Electrical Transport Properties of Metal-Ammonia and Metal-Amine

7 Electrical Transport Properties of Metal-Ammonia and Metal-Amine Solutions

Downloaded by NANYANG TECHNOLOGICAL UNIV on June 7, 2016 | http://pubs.acs.org Publication Date: January 1, 1965 | doi: 10.1021/ba-1965-0050.ch007

DONALD S. BERNS Division of Laboratories and Research, New York State Department of Health and Department of Biochemistry, Albany Medical College, Albany, Ν. Y.

The electrical transport properties of alkali metals dissolved in ammonia and primary amines in many ways resemble the properties of simple electrolytes except that the anionic species is apparently the solvated electron. The electrical conductance, the transference number, the temperature coefficient of con ductance, and the thermoelectric effect all re flect the presence of the solvated electron species. Whenever possible the detailed na ture of the interactions of the solvated electrons with solvent and solute species is interpreted by mass action expressions.

J h e properties of electrolyte solutions have presented an intriguing prob lem ever since they were first explored by Faraday over a century ago. Electrolytes are distinguished from molecular compounds i n solution by the fact that the compounds dissociate into atoms or groups of atoms bearing electrical charges, termed ions. These ions possess properties that permit us, b y means of an external field, literally to reach into the solution and move them about, and i n that way learn something about their interactions. In part, the origins of physical chemistry rest with the early investigation of the conductance of electrolyte solutions. Similarly, the origins of the studies of solvated electrons or metals i n amine solvents lie i n the precise electrical transport experiments begun over 50 years ago by Kraus (24). Kraus* early experiments and sub sequent interpretation made it quite clear that, in solutions of alkali and alkaline earth metals i n ammonia and the amines, one ionic species, the cation, is a metal ion, and the anionic species is some form of solvated electron. More sophisticated experimental techniques have greatly added to our knowledge of solvated electron properties; however, it is 82

Hart; Solvated Electron Advances in Chemistry; American Chemical Society: Washington, DC, 1965.


7.

BERNS

83

Electrical Transport Properties

still worthwhile to glean information from low field conductivity, trans ference number, and Wien effect experiments. F r o m these experiments one may learn something about the interactions of the electron with the solvent and the rest of the solute species. In attempting to clarify most of the electrical transport experiments, a dilemma arises that may be considered inherent in the properties of these solutions. Metal-ammonia and metal-amine solutions in the dilute concentration region exhibit quasi-electrolyte behavior. T h i s is ad vantageous i n applying much of current electrolyte theory to explain the behavior of these systems and, therefore, we apply the current mecha nisms of electrolyte solutions to our treatment of metal-ammonia solu tions. T h e ideal behavior of dilute electrolyte solutions, namely the i n teraction of the ionic species, is generally explained by using the DebyeHuckle theory. T h e deviation from ideal behavior by the properties of a dilute electrolyte solution is generally attributed to the association of ions into pairs or larger clusters. T h e definition of an ion pair or higher cluster, or for that matter the tacit assumption that these species exist, can tend to restrict the interpretation of the transport properties of the solutions. In our discussion we shall knowingly step into this trap at some times because we believe that the proposed complex is realistic and at other times because the author's ideas on this subject are not suf ficiently sophisticated to propose a convincing alternative. T h e question should be kept in mind, however, whether ion pairs and other association complexes exist or whether the observed results can be explained on the basis of interactions without using complexes. B o t h methods should be considered (17). Conductance T h e equivalent conductance vs. concentration curve i n the dilute region for salts (binary electrolytes) is quite similar to that for the alkali metals in ammonia and amine solvents i n the dilute region. F o r example, if we examine the data below 10 ~ Ν for solution of sodium i n N H , potassium i n ethylenediamine, lithium i n methylamine, and at the same time solutions of K N 0 in N H and K N 0 i n H 0 , we may begin to under stand the similar behavior of solutions of the salts and metals. A plot of Λ / Λ ο is particularly revealing for making this comparison (Figure 1). T h e behavior of all of these solutions is somewhat analogous. Certainly the equivalent conductance of each is significantly different (Table I) owing to the differences in anionic transport species and in viscosity of the solvents as a result of temperature and type. T h e A / A behavior for salts and metals in the several solvents is, however, quite similar. A more accurate comparison of limiting conductance may be made by using the Walden product λο η for the cationic species involved. T h i s is not a straightforward calculation and requires some gross assumptions. A t present we can only compare Λ0770 for metallic species from solvent to solvent. E v e n the reported conductance studies of Cafasso and Sundheim (5) for potassium in 1,2-dimethoxyethane are apparently comparable 2

8

3

3

8

2

0

+


SOLVATED ELECTRON

84

to the Uthium in the C H N H system. However, the work is on saturated solutions, and the system is reported to be diamagnetic, arousing i n teresting speculation concerning the conduction process. Further i n vestigation of this type of system is necessary.


3

2

Λ A,

Figure 1. A / A as a function of C . A comparison of the behavior of alkali metals in amine solvents with the behavior of salts in aqueous and nonaqueous systems. KNOs in H 0 and in NH (48); sodium in NHz (15); potassium in NH CH CHJSfH (11); lithium in CHzNH (3). 1 / 2

0

2

3

2

2

2

2

Table I. Metal or Salt Na

Mass Action Constants and Conductivity Data

Solvent N H , (-33°) NH (-37°)· N H (-37°)» C H N H (-78°) N H , (-72°) NH CH CH NH NH CH CH NH NH CH CH NH N H (-40°) N H (-33°) N H (-33°) H 0 (25°)

Ao 1022

3

3

3

Li Li Κ Rb Cs KN0 KC1 NaBr KN0

3

8

8

8

2

2

2

2

2

2

2

2

2

2

2

2

2

8

3

8

2

Ki 7.23 Χ 10" 9.2 Χ 10" 9.6 Χ 10" 5.5 Χ ΙΟ" 1.28 Χ 10" 1.54 Χ 10"< 1.69 Χ 10" 1.44 Χ 10" 3

228 555 139 117 204 331 348 314 145 c

c

c

6

3

4 4

8.7 Χ Ι Ο " 2.9 Χ 10"

K 27.0 18.5 23.0 5.42 4.33 2

Λ»/ο 2.61 0

2.07 2.8 2.14 1.80 3.14

4 3

References (15) (3) (34) (U) (U) (U) (48) (19) (19) (48)

a Transference data b Activity data c Room temperature

T h e general shape of the equivalent conductance vs. concentration plot for metal-ammonia solutions is shown by the behavior of sodium i n N H at — 3 3 ° C . i n Figure 2. T h e conductance behavior of metal-am monia solutions is quite analogous to the behavior of electrolytes i n solvents of low dielectric constant. T h e dilute region equivalent con ductance decreases with increasing concentration, eventually goes 8


7.

BERNS

85

Electrical Transport Properties Λ 1800 1600 1400 1200


1000 800 600 400

ιο"

β

id

5

ιό

4

ιό

3

GRAM ATOMS Na/LITER

ιό

2

NH

ιο"

1

3

Figure 2. Equivalent conductance as a function of concentration for sodium in NHi at —34° from the data of Kraus (26, 30). through a minimum, and begins to increase (22, 29). T h e large, almost exponential increase in conductivity observed i n the concentrated metalammonia solutions is peculiar to these solutions and is caused by the ap pearance of essentially metallic conduction properties. B o t h metalammonia and amine solutions and simple electrolyte solution conductance properties in the dilute region can be analyzed by comparing the experi mental slope to that predicted from the Onsager-Debye-Hiickle equation. A l l the experimental curves for those metal-ammonia and metal-amine systems investigated he between the slopes predicted for 1-1 and 1-2 electrolytes. Such behavior is generally interpreted i n terms of electro static interactions leading to the formation of uncharged, nonconducting species—ion pairs. T h e appearance of the conductance minimum in simple electrolyte systems has been explained by K r a u s and Fuoss (29) in terms of charged ion triplets. T h e competition for removal of charged species by ion pair formation and the appearance of other charged species from ion triplet formation are used to explain the appearance of the conductance minimum. In metal-ammonia solutions the appearance of the conductance minimum has been explained i n terms of competition between the metal ion-electron interaction which removes conducting species and the onset of metallic type behavior. T h e formation of higher aggregates that conduct can also be invoked. Sukhotin (50) and K e n a u sis, Evers, and K r a u s (22) have each contributed information for elec trolyte solutions that seriously questions the ion triplet interpretation of Kraus and Fuoss (29). Evers and co-workers suggest that under normal circumstances in the dilute concentration region the energy of the C o u lombic interaction of anion and cation in forming an ion pair is sufficiently


SOLVATED ELECTRON


86

greater than the energy supplied by solvent collisions with the ion pair; hence, the relative stability of the ion pair. If an ion pair happens to lie i n the field of an ion, it may be dissociated by the thermal impact of a molecule whose energy is less than that required i n the absence of a field. B o t h ions of the pair are acted upon by a force of the same order of magnitude but of opposite direction, and the Coulombic force holding them together is reduced. T h i s effect may be viewed i n somewhat the same way as the Wien effect (52) when an external field is applied to an electrolyte system. Wien has shown that when an external field is applied to a solution of a partially associated electrolyte, the degree of association decreases with an increasing field. T h e dissociation of ion pairs i n the field of ions has been aptly termed the "microscopic Wien effect of the ion fields." T h e frequency of this effect necessarily increases with, and is roughly proportional to, the concentration of ion pairs and free ions, and thus would be properly considered i n the concentration region of the observed conductance minimum i n electrolyte systems. T h e suggestion of a "microscopic Wien effect" in explaining the conductance minimum i n metal-ammonia solutions is a very inviting speculation. T h e effect could indeed be more pronounced i n these systems. Further investigation along this line is certainly worth pursuing since it is an appealing physical explanation which may be useful i n explaining other phenomena i n the concentration region where it is applicable (14). T h e deviations of experimental slopes for dilute solutions of metalammonia and metal-amine conductance plots from the limiting Onsager slope can be used conveniently to calculate association constants for the postulated, nonconducting species responsible for the deviations. Precise association constants of sufficient magnitude may be taken somewhat as confirming the probable existence of the postulated complexes. H e a v y reliance on precise association constants calculated from this type of interpretation can be dangerous as indicated b y the recent work of K a y and D y e (21) in simpler systems. Indeed, interactions not accounted for i n the conventional Debye-Huckle sense are indicated; however, chemical complexes are not to be taken as a unique interpretation. Analyzing the conductance data for sodium in N H solutions led to the original mass action equilibria postulate of K r a u s (27). It was originally proposed that metal i n N H solutions behave as weak electrolytes or ionogens, producing metal ions and solvated electrons in accord with the equation 8

8

XNH

8

+ M = M + + e-(NH )x 8

(1)

T o explain the properties of more concentrated solutions it was suggested that solvated electrons were released to form electrons essentially free in the metallic sense. e - ( N H ) x = (e~) free + X N H 3

3

(2)

More sophisticated extensions of the original Kraus theory are numerous and are dealt with i n other chapters. Another equiUbrium is necessary to


7.

BERNS

87


explain the magnetic properties that in more concentrated solutions i n dicate the presence of diamagnetic species. One can represent the two necessary equations agreed upon by most theories as M+(am) + c-(am)


M (am)

M (am)

1/2 M ( a m )

Ki

(3)

K

(4)

2

2

with no further comment as to the fine structure of the species. W e are concerned merely with the evidence from electrical transport experi ments that such postulated equilibria are indeed feasible. Assuming that electrolyte theory is applicable to metal solutions permits the calculation of the equilibrium constants K\ and K from the conductance data. It is also instructive to evaluate the mass action constants K for salts in N H and i n amines for purposes of comparison (Table I). F o r N a C l (28) in N H , Κ = 1.45 X 1 0 " as compared to 7.23 X 10 ~ for sodium i n N H (15). F o r tetra-zi-butylammonium picrate in methylamine (16) at —78° C , Κι ~ 8 X 10 compared to Ki = 5.8 X 10 " for Uthium in C H N H at - 7 8 ° C . (3). Note that i n both cases the constants for salts and metals are of the same order of magni tude. T h i s is apparent support for applying the same mechanism i n both systems, namely, ion association. T h e appUcabiUty of electrolyte theory to metal solutions should be considered i n Ught of the evidence that elec tron conduction i n these solutions probably does not proceed by what may be called classical ion transport. We may refer to the solvated electron as possessing an excess mobiUty over that of ordinary ion transport. M a n y mechanisms have been invoked with a tunneUng mechanism most favored. F o r our purposes (to be compatible with electrolyte theory) we may hope that, i n the concentration region under consideration, the conduction mechanism remains essentially unchanged with changing concentration; consequently, a mobiUty decrease with increasing con centration wiU be experienced by the electron largely owing to the inter action with metal ions i n accord with the Onsager equation. One can gather support for this proposal from the fact that Onsager's equation may be used i n connection with aqueous solutions of acids, where the protons probably migrate by a " G r o t t h u s " mechanism. Further sup port is found i n the fact that the Walden product A7?o for sodium in N H solution at —34° C , 2.6 is very close to that obtained with solutions of Uthium i n methylamine at — 7 8 ° C . (3). Longo (34) and Evers indicate for the Uthium in N H system, at — 71 ° C , A N a + + e~; AG° = + 1.0 kcal. m o l e " > K + + e " ; AG° =

- 2 . 1 kcal. m o l e "

(5)

1

(6)

1

These values refer to the hypothetical one molal standard state for the ions. T h e difference between these values, 3.1 kcal. m o l e , should rep resent the difference i n free energy of formation of N a and Κ . T h e value given b y Jolly (20) and Coulter (6\ 7) for this difference is + 3.4 kcal. mole"" i n good agreement as is the value of AG° = — 3.14 very accurately determined b y Schug and Friedman (47) for the reverse re action. Russell and Sienko (45) used the cells - 1

+

+

1

N a (s) I N a I ( C H N H ) | N a | (Hg) 3

2

N a (Hg) j N a l ( N H s ) j N a ( N H ) | inert metal 8

to estimate the standard potential and AG° for the N a process. Drastic assumptions were necessary to complete the calculation; however, their value, AG° = 2.4 kcal. m o l e " , agrees well with the other quoted values. A chronapotentiometric method has been used b y Gordon and Sundheim (18) to measure the anionic diffusion for the potassium i n N H solution i n the presence and absence of salt. T h e limiting value of dif fusion coefficient for the electron is i n the region of 14.5 X 10 " cm. /eec. T h i s agrees with values derived from D y e (12). T h e anionic diffusion coefficient is several times larger than that for the neutral metal species. Gordon and Sundheim, however, do not believe that there is necessarily 1

8

δ


2

7.

BERNS


93

any indication of a different diffusion mechanism for anion and neutral species. Similarly the activation energies for diffusion are not greatly different for the electron and neutral combination. T h e activation energies are pot consistent with the proposal of a conduction band i n dilute solutions, and interpretation in terms of tunneling is not necessary. A unique diffusion process is not involved.


Thermoelectric

Effect

Dewald and Lepoutre (8, 9, 33) have investigated the thermoelectric power of metal-ammonia solutions i n great detail. Concentrated solutions exhibit thermoelectric properties i n very good accord with predictions based on a degenerate electron gas model. T h e thermoelectric power of these solutions is of the order of magnitude characteristic of typical metals (1-10 Mvolts/degree). It was also noted that the thermoelectric power increases monotonically as the concentration decreases, indicating that conduction i n these solutions is by means of electrons ( " N " type conduction). In dilute solutions, the thermoelectric power for both potassium and sodium at — 33 ° C . also indicates a monotonie i n crease with decreasing metal concentration. Furthermore, below 0 . 1 M the data for potassium and sodium are indistinguishable within experimental error. T h e behavior is analogous to measurement results on silicon and germanium as well as dilute aqueous work. T h e variation of thermoelectric power with concentration i n the dilute region is almost logarithmic. T h e slope is about twice what would be predicted from the Nernst equation. Dewald and Lepoutre claim that deviations i n the aqueous work and the metal-ammonia work are not entirely accountable in terms of activity and association effects. Thermoelectric effect experiments of Lepoutre and Dewald with mixed salt-metal-ammonia solution do, however, indicate a pronounced ion-electron interaction. A t —78° C . the same positive temperature dependence and fairly large logarithmic dependence on metal concentration are present. A significant difference between the potassium and sodium curves over the entire concentration region is noted. T h i s suggests that the specific nature of the positive ion has an important bearing on the electronic properties of the dilute solutions. One would expect increased ion-electron interactions at lower temperatures owing to less thermal energy; however, a higher dielectric constant at lower temperatures would tend to decrease the interaction. T h e effect of salt added to sodium i n N H $ solutions at —33 ° C . is noteworthy. There is a very large increase in thermoelectric power, which is the direction expected if the salt were decreasing the concentration of electrons by a common ion effect. T h e observed change is about ten times that predicted simply by such a mechanism. It is concluded that the very large salt effect may be found i n changes i n heats of transport of the charged species. T h e suggestion is made that the apparent anomaly of large dilution effects and large salt effects can be explained by assuming that the transport heat of the electrons is negative; that is, when the electrons move through a solution, heat is transported


SOLVATED ELECTRON

94

in the opposite direction. T h i s could arise if the primary mode of elec tron transport were b y a tunneling or jump mechanism; i n making tran sitions from one polarization center to another, a n electron would then leave the dipole polarization energy of the center behind and, with the electron gone, this energy would eventually be transformed into heat. In its " n e w " environment the electron would polarize the solvent at the expense of thermal energy of the surroundings. Polarization energy would thus flow i n a direction opposite to electron flow. It should be noted that this postulate leads to S for the electron i n ammonia in strong disagreement with estimate b y other methods (33). T h e common facts that obviously permeate all aspects of the trans port properties of metal-ammonia and metal-amine systems are:


0

(1) Phenomenologically, i n dilute solutions the metal-ammonia and metal-amine solutions have transport properties quite similar to elec trolytes i n similar solvents, and consequently the types of interactions involved are probably similar. (2) T h e anionic species is responsible for the anomalously large con ductance and thermoelectric effects, and is definitely some form of sol vated electron. (3) T h e further anomaly i n all measurements is that the electron appears to have a n "extra m o b i l i t y " over that normally associated with anionic or cationic species. T h i s behavior resembles that associated with the "solvated p r o t o n " i n aqueous solutions. Acknowledgment We wish to thank D r . E . C . Evers and D r . H . Friedman for reading the manuscript and making valuable suggestions. Literature

Cited

(1) Arnold, E., Patterson, A . J . , "Calculation of Conductivity in Sodium-Liquid Ammonia Solutions" in "Metal-Ammonia Solutions, Physicochemical Proper ties," Colloque Weyl, G . Lepoutre, M . J . Sienko, eds., p. 160, Benjamin Press, New York, 1964. (2) Ibid., p. 285. (3) Berns, D . S., Evers, E. C., Frank, P. W., Jr., J. Am. Chem. Soc. 82, 310 (1960). (4) Berns, D . S., Lepoutre, G . , Bockelman, Ε. Α., Patterson, A . J., J. Chem. Phys. 35, 1820 (1961). (5) Cafasso, F . Α., Sundheim, B . R., J. Chem. Phys. 31, 809 (1959). (6) Coulter, L . V., J. Phys. Chem. 57, 553 (1953). (7) Coulter, L. V . , Maybury, R. H . , J. Am. Chem. Soc. 71, 3394 (1949). (8) Dewald, J. F . , Lepoutre, G., J. Am. Chem. Soc. 76, 3369 (1954). (9) Ibid., 78, 2956 (1956). (10) Dye, J. L . , "Electrochemical Properties of Metal-Ammonia Solutions: E . M . F . and Transference Numbers" in "Metal-Ammonia Solutions, Physicochemical Properties," Colloque Weyl, G . Lepoutre, M. J. Sienko, eds., p. 137, Benjamin Press, New York, 1964. (11) Dye, J. L., Dewald, R. R., J. Phys. Chem. 68, 135 (1964). (12) Dye,J.L.,Sankuer, R. F . , Smith, G . E., J. Am. Chem. Soc. 82, 4797 (1960). (13) Dye,J.L.,Smith, G . E., Sankuer, R. F . , J. Am. Chem. Soc. 82, 4803 (1960). (14) Evers, E. C., University of Pennsylvania, Philadelphia, private communica tion. (15) Evers, E. C., Frank, P. W., Jr., J. Chem. Phys. 30, 61 (1959). (16) Filbert, G., unpublished observations, University of Pennsylvania.



7.

BERNS


95

(17) Friedman, H . L., Ann. Rev. Phys. Chem. 12, 171 (1961). (18) Gordon, R. P., Sundheim, B . R., J. Phys. Chem. 68, 3347 (1964). (19) Hnizda, V . F . , Kraus, C. Α., J. Am. Chem. Soc. 71, 1565 (1949). (20) Jolly, W . L . , Chem. Rev. 50, 351 (1952). (21) Kay, R. L . , Dye, J . L . , Proc. Natl. Acad. Sci. 49, 5 (1963). (22) Kenausis, L . C., Evers, E. C., and Kraus, C . Α., Proc. Natl. Acad. Sci. 48, 121 (1962). (23) Klein, Η. M., P h . D . dissertation, University of Pennsylvania (1957). (24) Kraus, C. Α., J. Am. Chem. Soc. 30, 1323 (1908). (25) Kraus, C. Α., J. Am. Chem. Soc. 36, 864 (1914). (26) Kraus, C. Α., J. Am. Chem. Soc. 43, 749 (1921). (27) Kraus, C. Α., J. Franklin Inst. 212, 537 (1931). (28) Kraus, C . Α., Bray, W . C., J. Am. Chem. Soc. 35, 1315 (1913). (29) Kraus, C. Α., Fuoss, R. M . , J. Am. Chem. Soc. 55, 21 (1933). (30) Kraus, C. Α., Lucasse, W . W., J. Am. Chem. Soc. 43, 2529 (1921). (31) Kraus, C. Α., Lucasse, W . W., J. Am. Chem. Soc. 44, 1941 (1922). (32) Kraus, C. Α., Lucasse, W . W., J. Am. Chem. Soc. 44, 1949 (1922). (33) Lepoutre, G., Dewald, J. F., J. Am. Chem. Soc. 78, 2953 (1956). (34) Longo, F . , P h . D . Thesis, University of Pennsylvania, 1963. (35) Marshall, P. R., J. Chem. Eng. Data 7, 399 (1962). (36) Naiditch, S., "Electrical Conductivities of Sodium-Ammonia Solutions," in "Metal-Ammonia Solutions. Physicochemical Properties," Colloque Weyl, G . Lepoutre, M. J. Sienko, eds., p. 113, Benjamin Press, New York, 1964. (37) Panson, A . J., Evers, E. C., J. Am. Chem. Soc. 82, 4468 (1960). (38) Pleskov, V . Α., Acta Physicochim., URSS 6, 1 (1937). (39) Pleskov, V . Α., Acta Physicochim., URSS 21, 235 (1946). (40) Pleskov, V . Α., J. Phys. Chem. (USSR) 9, 12 (1937). (41) Pleskov, V . Α., J. Phys. Chem. (USSR) 20, 163 (1946). (42) Pleskov, V . Α., Monoszon, A . M., Acta Physicochim. (URSS) 2, 615 (1935). (43) Pleskov, V . Α., Monoszon, A . M . , J. Phys. Chem. (USSR) 6, 1290 (1935). (44) Pleskov, V . Α., Monoszon, A . M., J. Phys. Chem. (USSR) 6, 1286 (1935). (45) Russell, J . B., Sienko, M. J., J. Am. Chem. Soc. 79, 4051 (1957). (46) Schettler, P., Patterson, A . J., J. Am. Chem. Soc. 87, 392 (1965). (47) Schug, K., Friedman, H . , J. Am. Chem. Soc. 80, 45 (1958). (48) Shedlovsky, T . , J. Am. Chem. Soc. 54, 1411 (1932). (49) Smith, G . E., Ph.D. Dissertation, Michigan State University, 1963. (50) Sukhotin, A . M . , Zhur. Fiz. Khim. 33, 72 (1959). (51) Sukhotin, A . M . , Zhur. Fiz. Khim. 34, 63 (1960). (52) Wien, M . , Schiele, J. Physik. Z. 32, 545 (1931). R E C E I V E D April 27, 1965.


Electrical Transport Properties of Metal-Ammonia and Metal-Amine

Recommend Documents