ELECTROLYTES .\SD ELECTRO('HEMIC.\L
('OUPLES
505
age latent heat of evaporation over this temperature range of 6OG8 g.-cal., 'mole. This is in good agreement with equation 2 of Trautz and Gern-ig.
We wish t o thank Dr. R . F. Hudson for help with the apparatus employed in rhis study-. XEFERESCES : 1 I BRISER,lf,E . , ASD P Y L K O F FZ,. : J:chim.
phys. 10, 610 (1912). RAE,W.S.: Physicochemical X e t h o d s , T-01. 1, p. 564. lIethuen and Company, L t d . , London (1943). , 3 i SCOTT,R . B.: J. Research S a t l . Bur. Standards 25,459 (1940). (4) TRAL-TZ. AI,, . ~ S D GERWIG,IT.: Z . anorg. Chcm. 134,409 (1924). ('7; REILLY,R.,
ASD
SOME FUXDA;\IESTAL STUDIES OF ELECTROLYTES ;1SD ELECTROCHEhlIC,IL COUPLES OYER THE TEXIPERATVRE R A S G E 25°C. TO -75°C. -1.B. G-IRRETT, JhT WELSH, SAMUEL WOODRITFF, ROBERT COOPER, ASD J O H S HEIICS Department of Chemisiry, The Ohio S t a t e L-niueisity, Columbus 10, Ohio
R e w i r e d July l Y , 1948
The purpose of this paper is to give a preview of the results 1.i-hichhave been obtained to date in our invest'igations of the characterist'ics of electrolytes and These electrochemical couples over the temperature range of 25°C. to -75°C. investigations include: ( 1 ) a study of the physical properties of solvent-electroIyte systems that' remain liquid over this temperature range; (2) the measurement' of voltage and flash current, of several hundred electrochemical couples; (31 the measurement of the polarization of numerous metals and oxidizing agents; and (4)the measurement of the rate of solution of zinc, cadmium, and magnesium in acid solutions. One of t,he main objectives of this ~ v o r k is l to describe the parameters involved ;n developing a primary battery that, i d 1 operate at -T3"C. SOLTEST-ELECTROLYTE SYSTEMS
The systems chosen for study here vere those which would have ( 1 ) a lo~v freezing point, (2) a reasonably high clielect'ric constant, and electrical conducI ii-ky, ( 3 ) lon- volatility at ordinary temperature, and ( 4 ) good solvent characieristics. These qualifications practically restrict the choice of the solvent to :t -rater system. The systems given in tahle 1 are found to offer good possibilities. This ~ o r was k initiated and is supported by t h e Battery Branch of t h e Squier Laboriiof t h e Signal Corps of the P n i t e t l States Army :it Fort lIonniout11, SCVI. Jersey. The r r s u l t s will be published in niow detail later. This pnper w~ispresented a t t h e Sy-niposium o n Gnlvanic Cells a i d Batteries which W:LS Lcld under the auspices of the Division of Physic:il and Inorganic Chemistry a t the 113th J l r e t i n g of the .iniericon Chemical Society, C'hicago, Illinois, dpril, 194'3. tiirg
506
s.
A, B. GARRETT, J. WELSH,
WOODRUFF, R. COOPER AND J. HEIKS
The density, viscosity, and electrical conductivity of several of these systems are given in figures 1, 2 , and 3 inclusive. Of these systems, eutectic HCl-HgO probably has the best all-round characteristics and the systems in group 1 (table 1) have better all-round characteristics than those of group 2 . To be
-
__
- I _
"C
-86
Eutectic HCI-H,O 42% H B F r H z O LiCI-CH3NH2 HCl-H20 LiBr-MgBr,H?O LiBr-CaBrZ-H?O
ca. -75
+
,
Below -75 ca. -59 Below -60
-
~
~I1
HCIO,-H20 H2S04-H20 I SHI-H?O I Fe2C15-H20 CH30H-H20-HC1
1
--
"C
-GO -75 c a . -80 CQ -55 ca. -8.5 CQ
1,340
> 1.280 4
0.820 0.800 0.780
20
IO
-10 -20 -30 -40 -50 -60 -70 T E M P E R A T U R E "C.
0
FIG.1. Ileiisity of several of the systems given in table 1. Curve 1, 43 per cent fluoboric acid; curve 2 , 40.5 per writ pcichloric acid; curve 3 . 36 per ccnt sulfuric acid; curve 4, 24.S per cent hydrochloric ucid; curve 5 , 11 per ccnt lithium chloride43 per cent methylammonium chloridc-56 1)c.r c r n t water; curve 6. 100 pcr ccrit methanol.
noted in the data shoun in figures 1, 2, and 3 are: (1) the viscosity increases about twentyfold over this same temperature range with u sharp increase in the region of the freezing point of the system; (2) the electrical conductivity decreases about twentyfold for most of these system.: over the temperature range of +25"C. to -75°C. Fdr purposes of comparison the electrical conductivity of 1 normal potassium chloride is shorn1 on figure 3 ; it is of interest to note that the conductivity of
ELECTROLYTES AXD ELECTROCHEMICAL COUPLES
507
TEMPERATURE "C. FIG.2 . Ilelative viscosity of several of t h e systems given in table 1. Curve 1, 36 per cent sulfuric acid; curve 2 , 11 per cent lithium chloride-33 per cent methylammonium chloride46 per cent water; curve 3, 43 per cent fluoboric acid; curve 4, 40.8 per cent perchloric acid: curve 5 , 24.8 per cent hydrochloric acid; curve 6, 100 per cent methanol
3000-
o-
20
-10 -20 -30 -40 -50 -60 -70 T E M P E R A T U R E "C.
10 0
FIG 3. Specific conductivity of several of t h e systems given in table 1. C'urve 1, 36 per x n t sulfuric acid; curve 2 , 24 8 per cent hvdrochloric acid; curve 3, 43 per cent fluoboric icid; curve -1, 40.5 per cent perchloric acid; curve 5 , 11 per cent lithium chloride-33 per .cnt methvl:immonium chloride46 pcr cent water; curve 6. 1 S potassium chloride.
508
.I, B. G-IHRETT, J. TVELSH, S. \\.OODIITFF,
K . COOPEII AXD J. HEIKS
thebe >y>t.emsstudied a t -75°C. is of the order of magnitude of that of 1 normal potassium chloride. -4system with a relatively high electrical conductivity a t low temperatures is deemed necessary for primary cells to avoid high internal resistance a t Ion- temperatures.
hole a t b o t t o m
for m e t a l l e a d
weighed quantity of d e p o l a r i z e r m i x e d with graphite
( o r e a o f bottom surface .:13 cm?)
p a p e r separator surface
= 1.3cm?)
FIG.4. h p p a r a t u s ior t h e study of depolarizers
6-
20
40
60
80
100
120
140
160
180
200 220
TIME ( m n . )
FIG.5 Discharge curve for cell
C'urve 1,20'C.; curve 2 , -40°C.; curve 3 , -60°C.; curvc 4 , - i 3 " C . cnpacity s l i o ~ ~byn arrows. ELECTKOC€IIXIC.iL
Per cent cli theoretic:il
COUI'LIh
Over 200 electrochemical couples that seemed t o have possibilitie> for lon temperature operation have been investigated n-ith reference to (1 i stability, ( 2 ) open-circuit voltage, and ( 3 ) flash current over the temperature range of +2.3"C. to -73°C. Here, too, n-e noted a drop of from fivefold to twentyfold in the flash current of the cells as the temperaturc \ \ a s ieduced t o -75°C. Many of the couples gave no appreciable current at - T5"C. The next steps in this program n-ere t o study the polarizability u i n series oi anodic and cathodic materials that may be used in electrochemical couples and which :ippear from the above study to be the more promising materials.
ELECTROLYTES A S D ELECTROCHEMICAL COGPLES
609
A . oxidizing agents-cathode; "depolarizers" The efficiency of a number of oxidizing agents (commonly and- incorrectly called depolarizers) was studied by means of the apparatus shoum in figure 4. C:idmiiini \\-aschosen as the standard reference electrode because it shoived good
12
8 -
I
100
200
400
300
TIME i m i n )
FIG.6. Discharge curve for cell
Curve 1,20'C.;curve 2, -40°C.; curve 3, -6O"C2.; :apacity shon t i by arrows.
g
curve 4, -73°C.
-- --
~ ~ o - - ; 2 0 . ! ~ ! 2 ! : : 2
- __ e o z c ~ 1 2 o o-~2r
TIME
FIG.7 . Discharge curve for cell
Per cent of theoretical
bin)
I I
437, AgCI on C'd HBFa carbon
h r v e 1, -4O'C.; curve 2, -60°C.
Per cent of theoretical capacity shown.
11-round characteristics :IS anodic material. The cell was so constructed hat a weighed amount of the oxidizing agent could be placed in the electrode in uch a manner that a known area of the active material TI-as exposed. The cells -ere then discharged through a 200-ohm resistance at several temperatures and i e voltage characteristics us. time were determined; some of them :ire sho\rn 1 figures 5 to 9.
510
-4. B. GARRETT, J. WELSH, S. WOODRUFF, R. COOPER k h D J. HEWS
i
I.O
100% cop.
.e I
I
400
300
200
I00
T I M E (min.)
FIG.8. Discharge curve for cell
Curve 1,20"C.; curve 2, -40°C.; capacity shown by arrows.
.
20
curve 3, -60°C.; curve 4, -73°C.
40
60
80
100
I20
Is0
Per cent [of theoretical
160
TIME (rninJ
FIG.9. Discharge curve for cell
Curve 1,20°C.; curve 2 , -4c1"C.; curve 3, -60°C.;curve 4, -73°C. cal capacity shown by arrows.
Per cent of theoreti-
ELEC'TII0LYTI:S
1 S D ELC('TAOCHEVIC.1L
511
COT-PLES
A study of the polarization of the metal anodes as made with the apparatu!: shown in figure 10, which made it p o d d e t o pass a constant amount of ciirrent
reference electrode (MnOe + gtaphite) u e p e n i n g i n g l a s s tube
FIG.10. Anodic polarization apparatus
I
,5
lP5 ' O9 I
I
14
/7
I
6
I
100
200
300
400
500
600
700
800
TIME lsec)
FIG 11. Polarization curves i n 25 per cent hydrochloric ncid a t -73°C. Curves 1, 11, and 21, Lint a t 6, 4, and 2 ma./'cm.2; curve 2, zinc amalgam a t 6 ms./cm.2; curves 3, 13, and 23, cadmium a t 6, 4, and 2 ma./cm.2; curve 4, tin a t 6 ma./cni.*; curves 5 and 25, lead a t 6 a n d 2 ma icm ? ; ciirvc 6 , coppcr a t 6 ma./cni.?; curves 7 , 17, and 27, iron at 6, 4, and 2 ma cm *
through a. cell composed of a reference electrode containing manganese dioxide and an electrode whose polarization was being studied. A third electrode composed of manganese d i o d e was inserted t o determine the potential differ-
512
A. B. CLQRRETT, J. WELSH, S. WOODRUFF, R. COOPER A S D J . HEIKS
1.2
-
-
10-
w
0
> 5 .0c
e
-E
.6-
0
0
.4
-
.2
-
x
I
I
2
4
I 6
FIG. 100 sec.
taken a t
I .2
1
.o
. 1
c
> e
.a c e
5-2 .6 0
0
.4
.2
BO
60
40
10
0
-20
-40
-60
-80
-00
T'F
FIG. 13. Change of polarization with temperature a t a current density of 6 ma. cm.2 Reading taken at 100 sec. Electrolyte, 43 per cent fluoboric acid. 0 , cadmium; X, lead.
ence due t o polarization. Data were obtained over the temperature range f25PU. t o -75°C. at current densities of 2 , 4 , and G milliamp. per square centimeter. A sample of the data is shown in figure 11 and is summarized €or two metals in figures 12 and 18. These data indicate that polarization increases
ELECTROLYTES .4SD ELECTROCHEMIC.IL COUPLES
513
rapidly :is the load i n c r e ~ e aand as the temperature drops; this is very noticeable at tlic IOU temperatures. These observations prove t o be some of the most important ones of this entire investigation. They indicate that cell design is an importunt factor in ctrveloping A cell for loir-temperature operation.
2 0.
1.6-
x
1.2.
6-
4.
400
800
1200
1600
2000
2400
2800
R P.M
FIG.14 lliswlution of zinc in 90 per cent methanol-hydrochloric acid-water solutions. Rrncrion \.clocitv constant ( k ) for several tcmperntures plotted against stirring speed.
C. Srtntnaary of the better couples -1s A reaiilt of this study we have selected the following couples that appear more promising from the viewpoint of stability, flash current a t -73"C., and capacity 011 drain: Cd 43% HBF4 I 1 2 (with .IgCl) Cd j 43% HBFl I 31nOz Cd 1 43% fIBF4
I 1
,4gC1
Pb I 43% HBF4 PbOz Zn I CH3NH2.HC1-LiC1 I 31n02 3Ig I LiBr-11gBr2 I 1lnOz DISSOLUTION DATA
'The rates of reaction of zinc, cadmium, and magnesium in the methanol:\-at,er-hydrochloric acid system have been determined (1) over the temperature range of -!-25"C. to -6O"C., ( 2 ) under static and dynamic conditions, and ( 3 ) with and without "depolarizers." These data are summarized in figures 14 Lo 22. In these figures the term "velocity constant," I;, is used t o indicate the rate oi lissolution of the metal with time, temperature, concentration of acid, and rate ?f stirring. The velocity constant, 12, is defined here by the standard kinetics -qiintion for A first-order reaction as
514
B. GARRETT, J . WELSH, S. WOODRUFF, R. COOPER .4ND J. HEIKS
.4.
2.0.
1.6.
.
- 30'C. -60%.
FIG. 1.5. Dissolution of magnesium in 90 per cent methanol-hydrochloric acid-water solutions. Reaction velocity constant (12) plotted against stirring speed at vnriotis temperat urea.
.8 -
.7
-
.6
-
Y
R PM
FIG. 16. Dissolution of cadmium in 90 per cent methanol-hydrochloric acid-water solutions Reaction velocity constants ( k ) for several temperatures plotted against stirring speeds ___
1;
=
2.303V -log a At (a: ~
~
____
4
where A is the area of the metal in square centimeterh, 1- is the volume of the solution in milliliters, t is the time of the reaction in minutes, a is the potential reacting capacity of the original solution, and x is the amount of reactive mate-
ELECTROLYTES -4WD ELECTROCHEMICAL COUPLES
515
FIG. 17. Dissolution of zinc in 90 per cent methanol-hydrochloric acid-water solutions. Reaction velocity constant ( k ) plotted against temperature. Stirring speed, 1300 R . P . Y . ,031
.OSi
,026
Y
,021
.OI 4
,001
TEMP.
FIG. 18. Dissolution of cadmium in 90 per cent methanol-hydrochloric acid-water solutions. Reaction velocity constant (k) plotted against temperature. Static eondit ions,
516
A. 33. GbRRETT, J. WELSH, S. WOODRUFF, R. COOPER A S D J. IiEIIid
1.2.
t
.6-
.4-
.2
-
01
8
8
‘
I
+ 2 0 *I0 0 -10 -20 -Temp.%
1
”
$ 1
FIG. 19. Dissolution of iiisgnesiuni in 90 per cent methanol-hylrochlor ic acid-water solutions. Reaction velocity constant plotted against temperature. Stirring speed, 1620 R.P.M.
2
I
35
FIG. 20. Dissolution of cadmium. energy, 3770 csl./mole.
+. -lop
40
,
45
110-3
k 2’s. 1 / T , Static conditions. .lctivntion
517
ELECTROLYTES A S D ELECTROC'HEMICIL COUPLES
rial consumed in the reaction in the same units as a. The use of this type of expression is especially useful in a dissolution study over a Tvide temperature
I
* -7
* 3.5t
,
*,$"
*
-log k us. 1, 7'.
$.XG. 'Ii. Dissoiution oi zinc. cnergy, 3610 C R ~ .mole.
'
' 4.5
Stirring speed, 1.300 R.P.M. -4ctivation
- 125O I ~
TIG. 12. 1haolutioii o t 1n:igiiesium. tivation c n c r p . 4950 tal. molt.
+.lo3
-log k u s . 1 .1' Stirling speed, 1620
R.P 11
-1r-
range, since neccssary variations in the time of the runs and/or the concentration of the solutions are incorporated in the evaluation of k . T o test the validity of the application of the velocity constant t o dissolution rate studies, it is necessary only to show that 1; remains constant when the time of the run and/oy the eoncent.ration of the solution is varied, other conditions remaining constant.
518
A. B. GARRETT, J. WELSH,
S. WOODRUFF, R. COOPER AND J. HEIKS
Specific tests of this type were conducted for zinc, cadmium, and magnesium in dilute hydrochloric acid solution containing 90 per cent methanol, and in all cases IC was found to remain constant within the limits of experimental error. Several features of these data are of interest to LIS here: (I) the dissolution rate of these metals at -GO"C. is only about 5 per cent of that at room temperature, ( 2 ) the activation energy of dissolution of these metals under depolarized dynamic conditions is about 4000 cal., and ( 3 ) thc temperature coefficient is in the order of 1.2. These observations suggest t u o conclusions that may be drawn: First, the low activation energy and temperature coefficient are evidence that a physical process (probably diffusion) is the rate-determining step in the dissolution of the metal. If one assumes that the reaction in a cell involves diffusion, we can expect this t o be nn important parameter in cell reactions. Second, the 95 per cent drop in the reaction rate of metals with acids may also be expected to occur in a cell reaction; if so, it sets a very decided limit upon the rate at which current can he drawn from I: cell in which metals are used. COXCLUSIONS
The parameters involved in developing :i primary cell that n-ill operate over the temperature range of f25"C. t o -75'C. non hecome evident. The limited number of buitable sol.r.ent-electroly~e systems, the large increase in viscosity of qolvent-electrol~te, the lo\\. H a & current :it -73"C'., the high polarization of the electrodes at lo\\, temperatme., and the Ion- reaction rate of dissolution of metals a t low temperatures are the ohviou- factors that must be conbidered. Furthermore, the low activation energy and temperature coefficient of cliwolrition of metals would indicate that the iate-determining step is probahly a physical process, i.e., diffusion of the ions. This \\-auld produce concentratiou polai ization effects, especially a t Ion- tempmttuiw, and account for the high polarization observcd in thi:, n.orli. Moreover a diffusion-controlled process would probahly he viscosity controlled. Here, too, TI e find a correlation hetween the high viscosity and high polarization at the lon- temperature In addition, if we consider the viscosity as a rate proceis, n-e can calcuhte the :ICtivation energy of viscosity to be about 4000 cal.; this correlates Tvith thc order of magnitude of the activation energy of the dissolution proce The low rate of reaction at low temperature cannot be altered Imt the polarization process resulting from the variety of factors, i.e., high viscosity. slow diffusion rates, and possible surface coatings, can probably be overcome to some extent by a change of cell design and proper selection of the cell ingredients.
