A simplified method for titration of carbonate and hydroxide in alkaline battery electrolyte is presented involving a saturated KSCN solution as a complexing agent for zinc. Both hydroxide and carbonate can be determined in one titration, and the complexing reagent is readily prepared. Since the pH at the end point is shifted from 8.3 to 7.9 - 8.0, m-cresol purple or phenol red are used as indicators rather than phenolphthalein. Bromcresol green is recommended for determination of the second end point of a pH of 4.3 to 4.4

Reid, M. A.

English

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NASA Technical Memorandum 78961
(NASA-TM-78961) AN IMPROVED METHOD FOR
	
N78-28607
ANALYSIS OF HYDROXIDE AND CARBONATE IN
ALKALINE ELECTROLYTES CONTAINING ZINC (NASA)
14 p HC A02/MF A01	 CSCL 10C	 :, aclas
G3/44 27106
AN IMPROVED METHOD FOR ANALYSIS OF
HYDROXIDE AND CARBOr^ATE IN ALKALINE
ELECTROLYTES CONTAINING ZINC
Margaret A. Reid
Lewis Research Center
Cleveland, Ohio 44135
;Uu 1978
	
ti	 1 r.
r.•	 n ,
TECHNT_CAL PAPER to be presented at the	
^l 1 '	 4L^''
One-hundred fifty-fourth Meeting of
The Electrochemical Society, Inc.
Pittsburgh, Pennsylvania, October 15-20, 1978
AN IMPROVED METHOD FOR ANALYSIS OF HYDROXIDE AND CARBONATE
IN ALKALINE ELECTROLYTES CONTAINING ZINC
Margaret A. Reid*
National Aeronautics and Space Administration
Lewis Research Center
Cleveland, Ohio 44135
N	 ABSTRACT
A simplified method for titration of carbonate and hydroxide in alkaline
W
battery electrolyte is presented involving a saturated KSCN solution as a com-
plexing agent for zinc. Both hydroxide and carbonate can be determined in one
titration, and the complexing reagent is readily prepared. Since the pH at the
end point is shifted from 8.3 to 7.9 - 8. 0, m-cresol purple or phenol Feu are
used as indicators rather than phenolphthalein. Bromcresol green is recom-
mended for determination of the second end point of a pH of 4.3 to 4.4.
INTRODUCTION
Silver-zinc batteries are being used for terrestrial and space applica-
tions, and nickel-zinc batteries are under serious consideration for electric
vehicle propulsion and other terrestrial applications. Greater economy and
usage of these batteries would result if the battery life were extended. System-
atic improvement is possible if the modes of failure are known and understood;
this often requires a determination of the hydroxide and carbonate content of
the electrolyte as well as that of the zinc. The hydroxide and carbonate content
of an alkaline sample are usually determined by titration with a standard acid.
Two end points are used. One is at a pH ^-8. 3 and the other at a pH -4.3. In
one method excess Ba+2
 is added since Ba(OH) 2 is soluble at pH 8 and BaCO3
*Electrochemical Society Active Member.
Keywords; alkaline electrolytes, batteries, chemical analysis, acid-base
titrations, hydroxides, carbonates, thiocyanate complexes.
2
is insoluble. Thus only OH is neutralized at the first end point, and only
CO3-2 is neutralized between the two end points. If no Ba+2 is present, the
CO3-2 is converted to HCO 3 at the first end point. The HCO 3 is neutralized
to carbonic acid between the two end points. The details can be found in most
college textbooks of analytical chemistry.
In the titration of battery electrolyte the presence of dissolved zinc species
from the zinc electrode introduces errors at the first end point due to the for-
mation of zinc hydroxides and carbonates, both of which are insoluble at pH N 8.
A previous method developed at NASA Lewis Research Center (ref. 1) utilized
the barium salt of EDTA (f-thylenediaminetetraacetic acid) in an excess of BaC12
to complex the zinc ions by the reaction
Zn+2 + BaY 2 = ZnY-2 + Ba+2
where Y-4  is the anion of ethylenediaminetetraacetic acid. In the titration to
pH N8, all of the hydroxide will be neutralized and none of the carbonate. A
separate sample is taken for determination of total alkalinity by titration to
pH —4.3. Since both Zn(OH) 2 and ZnCO 3 are soluble at this pH, there is no in-
terference by the zinc at the second end point.
This procedure has the advantage of determining OH- without any interfer-
ence from the zinc species. However, a second sample without the BaEDTA
reagent must be titrated to determine the CO3-2 content because EDTA is a
weak acid. This means that the reagent would also significantly affect the
amount of standard acid used in a titration if its initial pH were different from
that of the end point. In addition, the BaEDTA solution must be adjusted to a
pH of 8 just before the titration since a precipitate, presumably BaEDTA,
forms within a few hours of this pH. For further details, the reader is re-
ferred to the original publication. For these reasons, a second method has
been developed which utilizes KSCN as the complexing reagent instead of EDTA.
3
Thiocyanate was investigated as a complexing agent for Zn +2 because thiocyanie
acid is a strong acid and the presence of its salts would not affect the titration
to either end point significantly.
THEORY
The interference due to Zn+2 can be avoided by the use of SCN- if the com-
plex does not permit the precipitation of ZnCO 3 or Zn(OH) 2 at pH N8. The pos-
sibility can be investigated by using the constants for the various reactions which
are involved. Table I lists the pertinent ones.
TABLE I. - EQUILIBRIUM CONSTANTS FOR PERTINENT REACTIONS
IN THE TITRATION OF ALKALINE BATTERY
ELECTROLYTES CONTAINING ZINC
Reaction Constant Reference
ZnCO 3 = Zn+2 + CO3 2 2x10-10 2
Zn(OH) 2 = Zn+2 + 2 OH 4.7x10-17 2
Zn(OH) 4-2 = Zn+2 +40 2.8x1015 2
HCO 3 = H+ + CO3-2 5.6x10-11 3
Zn+2 + SCN = ZnSCN+ 2.4x10-4 4
The SCN complex is rather weak. However, the solubility of KSCN is
large enough (-10 M) so that enough can be added to prevent the precipitation
of ZnCO 3 and possibly Zn(OH) 2 at the first end point although both may pre-
cipitate in the intermediate stages.
We will assume for purposes of calculation that the initial battery elec-
trolyte has the following approximate concentrations; 
	 = 10 M,
ICO
3-2 = 1 M, and ZnI = 1 M. These are typical concentrations for actual
e 	 If a 0.5 cm3 sample is taken for analysis, diluted, and aliquots
amounting to 1/5 of this solution analyzed, the amounts of OH - , CO3 2, and
Zn+2 in each sample will be 1 mmole (millimole), 0.1 mmole, and 0.1 mmole,
4
respectively. If the sample is diluted to a final volume of 100 cm 3, the respec-
tive concentrations of these ions will be 10-2 , 10-3, and 10-3 M.
If the concentration of SCN in the solution at the start of the titration is
5 M, since
[Zn+ I [-';C N] _ 2.4x10-2
1ZnSCN1
then
^Zn+2^
^ 5x10-3
1ZnSCN]
i.e., almost all of the zinc will be complexed with the SCN . The concentra-
tion of the uncomplexed zinc will be about 5x10 -6 M, and the concentration of
hydroxide will be 10 -6 M at the end point. Therefore, the ion products for
Zn(OH) 2 and ZnCO3 will be 5x10-18 and 2.8x10-11 , less than the solubility
product constants, and both will be soluble. If the zinc concentration had not
been reduced below 10 -3 M by complexing with the thiocyanate, both would be
insoluble at this point.
Using the above constants, one finds that in the initial battery electrolyte,
zinc is soluble due to the formation of the tetrahydroxyzincate ion, but that
Zn(OH)2 precipitates when the sample is diluted before taking of aliquots if the
OH concentration falls below about 1 M. However, Zn(OH) 2 forms supersatur-
ated solutions readily, especially in the presence of saturated KSCN, so that ii
the aliquots are taken within about an hour, no precipitation usually occurs.
Zinc carbonate will also precipitate in the intermediate stages even in the
presence of KSCN, but as shown above, both precipitates will dissolve just be-
fore the end point is reached.
5
EXPERIMENTAL
Materials and Apparatus
All reagents were C. P. chemicals. Na 2CO3 was prepared by heating
NaHCO3 to 2700 to 3000 C for several hours. The zinc stock solution was pre-
pared from ZnSO4
 • 7 H2 O and standardized by titration with EDTA.
Titration curves were made using an automatic potentiometric titrator
having a derivative mode as well as a pH mode. In addition, the rate of de-
livery of titrant is automatically adapted to the slope of the titration curve,
and the rate of adaptation can be varied.
The indicators used are described below.
Indicator Range Color
Acid Base
Phenolphthalein 8.2 to 10.0 Colorless Pink
m-cresol purple 7.4 to 9.0 Yellow Purple
Phenol red 6.8 to 8.4 Yellow Red
Methyl orange 3.1 to 4.4 Red Orange-yellow
Bromcresol green 3.0 to 4.6 Yellow Blue
Titration Curves
In order to check the experimental method for precision and accuracy, so-
lutions were prepared to simulate typical compositions obtained when samples
of battery electrolyte were diluted for titration. Figure 1 shows a titration
curve for a solution containing 0.859±0.003 mmoles(millimoles) KOH, 0.012
mmoles of K 2CO3 (as impurity in the KOH), 0.515010.0010 mmoles of
Na2CO3, and 0.1274±0, 001 mmoles of ZnSO4 with and without the addition of
28 ml of saturated KSCN, which is the equivalent amount added if the battery
electrolyte is diluted as described below. Curve A is the titration curve
without th,.,  addition of KSCN. The inflection points are not sharp and do not
6
occur at the stoichiometric end points. In curve B, saturated KSCN has been
added to the same solution. This shows satisfactory inflections at the proper
equivalence points. Curve C is the derivative curve for the same solution as
curve C. A small amount of acid, 0.021 mmoles in this case, is consumed in
titrating the KSCN to the second end point and for precision a blank correction
should be made. The calculated mmoles of HCl to the first and second end
points are 1.386±0.003 and 1. 9344.003 including the KSCN correction. For
five determinations we find 1.386±0.004 and 1.937-+0.006 mmoles.
In figure 2 similar titration curves are given for lower levels of car-
bonate and zinc. The calculated values to the first and second end points
are 1.094±0.004 and 1.341±0.004 mmoles, and the experimental valves
1.099±0.003 and 1.349±0.004. In figure 3 a curve is shown for a low level
of carbonate showing the utility of the derivative curve. The first end point
is much more readily obtained when the derivative is plotted rather than the
pH when low levels of carbonate are present.
The reason for the necessity of a blank correction for titration of KSCN
to the second end point is not clear. Several authors have found no evidence
for association of SCAT with H+ , but one aut1hor has found a slight degree of
association (ref. 5). Small amounts of impurities in the highly concentrated
solutions used could account for the blank. It is advisable to determine the
blank correction with a sample of the actual KSCN solution to be used in the
analyses.
Determination of the End Points by pH Meter or Indicators
If an -automatic titrator is not available, a pH meter or indicators can
be used to determine the end point. In the standard method for determination
of hydroxide and carbonate, phenolphthalein is used to detect the first end
point (pH -8.3) by titration to a very pale pink color. Methyl orange or
bromcresol green are use-1 to determine the second end point at a pH of
7
around 4.3. In the method described here, the pH at the first end point is
shifted from 8.3 to 7.9 - 8.0. The reason for this is not known, but possibly
is due to the effects of the high concentration of KSCN (about 5 M in the final
solution) on th activity coefficients of the ions present. If the end point is de-
termined by a pH meter, the titration to the first end point should be carried
out to a pH of about 7.9 to 8.0. If indicators are to be used, phenolphthalein
is no longer satisfactory since it changes from pink to colorless before the end
point is reached. Phenol red or m-cresol purple have been found satisfactory.
In both cases, titration is carried out to the appearance of a faint yellow color
in the solution. Metauresol purple is preferable due to a more distinctive
color change. For determination of the second end point, either methyl
orange or bromcresol green can be added after the first end point is reached.
Due to the presence of the yellow form of m-cresol purple or phenol red in
the solution, the color at the end point will be somewhat different from the
color of these indicators alone. If methyl orange is used, the second end
point is reached when the solution becomes a peach color, and if bromcresol
green is used, the end point is reached when the solution becomes completely
yellow with no tinge of green. Bromcresol green is preferable, again since
it has a more distinctive color change. If a pH meter is used, titration to a
pH of about 4.3 is satisfactory.
Recommended Analytical Procedure
A 0.5 cm3 sample of electrolyte is withdrawn from the battery with a
Syringe or pipet and placed in a 25 cm 3 volumetric flask. Fifteen cm 3 of
saturated KSCN are added and the solution diluted to the mark with distilled
water. The solution is immediately shaken and 5 cm 3 aliquots withdrawn.
Zinc hydroxide will precipitate at this pH, but ordinarily no precipitate forms
for at least a half-hour since the hydroxide readily forms supersaturated so-
lutions. The addition of the saturated KSCN seems to help retard the onset
8
of precipitation. If a precipitate does form, it does not settle for some time so
that accurate aliquots can still be taken.
Since the battery electrolyte is quite viscous, there is a question as to the
accuracy of sampling with a syringe or pipet. Experiments in our laboratories
indicate +2% accuracy using a syringe, : k1% with pipets; this is sufficient for
most purposes. Greater accuracy can be obtained by weighing samples into a
10 cm3 volumetric flask before dilution with KRI'N. In this case, 2 cm 3 ali-
quots would be taken.
Twenty-five cm 3 of saturated KSCN are added to each aliquot and the so-
lution is titrated with standardized HCl of about 0.05 M concentration. If an
automatic titrator is available, it can be used in either a pH or derivative
mode. The latter is especially useful when there are very small quantities
of carbonate in the solution, as shown in figure 3. In the absence of an auto-
matic titrator, a pH meter can be used to titrate to pH values of 7.9 to 8.0
for the first end point and 4.3 to 4.4 for the second end point. Alternatively,
either m-cresol purple or phenol red can be used to determine the first end
point and either bromcresol green or methyl orange to determine the second.
CALCULATIONS
This titration gives the CO3 2 content of the electrolyte as well as the
amount of OH- from both the KOH and the Zn(OH)4 2 in the solution. Zinc
oxide will dissolve in KOH by the reaction ZnO + H2 O + 2 KOH = Zn(OH)4-2.
At the first end point, the zinc will be complexed as ZnSCN+ so that each
mole of dissolved zinc will provide 2 moles of hydroxide for neutralization.
Mmoles to first end point = (om3 HCI) (MHC1)
= mmoles OH- + mmoles CO3-2
8Additional mmoles to second end point = (additional cm 3 HCl) (MHCI)
= mmoles CO3-2 + mmoles for KSCN blank
For the procedure described above, 28 cm 3 of saturated KSCN should be titra-
ted to determine the blank (3 ml KSCN are contained in the aliquot and 25 ml
more are added) . If samples are diluted in a different manner, appropriate
quantities should be taken. If the carbonate content is large and accuracy is
needed of only a few percent, the blank can be neglected.
Since 1/5 aliquots are titrated,
[OHJ iii electrolyte = 5 x mmoles OH-
Sample volume
5 x mmoles CO3-2CO3 j in electrolyte = Sample volume
If weighed samples are used,
mg K2CO3 = mmoles CO3-2 x 138.2
wt % K2CO3 = mg K2CO3 x 100/mg sample
mg KOH = mmoles OH- x 56.11
wt % KOH = mg KOH x 100/mg sample
If one wishes to express the results as ZnO + KOH + K2CO31 it is necessary
to analyze a separate portion of electrolyte for zinc, probably by titration with
EDTA. See appendix for further details. In this case,
mmoles KOH = mmoles OH- - 2 (mmoles Zn)
10
CONCLUDING REMARKS
The data presented here show that battery electrolyte containing zinc can be
analyzed conveniently and accurately for hydroxide and carbonate using a satu-
rated solution of KSCN to complex the zinc at the first end point. Accuracy ob-
tained is 3 to 4 parts per 1000. The titration then differs from the usual deter-
mination of hydroxide and carbonate in only two aspects. First, the first end
point is shifted from 8.3 to 7.9 - 8.0, so that if an indicator is used to determine
the end point, m-cres,,l purple or phenc,l red must be used instead of phenol-
phthalein. Secondly, a small but noticeable correction must be made for the
amount of acid required to titrate the KSCN ad 3ed between the two end points.
The simplification of the ,analysis for alkaline battery electrolyte enables
identification of the modes of failure of these batteries to be determined more
readily and is useful in research to lengthen the life of alkaline batteries.
11
d pPENDIX
For the convenience of the reader a few details of our method of titration
for zinc are included. An aliquot of the solution of battery electrolyte that has
been diluted with KSCN as described above is added to 5 cm 3 of pH 10 buffer
(-8 M NH3, 1.5 M NH4.vO3). The aliquot should contain 0.05 or 0.10 8m3
of the original electrolyte. The solution is titrated with 0.005 M EDTA. The
added KSCN does not interfere since the EDTA complex of zinc is stronger than
the SCN complex. Colorimetric determination of the end point can be used
(ref. 6), but if an automatic titrator is available, it is convenient to use am-
perometric titration with dual polarized amalgamated silver electrodes (refs.
7, 8, and 91. A potential difference of about 200 mV is maintained between the
electrodes, and the titration is carried out until a sharp rise in the current is
noted. Ammonium nitrate is used in the buffer rather than ammonium chlor-
ide to avoid interference of the chloride with the mercury electrode.
12
REFERENCES
1. D. A. Otterson, "Determination of Hydroxide and Carbonate Contents of Alka-
line Electrolytes Containing Zinc," NASA TN D-8007 (1975).
2. W. M. Latimer, "Oxidation Potentials," Prentice-Hall, New Jersey (1956).
3. "Handix ok of Chemistry, ' N. A. Lange, Editor, McGraw Hill Book Co.,
New York (1961).
4. K. B. Yatsimerskii and V. P. Vasil'ev, "Instability Constants of Complex
Compounds," D. A. Paterson, Tranal. , Pergamon Press, New York
(1960).
5. L. G. Sillen and A. E. Martell, "Stability Constants of Metal-Ion Com-
plexes," The Chemical Society, London (1964).
6. F. J. Welcher, "The Analytical Uses of Ethylenediaminetetraacetic Acid,"
Van Nostrand, New Jersey (1957).
7. G. Guerin, J. Desbarres, and B. Tremillon, Electroanal. Chem. 1, 226
(1960).
8. G. Charlot, J. Badoz-Lambling, and P. Tremillon, "Electrochemical Reac-
tions," Elsevier, New York (1962).
9. "The Complexometric Titration of the Alkaline Earth Metals and Aluminum
with Amperometric End-Point Indication," Metrohm Application Bulletin
75e, Switzerland (Nov. 25, 1970). Available from Brinkman Instruments,
Inc., New York.
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An improved method for analysis of hydroxide and carbonate in alkaline electrolytes containing zinc

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