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Research Article
Feasibility of Using an Electrolysis Cell for Quantification of the
Electrolytic Products of Water from Gravimetric Measurement
Samuel Melaku , Zewdu Gebeyehu , and Rajeev B. Dabke
Department of Chemistry, Columbus State University, Columbus, GA 31907, USA
Correspondence should be addressed to Rajeev B. Dabke; dabke_rajeev@columbusstate.edu
Received 8 August 2017; Revised 19 November 2017; Accepted 12 December 2017; Published 5 February 2018
Academic Editor: Pablo Richter
Copyright ©2018 Samuel Melaku et al. is is an open access article distributed under the Creative Commons Attribution License,
which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited.
A gravimetric method for the quantitative assessment of the products of electrolysis of water is presented. In this approach, the
electrolysis cell was directly powered by 9 V batteries. Prior to electrolysis, a known amount of potassium hydrogen phthalate
(KHP) was added to the cathode compartment, and an excess amount of KHCO
3
was added to the anode compartment
electrolyte. During electrolysis, cathode and anode compartments produced OH
−
(aq) and H
+
(aq) ions, respectively. Electro-
lytically produced OH
−
(aq) neutralized the KHP, and the completion of this neutralization was detected by a visual indicator color
change. Electrolytically produced H
+
(aq) reacted with HCO
3
−
(aq) liberating CO
2
(g) from the anode compartment. Concurrent
liberation of H
2
(g) and O
2
(g) at the cathode and anode, respectively, resulted in a decrease in the mass of the cell. Mass of the
electrolysis cell was monitored. Liberation of CO
2
(g) resulted in a pronounced eect of a decrease in mass. Experimentally
determined decrease in mass (53.7 g/Faraday) agreed with that predicted from Faraday’s laws of electrolysis (53.0 g/Faraday). e
ecacy of the cell was tested to quantify the acid content in household vinegar samples. Accurate results were obtained for vinegar
analysis with a precision better than 5% in most cases. e cell oers the advantages of coulometric method and additionally
simplies the circuitry by eliminating the use of a constant current power source or a coulometer.
1. Introduction
e quantity of a substance produced at the electrode and
the quantity of electric charge passed are linked with Far-
aday’s laws of electrolysis. Many reports present coulometric
methods of analysis of a variety of reagents [1–17]. A
number of original reports [3, 4, 11, 15–17] and reviews
[1, 2, 5–10, 12–14] on coulometric methods were published
over the past decades. Coulometric method has been applied
to analyze a large variety of inorganic and organic com-
pounds [1, 6, 13]. Standardization of reagents is not required
in coulometric analysis as the quantity of reagent prepared
is directly determined from the charge passing through
the cell.
Alternatively, the quantication of a reagent produced in
an electrolysis cell without direct monitoring of charge is
also well known [5]. is work includes early research on
estimating the charge passing through the cell from mea-
suring the mass of silver deposited on the platinum electrode
[18, 19]. Similarly, the measurement of color change has
been applied to quantify the charge passing through the cell
[20]. is method was employed for monitoring the re-
duction of permanganate ions at platinum electrode and the
oxidation of copper to blue cupric triethanolamine complex.
Although these methods sidestep the use of a coulometer or
a constant current source, they had practical constraints for
routine analytical purposes due to the involvement of large-
size precious metal electrodes or a rotating electrode placed
in a special type of electrochemical cell.
Simple measurement of the volumes of hydrogen and
oxygen gases produced at the cathode and anode in an
electrolysis cell oers advantages of coulometric method
without actually using a coulometer or any electronic
equipment. In fact, early studies indicated that hydrogen-
oxygen coulometer was easy to assemble and capable of
measuring the quantity of electric charge with an accuracy of
±0.1% or better [4]. By choosing a proper size buret for
monitoring the volume of gases, the instrument was readily
Hindawi
Journal of Analytical Methods in Chemistry
Volume 2018, Article ID 2681796, 5 pages
https://doi.org/10.1155/2018/2681796
adapted to measuring diverse quantities of electricity down
to about 10 coulombs [5]. Enhanced current eciency was
noted for the hydrogen–nitrogen coulometer [21] in which
hydrazine sulfate was used as an electrolyte responsible for
producing nitrogen gas instead of oxygen gas at the platinum
anode. An electrolysis cell was used to measure the volume
of O
2
(g) and monitor an acid-base titration [22].
As stated in (1), electrolysis of water produces OH
−
(aq)
in the cathode compartment. In a 1 : 1 stoichiometric ratio
(2), electrolytically produced OH
−
(aq) neutralizes the weak
acid (e.g., potassium hydrogen phthalate (KC
8
H
4
O
4
H) or
C
8
H
4
O
4
H
−
in an anionic form) added to this compartment:
2H2O(l) + 2e→2OH−(aq) + H2(g) (1)
C8H4O4H−(aq) + OH−(aq)→C8H4O2−
4(aq) + H2O(l)
(2)
Concurrently, anode produces H
+
(aq) the following
equation:
H2O(l)→2H+(aq) + 1
2O2(g) + 2e (3)
Electrolysis of water in cathode and anode compart-
ments (1) and (3) also produces H
2
(g) and O
2
(g) at the
respective electrodes. A chemical reaction (4) between
electrolytically produced H
+
(aq) and externally added
KHCO
3
(aq) produces CO
2
(g):
H+(aq) + HCO−
3(aq)→H2O(l) + CO2(g) (4)
When a known quantity of KHP (a primary standard
substance) is added to the cathode compartment, the end-
point of the neutralization reaction between electrolytically
produced OH
−
(aq) and KHP can be determined visually
with use of phenolphthalein as an indicator. Electrolysis is
promptly stopped at the endpoint. Prior to electrolysis, an
excess amount of KHCO
3
is added to the anode
compartment.
Concurrent liberation of chemically produced CO
2
(g) in
addition to H
2
(g) and O
2
(g) amplied the eect of de-
creasing in the mass. Mass of the cell is continuously
monitored before, during, and after electrolysis. From the
mass change experienced by the cell, the amount of KHP can
be determined to assess the stoichiometric relations between
the electrolytic products of water. In this study, the feasibility
of monitoring the mass change experienced by an elec-
trolysis cell due to the liberation of gases and quantifying the
electrolytic products of water was assessed.
2. Materials and Methods
2.1. Chemicals and Materials. Reagent-grade potassium ni-
trate, KHP, potassium hydrogen carbonate, and agar were
obtained from Fisher. Millipore deionized water was used
for preparation of reagents. Alfa Aesar platinum wire
(0.762 mm diameter) was used as purchased. 0.1% phe-
nolphthalein was used as a visual indicator to detect the
completion of neutralization. Ohaus Pioneer balance
(Model: PA114, 0.0001 g readability, and ±0.0002 g linearity)
was used for mass measurements. Vinegar samples were
locally purchased from a grocery store.
2.2. Electrolysis Cell. Electrolysis cell was made of spec-
trophotometer plastic cuvettes. A 3 mm hole was drilled on
a side near the base of each cuvette, and the cuvettes were
glued together using an epoxy. Bottom portions of two
additional cuvettes were cut and glued on top (Figure 1).
e extra depth oered by these cuvettes prevented the mass
loss due to the spattering of electrolyte mist formed due to
liberation of gas bubbles during electrolysis. A mixture of
agar and potassium nitrate (3% by mass and 1 M, resp.) was
added to water, and the mixture was heated to 90°C. A few
drops of this mixture were placed on the bottom of the
cuvettes. A gel was formed on the bottom of the cuvettes,
serving as a salt bridge. KNO
3
(aq) was used as an electrolyte
in both compartments. KHCO
3
was added to the anode
compartment. Eective concentrations of KNO
3
and
KHCO
3
were 1 M each. Platinum wires served as cathode
and anode and were directly powered by two 9−V batteries
connected in parallel. Approximately 3.5 cm platinum wire
was dipped in the electrolyte in each compartment. e
current passing through the cell was in the range of 55 mA
to 90 mA. A multimeter was used to display the current
passing through the electrolysis cell. Two miniature mag-
netic stir bars were placed in the compartments, and an
external stir bar retriever was used to maneuver the internal
stir bars in up and down directions. e movement of stir
bars helped dislodge the gas bubbles adhering to the
Electrolytically
produced
H+(aq), O2(g),
& chemically
produced CO2(g)
+−
Pt Pt
9V
Electrolytically produced
OH−(aq) and H2(g)
KHP (analyte) added
Agar salt bridge
Balance
Figure 1: Schematic diagram of the water electrolysis cell.
KNO
3
(aq) was used as an electrolyte in both compartments.
KHCO
3
(aq) was added to the anode compartment. Eective
concentrations of KNO
3
and KHCO
3
were 1 M each. Desired
number of moles of KHP(aq) and a drop of 0.1% phenolphthalein
were added to the cathode compartment. e volumes of elec-
trolyte in each compartment were approximately 3.2 mL. Platinum
wires were immersed in the electrolytes and directly connected to
two 9 V batteries connected in parallel. e cell was placed on
a digital balance housed in a glass cabinet.
2Journal of Analytical Methods in Chemistry
electrodes and sidewalls and stir the electrolytes in both
compartments. Additionally, electrolyte stirring assisted
uniform mixing of the contents of cathode compartment.
is action was periodically performed, particularly before
mass measurement.
2.3. Determination of Mass Correction Factor. Mass cor-
rection factor was determined to quantify an error associated
with the mass loss due to the evaporation of electrolytes. In
this experiment, mass of an idle cell (no electrolysis) was
monitored at an interval of 5 minutes. Mass of the cell was
monitored for 60 minutes (Figure 2).
2.4. Electrolytic Titrations of KHP. Desired moles of KHP
were added to the cathode compartment. Mass of the cell
was monitored at a time interval of one minute for the rst
ve minutes before electrolysis. After this initial rest period,
electrolysis was started. Electrolysis was paused at a two-
minute time interval, stir bars were maneuvered in up and
down position, and the mass was recorded after a two-
minute rest period. Electrolysis was promptly stopped when
the cathode compartment electrolyte turned light pink in-
dicating an endpoint of the titration. Mass measurement was
continued for additional eight minutes. e results of this
experiment are presented in Figure 3. Electrolytic titration of
KHP was repeated to determine the mass change experi-
enced by the cell for various concentrations (0 to 6.0 ×10
−4
)
of KHP (Figure 4).
2.5. Coulometric Measurements. In view of conrming the
stoichiometric relations between the moles of OH
−
(aq) and
the mass loss due to the liberation of electrolytically pro-
duced H
2
(g) and O
2
(g) and chemically produced CO
2
(g), we
monitored the charge passing through the electrolysis cell on
Faraday MP potentiostat. Electrolysis was paused at 10
coulombs interval, stir bars were maneuvered in up and
down position, and the mass was recorded after one-minute
rest period (Figure 5).
2.6. Electrolytic Titrations of Real Sample. A known volume
of commercial vinegar sample was added to the cathode
compartment and titrated against electrolytically produced
OH
−
(aq). e quantity of acid in a sample was determined
from the mass change experienced by the cell.
y=−(1.17 x 10−4)x + 19.1668
R2=0.9993
19.1595
19.1605
19.1615
19.1625
19.1635
19.1645
19.1655
19.1665
19.1675
0.0 10.0 20.0 30.0 40.0 50.0 60.0
Mass of the cell (g)
Time (min)
Figure 2: Monitoring the mass of an idle cell (no electrolysis).
17.8624
17.8664
17.8704
17.8744
17.8784
17.8824
17.8864
0.0 10.0 20.0 30.0 40.0
Mass of the cell (g)
Progress of electrolysis (min)
Electrolysis OFF
Electrolysis ON
Figure 3: Monitoring the mass of the electrolysis cell before,
during, and after electrolysis.
y=54.9x
R2=0.9936
0.0000
0.0050
0.0100
0.0150
0.0200
0.0250
0.0300
0.0350
0.00000 0.00020 0.00040 0.00060
Mass change (g)
Moles of KHP
Figure 4: A plot of the mass change experienced by the cell versus
number of moles of KHP neutralized in the cathode compartment.
17.4440
17.4540
17.4640
17.4740
17.4840
17.4940
0.0 20.0 40.0 60.0 80.0
Mass of the cell (g)
Charge (coulombs)
y=−(5.91 x 10−4)x + 17.4932
R²=0.9948
Figure 5: A plot of the mass of the electrolysis cell versus the charge
passing through the cell.
Journal of Analytical Methods in Chemistry 3
3. Results and Discussion
3.1. Mass Correction Factor. e mass of an idle cell linearly
decreased with time (Figure 2). A linear decrease in mass
indicated uniform mass loss resulting from the evaporation
of electrolyte. e slope of this plot quantied the mass
correction factor of 1.17 ×10
−4
g/min or 1.95 ×10
−6
g/s. is
factor was used to determine the mass loss due to evapo-
ration during the period for each trial of the experiment.
3.2. Quantitative Assessment of the Electrolytic
Products. Faraday’s laws of electrolysis and (1) and (3) in-
dicate, on passing of one Faraday charge through water, the
electrolysis cell produces one mole of OH
−
(aq) and H
+
(aq)
in their respective compartments. Additionally, 0.5 moles of
H
2
(g), 0.25 moles of O
2
(g), and 1 mole of CO
2
(g) are lib-
erated. Considering molar masses of these gases, a decrease
in mass by 53.0 g was expected to result from passing one
Faraday of charge.
Figure 3 presents data on the mass of the cell before,
during, and after electrolytic titration. Desired moles
(4.0 ×10
−4
moles) of aqueous KHP were added to the
cathode compartment, and the charge was passed through
the cell. Mass of the cell signicantly decreased during the
course of electrolysis. Mass of the cell slightly decreased
prior to electrolysis, resulting from the evaporation of
electrolyte, and it continued to decrease after electrolysis.
Mass correction factor was applied to minimize the error
associated with the evaporative mass losses during the
course of electrolysis.
e electrolytic titration experiment presented in Figure 3
was repeated for various concentrations of KHP, and the
mass correction factor was independently applied to each
trial of experiment. A linear response of mass change versus
moles of KHP conrmed the quantitative relationship be-
tween the moles of KHP and the drop in the mass due to the
liberation of gases (Figure 4). Slope of the plot indicated
a mass drop of 54.9 g due to the liberation of gases. is mass
accounted for the electrolytic neutralization of one mole of
KHP after passing one Faraday charge. is quantity of mass
obtained from the slope matched with the estimated mass
loss of 53.0 g/Faraday. A blank trial (a data point corre-
sponding to zero moles of KHP in Figure 4) conrmed
minimal contribution of acidic impurities and dissolved
CO
2
(g) present in the electrolyte. Literature values [23] of
solubility of H
2
(g), O
2
(g), and CO
2
(g) at 1 atm pressure and
at 25°C temperature also conrm a minor error in mass
measurement (<0.15%) associated with the dissolution of
gases produced in the cell. e trials presented in Figure 4
were repeated to determine the relative standard deviation in
the slope. e RSD value (0.93%, n�3) signied a reliable
quantitative relationship between the mass change and the
moles of KHP.
A linear response of the plot (Figure 5) conrmed the
mole stoichiometry presented in (1), (3), and (4). A slope of
this plot indicated that 5.91 ×10
−4
g mass drop occurred per
coulomb charge passed through the cell. Accounting for the
mass loss due to evaporation of the electrolyte, a correction
factor (1.95 ×10
−6
g/s) was applied to the slope. Considering
an estimated time (17.5s) required to pass one coulomb
charge, mass loss due to evaporation during the passage
of one coulomb was 3.41 ×10
−5
g/C. e mass-corrected
slope after deducting this factor was 5.57 ×10
−4
g/C or
53.7 g/Faraday. is mass drop was in agreement with the
estimated mass drop of 53.0 g/Faraday determined from the
stoichiometric relation presented in (1–4). A linear relation
between the mass loss due to liberation of gases and the
charge passed through the cell (Figure 5) indicated negligible
interference of dissolution of gases produced during elec-
trolysis. e slope of the plot presented in Figure 5 was
conrmed from three independent experiments (RDS
0.49%).
3.3. Analysis of Real Sample. e ecacy of the electrolytic
titration method presented in this paper was tested for the
analysis of acid content in household vinegar samples. A
known volume (500 μL) of vinegar was added to the cathode
compartment and titrated against electrolytically produced
OH
−
(aq) as presented in the experimental section. e
molarity and percent acid content in the sample were de-
termined from the mass change experienced by the cell
(Table 1). e mass correction factor was independently
applied to each trial. Quantity of acid determined from
electrolytic titration was in agreement with the manufac-
turer’s label, and the quantities were consistent with three
household vinegar samples.
4. Conclusions
In this work, the feasibility of using an electrolysis cell for
quantication of the electrolytic products of water from
gravimetric measurement was tested. e cell presented in
this paper enables in situ production of reagents and their
direct quantication and does not require standardization of
reagents. e electrolysis cell directly powered by 9 V bat-
teries eliminates the requirement of a constant current
source or a coulometer, yet oers advantages of coulometric
method of titration. e cell utilizes minimal volume of
reagents (3.2 mL or less in each compartment). e elec-
trolysis cell is simple, transparent, and easy to fabricate. e
Table 1: Results of titration of an acid in commercial vinegar samples by electrolytically produced OH
−
(aq).
Percent acid content stated on
manufacturer’s label
Mass change in
electrolysis cell (g)
Moles of acid in
(mmol)
Average
molarity
Percent acid content determined
from this study (95% CI, n�3)
Percent
RSD
4 0.0165 0.311 0.623 3.74% (±1.32) 14.3
5 0.0222 0.419 0.838 5.03% (±0.53) 4.3
7 0.0305 0.575 1.15 6.91% (±0.52) 3.0
4Journal of Analytical Methods in Chemistry
cell eliminates the requirement of an external salt bridge or
a fritted glass membrane. Linear response of the decrease in
the mass of the cell to the moles of KHP added and to the
charge passing through the cell validates the applicability of
the method. An agreement between the estimated drop in
mass (determined from Faraday’s laws of electrolysis) and
the experimentally determined drop in mass underlines the
feasibility of using the electrolysis cell for quantication of
the electrolytic products. With acceptable values of relative
standard deviation (better than 5% in most cases), the mass
drop experienced by the cell quantitatively responds to the
acid content in household vinegar samples.
Conflicts of Interest
e authors declare that there are no conicts of interest
regarding the publication of this paper.
Acknowledgments
e authors thank James O. Schreck and reviewers of the
manuscript for their helpful comments and suggestions.
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