower composition range because of extensive decomposition at the higher temperatures.

A one to one compound definitely

is observed, melting (with an extremely flat curve) at 205.7° containing 41.65# sodium nitrate (theoretical =

41.08). The first complete ternary isotherm involving sul­ famates to be mentioned in the literature was studied by the present writer4 ;

the system sulfamic acid - ammonium

sulfamate and water at 25° with an invariant point occurring at 6 2 .2# salt and 1 5 .2# of acid.

The system

is simple and is noteworthy mainly as an indication of the feasibility of aqueous equilibrium studies involving sulfamates, even in strong acid solutions, and also be­ cause of the unusual straight solubility curve of the acid which dropped only three and a half per cent (absolute) as the ammonium sulfamate increased from 0^» to 62#.

General Discussion The ternary systems reported in this thesis are salt systems with a common ion plus water and it is of interest to study these systems and those with a higher number of components to obtain information concerning their inter-relations.

At a particular temperature these

studies yield information on the possibility of occurrence of a compound or solid solution between the salts, which has bearing on purification and analytical processes in­ volving these salts; and at various temperatures, the polythermal relations are obtained, with a more complete knowledge of the system.

The addition of a third com­

ponent with the aid of graphical22 or mathematical23 ex­ trapolation of the equilibrium tie-lines enables the com­ position of a binary (or ternary) compound to be in­ directly determined^" and sometimes the (possible) presence of a binary phase undetected in the binary study may be discovered through a ternary study in which system the binary compound is stabilized2^ . The Phase Rule applied to a condensed isothermal2? three component system forbids the coexistence of more than three phases in equilibrium, since F * C - P.

Ex­

amination of a typical isotherm of the simplest type dis­ closes an area of under saturated liquid which is bi­ variant; two areas in which liquid is saturated respec­ tively .with either salt A or B and therefore univariant; and finally an area in which all complexes separate into the invariant liquid and its saturating phases A and B.

16

The invariant liquid is saturated with respect to both A and B and in fact occurs at the crossing of both solu­ bility curves, each formed as more of the other salt is added to the respective saturated binary liquid.

There­

fore, this most concentrated invariant liquid will ex­ hibit the lowest vapor pressure of any liquid in the simplest system and will be the drying up liquid for the system.

That is, all liquids will ultimately yield this

one before the liquid phase disappears, as seen by follow­ ing the transitions occurring as water is removed from any ternary complex. Solid Solution Formation A type of binary interaction of interest here is the formation of an often continuous and therefore homogeneous solution of the component solid phases.

In

17

order for the solid solution to be continuous over the entire composition range, the lattice dimensions of both solids must be in close enough agreement to allow complete interchange

The relative distribution of the two salts

in their solid solution and the aqueous liquid solution (which constitutes a ternary system) gives information on the nature of the solids and of their solid solution, and it is possible to classify these behaviors both for the binary and their related ternary systems.

It is assumed

in all of the systems under discussion here that the solid solutions are homogeneous in composition and uniform throughout the entire solid (sometimes termed a Nernst distribution) as opposed to a non-equilibrium type of distribution in which the composition of the "mixed-crystals" varies as the composition of xhe liquid varied during the formation of the precipitate (termed a iJoerner-hoskins distribution^).

The former type of crystal is formed

either slowly or after considerable “equilibration", i.e., stirring in the presence of the mother liquor and the distribution describes the system under thermodynamic equilibrium.

The latter type of distribution describes the

crystals when first formed and will apply less the longer the crystals are allowed to remain in contact with the mother liquor. Roozeboom grouped the binary systems involving solid solution into five general types31 as presented in Figure the first three of which are continuous, and the latter two

Figure 1

THE FIVE TYPES OF SOLID SOLUTION BINARY

POLYTHERMS

n

T»f« r

T O

l

I

T t-i.

it.

A

%

Z

is

B

TERNARY

ISOTHERMS

m

TERNARY x I

y

ISOTHERMAL m

IT

iv

DISTRIBUTION IV

19

discontinuous.

Each type is derivable from the zeta

"?

p o t e n t i a l o r free energy function) versus composition family of curves, obtained as the temperature is changed®®*31-56#

The free energy of the two continuous

phases solid and liquid are depicted^tnd the changes in their relative positions are followed as the temperatures of the systems are changed.

For systems of type I, the

zeta potentials for both phases are convex toward the composition axis and regular; for those of type II and III, the same, with the curve for one phase more inclined than the other; and for those of types IV and V, they would be irregular with the zeta curve of the solid phase ex­ hibiting a plait, thus giving rise to an immiscibility in that phase. It is possible for systems to show positive or nega­ tive deviations from Raoult's Law (or Retger's Law, in the case of solid solutions) in which event the fusion diagram for the system would tend to be irregular and would in most cases lead to a system of type II or type III respectively as shown in the next drawing of Figure

In

order to relate more clearly those binary and ternary systems which exhibit a behavior resulting from the same underlying phenomenon yviz.. positive deviations for type II and negative deviations for type III), this writer has taken the liberty to alter the numbering of the binary types so that they will correspond to the numbers of their related ternary systems as originally used by Roozeboom^.

20

Thus Roozeboom's binary type III will be called here type II and vice versa; and Roozeboom's binary type V will be called here type IV and vice versa.

In systems

of type IV and V, the tendency toward immiscibility is sufficiently pronounced to cause the binodal curve of the solid - solid equilibria to occur at high enough tempera­ ture to interfere with the solid - liquid equilibria. Type IV is the familiar eutectic type in which liquid may not appear at temperatures below the invariant point, while type V is called the peritectic type in which liquid may exist below the invariant temperature but not below the melting point of the lower melting component.

Type V

can be arrived at from type I as the binodal region of immiscible solid solutions approaches (say, at different pressures) and interferes with the liquid - solid region, while type IV can be arrived at from type II if the maxi­ mum of the binodal curve coincides with the minimum of fusion curve. These binary salt systems form, with water, ternary systems which also fall into distinct

c a t e g o r i e s ^ ,

38 which

categories will now be related to the type of binary system involved.

It is to be noted that in the binary systems

at temperatures at which only solid(s) would exist, types I, II and III are indistinguishable on the ordinary phase diagram (all are homogeneous solid solutions) as are also types IV and V (both involve immiscible solid solutions).

All these types are distinguishable (in a

probable sense) at these lower temperatures when one studies their corresponding ternary systems.

It must be

kept in mind that the binary types refer actually to the fusion behavior of the systems and that therefore, the behavior of the binary system at the (usually) very much lower temperature of the ternary isothermal study need not necessarily be caused by the same forces.

Strictly,

one can only say that if the fusion phenomena of a binary system exhibit positive deviations from Raoultfs Law, then the behavior of the ternary system formed at a lower tem­ perature by this binary system with water will probably exhibit positive deviations too.

Often, as in the case

of the systems reported upon in this thesis, the salts would undergo considerable decomposition at the tempera­ ture of their fusion phenomena, thereby making the resuits of a binary investigation often questionable (refer p.13) if not unattainable.

Since it is important to classify

these ternary systems and also to relate them to the corresponding binary systems in order to obtain a better understanding of the intermolecular forces involved, the five types of ternary systems involving solid solutions will now be discussed and are presented in Figure 2 under the corresponding binary systems of related type.

It

should be remembered that the ternary systems plotted are all at a (much) lower temperature than are those described by the binary plots and also that they are isotherms rather than polytherms.

Two different methods of

22

describing the systems are shown:

the more familiar

phase diagram using an equilateral triangle in which the solubilities (i.e.. concentrations) of each component in each phase are depicted; and the method introduced by Roozeboom37 (and used by Blasdale^®, Hill3 9 } Ricci^®, and others) utilizing a square diagram in which is plotted the proportion, y, of one of the salts (say A) to the total salt present in the liquid solution versus the pro­ portion, x, of the same function in the solid solution. The absolute values of the solubilities can not be rep­ resented by the latter method. The phase diagram for systems of type I shows a con­ tinuous and regular solubility curve across the diagram with tie-lines all regularly displaced.

These lines will

all point to one side of the water apex, to the side of the species whose saturated solution will have the lowest vapor pressure and water activity (in the example given, toward compound A).

Upon drying any saturated solution of

this system, equilibrium solid solution will precipitate which is richer in B than is the liquid.

Therefore the liquid

composition will move toward a saturated solution of pure A; there will not occur a congruent or drying up point within the ternary system.

It is obvious that the value of A per

unit of salt will at all concentrations be greater in the liquid than in the saturating solid solution phase and therefore the distribution curve for this system (presented)

below the phase diagram in Figure 2 ) will progress smoothly and will always lie above the 45° diagonal. The binary system of type II, in which positive deviations from flaoult's Law (and a consequent minimum melting point), occurred, would be expected to lead to the ternary system of type II in the presence of water.

The

positive deviations are in evidence here by the spreading apart, i.e.. divergence of the tie-lines on the(binary) base line, indicating a tendency toward immiscibility of solid solutions and by the bunching of the tie-lines on the liquid solubility curva- indicating a tendency toward an invariant liquid.

Although the solubility curve remains

continuous, xhere is a congruent and drying up point corresponding to a minimum or water vapor pressure.

This

point corresponds in the triangular diagram to the pair of coexisting phases which have a connecting tie-line pointing directly at the water apex, and in the square diagram to the point of crossing of the diagonal.

At

this point, the tendency for the liquid composition is to remain constant as evidenced by a bunching of the tie-lines toward this liquid, while the solid composition varies continuously.

Solids richer in A than this one are in

equilibrium with liquids poorer in A than is the solid (the curve is below the diagonal) and so upon evaporation, the liquid will move toward and stop at the crossing point. The reverse situation holds for solids of smaller concen-

trations of A than that of the crossing point (the curve is above the diagonal) and now, upon evaporation, the liquid becomes richer in A finally stopping at the same point as that mentioned above, since B is being removed from the liquid to a greater extent than is A. Type III binary and ternary systems are those in which negative deviations from Raoult's Law occur.

There

is no change in phase in the binary system and so the (ternary) aqueous solubility curve would also be expected to be continuous.

However there is a tendency toward

compound formation as evidenced by a bunching of tie-lines down to the base in the neighborhood of solid which tends to remain constant in composition. diagram in Figure

Hotted on the square

the distribution curve at low values

of A is seen to lie below the diagonal, indicating that upon evaporation of such solutions, solid which is richer per unit of total salt in A than is the liquid will pre­ cipitate.

The result is that the solution will move toward

a saturated solution of pure B.

The distribution curve

starts out at low values of A below the diagonal, then crosses the diagonal, where the solid composition is seen to vary only little while the liquid composition varies considerably, and finally ends above the diagonal.

Since

the distribution curve at high values of A is seen to lie above the diagonal, the concentration of A per unit of total salt is greater here in the liquid than in the solid solution.

Evaporation of these solutions causes solid to

precipitate which is richer per total salt in B than is the liquid, and conjugate liquids will thus change in composition toward a saturated solution of pure A.

Thus

there is no ternary solution which has a minimum in vapor pressure and so there will be no regular ternary drying up point and no regular congruent point.

Since the dis­

tribution curve (in the square plot) crosses the diagonal, there must be some ternary complex which separates into a solid solution and its saturated aqueous solution, both of which contain the same ratio of salts.

Although

saturated solutions of composition to either side of this one will move away from it upon drying, the tie-line for this saturated solution points directly to the water apex and so no changes in salt ratio occur when this particular (but no other ternary) solution dries up.

Thus there is

seen to occur a fortuitous drying up and congruent point in this system. drying up point.

In fact, any congruent point will be a In this particular type of system how­

ever, this particular solution will exhibit a maximum in the vapor pressure and therefore all other saturated solutions will move away from it; it is not a drying up or crystallization end point. It is to be noted that certain systems when plotted on the square type of diagram may appear to be of type I but actually may have a distribution curve crossing at extremely high concentrations to become type II or III^l. If the repulsive forces in the solid solution become

sufficiently great, the solid will separate to form two solid solutions,

Thus, some ternary systems of type II

may become, with a change (usually lowering) of tempera­ ture, type IV.

The solubility curve is discontinuous

with the occurrence of a liquid of minimum vapor pressure similar to type II but which is now invariant (rather than tending toward invariance) over the range of composi­ tion of two solid solution mixtures.

For the liquid to

be invariant, the coexistence of two solid phases along with the liquid is required by the Phase Rule.

The liquid

is congruently saturated and also is a crystallization end point.

The square diagram (Figure 5 ) shows the distribu­

tion curve crossing the diagonal but with a discontinuous flat portion, indicating the coexistence of two mutually saturated solid solutions one of which contains more of A and the other more of B per unit of total salt than does the liquid. Lowering the temperature of systems of type I may lead to systems of type V in which there occur

(as in

type IV) two solid solution phases, an indication of pro­ nounced positive deviations from Raoult's Law in the solid. The solubility curve is here also broken with the appear­ ance of an invariant liquid but this liquid is neither congruently saturated (and therefore it is not a drying up point) nor is it a crystallization end point for it does not have the lowest vapor pressure of the system. The square diagram for this system also shows a flat

27

portion (similar to type IV) but the entire distribution curve is above the diagonal.

Blasdale^ claims that

systems of this type are very rare but the system with ammonium and potassium sulfamates (reported later) is of this type as possibly are others which are not recognized as such. Quaternary Systems A maximum of six phases may coexist in equilibrium in any four component system, and this coexistence is re­ quired by the Phase Rule for the establishment of real invariance.

Of course, for a condensed isothermal system,

the coexistence of four phases suffices for invariance and will occur at only one point in the (simplest) .system of type a, below, but this "invariant" point will change with temperature and in fact (albeit slightly) with pressure.

As

all the systems discussed here are considered condensed sys­ tems, three solids may simultaneously saturate only one par­ ticular liquid, viz.. the invariant liquid, at any defined temperature.

For salt - water systems, we may consider

two general types:

(a) three salts which have a common

ion and constitute a ternary system; (b) a reciprocal salt pair which undergoes metathesis thereby comprising a ternary system.

Each of these ternary systems forms with

water a quaternary system, the first type being represented by a symmetrical tetrahedron or equilateral triangularly based pyramid (which representation is also used for quaternary alloy systems) and the latter being represented «

28

by a square based pyramid. Since, even after the limitations of an isothermal "condensed system" have been placed upon a quaternary system, the system still requires three dimensions for complete depiction, much of the problem concerns the attempt to describe completely and unequivocally an isothermal condensed four component system (i.e.. a three dimensional figure) in two dimensions.

Some of the more

common methods require two or more separate diagrams.

In

addition to these, it is possible to superimpose two views of the three dimensional model on to one graph and thereby obtain a more compact as well as complete image.

Since a

quaternary system involving a reciprocal salt pair was in­ vestigated during the course of this research, this type of system is used for discussion although everything in general that will be said about this type of system will apply to systems involving three salts with a common ion as well. When considering a reciprocal salt pair, it is con­ venient to express concentrations on an equivalent basis so that the model of the system will be regular, here, a regular square based pyramid.

The corners of the pyramid

represent pure salts at the base and pure water at the apex; the edges, all of unit length, represent binary systems; the faces represent ternary systems; and the solid model represents the quaternary system.

The solid model

may be considered to be latticed in space and then the height of any quaternary point gives the equivalents of

water per equivalent of sample.

The equivalents of salt

per equivalent of sample are obtained by difference.

Thus

a mixture of equivalent quantities of the three salts will be located on the base of the pyramid:

the exact location

is found by measuring, for the two salts at the end of the diagonal, the appropriate (here equal) distance from the face opposite each salt, in a direction parallel to the sides of the figure.

Or the salt compositions may be ex­

pressed in terms of a positive and a negative radical, each side of the square representing the appropriate ion.

A

complex composed of equivalent quantities each of the three salts and water will-be located at a distance up from the base equal to one quarter the total distance from the base to the water apex.

At this level, the salt proportions are

expressed similarly as done on the base of the pyramid. Since, at this level, any complex contains only 75# of salt, the sides of this square will be 75# as long as the sides at the base. Figure 2 is a schematic model of the quaternary system ammonium sulfamate - potassium sulfate - water at 25°C, presented in detail later.

Points C, D, E and F represent »

saturated solutions of the pure salts in water.

These

points are extended to curves in the ternary systems which meet the other curves at the ternary invariant points A, B and G.

These invariant points now extend into the

diagram to become curves which meet at the quaternary invariant liquid q.

A binary liquid (e.g.. E) is

Figure 2

o

31

saturated with respect to a single pure solid, a ternary curve (EA) represents saturation with respect to a single solid while the third component is being added, a ternary invariant liquid (A) is saturated with respect to two solids, a quaternary curve (Aq) represents saturation with respect to two solids as the fourth component is added, and the quaternary invariant liquid (q) is saturated with respect to three solids. The saturation surfaces divide the model into regions of undersaturated and saturated liquid.

Since ammonium

sulfate forms.a continuous solid solution with potassium sulfate, it is seen that the ternary triangular area EAe, in which liquids are saturated with respect to ammonium sulfate, extends continuously through the pyramid, defining a pentahedron, and ending with the area FBf in which liquids are saturated with respect to potassium sulfate.

The top

surface of this pentahedron EFBqA represents liquids which are saturated with respect to a single solid phase of continuously variable composition.

Along the opposite side

of the diagram, however, it is seen that the solid solution between ammonium and potassium sulfamate is limited.

Thus

the corresponding triangle CAc, in which liquids are saturated with respect to ammonium sulfamate extends only as far as qGx to form a smaller pentahedron.

The surface

ACGq represents solutions saturated with respect to a single solid, the composition of which varies continuously but only as far as point x.

This pentahedron is separated

from the adjacent pentahedron by a tetrahedron.

In this

adjacent pentahedron, the saturating phase, which starts at y and ends at pure potassium sulfamate, saturates liquids which are found at the surface GDB.

In the intervening

tetrahedron, liquid saturated with respect to solid solu­ tions x and y, changes in composition from & to q, as the fourth component is added.

The rest of the diagram is

composed of two irregular pentahedral pyramids, Aecxzq and qzydfB, which are again separated by a tetrahedron.

How­

ever, the tops of these pyramids will not be points but will be the curves Aq and qB respectively.

These curves

represent liquids saturated with respect to a sulfate and a sulfamate solid solution, ammonium rich for curve Aq, and potassium rich for ourve qB.

Finally, the intervening

tetrahedral volume qxyz completes the model and represents the quaternary invariant liquid q, saturated with respect to the three solid solutions x, y and z. Thus the schematic model presented in Figure 2 of the condensed isothermal system of the reciprocal salt pair plus water is composed of the following volumes:

a tri­

variant region of undersaturated liquid; three bivariant regions of liquid I, II and III, each saturated with respect to a single but variable solid solution; three univariant regions of liquid Aq, Bq and Gq, each saturated with respect to two solid solutions; and finally an invariant region of liquid q, saturated with respect to the three solid solu­ tions x, y and z.

33

If the solid solution of the sulfamates were continuous ov.er the entire range, the regions of saturation could he divided into three groups:

the first and third regions

would be inverted pentahedral pyramids pointing down to the line (rather than point) of the saturating solid phase of variable composition; the intervening pentahedral pyramid would point up to the curve of solution, saturated with respect to both solids as the compositions of all phases vary continuously across the diagram.

For the system pre­

sented here, however, one of the solid solutions is dis­ continuous, and so the second and third pentahedral pyramids are each interrupted by a tetrahedron, in each of which the invariant liquid, q, occupies one corner. We may now consider the problem of the useful rep­ resentation of this three dimensional model in two di­ mensions, i.e.. the projection of these saturation curves and surfaces onto the base of the model. For the projection of the saturation surfaces of the system from above onto the base of the diagram, two essentially different kinds of projection are generally used, each of which produces a different result than the other.

One type of projection, called an orthogonal or

orthographic (or Schreinemakers) projection, involves parallel lines of projection and tends to preserve the shape or perspective of the original surface.

The other

type, which is known as a radial or clinographic or Janecke type of projection (after the originator of the

method44^ although Le Chatelier proposed the idea in­ dependently^ )^ involves lines of projection which radiate from the water apex, which therefore are not parallel, and which tend to produce an image which is larger than the original surface.

If these surfaces were represented merely

by their limits, which in the three dimensional model could be shown as wires, and if the entire model were otherwise "open", or transparent, then the shadow of these wires cast by a light from above will give the projections mentioned above.

It is obvious that there can be only two positions

for the light source which will give useful "shadows", viz.. the light is held at infinite distance away to produce parallel lines of projection (i.e.. the orthographic type) and the light is held at the water apex to produce lines of projection which radiate from the apex and which are inclined away from each other (i.e.. the clinographic or radial projection).

On the base of Figure 2 there are

shown both the orthogonal and Janecke projections, in that order, with small primed and unprimed lettering,respectively which correspond to the capital lettering of the solid model In the orthographic projection, then, concentrations expressed in terms of equivalents of salt per equivalent of saturated solution are plotted on the base.

The corners

of the projected -figure represent the pure component salts, the center represents pure water (lying, of course, above the saturation surfacesj, and four lines may be drawn connecting the corners to the center and dividing the base

35

into four triangles,

compositions of saturated binary

and ternary liquids may be found at the outer limits of the projected surface and may be read back from the curve on one of the dividing lines or in the appropriate triangle formed by these lines.

However, compositions

of saturated quaternary liquids found within the projected surface can not be read back from the figure since each point is determined by the difference in the concentrations of the reciprocal pair of salts concerned. In the Janecke method, the projected surface occupies the entire base since the view point is the water apex. Concentrations of saturated solutions are expressed in terms of equivalents of salt per equivalent of total salt in the solution (rather than per equivalent of solution) or in terms of equivalents of a positive and of a negative salt ion per equivalent of salt in the solution.

The four

corners of the square here represent equivalent salt pro­ portions of saturated binary liquids, the edges represent equivalent salt proportions of saturated ternary liquids and the points inside the square represent equivalent salt proportions of saturated quaternary liquids.

Here, the

salt concentrations in all the saturated liquids may be read back from the projection of the saturation surface. Since the lines of projection radiate from the water apex, it is not possible to represent pure water and in fact it is impossible to depict differences in water concentration. This is a consequence of the manner in which concentrations

are expressed:

equivalents of salt per equivalent of total

salt, water being excluded.

The water concentration may be

expressed as equivalents of water per equivalent of total salt and may then be plotted as an elevation projected planes corresponding to (separately) onto/two faces, or two diagonals, or a face and a diagonal to give a more complete picture of the system. Utilizing these elevations, it is possible to determine the composition of a solid indirectly by extrapolation of these (twice) projected tie-lines.

The method is equally appli­

cable to itregular based figures47 as well as regular ones. Elevations of the same height may be connected in a three dimensional model elevated over the base to give planes of equal water content known as isohydrores.

The curves formed

as these planes cut the saturation surfaces would form a family of contours describing the spacial relations of tbs system, which contours could also be projected to the base of the diagram. Blasdale has plotted on one triangular49 (and also square50) diagram both a Janeeke and a Lowenherz projec­ tion of a quaternary system.

By the Lowenherz method48t

salt concentrations are plotted in an open model as equiva­ lents per thousand mols of water and the resulting figure is projected down orthogonally.

It also is possible to use

the Schreinemakers with the Janecke projection on one diagram as has been suggested recently51.

it is obvious

from .Figure 3 that, as long as the units used are not confused, any orthogonal projection may be used in con-

>

Figure 3

ff

j7

v

38

junction with a radial projection to obtain complete in­ formation about the quaternary system with regard to the quantities of phases in equilibrium.

The J&necke pro­

jection of point P in Figure 3 is J, the Schreinemakers projection is S, and both projections for pure water give 0.

These projections form two similar triangles such that

WO/OJ (for the large triangle) equals PS/SJ (for the small triangle).

This means that the ratio of one hundred

(for pure water or WO) to the length of the radially pro­ jected point from the center (OJ) equals the ratio of the per cent water in P (PS) to the difference between the distances of the orthogonally and the radially projected point from the center (SJ).

Therefore the per cent water

in point P is given by the ratio of SJ/OJ (obviously, since the closer to the water apex is the original point, the greater is line SJ), and the per cent salt in point P is given by OS/OJ.

EXPERIMENTAL Preparation of Salts Much of the preparatory work was similar to that which was reported in an earlier report^. Ammonium Sulfamate This salt was prepared by neutralization of the acid with excess aqueous ammonia followed by several recrystal­ lizations, and also by several recrystallizations of the technical grade salt in slightly ammoniacal solution. The excess ammonia is added to reduce hydrolysis, since sulfamate ion is hydrolyzed in the presence of acid as are all of the salts in amine sulfonate series1^ to yield in this case ammonium and sulfate ions.

The crystals

were steeped with warm water and then cooled, sucked dry and washed with cold water, until a sulfate test made on the mother liquor as well as on the crystals was negative. The amount of sulfate which was present in the "least pure" preparation was estimated by comparison to turbidi­ ties produced by samples of known sulfate concentration and found to be in the order of a few parts per million. The white odorless crystals were then washed with cold alcohol and then with ether and finally air dried (to form a cake), ground in a mortar to a fine powder and stored in thin layers over Anhydrone.

The salt melted

without decomposition at 131°C, in agreement with litera­ ture values^, and was found to be at least 99»9% pure by analysis of both ammonia and total nitrogen content.

40

It was not possible to use the commercially available •‘pure" grade without further purification since it con­ tained approximately a half per cent of sulfate and was visibly wet.

The commercially available “pure" sulfamic

acid had an even slightly higher percentage of sulfate. Certain points of interest concerning the salt may be mentioned.

The salt crystallized from water to form large

rectangular plates which could be split to give thinner crystals, indicating a sheet-like internal structure. Like all the other sulfamates reported here, the crystals exhibited parallel extinction.

Unless the stirring was

continuous and moderately intense, the crystals which formed were always at least several millimeters to a centimeter in length.

The vapor pressure of a saturated

solution of the sal-6 must be close to that of the partial pressure of water in the atmosphere for on some especially humid days, particularly in the summer, the dry salt would deliquesce rapidly.

On some occasions the salt

would practically dissolve before the weigher's eyes as soon as it was removed from the desiccator and it was then impossible to work with it.

A solution of the salt

always gives a very slippery feeling to the touch.

The

solution would not wet glass and for this reason the technique of determining the densities of solutions by use of a calibrated delivery pipette had to be abandoned and changed to the technique of using a containing pycnometer.

In fact, the walls of solubility tubes in

a which solutions of sulfamates had been stirred for some time were not wet either by water or by the contained solution.

This non-wetting property was noticed for all

the sulfamate solutions but particularly for that of the ammonium salt.

Also noticed for all sulfamate salts but

particularly for the ammonium salt was the pronounced supercooling of the "melt" of the salt.

Saturated solu­

tions of ammonium sulfamate and of this salt with other sulfamates did not give a solid when dried at 80° or 90°, nor did the solid crystallize readily upon cooling in a desiccator to room temperature.

(The material was quite

anhydrous, however, because the weight corresponded to the theoretical or expected weight and it did not change on further heating.)

After several minutes of cooling,

particularly if the liquid were disturbed by tilting and rotating the containing weighing bottle, the liquid would become somewhat opaque.

This was accompanied by the

appearance of long striations and the emission of cracking noises of the type heard when ice cracks under pressure. Although solutions of the sulfamates are unstable and undergo hydrolysis in acid medium, it is of interest to note that after approximately one and a half years of stirring a ternary complex containing the ammonium and potassium salts dissolved in water gave no evidence of even a trace of sulfate in solution.

This stability of

the sulfamate ion in neutral solution is in agreement with the observation that the uncatalyzed hydrolysis for

the less stable amine disulfonate ion has an undetectable rate compared to the acid catalyzed reaction52.

in -the

system of ammonium and potassium sulfamates and water, it was noticed that solutions rich in the ammonium salt were neutral to methyl red (color was orange) while those rich in the potassium salt were alkaline to methyl red.

As

mentioned earlier, the ammonium salt is the least stable of the sulfamates studied here and its saturated solutions could not be dried at temperattires over 80°C. without some loss in weight and some hydrolysis as evidenced by a positive sulfate test.

On the other hand it took too

long to dry in an open oven at 80° and therefore a vacuum oven was used.

Utilizing the full power of an oil vacuum

pump at 80° (through a dry ice trap to prevent erratic weight changes), after slowly bringing down the pressure to minimize the occurrence of bumping, did not cause any weight loss of ammonium sulfamate and enabled solutions of this salt to be dried successfully. Potassium Sulfamate This salt was prepared by neutralization of the acid with the hydroxide or carbonate followed by several re­ crystallizations, and also by several recrystallizations of the technical grade.

The white square (but not cubic)

crystals were steeped until shown pure as with the ammonium salt and then treated and stored as done with the ammonium salt.

It was found also to be at least 99*9%

pure by nitrogen content.

Solutions of the salt also

did not wet glass but the effect was not as pronounced. The salt was never observed to deliquesce.

Addition of

ethanol to the mother liquor in a recrystallization caused a marked precipitation of the salt which was now found to be quite free of sulfate.

The salt also under­

goes marked supercooling but the effect is not as pro­ nounced as with the ammonium salt.

A solution of the

potassium salt is alkaline to methyl red which is in line, no doubt, with the distinctly greater stability of the potassium over the ammonium salt.

The salt was generally

dried at 80° under vacuum as was done with ammonium salt but it was observed that five successive ten minute heat treatments at the bottom of an oven 1200°) on two dried samples of the potassium salt gave no loss in weight and no test for sulfate.

It was also observed that quad­

ruplicate samples of mixtures of potassium sulfamate and iodide were heated at the bottom of an oven for three one and a half hour periods without loss in weight.

The salt

exhibits parallel extinction. Sodium Sulfamate Monohydrate As reported earlier^, the anhydrous salt was made fortuitously, at first, by addition of sodium carbonate to the acid and was even re crystallized from water in the metastable anhydrous form.

However, after the stable

hydrate phase did appear, it was impossible to prepare

the anhydrous form again as a metastable phase.

Nonethe­

less, the anhydrous form was prepared once in this re­ search, according to the suggestion of van der Meulen^l. It was 'unfortunately not possible to repeat this prepara­ tion.

The method is as follows and presumably is based

on the lowering of a transition point by the addition of another component, just as the melting point of a solid is lowered by the addition of another component.

A

solution of 175 grams of sodium sulfamate in 100 grams of water is warmed to 55°, methanol is added until a permanent precipitate of the anhydrous salt is obtained, the system is cooled to room temperature, and the anhydrous crystals (now stabilized by the methanol) are easily filtered and washed with methanol.

The difficulty

of obtaining the anhydrous crystals again by this method might be found in the nature of the hydrate.

When the

anhydrous salt takes up water of hydration, it becomes sticky and unworkable.

If the hydrate forms from (even

a large amount of) solution, it forms long, soft, silky, fibrous crystals which spread out through the entire liquid and immobilize it, somewhat like a gelation. forms extensively supersaturated solutions. possible to recrystallize the hydrate.

It

It was im­

Attempts to wash

the hydrate at temperatures above its incongruent melting point (38»3°) were unsuccessful, with the sticky hydrate always appearing to block the passage of liquid.

Washing

the damp solid with methanol also produced caking and did

45

not remove any sulfate.

Four successive recrystalliza­

tions of the salt by dissolving it in water, adding methanol until solid precipitated and centrifuging did not purify the salt of sulfate.

Addition of ethanol,

isopropanol, dioxane or acetone as salting out agents to an aqueous.solution of the salt caused the formation of two immiscible liquids.

Addition of a solution of picric

acid or hydrochloric acid to vary the dielectric constant (and therefore perhaps stabilize the anhydrous salt) pro­ duced only sodium picrate and common salt, respectively, as precipitates. Finally, the pure anhydrous salt was prepared as follows.

Pure sulfamic acid was prepared by washing the

technical grade acid by decantation and with much stirring until free of sulfate.

The pure acid was then

washed with alcohol and ether, and air dried.

To a mol

of the acid was added slightly less than an equivalent of Reagent grade sodium carbonate decahydrate.

(The result­

ing mixture became cold rather than hot, probably be­ cause of the heat of volatilization of the carbon dioxide.) When all the carbon dioxide had been expelled, the solution was made neutral to litmus with a solution of sodium hydroxide and then filtered.

Methanol was added, thereby

precipitating the anhydrous salt which was filtered and stored in the open.

However, after using part of this

material in the first few ternary complexes, it was placed in an oven at 95° for from two to four hours with

occasional stirring (after which it lost 3 parts per 10 thousand of weight in the heating of test samples at 200°C for half hour periods), and then stored over Anhydrone.

This material was found to be free of sulfate,

pure by nitrogen content, and stable at 200° for short periods (thirty minutes) of heating. Both the anhydrous and the hydrate crystals exhibited parallel extinction.

The property of creeping when the

hydrate was exposed to the atmosphere, with a bristling formation of anhydrous crystals was observed under the microscope.

The small, even, well defined, clear hydrate

crystals were seen, on exposure to air, to thicken and become distinctly opaque and white as the salt rose up out of focus of the microscope.

The resulting anhydrous

crystals crumpled when pressed between the fingers and were quite dry as opposed to the long, silky, shiny, sticky fibers usually obtained as the hydrate is formed and to the tiny translucent well-settling needles of the hydrate obtained under carefully controlled conditions discussed later. Additional Reagents The potassium chloride used in phase studies and as a standard for argentimetry, a Reagent grade, was ground, and dried at the bottom of an oven (at approximately 200°). Reagent grade sodium carbonate which was dried at the bottom of an oven (200°) was used for standardization of

47

acid.

C, P* grade ammonium chloride was recrystallized

until it sublimed, and used in early work as a standard source of chloride and ammonium ions, giving the same value as did potassium chloride.

Reagent grade potassium

acid phthalate was used as a standard for the alkali titrating solution.

After one hour of drying in an open

dish with occasional stirring, both salts gave essentially the same value.

C. P. crystals of potassium iodide were

ground and dried at the bottom of an oven and used as such.They apparently contained .2 to .3% water which was lost only after fifteen hours of heating in a dish at 550° (after days at 350°).

Reagent grade sodium chloride

was heated at the bottom of an oven and then used after cooling.

The potassium sulfate was a O.P. powder and was

used directly from the bottle after duplicate samples, heated for forty hours at the bottom of an oven, lost only 0.02% of their weight.

Reagent grade ammonium sulfate was

ground and dried at 110° for a few days and was used as the standard for ammonium analysis in all the work using formaldehyde as well as the source of salt in the phase studies. this salt.

It was difficult to dry solutions saturated with Weighing bottles containing the salt and its

solution could be kept for days in a 110° oven without apparent change as a result of the crust which always would form on the top of the solution.

It was necessary to

shake the bottles frequently and thus break the crust to cause noticeable drying.

37# Reagent grade formaldehyde

was used in this work.

The presence of methanol in the

formaldehyde did not noticeably affect the colors of the end point nor the results obtained. Methods of Analysis The Vo1hard method, with digestion and filtration of AgCl before

titration was used for determining chloride

in systems where chloride is one of the constituent ions. Sulfamate does not interfere with the determination.

In

the standardization of the 0.2 N HC1, using sodium car­ bonate, methyl red was used as an indicator and the solution, after slightly exceeding the end point, is boiled to expel CO2 and back-titrated with a small amount of 0.2 N NaOH.

Iodide was determined by the clear-point

method utilizing eosin as an aid or preliminary indicator. Since the system can not be back-titrated, the success of the method, which involved the sharp clearing up of the solution at the end point, depends on the vigor with which the flask is shaken. The early systems were studied by means of deter­ mining ammonia by the Kjeldahl method.

As previously

mentioned^-, concentrated alkali was added to the closed system and the liberated ammonia was steam distilled into a saturated solution of boric acid.

This is a strong

enough acid to fix the ammonia but weak enough not to interfere with the titration.

The ammonium borate was

titrated with 0.2 N HC1, using a mixed methyl red -

49

methylene blue indicator to a color of pale blue or gray. Total nitrogen was determined by the same method on a sample which had previously been digested with concen­ trated sulfuric acid,

Atempts to improve the reliability

of the method were as follows:

the use of an all glass

apparatus including ball socket joints (but which some­ times allowed for the introduction of grease into the system, swept over by the steam); the use of a two-holed rubber stopper on the receiving flask, thereby sealing it from the atmosphere; the use of a tiny trap filled with boric acid on the outlet from this rubber stopper to insure trapping all the ammonia; the sweeping out with steam of all non-condensable gas (i.e,. air) prior to addition of alkali to prevent any carrying away of ammonia by this gas; et cetera.

However, because of the erratic

nature of some of the later results and because of the considerable care and time required for the distillation, it was deemed advisable to find a more dependable and a speedier method for determining ammonia.

Such a method,

utilizing the reaction between formaldehyde and ammonia, as reported by KoltKoffS?^ was found and is discussed below. The reaction between an ammonium salt and formalde­ hyde to produce hexamethylenetetramine (hmt) and to liberate the acid has been known for nearly a century54. Early studies55 of the mechanism of the reaction indicated that weakly basic methyleneimine was formed as an inter­ mediate according to the following reaction

50

NH4C1 -f CH20 -^NH^OH^OH (HOI) — ?H^0 -f- CH2:NH(H01) It was suggested that the imine is stable in acid and that hmt is formed only after the acid is neutralized or is removed.

The experimental observations made by

the present writer do not verify this conclusion and, in fact, indicate that the acid is liberated immediately after the addition of the formaldehyde.

The reaction

between aqueous formaldehyde and gaseous ammonia liberates 81 kcals per mol of hmt formed5b.

studies of the reaction

rate in solution indicate that the heat of the reaction is only 19«5 calories57.

The reaction was studied at various

temperatures and the rate was best represented as third order.

The first two reactions, viz.. the formation of

three mols of methyleneimine from three mols each of formaldehyde and ammonia, and their condensation to form the cyclic six-membered ring compound trimethylenetriamine are believed to take place rapidly.

The subsequent re­

action with three mols of formaldehyde to give trimethyloltrimethylenetriamine is comparatively slow and is the measured reaction.

The final reaction with a mol of

ammonia to give hmt may be slow or fast.

The rate deter­

mining step plus the first two steps result in the con­ sumption of formaldehyde and ammonia in the ratio of two to one; and with the last reaction, the over-all con­ sumption is in the ratio of three to two.

The formation -

decomposition curves at different temperatures and different values of pH show that curves of the

51

velocity versus the distance from the equilibrium con­ centration are broken, that decomposition is more rapid than formation at room temperature, but that the dis­ continuity diminished at 50° to just a hint of discon­ tinuity.

This fact was utilized by the present writer

and will be mentioned again.

The equilibrium constant,

(claimed by Baur57 to be a function of pH), for the reaction 6CH20 + 4NH4* a (CH^) (jN4 + 4H* -t 6H*0 is given by 1/K a (formaldehyde)b (ammonia)4/(hmt)(H )4 At pH = 5.4 Baur claims the values of 1/K x 10“10 are 1.58 at 25° and 21.8 at 50°.

This displacement of the

equilibrium point with temperature is in the expected direction for an exothermal reaction.

However, similar

systems employed in this research were heated to 50° for reasons discussed later.

The application of the above reaction toward the analysis of ammonium salts was first made by Kolthoff58. The product, hmt, is such a weak base (K^ equals 8 X 10_10) that its aqueous solution is not alkaline to phenolphthalein and so ammonium salts may be titrated sharply using this indicator, according to Kolthoff.

On p. 158

of his latest book53t Kolthoff says that "the reaction with formaldehyde proceeds rapidly enough to permit direct titration with alkali" while on p. 217 of the same book he states that "the reaction between formaldehyde and ammonia does not take place instantaneously; at least an hour is required before titration".

Although

52

this hardly affords a clear understanding of the reaction, in a sense, both statements are correct.

According to the

findings of the present writer, the reaction is initially rapid but not complete at room temperature for at least an hour.

Recently59, the method has been applied to the

analysis of nitrogen containing organic compounds.

The

digested samples were neutralized to a methyl red end point and titrated to the first pink color' of phenolphthalein (as also stated by Kolthoff), with the results being low on the average by -0.04# nitrogen even though a blank of .2 or .3 ml. was applied.

Below is described the method of anslysisfor

ammonium nitrogen and for total amino and ammonium nitrogen as worked out and used by the present writer in the phase studies reported here.

It affords a rapid, accurate, direct

and fairly simple method for determining nitrogen without the time consuming and laborious distillation used in the Kjeldahl method.

It is applicable to ammonium salts, to

amino compounds and presumably to nitrates (reduced in neutral medium with Arnd's alloy53) both separately and together and presumably also to the volumetric determina­ tion of bismuth as the chromium hexa-ammonia complex?5.

its

limitations will become apparent during the ensuing des­ cription and will be summarized later. The method may be applied directly to an ammonium salt in the absence of excess base or acid, or to a solution con­ taining either excess acid (or base) after neutralizing with base (or acid) to the methyl red end point.

The analysis

discussed now is for nitrogen in an ammonium salt

referred to as "ammonium” for our purposes, modifications being required (as outlined later) for the determination of ammonium in the presence of an amine and for the deter­ mination of total ammonium and amine.

In essence, the

method involves for ammonium salts the liberation of the acid of the ammonium salt by the addition of formaldehyde, followed by the titration of that acid with alkali.

Since

formaldehyde always contains some formic acid, it must first be neutralized and this was always done by diluting with water one to one, adding one drop of 1% phenolphthalein for each 10' cc. of 37% formaldehyde taken and then treating with HaOH solution to a pale pink color. Shis freshly prepared formaldehyde solution was adjusted with NaOH to such a point that dilution of a sample of the same size as used in an analysis to the same final volume as attained in a titration would give the same color when .3 to .4 ml. of 1% phenolphthalein are added as was obtained at the end point of the titration.

This

statement implies correctly that addition of the phenol­ phthalein caused a color change.

Also, if the phenol­

phthalein were added before dilution in these ’blanks*, then dilution with water would cause a color change. All the analyses were carried out in rubber stoppered 125 cc. Erlenmeyer flasks.

A weighed sample of ammonium

sulfate, of sufficient size to require approximately 40 cc. of .2 N NaOH solution, is introduced and diluted to approxi­ mately 50 cc..

A drop of methyl red is added, followed

by addition of 10 cc. of freshly prepared formaldehyde solution with .3 to .4 ml. of 1% phenolphthalein (in alcohol).

Occasionally, if the phenolphthalein were

added before the formaldehyde, the extreme dilution caused it to come out of solution as a white colloid and so the phenolphthalein is added to the required sample of formaldehyde prior to each titration and they are then added simultaneously.

The presence of the formaldehyde

kept the phenolphthalein in solution.

The addition of

the alkaline formaldehyde causes the pale orange - pink color of the aqueous ammonium sulfate system to become instantly yellow.

However, within a second, it changes

to red, indicating that the reaction between formaldehyde and the ammonium salt to liberate the acid takes place virtually as soon as the reagents are brought together. The system is now titrated rapidly (for convenience) with base and the color of the solution gradually becomes yellow beyond the middle of the titration.

The titration

is halted when a "permanent” pink color (of phenol­ phthalein) is obtained.

The flask is stoppered, placed

upon a hot plate and maintained at 50° for at least five minutes, during which time the color fades to the yellow of alkaline methyl red.

After a period of from five

minutes to three hours of maintenance' at 50°, the samples are titrated finally and directly to the point of greatest color change of phenolphthalein which is not very easy to detect.

At the end point, the system is best described

55

as appearing to flush. as deep rose.

The color may be best described

The end point occurs approximately five

drops of .2 N NaOH solution beyond the first appearance of a pink color.

Beyond this point, the darkening of the

solution occurs at a much lower rate.

Potentiometric

fixing of the end point would probably eliminate this difficulty.

This end point will not fade except very

slightly, over a period of several days to a week.

Con­

sequently, to make easier and more rapid the fixing of the end point, two weighed samples of ammonium sulfate were always analyzed before each regular set of analyses. These samples then could be used as reference color standards in the event of doubt of going beyond the end point. During earlier work, the above mentioned early stopping point (without heating) was believed to be the end point. It was noticed at that time that stoppered titrated samples, which had been placed aside, kept fading. point was readjusted, it would slowly fade.

Thus as the end It was sus­

pected that this was caused by formic acid which was slowly being formed from the formaldehyde.

However, on

readjusting the end point of some titrated samples which had remained stoppered overnight, a color was obtained which did not fade.

At this point, a literature search

disclosed the aforementioned research57 which revealed that the forward reaction was slower than the reverse reaction at room temperature but almost completely equally

reversible at 50°.

It was for this reason that systems

which are close to the end point are heated to 50°.

How­

ever, samples which were preheated to 50° and then ti­ trated rapidly gave low results and faded slightly.

A

minimum time period of five minutes was necessary for com­ pletion of the reaction at 50°, whether the original sample was titrated at 50°, or at room temperature and then heated to 50°. Variation of formaldehyde concentration showed that while more formaldehyde had no pronounced effect on the reaction, half the amount of formaldehyde definitely slowed the reaction.

Thus, samples to which only 5 cc.

of formaldehyde had been added gave an end point which kept fading as it was readjusted over a period of thirty minutes, after which time the correct and constant value was obtained.

Por this reason, the concentration of

formaldehyde was kept under control but it was hardly as critical as when amine is present. Analyzing samples of pure ammonium sulfate of such weight that 10, 20, 30 and 40 cc. of .2 N HaOH solution were used gave the same results, thereby indicating that any internal blank was negligible.

If the final end

point were exceeded, the sample was successfully backtitrated only if done immediately; otherwise high results were always obtained. It was noticed, even before the technique of heating the system to 50° near the end point was developed, that

57

results of analyses on samples of pure ammonium sulfamate were high by several per cent.

Apparently the presence

of sulfamate interferes with the titration.

The same

effect was noticed in analyses of samples of ammonium sulfate to which pure potassium sulfamate was added.

As

noted, the concentration of formaldehyde did not affect the titration of ammonium sulfate as long as at least 10 cc. of prepared formaldehyde were used for the specified time. This was not true in the presence of sulfamate for the greater the formaldehyde concentration, the higher the results.

Varying the

quantity

. of potassium sul­

famate from .1 to 3 or 4 grams did not appreciably affect the results, as long as the formaldehyde concentration is doubled for sulfamate samples of from .1 to .3 grams. Apparently formaldehyde or alkali react with sulfamate in some way which is dependent on the concentration of the formaldehyde but not appreciably on that of the sulfamate, to produce an acidic product.

In none of the reactions of

formaldehyde with amines56 is acid produced or used up and so probably either of the following two processes takes place:

a) formaldehyde increases the ionization of the

hydrogen from the amino group and thus causes sulfamate to act acidic; or b) formaldehyde speeds up the hydrolysis of sulfamate to give the corresponding ammonium salt which then reacts with the excess formaldehyde. It was further noticed that this interference was sensitive to changes in room temperature.

As already

50

stated, the titration of ammonium sulfamate always gave high results by several tenths to several per cent.

The

end point would fade in several minutes and refixing the end point would again not give a permanent end point. This fading was pronounced on a summer day while not too noticeable in the wintertime.

Therefore, to standardize

the procedure for analyses in the presence of sulfamate, the samples were always diluted to oO cc., and 15 cc. of formaldehyde was added to each sample which had always been preheated to 35°, presumably to allow completion of the main reaction and not speed up too much the inter­ fering reaction. possible:

Samples were titrated as rapidly as

oO to 10 seconds was the usual time required

between the addition of the combined foimaldehyde and in­ dicator to the fixing of the end point at the "flushing" point.

The standard color in the known ammonium sulfate

samples was referred to, as a check.

As mentioned earlier,

standard ammonium sulfate sample s , both pure and with added potassium sulfamate, were always analyzed prior to the analysis of any set of unknown samples. The determinations of total ammonium and amino nitrogen involved the same titration as described for an ammonium salt with the variation that the sample was first digested with sulfuric acid.

The sulfamate ion undergoes

hydrolysis rather easily under the conditions used.

The

15 to 20 cc. samples, after addition of 1 to 2 cc. of concentrated sulfuric acid, were heated on a hot plate

59

until fumes of SO^ were evolved for several minutes. Actually this stringent treatment was not required for it was found that maintenance of a sample on the hot place for an hour but without fuming was sufficient to complete the hydrolysis although a forty-five minute treatment gave results which were low by .2%.

Both samples had had 3 cc.

of concentrated sulfuric acid added and were diluted to 50 cc..

Also, complete hydrolysis was effected when 5 cc.

of concentrated sulfuric acid were added to samples of these dilutions and kept in a 110° oven overnight (18 hours).

To minimize the amount of neutralization required

and therefore amount of salt introduced and yet to insure complete hydrolysis, 1 to 2 cc. of concentrated sulfuric acid was added and the samples brought to flaming for several minutes.

Samples which fumed freely, i.e.. with no watch

glass cover, for 13 minutes did not lose ammonia.

On

cooling, the sample is diluted with water, cooled in ice, and a drop of methyl red is added.

The sample is nearly

neutralized with 30% NaOH, then neutralized with 1 N NaOH and finally made acid with HC1. to evolve CO2 .

The sample is then boiled

Long boiling did not alter the results but

some boiling is essential or else the system will be buffered at the methyl red end point.

The system is cooled,

and its pH adjusted to the neutral point of methyl red, which also corresponds to the color toward methyl red of a solution of the check samples of ammonium sulfate.

The

ensuing procedure was identical with that as described for

60

the pure ammonium salt.

All the results obtained were

reproducible to within a part per thousand. The advantages of the reaction may then be summarized as follows: a.

The analysis is very much more rapid than the

older methods. b.

No special equipment is required, e.g.. distillation.

c.

No transfer of sample is required.

d.

The method is direct, requiring no back-titration.

e.

The reaction is versatile, as applied to ammonium,

amino, and presumably to nitrate and even bismuth analysis. The limitations of the reaction may be summarized as follows: a.

The lack of sharpness of the color change at the

end point requires the eye to be somewhat trained.

This

difficulty might be eliminated by the use of thymolphthalein which shows a sensitive change at a higher pH than does phenolphthalein.

Potentiometric titration might

eliminate the problem entirely. b.

A waiting period of five minutes is necessary

after the bulk of the titration has occurred, for accurate work, although this period might be eliminated by applying a blank.

Actually, no time is wasted, for several other

samples are titrated while the preceding; c.

one is. heating.

The interference of amino groups imposes the

requirement of fairly strict temperature, volume and time control.

fo

d.

For digested samples, a careful neutralization

must be included, (which introduces the possibility of loss of ammonia, minimized by cooling with ice), slowing up the procedure, and introducing the slight possibility of certain errors because of a second end point. e.

The absence of any other ions which will be

titrated between these end points must be insured, e.g.. carbonate, phosphate. Solubility Measurements The solubility measurements were made in the usual way as described in earlier papers.

Known complexes were

made up in rubber or glass stoppered solubility tubes which contained several marbles to speed equilibration rates.

The mouth of the tube was protected with a glass

cap held in place, sometimes, by a rubber stopper which gripped the tube.

Almost all the tubes were shaken

vigorously after compounding and placed directly into their respective baths for tumbling.

Therefore, since

pronounced cooling occurred on shaking, most of the solubilities reported were approached from undersaturation. The temperatures of the baths were respectively 9.0° - .05° 0. (because that temperature was required in connection with another problem) 25.00° t .03° C. (as checked with a Bureau of Standards thermometer) and 45.0° - .05° C..

In the simple systems, equilibrium was

attained in several hours.

In the systems involving

solid solution, equilibrium was established within a few days (probably not more than twenty-four hours being re­ quired).

Saturated solutions were removed through filter

paper by calibrated weighed pycnometers for density deter­ mination and then, after that determination, were trans­ ferred to weighed bottles, reweighed, heated usually overnight or for eight hours in an 80° oven and then dried to constant weight in a vacuum oven overnight at 80° with full pumping (through a dry ice trap).

Filtered liquid

samples were removed by pipettes (preheated for 45° work) and transferred to weighed rubber stoppered 125 cc. Erlenmeyer flasks (to which a known weight of water had been added for the 45° work in order to minimize the evaporation of water from the warm sample) and diluted and aliquoted to give two samples each for titrations of ammonium and of total ammonium and amino nitrogen.

From

these data, all necessary information could be extracted for the plotting of four component (and simpler) systems. All the salts settled well from solution with the exceptions of potassium sulfate, which was a fine powder requiring hours to settle, and of sodium sulfamate hydrate, which is discussed below.

It was noticed that the solid

solution, in the system of ammonium and potassium sulfamates at 25°, would enclose the marbles on long standing to form an apparently continuous solid which had to be dissolved away with hot water and which could not be broken up by chopping.

Also noticed in this system was

the formation of a cake or persistent plug at the solid solid interface, when water was added in making up a complex in the solid solution region.

Qualitative tests

on this system and on the corresponding sulfate system, however, did not show this phenomenon to be repeated and indicated that it probably occurred because the complex was not stirred at all for some time after it was completed, thereby allowing the solid reaction or solution to occur on a more than local scale. With regard to handling, the sodium sulfamate hydrate gave a great deal of difficulty, making it virtually im­ possible to make up mobile ternary complexes.

Complexes

made up in the region where the monohydrate did not appear as the saturating phase could be made up without difficulty. The procedure which v#as evolved for compounding complexes in which the monohydrate was the saturating phase was as follows.

Ternary complexes were made up in which the

monohydrate should be the saturating phase and, utilizing a knowledge of the solubility in pure water, they were made to lie close to the saturation curve.

Thus the amount

of troublesome monohydrate with which it was necessary to work was kept as low as possible.

Then, after the complex

had been made up (of the anhydrous salts and water), the tubes were rotated at about 30°, at which temperature the salts were completely dissolved or, if the complex lay in the invariant region, sodium chloride was the solid phase. They were then cooled at a very slow rate with much shaking.

Sometimes they had to be seeded if the monohydrate would not form when stirred in a 10° bath.

With much attention

and vigorous agitation, and by warming the tubes with the hand (i.e.. very slightly) as soon as the mass became immobile in the 25° bath, the complexes were finally made so that they remained very slightly mobile in the 25° bath. (Approximately a full day’s work was individually required for each complex.)

By the next day, the complex was

either immobile again (most commonly) or slightly more mobile.

After another day of stirring in the 25° bath,

the complex was mobile enough to be considered satis­ factory, although after several weeks of stirring at 25° of a tube of the hydrate in water, the crystals became quite uniform, small and needle-like and very nicely settling; and after several months of stirring of this tube, the solid became considerably more translucent and seemed to have attained a density close to that of the saturated solution.

65 The.System NaSOjNH^, - NaCl - H.,0 at 25°C The results of the investigation of this system are presented in Table 1 and are plotted in Figure 4.

NaCl

was determined by the Volhard method using .2N reagents and Na sulfamate was obtained as the difference from total solid as determined by drying in an open oven at 95° over­ night.

Further drying at 110°, or for half hour periods

at 200° (i.e.. the bottom of an oven) showed no change in weight.

The system is simple in that no solid solution or

compounds are observei to occur between the solid phases. In general, extrapolations of tie-lines are made to a line which, in most of the systems reported here, is located at the comer.

Specifically extrapolations are

made to the line of 100% of the salt at that corner and the extrapolated error is expressed as per cent of the second salt at 100% of the first.

Thus in extrapolating to the

100% NaCl line, the following formula is used [A s differ­ ence in composition between the complex and solution): X (in % sulfamate) * aNall (100 - % NaCl in complex) Then, % sulfamate in complex - X = error in % sulfamate at 100% NaCl. Extrapolations of tie-lines toward Na sulfamate monohydrate are expressed as per cent water and are obtained as follows: X «

(% NaCl in complex) ANaCl

Then % H^O in complex - X * % H^O in hydrate.

The extrapolated values obtained are then compared to the per cent water theoretically present in the monohydrate, viz,. 13.1$.

At 100$ NaCl, the average

error obtained by extrapolation is less than 1$ NaS, and at the monohydrate, the errors for the two complexes are 0.8$ and 0$ water.

As described on pp. 53 and 72, it was

very difficult to work with complexes in which the hydrate was the saturating phase.

Complex 3 was made by adding

NaCl to a clear saturated analyzed solution of the hydrate to avoid working with a large amount of the hydrate, but although it is obvious that only a relatively sBall weight of monohydrate could have been precipitated, the amount was sufficient to immobilize the entire complex.

It was

necessary to open the tube and stir several times with a rod, thereby removing material, and voiding any ex­ trapolation through the original complex.

Invariant

Complex 12 was made by adding 5 g< of NaCl to the remains of Complex 11.

Complex 4 gave rise to a liquid metastably

saturated with respect to the small well-settling, needle­ like crystals of anhydrous sodium sulfamate.

(The ex­

trapolation, of limited significance because of the closeness of solution and complex, is to the anhydrous sulfamate coiner.)

This complex remained in metastable

condition throughout two samplings and over a period of six days of tumbling in the bath, but shortly thereafter, on standing in the bath, it underwent the distinct transi­ tion of hydration.

Unfortunately, the vigor of shaking,

97

required in the attempt to make the mixture mobile, forced the marbles through the tube.

The complexes in which the

hydrate appears as a saturating phase were made up so that equilibrium was approached from a higher temperature while the others were equilibrated from a lower temperature.

©8

© —

%

ra

soi t m’

S3 M « 2 »* CM CM

°«

"It

«

co

co

10

© «P a)

I

in CM

**

nJ

® H 8 % •Sw.3 R.S3

®

"S

Eh

in

>. m

CM

Vi

»d ,d

03

O

CO •

>

to

• rH

01

d 0 E d •H

V.

*

cS

1 aJ

c^ in

to

© •P

!3 I

o

rt

oo

0

0 0 w 1 I— I o

£5

CO • CO

§ n

to

JS

O)

IO • rH

$

03 IO • rH

m a*

cn in

• rH

• rH

• rH

• rH rH

CTi

CO

03

in ti

,lS &

CM pi i-t • p •

in •cj cm in in

cn 0

•

cn

^

0

to

• o-

S3 CM m

CO to • cn

03 •

to

O

t o CM CO

03

03

LO

03

03

£ .o rH rH •

o'

P3

« rH

flD

d Jfi Jh CD H-l T5*IrOH • „

rt= X

®

X

+3

03 5!

O-

„

0

rH

rH

O

in

0-

in •

rH •

to

rH

to CO

O • Q 03

cn

GO • 03

CO

•P

& • e 43 o £s

03 O

o

• 03 in

in

0•

CO

0

O • CO sjl

in

• 03

in

m

•

0

in

3 *H

u

aj

>

d

n

rH

. CO • rH CO

to

•

cn

•

Vi » ,d CO o >»

o

TjC rH •

o

0-

CD d

d

03 • -

o « o

to CO

-P

CO

CO

fr•

H

o

CO

o

CD

a 03 «

03

CD

O«

i —I C M CO

03 • i—1

7Jf> The System KSO^H,, - KOI - H O at 25°C The results of the

investigation of this system are

presented in Table 4- and plotted in Figure 7 . 'KOI was determined by the Volhard method and total solid was determined by drying at 80° in a vacuum oven overnight, the weight obtained being constant.

The system is simple,

with errors obtained by

extrapolation averaging

-.1$ KOI at 100$ KS and

+ .3$ KS at 100$ KOI.

Point 5

is seen to lie exactly on the border of the invariant univariant regions.

78

®) —M 0 +» ri ro ri^3 M OT Ph

103 M

u

k M

IC/3

=

U X. in o

+>

H

S\a

•H W

CO

£ 0 Q

3 • I

00 CO • rH

> Fl

f» C

H

n

M in CO • rH

CO in CO • rH

W

lO) X. as CO 03

s

f*

rH

i

S CO

cjl 03 CO • rH

CO CO • rH

rH CO 02 • rH

LO LO • IN rH

to IN •

to lO •

CO

rH

8

02 LO • CO CO

0cr> • CO 03

CO o •

to o• a? rH

O00 •

O

S4• O •a

03 O• CO

rH LO • to CO

CO to • LO

rH

oon •

rH rH •

$

8

rH rH • O 03

rH rH • CD 03

8•

in

in •

oco

CO

O

03

03

d O P4 +>

CO CD •

o rH

to to •

rH CD

CO

03

•H ?H >

rH

•

•rH

03

03

CO

CO

• in

CO

9

in

•

03

CO

• CO 03

d

to • in

to • CO

O-

10

Ico

rH

rH LO • > rH

0-

LO LO •

-

in

CO

02

to

• O

Aver.

•p

O•

CO • to rH

I

o

tn

ii

in the system KS - KOI - HgO et 25°C Equilibria

Table Ko. 4

K PI

C3

IN

to

in

• to 03

_

ICO

79

*

The System KSO^HH^ - KI - H20 at 25°C The results of the investigation of this system are presented in Table 5 and plotted in Figure

8.

KI was

determined directly by the clear point method, sulfamate being found not to interfere.

Total solid was obtained

by drying overnight at 80° in an open oven.

Increased

heating was not necessary although tests were made (in­ cidentally, showing that potassium sulfamate is stable at the bottom of the oven).

The system is a simple one, the

errors obtained by extrapolation averaging -.1% KI at 100% KS and -.4% KS at 100% KI respectively.

Point 3 is

seen to lie exactly on the border of the invariant univariant regions.

81

Iw tifl® in •p nj c v S m m

-

\xn

'M * >* O O • o

u -p W

=

=

l!

M w

M lio *: H ^

• o

•

8 • 1

03 rH » +

d ►h

• t> d i—i

• to > Cl S • pi 1w

CO • 1

25°C

re

at - Hgp - KI

d D PI

d

o

•H

Os CO •

I—I O

vf‘ CO «

CO

o

CO

CO

• CO CO

to 03 • CO CO

m o • o03

to O• 03 03

o o

in cr>

03 to • c^-

1—[

CT> o •

to

cr>

o

in

rH •

o in

rto

oto

m

c

c*• in in

03 O • in rH

iH O • in rH

to 03 • to

CO in •

-P d O 04 •P

to pH • to to

• o

in

•

in

system

KS

CO

•

Iw

in •

CTi

»-l

o • in H

IN cr> • iH

the

W

03 r• CO

in

•

c

•

rH

pH

pH

O

03

’tJ1 • 003

o in

to

CO

pH in

•

rH •

to

rH

P< •

& *5 O o

Jw >4

CO • CO in

00 • 03

in

03

CO • CO CO

CO in • CO

o •

in to • pH 03

o pH

a

CO

to

CO

03 o•

•H

§ u

•H

> d

I

Equilibria

No.

cs

lO

Aver.

5

tO

Vi

CO

ysble

i n c^-

CO

-p •H

a

•3'

in

LTj

CJ

pH O•

o • o

to

85 The System KSOjNIL, - KgSO^ - HgO at 25°C The results of the investigation of this system are presented in Table b and plotted in Figure .'.9 .

Sulfamate

was determined on digested samples sometimes by Kjeldahl distillation (noted with the letters Kj) and sometimes by the method employing formaldehyde as described earlier. The agreement between the methods is fairly good consider­ ing the difficulties being experienced at the time with the distillation.

Total solid was determined by drying

in an open oven overnight at 80°. stant.

The weight was con­

The system is simple with errors obtained by ex­

trapolation averaging +-.4% KS at 100* K . ^ and -.1% KOI at 100% KS.

84

o %

W 01fAaj

Itn

tn ^

M +

■

a =

to CV3 W

S

IwO

p< 5S it

t£) o

+3 M W

*

s &

t-H

t> a

£

t n CM

H

ro

CO *

O o in CM

to|

a)

rH 'S

EH

CM M I |tf> M E a> -p u> >* v)

£

cn rH

in

CO CO • rH

in CO CO

« rH

03

in CO CO

• rH

03

rH

03

03

00 03

•n> M 'ao C-3 03 l-t O • • • rH 03 03

03

in o CO

o CO CM

cn to

rH

• rH

rH • rH

rH • rH

om c4

• 03

i—i

rH

•

•r-3

W

•P d

CO

IS • CM

n

to

•

in

in CO • CO

er) CO • CM

CM

cn • cn CO

d

*H U

d >

In

o tn

d o •H

-P •»H w d a) P

in

CM O •

in

rH

A ver.

*P d O _ 01 W f

o

c

rH

in rH

CO

a>

t£)

in Gl

o

to

to • IS

in

is

The System NH^SC^NHg - KSOjNHg - HgO The results of the investigation of this system at 25° are presented in Table 7 and are plotted in Figure 10. Determinations were made for total nitrogen and for ammonia by the method using formaldehyde and usually for total solid at 80° overnight in a vacuum oven.

The

analyses were always done (as in most of the systems in­ vestigated) in duplicate on aliquot portions of a single sample and then repeated on new samples after further stirring.

Total solid was determined on the samples used

for density determination.

Whereas it was indicated that

equilibrium was established within a half day in the simple systems of sulfamates and other salts, here equilibrium was established within a day in systems saturated with the potassium rich solid solution phase and within three days, possibly less, for the rest of the system.

Complex 3 was made up so close to the saturation

curve that only a very small amount of solid was left un­ dissolved.

The amount was estimated by volume to be .1 g.,

the tie-line was extended the appropriate distance using a base point defined (approximately) by the neighboring com­ plexes, and the values were read from the graph as listed in i'able 7.

complex 10 formed a cake that was exceedingly

difficult to break and it was suspected that a compound might be formed here,

however, tests on small complexes

made up along this region showed no repetition of this effect.

Apparently the solids had proceeded to form their

3?

solid solution while the tube was standing, prior to shaking.

Solid composition is determined indirectly by

extrapolation of tie-lines according to the following relation: X then

s

. (HyO in complex) A H2O *

_ #'KS in the solid solution » X + % KS in complex.

88

0K

o

5 cn

CD

CO

S3 00

to in

t o r—i

fr-

•H

•

CO

o

rH

in

CO CO

rH

CO

to

in

to O l o

^ to p-

p-

CD

CD

to rH • rH in

•

in

«

rH

in

o

o

rH

CO

CO

•

CO CO

rH «

fr­ ee

CD

o•

f*

•P

W •H r

cn

- 'fcR, .-I Mot of M OT5=

rH

to • to

m

CD

CO

S 3 03

to fr• CO

CO

CO •

CO *

I —I

8

CO CO

03 03

IO

rH

fr• CO CO

00 • in

to ^

to CO P-

s

• 07 in

cn

to

p-

p-

co

pto

in co

00 o

fr*

frO •

fr-

co

p-

p-

to co p-

rH

•

• rH to

fr-

CO

rH CD

p-

p

*

in

in

to

3

3

3

tp

tf

oo

in fr•

PJ

CD

to •

•

03

O c

>>

IS

P

-p

CO £

aJ

©

rH CD

CO • i— I

CO o • rH

•

CO

rH

• rH

• rH

fr-

in

CO

cn

pts
= a>

A p

to

£

“

W 13 S• p

*

tn

>

&

S*

00

CO

til

>*

P

t

• C 03

CO

•rt • 0 CU a* -H

o

03

to

• CO ClQ

cn

'S

oi

o

o

CO

So

>

til CO

a

&

i—t

in H

CO

$

> CO

CO

co

>

0

to

09"

to to o •

to o •

CM CM •

c^

to to •

8

CD •

CD

cn

03 05

03 CD ■tf •tf 3 P J a)

CM to

to Q

• rH

•

in

cn O • 03 03

to o rH •

CO

in

'w

CM r•

P'-

03

to

05

o

05

to

>5 •H P

cn

R

to p>

8 ^•>da>

o p R •H c

rH

—

CO

03 in $ to

tfi^

• rH M CD tO—0- -Jtin

CO

o cn cn • co

03 •

CO

•

rH

tf* D00

00 CO CD

CO

O • c

CO P O-

•

D* CO

LO rH

*

to ^3 03

CO

•

rH

CO

CO

t> LO •

rH

o

o o

GO • CD CD

o-

O

co

GO • rH

^

CO

to

t(

f i

in CM • in CO

03 in •

03 c• CO

CO rH rH

t-s •

co co

M3 cn in

o Mo

rH a •»H 'd in

P

8

O

CM

. •

^1

■P 03

-p

CO

o co o o CO P

to

co 03 • cn

tO

CO

CO

•

03

tO

P

in

p

cn •

s CM

CD CO

IN

CM •

rH

rH

rH

•

o

to o

•

P

> |W|

to CD • o

CO

• 03

CD CD

•

CD CO

5? -i 4 03 1o •P rH O 43

CO

•

o •

o •

CM CO

CO

03

00

$

oo

IS

co 03

CM

o>

&

CO

to

•t*

CM

CD CD • 00

V) cn I

M

tOI RI

O

CO

a

rH

00

CD

W

k

h |w

CO

CO

03

cn

H

CO

rH

01 05

P

to

rH IO CM

0

CM •

CO W

• H-J

CO

tj*

CO

CO

O

lO

rH

03 rH

•Ji

•H

•

58

o o

rH

S3

•

• 4*

03 cn o

o

CM CM

5 8 O

in W in w in in co co m to • • to COrl H W to CD rH rH CO n co co d 03 (S 03 03 •rH •rH & >» u 'd 'd •S-, >•o- p CO > » o p t~ P IS 03 CO 03

H }t

8

in

rH rH

00

o

to

g

CO 8

03 o

CO

The results of the investigations of the 9° and 45° isotherms are presented in Tables 8 and 9 respectively and are plotted in Figures 11 and lii respectively.

At

the lower temperature, the pure salts exhibited a solu­ bility in nine hours of stirring that was within one part per thousand of the reported solubility and so it is felt thet the minimum seven day stirring period for ternary complexes was more than ample to insure equilibrium.

The

results of the 45° study indicate that equilibrium is established within two days.

After a period of stirring

of a week at 45°, a solubility tube of the pure potassium salt showed no sulfate while that of the ammonium salt gave a positive test.

However, the amount of hydrolysis

to which this corresponded was negligible.

92

-> •H

to

o> GO •

in is •

3 to •

s

(O CO CD

Hj* CO 00 •

03 rH 00

p*

P•

rH to

CD

CO to

03 CO

q>

o

CO

in

CM CM

in•

rH CM CM •

03

rH

G>

8

Oi

o

8

8

O

rH PP-

CO p* IS •

O

GO CO p-

rH O P•

to in

to

GO o>

o 03

03

to 03

O

CD

Pi

■P CD

o

w 49W • CO

+» rH

00

03 & LO to

CO

03

EH CO

s CO

a

®

CO

to

tO

t» ■P CD

• Oi

Oi • rH

•

O O

o

to

in is CO • rH

rH

00 •

•

in

•

rH

to

Oi

•

in

rH

in

• rH

to

IS IS

CO frfr•

03

to to

•

o• 03 m

CO

• rH

to

pp-

•

• rH

•

o

00

8

• rH

IS •

•

rH •

a*

oi

Oi

O)

• rH Oi

rH GO •

to to

03 to

pco •

GO Oi •

o

tO

in

« rH

•

o

03

• 00

to

to

in

p-

in

O

8 rH

in to CO

8 CO

rH

rH

rH

rH

rH

rH

rH

CO

$

PIS

CO to •

in to

to

•

r—1

in

Oi •

in

in

s

O CO

o

«4< CD

in

tn

in

•tf

in

Hj*

Si CM O •

'I*

Oi •

Cl

0-

IO

rH rH •

O rH

s

8

cr>

oto

CO IO in

rH rH

•H +» to •rt R

►»

T-t Jw Of M CO 5= •

CO

03 00

c*

IS C^

• to

0> tO

>

• C\J

rH •

a

Cl

•

in

in

in

Cl to • to 03

rH

CM oc •

co

0

0

to

• Cl IS

to

to O i tO

to

CM

Oi

OJ

O

O

CO

•

• in

CJi

00

to

•

O

•

d

•

IS

ci

O o in

'w

cm

c

(O

01

c^ •

01

CO

8

in 02

CO

■S %

0

tD 03

O/ • to 03

Oi in •

s

in

CO •

in

0

0

0

0

03 • tG 0)

rH •

3

•

8

0 « « CD CM

CO CO

to

8

to

>

cn 0)

*p

(X)

•cj*

Cl

g in

gi

•>*

O O-

in

00

> In

to

CT**rl

CM

8

I S rH

00

o>

CM O

Oi 00

Cl

CO CM

CM CO

03 GO • CO

rH

0 0 o^ o>

00

d-

co

o

o

6 rH

O• 03 to

Cl Cl

CO 0 •

Cl

03 00 IS 0

rH

00

® R •H rH

I

O

|co

r

H M ©\P.

rH &

°

•

rH

• rH CO

CD

CO

CM O • IS CO

4-> a •rl O Ph

cn 01

oco

rH

cn a • 0)

8•

*

oGO •

i«2

Cl

to

rH

O • 03 tO

in

O

Cl

•

•

•

O in

Cl

cn

00 Cl

•

0

•

ri

> a

H

0 0

•

to Cl •

CO CO

to

rH

CM

CO

tO

to

CM

CO

0

01

IH

© ■5

CD I M

CM

tiI

«

CM

g o 52;

in

I

CM

CO

o>

g

g in

g 00

0

■d* • to in

ir'*r.jruiur c-o •

96

The component solids in the system ammonium sul­ famate - potassium sulfamate - water are seen to undergo mutual solution to a limited extent.

The limits of the

two solid solution series at the temperatures studied, as obtained by mathematical extrapolation of tie-lines, are as follows: at

9°

50.8# KS

and

8.5# NHAS

25°

5*.5# KS

and

9.1# NH4§

45°

52.7# KS

and

9.8# NH 4S

The mutual solubilities of the solids are seen to increase with increasing temperature, as is expected.

Since the

first series seemed to end at 25° close to the value for the 1:1 compound (» 54.2# KS), the system was also studied at 96 and 450 to see if this limiting composition remained constant.

This series is seen to increase from

9° to 25° but then to remain approximately constant over the next 20° at a value somewhat short of that of the 1:1 compound.

However, since the limit is short of the com­

pound, the compound is ruled out.

An estimate of the solid

solution limits, by judging, for each temperature, the intersection of each of the two curves with the fixed horizontal line in the distribution plot of Figure 13, does not give appreciably different values from those obtained by tie-line extrapolation.

If the limit had been at (or beyond)

the compound, one would expect -- if the compound were actually involved — value.

a convergence of tie-lines around the 1:1

However, there was found no evidence of the

occurrence

of

a

compound

at

97

•the 1:1 value, as would be shown by the bunching of tielines at the composition of the suspected phase., e.g.. in the system AgBr05 - NaBrOj - HgO at 25° 61.

Tie-lines just

within the solid solution region are almost parallel to the limiting tie-line and point quite clearly away from the limiting value at all temperatures.

Further, if the 1:1

compound were the limit to the solid solution, then this would require that that compound was decomposed as the temperature is lowered.

This could be the case if one of

the components were volatile so that lowering the tem­ perature would lower the partial pressure and therefore the activity of that component to a value insufficient for the formation of the compound, e.g.. the Pd - H2 system^. This is not the case in the sulfamate system.

The decompo­

sition of the racemate formed by d and 1 potassium acid tartrates as the temperature is lowered (investigated by ±&steur) is a very rare phenomenon.

A similar effect is

observed as the compound Fe2Sn is decomposed below a a _ as is CaCl^-KCl43 certain t e m p e r a t u r e A n o t h e r reason for the belief that the solid solution is not halting at the 1:1 compound is the shape of the solubility curve, viz.. convex down. This is the kind of curve one expects as the tie-lines spread out on the base.

Thus, the solid is not tending to

remain constant in composition. The similarity of the * over-all shape of the curve to the curve in the system IfeClOj - NaBrOj. - H^O at 25° is noteworthy^4. The relative distribution of the salts between the liquid and solid solutions is interesting, even though

98

discontinuous, and is presented for the three temperatures in Figure 13.

y is the ratio of mols of NH^S per mol of

salt in the liquid and x is the same ratio for the solid. It is seen that the distributions are regularly displaced for the three temperatures, at each of which temperature there is formed a well-delineated plateau indicating the coexistence of two solids thereby holding the liquid composition and in fact the system constant, as required by the Hiase Rule.

Horizontal lines indicate the coexist­

ence of two solids accompanying liquid invariance, vertical portions would indicate constancy of solid composition, while inclined portions indicate change in both liquid and solid composition.

The values of y and x for the various

temperatures are as follows:

TOC

y

9°

0.776

0.094

0.535

*5°

0.736

0.100

0.520

45°

0 .095

0.125

0.512

XI

x2

As is normally expected, the mutual solubilities increase with increase in temperature and so the plateaus are seen to decrease witn increasing temperature until they would presumably disappear at a higher temperature, as shown schematically by the dashed curve in Figure 13* It is clear from the earlier discussion and from ex­ amination of the distribution curve that this system be­ longs to type V in which the positive deviations in the solid state have led to partial miscibility in the solid

j 5

o

CM

Q

(/> v/ ^ |

O O O o o o O) ID in CJ

. O n

•H LU

n

2

o

o

CM

*o

100

phase.

At a higher temperature, the system would pass

probably into type 1, or into type II, i.e.. if decomposi­ tion and melting did not interfere.

It is also apparent

that the invariant liquid is incongruently saturated, a characteristic of type V ternary invariant liquids.

Thus

if g represents the composition of this invariant liquid and x and y represent the compositions of the ammonium and potassium rich saturating solid solutions, respectively, then it is seen that the ammonium sulfamate salt proportion will be greater for the liquid g than for the solids x and y. Isothermal equilibrium evaporation of liquids with ammonium sulfamate salt proportions smaller than y will yield solid solutions of the potassium rich phase.

The compositions of

these solid solutions will vary as evaporation proceeds, becoming richer in the ammonium salt.

The liquids saturated

by these solids will move toward g, although they will dry up before reaching g.

The limit for this group of liquids

will be liquid of salt proportions equal to y:

this system

will dry up with liquid at g and with y as the single saturating phase.

Liquid of salt proportions between x and

y will also dry up at g but there will be two saturating phases:

x and y.

On drying, these liquids will first

precipitate a potassium rich solid, the composition of which will move toward y (as drying proceeds and) as the liquid composition moves toward g.

When the liquid reaches g and

the solid reaches y, the system will be held invariant as the new phase, solid x, appears.

The reaction on drying

101

this system is as follows:

g plus y gives x plus water vapor.

Thus solid y is converted to solid x.

Liquid of salt pro­

portions lying between x and g will go through exactly the same transitions:

a potassium rich phase appears which

changes to y; then x appears; and then all of y is converted to x.

However, invariant liquid g is not entirely consumed

and ao drying proceeds as g moves toward a saturated solu­ tion of ammonium sulfamate (called c) and as the solid phase becomes richer in the ammonium salt.

Liquid of salt pro­

portions between g and c dry up along this curve, moving toward c as the single solid phase becomes richer in the ammonium salt, but there appears only the single solid solution a3 a saturating phase.

Since, from the inclina­

tions of the tie-lines, the salt proportions for all ternary liquids in this system are seen to be higher in the ammonium salt than for the corresponding saturating solid solutions, it is seen that all ternary liquids become richer in ammonium sulfamate on isothermal equilibrium drying.

Thus

it can be deduced that the vapor pressure of the invariant liquid, g, is lower than that of liquids saturated with respect to the potassium rich phase, and that liquid of the lowest vapor pressure in this system will be a solution of pure ammonium sulfamate, c.

It also follows from the

earlier discussion that the binary system NH^S - KS probably would exhibit peritectic behavior on fusion. Finally it should be pointed out that although the mutual solubility of the solids increases with temperature

rise, the intervening (ternary) invariant region need not necessarily disappear, if the two salts are of different crystalline structure.

The different shapes of the

solubility curves for this system indicate that the solubility curves might not be related and that therefore the salts may be different in species.

As would be ex­

pected for two hydrates (and therefore of different crystalline structure) and has been shown*^ the solid phase is separated by a transition gap.

Thus it is not

required, from the observation of mutual but limited solid solution, that the two salts have the same crystal structure.

103

The System (NH4)2SC>4 - K2S04 - H20 at 250c The data for this system are presented in Table 10 and plotted in Figure 14.

The data are those of Weston°9 and

of Hill and Loucks^® with a few points obtained by the present writer.

Hill proved attainment of equilibrium in

the systems he studied, within three days; the present writer proved attainment of equilibrium within one day by constancy of results on repeated analyses after a longer period of stirring; whereas Weston stirred his complexes "for some hours" and then analyzed the liquid and the settled residue.

All the data for the liquids fail on the

same smooth saturation curve.

The salts are seen to form

a continuous solid solution over the entire composition range.

Tie-line inclinations indicate that the system be­

longs to type I, with ammonium sulfate salt proportions being greater in the liquid than in the saturating solid for all the ternary complexes of this system.

Thus, on iso­

thermal equilibrium evaporation, all liquids and solids be­ come richer in the ammonium salt.

The distribution curve for

this system lies above the diagonal over the entire range, as it does for the corresponding sulfamate system (shown in Figure 13), but here it is smooth rather than broken.

Also

from the tie-line inclinations, and (the same thing) from the knowledge that saturated liquids will always become richer in the ammonium salt upon isothermal equilibrium evapo ration, one may deduce that a saturated solution of the pure ammonium salt will exhibit the lowest vapor pressure.

104

o o M05

CO

to

25°0

03

CO

CO

CO

CO CM

CO

o

CO

ot£

oo t

D*-

CO

0

O

CO

•H

10

^+t

to

CO

1C

CO

03

m

cr>

01

O

O

o>

LO 0•-

O

^pP

CO tJ o •r! CO

(11114)gS0A ~ In the .system Equilibria

Table

Bo.

10

~

rH

O'

CM

CO

+3 3*' P d o)

8 rH

CO CO

O

CO

Ifj 0} c o CD

to o: 03

to

03 rH

CO

o-

to

10

lO

to

o-

CO

0

0

0CO

rH CO t CO

CO

O

03

• to rH

CO

to

H

• 0

03 to • 03

to

05

03 CO

CO • eg

0 0

•

cc

co • CO

0

•

03 •

•

•

IfJ •

03

•

eg *

•

0

O

CO

n •3 HOrl* P CO D< 05 •H * £

•

CO < {

to • 0 ) CO

®

U

•4t3 o

rH

o o0 0

CJ*"H « n

rH

ten

tO

CO

s

CO

03

o CO in

CD

Oi

m

8

in

8

rH in

CM CO

CM

in *9

O rH

rH s co

rH

rH

CM 00

rH

CO

in

rH CO

0 is

O

03 Q

CO CO CO

d
4 43 •H CO fl 0)

O

CO rH

3 rH O CO

%

in

o

CO

•

o

IN

m •

O

IN

CM

•

o N

• Oi CO

in

•

o

c^

rj*

in

•

o•

CO Oi

tD

to

Oi

Oi to

Oi to

00

in

o in d»

03 in d


CO

rH

CO

rH

rH

§

£

►H O

§

03

-4 03

03 03

CO

«

03

a CO

9

O in

9

in

03

\

110

tJ d

cdr* o

&

=

t* to to •

oto o • in

60

O

o

W

>

3f S*M

8

Cv3

to

til .

t}l .

0


01

t i ID

•d O C

ft ■P d rH O

03 CO rH

(D

■s° YP. rH tO

Oi «H

• +> rH • PO O

^ tl OT

Ol

0-

Ol tC

tO

to rto m

o o • 00 in

00 • to in

• to in

o • 03 in

tO

rH

CO

CO to 5f«

in to

O i

9

s a;

CD rH

a O i

si

O i

ro in

2

•

•

03 tli

9

oco

to CO

o03 in co

rH

Oi rH

in 00 • in

in • in

•

« to

O i

in

00

8 CO

o I—I CO

to •

O i

-p

•H « d Q> P

3 CO

to

CO

tjl

&

to

in

03 03

to

Ol

in 03

in o-

03

l.SHC,

rH

in

1*00

>

in rH

►

d

o

&

CD

to

in

n

oo

to

«4

oo

CO

d

o

n o

CO

CO

fH

1

lll'i

«d d

$

cS o co co

to o o> • «

-Si rj W CS 01

03

m 03

rH

03

rH

D-

X

Os

00 •

03 •

O *M

© d rH O

C^ . 03

&

tC o)

i—Id •i p

OS

4*

> CO

Va1pw

..

03

tO OS

in ■ • 00 to

in 00 • 00 to

Os

os

to

CO os os

rH

h;

in in 00 •

m 00 •

to CO to

d

ww

0

— '

0)

ry

a ? &

0> CO O

02

C)

in

01 4J :J

COl tu 1 03

to to o

£

o

to to o

8

i—t CD O

in m o

o in o

in in o

in

o■ci*

rH

CD m

oo

tp -cF c5

s

03 m o

o ■dO

i—1 CO o

to

CD O

C2

CO CD c

CD O

00

in oo

CO

cn

03 O

to

o

o

8

S3

03

o

•H ■P 2

r~f O

w

O VI

02

El

Vi

1

•P 6 a>

ci to

S e o +3 e> co

d

EH

J 3 +j Pi

0

col

o o

CO 1 —1

03

03

o

03

f—1

o

CO

o

in

O

o

O

CO

in

o

o

en o-

03 o-

03 CO O

o CD O

CD O

rH cd

>

•H 2 01

Jxl W

I o

si

in c\j 02

03

C‘

00

e'­

oo

03

00

to 03 O-

$ 02

?3 02

03

< 1 Ol

r -1

03

"3


rt o

£ Q H Cd

o O-

K

0-

cC

o-

03

to

tO

CO

CO

CO

CO

in

0-

o-

in 03 03

03 03 CQ

o o

o

P P tO

o

CP in rH

to

c-

r-

CD

>

o

HQ Ti

system, is impossible.

Therefore, solid, y is consumed

before solid 1 is consumed, either a) before the liquid reaches Q or b) after it reaches Q. a) The solid y is consumed before Q is reached, if P is in Qgx.

This happens when P (near P^) comes xo be

swept or crossed by the line lx of the three-phase triangle lxy. alone.

For P^>, this leaves 1^ saturated with x

But P^, being in cxze, must end up as a mixture

of solids, one on cx and the other on ez.

Continued

evaporation therefore causes the single solid to move from x toward c while the liquid crosses the II field to curve aQ but side-passes the invariant Q. When liquid reaches aQ, a solid on ez appears, giving rise to a three-phase triangle 1 (on aQ), s (on cx), and s ’ (on ez), such that the line Is passes through P^.

With continued evaporation,

1 moves from Q toward a, s moves toward c, and s ’ moves toward e.

When the line s - s' passes through P^, liquid

vanishes to leave two solids. b) Q is reached before y is consumed if P is in Qxy. 3)

Equilibrium isothermal evaporation of liquids on

curve aQ causes the precipitation of a mixture of two solid solutions, one on cx and the other on ez. three-phase triangle starts as

xqz

The

and ends as cae.

aQ is reached by all liquids with P in axze.

Curve

If P is in

127

Qhz, the liquid first reaches curve by as explained under 1). disappears. consumed.

At Q, the liquid is invariant, and solid y Then the liquid travels along aQ until 1 is If P is in Qgx, the liquid first reaches aQ,

cfosses it (without reaching

q ),

and then reaches aQ.

If

P is in Qxh, the liquid first reaches gQ; then Q, where y is consumed; and then moves along aQ toward a.

If P is

in aQze, the liquid reaches aQ directly from below.

If

P is in field II, the liquid reaches aQ directly from above. In every case, on every curve, the liquid is consumed when the solid - solid line of the three-phase triangle sweeps through the fixed point P. 4)

Point Q is reached by all solutions with P in

the invariant quadrangle Qxyz.

Those in Qxy reach it from

curve gQ carrying solids x and y; those in Qyz reach it from curve bQ carrying a solid on yd and one on zf.

When

liquid reaches Q the invariant reaction is liquid (Q) plus y to give x plus z plus water vapor. For 1hose liquids with P in Qxz, y is consumed before the liquid, leaving 1 plus x plus z and then liquid moves from Q toward a until it dries up.

Por -those liquids in xyz

liquid is consumed at Q leaving x plus y plus z.

The diagonal

Qy is significant for the direction of approach to Q, the diagonal xz is significant for what happens when either 1 or y is consumed.

Point Q is therefore the incongruent

drying up point for solutions with original salt proportions

128 In xyz.

Por liquids in Qxz, it is only a transition point.

Now let us consider what happens bn mixing NH^S and K2S0^ in equivalent proportions (point o), or (the same thing) (NH^)2S0^ and K$ at point o; then dissolving completely and evaporating to dryness.

Por the position of z as assumed,

the reactions are as follows: on zf since point o is in Qbfz.

The first solid to appear is Then the curve Qb is reached,

and the^second solid, on yd appears.

Then Q is reached, when

the first solid reaches z, the second reaches y, and a third solid x appears.

Then the reaction; liquid plus y to give

x plus z, proceeds until y disappears to leave x plus z+liquid. *vA so(4 wvowf> ^»»»>^

frotw

Drying proceeds as -Qvmoves Qtoward a, q| toward om andAz toward e.

The system finally dries up with the last liquid on aQ

in equilibrium with two solid solutions, one on jjc and the other on ez.

But the position of z is not definitely known,

and the drying up behavior of an equivalent mixture of a reciprocal salt pair with water (i.e.. at o) depends upon the position of o relative to the diagram: yd|z, xyz, qxz or aQze.

whether in

If z were to the left of the exten­

sion of yo and therefore o were in ydfz, the process would never involve more than two solids, one on yd and one on zf, and the liquid would dry up on bQ before reaching the in­ variant liquid Q, as did point

(p.124).

If z were to the

right of the extension of yo and to the left of the extension of xo and therefore o were in xyz, then liquid would dry up at q as it reacts

with y to give x plus z, until all the

liquid is gone, leaving the three solids x plus y plus z. If z were to the right of the extension of xo and (therefore) o were in Qxz as is drawn in Figure 17, the system would dry

up as described above with liquid on aQ and a solid on go and on ez, but the system would go through the "vapor pressure .halt", so to speak, caused by the reaction given above.

If

z were to the right of the extension of Qo (just barely possible) and therefore o were in aQze, liquid would, as in the preceeding case, dry up on aQ (but now closer to a) with a solid solution on gc and on ez.

However, now, as opposed

to the preceeding three cases, Ihe first solid to appear would be on ez, and the next solid to appear would be on gc; d-y- ae opposed to tho •proooeding two cases,— Q-ncver appoare)and, as opposed to the preceeding case, therefore, there occurs no invariant conditions for a "vapor pressure halt", as would be caused by the coexistence of four phases. It may be mentioned that the position of z as assumed in Figure 17 is probably approximately correct for it was estimated as follows:

The salt proportions of sulfamates in

the invariant liquid Q is seen to be approximately the same as for the ternary invariant g.

It is assumed that the sul­

fate salt proportions in Q are' the same as in the ternary sulfate liquid, say m, saturated by z in the absence of any sulfamate; i.e.. that m has approximately the same ammonium to potassium ratio as Q and as g; or that mQg forms approxi­ mately a straight line, parallel to the sides of the square. The value of m being assumed then to be approximately 25 equivalent per cent of potassium, this is converted to a weight basis (s 31% K^SO^), a line is drawn in Figure 14 from this basal point to the water axis, cutting the solubility curve at 8% K*S04 and 18% (NH^J^SO^.,

The saturating solid

phase is estimated from the slope of contiguous tie-lines to

130 contain 77 weight per cent or 72 equivalent per cent of K 2SO4. This was the basis of locating point z in Figure 17. The system just presented could be considered as one in a series of quaternary systems involving solid solutions. The J&necke projection for the first of this series, in which all the. salts are completely miscible in each other, would contain no curves or lines of saturation with respect to more than one phase.

In the second of the series, in

which there occurs two series of complete solid solution between two neighboring salt pairs, there would be seen (Figure 18a) a single curve smoothly crossing the diagram which represents liquid saturated with respect to two solids of continuously variable composition (top p.33).

In the

third of the series, in which there occurs a complete and a pair of limited solid solutions, there will be seen three curves which meet at a quaternary invariant liquid.

If this

invariant liquid lies outside the triangle xyz, as for the I system just reported, it is incongruently saturated (Figure 18b), whereas if it lies within the triangle, it is congruently saturated (Figure 18c).

Also, incongruently

saturated ternary invariant liquid may become congruentely saturated in the quaternary system (see Figure 18d and p.121); and congruently saturated ternary invariant liquid may become incongruently saturated in the quaternary system (Figure 18e).

J&necke projections of the fourth of this

series of systems, in which there would exist four limited solid solutions, would look somewhat like a projection for the simplest reciprocal salt pair.

The saturation of the

invariant liquids may be both congruent (Figure 18f),

\

130a

both incongruent (figure 18g), or one congruent and the other incongruent (Figure 18h), and may go through any of the transitions illustrated in Figure 18d and 18e.

It is

clear that neither for the system presented above nor for any of the systems of the series just described can there coexist a reciprocal salt pair.

In the last of the series,

quaternary liquids are saturated with respect to pure salts, and there would occur two invariant liquids which are separated by a 'curve which in turn represents liquids saturated with respect to a reciprocal pair of pure salts. In Figure 18i, only one of the quaternary invariant liquids is congruentely saturated while in Figure 18j, both liquids are congruently saturated.

Figure 18

a

c

*153,’ ''^

i3a Summary « A systematic investigation of the aqueous behavior of the sulfamates of sodium, ammonium and potassium in the presence of halide and sulfate indicates that these systems are simple, with no solid solution or double salt occurring.

Sulfate and sulfamate ions therefore are

found not to replace each other in the solid structure. Partial solid solution is found to occur between ammonium and potassium sulfamates, studied in aqueous medium at 9°, 25° and 45°.

The distribution of the salts

in the liquid and solid solutions is such that the system belongs to type V.

Prom this, it is shown that the binary

system is probably peritectic. The reciprocal salt pair ammonium sulfamate and potassium sulfate with water is studied at 25°.

Equilib­

rium isothermal evaporation is discussed for the quaternary system and for its component systems, but in a limited manner because the exact composition of the saturating phases was not determined.

Por liquids undergoing evapora­

tion, curve gQ is shown to be a transition curve.

Because

of the complete solid solution of the sulfates (type I) and the limited solid solution of the sulfamates (type V), there does not exist a stable salt pair in this system. Rather there is observed the coexistence of at least two solid solutions.

The three curves, each representing

saturation with respect to two solid solutions, meet at the single incongruently saturated invariant liquid.

133 Bibliography 1.

M.E.Cupery & Iff.E.Gordon, Ind. and Eng.Chem. ]54, 792 (1942)

2.

L.P.Audrieth, M.Sveda, H.H.Sisler & M.J.Butler, Chem. Rev. 2o, 49 (1940)

3.

P.Oberhauser B. & H.E.Urbina G., Chem.Abs. 41, 1944f (1947)

4.

J.E.Ricci & B.Selikson, J.A.C.S. 69, 995 (1947)

5.

V.Mellor, "Comprehensive Treatise on Inorganic & Theoretical Chemistry" Vol.8 p638

' 6.

M.E.Cupery, Ind. and Eng.Chem. 30, 627 (1938)

7.

C.J.Brown & E.G.Cox, J.Chem.Soc. 1-1. 1 (1940)

8.

E.Divers A T.Haga, ibid. 69. 1634 (1896)

9.

M.J.Butler, G.E.Smith A L.P.Audrieth, Ind. and Eng.Chem., Anal.Ed. 10, 690 (1938)

10.

P.Oberhauser B. A H.E.Urbina C., Chem.Abs. 41, 1944g (1947)

11. E.Berglund, J.Chem.Soc. 34, 043 (1878) 12. M.J.Butler & L.P.Audrieth, J.A.C.S. 61, 914 (1939) 13.

G.B.RLng A J.P.Hooper, J.Phys.Chem. 4j>, 938 (1941)

14.

A.P.SchmeMLe A J.E.Westfall, ibid. 48. 165 (1944)

15.

E.M.Baker, J.A.C.S. 71, 333b (1949)

lo.

W.E.Gordon A M.E.Cupery, Ind. and Eng.Chem. 31, 1237 (1939)

17.

S.Uchida, J.Chem.Soc.Japan 6£, 504 (1942)

18.

V.Mellor, "Comprehensive Treatise on Inorganic & Theoretical Chemistry" Vol.8 p641

19.

D.M.Yost & H.Russell, "Systematic Inorganic Chemistry", Prentice Hall, U.Y. (1944) p99

20 .

Ketelaar & Heilman, Z.Krist 103. 41 (1940)

21. S.H.Laning & P.A.van der Meulen, J.A.C.S. 6^, 1828 (1947)

134

22.

F. A. H. Schreinemakers, Z.physik.Chem. 11, 76 (1893)

23.

A.E.Hill & J.E.Ricci, J.A.C.S. ££, 4306 (1931)

24*

G.Beck, J.prakt. Chem. 156. 227 (1940)

25.

J.A.Thelin

& P.A.van der Meulen, J.A.C.S. 70,1796 (1948)

26.

S.H.Laning

& P.A.van der Meulen, ibid.70. 1799 (1948)

27.

J.H.van't Hoff, quoted in S. Glasstone, "Textbook of Physical Chemistry" McGraw-Hill N.Y. (1940) p734

28.

A.E.Hill &

C.M.Loucks, J.A.C.S., 5£, 2094 (1937)

29. A.Friedman,i-h.D. Thesis, Hep't.of Chem.,

New York

University (Heights),N.Y., (1948) 30.

H.A.Doerner & W.H.Hoskins, J.A.C.S. ££, 662 (1925)

31.

B.Roozeboom, Z.physik*Chem. 22., 385 (1899)

32.

J.C.Slater, "Introduction to Chemical Physics", McGraw-Hill, N.Y. , (1939)

33.

R.W.Gurney, "Introduction to Statistical Mechanics", McGraw-Hill, N.Y., (1949)

34.

J. A. Y. Butler, "The Fundamentals of Chemical Thermodynamics", Part II, MacMillan, London, 1934 (but not second edition (1949))

35.

J.A. Y.Butler, "Commentary on the Scientific Writings of J. Willard Gibbs", Vol.I, Yale University Press, New Haven (1936)

3o.

P. A.H. Schreinemakers, ibid.

37.

B.Roozeboom, Z.physik.Chem. 8, 521 (1891)

38. . W.C.Blasdale "Equilibria in Saturated Salt Solutions" Chemical Catalog Co., N.Y. (1927)

135

39.

A.E.Hill & N.Kaplan, J.A.C.S. 60, 550 (1939)

40.

A. E.Hill & G.S.Durham & J.E.Ricci, ibid. 62,2723(1940)

41.

J.E.Ricci & S.H.Smiley, ibid. 66, 1011 (1944)

42.

W.C.Blasdale, Ref.38, pl46

43.

G.0.As3arsson, J.A.C.S. 72, 1433 (1950)

44.

E.Janecke, Z.anorg.Chem. jjl, 132 (1906); %2, 358 (1907); £3, 319 (1907)

45.

H.LeChatelier, Compt.rend. 172, 345 (1921)

46.

W.C.Blasdale, Ref.38, pl46

47.

J.E.Ricci & C.M.Loucks, J.Chem.Ed. 15, 329 (1938)

48.

Iiowenherz, Z.physik.Chem. 13, 459 (1894)

49.

W.C.Blasdale, Ref.38, pl23

50.

W. C.Blasdale, ibid. pl41

51.

E.V.Sayre, J.A.C.S. 71, 3284 (1949)

52.

G.J.Doyle & N.Davidson, ibid. 71, 3491 (1949)

53.

I.M.Kolthoff & V.A.Stenger, "Volumetric Analysis" Vol.II, Interscience, N.Y. (1947)

54.

A.Butlerov, Ann. 115. 322 (1860)

55-

E.A.Werner, J.Chem.Soc. 111. 844 (1917)

56.

P.J.Walker, "Formaldehyde", Reinhold, N.Y.(1944)

57.

E.Baur & W.Ruetschi, Helv.Chim.Acta, 24, 745 (1941)

58.

I.M.Kolthoff, .Eliarm.Weekblad, £8, 1463 (1921)

59.

K.Marcali & W.Rieman III, Ind. and Eng.Chem.,Anal.Ed. 18, 709 (1946)

60.

E. Janecke, "Gesattigte Salzlbsungen von Standpunkt der Phasenlehre", W. Knapp (1908) pl3o

/

13$

bl.

J.E.RicOi & J.J.Aleshnick, J.i.C.S. 06 , 980 (1944)

62.

A.Findlay & A.N.Campbell, "The Phase Rule and its Applications", 8th Ed., Dover, Bf.Y. 1945 Chap.VI

63.

C.J.Smithells, "Metals Reference Book", Interscience, H.Y., 1949, P321

64.

T.Swenson & J.E.Ricci, J.A.C.S. 61, 1974 (1939)

05.

Stortenbecker, Z.physik.Chem. lb, 256 (1895)

DO.

J.W.Gibbs, Trans,Conn.Acad. 3, 152 (1876) and "Collected Works", Vol.I, Longmans (1928)

07.

B.Roozeboom, "Die heterogenen Gleichgewichte vom Standpunkt# der Phasenlehren, P. Vieweg,(1901-1918)

b8.

A.C.D.Rivett, "The Hiase Rule", Oxford University Press, London (1923)

69.

A.Weston, J.Chem.Soc. 121. 1223 (1922)

LIBBARZ Of HEW T O W UNIVERSITY UNIVERSITY HEIGHT!5