BEHAVIOR O F ALUMINUM
IN
2057
ALKALINE SOLUTIOSS
The dashed curve for [fii111F6]4in Fig. 2 indicates that here the influence of the five ligands dominates a t small ;values of 12. Contrary to the configuration
0xide-Coated Electrodes. 11.
of the dipoles calculated for the other four complex ions, [y111F6]4-here shows the positive pole of the induced dipole of an F- ligand faces the central ion.
Alumiiium in Alkaline Solutions
and the N,ature of the Aluminate Ion'
by Robert C. Plumb and James W. Swaine, Jr. Worcester Polytechnic Institute, Worcester, Massachusetts
(Rewined October 1 , 1963)
An investigation of the electrode potentials of evaporated aluminum films as a function of oxidation time has given the reversible electrode potentials for the aluminum-aluminate ion system. The potentials ha,ve been studied over the pH range from 10.5 to 13.8. Using the observed dependence of electrode potential upon the over-all concentration of aluminum in the solution and the pH, it is shown that the electrode reaction below a pH of 12.4 consists of the addition of aluminum ions from the metal to the aluminzte anion. Above a pH of 12.4 the reaction is directly between hydroxyl ions in solution and the metal ions from the metal. Possible structures for the aluminate anion are discussed. The aluminate ion is shown to be a polymeric anion of composition (OH-)2[A1(OH)4-], in which the Al(OH)4-units are, most likely, linked together by two hydroxyl bridges and each aluminum ion is surrounded by six hydroxyl ions. The chain length is of the order of n = 40 to n = 100.
Introduction A technique for determining the effects of oxide films upon electrode potentials was described in a previous publication.2 'Those studies showed that the electrode potentials of alumin.um in acidic and neutral solutions varied in a systematic way with the period of aiir oxidation. Electrode potentials measured after known and variable periods of air oxidation could be extrapolated to zero oxida,tion time to obtain potentials characteristic of the electrode in particular solutions. Using (1) an assumed electrode reaction and (2) wellestablished free energies of formation and of reaction determined by thermochemical techniques, one can predict a priori thrlee quantities : (1) the standard electrode poten.tia1 far the reaction, (2) the variation. of electrode potential with pH, and (3) the variation of
electrode potential with metal ion concentration in solution. The potentials obtained by extrapolating the experimentally measured potentials to zero oxidation time agree quantitatively with those potentials predicted a priori from thermodynamic data in all three respects showing that the observed potentials are thermodynamic potentials. Accepting the capability of measuring reversible potentials, one can anticipate using these potentials to study the nature of ions in solution. This communication describes our findings on the behavior of the
(1) From a dissertation submitted by J. W. S. in partial fulfillment of the requirements for the degree of Doctor of Philosophy in chemistry.
(2) R.
c. Plumb, J . P h y s .
Chem., 66, 866 (1962).
Volume 68, Number 8
A u g u s t , 1964
2058
aluminum electrode in alkaline solutions and the nature of the aluminate ion. The nature of the aluminate ioii has been the subject of much discussion and speculation. It has generally been assumed to be monoiiuclear, such as A102-, or a hydrated form of this species. Using potentiometric titrations, in which the number of “bound” hydroxyl ions per aluminum ion is determined from the measured pH and the known additions of base, Sillen and co-workers3have concluded that, at least at high pH, the aluminate ion is niononuclear and of the form Al(OH)4-. There seems to be little doubt about the correctness of the assertion that the ligands contribute approximately four negative charges for each tripositive aluminum ion. It is not clear that the potentiometric titration is sufficiently sensitive to determine whether or not the ligand-metal ion ratio results from an associated species containing a larger number of ligand and metal ions or from the mononuclear species proposed. Bode4 found no evidence for colloidal particles in sodium aluminate solution in an investigation using light scattering measurements. He found that the mobility of the aluminate ion was comparable with that of the lithium cation and the acetate anions, and on the basis of conductivity measurements found no evidence for the existence of ionic species other than AIOz-. Contrary to these observations Brintzingers found, by dialysis in KOH solutions, the molecular weight of the aluminate ion to be consistent with either A12(OH)lo-4 or A1202(OH)s-6. Jahrs found it necessary to assume the existeiice of both polymerized and monomeric forms, depending on the pH, in order to interpret diffusion in aluminate solutions. His results indicated that aluminum had a coordination number of four or six depending upon the conditioiis prevailing in the solution. Recently, Lippincott7 and co-workers measured the Rarnan spectra of the aluminate ion and found it consistent with T d symmetry even though us, of symmetry species Fz, which should be active in the Raman effect, was not observed. K.m.r. studies of the aluminum cation in aqueous solutions have not been capable of distinguishing between a coordination number of four and six.8,sa Baldwin and Taubeg concluded that the hydration number of the aluminum ion in aqueous solution was about six from isotopic dilution measurements in a flow system. It is apparent that the nature of the aluminate ion is still not clear. We will show, in the discussion to follow, that the reversible potentials of aluminum in aluminate ion solution cannot be interpreted in terms of electrode reactions involving the collventiorial niodel of a tetrahedrally coordinated aluminate ioii. The potentials
T h e J o u r n a l of Physical Chemistry
ROBERT C. PLUMB ASD JAMES W. SWAINE, JR.
may be interpreted quantitatively (both with respect to dependence on metal ion concentration and dependence on pH) by a model in which the metal ions are octahedrally coordinated and linked together in a polymeric anion structure by hydroxyl bridges.
Experimental The General Experimental Technique. The experimental apparatus used for studying the effects of oxide films upon electrode potentials has been described in a previous publication2 and only the essential features of the experiments will be reviewed here. Initially oxide-free metal surfaces are generated by evaporating the metal from a filament onto a glass microscope slide under vacuum. In a rapid sequence of events the vacuum is broken with ambient atmosphere, the sample is removed from the vacuum chamber, and, after a predetermined lapse of time, an electrolyte in contact with a reference electrode is brought in contact with the metal surface and the transient in potential is recorded on an oscilloscope. The lapse of time between the first exposure to the atmosphere and the contact with the electrolyte is variable upward from a minimum of 0.2 sec. There are brief initial phase charging transients which may be corrected for by extrapolation, there are changes in potential as a result of oxidation or dissolution of oxide in the electrolyte, and there are effects upon the potential as a result of the oxide film present when the cell is first completed. The only major modification of the transient potential apparatus has been a change in the cell design to increase the capabilities of the equipment. Whereas in the previous design one data point, at a particular oxide film thickness and with a particular electrolyte, was obtained, it is now possible to obtain three separate data points on a particular metal film evaporation. This has been achieved by forming three separate miniature cells with the single common metal film. The Transients in Potential. Two distinguishable types of transients were obtained with aluminum in basic solutions. The type of transient observed in the pH range from 10.5 to 11.1 is shown in Fig. la. In this (3) (a) C. Brosset, Acta Chem. Scand., 6 , 910 (1952); (b) C. Brosset, G. Biedermann, and L. G. Sill&, ibid., 8, 1917 (1954). (4) H. Bode, 2. anorg. allgem. Chem., 269, 44 (1952). (5) H. Brintainger, ibid., 256, 98 (1948). (6) K. F. Jahr, Naturwissenschaften, 38, 302 (1961). (7) E. R. Lippincott, J. A. Psellos, and M. C. Tobin, J. Chem. Phys., 20, 536 (1952). (8) J. 9.Jackson. J. F. Lemons, and H. Taube, ibid., 32, 553 (1960). (Sa) NOTE. ~ D D E D I N PRooF.-After this paper was submitted, n.m.r measurements by Connick and Fiat indicating a coordination number of six for the aluminum ion were published [ibid., 39, 1349 (1963) 1. (9) H. W. Baldwin and H. Taube, ibid., 33, 206 (1960).
BEHAVIOR OF ALUMIXUM IN ALKALINE SOLUTIONS
2.00
d
T
I
1
1
I
I
I
I
1.90
-5 d 6
m
c
1.80
3
m
5z 1.70 e.r
$
2059
The small saturated calomel electrodes required for the miniaturized cells exhibited junction potentials as a result of contact with the alkaline aluminate electrolytes. These junction potentials amounted to 20 or 30 mv. and were reproducible for a given electrolyte. They were determined by comparison to a reference calomel electrode and the data were corrected. The potential measurements, after correction for the junction potentials, were reliable to + 5 mv. except where the transients posed special problems in analysis. Solutions Studied. Most of the solutions were prepared from H2S04 with Alz(S04)3 and the addition of NaOH to the appropriate pH. The concentration of the sulfate ion was 0.0057 M 1.5 MAlin most cases. The concentrations of the aluminum ion were 0 0072, 0.00072, and 0.000072 M . Sulfate was chosen as the principal anion because it was found to have negligible tendency to coordinate with the aluminum in comparison with the tendency of the hydroxyl ion present in the solution. The p H values of the solutions were determined with a Beckman Type E2 glass electrode using Ca(OH)z standards having a pH of 12.45 (based on NBS measurements supplied by Beckman Instruments, Inc.). Corrections to the p H were made for Na+ ion concentrations. Potentials were also measured with a phosphate electrolyte a t a pH of 11.5 but these did not agree with the sulfate electrolyte. It was demonstrated that the sulfate ion concentrations did not affect the electrode potentials by measurements a t a pH of 11.6 with an electrolyte consisting of 0.0072 M sodium aluminate which had been prepared free of sulfate. In the high pH region it was found that the potentials were independent of the aluminate ion concentration. Measurements a t high pH were made over a range of three decades of aluminate ion concentration with the ordinary sulfate electrolyte and also with a pure 1 M sodium hydroxide solution containing no sulfate or aluminate. Experimental Results. The experimental results are shown in Fig. 2. It is seen that in the lower pH region the electrode potentials depend in a systematic way upon the aluminate ion concentration and upon the pH. At a pH of about 12.4 the potential becomes less sensitive to the aluminate ion concentration and above this pH' the potential is independent of the aluminate ion concentration and changes less drastically with the pH. In the pH range from 10.5 to about 12.4 the experimental data can be fitted with the empirical relationship
+
1.60
1.50 2
4
6
8 10 Time, sec.
12
14
16
18
Figure 1. The bransients ;n potential observed when alkaline solutions are brought into contact with an evaporated aluminum film after three different periods of air oxidation. (a) Type of transient observed in pH range from 10.5 to 11.1. ( b ) Type of transient observed in pH range above 11.1.
type of transient the potential reaches a maximuim shortly after contact with the electrolyte and then dlecreases as further oxidation of the metal surfaces takes place in the solution. The initial portion of the transient is probably a phase charging transient. The time constant for this increases with increasing oxide film thickness. Extrapollation of the transient through the phase charging region to the moment of contact of the electrolyte with the metal surface gives a potential which is found to be independent of the oxidation time. The type of transient found in the p H range from 11.1 to 13.8 is shown in Fig. l b . The potential rises to a maximum which is independent of both the oxidation time and the continued contact of the electrolyte with the metal film. The time constant for the initial transient increases with increasing oxide film thickness, but the steady value of the potential is independent of the initial oxide film thickness. Probably the initial transient is a mixture of effects caused by phase charging and oxide dissolution. Below a p H of about 10.5 the transients are complicated and difficult to interpret. I n this p H region there is considerable tendency for the oxide film to grow in the electrolyte so that the metal becomes more and more passive. However, the competition between oxide dissolution and oxide growth seems to be critically balanced and we have found it difficult to get reproducible measurements.
Eaat
cal.
=
0,5099 t 0.0591 log [aluminate ion] -I- 2(0.0591) pH 3 Volume 68, Number 8
August, 1964
ROBERTC. PLUMB AND JAMES W. SWAISE,JR.
2060
2.10
0.059 E=,@"---log [Al(OH),-] 3
+ 43 (0.059) pH -
2.06
2.00
d
2
1.95
3
i 4" 1.90
; 5 I
1.85
1.75
1.7n
LL/
11
12
13
14
PH.
Figure 2. The potentials obtained by extrapolation t o zero oxide film thickness for an aluminum electrode in a variety of electrolytes. Points experimental. Curves according to empirical relationships given in text. 0 , NaOH; A, 7.16 X 10-3 M ~ 1 3 + 0, ; 7.16 x 10-4 M ~ i a + ;+, 7.16 x 10-6 M ~ 1 3 + .
The constant, 0.5099, was chosen to provide the best fit to the data in 0.00716 M solutions. A surprising feature of the results in this pH region is that the potential tends to increase with increasing aluminate ion concentration rather than decrease. The curves drawn on Fig. 2 represent this empirical equation. In the pH range from 12.4 to 13.8 the experimental data may be fitted by the equation Esat.tal.
=
0.962
+ 43-(0.0591) pH
I n this region the potentials are independent of the aluminate ion concentration.
Interpretation of Results The most obvious electrode reaction to test against the experimental data is obtained assuming that the anionic species in solution is the monomer Al(OH)d-.
+
AlMe3+ 4Hz0 -+ A1(OH)4-
+ 4H+
The potential for this reaction should be given by The Journal of Physical Chemistry
The experimental results do not agree with this expression. In the high pH range the dependence upon pH is correct but there is no observed dependence upon the concentration of the aluminate ion. In the low pH range the potential varies a t the rate of 0.059/3 v. per decade of aluminate ion concentration but the potential is observed to increase with increasing aluminate ion concentration whereas the equation predicts a decrease. I n the low p H range the observed dependence on pH (2[0.059] v. per pH unit) does not agree with that predicted by the equation. Thus, the above hypothetical electrode reaction cannot be correct. In order to explain the fact that the potential increases with increasing aluminate ion concentration it is necessary to assume that the aluminate ion is a reactant in the electrode reaction. Thus, an aluminum ion from the metal, entering the solution, must react with aluminate ions forming some type of polynuclear complex. In the high pH range, in order to explain the lack of dependence of the potential upon aluminate ion concentration it is necessary to assume that the product of the electrode reaction is in equilibrium with the aluminate ion concentration in such a way that the concentration of the product of the electrode reaction is quite independent of the over-all concentration of the aluminate ion. One may use the observed pH dependence to arrive a t the number of protons displaced by an aluminum ion entering the solution from the metal. Thus, in the low pH range six protons must be released when an aluminum atom reacts at the electrode. In the high pH range four protons must be released per aluminum atom. One may use these observations as a starting point in deducing the nature of the electrode reactions. The behavior of the electrode potentials with respect to the concentration of the aluminate ion in solution and the pH is explicable if one assumes that the aluminate ion in solution exists as a polymeric anion. Several alternative structures for the polymeric anion are possible. We will first describe the general characteristics which are necessary in order to be consistent with the experimental observations. Since the aluminate ion bears a negative charge, it is not surprising that it would react with the positively charged aluminum ion entering the solution from the metal. I n the low pH range the reaction can be considered as that of an aluminum ion from the electrode reacting with a site of a polymeric aluminate anion to produce a branched polymeric structure. The number of active sites which the polymer will present to the ion
BEHAVIOR OF ALUMINUM IN ALKALIXE SOLUTIONS
will be proportional l,o the over-all aluminate concentration in the solution. Thus, the observed dependence of the electrode potential upon aluminate ion concentration will result. The reaction of an aluminum ion with an active site must be of such a form that a net of six protons are displaced for each aluminum ion reacting. It is attractive to think that the same polymeric structure persists a,t high pH values. The observed results a t high pH may be explained if an aluminum ion from the electrode reacts with hydroxyl ions to form Al(OH),- which then reacts with a polymer chain end without any further proton displacemedt. Such a reaction will ,produce potentials which are dependent in the observed manner on the hydrogen ion concentration and, if the chains are long enough, will be approximately independent of the over-all aluminate ion concentration. The change in the nature of the reaction observed in the vicinity of p H 12.4 is reasonable. The hydroxyl ion concentration a t p H of 12.4 is of the order of the magnitude of 0.01 M . Below this pH the negativelg charged sites on the aluminate ion outnumber the hydroxyl ions in the solution and one would expect that a reaction of aluminum with the polymer directlg would predominate. Above the pH of 12.4, hydroxyl ions are outnumbering the negatively charged active sites on the polymer and one would expect that aluminum ions would react directly with hydroxyl ions. One would expect that the transition from one reaction to the other would take place a t lower pH values for dilute aluminate ion solutions and higher pH values for concentrated solutions. There is some evidence in the experimental data that such is the case. The Selection of Possible Polymeric Anion Structures. The ratio of bound hydroxyl ions to aluminum ions is useful information to establish the nature of the polymeric anion in more detail. From Brosset's data3"one finds that in the p H range above 10.5 the number of bound hydroxyl ions per aluminum ion is 4.10 f 0.05. The conditions with which possible polymeric structures must comply are that the building of bridged chains must produce six protons, and addition to chain ends must produce four protons. It is reasonable d o assume that the coordination number of the aluminum must be four or six. We will refer to skeletal structures for the polymer in which we will neglect explicit mention of the positions of hydrogen atoms. There are four simple skeletal structures having aluminum with a, coordination number of four. These involve double bridges between aluminum atoms, single bridges between aluminum atoms, alternate single and double bridges, and alternate single and triple bridges. It can be demonstrated that building branched chains from
2061
any of these four structures by reaction of aluminum ions from the metal with active sites will displace only four protons instead of the necessary six protons so that aluminum with a coordination number of four does not appear to be a reasonable possibility, a t least in the low pH region. It is possible that aluminum has a coordination number of six in the low p H region and changes to a coordination number of four a t high p H values but this does not seem likely. For octahedrally coordinated aluminum the possible skeletal structures are shown below.
It is reasonable to assume that the bridging oxygen atoms may be either 0 or OH but oxygen atoms in similar bonding positions should have the same number of hydrogen atoms attached to them. It also is reasonable to assume that all the rionbridging oxygen in a particular structure must be of one type, either OH or HzO. It can be seen that skeletal structures (a) and (d) must violate these assumptions in order to have an over-all composition of A1(OH)4-. Using the observed dependence of potential upon pH and the number of protons involved in the electrode reactions one can systen~aticallyeliminate many of the possible structures. Seven polymer structures which are consistent with the experimental facts are
r
Volume 68, Number 8
-
August, 1964
c. P L U M B A4NDJAMES w.SWAINE, JR.
2062
ROBERT
The Electrostatic Energy of a Polymeric Anion. It fs useful to estimate the electrostatic energy of a polymeric anion of the type which me have proposed because, without a priori knowledge, one might expect that the electrostatic energy would be so great that the ion could not exist. One is interested in the difference in electrostatic energy of a linear array of negative charges in an atmosphere of an equivalent number of positive charges and an array of individual positive and negative charges which are not localized as in the polymer. A crude estimate of the energy of the ionized mononieric species is obtainable by assuming a b.c.c. lattice of alternating positive and negative charges a t distances predicted from the over-all concentration of the solution (at loF3M the distance between charges is about 100 ANumber 8
Augzlst, 2964
D. R. MERKERAND B. F. DAUBERT
2064
A1(OH)4-
Psite
---f
Pbranoh
+ 2H+
that the free energy for formation of a branch is f19.4 kcal./mole. Similarly, for the reaction a t high pH one may divide the reaction Ah3+
+ 4HsO +
Pend
--+
P
+
Pend
+ 4H+
A F o = -49.84 kcal./mole
into parts and obtain for the association of the polymer Al(OH),-
+
Pend
--+ P
+
Pend
AFO = -11.84 kcal./mole
Acknowledgments. This work was supported by the U. S. Atomic Energy Commission under Research Contract AT(30-1) 2479.
The Molecular Structure in Surface Films of Unsaturated 1-Monoglycerides
on Water as Related to Three-Dimensional States'
by D. R. Merker and B. F. Daubert Department of Chemistry, University of Pittsburgh, Pittsburgh, P e n m y h a n t h (Received November $0,1969)
The relationship of the molecular structure in monolayers of unsaturated 1-monoglycerides to three-dimensional states is discussed. Monolayers of cis-monoglycerides were in the expanded state under all the conditions studied and extremely stable. The physical properties of the monolayers in this state resemble in certain respects the physical properties of three-dimensional liquids. An expanded monolayer of 1-monoelaidin was transformed under pressure to a condensed state with a limiting area of 24 This area was in close agreement with the cross-sectional area calculated for a hypothetical liquid crystal from density and X-ray data. The correlation indicates that the orientation of the molecules In the condensed monolayer of 1-monoelaidin approximates the orientation in the anhydrous three-dimensional crystal.
Introduction
This investigation was undertaken in part to ascertain if either of the above-mentioned theories is applicable to the unsaturated 1-monoglycerides. It was
The effect of unsaturation in monolayers of fatty acids on water has been studied extensi~ely.~-7 The unsaturated CI8 acids occupy much greater areas per 'molecule than does stearic acid a t the same tem(1) The authors extend their appreciation to Swift and Co. for a perature. Hughes4 has suggested that expansion is research grant that made this work possible. caused by the attraction of water for the double bonds. (2) E. K. Rideal and R.K. Sohofield, Proc. R o y . SOC.(London), 110, Sneider, et a1.,6 have found that the area of the cis C I ~ 170 (1926). (3) E. K. Rideal and A. Hughes, ibid., 140, 253 (1933). acids a t negligible pressure increases as the unsatura(4) A. Hughes, J . Chern. Soc., 338 (1933). tion increases. They have suggested that this effect (5) J . Marsden and E. K. Rideal, ibid., 1163 (1938). and the stability of the expanded monolayers are (6) V. L. Sneider, R. T. Holman, and G , 0. Burr, J , P h y s . Colloid caused by a progressively greater curl of the unsymChem., 53, 1016 (1949). (7) E. D. Goddard and A. E. Alexander, Biochem. J., 47, 331 (1950). metrical molecules. ~~
T h e Journal of Physical Chemistry
