SPECTROPHOTOMETRIC IN V E S T IG A T IO N S O F
ALK A LI METALS IN L IQ U ID A M M O N IA

BY
Robert C .

Douthit

A THESIS

Submitted to the School for Advanced G raduate Studies of M ichigan
State University of A griculture and Applied Science
in p a rtial fu lfillm e n t of the requirements
for the degree of

D O C T O R OF PHILO SPHY

Department of Chemistry

1959

ProQuest Number: 10008633

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ACKNO W LEDG M ENT

The author wishes to express his sincere appreciation to Professor
James L. Dye for his guidance and assistance throughout the course of
this investigation7 to Professor Andrew Timmick for his many helpful
suggestions on instrumentation, and to the Atom ic Energy Commission for
a grant subsidizing this research.

V IT A

Robert C . Douthit
candidate for the degree of
Doctor of Philosophy

O ra l Exam ination:

O ctober 16, 1959

Dissertation:

Spectrophotometric Investigations of A lk a li Metals
in Liquid Ammonia

O u tlin e of G raduate Studies:
M ajor subject:

Physical Chemistry

M inor subjects:

Inorganic Chem istry, Physics

Biographical Items:
Born, December 2 9 , 1932, Camden, N e w Jersey
Undergraduate Studies, C lark U niversity, W orcester,
Massachusetts, 195 0 -5 4
G raduate Studies, M ichigan State U niversity, 1 95 4 -5 9
Experience:

G radu ate Teaching Assistant, M ichigan
State University, 1 95 4 -5 6; Special
G raduate Research Assistant, M ichigan
State U niversity, 1 9 5 7 -5 9 .

Member of American Chem ical S ociety.

Ill

SPECTROPHOTOMETRIC IN V E S T IG A T IO N S OF ALKA LI
AAETALS IN L IQ U ID A M M O N IA

BY
Robert C . Douthit

A N ABSTRACT

Submitted to the School for Advanced G raduate Studies of M ichigan
State University o f A griculture and A pplied Science
for the degree of

D O C TO R OF P H ILO S O P H Y

Department of Chemistry

1959
Approved

/

/

,/

\

j

,

ABSTRACT

The absorption spectra of d ilu te solutions of sodium and potassium in liquid
ammonia were obtained as a function of concentration and tem perature.

The line shapes

for the absorption curves were found to be id e n tic a l, and independent of concentration
for both sodium and potassium solutions.
curve is shifted to lower energies.

W ith an increase in temperature the absorption

These results indicate that the absorption process

in the near infrared involves the e xcitatio n o f an electron in a cavity to a higher energy
state and is the same for d ilu te solutions of both metals.
Sodium-ammonia solutions obeyed Beer's la w , with respect to total m etal, over
the concentration range covered (c a . 3 x 10” 4 to 4 x 10"^ M ) w h ile potassium solutions
(c a . 1 x 10“ ^ fo i x 10" ^ M ) gave a negative deviation from Beer's law .
is explained in the light of current models for these solutions.
io n -p a ir formation in d ilu te solutions for both metals.

This deviation

The model preferred involves

In the more concentrated solutions

exam ined, dimer formation is presumed to occur for potassium, but not for sodium.

This

is in agreement w ith the results of paramagnetic resonance studies for potassium solution
but not for sodium solutions.

The molar absorbancy index of sodium solutions and of

in fin ite ly d ilu te potassium solutions was found to be 4 5 ,0 0 0 ± 2 ,0 0 0 1. mole"^ c m .” ^ .
A general model is presented to explain the magnetic and op tical properties of
solutions of metals in both m ethylamine and ammonia.

The properties of these solutions

cannot be com pletely explained in terms of the existing models.

The new model includes

competing e q u ilib ria between diatom ic m olecules, solvated electrons, solvated metal

V

ions, and solvated dimers.

The direction in which these e q u ilib ria shift depends upon

the physical properties of both the metal and the solvent.

The absorption bands in the

v is ib le region are a ll attributed to diatom ic molecules w h ile the absorption bands in the
near infrared are attributed to solvated electrons.
to absorb in these regions.

The solvated dimer is presumed not

Q u a lita tiv e experiments on the absorption spectrum of

cesium -m ethylam ine solutions showed the peak to be in the region p redicted.

H ow ever,

the existence of two additional peaks for potassJum-methylamine solutions cannot be
adequately explained by this model.

VI

TABLE O F C O N T E N T S
Page
I.
II.

IN T R O D U C T IO N ...........................................................................................

1

H IS T O R IC A L

4

............................................

III.

T H E O R E T IC A L ................................ .............................................................

IV

E X P E R IM E N T A L

V.
V I.
V II.

V III.

. ............

. ..............................

19
29

RESULTS.............................................

53

D IS C U S S IO N ....

82

...................................................................

PRESENTA TIO N OF N E W M O D E L FOR M E T A L -A M IN E
S O L U T IO N S
.....................................................
S U M M A R Y ........................................

V II

90
100

LIST OF TABLES

Table
I.

II.
III.
IV .
V.
V I.
V II.

V III.
IX.

X.

X I.

Page
A B SO R PTIO N SPECTRA OF DILUTE S O L U T IO N S O F METALS IN
A M IN E S O L V E N T S .
......................... . ................ . . . . . . . . ...................

5

VALUE CHART FOR TH ERM O REG ULATO R............................................
A B SO RPTIO N SPECTRUM O F A N H Y D R O U S L IQ U ID A M M O N I A . .

38
49

THERMISTOR C A L IB R A T IO N .............................................................

51

D A TA FR O M A TY P IC A L S O D IU M -A M M O N IA ABSO RPTIO N CURVE

56

EFFECT O F C O N C E N T R A T IO N O N S O D IU M SPECTRUM AT - 6 5 ° C .

60

M O LAR ABSO RBANCY IN D E X FOR S O D IU M -A M M O N IA
............................................................
S O L U T IO N AT - 6 5 ° C .
EFFECT O F TEMPERATURE O N A SO D IU M - A M M O N IA SPECTRUM.
M O LAR ABSORBANCY IN D E X FOR P O T A S S IU M -A M M O N IA
S O L U T IO N S AT - 6 5 ° C ........................................................................

62
67

71

EFFECT O F C O N C E N T R A T IO N O N THE PO TA SSIU M SPECTRUM
AT - 6 5 ° C ..............................................................................
EFFECT O F TEMPERATURE O N THE P O T A S S IU M -A M M O N IA
SPECTRUM......................................................................

X II.

A B S O R P TIO N PEAKS FOR AAETALS IN M E T H Y L A M IN E A N D FOR
G A S PHASE MOLECULES
..................................................

V I II

91

LIST O F FIGURES
Figure

Page

1.

Isometric V ie w of Assembled Ves s el . . . . . . . . . . . . . . . . . . . . . . . . . . . .

31

2.

Cross-section of Figure 1 Through A - A , . . . . . . . . . . . . . . . . . . . . . . . . . .

32

3.

Cooling Arrangement

35

4.

Thermoregulator C irc u it. . . . . . . . . ___. . . . . . . . . . . . . . . . . . . . . . . . . .

37

5.

Solvent Purification System . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

42

6.

......................................

.....................

Left: Solution M a k e -u p C e l l .
Right: M e ta l Degassing Tube

7.

............ ..

Top: C alib ratio n of Flame Photometer

44

...................................

Bottom: Absorption Spectrum of Anhydrous Ammonia . . . . . . . . . . . . .

52

8.

Thermistor C a lib ra tio n .

9.

Typical Sodium-Ammonia Absorption Curve (W avenum ber).. . . . . . .

58

10.

Typical Sodium-Ammonia Absorption Curve (W av e le n g th )....................

59

11.

Effect of Concentration on Sodium Spectrum at - 6 5 ° C . . . . . . . . . . . .

61

12.

Absorbancy versus Concentration for Sodium-Ammonia Solutions
at - 6 5 ° c . . 7 7 7 7 7 7 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

63

Temperature Dependence of the Absorption Curve for SodiumAmmonia Sol ut i ons. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

68

Absorbancyversus Concentration for Potassium-Ammonia Solutions
at - 6 5 ° C . . 7T777 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

72

M o la r Absorbancy Index versus Concentration for Sodium and
Potassium Solutions at - 6 5 ° C . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

73

Position of Peak Maximum for Solutions of Sodium and Potassium
versus Concentration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

74

13.

14.

15.

16.

.......................

50

IX

LIST O F FIGURES (c o n t.)

Figure
17.

Page
Position of Peak Maximum for Solutions of Sodium and Potassium
versus Square Root of A bsorbance.
................

75

18.

Effect of Concentration on the Potassium Spectrum at -6 5 ° C

78

19.

Temperature Dependence of the Absorption Curve for PotassiumAmmonia S o lu tio n s .
................

80

Comparison of the Shapes of Sodium and Potassium Spectra in
Ammonia at -6 5 ° C
.......

81

Absorption Curves for M etals in Me t h y l a mi n e . . . . . . . . . . . . . . . . . . .

92

20.

21.

X

...........

I. IN T R O D U C T IO N

Some of the com plexity in the study of m etal-am m onia solutions is reflected in
a quotation of C . A . Kraus regarding the sodium-ammonia system alone (1 ).
"The prosecution of this investigation has made great demands on both
time and patien ce and progress at times has seemed p a in fu lly slow.
N o t only was it necessary to develop an extensive technique in order
to m anipulate successfully an extrem ely reactive solute and a highly
v o la tile solvent, but the number of experiments has to be m ultiplied
many times in order to exclude certain disturbing factors which have
their origin in a slow reaction taking place between the two com­
ponents . "
The solutions of a lk a li metals in liquid ammonia are characterized by certain
unusual properties w hich are generally sim ilar for this group of metals.

This unique

system of a meitbl disolved in a non-m etal has been of interest to chemists for nearly
one hundred years (2 ).

Although there is s till much uncertainty regarding the exact

structure of these solutions many of their physical properties have been determ ined.
The literature in this fie ld is extensive and the reader is referred to several review
articles for a more complete survey of the fie ld ( 3 , 4 , 5 ) .
A few of the physical properties of m etal-am m onia solutions w hich must be
exp lained by any theo retical model are listed below:
1.

D ilu te solutions are blue w h ile more concentrated solutions e x h ib it a
copper-bronze color (4 ).

2.

The solutions are less dense than the solvent (6).

2.

3.

C onductivities appear to be e le c tro ly tic in nature in d ilu te solutions
and ele c tro n ic in very concentrated solutions (7 ).

4.

The negative species carries approxim ately 90 per cent of the current in
d ilu te solutions and approaches 100 per cent in more concentrated
solutions ( 8 ,9 ) .

5.

A ll d ilu te solutions in liquid ammonia e x h ib it an increasing absorption
throughout the visible spectrum and reach a maximum in the near Infrared
at about 1 .5 microns (1 0 ,1 1 ).

6.

Paramagnetic resonance measurements show that the molar susceptibility in
d ilu te solutions approaches one Bohr magneton per metal atom , w h ile in
more concentrated solutions the molar susceptibility indicates the solutions
are only slightly param agnetic ( 12),

Theoretical speculations concerning the structure of m etal-am m onia solutions stem
from the classical investigations o f C . A . Kraus which commenced over fifty years ago.
His studies included vapor pressures, densities, apparent molecular w eights, conductivities,
phase studies and re la tiv e transport numbers of the conducting species. As a result of these
studies Kraus (13) proposed a m odel, w hich consisted of an equilibrium between metal
atoms,

metal ions, and solvated and unsolvated electrons; which q u a lita tiv e ly accounted

for the properties observed up to that tim e.

N o significant theo retical advances were made

for about th irty years until the magnetic properties were investigated by Huster, (14) and
by Freed and Sugarman (1 5 ), using static magnetic fie ld s , and later by Hutchison (12) using

3.

the electron param agnetic resonance technique.

These magnetic properties along with

conductivities form the basis of the currently popular models for d ilu te solutions.
The currently popular models are the "electron c a v ity " model first postulated
by O g g , (16) and the "cluster model" proposed by Becker, et a l . , (1 7 ).

Both of these

models exp lain some of the physical data ad eq u ate ly , but neither explains a ll of the
facts.

O ne of the important physical properties which promises to give structural

inform ation is the absorption spectrum.
There have been many q u a lita tiv e investigations of the absorption spectra of
lith iu m , sodium and potassium in amine solvents and liquid ammonia.

The only

q u a n tita tiv e work reported in m etal-am m onia solutions involves measurements of
absorption near the ta il-e n d of the absorption band (1 8 ,1 9 ).

In other solvents the

only q u a n tita tiv e work reported deals w ith solutions of lithium in m ethylam ine (2 0 ) .
The purpose of this research is to extend the q u a n tita tiv e measurements to the
maximum in liquid ammonia (1 .5 microns) and to compare the spectra of sodium and
potassium w ith regard to line shapes and molar absorbancy indices.

P articular attention has

been given to the purity of the solutions used, to their s ta b ility and to the variation of
absorption w ith tem perature.

H IS TO R IC A L

The supposition that the absorption spectra of the a lk a li and a lk a lin e -e a rth metals
in liquid ammonia are id e n tic a l, has often been used to support the various theories con­
cerning these solutions, p a rticu la rly w ith respect to the identity of the species present.
Examination of the a v a ila b le spectroscopic data for d ilu te solutions of the a lk a li
and' a lk a lin e -e a rth metals in ammonia, amines and mixed solvents often shows the presence
of a characteristic band in the infrared at approxim ately 6800 cm.
bands are observed in the visible (1 2 ,5 0 0

^ w h ile in some systems

to 1 4 ,5 0 0 c m ."^ ) w hich vary from metal to m etal.

The existing d a ta , together w ith the results obtained in the present study, are summarized
in Table I, and w ill be discussed in more d e ta il.
Gibson and Argo (1 0 ,2 1 ) made the first extensive studies of the absorption spectra
o f the a lk a li metals in liquid ammonia approxim ately forty years ago.

They made visual

estimations of the extinction using polarizers, thus lim iting their work to the visible region
and enabling them to measure only the ta il end of the absorption curves for the m e ta lammonia solutions.
the absorption c e ll.

They prepared their samples by placing a piece of freshly cut metal into
Due to decomposition they w ere forced to extrapolate their measure­

ments back to a common tim e . They measured approximate concentrations for only two of
their solutions.

Their results, plotted as extinction versus w avelen g th , and giving p a ra lle l

lines for d iffe re n t m etals, led them to conclude that the coloring species is the same for
lith iu m , sodium, potassium, and magnesium in liquid ammonia w ith slight deviations for

5.

TABLE I
A B S O R P TIO N SPECTRA O F DILUTE S O L U T IO N S O F
METALS IN A M IN E SO LVENTS

Tem perature,
Solvent

nh

3

^

M e ta l

Range Investigated, Absorption M a x .

°C

c m .” ^

cm.

^

Li

- 6 5 to 25

1 3 ,3 3 3 - 2 5 ,0 0 0

Li

—

5 , 0 0 0 - 2 5 ,0 0 0

5 ,5 5 0

22

Li

-7 0

5 , 0 0 0 - 2 5 ,0 0 0

6 ,7 0 0

11

Li

-2 5 3

4 , 0 0 0 - 2 5 ,0 0 0

—

8 , 0 0 0 - 1 7 ,0 0 0

Reference
10

34

Li

- 7 0 to 25

1 0 , 0 0 0 - 2 5 ,0 0 0

—

18

Na

- 6 5 to 25

1 3 ,3 3 3 - 2 5 ,0 0 0

—

10,21

Na

—

5 ,0 0 0 - 2 5 ,0 0 0

5 ,5 5 0

22

Na

—

4 , 0 0 0 - 1 0 ,0 0 0

6 ,6 6 7

25

Na

-6 0

5 , 0 0 0 - 2 5 ,0 0 0

6 ,8 0 0

11

Na

-2 5 3

4 , 0 0 0 - 2 5 ,0 0 0

Na

- 7 0 to -3 5

4 ,0 0 0 - 2 0 ,0 0 0

Na

-7 5

1 4 ,0 0 0 - 2 5 ,0 0 0

—

19

Na

- 7 5 to 25

1 0 , 0 0 0 - 2 5 ,0 0 0

—

18

K

- 6 5 to 25

1 3 ,3 3 3 - 2 5 ,0 0 0

K

-7 0

5 , 0 0 0 - 2 5 ,0 0 0

K

-7 5

1 4 , 0 0 0 - 2 5 ,0 0 0

K

- 7 0 to -3 5

4 ,0 0 0 - 2 0 ,0 0 0

K

-7 5 to 25

1 0 ,0 0 0 - 2 5 ,0 0 0

8 , 0 0 0 - 1 6 ,9 5 0
6 ,7 8 0

—
6 ,6 6 7
—

6 ,7 8 0
—

34
*

10
11
19
*
18

6.

TABLE I (continued)

Tem perature,
Solvent
NH3

CH3:NH2

Range Investigated,
cm” ^

Absorption M a x .
1^eferenc

Aetal

°C

Mg

- 6 5 to 25

1 3 ,3 3 3 - 2 5 ,0 0 0

--------

10,21

Mg

25

1 0 ,0 0 0 - 2 5 ,0 0 0

--------

18

Ca

-6 5 to 25

1 3 ,3 3 3 - 2 5 ,0 0 0

--------

10

Ca

25

1 0 , 0 0 0 - 2 5 ,0 0 0

--------

18

Ca

-7 0

4 , 0 0 0 - 2 5 ,0 0 0

Ba

25

4 , 0 0 0 - 1 0 ,0 0 0

Li

-6 0

5 , 0 0 0 - 2 5 ,0 0 0
6 , 2 5 0 - 2 5 ,0 0 0

Li

cm ."^

7 ,8 0 0

18

—

7 ,4 0 0
6 ,6 5 0

11

- 1 4 ,0 0 0

11
28

Li

-7 5

1 3 ,3 3 3 - 2 5 ,0 0 0

1 4 ,5 0 0

20

Na

-6 0

5 ,0 0 0 -2 5 ,0 0 0

1 5 ,6 0 0

11

Na

- 7 0 to 25

1 3 ,3 3 3 - 2 5 ,0 0 0

1 5 ,4 0 0

10,21

6 , 2 5 0 - 2 5 ,0 0 0

1 4 ,5 0 0

28

Na

K

-6 0

5 ,0 0 0 - 2 5 ,0 0 0

1 2 , 2 0 0 - 1 5 ,0 0 0

11

K

- 7 0 to 25

1 3 ,3 3 3 - 2 5 ,0 0 0

1 2 , 0 0 0 - 1 5 ,4 0 0

18

Cs

-6 0

4 , 0 0 0 - 2 5 ,0 0 0

1 0 ,0 0 0

Ca

-6 0

5 , 0 0 0 - 2 5 ,0 0 0

7 ,2 3 0

*

11

7.

TABLE 1 (continued)

Solvent

M e ta l

Tem perature/
°C

Range Investigated/ Absorption M a x .
cm. ”
cm” ^

Referei

Li

-4 0

5 / 0 0 0 - 2 5 /0 0 0

6 ,7 0 0

11

Li

—

4 2 5 0 - 2 5 ,0 0 0

1 6 ,7 0 0

28

Na

—

6 ,2 5 0 - 2 5 ,0 0 0

1 4 ,7 0 0

28

-6 0

5 , 0 0 0 - 2 5 ,0 0 0

1 5 ,5 0 0

11

Na

25

1 0 ,0 0 0 - 2 5 ,0 0 0

1 4 ,9 0 0

26

K

25

1 0 , 0 0 0 - 2 5 ,0 0 0

1 4 ,9 0 0

26

Na

25

1 0 ,0 0 0 - 2 5 ,0 0 0

1 5 ,0 0 0

26

K

25

1 0 ,0 0 0 - 2 5 ,0 0 0

1 5 ,0 0 0

26

1 0 -2 0 % N H 3

Li

25

6 , 2 5 0 - 2 5 ,0 0 0

30% N H 3

Li

25

6 , 2 5 0 - 2 5 ,0 0 0

7 ,7 0 0

28

10% n h 3

Na

25

6 , 2 5 0 - 2 5 ,0 0 0

1 4 ,3 0 0

28

2 0 -3 0 % N H g

Na

25

6 ,2 5 0 - 2 5 ,0 0 0

7 ,6 0 0 - 8 ,0 0 0
1 3 ,9 0 0

28

4 0 -1 0 0 % N H 3

Na

25

6 , 2 5 0 - 2 5 ,0 0 0

7 , 7 0 0 - 8 ,0 0 0

28

50% N H 3

Na

-5 0

5 , 0 0 0 - 2 5 ,0 0 0

7 ,5 0 0 - 1 5 ,0 0 0

11

5 0% N H 3

K

-7 0

5 , 0 0 0 - 2 5 ,0 0 0

C2 H 5 N H 2

K
(• C H 2 N H 2 )2

(N H 2 - C H C H 3 C H 2 N H 2)

N H 3- C H 3 N H 2
7 , 7 0 0 - 1 3 ,5 0 0

7 ,5 0 0

28

11

TABLE I (continued)

Tem perature,
Solvent

M e ta l

°C

Range Investigated, Absorption M a x .
cnrT^

cm” ^

Reference

NH3-C2H5NH2
1 0 -3 0 % N H 3

Li

25

6 ,2 5 0 - 2 5 ,0 0 0

50% N H 3

Li

25

6 , 2 5 0 - 2 5 ,0 0 0

8 ,0 0 0

28

1 0 -2 0 % N H 3

Na

25

6 , 2 5 0 - 2 5 ,0 0 0

1 4 ,3 0 0

28

3 0 -4 0 % N H 3

Na

25

6 ,2 5 0 - 2 5 ,0 0 0

50% N H 3

Na

25

6 ,2 5 0 - 2 5 ,0 0 0

* This Research

7 ,7 0 0 - 1 4 ,1 0 0

7 ,7 0 0 - ^ 1 3 7 0
8 ,0 0 0

28

28
28

Since the color is obtained regardless of the metal dissolved in liquid ammonia
they exam ined three possible sources of color:
1.

U n -io n ize d molecules or atoms of the m etal.

2.

Unsolvated electrons.

3.

Solvated electrons.

From the electron theory of m e ta llic dispersions they calculated the number and m obility
of unsolvated electrons to be expected in solution based upon the absorption spectrum.

The

calculations gave erroneous results so they elim inated the unsolvated electron as the source
of c o lo r.

Since they found the absorption spectra of a ll a lk a li metals in liquid ammonia to

be the same over the region examined they elim inated any u n -io n ize d metal atoms as being
the main source o f color and concluded that the color was due to solvated electrons.
Absorption measurements of metals dissolved in m ethylamine resulted in a series of
curves which had absorption maxima at approxim ately 650 m illim icrons which they concluded
to be id e n tica l for the various metals.
It was not until 1939 that the investigation of the absorption spectra of a lk a li metals
in liquid ammonia was resumed by V o g t (2 2 ).

He measured the spectra of sodium and

lith iu m , presumably at room tem perature, and found absorption maxima at 1 .8 microns.
experim ental details were g iv e n .

No

Vogt also mentioned that the visible spectra were the same

for the two metals measured and concluded that the electron is com pletely dissociated from
the m e ta l.

He interprets the absorption maximum at 1 .8 microns ( 0 .7 e . v . ) as the minimum

energy for the bonding of the e le c tro n , probably w ith the solvent.

10.

Ogg and his students (1 8 ,2 3 ,2 4 ) have investigated the absorption spectra of
solutions of L i, N a , K , M g , Ca and Ba in liquid ammonia.

They used a modified

Beckman DU Spectrophotometer w ith a specially prepared c e ll holder to perm it cooling
w ith a stream o f nitrogen, w hich had been precooled w ith liquid a ir. Since the path length
was fix ed at 1 cm ., they were able to exam ine only very d ilu te solutions. W ithin the
te n -fo ld concentration range that could be studied, the solutions seemed to obey Beer's
law .

Withito their considerable experim ental error, at a given w avelength and tem perature,

the molar absorbancy indices for lith iu m , sodium and potassium w ere the same w h ile
magnesium, calcium and barium were likew ise id e n tic a l, but tw ice the values found for
the a lk a li metals.
The spectra showed continuous absorption w ith increasing w avelength and a broad
maximum was found between 7 00 and 900 m illim icrons. Since it has been shown that the
absorption maximum for these solutions is at 1 ,5 0 0 m illim icrons, the maximum found by
Ogg probably can be attributed to a decrease in sensitivity of the phototube used.

Ogg

also noted that the temperature dependence of the light absorption of these solutions was
g reater than one would expect by solvent contraction alo ne.

H e found that the

absorption c o e ffic ie n t for "blue lig h t" was but slightly affected by tem perature.

Ogg

observed that by shifting the curves at various temperatures - 2 . 5 x 10"^ e . v . degT^,
a ll curves could be made to co in cid e .

He interpreted his results in terms of two colored

constituents w hich he postulated to be paired and unpaired electrons in cavities w ith in the
liquid

am m onia. He assumed that the equilibrium between the paired and unpaired electrons

was tem perature dependent. H e considered the unpaired electrons to absorb at longer
w avelengths than the paired electrons.

11.

Eding (18) also measured the temperature dependence of the amide absorption
in liquid ammonia and found that the amide peaks at 3 50 m illim icrons could a ll be made
to coincide by shifting a ll curves - 2 . 5 x 10 ^ e . v . deg. ^.

This would seem to indicate

that the absorption process is sim ilar for both amide and metal in ammonia and that this
large temperature dependence is not characteristic of just electrons in ammonia.
J o lly (25) measured the absorption spectra of sodium-ammonia solutions in a 1 m illim eter
c e ll.

The measurements from 3 5 0 -1 1 0 0 m illim icrons were made w ith a Cary Recording

Spectrophotom eter and a Beckman D U Spectrophotometer.

The range from 1 . 0 - 2 .5 microns

was covered w ith a Perkin-Elm er Infrared Spectrophotometer.

His solutions were made by

simply dropping a piece of sodium into a c e ll p re fille d w ith liquid ammonia.

The

decomposition rate was about 30 per cent per!hour.
His data in the visible region were the same as the data of Gibson and Argo (1 0 ,2 1 ),
but he found an absorption maximum at 1 .5 microns, 0 .3 microns lower than that reported
by V o gt (2 2 ).

Jo lly interprets this maximum at 1 .5 microns (0 .8 3 e . v . ) as the exc ita tio n

energy of a solvated (unpaired) electron to a conduction band where it would be a free
or unsolvated e le c tro n .
Blade and Hodgins (11) exam ined the absorption spectra of Li, C a , K , and N a in
liquid am m onia, m ethylam lne, and ethylam ine as w e ll as several binary mixtures of these
solvents.

A ll solvents and metals used were of high p u rity.

The spectra were determ ined

in two stages; from 300 - 1000 m illim icrons on a Beckman D U Spectrophotometer and from
1 . 0 - 2 . 5 microns on a Beckman IR—2 Infrared Spectrophotometer.
was u tiliz e d for temperature measurements as low as - 7 5 ° C .

A special cell holder

The absorption cells used

w ere made by the American Instrument Company and had light paths of 1 0, 1, and 0.1
m illim eters.

They found that a ll metals in liquid ammonia had identical spectra w ith absorption
maxima at 1 .5 microns.

H ow ever, they made no attempt to measure the concentration

of these solutions so one can conclude that the spectra are id e n tica l only w ith regard to
the shape and position of the absorption maximum.

In methyfamine d iffe re n t bands were

observed for these metals which were dependent upon the m e ta l.

Lithium and calcium

showed an absorption maximum at 1 .3 microns, sodium a maximum a t 540 m illim icrons and
potassium two bands around 830 m illim icrons.
Blade and Hodgins found that when sodium is dissolved in a 50 per cent m ixture
of ammonia and m ethylam ine, two peaks appear, one at 650 m illim icrons and another
at 1 .5 microns.

This system is very sensitive to temperature changes.

As the temperature

is lowered the intensity of the 650 m illimicrons peak increases w h ile the intensity of the
1 .5 micron band decreases.

This temperature dependence seems to be reversible and

might be explained w ith the "electron c av ity " model as being the temperature dependent
equilibrium between paired and unpaired electrons, or by the Becker, Linquist, A lder
model (17) as the temperature dependence of the equilibrium between monomers and dimers.
Blade and Hodgins interpreted the absorption maxima to be due to trapped
electrons since neither the solvents nor the metals show absorption in these regions.
further elaborated on the nature of these traps.

They

They considered the 1 .5 micron peak to

be an amine type trap arising from the symmetry of protons in the ammonia m olecule.

They

considered the 650 m illim icron peak to be an a lip h a tic trap arising from nonsymmetry of
protons in m ethylam ine.

In this type of tra p , the protons forming the trap for the electrons

in the solution are the protons of the methyl group.

13.

In a re -e xa m in atio n of the theory of Blade and Hodgins on "a lip h a tic " and
"am ine" type traps, and in an attem pt to correlate the existing spectral results w ith magnetic
data (1 2 ), Fowles, M cG regor, and Symons,(26) have proposed that rather than being caused
by configurations of solvent molecules about the h o le , the absorption bands in the visible
( 650 m illim icrons) and near infrared (1 .5 microns) are due to e2 species and e j

species

residing in cavities sim ilar to Ogg°s model. Fowles, et a I . , suggest that the infrared band
( 1 .5 microns) arises from e]
centers.

centers w h ile the visib le band (650 m illim icrons) arises from e2

They made absorption measurements on sodium and potassium in ethylenediam ine and

propylenediam ine and found only the visible band at 670 m illim icrons.

M a g n e tic studies of

these same solutions indicated that the solutions were not param agnetic, although the
to tal metal concentration in these solutions (c a . 10

—3

M ) would give an unpaired electron

concentration many hundred of times greater than the minimum w hich could be d e tected .
When both bands are observed or only the infrared band, the solutions are found to be
param agnetic.
FowleS, et a I . . postulated that the two absorption bands found in m etal-am ine
solutions do not correspond to the e je c tio n of an e le c tro n , paired or unpaired, from its
c a v ity into a conduction band, but rather to its excitatio n to a higher energy orbital s till
w ith in the same c a v ity .

This is sim ilar to the postulate of Lindschitz (2 7 ), who studied

the effects of irradiation on lithium in glasses containing e th e r, isopentane, methylam ine
and trim eth ylam ine.

14.

Because small changes in tem perature, concentration and the nature of the metal
and solvent appear to a ffe c t the location of theiband maximum, single and paired
electrons are postulated to have sim ilar energies.

It was observed that lowering the

tem perature, or increasing the metal concentration favored the absorption in the v is ib le .
Therefore, this absorption was interpreted as being due to paired electrons.

According

to this theory, solvents w ith re la tiv e ly high d ie le c tric constant (N H g =22) appear to
favor e j form ation irrespective of the m etal, but in amine solvents e"j formation appears only
when the metal ion has a high surface-charge density.
m ethylamine appear to form e l centers.

Thus lithium and calcium in

These trends suggested, that in amine solvents,

io n -p a ir formation could take place so that the metal ion is close to the solvated electron
and has n oticeable influence on the type of cav ity form ed.

This could probably be

attributed to the combined e ffe c t of the higher charge-density per cavity and the lower
d ie le c tric constant of the amine solvents which makes io n -p a ir formation significant
despite the low metal concentration.
Hohlstein and W annagat (28) have recently made absorption studies in m etal-am in e
systems.

In connection w ith another problem they synthesized w hat they considered to be

am m onia-free m ethylam ine.

N o tin g that sodium was only slightly soluble in m ethylamine

they postulated that the blue coloration achieved In amine solvents was due solely to traces
o f ammonia in the amine solvents.

Their work led them to measure the absorption spectra of both

sodium and lithium in sealed ampoules at room tem perature, in various mixtures of ammonia
and m ethylam ine.

They concluded that the ammonia complex is responsible for the color

o f the solutions and they discarded the "electron c a v ity " p rinciple of solvation.

In the

case o f pure m ethylam ine, they proposed the follow ing disproportionation to take p la c e ,
c ata ly ze d by the dissolved m etal.

15.

2 CH3 N H 2

=

(C H 3 ) 2 N H

+

NH3

Shatz (20 ) on the basis of thermodynamic calculations showed that the above reaction
in not fe a sib le .
Meranda (19) measured the absorption spectra ot solutions of sodium, potassium
and calcium in liquid ammonia.
microns.

Measurements were made in the range 4 0 0 -7 0 0 m illi­

Meranda°s solute and solvent appeared to be quite pure. The absorption cell

he used was specially constructed for the purpose and had a light path of about 0 .3
m illim e te r.

His c e ll had no provision for stirring the solution, and his method for

a tta in in g the tem perature o f measurement, - 7 5 ° C , was one of refluxing and condensation,
and could have given rise to serious concentration gradients w ith in the c e ll at the time
o f measurement, w hich would lead to erroneous results.

His spectrophotometrlc data were

obtained in the form of photographs of the oscilloscope tracings generated by the output
o f the phototube. Concentrations were determined by standard acid-base titra tio n
techniques.

H e assumed the density of his solutions at - 7 5 ° C to be equal to the density

o f liquid ammonia a t - 3 3 . 5 ° C , w hich introduces at least a 5 per cent error.

M eranda

presents a tab le on page 119 of his thesis giving his experim ental d a ta , but nowhere can
one find the concentrations of the various solutions measured except for a single sodium
solution of 0 .0 3 2 8 5

m olar.

His "average*1 extinction c o e ffic ie n t could be the

average of three separate solutions of d iffe re n t concentrations or could be the average of
three d iffe re n t oscilloscope tracings of one solution.

His data are not in agreerrent

w ith the g e n era lly accepted statement that metals of equal concentration in liquid ammonia
show id e n tic a l spectra.

H e finds great sim ilarity in the absorption curves of sodium

16.

and

pptassium,

but the molar absorbancy indices for these solutions d iffe r by

almost 5 0 per cent w ith potassium higher than sodium, in contradiction w ith the
findings o f this research.

H e also found that the shape of the absorption curve for

calcium was not the same as that of sodium or potassium as was e a rlie r concluded
by Gibson and Argo (1 0 ).
M eranda concludes that his data are not sufficient to determ ine whether
d iffe re n t species are the cause of the color since he obtained no absorption maxima.
H e mentions that his data may be accounted for by the presence in each solution of
the metal ammonium.

The small dissim ilarities could then be ascribed to differences

in the interaction of the solvent w ith c o llo id a l particles of the m e ta l, the metal ions
or to an unknown cause.

It appears that his experim ental work is open to question.

Shatz (2 0 ), using a Bausch and Lomb grating monochromator and a s p ecially
constructed c e ll of V ariable light path measured the absorption of lithium in m e th y lam ine as a function of concentration (c a . 1 x 10 ^ to 5 x 107$) at - 7 8 ° C .

His

c e ll was designed so that the solution could be mixed before and after each optical
measurement thus e lim in atin g concentration gradients. His apparatus seems to be
ideal for such measurements except for the fa c t that his measurements were restricted to
the visib le region using his sp ecially constructed optical equipm ent.

A m odification

of his apparatus to f it the c e ll compartment of a Cary Recording Spectrophotometer,
for exam p le, would c ertain ly be helpful in a complete investigation of m etal-am ine
systems.

17.

Shatz measured the metal concentration by titra tio n of the residual lithium after
rem oval of the m ethylam ine.

He found that solutions of lithium in m ethylamine did not

obey Beer“s law over the concentration range exam ined.

Extremely d ilu te solutions appeared

to obey BeerQs la w , but the more concentrated solutions tended to absorb more than they
w ould if they obeyed Beer“s law .
et

He treats his results in accord w ith the theory of Becker,

a l . , and calculates extin ctio n coefficients for solvated electrons, unsolvated electrons,

and free electrons.
O thers who have attempted to account for the absorption spectra of m etal-am m onia
solutions are Jaffe (2 9 ), Gordon and Broude (3 0 ), Davydov (3 1 ), Deigen (3 2 ), D anilova
(3 3 ), Bosch (3 4 ), Kruger (3 5 ), and Jortner (36).
Jaffe (29) measured the spectral v ariatio n of light reflected from the m e ta llic solutions as
a function o f the angle of in cid en ce.

He interpreted his results as consistent w ith his

th eo ry, which requires free electrons dispersed in the solutions as w e ll as in the compounds
L i ( N H 3) 4

and N a ( N H 3 )x »

Gordon and Broude (30) found that the blue streaks observed during the electrolysis
o f sodium amide in liquid ammonia were n e g ative ly charged, that is they moved toward
the anode and were deflected in the proper direction by a magnetic fie ld .

They

interpreted their results as in agreement w ith the polaron theory, e v id e n tly , the e quivalent
in Russian literatu re of the "solvated e le c tro n ".

Davydov (31) gives a quantum -m echanical

treatm ent of the polaron which he interprets as being consistent w ith the absorption spectra,
and w ith the observation of Gordon and Broude.

Deigen (32) employs the polaron theory

to account for the properties of the m e ta llic solutions, and also introduces the additional
concept of single and double color centers (perhaps eq uivalent to the monomer and dimer of
the Becker, Lindquist, and A lder m odel), and treats the solutions as quasi-crystaline
d ie le c tric s .

18.

D an ilova (33) measured the absorption spectra of N a , K , Rb, and Cs absorbed on
crystalline ammonia at 183°C„

H e interpreted his results in terms of the polaron theory

described above.
By condensation methods Bosch (34) obtained thin layers of ammonia containing
small amounts of sodium, lith iu m , or potassium at 9 0 ° K . and 2 0 ° K .
of the metal added a ll films showed light absorption at 1. 15 microns.

A t 2 0 ° K . regardless
Also bands characteristic

of the metal appeared w hich were attributed to neutral atoms or colloids at the surface
boundary.

From this an electro n ic state of solid ammonia is deduced whose electrons are

supplied by the evaporated m etal.
Kruger (35) measured the absorption spectrum of sodium-ammonia solution at
- 1 8 3 ° C and found a maximum at 589 m illim icrons.

He attributes this peak to F-centers.

The calculations of Jortner (36) w ill be discussed in d e ta il la te r.

19.

II.

THEORETICAL

A number of models have been proposed fro accounfr for frhe properties of mefralargmonia solutionsc N one of these offers a complete explanation of the experim ental fa c ts „
It is g en erally agreed that the concentrated solutions are m etal-1 ik e , but the nature of the
d ilu te solutions is not c le a rly understoodinto two groups; A .

The metal independent or "electron c a v ity " model of Ogg (2 4 ),

Lipscomb (3 7 ), and Kaplan and K itte l (3 8 ).
first proposed by D eigeryet
A.

The most popular current theories can be d ivid ed

B„

The metal dependent or cluster theory,

a h , (39) and independently by Becker, et

a L , (1 7 ).

The "Electron C a v ity " Model
The "electron c a v ity " model was first proposed by Ogg (16) in 1946.

fo llo w in g features:

It has the

1) the metal is com pletely ionized into metal ions and electrons;

2) the electrons are located in cavities in the liq u id , the cavities having approxim ately
the volume of two to four ammonia molecules; 3) the electrons in the cavities are in
equilibrium w ith respect to the fo llo w in g reaction:

2 eT

~ e2 i

6 F° = - 0 . 2 e Dv .

where ©2 represents two electrons in one cavity w ith a n ti-p a ra lle l spins, e2

centers being

e n e rg e tic a lly stable w ith respect to e ] centers; 4) conduction at low concentrations is
essentially e le c tro ly tic since the measured m obility of the electron is some 200 fold less
than the minimum m obility expected for untrapped thermal electrons in a conduction band;
5 ) the conduction of concentrated solutions is essentially ele c tro n ic and probably represents
tunneling between adjacent centers-

20 ,

Assuming the electrons to be in a spherical c a v ity , Ogg calcu lated that the
radius of this cav ity should be about 10 ^ .

O g g “s theoretical value was obtained

using the p otential energy function

T/~|L =

~e
2a

/i _ 1 v

e = e le c tro n ic charge
a = radius of cavity
K= d ie le c tric constant

and neglected surface tension, electrostriction and electro n ic p o la riza tio n in the
d ie le c tric surrounding the spherical c a v ity .
More d e ta ile d calculations by Lipscomb (37) have shown that the cavities w hich
form the traps would be expected to have sizes about the same as that of the N H 3
m o lecule,

Lipscomb calculated the radius of the c a v ity to be 4 , 8

X

in fa ir agreement

o
w ith the experim ental value of 3 ,2 A obtained from density measurements.
Kaplan and K itte l (38) elaborate further on the electron c a v ity and show it can be used
to v e rify th e o re tic a lly the experim ental results obtained by Hutchison (12) on the para­
m agnetic resonance absorption of potassium ammonia solutions. They consider the electrons
to occupy m olecular orbitals on the protons of the ammonia molecules adfacent to the
c a v ity .
Coulter and students (40) have measured the heat of the follow ing reaction:

1 /2

e2 (am) + N H 4 = 1 /2 H 2 (g) + N H g (1)

They found a common heat of reaction per equ ivalen t of metal for both a lk a li and a lk a lin e
earth metals w hich indicates one and two electron ionization of the a lk a li and a lk a lin e earth
respectively in liquid ammonia.

Dew ald and Lepoutre (41) have included the e ffe c t of electron pairing and
io n -p a ir interactions in thermodynamic equations derived for the therm oelectric power
of metal ammonia solutions.

They conclude that these interactions do not account for the

anomalies in the therm oelectric behavior of metal and m e ta l-sa lt solutions in liquid
ammonia.

They are able to account for the observed anomalies by assuming th at the

electrons in solutions have a large negative heat of transport ( - 0 . 7 e „ v .)„

This large

negative heat of transport is accounted for by the assumption that the electrons move through
the solutions, even at high d ilu tio n , by a quantum tunneling process, from one cavity
to another or to a random "hole" in the solvent, rather than by ionic migration or a
conduction band process.
Jortner (36) attempts to interpret the physical properties of extrem ely d ilu te metal
ammonia solutions on the basis of Landau"s (42) model for electron binding by p o larizatio n
of the d ie le c tric medium.

It is assumed that the electrons are in a spherical w e ll created

by orientation of the permanent dipoles of the medium.

The part of the solvent p o larizatio n

e ffe c tiv e for binding is given by

PD = e/471" r2 ( i / D o - 1 /D s)

where r is the c av ity radius, DQ is the optical d ie le c tric constant, Ds is the static
d ie le c tric constant, and Pq is the part of the total polarization which cannot fo llo w
the motion of the e le c tro n .

The electrostatic potential for a spherical c av ity arising

from this p o la riza tio n is given by

22 .

Jortner assumes that t f \s a continuous function up to RQ, the c a v ity radius, and
that w ith in the c av ity ^

is constant. He does not include the effects of proton-electron

exchange forces and he assumes that the electro n ic wave function may extend beyond the
c a v ity radius.

He carried out his calculations using the variation m ethod, and assuming

a hydrogen-1 ike wave function for the ele c tro n .
The follow ing equation represents his energies for the Is state:
W |s ~ (h2/< 2 /8 z r 2m) - (Be2/ k Q) + (Be2/ k Q) 0 -/< R o) e x p (-2 //R 0 )
B

= ( 1 /D q - 1 /D S) ; A = l / r g S ; rgS = Bohr radius

The energy for the 2p state is given

W 2p =

by:

m) ” ( B e X ) + (Be2/ R Q)(1 + 3 / 2 ^ RQ + < *2r 2

+ 1 /3 <7^^Rq ) exp(-2c^R0 )
A

B are constants, and B = ^ 5 / V .
The value of R0 and the form of
and 2p state.

the potential w e ll are assumed to be the same for the Is

These calcu lated energies are further lowered by the e ffe c t of electro n ic

p o la riz a tio n . Jortner

calculated the Is and 2p energies for various values of RQ andfound

that values o f RQ between 3 .2 and 3 .4 5 A gave the best agreement

w ith experim ental d a ta .

Jortner interprets the near infrared absorption in m etal ammonia solutions to arise
from 2 p -i -ls transition where
H v = E 2p - E|s

H e also derives expressions to explain the temperature dependence of the position of
the band maximum.

The energy shift represents the difference between the change in

energy of the ground and the excited state.

d ( h ^ ) /d T

= dE2p/d T

-

dEls/d T

His fin a l equation for the tem perature dependence is

d (h /)/d t

where B = ( l / D 0 - 1 /D S)

= 0 .8 8 (dB /dt) -

, R0

0.2 7 7(d R o/d t )

= cavity radius.

The temperature dependence of the energy levels is attributed to the temperature
dependence o f the properties of the d ie le c tric medium. Jortner found no data a v a ila b le for
the tem perature dependence of D0 and R0 .

Assuming that the temperature dependence

o f Ds is much larger than that of DQ/ and assuming further that RQ is a linear function of
tem perature, Jortner presents the follow ing equation to represent the temperature dependen
of the e le c tro n ic e xc ita tio n energy.

h V (T) =

h v (240) + 1 .3 3 X 1O' 2 [e x p (2 4 0 /2 1 7 ) - exp (T /2 1 7 )]
+ 0 .2 7 7 (dR (/clO (240 - T)

Comparing his th eo retical calculations w ith the experim ental results of Blade and
Hodgins (11) the best agreement is obtained for RQ = 3 .2 5 A .

The temperature c o effic ie n t

o f the c a v ity radius (dR o/dt) is found to be approxim ately 3 x 10- 2 A /d e g .

24,

Experim ental density data of metals and th eir salts in liquid ammonia at various
temperatures should g ive information on the temperature dependence of the c a v ity radius.
C urrently Jortner is carrying out a more refined treatm ent using the self-consistent fie ld
method.
W h ile the c av ity model accounts q u a n tita tiv e ly for the e ffe c t of temperature and
concentration on the magnetic susceptibility o f d ilu te m etal-am m onia solutions, it only
p a rtia lly explains the results of absorption studies, and does not exp lain the q u a n tita tiv e
e ffe c t o f temperature and concentration on e le c tric a l conductance.

Recent models stressing

the role of the metal ion and containing much o f the essence of Kraus1 and Freed's model
have been proposed by Deigen (39) and by Becker, et a l . , (1 7 ).

Both models have the

same features but only the latte r w ill be discussed in d e ta il.
B.

The Becker Lindquist-A lder M odel
This model considers the solute to consist of four species:

1) a neutral species

c a lle d a monomer, consisting of an a lk a li metal ion surrounded by approxim ately six
oriented ammonia molecules; 2 and 3) a solvated metal ion and a solvated electron
arising from the dissociation o f the monomer; 4) a dimer consisting of two monomers bound
p rin c ip a lly by quantum -m echanical exchange forces.
In the monomer, the six ammonia molecules are presumed to be oriented on the
average w ith their nitrogens towards the metal ion so that the ion polarizes the ammonia
m olecule and further increases the electron a ffin ity of the protons.
in a roughly spherical orbit about the protons.

The electron circulates

Dissociation occurs when the electrons

overcome the coulombic attraction of the cation and become associated w ith the protons

2 5.

o f the bulk ammonia m olecules„ This dissociation is treated as being analogous to the
behavior o f an ion-pair„

W ith Increased concentration it is possible for the monomer to

dim erize or form higher clusters.

The energy for this binding arises from both exchange

forces and V an der W aals attractio n *

The proposed model calls for a consideration of the

com petition of two e q u ilib ria involving solvated electrons, (e " ), solvated cations, (M + ) ,
monomers, (M ), and dimers, (M 2 ) .

M =

M + + e“

Ki

= (M + )(e~)

(M)
2M =

Mo
2

K2 =

(M 2)--------(M )2

This model accounts q u a lita tiv e ly for the magnetic susceptability, the conductance
in d ilu te solutions, the densities, and also for the small variatio n o f these phenomena w ith
d iffe re n t a lk a li metals.

The absorption spectrum might contain the fo llo w in g e le c tro n ic

transitions:
1.

The solvated, ionized electron to a free e le c tro n .

2.

a) Radial transitions of the electrons associated w ith the monomer.
b) Change in angular momentum of the electron associated w ith

3.

a) Angular momentum and radial transitions of the dim er,
b) S inglet to trip le t transition of the dim er.

These transitions w ill be review ed again in the discussion.

the monomer.

2 6.

M cC onnel and Holm (43) have measured the Knight shift ( : H ) of the

and N a ^

n uclei in d ilu te sodium-ammonia solutions by nuclear magnetic resonance technique.
1A

discuss th eir results in terms of P (N

They

OO

) and P (N a

) , the average hyperfine contact densities

1A
OQ
o f N 14 and N a
nuclei a t an unpaired ele c tro n . They conclude that the odd electrons moved
in h ighly expanded orbitals about the sodium ions sim ilar to the results found for phosphorous
ion in P-doped s ilic o n .

N o Knight shift was found for the H^ n u c le i.

Their data are

consistent w ith the Knight shift to be expected for the monomer in the theory of Becker,
et a l . ,

(1 7 ).
Blumberg and Das (44) have obtained two w ave functions for the electrons in sodium-

ammonia solutions, from two potential wells based on the model proposed by Becker, et a I .
Their first w ave function is computed for the follow ing model: The monomer is considered
to be a unit point charge at the position of the N a ion and point dipoles and quadrapoles
a t the positions of the charge centroids of six ammonia molecules arranged in a regular
octahedron about the sodium ion.

In their second w ave function the p o te n tia l for the

electron is computed by taking the potential of the N and H nuclei as point charges
superimposed on a potential obtained from the e le c tro n ic bond structure of ammonia „
They use these w ave functions to calc u late the expected Knight shift in d ilu te sodiumammonia solutions by means of an n -e le c tro n formalism previously found successful in
treating the hyperfine interactions of F centers.
The results have been compared w ith the experim ental values of M cConnel and
H olm (43) for the Knight shift of N ^ and N a ^ in sodium-ammonia solutions, and agree
q uite fa v o ra b ly .

The proton shift is calculated to be less than 10“ ^ for a ll concentrations

27.

less than 0 .5 m olar, thus explaining why it was not observed by M cConnel and H olm .
Dependence of the Knight shift upon concentrations is p red icted , for concentrations
less than 0 .0 5 m olar, but these concentrations have not y et been studied e xp e rim en tally .
Evers and Frank (45) have derived a conductance function which gives good
agreement w ith experim ental values up to a concentration of 0 .0 4 N .

Their function

is based upon the mass action concept involving species described in models by
Deigen (3 9) and Becker, et a l . , (1 7 ).

Since conductance measurements have been

made w ith high precision in d ilu te sodium-ammonia solutions they fe lt justified in
introducing a c tiv ity coefficients for the charged species and treat the conductance data
according to methods developed for solutions of normal e lectro lytes.
The two competing reactions of the model are:

M

= M+ + e-

Kj = (M+)(e)
(M )-

M

K2 = (M2 ) 1/2
(M )

where the brackets indicate a c tiv itie s .
m e ta l, / j ,

Then

the ratio

K,

=

They let C be the to tal concentration of

of metal ions to total m etal, / 2 rQl‘‘0 of dimers to total m e ta l.

_______

c a - f t - 2*2)

where f is the a c tiv ity c o efficie n t of ionized m e ta l.
may be expressed in terms o f ^ \ .

K„ =

V

^

C ( l - / , - 2/ 2)

Upon dividing K-| by K2

2

They assume th at the conductance of the ionic species

follows the Shedlovsky equation (46) and they w r ite ^ j = ^ S ( z ) / t 0.
m o bility corrections.

2

where S(z) includes

Upon substitution and rearrangement they obtained a fin a l equation

w here the bracketed term accounts for the presence of dimers.
From the data of Kraus on the conductance in sodium-ammonia solutions they are
able to obtain values of K^ and ^

equal to 7 x 10“ ^ and 2 7 , resp ectively.

Values for

the same constants obtained from Hutchisons param agnetic data on potass I um-ammonia
solution are 3 x 10” 2 ancj 9 9 ^ respectively. Both sets of constants are of the same order
o f magnitude although they d iffe r by a factor of four.
Further support for the theory of Becker, et

a l . , comes from the studies o f small

angle X - r a y scattering in m etal-am m onia solutions by Schmidt (47) in approxim ately
one molar solutions.
o
o f the order of 15 A .
the metal ion present.
order of 8 %
present.

The data indicate the presence of scattering centers w ith dimension
Schmidt also found that the type of scattering center depended upon
H e found no evidence for the presence of cavities w ith diameters of the

C al culations indicate that these cavities would have been observed if

H e interprets his results to be in q u a lita tiv e agreement w ith the theory of

Becker, §t__

29.

ML

EXPERIMENTAL

Instrumentation
The Beckman Model D K -2 Spectrophotometer is suitable for measuring the
absorption spectra of d ilu te solution of a lk a li metals in liquid ammonia in both the
v is ib le and near infrared regions.

The only change that was found necessary was the

construction of a special apparatus for cooling.
Passerini, et

a I .,; (4 8 ).

ft is a m odification of the apparatus of

The apparatus, s p e c ific a lly designed for use w ith the Beckman

D K -2 Recording Spectrophotometer is suitable for qu an titative measurements of solution
spectra between room tem perature and dry ice temperature (- 7 5 ° C ) w ithout modifying
the spectrophotometer.
The general requirements for an op tical cell for low temperatures are listed
by Beale and Roe (4 9 ), and the present apparatus incorporates many of these features,
ft consists essentially of a rectangular metal insulated vessel in w hich the Pyrex
absorption cells can be p la ce d . When in use the Pyrex cells are in contact w ith metal
w hich is surrounded by the cooling bath.
Description of the Apparatus
The d e ta ile d description of the apparatus is shown In Figures 1 and 2 .

The

dimensions included are those w hich are necessary for the Beckman D K -2 Recording
Spectrophotom eter.
rectangular can.
cooling them .

The vessel consists of two detachable parts, an inner and an outer

The inner can is used for holding the absorption cells as w e ll as for

The outer can serves to insulate the inner can.

30„

The inner container consists of two rectangular boxes w ith the follow ing
dimensions:

1) 2 ,4 6 " x 2 , 4 6 s1 x 5 s",

2) 2" x 0 .7 " x 6 11. They were made by bending

0 .0 4 0 inch copper sheet metal into the required dimensions and silver-soldering the loose
ends together.

Two 1 1 /1 6 " holes were cut 1 .1 8 inches a p art, and 1,0 2 inches from the

bottom in each can to a llo w the lig h t beam to pass through.

Then two 11/16" brass

tubes w ere extended through both boxes and soldered to each box.

The brass tubing

was then m illed out of the innermost box to provide room for the absorption c ells.
Construction details are shown belo w .

/y V

AV

o

O
Front

Bottom

N e x t a bottom piece was cut for the innermost container.
d rille d in this piece 0 .3 inch apart.

Bottom

Then two 1 /1 6 " holes were

Two pieces of 1 /1 6 " copper rod 1 cm, long were

soldered to this bottom p la te , as shown below , to provide a method for positioning the
cells in the innermost container.

Bottom

+1
0 .3 "
Side

(Windows and Insulation
om itted)
Figure 1. Isometric V ie w of Assembled Vessel

32.
Plastic 1 /2 0 " th ic k

1 /4 " O D —

3 /8 " O D

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Figure 2 . Cross-section of Figure 1 through A - A

Pn

T

0 . 68 "
_

t

0 .4 "
A
i

3 3.

N e x t the bottom was soldered to the innermost container. Then a longer piece of
copper sheet was cut and soldered to the bottom of the next larger container . The
top piece of this container differed from the bottom piece in having holds cut for the in le t
and o u tle t flow o f the cooling medium as w e ll as a hole for the extension of the innermost
container above this container.

i
a /v

ii

!'1
1

i1 i 1

i 1

i

i

A /V

Side
When the top was soldered in p la c e , a 1/ 8 M copper tube was extended through the
in le t hole about 3 inches and soldered in p la c e .
through the o u tle t hole flush w ith the cover.

A piece of 3 / 8 '1 tubing was soldered

A ll solder joints were checked for leaks

by submerging the container in w ater w ith about 2 atmospheres of H eliu m pressure
on the inside.
The outermost container is essentially a rectangular box to provide insulation
for the inner container.

It was made by bending 0 .0 4 0 inch copper sheet metal and

soldering the ends and bottom together. Two holes were cut in each end to provide
for the lig ht path.
be two 1 cm.

It was decided that the simplest type of window for this outer can would

Pyrex absorption c e lls , one on either side.

Two brass strips w ith 1 cm. square

holes were cut and soldered over the holes in the outer can. The outer container was
mounted on a m illin g machine and the brass strips were m illed [ust enough to insure that
the outer edges w ere p a ra lle l.
w ith 3 - M brand cem ent.

N e x t the 1 cm. Pyrex absorption cells were glued on

34.

To provide insulation between the two detachable cans, holes were cut in
styrofoam blocks. There was approxim ately 1 inch of styrofoam between the two cans at
a ll points except the bottom where is was approxim ately 1 /2 inch th ic k . This lesser
amount of insulation was necessary to a llo w the light path to be continuous through
the apparatus and was a lim itation imposed by the spectrophotometer used.
Cooling Arrangement
The cooling cycle is essentially that shown in the block diagram in Figure 3 .
Absolute ethyl alcohol was used as the coolant.

The system was fille d by introducing

a T tube into the line connected to a dropping funnel fille d w ith a lc o h o l, w hich acted
as a reservoir to make up for the contraction of the alcohol w ith cooling.

Trapped air

could also escape through this T tube.
The alcohol passed through the pump into 3 / 8 inch copper coils immersed in a
dry ic e -a c eto n e bath, then into the heaters and metal Dewar and back to the pump.
The rough heater consists of a short piece of a heating elem ent of an e le c tric stove
enclosed in a piece of one inch copper tubing w ith 1 /4 inch in le t and ou tlet tubes.
The fin e heater consists of a k n ife -e d g e heater enclosed in the same manner.
tubes for both heaters are insulated w ith styrofoam.

The copper

The rough heater was controlled w ith

a v aria c w h ile the fin e heater was controlled by a thermistor elem ent w ith in the metal
Dewar.

35.

METAL
DEWAR

CONTROL
HEATER

ROUGH
HEATER

COLD
BATH

Figure 3 . C o d in g Arrangement

36.

Thermoregulator
The c irc u it shown in Figure 4 is an adaptation of a thermoregulator designed by
B urw ell, et g | .

(5 0 ).

The temperature sensitive elem ent is a Fenw all thermistor

in a Wheatstone bridge c irc u it, w h ile the heater current is controlled by means of a
saturable reactor in series w ith the heater.

W ith this regulator, the temperature in

the cell compartment could be m aintained w e ll w ithin ± 1°C at any point between
- 3 0 ° C and - 7 0 ° C .
The actual temperature in the v ic in ity of the absorption cells was measured
using another thermistor. This thermistor is located in a small piece of aluminum block
w hich served to position the cells.
C e ll H older
This small block of aluminum has holes cut in the bottom which a llo w it to fit
over the pins in the bottom piece of the cell compartment.

The sides were m illed

f la t and two pieces of aluminum sheet w ith holes cut for the light path are held
against the outside edge by small screws.

0 0

Bottom

o

o o
Front

c

'

1 °
Top

1

37.

•vJiAJU p>

circu it

TOP ''

v- H

-|I

4.

W

Figure

h-

Thermoregulator

•A A A /

( Icvl

N

OD

)

IM IU

IHMI

f

to

X

o

.0000.

I*
0(0

-0 > =

38.

TABLE II
VALUE CHART
FOR THERMOREGULATOR

Symbol

V a lu e

Symbol

V alu e

C -l

0 .0 5 M FD

600V

R8

1 K

1 /2 W

C -2

0 .2 5 M FD

2$ V

R9

5 K

1 /2 W

C -3

20 M FD

25V

R10

1 M

1 /2 W

C -4

0 .0 5 M FD

600 V

R ll

1 M

1 /2 W

C -5

0 .2 5 M FD

600V

R12

0 .1 5 M

1 /2 W

C-6

0 .0 5 MFD

6 00 V

R13

27 K

1 /2 W

C -7

2 0 M FD

2 5V

R14

5 K

1 /2 W

C- 8

0 .0 5 M FD

600V

R15

10 K

POT

C -9

0 .5

M FD

600V

R16

15 M

HELIPOT

C -1 0

0 .5

M FD

6 00 V

R 17

10 K

2 W

c-11

10 M FD

450V

R18

22 K

2 W

C -1 2

10 M FD

450V

R19

22 K

2 W

C -l 3

4 0 M FD

150V

R20

2 M

C -1 4

20 M FD

3 50 V

R21

0 .0 4 7 M

C -1 5

10 M FD

300V

T1

6 .3 V -3 A M P - 5V -2 A -3 5 0 -

T2

PRI. 110V

1 /2 W
1 /2 W

SEC - 3 V

39.

TABLE II (continued)

Symbol

V alu e

V alu e

Symbol

110V .

OUT

R1

0 .5 K

1 /2 W

T3

110V . IN

R2

2 .2 M

1 /2 W

T4

Saturable Reactor Ironcore

R3

2 .2 M

1 /2 W

VI

6SJ7

R4

0 .5 M

1 /2 W

V2

6SJ7

R5

1M

1 /2 W

V3

6 J5

R6

5 K

1 /2 W

V4

6 AC7

R7

1 M

POT

V5

80

40.

The cells are held in position w ith this cell holder which insures that th e / are
perpendicular to the light path. Small fe lt strips were glued to the aluminum strips to
reduce the danger of breaking or scratching the cells.
p lace by means of a hollow hard rubber tube.
into the top of the aluminum b lo ck.

The cell holder was lowered into

The end of this rod is threaded and screwed

The temperature measuring thermistor and wires are

contained in this hollow rod w ith the thermistor just protruding from the end into the
aluminum b lo c k .
Preparation and P urification of Solvent and Metals
The sodium (Baker’s A n alyzed ) and potassium (M allin ck ro d t) came in the form
o f rods, and pieces were cut as needed.
sodium and potassium.

The same method was used for purifying both

The metal was cut into small pieces under ether and placed in

the degassing tube shown in Figure 6 .

A fte r the tube was evacuated, the metal was

m elted and allow ed to run through the constrictions leaving most of the oxide at these
points. It was then degassed by gently heating the molten metal until the formation of
bubbles ceased.

When c o o l, the tube containing the metal was broken and q uickly

transferred to the solution makeup vessel where it was degassed further and then d is tille d
under high vacuum (c a . 2 x 10“ ^ mm).
Ammonia
The liquid ammonia used as the solvent in these experiments was obtained from
O lin Matheson Company In c .

The ammonia was distilled from the commercial cylinder into

a smaller evacuated tank w hich had been previously charged w ith sodium. Condensation
o f the ammonia was e ffected by means of a dry ic e -a c eto n e bath. This storage tank holds

41.

about four liters of liquid ammonia and was kept at room tem perature.

The system

used for further purificatio n is shown in Figure 5.
A fte r complete evacuation of the system, ammonia was d is tille d from the storage
tank into container C w hich had previously been charged w ith sodium and outgassed by
heating under high vacuum (10” ^ m m .).
the co olant.

Dry ice -tric h lo ro e th y le n e baths were used as

When the container was f u ll, the cylinder v alv e A and stopcock B were

closed and the ammonia was allow ed to stand over the sodium at approxim ately - 5 0 ° C
for about th irty minutes.

Stopcock B was then opened slowly for a short tim e to allo w

any residual gases to be pumped o ff.
opened.

Then stopcock B was closed and stopcock E was

The cold trap was lowered to a llo w a b u ild -u p of ammonia pressure, which forced

the metal ammonia solution through the fritted glass tube D , and over into container F which
was cooled by a dry ic e -tric h lo ro e th y le n e bath. Care was taken not to a llo w the liquid
lev e l in container C to get below the fritte d glass tube.
gases present in C would also have gone over into F .

Otherw ise any non-condensable

When transfer was com pleted, stopcock

E was closed and the cold trap was again raised around C .

The ammonia could then be

d is tille d from this blue solution into other parts o f the system when needed.

Blue so lut ions

have been stored in F up to two weeks at - 7 5 ° C and showed no visible signs of
decom position.
Experim ental Procedure
The first attempts to obtain a sodium-ammonia spectrum in a 1 cm quartz absorption
c e ll met w ith

no success.

The apparatus used in making up these solutions was a make-up

vessel (9) w ith the absorption c ell attached by a ball jo in t.

The extrem ely d ilu te solution,

Figure

UJ
5. Solvent Purification

Syste

HIGH

VACUUM

4 2.

UJ

2

H
CO

>

<

X

43.

c a . 5 x lO -^ M , decomposed w ith in 10 minutes after transfer Into the absorption c e ll.
Th is decomposition was probably due to traces o f w ater absorbed on the w alls of the
c e ll w hich could not be outgassed as w e ll as any impurities picked up w h ile the solution
was being transferred through numerous stopcocks and ball joints.
The next type of apparatus used was one in which it was possible to d is till both
the sodium and the ammonia d ire c tly into the absorption c e lL

This met w ith little success

for most of the solutions were too concentrated to a llo w measurement of the spectra,,
F in a lly an a ll glass apparatus, Figure 6 was \constructed which to a certain
degree allow ed one to adjust the concentration by adding ammonia in various amounts up
to 150 cc. so that the absorbance would lie between 0 and 2 .
Description of a Typical Run
The description of a typical determ ination is given below .

The cell is rinsed w ith

hot dichromate cleaning solution and then rinsed about 6 times w ith d is tille d w a te r. The
fin a l two aqueous rinses are made w ith dem ineralized w ater. This dem ineralized w ater
is from the same b ottle as the w ater used for making the standard solutions for the flam e
photom eter.

The c e ll is then dried overnight in an oven at 11 0 °C .

Several small pieces of metal are cut from reagent grade sodium and placed in the
degassing tube shown in Figure 6 . A fter the tube is evacuated it is warmed g ently until
the metal melts and runs down through the constriction.
is heated a t intervals until a ll dissolved gases escape.

Then the molten sodium metal
W h ile this tube is cooling the

c e ll is connected to the high-vacuum system by means of a 1 2 /5 fem ale b a ll joint using
A piezon

"W " W a x .

When the degassing tube is cool, the lower end containing the metal

44.

35

SOLUTION
MAKEUP
CELL
METAL
DEGASSING
TUBE

Figure 6
Left:
Solution M a k e -u p C e ll
Right: M e ta l degassing tube

45.

is broken off and placed upside down into the 2 4 /4 0 male taper.

The 2 4 /4 0 taper is

e
then capped and the system evacuated to approxim ately 10

mm. H g .

The lower part

of the standard taper is warmed gently until the metal melts and runs through the two
constrictions into bulb B shown in Figure 6 ,

The standard taper then is removed by m elting

the glass under vacuum at the lower constriction.
d is till back and forth in the bulb.

The metal is heated further causing it to

This is continued until the pressure, as indicated by a

V eeco ion tu b e , remains below 2 x 1 0

“5

mm. Hg.

N e x t the entire c e ll, except for the bulb containing the m e ta l, is covered w ith
aluminum fo il and asbestos sheet and heated to 3 5 0 °C for about 6 hours to remove traces
o f w ater adsorbed on the w alls .

W h ile this heating takes place liquid ammonia is purified

as previously described.
Approxim ately 1 -2 mg of metal is distilled into tube D shown in Figure 6 .

Then

a dry ic e -a c e to n e trap is raised around the tube D and ammonia is condensed until a
flashlight held behind the 35 mm tube can barely be seen.
the absorbance usually lies between one and tw o.

When the light is just visib le

M ore ammonia can be added if more

d ilu te solutions are to be studied.
The tube and absorption c e ll are then sealed o ff from the rest of the apparatus
by heating the h e a v y -w a lle d glass tubing at C , Figure 6 , w ith a hot fla m e.

Dry

ice is added to about 15 liters of waste alcohol in a large glass container until the
tem perature is about - 7 5 ° C .

The tube w ith the m etal-am m onia solution and attached

absorption c e ll are then cooled in this bath until a uniform temperature is o btained.

By

tiltin g the tu b e , the blue solution flows over into the absorption c e ll. Then by tiltin g it

46.

in the other direction it is run out of the absorption cell back into the tube.

The

absorption c e ll is rinsed about six times in the above manner and then fille d w ith
about 1 /2 cc of blue solution.

Then the absorption cell is sealed off q u ickly by heating

the h e a v y -w a lle d glass tube a t E.

It is stored in a dry ice -a lc o h o l bath until the time

o f the op tical measurements.
A N A LY S IS
Ammonia in the C e ll
The ammonia from the absorption c e ll was slowly distilled into a 75 cc a lliq u o t
o f 0 .4 4 8 5 N H C I by surrounding the c e ll w ith an acetone bath slightly warmer than the
boilin g point of the ammonia.

O ccasionally there would be some bumping w hich would

causethe loss of some of the m etal/ but this could be m inim ized by careful control of the
tem perature of the external b ath . Care was required/ when the ammonia finished d is tillin g ,
to prevent the aqueous acid solution from being forced back into the c e ll.
a cid was titrated to a m ethyl-red end point w ith 0 .0 9 8 4 N

The rem aining

NaO H.

Ammonia in the Tube
The ammonia from the large tube (approxim ately 5 0 cc) was d istilled into a stainless
steel bomb.

This bomb was constructed by d rillin g a 0 .8 inch hole in a one-in ch stainless

steel rod leaving approxim ately 0 .1 inch sides and bottom. A stainless steel Hoke (Hap)
v alv e w ith a steel 1 0 /3 0 male standard taper was w elded to the open end.

This bomb was

designed to be strong enough to hold the ammonia at room temperature so that it could be
a cc u ra te ly w eighed.

47,

A fte r the ammonia in the large tube was frozen w ith liquid nitrogen the side tube
was broked o ff and connected to the stain lesS'Steel bomb,

A mercury manometer was

included in the system. W h ile the ammonia in the tube was still frozen the a ir in the
system was evacuated w ith a vacuum pump.

A fter the pump was shut off from the system,

the ammonia was allow ed to warm to its boiling point and was d istilled into the stainless steel
bomb, cooled w ith a dry ice -a c eto n e bath.

The bomb was reweighed after warming to room

tem perature.
M e ta l
The first few analyses for sodium were carried out using the trip le -a c e ta te .method.
H o w e ve r, it was necessary that in some of the tubes the total sodium content be less than
-3
10
grams.

-3
Since 10
grams o f sodium gives only about 72 milligrams of sodium zin c

uranyl a c e ta te , it was fe lt that this method would not give sufficiently accurate results.
Also it would be impossible to check the sodium content in the absorption c e ll where the total
w eig ht was on the order of 10” ^ grams of sodium.
A fte r the prelim inary runs, the concentration of metal was found w ith the aid of a
Beckman M odel B Flame Photometer. Standard solutions of sodium were made by w eight
from sodium chloride w hich had been re crystallized from conductivity w ater.

The salt was

dissolved in conductivity w ater and aliquots were diluted to give one lite r solutions containing
0 , 2 , 4 , 6 , 8 , and 10 parts per m illion of sodium, respectively.

F ifty m l. of a 10s 1 H C I~

C a C 0 3 solution were added to each lite r of the standard solutions to e lim in a te effects of pH
on flam e in tensity.

Standard potassium solutions were made up in a sim ilar manner. Both

4 8.

sodium and potassium standard solutions gave a straight line when per cent transmission was
plotted versus concentration as shown in Figure 7 .
U su ally, when the metal in the tube was taken up in 200 cc of w ater it gave a reading
w h ich was on s cale.

O ccasionally a solution had to be further d ilu te d .

The metal in the

absorption c e ll could be taken up in 10 cc water and s till give a reading which was on s cale.
The to tal w eight of metal in any solution was found from the scale reading and the calib ratio n
curve.
As a check on the w avelength scale incorporated in the Beckman D K -2 Recording
Spectrophotometer, measurements of the absorption spectrum of pure liquid ammonia In a 0 .1 mm
absorption c e ll were made.

The absorption bands found are in good agreement w ith those of

J o lly (25) as shown in Table I I I .
Figure 7 .

The absorption spectrum of pure liquid ammonia is shown in

N o detectable difference was noted in the spectrum as the temperature was varied

from -3 5 ° C to - 6 5 ° C .
Temperature Measurement
For making temperature measurements a Fenwall type GB32p8 thermistor was
calib rated against a platinum resistance thermometer.

The variation of the thermistor

resistance w ith temperature is given by the follow ing equation

Log R = A + B /T
where A and B are constants characteristic of each thermistor R is the resistance of the
thermistor at the absolute temperature T.

The data for the calibration are given in Table IV

and plotted in Figure 8 as Log R vs. 1 /T .
The resistance of the thermistor was measured by means of a Wheatstone bridge.
decade

r e s is ta n c e

The

box used in the Wheatstone bridge was calibrated against a m odified Jones

Conductance bridge.

49.

TABLE III
A B SO R PTIO N SPECTRUM OF A N H Y D R O U S L IQ U ID A M M O N IA

W avelength
(microns)

Absorption Intensity
0.1 mm C ell

J o lly (25)

2 .2 4 0

very strong

2 .2 4

1 .9 9 5

very strong

2 .0 0

1 .6 4 0

weak

1 .6 4

1 .535

strong

1 .53

1 .3 1 5

weak

1 .3 0 5

1 .2 7 5 *

weak

1 .2 7 0

1 .2 2 5

weak

1 .2 3 2

1 . 200*

shoulder

1 .20

1 .0 4 0

weak

1 .0 4

1 . 010*

shoulder

1.01

0 .8 0 0 *

weak

0 .8 0 2

* D etectable only w ith 1 mm c e ll.

50.

o
o

-i o

pH

CO

Photometer
of Anhydrous Ammonia

Cl

o
o

00

t>

of Flame
Spectrum

o
o
o

s

o

o
o
o

Calibration
Absorption

_

rH

o
o
©

ov

CO

120

o
o

o

o

o

o

o
rH

aouB^q-TuisuBax

.iad

a o u B q jo s q v

to

o

Figure

o
o
o

7„

Top:
Bottom:

Cl

51.

TABLE IV
THERMISTOR C A LIB R A TIO N

Tem perature,
°C

1q3 /T ^ s o lu te

R in (Ohms)

Log R

2 5 .0

3 .3 5 0

2,020

3 .3 0 5

1 3 .0 2

3 .4 9 4

3 ,2 2 8

3 .5 0 9

-1 .0 5

3 .6 7 5

6 ,1 0 7

3 .7 8 6

- 1 1 .9 2

3 ,8 2 8

10, 012

4 .0 0 1

- 1 9 .7 8

3 .9 4 6

1 4 ,8 1 0

4 .1 7 0

- 2 6 .7 2

4 .0 5 7

2 1 ,2 6 3

4 .3 2 8

- 3 2 .0 9

4 .1 4 8

2 8 ,3 5 3

4 .4 5 3

- 3 8 .8 9

4 .2 6 8

4 1 ,9 2 3

4 .6 2 1

- 4 6 .3 2

4 .4 0 8

6 5 ,0 7 0

4 .8 1 3

- 5 3 .4 3

4 .5 5 1

101 ,16 9

5 .0 0 5

- 6 0 .4 0

4 .7 0 0

158 ,72 4

5 .1 8 9

- 6 5 .5 8

4 .8 1 7

2 2 4 ,9 8 2

5 .3 5 2

- 7 1 .8 6

4 .9 6 7

348,471

5 .5 4 2

-7 5 .2 1

5 .0 5 1

4 4 2 ,7 5 5

5 .6 4 6

-7 8 .2

5 .1 2 8

5 4 4 ,7 7 0

5 .7 3 6

51.

TABLE IV
THERMISTOR C A LIB R A TIO N

Tem perature,
°C

.
10

absolute

R in (Ohms)

Log R

2 5 .0

3 .3 5 0

2,0 20

3 .3 0 5

1 3 .0 2

3 .4 9 4

3 ,2 2 8

3 .5 0 9

-1 .0 5

3 .6 7 5

6 .1 0 7

3 .7 8 6

- 1 1 .9 2

3 .8 2 8

10, 012

4 .0 0 1

- 1 9 .7 8

3 .9 4 6

1 4 ,8 1 0

4 .1 7 0

- 2 6 .7 2

4 .0 5 7

2 1 ,2 6 3

4 .3 2 8

- 3 2 .0 9

4 .1 4 8

2 8 ,3 5 3

4 .4 5 3

- 3 8 .8 9

4 .2 6 8

4 1 ,9 2 3

4 .6 2 1

- 4 6 .3 2

4 .4 0 8

6 5 ,0 7 0

4 .8 1 3

- 5 3 .4 3

4 .5 5 1

1 0 1 ,1 6 9

5 .0 0 5

- 6 0 .4 0

4 .7 0 0

1 5 8 ,7 2 4

5 .1 8 9

- 6 5 .5 8

4 ,8 1 7

2 2 4 ,9 8 2

5 .3 5 2

- 7 1 .8 6

4 .9 6 7

348,471

5 .5 4 2

-7 5 .2 1

5 .0 5 1

4 4 2 ,7 5 5

5 .6 4 6

-7 8 .2

5 .1 2 8

5 4 4 ,7 7 0

5 .7 3 6

5.4

Figure

8.

Thermistor

Calibration

52.

53,

IV ,

RESULTS

The absorption data were treated in the conventional manner.

The laws of

Bouger and Beer deal w ith the decrease in intensity of monochromatic light when the
lig h t passes through an absorbing medium,

Bougerls law states that when a beam of mono*"

chrom atic light is passed through an absorbing medium, each infinitesim ally small layer
of the medium decreases the intensity of the beam entering that layer by a constant
fra c tio n .

Beer’s law states that the rate of decrease of the intensity is d ire c tly proportional

to the concentration of the absorbing species.
The fin a l form of the combined law is given below:

I =

lQ 10' abc

w here
I

is the intensity of the transmitted lig h t.

1 is the intensity of the incident lig h t,
a

is the molar absorbancy index,.

b

is the distance the light travels through the solution in centimeters,

c

is the concentration in moles per lite r.

It follows that

A =

log I q/ | ~ abc

where A is c alle d the absorbance.
from measurements o f A , b , and c .

The molar absorbancy index, a , can be calculated

54.

The absorbance A , for each m etal-am m onia solution was measured d ire c tly
w ith a Beckman D K -2 Recording Spectrophotometer.

The c ell lengths, b for the cells

used were calcu lated from the measured absorbance of a standard solution of potassium
chromate in 0 .0 5 N potassium hydroxide (5 1 ).

The concentration of this standard solution

was adjusted to give an absorbance of 0 .7 5 at 375 millimicrons for a c ell length of 10”^ cm.
The light paths of the three cells used for m etal-am m onia solutions were found
to be 1 .0 1 5 x 10"^ c m ., 1 .0 9 5 x 10 ^ c m ., and 1.21 x 10 ^ c m ., w h ile the reference
c e ll was 1 .1 0 x 10“ ^ cm.

The above values were used in a ll calculations.

The concentrations of the m etal-am m onia solutions were expressed in moles per
lite r.

The assumption was made that the density of these d ilu te solutions (10“3 m ) ;s th e

same as the density of pure ammonia.

The density of liquid ammonia at - 6 5 ° C was obtained

by extrapo lation of the data of Cragoe and Harper (52) to - 6 5 ° C .

The value used

was 0 .7 2 0 gram per cc.
T yp ical Sodium-Ammonia Absorption Curve
The data from

a ty p ica l sodium-ammonia absorption curve are given in Table V .

The data are plotted as wavenumber (Figure 9 ), and wavelength (Figure 10) versus A / A ^
where A is the absorbance at any w avelength and A |^ ax

Is the absorbance at the peak.

Since this type of graph always has a maximum value of u n ity , it is useful for examining
the effects of concentration and temperature on the shape of the absorption curve.

The

shape o f these reduced curves should not be altered by small amounts of decomposition.
Figures 9 and 10 show that the absorption curve is nearly symmetrical when plotted on a
w avelength s c a le , but is asymmetrical on a wavenumber or energy scale.

qx

55.

E ffect of Concentration on Sodium Spectrum
Absorption d a ta , A / A M ax , for three concentrations of sodium-ammonia solutions
ot - 6 5 ° C are given in Table V I and plotted in Figure 11 „ W ithin experim ental error the
shape of the absorption curves for sodium-ammonia solutions is independent of concentration„
The slight scatter of the points is not a systematic function of concentration and
probably arises from the technique used in measuring the spectra„

In the present experim ental

set-u p , there was no satisfactory method for measuring the 100 per cent transmission (or zero
absorbance) curve as a function o f w av e len g th . The general instrument response over the
range covered was measured by scanning w ith a ir in both the sample and reference beams„
This gave a horizontal straight lin e .

The cells containing the m etal-am m onia solutions

w ere stored in a dry ice -a lc o h o l bath. When the absorbance of a solution was to be measured,
the sample and reference cells were removed from the alcohol bath, and the excess alcohol
was w iped o ff w ith tissue paper. The cells, retaining a thin film of alcohol over the windows,
w ere placed in the cell holder.

It was found that the thickness of the film affected the

absorbance by the same amount for a ll wavelengths. This made it impossible to determine
an absolute zero absorbance line using pure ammonia in both sample and reference cells as
the film thickness could not be reproduced.

H ow ever, the alcohol film can only cause a

v e rtic a l displacement of the whole curve. A t wavelengths shorter than 700 m illim icrons,
the absorbance approaches a constant v a lu e .

In this research it is assumed that the absolute

absorbance at these wavelengths is essentially zero. Calculations based upon the molar
absorbancy index obtained by M iranda (19) at 500 millimicrons indicate that this assumption
would introduce an error of not more than 1 .5 per cent, which is below the error introduced
by decomposition and uncertainty in the analysis of these solutions.

56.

TABLE V
D A TA FR O M TY P IC A L S O D IU M -A M M O N IA ABSO RPTIO N CURVE

Na

N o . 25

W avelength (m/<)

C

=

1.631 x 10' 3 m /l

Peak

=

1470 rryj. = 6800 cm” ^

A M ax.
T

=
=

° ‘ 78
- 6 6 °C

W ave Number (cm

A-

A /A M

qx

600

1 6 ,6 6 7

0.000

0.000

650

1 5 ,3 8 5

0.002

0 .0 0 3

700

1 4 ,2 8 6

0.010

0 .0 1 3

750

1 3 ,3 8 3

0 .0 2 8

0 .0 3 6

800

1 2 ,5 0 0

0 .0 4 5

0 .0 5 8

850

1 1 ,7 6 5

0 .0 6 8

0 .0 8 7

900

11,111

0 .0 9 8

0 .1 2 6

950

1 0 ,5 2 6

0 .1 3 5

0 .1 7 3

1000

10,000

0.1.80

0.231

1050

9 ,5 2 4

0 .2 4 0

0 .3 0 8

1100

9 ,0 9 1

0 .3 0 0

0 .3 8 5

1150

8 ,6 9 6

0 .3 8 0

0 .4 8 7

1200

8 ,3 3 3

0 .4 7

0 .6 0 3

1250

8,000

0 .5 6 0

0 .7 1 8

1300

7 ,6 9 2

0 .6 4 0

0.821

1350

7 /4 0 7

0 .7 2

0 .9 2 3

5 7.

TABLE V (continued)

W avelength (mAJ)

W ave Number (cm~^)

A _________

M ax

1400

7 /1 4 3

0 .7 6 0

0 .9 7 4

1450

6 ,8 9 7

0 .7 8 0

1.000

1500

6 ,6 6 7

0 .7 7

0 .9 8 7

1550

6 ,4 5 0

0 .7 4

0 .9 4 9

1600

6 ,2 5 0

0 .7 0

0 .8 9 7

1650

6 ,0 6 0

0.6 1

0 .7 8 2

1700

5 ,8 8 0

0 .5 4

0 .6 9 2

1750

5 ,7 2 0

0 .4 6

0 .5 9

1800

5 ,5 5 0

0 .3 8

0 .4 8 7

1850

5 ,4 0 0

0 .3 0

0 .3 8 5

1900

5 ,2 6 0

0 .2 3

0 .2 9 5

1950

5 ,1 3 0

0 .1 7

0 .2 1 8

2 00 0

5 ,0 0 0

0 .1 1

0.141

58.

Na No. 25
A/AMax. I I

cm 1

T = -66 °C

0.6

0.4

17,000

15,000

13,000

11,000

9,000

7,000

5,000

Cm-1
Figure 9.

Typical Sodium-Ammonia Absorption Curve

(Wavenumber)

Figure

10.

Typical

Sodium-Ammonia

Absorption

Curve

(Wavelength)

59 o

o o

60.

TABLE V I
EFFECT OF C O N C E N T R A T IO N O N S O D IU M SPECTRUM AT - 6 5 ° C

N a N o . 25

N a N o . 28

c = 1 .2 6 x 10“ $,

c = 1.61 x 10“ $ , y K

b = 1.21 x 10- 2 cm
a = 4 4 ,5 0 0 1. mole -1 cm-1

b = 1 .0 1 5 x 10~2Cm.
a = 5 6 ,0 0 0 1. mole- 'cm

c = 3 . 89 x 10" i n , / l .
b = 1 . 015 x 10“ 2 cm.
a = 46 i, 0 0 0 1. moIe“ ^cm

%
<

^M ax.

: 0 .6 6 5 a t 6 ,9 0 0 cnrf^

I!

i

N a N o . 10

i ™
cm. - 1

A

^^M ax

.

A

0 .7 8 at 6 ,8 0 0
cm“ l
^ M a x .

A M a x .= 1 *83 a t‘6 ,8 0 0 cm"
A

^^M ax

11,110

0 .0 9

0 .1 4

0.11

0 .1 4

0 .2 8

0 .1 5

10,000

0 .1 6

0 .2 4

0.20

0 .2 5

0 .4 5

0 .2 5

9 ,1 0 0

0 .2 7

0 .4 1

0 .3 3

0 .4 2

0 .7 5

0.41

8 ,3 3 0

0 .4 2

0 .6 2

0 .4 9

0 .6 2

1.10

0 .6 0

7 ,6 9 2

0 .5 6

0 .8 4

0.66

0 .8 5

1 .4 8

0.81

7 ,1 4 3

0.66

0 .9 9

0 .7 7

0 .9 9

1 .7 8

0 .9 7

6 ,6 6 7

0 .6 4

0 .9 6

0 .7 7

0 .9 9

1 .8 2

0 .9 9

6 ,2 5 0

0 .5 8

0 .8 7

0 .6 9

0 .8 9

1 .6 9

0 .9 2

5 ,8 8 0

0 .4 5

0.68

0 .5 4

0 .7 0

1.41

0 .7 7

5 ,5 5 0

0 .3 2

0 .4 8

0 .3 6

0 .4 6

0 .9 7

0 .5 3

5 ,2 6 0

0.22

0 .3 3

0.21

0 .2 7

0 .6 3

0 .3 4

5 ,0 0 0

0.11

0 .1 7

0.10

0 .1 3

0 .3 5

0 .1 9

.

Figure

11,

Effect

of

Concentration

on

Sodium

Spectrum

at

-65

61.

00

©

o

©

’X% / V

o

o
©

62.

TABLE V I I
M O LA R A B SO RBA NCY IN D E X FOR S O D IU M -A M M O N IA S O L U T IO N AT - 6 5 ° C

Solutlon

oncen ra ,ono
(M o la rity x I f f - 3)

A
(for b = 1Cr2 c m .)

M olar
Absorbancy Index
(I. mole- 1 cm?)

Peak
Position
(c m "1)

6

0 .2 7

0 .1 1 3

4 1 ,8 0 0

7080

7

0.68

0.220

3 2 ,4 0 0

7020

10

1 .2 6

0 .5 5 0

4 3 ,6 0 0

6900

13

1 .2 2 5

0 .5 3 0

4 3 ,3 0 0

6900

14

0 .9

0 .2 1 5

2 3 ,9 0 0

7150

15

1 .0

0 .2 8 0

2 8 ,0 0 0

7000

17

2.1

0 .8 8 3

4 2 ,1 0 0

6800

18

0 .9 3

0 .3 5 5

3 8 ,2 0 0

7020

19

1 .5

0 .7 7 0

5 1 ,2 0 0

6850

23

1.81

0 .9 5 0

5 2 ,5 0 0

6780

24

1 .8 5

0 .8 1 0

4 3 ,7 0 0

6850

25

1 .6 3

0 .7 7 0

4 7 ,2 0 0

6800

26

1.61

0 .6 2 0

3 8 ,5 0 0

6800

27

1 .7 3

0 .7 5 0

4 3 ,3 0 0

6800

28

3 .8 9

1 .8 0 0

4 6 ,3 0 0

6800

29

3 .2 0

1 .0 8 0

3 3 ,8 0 0

6750

30

2 .0 9

1 .0 6 0

5 0 ,6 0 0

6820

31

3 .9 0

1 .9 6 0

5 0 ,3 0 0

6820

32

0 .9 1 6

0 .4 3 3

4 7 ,3 0 0

7000

33

2 .5

1.110

4 4 ,4 0 0

6850

63.

O

u

CD

Pi
C
O

<d
rH

O
£

<N
O

O
N

X
P
o
■H
+->
03

U
+->

CD

a>
o
c
o
u
£

M

00

"CS
0
01

2.0

Figure

TP

•H

Sodium-Aimnonia

N

for

00

Concentration

+->

•H

versus

u
<D

Absorbancy

<N
CO

12.

CD
CO

Solutions

at

-65

o

ft
iH

rH

o

(ifead) Xoueqjosqv

o

o

64.

M o la r Absorbancy Index for Sodium-Ammoniq Solutions
The absorbance readings of m etal-am m onia solutions appeared to be quite
reproducible at - 6 5 ° C , even after long periods of standing.
decom position, this is not the case at higher temperatures.

U nfortunately, due to
For this reason the values

reported for the molar absorbancy index are those at the lowest tem perature.

These data

for sodium are given in Table V II and plotted in Figure 12.
Figure 12 shows that w ithin experim ental error, a plot of A versus concentration
gives a straight line passing through the o rig in .

The absorbance for various cells are

corrected to 10- ^ cm. by dividing the measured absorbance by the c e ll length.
The average calculated molar absorbancy index is 4 2 ,1 5 0 1 . mole ^ cm.

\

O m itting runs 7 , 14, 15, and 19, for which the deviation is greater than four times the
average d e v ia tio n , the average molar absorbancy index is 4 5 ,2 6 8 1 . mole ^ cm.

^ o Since

the data appeared to fa ll on a straight lin e , a least squares treatment was performed to
find the best equation for this straight line in terms of it slope and in tercep t.

Using the

— 1
“1
data for a il sodium runs the slope was found to be 4 7 ,5 0 0 1. mole-1 cm.
; the intercept

to be - 6 . 2 x 10- ^ .

T heoretically the intercept would be zero if the solutions follow ed

Beer's law and a ll the data were w ithout error.
A least squares c a lc u la tio n , u tiliz in g only the sodium runs which gave a calculated
molar absorbancy index between 4 0 ,0 0 0 and 5 0 ,0 0 0 1. mole” ^ cm.

^ ogives a slope of

4 5 ,5 0 0 1. m ole- ^c m . a n d an intercept of - 1 . 5 x 10“ ^ . It appears that the true molar
absorbancy is quite close to 4 5 ,0 0 0 1. mole- ^ cm.
4 0 ,0 0 0

1. m oIe- ^cm.

^ estimated by Jolly (2 5 ).

\

which is slightly higher than the value

65.

Discussion of Possible Error
For a few of the early sodium runs the analysis for sodium was carried out by the
zin c uranyl ace ta te method.

H ow ever, this method is not sensitive enough to measure the

small amount of sodium in the absorption cell (ca. 10 ^ moles).

For the later runs, the

sodium was taken up in aqueous solution and analyzed w ith the aid of a Beckman Model B
Flame Photometer.

The prelim inary data indicated that the apparent concentration was

much higher in the absorption c ell than in the large tube (ca. 10-35 per cen t).
questions about the method of transfer from the large tube into the c e ll.

This raised

Aqueous sodium

hydroxide solutions were transferred using the same technique, and the results were again
higher in the absorption c ell indicating the p ic k-u p of sodium ions.

In order to test whether

the transfer was qu an titative the m ake-up vessel was modified for several runs by connecting
three absorption cells to one large tube.
ammonia solutions.
the cell lengths.

Absorption spectra were obtained for the m e ta l-

It was found that the observed absorbance values were in the ratio of

Thus it was concluded that the transfer technique was indeed q u a n tita tiv e .

The most probable explanation for the high results in the analyses of the cell contents
involves the extrem ely small amount of metal present (ca. 10“ ^ moles of sodium).

Small

traces of sodium salts introduced when taking up the aqueous solutions could contribute
nearly as much sodium as was in itia lly present.

It is also possible that small amounts

of sodium could be picked up by the action of the a lk a li solution on the glass.
The anomalously low molar absorbancy indices for several of the runs could be due to
decomposition to the am ide.

This would lower the measured absorbance but would not a ffe c t

the to tal sodium content and hence would give low calculated molar absorbancy indices.

66.

The decomposition rate is much higher at - 3 3 ° C than at -6 5 °C „

This Is shown by the

n o n -re p ro d u c ib ility of the molar absorbancy indices at the higher tem perature.

In many

cases a fte r the solutions were measured at approxim ately - 3 3 ° C they were cooled and the
absorbance again measured at -6 5 °C «

O fte n the absorbance at these low temperatures was

1 0-15 per cent lower than before the measurements were made at - 3 3 ° C .

H ow ever, it was

found for some solutions that upon standing overnight the absorbance returned nearly to the
orig in al v a lu e .

This can be explained by a faster rate of decomposition between the windows

than in the bulk of the solution.

Upon standing overnight the solution could then come to

e q u ilib riu m , returning the absorbance almost to its original v a lu e .
observed only for a few solutions w ith the m ajority continuing to

This phenomenon was
decompose andnever

returning to th e ir original absorbance values.
The high values of the molar absorbancy index could arise if some of the metal were
lost before analysis w h ile transferring from one container to another.
The Effect o f Temperature on the Sodium Spectra
Table V I II gives the data for the temperature dependence of the absorbance for a
ty p ica l sodium-ammonia solution.

These data are plotted in Figure 13.

It can be seen that

the w hole absorption curve appears to be shifted to lower energies as the temperature is
increased.

For N a N o . 17 and N a N o . 33 the area under the curves from 4 000 to 1 2 ,5 0 0 cm.- ^

was estimated for both high and low temperatures. In both cases, the

area

underthe curves at

the higher temperatures was approxim ately 5 per cent larger than the

area

at - 6 5 ° C .

e ffe c t of tem perature is not the same for a ll portions of the absorption curve.

The average

v a lu e , for eig ht solutions, o f the shift on the high energy side of the peak (c a . 8 ,5 0 0 c m " ')

The

67.

TABLE V IM
EFFECT OF TEMPERATURE O N A S O D IU M -A M M O N IA SPECTRUM

N a N o . 17
C (-6 7 ° C )= 2 .1 x 10“ 3 m /l
T = -3 7 °C
A m q x . = 0 .9 0 5 at 6 ,5 3 0 c m .“ ^
A /A M a x .

A mqx
A

= 0 -9 7 5 at 6 ,8 4 0 cm.

Cm" 1

A

1 2 ,5 0 0

0 .0 6 5

0 .0 7 2

0 .0 7 5

0 .0 7 7

11,110

0 .1 1 5

0 .1 2 5

0 .1 3 5

0 .1 4

10,000

0 .1 9

0.21

0 .2 3

0 .2 3 6

9 ,1 0 0

0 .3 1

0 .3 4

0 .3 8

0 .3 9

8 ,3 3 0

0 .4 8

0 .5 3

0 .5 8

0 .5 9 5

7 ,6 9 2

0 .6 7 5

0 .7 4 5

0 .7 7 5

0 .7 9 5

7 ,1 4 3

0 .8 4 5

0 .9 3 5

0 .9 4

0 .9 6 5

6 ,6 6 7

0 .9 0

0 .9 9 5

0 .9 7

0 .9 9

6 ,2 5 0

0 .8 9

0 .9 8 5

0.88

0 .9 0

5 ,8 8 0

0 .7 9 5

0.88

0 .6 9 5

0.71

5 ,5 5 0

0 .6 7

0 .7 4

0 .5 1

0 .5 2

5 ,2 6 0

0 .5 1

0 .5 6 5

0 .3 3

0 .3 3

5 ,0 0 0

0 .3 5

0 .3 9

0 .1 9

0 .1 9 5

4 ,7 6 0

0 .2 7 5

0 .3 0 5

0 .1 4 5

0 .1 5

4 ,5 4 0

0 .1 8

0.20

0 .0 7

0 .0 7 2

A /A M ax.

Figure

13.

Na

17

Temperature

No.

Dependence

of

the

Absorption

Curve

for

Sodium-Ammonia

Solutions

68.

ex e W v /v

69.

is - 9 cm.

1 per degree which is about o n e -h a lf the value reported by Eding (18) who

exam ined the same region.

As the temperature is increased, the peak shifts to lower

energies and also tends to fla tte n out.
of the peak is - 9 , 7 cm.

The average value for the shift in location

^ per degree, which is in good agreement w ith the value of

- 9 cm„“ l per degree described by Jortner (36) using the data of Blade and Hodgins (1 1 ).
The average value for the shift on the low energy side of the peak (ca. 5 ,5 0 0 c m .*^ ) is
- 1 2 . 5 cm ."^ per degree.

Fart II PO TASSIUM
The data for the absorption spectra of potassium-ammonia solutions are given in
Table IX . Figure 14 is a plot of A versus concentration for a ll potassium-ammonia solutions
exam ined .

W h ile sodium solutions appear to obey Beer’s law w ith respect to total dissolved

metal over the concentration range examined (ca. 3 x 10"4 to 4 x 10“ ^ M . ) , it can be
seen that potassium solutions deviate from Beer’s law at higher concentrations.
Figure 15 is a plo t of molar absorbancy index for both sodium and potassium
solutions as a function of concentration.

The data are taken from Tables V II and IX . It

can be seen that for potassium solutions an extrapolation of the absorbancy Index curve
to in fin ite dilu tio n yields a lim iting absorbancy index which is very close to the value of
4 5 , 0 0 0 I. mole"^ cm ."^ obtained for sodium solutions.
Figure 16 illustrates the change in position of the peak maximum as a function of
concentration for both sodium and potassium.

The data are taken from Tables V II and IX .

An increase in concentration of sodium up to 2 x 1CT3 M . , appears to shift the spectra
slig htly to lower energies. Above this concentration, which includes most of the potassium

70.

solutions, there is very little e ffe c t on the positions of the absorption maximum w ith
further increase in concentration.

The same types of experim ental errors are present when

working w ith potassium solutions as with sodium solutions.

The main source of error is the

uncertainty in obtaining precise concentrations of the solutions used.

Figure 17 shows a

graph of the peak position versus the square root of absorbance for both metals.

The

significance o f this graph w ill be discussed la te r.
The absorption spectra for solutions 4 , 5 , 6 , 1 4 , 1 5 , and 16 could not be measured
im m ediately after the solutions were prepared due to a breakdown of the spectrophotometer.
These solutions were stored in the large dewar a t dry ice temperatures for nearly a week
w h ile the spectrophotometer was being repaired.

Since the decomposition rate is faster in

more d ilu te solutions it would be expected that solutions 6 and 16 would give extrem ely low
molar absorbancy indices if much decomposition had taken place during this w aiting period.
This is not the case as the absorbancy indices for these solutions are higher than most of the
other potassium solutions.
Effect of Concentration of the Shape of the Potassium Spectra
Absorption d a ta , A /A y ^ ax , for potassium at d ifferent concentrations in liquid
ammonia at - 6 5 ° C taken from Table X are plotted in Figure 18.

It appears that a th re e -fo ld

change in concentration has no e ffe c t on the shape of the spectra.
E ffect of Temperature on Potassium Spectra
Absorption d a ta , A / A m

qx

/ ^or potassium solutions as a function of temperature are

given in Table X I and plotted In Figure 19.

The results indicate that the absorbing species

in potassium solutions may be more susceptible to temperature variations than the corresponding
species in sodium solutions.

The average v a lu e , from six solutions, for the shift of the high

71

TABLE IX
M O LA R ABSO RBA NCY IN D E X FOR P O T A S S IU M -A M M O N IA S O L U T IO N AT -6 5 ° C

„
Concentration
Solution

(M o la rity x 10^ )

A
( for b = 1 0 cm)

M olar
Absorbancy Index
( 1. mole

✓ cm~^)

Position
(cm- ^ )

1

2.8

1 .0 7

3 8 ,3 0 0

6 ,7 8 0

2

2.66

0 .6 7

2 5 ,2 0 0

6 ,7 5 0

3

2 .1 4

0.81

3 7 ,8 0 0

6 ,9 5 0

4

3 .7 5

1.20

3 2 ,0 0 0

6 ,7 8 0

5

2 .0 5

0 .7 9

3 8 ,5 0 0

6 ,8 7 0

6

1 .2 4

0 .5 3 5

4 8 /2 0 0

6 ,8 3 0

7

3 .2 7

0 .9 4

2 8 ,7 0 0

6 ,8 0 0

8

10*

2 .0

20,000

6 ,7 0 0

9

6.20

1 .6 3

2 6 ,3 0 0

6 ,8 3 0

10

4 .4 0

1 .3 2

3 0 ,0 0 0

6 ,7 8 0

11

2 .4 8

0.81

3 2 ,6 0 0

6 ,8 5 0

12

too conc.

—

13

decomposed

—

14

2 .4 3

0.86

3 5 ,4 0 0

6 ,8 5 0

15

2 .9 3

0 .8 3

2 8 ,3 0 0

6 ,9 0 0

16

1 .8 4

0.68

3 6 ,9 0 0

6 ,9 0 0

* Concentration from cell analysis.

-----------

—

—

—

(sjead) iCouBqjosqv

*H

Sh
0
a
CO

-65
for

u
0
p

Sodium

and

Potassium

Solutions

at

73.

o
s

CO

o
X
a
o

Concentration

0

rH ,

05

h

P

G

0

versus

•H

-P

G

(

wo

axow

* i)

xapui

iC o u e q jro s q v . i e i o H

ofead

Figure^15.

Molar

Absorbancy

O
O

Index

O

Figure

16.

Position

of

Peak

Maximum

for

-6 5

Solutions

for

Sodium

and

Potassium

versus

Concentration

i
O

o

<N

o
o

17.

X

Position of Peak
of Absorbance

a
S5

Maximum

W

O

o

00

CO

(I- “3

o
o

o
o

CO

CO

2[B0d; jo UOTq.TSO(i

CO

o
o

o
o
1ft

CO

CO

Figure

for

Solutions

of

Sodium

and

Potassium

versus

Square

Root

75*

I

76

energy side of the peak is -1 0 cm“ l per degree.

The average value for the shift of the

position of the maximum was about -11 .5 cm“ ^ per degree.
shift of the low energy side was - 1 5 .3 cm"^ per degree.

The average value for the

Corresponding shifts for sodium

were - 9 cm” l per degree, - 9 . 7 cm” ^ per degree and - 1 2 . 5 cm” ^ per degree, respectively.
Comparison of the Shapes of Sodium and Potassium Spectra in Liquid Ammonia
Absorption data for sodium and potassium in liquid ammonia taken from Tables V I
and X are plotted in Figure 2 0 .

The data is plotted as A / A ^ qx

.

It can be seen th a t,

w ith in experim ental error the spectra of sodium and potassium in liquid ammonia have
id e n tica l shapes. This would be expected if the same absorbing species were present an
both solutions.

77.

TABLE X
EFFECT OF C O N C E N T R A T IO N O N THE PO TASSIUM SPECTRUM AT -6 5 ° C

K No. 5

K N o . 15

c = 2 .0 5 x 10" 3 m /I
0 = 3 8 ,5 0 0 I . mole- ^cm” ^
A m qx = 0 .8 1 at 6 ,8 7 0 cm- ^

cm” ^

A

A /^ M a x .

K No. 9

c = 2 . 9 3 x 10“ 3 m /I
a = 2 8 ,3 0 0 I. mole"^cm"^
^M ax =

a* 6 ,9 0 0 cm- 1

A

A /A M a x .

c = 6 .2 x 10" 3 m /l
a = 2 6 ,5 0 0 I. mole- ^cm
A M ax = 1 .6 6 at 6 ,8 3 0 cm- l

A

A /A m ,

1 1 ,1 1 0

0 .1 1 5

0 .1 4

0 .1 4

0 .1 5

0 .2 4

0 .1 4

1 0 ,0 0 0

0 .2 2

0 .2 7

0 .2 3

0 .2 5

0.4 1

0 .2 5

9 ,1 0 0

0 .3 5

0 .4 3

0 .3 8

0 ,4 0

0 .6 8

0 .4 1

8 ,3 3 0

0 .5 2

0 .6 4

0 .5 5

0 .6 0

0 .9 7

0 .5 8

7 ,6 9 2

0 .6 9

0 .8 5

0 .7 4

0 .8 1

1 .3 3

0 .8 0

7 ,1 4 3

0 .8 0 5

0 .9 9

0 .9 0

0 ,9 8

1 .6 0

0 .9 7

6 ,6 6 7

0 .8 0

0 .9 9

0 .9 0

0 .9 8

1.61

0 .9 8

6 ,2 5 0

0 .7 2

0 .8 9

0 .8 2

0 .8 9

1 .4 9

0 .9 0

5 ,8 8 0

0 .5 6

0 .6 9

0 .6 4

0 .7 0

1 .1 5

0 .6 9

5 ,5 5 0

0 .3 9

0 .4 8

0 .4 3

0 .4 7

0 .7 6

0 .4 6

5 ,2 6 0

0 .2 6

0 .3 2

0 .2 6

0 .2 8

0 .4 8

0 .2 9

5 ,0 0 0

0 .1 5

0 .1 9

0 .1 2

0 .1 3

0 .2 5

0 .1 5

78.

1. 0-

0 . 8-

0.6
X
C3

S
<!

11,000

1 0 ,0 0 0

9,000

8,000

7,000

6,000

5,000

Cm“ l
Figure 18.

Effect of Concentration on Potassium Spectrum at -65°C

79.

TABLE X I
EFFECT OF TEMPERATURE O N P O T A S S IU M -A M M O N IA SPECTRUM

K No. 3
C ( - 7 Q°C) = 2 .1 4 x 10-3

T
Am q x .

= -7 0 ° C
=

T

0 .8 2 at 6 ,9 5 0 crrT^

Cm -1

vM a x .

=

-3 1 °C

=

0 .5 8 at 6 ,4 1 0 cm"^

A /A

A /A M ax,

M ax,

1 2 ,5 0 0

0 .0 6

0 .0 7 3

0 .0 3

0 .0 5

11,110

0.11

0 .1 4

0 .0 6

0.10

10,000

0.21

0 .2 6

0.12

0.20

9 ,1 0 0

0 .3 5

0 .4 3

0 .1 9

0 .3 3

8 ,3 3 0

0 .5 2

0 .6 4

0 .3 0

0 .5 2

7 ,6 9 2

0 .7 1

0 .8 7

0 .4 2

0 .7 3

7 ,1 4 3

0 .8 2

0 .9 9

0 .5 3

0 .9 2

6 ,6 6 7

0 .7 9

0 .9 6

0 .5 7

0 .9 9

6 ,2 5 0

0 .6 9

0 .8 4

0 .5 8

0 .9 9

5 ,8 8 0

0 .5 1

0.61

0 .5 4

0 .9 3

5 ,5 5 0

0 .3 5

0 .4 3

0 .4 6

0 .8 0

5 ,2 6 0

0.21

0 .2 6

0 .3 6

0.61

5 ,0 0 0

0 .1 7

0.21

0 .2 6

0 .4 5

4 ,7 6 0

0 .0 8

0.10

0.21

0 .3 6

4 ,5 4 0

0 .0 6

0 .0 7

0 .1 4

0 .2 4

Figure

19.

o

Temperature

K No

Dependence

o

00

o

•

xvyi

V/V
the

Absorption

I
s
o

of

Curve

Potassium-Ammonia

Solutions

80.

81.

Na No. 10

X

co

S

<
\

<
0 .4

0.0
11,000

10,000

F ig u re 20.

9 ,0 0 0

8,000

7 ,0 0 0

6 ,000

Comparison o f the Shapes o f Sodium
and Potassium S p ectra in Ammonia a t - 6 5 °C

5 ,0 0 0

82.

V I . D IS C U S S IO N OF RESULTS

Figure 20 shows that w ithin experim ental error the shape of the absorption curves
for sodium and potassium in ammonia are id e n tic a l.
is the same in both solutions.

This indicates that the absorbing species

It would seem reasonable to expect that the monomers in the

model of Becker, et a l . , having sodium ions as the nucleus would absorb at different energies
than monomers w ith potassium ions.

In both the ‘electron c a v ity 1 model and the ‘cluster model'

the only species which is common to both sodium and potassium solutions is the solvated
e le c tro n .

Therefore, it seems reasonable to assume that it is the solvated electron which is

absorbing light in these d ilu te solutions.
The absorption peak at 6800 cm"^ corresponds to an energy of approxim ately 0 .8 3
electron volts.

Jo lly (25) and Ogg (24) propose that this energy corresponds to the exc ita tio n

of the solvated electron to a conduction band.
sodium-ammonia solution

Ogg (5 3 ), working w ith extrem ely d ilu te

(10"4 - 10”^ M o lar) obtained a work function of about 0 . 8 volts.

H o w ever, Hasing (54) and Teal (55) for more concentrated solutions place the work function
at approxim ately 1 .6 volts.
Ogg and Eding (18) calculated the energy of the 2p state for an electron in a cavity
and obtained a slig h tly positive v a lu e .

They concluded that an electron in this state would

be unstable and c alled it a conduction electron.

How ever, it seems quite probable that

the energy values obtained by Ogg are in considerable error due to some of the approximations
made.

83

As mentioned in the theoretical section, Jortner (36) has recently published a
paper on the energy levels of electrons in ammonia.

He calculated the energies for

the Is and 2p states and obtained e xc e lle n t agreement w ith the experim ental results
of Blades and Hodgins when comparing the energy required for the 2pVls transition.
If the 'electron c a v ity 1 model gives the correct picture for the state of the electron
in metal ammonia solutions one might expect that even in some of the d ilu te solutions
exam ined there would be an appreciable number of paired electrons. Since the energy
d ifferen ce between paired and unpaired electrons is considered to be 0 . 2 electron volts
one might expect to observe an additional peak due to paired electrons.

Several

concentrated solutions ( 10" 2 j 0 10” ^ M o lar) were examined and fa iled to show any
ad d itio n al peaks.
There is the possibility that the paired electrons might absorb on the high
frequency side of the main absorption band as might be pictured below .

cm

84.

If this were the case then a plot of

A /A j^ ax versus crrT^ for solutions of

d iffe re n t concentrations would be expected to show a large change in absorption
on the

high energy side of the p e a k

Figure I I , for sodium solutions shows that w ithin

experim ental scatter there is no change in the shape of the spectrum as the concentration
is changed. Figure 18, for potassium solutions could possibly be interpreted as changing
shape when the concentration is changed.

How ever, it can be seen that no systematic

trend is observed, as some of the points A / A M ax
for the lower concentrations.
zero lin e .

for higher concentraions are below those

This is probably due to the d iffic u lties in obtaining an absolute

It is concluded for both sodium and potassium solutions that there is no change in

the shape of the spectra as the concentration is varie d .

This also indicates that the absorption

process involves only one species, nam ely, the solvated e le c tro n .
Figures 13 and 19 indicate that the absorption curves for both sodium and potassium
are strongly dependant upon the tem perature.

As the temperature is increased the curve

tends to fla tte n out as w e ll as being shifted to lower energies.

O ne can see that the low

energy side of the peak is being shifted more than the high energy side.

This is probably

observed because the true absorption curve is a combination of a broadening effect
superimposed upon a shift of the peak to lower energies.
For ordinary electronic transitions one would expect only a broadening e ffe c t as the
tem perature is increased, because the absorbing species would be in higher vib ratio n al and
rotational states due to an increase in thermal energy.

H ow ever, if in a d d itio n , the whole

absorption curve is shifted to lower energies then it would be expected that the apparent

85.

high energy shift would be slightly less than the low energy shift.

This is illustrated

b elo w .
T2 greater than Tj

-

1.0

A
iax.

cm

‘M a x .

cm

cm

N o Shift

Shift

According to the cavity m odel, paired electrons are more stable at lower
temperatures than at higher temperatures.

It would be expected that as the temperature

is increased the paired electrons would dissociate into unpaired electrons.

Even if paired

electrons were to absorb at the same energy as unpaired ^l^ctrons, one would expect a
temperature dependence of the area under the absorption curve unless the molar absorbancy
index for pairs were exa c tly tw ice that for unpaired electrons,

although it was found that

the area did increase by approxim ately fiv e per cent for a thirty degree temperature change it
is believed this change was probably w ithin the range of experim ental scatter and not a
true change in a re a .
The experim ental shift in the position of the absorption peak w ith temperature in m e ta lammonia solutions is much greater than shifts observed in other systems over the same temperature
range.

The best explaination to date for this large shift is offered by Jortner (36) who

derives equations for this shift in terms of the temperature dependance of the static d ie le c tric
constant and the temperature dependence of the cavity radius.

He obtains a positive

86.

tem perature c o efficie n t for the cavity radius and obtains fa ir agreement w ith the
experim ental results of Blade and Hodgins.

The results for the temperature dependence for

the m etal-am m onia solutions obtained in this research are also in good agreement with
these workers.
In looking at Figure 12 which is a graph of A versus

C for sodium-ammonia

solutions it appears that sodium obeys Beer’s law over the concentration range covered.
W ithin experim ental scatter the points fa ll on a straight line which passes through ze ro .
The reason for the high value of the absorbancy index obtained from the least squares
treatm ent can be understood when one notes that the least-squares intercept is below zero .
This would give the line a larger slope then if the intercept were zero. Figure 14 shows
that potassium does not obey Beer's law as the points do not form a straight lin e .

H ow ever,

there might be some uncertainty in the data for potassium solutions due to decomposition.
This contradicts the general opinion that the rate of decomposition is higher in d ilu te
solutions than for more concentrated solutions.

If the absorbancy indices for sodium and

potassium are not e q u a l, as it appears, then the results are not in agreement w ith the
g en erally accepted statement that the spectra of these metals are id e n tic a l, for the absorbancy
indices would have to be equal as w ell as the line shape in order for the spectra to be
id e n tic a l.
O ne might explain why the absorbancy indices for potassium are less than for
sodium by considering the currently popular models for these solutions. There appear to be
four ways for removing electrons from the state in which they absorb light in the near
in fra re d .

87.

(1)

Electron pair formations

2e

(2)

(s o l.) = e 2 (s o l.)

Monomer formations

e'(s o l„ )

+ M + (s o l.) = M

Where M is the monomer of Becker, et a l.
(3)

lo n -p a ir formation analogous to electrolytes in solvents of low
d ie le c tric constant
e'( s o l.) + M + (s o l.)

(4)

M + (s o l.)’ e (s o l.)

Dimer formation

2 M + (s o l.) + 2 e " (s o l.) = M 2 (s o l.)
(1)

The electron might be removed by electron pair formation which would be

favored at high concentrations.

This would explain the decrease in molar susceptibility

observed in these solutions at higher concentrations and could be the reason for the decrease
in molar absorbancy index for potassium.
O ne would probably get a new peak if

~ were present.

H ow ever, the peak

might be out of range since the energy d iffe re n ce , 0 . 2 e . v . , is the o v e r-a ll free energy
change and not necessarily the energy level shift.

If this were the case, then sodium and

potassium should e xh ib it the same behavior because the e 2 = would be re la tiv e ly independent
o f the c a tio n .

Since these metals do show different behavior, the presence of an appreciable

concentration of e2 = centers independent of the metal seems to be ruled out.

88.

(2 ) Monomers could result in removal of electrons.

H ow ever, it is inconceivable

th at the spectrum of a monomer of the type postulated by Becker, et a I . ,
id e n tica l to that of a solvated, but dissociated, electron.

would be

Therefore, there probably are

no monomers in sodium-ammonia solutions at these concentrations (5 7 ).

Since dim erization

seems to adequately explain the results for potassium, there is no need to

introduce

monomers in this system e ith e r, although their presence cannot be ruled ou t.
(3) Having dispensed w ith monomers, it is evident from electrostatic consideration
alone that some type of ion-pairing must occur.

It is lik e ly that this involves only a

contact of the solvated cation with the cavity containing the electron.

Calculations

u tiliz in g the Bjerrum (56) theory of ion association indicate that the most d ilu te sodium
solution would be com pletely dissociated, w hile in the most concentrated solution studied,
the ion-pairs would be only about 70 per cent dissociated.

Due to the re la tiv e ly large

separation of the ions in the io n -p a ir, the energy levels of the electron
slightly shifted.

may be only

The e ffe c t of concentration on the position of the peak shown in

Figure 16 indicates a shift to slightly lower energies as the concentration is increased.
It appears that the energy shift depends upon the ionic strength of the solution.

If one

assumes the absorbance to be a measure of the free Ion concentration, and hence of the
ionic strength the results for sodium and potassium can be plotted together.

Figure 17 shows

a graph of the peak position versus the square root of the absorbance for both metals.
W h ile

the scatter is large, it is not unlikely that this is a straight lin e . Thus it appears

th at both the sodium and potassium solutions contain appreciable concentrations of ion
pairs, but that the e ffe c t on the spectra is slight.

89

(4)

M agnetic studies make it clear that the electrons do not remain unpaired

at a ll concentrations.

It is lik e ly that the pairing mechanism also causes the decrease in

molar absorbancy index observed for potassium.
the dimer suggested by Becker, et

Since we have ruled out

a l . , seems to be adequate.

~ centers,

Assuming the lim iting molar

absorbancy index for potassium to be the same as that for sodium, an estimate of the pairing
constant can be made.

This gives a value of 6 ± 2 x 10” ^ for the dissociation constant

compared w ith the value of 2 x 10“ ^ obtained in this concentration range by extrapolation
of the data of Hutchison (1 2 ). Considering the errors inherent in both methods, this
agreement is satisfactory.
The data presented for sodium in this thesis would indicate n egligible dimer
formation at these concentrations.

The data of Hutchison for sodium are much less

satisfactory in internal precision than for potassium.
that a t low

concentrations and temperatures, sodium would result in more pairing than

for potassium.
in order.

They seem to in d ic a te , how ever,

A repetition of the measurements of Hutchison for sodium certainly seems

In this Same connection, Evers and Frank (45) conclude on the basis of

conductance data that sodium exhibits jess pairing than potassium in agreement with
our observations.

90.

V II.

PRESENTATIO N O F A N EW M O D E L FOR M E T A L -A M IN E S O L U T IO N S
IN T R O D U C T IO N

It has often been stated that the a lk a li metals dissolve in liquid ammonia and
e x h ib it an absorption maximum at 1 .5 microns which is independent of the dissolved
m e ta l.

The absorption band is ascribed to an electron trapped in a cavity formed by the

p o la riza tio n of the solvent.

The results of the present research in d ic a te , that w h ile the

absorption process is essentially the same for sodium and potassium in liquid ammonia,
the molar absorbancy indicies for sodium and potassium are not id e n tic a l.

This can be

explained by removal of electrons from the cavities by mass action effects.
Absorption measurements of the a lk a li metals in methylamine show a visible
peak which is found to vary from metal to m etal.

Lithium and calcium in methylamine

both have a peak in the infrared, but shifted to slightly higher energies than the peak
observed in ammonia.
Figure 2 1 .

The observed data are summarized in Table X II and show in

Blade and Hodgin's attribute these visible peaks in methylamine to arise

from electrons in a lip h a tic traps.

An a lip h a tic trap involves cavities u tiliz in g the

protons on the methyl group rather than the amine group.

Paramagnetic resonance

measurements by Fowles, et a l . , showed that metal solutions which exhibited only the
visible peak were diam agnetic.

They concluded that the principle species in these

solutions were paired electrons.

This does not explain why the position of the peak depends

upon the m etal.

91

TABLE X II
ABSO RPTIO N PEAKS FOR METALS IN M E T H Y L A M IN E A N D
FOR G AS PHASE MOLECULE PEAKS

Gas phase
M olecu le peak
M -M

M ethylam ine
Peak
M e ta l

Li

)

cm” ^

7 ,8 0 0
1 4 .5 0 0

Na

1 5 ,0 0 0

K

1 2 ,2 0 0
1 5 .5 0 0

°c m - ^ (56)

1 4 ,0 2 0
( l f l r - H * p 2 0 ,3 9 8
1 4 ,6 5 9
(Ir ^ k p
2 0 ,3 0 2
1 1 ,6 7 0
V 1 5 ,3 6 9

(1

22,000
1
Cs

1 0 ,0 0 0

extended bands 8 ,0 0 0 to

11,000
Ca

7 ,8 0 0

unstable

92.

0 .8 —

Ca ( 1 )

0.0
Li (11, 20)

A

0.0
(11)

Absorbance

Na

0. 0

K (11)

0.7

V

o.o

18,000

Cs
This research

16,000

14,000

12,000

10,000

8,000

6,000

Cm-1
Figure 21.

Absorption Curves for Metals in Methylamine

93.

If the visible peak in methylamine were due solely to electron pairs it is d iffic u lt
to understand why a visible peak Is not observed in solutions of metals in liquid ammonia,
at least at higher concentrations, where according to the data of Hutchison significant
p airing must occur. This would seem to indicate that the species giving rise to the visible
absorption in methylamine are not present in ammonia solutions.
When ammonia is added to a solution of sodium in methylamine it is found that
the visible peak diminishes in intensity with a new peak appearing in the infrared.

At

high percentages of ammonia only the infrared peak is present.
M O DEL
The author would lik e to present the follow ing model based p rin c ip a lly upon
the op tical and magnetic data found in the literature for metal solutions in ammonia and
m ethylam ine.

In metal ammonia solutions the metal atom loses its electrons to the solvent

as a result of the high ion-solvent interaction.

In extremely d ilu te solutions the major

species present are the solvated metal ion and the solvated electron.
this situation is eleg an tly described by Jortner (3 6 ).
io n -p a ir formation takes
am m onia.
et a l . ,

The energetics of

As the concentration is increased,

place just as would be expected for any other e le c tro ly te in

The io n -p a ir is not the equivalent of the monomer in the theory of Becker,

but arises merely from electrostate interaction of the solvated ion w ith the electrons

in its c av ity as represented by

^ (sol.) +

6 (soL)

(sol.)

(sol.)

94.

A t moderate concentrations the electrons must be paired to explain magnetic data
and the dimer of the Becker, et

a l . , model seems to be adequate.

The dimer might

be represented as

T
j

MS S M j

5

s

T
or

S

{ MS M 3

s

The electrons are assumed to be in molecular orbitals involving the protons of the
solvent.
C alcium behaves in the same manner in methylamine as in ammonia.

This is

probably due to the high solvating power of the ion which has a large surface-charge
density.

Presumably the other a lk a lin e earth metals would also behave in the same

manner.

Lithium in methylamine also has a tendency to y ie ld solvated electrons

as evidenced by its infrared peak.

How ever, lithium shows in addition an absorption

peak in the visible which the author postulates to be due to the presence of U 2 molecules
which have only a very weak interaction with the solvent.
Sodium, potassium and cesium having lower surface charge-densities are
e le c tro s ta tica lly too weak to form strongly solvated cations in a solvent w ith as low a
d ie le c tric constant as m ethylam ine.

The author postulates that the principle species in

these solutions are N a 2 , K2 , and Cs2 molecules, respectively.
Table X II shows that there is good agreement between the absorption peaks of
these metals in methylamine and the gas phase molecule*

95,

The follow ing e q u ilib ria are proposed for the species present in m etal-am ine
solvents.
I
(solvent)
1 /2 M

-

M

=

2

M +(so() + e (so|)

=

_

3

jM +(so|) • e ‘(so| j = ' / 2 <M Sn)2

The above e q u ilib ria w ill be described in more detail for each individual metal
for which data are a v a ila b le .
Lithium in methylamine has two absorption peaks, one at 1 4 ,5 0 0 cm’ ^ and one
at 7800 cm ^.

In ammonicy lithium has one absorption peak at 6800 c m " !.

The data of Blade and Hodgins (17) indicate that the peak at 1 4 ,5 0 0 cm” ^ appears
only in extrem ely d ilu te solutions.

This contradicts the data of Shatz (20) who

q u a n tita tiv e ly measured the molar absorbancy index of the 1 4 ,5 0 0 peak as a function
of concentration at - 7 5 ° C .

Shatz found that in dilute solutions Beer's law was obeyed,

w h ile at higher concentrations the molar absorbancy index increased as the concentrations
Increased.

It is believed that these data are more re lia b le and the model for lithium

w ill be based upon these d a ta .
the p rin c ip le

For lithium and a ll of the a lk a li metals in liquid ammonia

species are presumed to be the solvated metal ion, the solvated e le c tro n ,

and the solvated metal dim er.

In dilute metal ammonia solutions the predominant species

are the solvated ion and e lectro n .
When lithium is dissolved in methylamine the principle species are presumed to be
the metal m olecule, the solvated metal ion and solvated electro n .

In d ilu te solutions the

equilibrium favors dissociation w h ile the more concentrated solutions favor molecule
form ation.

This would explain the increase in molar absorbancy index obtained by Shatz

for the visible peak at 1 4 ,5 0 0 cm-1 .

This peak is attributed to an absorption of light

96.

by the L12 m olecule.
show that the 1^

Data from the appendix in Herzberg's 'Diatom ic M o le cu le s 1 (56)

Ar

1 4 ,2 0 0 cm ^.

transition in the gaseous state gives an absorption line

Thus it can be seen that in methylamine the absorption occurs at

approxim ately the same energy as in the gaseous state. The peak at 7 ,8 0 0 cm“ l is
thought to be due to the absorption of an electron in a methylamine c a v ity .
The l £. u

transition for gaseous N a 2 molecules occurs at 1 4 ,6 5 9 cm "^ .

Solutions of sodium in methylamine have a single absorption peak at 1 5 ,0 0 0 cm” ^ which
is in the same energy range found for sodium molecules in the gaseous state.

Solutions

of sodium in liquid ammonia have an absorption peak at 6 ,8 0 0 cm"^ which is attributed
to the solvated electron.
Sodium does not have as high surface-charge density as lithium and thus the
solvating power of the sodium ions is not strong enough to overcome the dissociation
energy of the m olecule.
diatom ic molecules.

Presumably the main species present in these solutions are the

O ne cannot te ll from the magnetic and optical data a v a ila b le

whether solvated dimers are also present.

Since both of these species are diam agnetic the

best way to determine if both species are present would be a q u antitative study of the
absorption spectra to see if there were any derivation from Beer's law .

A t present these

data are not a v a ila b le in the lite ra tu re .
Potassium
The l £ u

+

•*"" l£ g

+

transition for the gaseous l<2 m olecule occurs at 1 1 ,6 2 0 cm

Solutions of potassium in methylamine have an absorption peak at 1 2 ,2 0 0 crrT^.

This

peak can probably be assigned to the K2 m olecule because an absorption occurs in the

_I

97.

same energy range for the gaseous m olecule.

There is an additional peak reported

■l

at 1 5 ,5 0 0 cm

and in extrem ely dilute solutions a third peak appears at 2 2 ,0 0 0 cm- 1 .

This model cannot adequately explain the appearance of the two higher energy peaks.
For the gaseous K -K molecule

1 ^ -'rl^.g ^an d (*£g+ ) ^ r l £ g + transitions are observed

a t 1 5 ,3 6 9 cm"^ and 2 2 ,9 5 4 cm’ \

respectively.

How ever, it is rather doubtful that these

transitions would be observed for potassium in methylamine since similar transitions
should also be observed for lithium at 2 0 ,4 0 0 c m " ^ sodium at 2 0 ,3 0 0 cm \

and cesium

a t 1 3 ,0 0 0 cm 1 and 1 6 ,0 0 0 cm "^.
M a g n e tic evidence by Fowles, et a l . ,

(26) indicates that solutions of potassium

in m ethylam ine are slightly param agnetic. The only species which could be present and
have an absorption band in this region are potassium atoms or the monomer of Becker/ et a l . .
H o w ever, potassium atoms would be expected to absorb at approxim ately 1 3 ,0 0 0 cm" 1 •
The absorption band for the monomer would be expected to lie at lower energies then for the
potassium atoms, and it is questionable whether the monomer would absorb at these energies.
In measurements on faded m etal-amm onia solutions, Eding (18) was able to measure
an absorption peak forthe>metal amides at approximately 2 9 ,0 0 0 cm 1 .

It might be

possible that decomposition to the methylamide had occurred in the extrem ely d ilu te
_ 1

solutions required to detect the absorption band at 2 2 ,0 0 0 cm
was due to the absorption of methylamide ion.

/ ( l l ) and that this band

In the present research several q u a lita tiv e

runs were preformed to determine the position of the absorption peak for cesium in
m ethylam ine.

O ne of these solutions in a I mm absorption cell was allow ed to decompose

and the absorbance was measured from 5 ,0 0 0 cm” 1 to 2 5 ,0 0 0 cm "1 .

There was not trace

o f a peak a t 2 2 , 0 0 0 cm" 1 even though the methylamide concentration was high.

98.

C alciu m
C alcium would not be expected to form a stable Ca2 m olecule.

It is observed

for solutions of calcium in methylamine that only a single absorption maximum at
approxim ately 7 , 8 0 0 cm 1 appears.
solvated by m ethylam ine.

This absorption band is probably due to electrons

The peak occurs essentially at the same energy as the near

infrared peak for lithium in m ethylam ine.

It would be expected that the calcium metal

would be com pletely ionized since it has a surface-charge density which is larger than
lith iu m , and would have a high solvating power.
might lose only one of its valence electrons.

There is the possibility that calcium

This could be detected in a quantitative

study o f the molar absorbancy index w ith concentration.

Although the data are not

a v a ila b le it would be expected that the other a lk a lin e earth metals would behave like
calcium in both ammonia and methylamine
Cesium
In order to provide a further test of the above theory three runs were made w ith

+
cesium in m ethylam ine.

According to Herzberg (56) the 1^

+
g

trans' i''on

gives an extended series of absorption bands between 8 ,0 0 0 cm 1 and 1 1 ,0 0 0 cm- 1 .
The three cesium -m ethylam ine solutions a ll possessed anabsorption peak at 1 0 ,0 0 0 cm-1
which is w ithin the

r e g io n

given by H erzberg.

Deviations from Model
W h ile the above model q u a lita tiv e ly explains many of the observed facts for
metal solutions in ammonia and methylamine it does not explain a ll of them .

There

is a t present no satisfactory explanation for the appearance of the higher energy absorption
bands at 1 5 ,5 0 0 cm" 1 and 2 2 , 0 0 0 cm" 1 for solutions of potassium in m ethylam ine.

99.

M a g n e tic studies by Fowles, et a t . ,

indicate that these solutions are w eakly para­

m agnetic which might be interpreted as possibly due to a metal atom or monomer.

Solutions

of potassium in ethylam ine (11) have a single absorption peak at 15/ 000 cm- ^ w ith no
peak a t 1 2 ,2 0 0 cm ^ .

M agnetic studies have not been reported for these solutions.

It

is fe lt that magnetic studies would be valuable for determining the species present.
Symons (57) has recently observed that additions of large quantities of sodium salt
to d ilu te solutions of sodium-ammonia solutions results in the appearance of a new band
at 1 2 ,5 0 0 cm“ ^ .

He postulates this band might be due to the monomer of Becker, et a l . ,

Symons (58) plans to publish more details about this e ffe c t in the near future.
The above model does not mention the possibility of a monomer for thesem etalamine solutions because the observed optical and magnetic data can be explained in terms
of other species.

If a monomer were to be added to the above model then the model would

become id en tical to the model of Becker, et a L , for m etal-am monia solutions.
The above model does not adequately describe the position of the metal absorption
peaks in some of the higher amines like ethylam ine and eth ylen ed iam in e.

100.

S UM M A RY

The absorption spectra for d ilu te solutions of sodium and potassium in liquid
ammonia were obtained as a function of concentration and tem perature.

The line shapes

for the absorption curves were found to be id e n tic a l, and independant of concentration
for both sodium and potassium solutions.
curve is shifted to lower energies.

As the temperature is increased the absorption

These results indicate that the absorption process in

the near infrared Involves the excitation of an electron In cavity to a higher energy state
and is the same for d ilu te solutions of both metals.
Sodium-ammonia solutions obeyed Beer's law , with respect to total m etal, over
the concentration range covered (ca 3 x 10“4 to 4 x 10- 2 M ) w hile potassium solutions
(ca 1 x 10“ 2 to 1 x 1CT2 M ) gave a negative deviation from Beer's law .
is exp lained essentially in terms of the current models for these solutions.

This deviation
The model

involves io n -p a ir formation in d ilu te solutions for both metals. In the more concentrated
solutions examined dimer formation is presumed to occur for potassium but not for sodium.
This is in agreement w ith the results of paramagnetic resonance studies for potassium
solution but not for sodium solutions.

The calculated molar absorbancy index of sodium

solutions and of in fin ite ly d ilu te potassium solutions was found to be 4 5 ,0 0 0 ± 2 ,0 0 0
i1,

i mole

l

cm" I
A general model is presented to explain the magnetic and optical properties of

solutions of metals In both methylamine and ammonia which cannot be com pletely explained

101.

in terms of the existing models. This model consists of competing e q u ilib ria between
diatom ic m olecules, solvated electrons, solvated metal ions, and solvated dimers.
The direction in which these e q u ilib ria shift depends upon the physical properties of
both the metal and the solvent. The absorption bands in the visible region are a ll
attributed to diatom ic molecules w hile the absorption bands in the near infrared are
attributed to solvated electrons.

The solvated dimer is presumed not to absorb in these

regions.
Q u a lita tiv e experiments on the absorption spectrum of cesium-methylamine
solutions showed the peak to be in the region predicted.

How ever, the existance of

two ad d itio n al peaks for potassium-methylamine solutions cannot be adequately explained
by this m odel.

102.

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