Recovery of Li and Co in Waste Lithium Cobalt Oxide-Based Battery Using H1.6Mn1.6O4

Abstract

H1.6Mn1.6O4 lithium-ion screen adsorbents were synthesized by soft chemical synthesis and solid phase calcination and then applied to the recovery of metal Li and Co from waste cathode materials of a lithium cobalt oxide-based battery. The leaching experiments of cobalt and lithium from cathode materials by a citrate hydrogen peroxide system and tartaric acid system were investigated. The experimental results showed that under the citrate hydrogen peroxide system, when the temperature was 90 °C, the rotation speed was 600 r·min−1 and the solid–liquid ratio was 10 g·1 L−1, the leaching rate of Co and Li could reach 86.21% and 96.9%, respectively. Under the tartaric acid system, the leaching rates of Co and Li were 90.34% and 92.47%, respectively, under the previous operating conditions. The adsorption results of the lithium-ion screen showed that the adsorbents were highly selective for Li+, and the maximum adsorption capacities were 38.05 mg·g−1. In the process of lithium removal, the dissolution rate of lithium was about 91%, and the results of multiple cycles showed that the stability of the adsorbent was high. The recovery results showed that the purity of LiCl, Li2CO3 and CoCl2 crystals could reach 93%, 99.59% and 87.9%, respectively. LiCoO2 was regenerated by the sol–gel method. XRD results showed that the regenerated LiCoO2 had the advantages of higher crystallinity and less impurity.

Full text

Citation: Wang, H.; Chen, G.; Mo, L.;
Wu, G.; Deng, X.; Cui, R. Recovery of
Li and Co in Waste Lithium Cobalt
Oxide-Based Battery Using
H1.6Mn1.6 O4.Molecules 2023,28, 3737.
https://doi.org/10.3390/
molecules28093737
Academic Editor: Krzysztof
Pielichowski
Received: 13 March 2023
Revised: 5 April 2023
Accepted: 19 April 2023
Published: 26 April 2023
Copyright: © 2023 by the authors.
Licensee MDPI, Basel, Switzerland.
This article is an open access article
distributed under the terms and
conditions of the Creative Commons
Attribution (CC BY) license (https://
creativecommons.org/licenses/by/
4.0/).
molecules
Article
Recovery of Li and Co in Waste Lithium Cobalt Oxide-Based
Battery Using H1.6Mn1.6O4
Hua Wang 1,2, Guangzhou Chen 1,2,3,*, Lijie Mo 1,2, Guoqiang Wu 1,2, Xinyue Deng 1,2 and Rong Cui 4
1Anhui Key Laboratory of Water Pollution Control and Waste Water Recycling, Anhui Jianzhu University,
Hefei 230601, China
2Anhui Key Laboratory of Environmental Pollution Control and Waste Resource Utilization, Anhui Jianzhu
University, Hefei 230601, China
3Anhui Research Academy of Ecological Civilization, Anhui Jianzhu University, Hefei 230601, China
4School of Environment and Energy Engineering, Anhui Jianzhu University, Hefei 230601, China
*Correspondence: chgzh5@163.com
Abstract:
H
1.6
Mn
1.6
O
4
lithium-ion screen adsorbents were synthesized by soft chemical synthesis
and solid phase calcination and then applied to the recovery of metal Li and Co from waste cathode
materials of a lithium cobalt oxide-based battery. The leaching experiments of cobalt and lithium
from cathode materials by a citrate hydrogen peroxide system and tartaric acid system were investi-
gated. The experimental results showed that under the citrate hydrogen peroxide system, when the
temperature was 90
◦
C, the rotation speed was 600 r
·
min
−1
and the solid–liquid ratio was 10 g
·
1 L
−1
,
the leaching rate of Co and Li could reach 86.21% and 96.9%, respectively. Under the tartaric acid
system, the leaching rates of Co and Li were 90.34% and 92.47%, respectively, under the previous
operating conditions. The adsorption results of the lithium-ion screen showed that the adsorbents
were highly selective for Li
+
, and the maximum adsorption capacities were 38.05 mg
·
g
−1
. In the
process of lithium removal, the dissolution rate of lithium was about 91%, and the results of multiple
cycles showed that the stability of the adsorbent was high. The recovery results showed that the
purity of LiCl, Li2CO3and CoCl2crystals could reach 93%, 99.59% and 87.9%, respectively. LiCoO2
was regenerated by the sol–gel method. XRD results showed that the regenerated LiCoO
2
had the
advantages of higher crystallinity and less impurity.
Keywords:
waste lithium cobalt oxide based battery; H
1.6
Mn
1.6
O
4
; citric acid; tartaric acid;
regenerated
lithium cobaltate
1. Introduction
With the wide application of lithium batteries in mobile communications, electric
vehicles and other fields, environmental problems caused by their e-wastes are becoming
increasingly prominent, especially heavy metal pollution caused by cobalt, lithium and
manganese in batteries [
1
]. Therefore, the recovery of valuable metals in waste lithium-
ion batteries has attracted the attention of many scholars. The main recovery methods
are pyrometallurgy [
2
] and hydrometallurgy [
3
]: (1) the first process was simple, but it
had the disadvantages of high energy consumption, complex temperature control and
easy-to-produce harmful dust and gas; (2) the second process had a high recovery rate
of valuable metals and high purity of products, which is the mainstream technology of
waste lithium-ion battery recovery at present. The wet process was mainly acid leaching,
in which metal ions were dissolved by acid and then selectively recovered from metals
such as lithium and cobalt. Among them, common acids could be divided into inorganic
acids and organic acids. Inorganic acids were usually hydrochloric acid [
4
], nitric acid [
5
],
sulfuric acid [
6
], etc. The above strong acids were used to destroy the structure of positive
electrode materials and achieved the purpose of recycling. However, strong inorganic acids
had defects, such as corroding machinery and producing waste gas and wastewater. On
Molecules 2023,28, 3737. https://doi.org/10.3390/molecules28093737 https://www.mdpi.com/journal/molecules

Molecules 2023,28, 3737 2 of 15
the contrary, organic acids had mild acidity, little corrosion and were easy to decompose,
so they had an excellent performance as leaching agents [
7
]. In addition, some alkaline
leaching processes were used to recover aluminum in batteries [8].
For metal ions after leaching, the precipitation method [
9
], solvent extraction method [
10
]
and other methods are usually used for separation. Due to the many separation steps
and difficult separation, the applicability was poor. Among the treatment methods of
heavy metal pollution in water, the adsorption method was often used. For example,
microplastics [
11
] and chitosan [
12
] had a strong adsorption capacity for metal ions. At
present, there is little research on metal ion separation by the adsorption method, the reason
being that the common adsorbents often could not adsorb a single metal ion alone. A
MnO
2·
0.5H
2
O lithium-ion sieve obtained by elution of a Li
1.6
Mn
1.6
O
4
precursor is currently
the adsorbent with the largest adsorption capacity of lithium ions [
13
]. Moreover, the Li-ion
sieve was highly selective in the adsorption of Li
+
, which provided a possibility for the
adsorption method to be applied to the separation of lithium in lithium cobaltate batteries.
In this study, citric acid and tartaric acid in organic acids were selected as acids in the
leaching agent, and the leaching experiment was carried out on the waste cathode material
of a lithium cobalt oxide-based battery. A lithium-ion screen with selective adsorption ca-
pacity was used to adsorb, separate and recover Li
+
from the leaching solution. Two routes
were tested for cobalt recovery: one was to reprecipitate to form cobalt chloride crystals,
and the other was to use recycled lithium as a lithium source to regenerate lithium cobaltate.
2. Results and Discussion
2.1. Leaching Rules of Positive Grade Materials
2.1.1. Powder Parameters of Cathode Material
The material came from the cathode material of a waste lithium cobalt oxide-based
battery, which was mainly composed of lithium cobaltate. During the disassembly process,
the aluminum shell was manually removed to fully remove the attached organic plastics
and adhesives. The element composition of the cathode material was determined, as shown
in Table 1. The main metals were 60.2% Co and 7.2% Li, respectively, while other metal
elements such as Fe, Al and Ni were all lower than 1%. Lithium cobalt oxide-based batteries
on the market might also contain Na, Cu, Zn and other elements, which were not detected
in the powder.
Table 1. Positive material group division table.
Type of Metal Li Co Fe Al Ni
Proportion (%) 7.2 60.2 0.02 0.01 0.01
2.1.2. Leaching Kinetics
The leaching efficiency of the valuable metals was affected by many factors, such
as leaching acid concentration, H
2
O
2
concentration, temperature, rotational speed and
reaction time. To control these factors, the authors conducted a lot of exploration, which
referred to other studies in the literature [
14
–
17
]. So, in this study, the temperature was
determined to be 90
◦
C, and the rotation speed was 600 r
·
min
−1
. With the increase in
reaction time, the metal leaching efficiency in the cathode material is shown in Figure 1,
where the leaching rates of Co and Li reached the highest of 86.21% and 96.9% around 4 h
(Figure 1a). Then, the data were fitted according to the control velocity equation of the
chemical reaction (Figure 1b), where X is the leaching efficiency (%) and t is the leaching
time (min). The determination coefficients R
2
of Co and Li fitting curves were 0.95334 and
0.99447, both above 0.95, indicating a high degree of fitting; that is, metal leaching was
controlled by chemical reaction.

Molecules 2023,28, 3737 3 of 15
Molecules 2023, 28, x FOR PEER REVIEW 3 of 16
0 50 100 150 200 250 300 350 400
20
40
60
80
100
Leaching rate(%)
t(min)
(a)
Li
Co
0 50 100 150 200 250
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
y=0.03487+0.00309t
R
2
=0.99447
(b)
Li
Co
1-3(1-X)
2/3
+2(1-X)
t(min)
y=0.04235+0.00199t
R
2
=0.95334
Figure 1. Curves of citric acid leaching performance (a) curves with time; (b) dynamic fiing curves.
When tartaric acid was used as the leaching agent, the metal leaching efficiency in
the cathode material was shown in Figure 2. The leaching rates of Co and Li basically
increased with the increase of time and reached the highest values of 90.34% and 92.47%
at about 5 h, respectively. However, they decreased when the time exceeded 5 h. Then,
the data were fied according to the control velocity equation of the chemical reaction.
The determination coefficients R2 of the fiing curves of Co and Li were 0.99027 and
0.7692, and the fiing curves of Co had a higher fiing degree, which was controlled by
chemical reaction, while Li was obviously easier to leach in tartaric acid. The leaching rate
reached 75.29% in the first 30 min. The fiing result of the control velocity equation of the
chemical reaction was poor, which was not suitable for Li leaching in a tartaric acid sys-
tem.
0 50 100 150 200 250 300 350 400
20
40
60
80
100
Leaching rate(%)
t(min)
(a)
Li
Co
0 50 100 150 200 250 300
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
y=0.20294+0.00146t
R
2
=0.7692
(b)
Li
Co
1-3(1-X)
2/3
+2(1-X)
t(min)
y=0.01939+0.00186t
R
2
=0.99027
Figure 2. Leaching performance curves under tartaric acid (a) change curves with time; (b) dynamic
fiing curves.
Cobalt trivalent could not be directly dissolved by acid, so hydrogen peroxide was
used as a reducing agent to make cobalt oxide accept an electron of hydrogen peroxide,
then was dissolved by citric acid to form cobalt citrate. The leaching reaction is described
by Equation (1) [18].
2HCit(aq)+2LiCoO
(s) + HO(aq) = 2Li (aq) + 2Cit (aq) + 2Co (aq) + 4HO + O(g) (1)
When tartaric acid was used as the leaching solution, it also played a reducing role.
So, the involvement of hydrogen peroxide was not required. The reaction with lithium
cobaltate is described by Equation (2) [7].
LiCoO+ 3C
HO→Co
(CHO)+LiC
HO+ 2H
O (2)
Figure 1.
Curves of citric acid leaching performance (
a
) curves with time; (
b
) dynamic fitting curves.
When tartaric acid was used as the leaching agent, the metal leaching efficiency in the
cathode material was shown in Figure 2. The leaching rates of Co and Li basically increased
with the increase of time and reached the highest values of 90.34% and 92.47% at about 5 h,
respectively. However, they decreased when the time exceeded 5 h. Then, the data were
fitted according to the control velocity equation of the chemical reaction. The determination
coefficients R
2
of the fitting curves of Co and Li were 0.99027 and 0.7692, and the fitting
curves of Co had a higher fitting degree, which was controlled by chemical reaction, while
Li was obviously easier to leach in tartaric acid. The leaching rate reached 75.29% in the
first 30 min. The fitting result of the control velocity equation of the chemical reaction was
poor, which was not suitable for Li leaching in a tartaric acid system.
Molecules 2023, 28, x FOR PEER REVIEW 3 of 16
0 50 100 150 200 250 300 350 400
20
40
60
80
100
Leaching rate(%)
t(min)
(a)
Li
Co
0 50 100 150 200 250
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
y=0.03487+0.00309t
R
2
=0.99447
(b)
Li
Co
1-3(1-X)
2/3
+2(1-X)
t(min)
y=0.04235+0.00199t
R
2
=0.95334
Figure 1. Curves of citric acid leaching performance (a) curves with time; (b) dynamic fiing curves.
When tartaric acid was used as the leaching agent, the metal leaching efficiency in
the cathode material was shown in Figure 2. The leaching rates of Co and Li basically
increased with the increase of time and reached the highest values of 90.34% and 92.47%
at about 5 h, respectively. However, they decreased when the time exceeded 5 h. Then,
the data were fied according to the control velocity equation of the chemical reaction.
The determination coefficients R2 of the fiing curves of Co and Li were 0.99027 and
0.7692, and the fiing curves of Co had a higher fiing degree, which was controlled by
chemical reaction, while Li was obviously easier to leach in tartaric acid. The leaching rate
reached 75.29% in the first 30 min. The fiing result of the control velocity equation of the
chemical reaction was poor, which was not suitable for Li leaching in a tartaric acid sys-
tem.
0 50 100 150 200 250 300 350 400
20
40
60
80
100
Leaching rate(%)
t(min)
(a)
Li
Co
0 50 100 150 200 250 300
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
y=0.20294+0.00146t
R
2
=0.7692
(b)
Li
Co
1-3(1-X)
2/3
+2(1-X)
t(min)
y=0.01939+0.00186t
R
2
=0.99027
Figure 2. Leaching performance curves under tartaric acid (a) change curves with time; (b) dynamic
fiing curves.
Cobalt trivalent could not be directly dissolved by acid, so hydrogen peroxide was
used as a reducing agent to make cobalt oxide accept an electron of hydrogen peroxide,
then was dissolved by citric acid to form cobalt citrate. The leaching reaction is described
by Equation (1) [18].
2HCit(aq)+2LiCoO
(s) + HO(aq) = 2Li (aq) + 2Cit (aq) + 2Co (aq) + 4HO + O(g) (1)
When tartaric acid was used as the leaching solution, it also played a reducing role.
So, the involvement of hydrogen peroxide was not required. The reaction with lithium
cobaltate is described by Equation (2) [7].
LiCoO+ 3C
HO→Co
(CHO)+LiC
HO+ 2H
O (2)
Figure 2.
Leaching performance curves under tartaric acid (
a
) change curves with time; (
b
) dynamic
fitting curves.
Cobalt trivalent could not be directly dissolved by acid, so hydrogen peroxide was
used as a reducing agent to make cobalt oxide accept an electron of hydrogen peroxide,
then was dissolved by citric acid to form cobalt citrate. The leaching reaction is described
by Equation (1) [18].
2H3Cit(aq)+2LiCoO2(s)+H2O2(aq)=2Li+(aq)+2Cit3−(aq)+2Co2+(aq)+4H2O+O2(g)(1)
When tartaric acid was used as the leaching solution, it also played a reducing role.
So, the involvement of hydrogen peroxide was not required. The reaction with lithium
cobaltate is described by Equation (2) [7].
LiCoO2+3C4H6O6→Co(C4H5O6)2+LiC4H5O6+2H2O (2)
According to the chemical equation, the complete reaction between citric acid and
lithium cobaltate should be 1:1, but in the actual process, only excessive citric acid could
ensure the complete leaching of metal in lithium cobaltate material. As shown in Figure 3a,

Molecules 2023,28, 3737 4 of 15
when the ratio of solid to liquid in the reaction system was low, the cobalt lithium element
could be well leached. However, when the ratio of solid to liquid increased, it became
difficult for the products on the surface of the material to diffuse into the solution with
higher concentration, so the efficiency decreased. According to the influence of the solid–
liquid ratio on the leaching rate, the best solid–liquid ratio was 10 g·1 L−1.
Molecules 2023, 28, x FOR PEER REVIEW 4 of 16
According to the chemical equation, the complete reaction between citric acid and lith-
ium cobaltate should be 1:1, but in the actual process, only excessive citric acid could ensure
the complete leaching of metal in lithium cobaltate material. As shown in Figure 3a, when
the ratio of solid to liquid in the reaction system was low, the cobalt lithium element could
be well leached. However, when the ratio of solid to liquid increased, it became difficult for
the products on the surface of the material to diffuse into the solution with higher concen-
tration, so the efficiency decreased. According to the influence of the solid–liquid ratio on
the leaching rate, the best solid–liquid ratio was 10 g·1 L−1.
12345
50
60
70
80
90
100
Leaching rate (%)
Solid-liquid ratio (g/100mL)
Co
Li
(a)
12345
50
60
70
80
90
100
Leaching rate (%)
Solid-liquid ratio (g/100mL)
Co
Li
(b)
Figure 3. Influence of solid–liquid ratio on leaching rate (a) citric acid; (b) tartaric acid.
In the same way, an excessive ratio of tartaric acid to lithium cobaltate could be used
for complete leaching. When the solid–liquid ratio was 20 g·1 L−1, the best leaching rates
of Co and Li were the highest, reaching 91.86% and 93.66%, respectively, which was sim-
ilar to the citric acid system. When the solid–liquid ratio increased, there was also a decline
in the tartaric acid system. When the solid–liquid ratio increased, it became difficult for
the products on the material surface to diffuse to the solution with a larger concentration.
The cathode material powder before and after leaching was characterized by XRD,
as shown in Figure 4. The main composition of the cathode material in the figure was
LiCoO2 (PDF#75-0532). Compared with the XRD paern of the leached material, it was
obvious that the characteristic diffraction peaks belonging to LiCoO2 basically disap-
peared, and only some stray peaks with poor matching degree were left, indicating that
the structure of lithium cobaltate was destroyed in the leaching process, and the valuable
metal was fully dissolved in the leaching solution.
0 102030405060708090
10 20 30 40 50 60 70 80 90
0
20
40
60
80
100
Intensity/a.u.
2θ/(
o
)
Intensity/a.u.
2θ/(
o
)
The anode material
Anode material after leaching
XRD standard card of LiCoO
2
(PDF#75-0532)
Figure 4. XRD images of the cathode material before and after leaching.
Figure 3. Influence of solid–liquid ratio on leaching rate (a) citric acid; (b) tartaric acid.
In the same way, an excessive ratio of tartaric acid to lithium cobaltate could be used
for complete leaching. When the solid–liquid ratio was 20 g
·
1 L
−1
, the best leaching rates
of Co and Li were the highest, reaching 91.86% and 93.66%, respectively, which was similar
to the citric acid system. When the solid–liquid ratio increased, there was also a decline in
the tartaric acid system. When the solid–liquid ratio increased, it became difficult for the
products on the material surface to diffuse to the solution with a larger concentration.
The cathode material powder before and after leaching was characterized by XRD, as
shown in Figure 4. The main composition of the cathode material in the figure was LiCoO
2
(PDF#75-0532). Compared with the XRD pattern of the leached material, it was obvious
that the characteristic diffraction peaks belonging to LiCoO
2
basically disappeared, and
only some stray peaks with poor matching degree were left, indicating that the structure of
lithium cobaltate was destroyed in the leaching process, and the valuable metal was fully
dissolved in the leaching solution.
Molecules 2023, 28, x FOR PEER REVIEW 4 of 16
According to the chemical equation, the complete reaction between citric acid and lith-
ium cobaltate should be 1:1, but in the actual process, only excessive citric acid could ensure
the complete leaching of metal in lithium cobaltate material. As shown in Figure 3a, when
the ratio of solid to liquid in the reaction system was low, the cobalt lithium element could
be well leached. However, when the ratio of solid to liquid increased, it became difficult for
the products on the surface of the material to diffuse into the solution with higher concen-
tration, so the efficiency decreased. According to the influence of the solid–liquid ratio on
the leaching rate, the best solid–liquid ratio was 10 g·1 L−1.
12345
50
60
70
80
90
100
Leaching rate (%)
Solid-liquid ratio (g/100mL)
Co
Li
(a)
12345
50
60
70
80
90
100
Leaching rate (%)
Solid-liquid ratio (g/100mL)
Co
Li
(b)
Figure 3. Influence of solid–liquid ratio on leaching rate (a) citric acid; (b) tartaric acid.
In the same way, an excessive ratio of tartaric acid to lithium cobaltate could be used
for complete leaching. When the solid–liquid ratio was 20 g·1 L−1, the best leaching rates
of Co and Li were the highest, reaching 91.86% and 93.66%, respectively, which was sim-
ilar to the citric acid system. When the solid–liquid ratio increased, there was also a decline
in the tartaric acid system. When the solid–liquid ratio increased, it became difficult for
the products on the material surface to diffuse to the solution with a larger concentration.
The cathode material powder before and after leaching was characterized by XRD,
as shown in Figure 4. The main composition of the cathode material in the figure was
LiCoO2 (PDF#75-0532). Compared with the XRD paern of the leached material, it was
obvious that the characteristic diffraction peaks belonging to LiCoO2 basically disap-
peared, and only some stray peaks with poor matching degree were left, indicating that
the structure of lithium cobaltate was destroyed in the leaching process, and the valuable
metal was fully dissolved in the leaching solution.
0 102030405060708090
10 20 30 40 50 60 70 80 90
0
20
40
60
80
100
Intensity/a.u.
2θ/(
o
)
Intensity/a.u.
2θ/(
o
)
The anode material
Anode material after leaching
XRD standard card of LiCoO
2
(PDF#75-0532)
Figure 4. XRD images of the cathode material before and after leaching.
Figure 4. XRD images of the cathode material before and after leaching.

Molecules 2023,28, 3737 5 of 15
2.2. Preparation of Lithium Ion Sieve
The precursor of the Li
1.6
Mn
1.6
O
4
lithium-ion sieve was difficult to obtain by direct
treatment of lithium manganese compounds with a traditional solid phase reaction in one
step, so in this paper, a Li
1.6
Mn
1.6
O
4
lithium-ion sieve was prepared step-by-step using the
sol–gel method, hydrothermal method and solid phase calcination method. The sol–gel
method was used to first add LiOH aqueous solution to MnCl
2
solution; at this time, Mn
2+
would generate Mn(OH)
2
precipitation under alkaline conditions. Then, 30% H
2
O
2
was
slowly added, making H
2
O
2
and Mn(OH)
2
react to
γ
-MnOOH; dropping speed was very
important. If H
2
O
2
drops were added too quickly, part of Mn(OH)
2
precipitates would be
oxidized directly into MnO
2
, making
γ
-MnOOH precipitates contain a large amount of
MnO2, thus increasing the impurity of the target product.
The
γ
-MnOOH reacted to the lithium source of the above reactor by hydrothermal
treatment, then formed LiMnO
2
. Finally, the precursor of Li
1.6
Mn
1.6
O
4
was obtained by the
calcination of LiMnO2[19].
The precursor of the lithium-ion screen needed elution before adsorption. Firstly, Li+
was stripped from the precursor by the eluent, that is, through the H–Li-ion exchange
process. A void filled with H
+
was formed in the material, and the eluted void played a
role in screening and memory of Li
+
. The XRD of the prepared lithium-ion before and after
sieve elution is shown in Figure 5. As seen from Figure 5, the diffraction peaks appeared at
2
θ
= 19.1
◦
, 37.0
◦
, 45.0
◦
and 65.5
◦
, respectively, corresponding to the (111), (311), (400) and
(440) crystal planes, and there were basically no mixed peaks. The XRD characteristic peaks
did not change significantly before and after pickling, except that the cells would become
slightly smaller. The reason for this was that during the elution, H
+
was the only ion
exchanged with Li
+
in Li
1.6
Mn
1.6
O
4
, so its basic structure did not change [
20
]. Additionally,
the XRD results could prove that our prepared Li1.6Mn1.6O4was highly stable and would
not be destroyed by acid washing.
Molecules 2023, 28, x FOR PEER REVIEW 5 of 16
2.2. Preparation of Lithium Ion Sieve
The precursor of the Li1.6Mn1.6O4 lithium-ion sieve was difficult to obtain by direct
treatment of lithium manganese compounds with a traditional solid phase reaction in one
step, so in this paper, a Li1.6Mn1.6O4 lithium-ion sieve was prepared step-by-step using the
sol–gel method, hydrothermal method and solid phase calcination method. The sol–gel
method was used to first add LiOH aqueous solution to MnCl2 solution; at this time, Mn2+
would generate Mn(OH)2 precipitation under alkaline conditions. Then, 30% H2O2 was
slowly added, making H2O2 and Mn(OH)2 react to γ-MnOOH; dropping speed was very
important. If H2O2 drops were added too quickly, part of Mn(OH)2 precipitates would be
oxidized directly into MnO2, making γ-MnOOH precipitates contain a large amount of
MnO2, thus increasing the impurity of the target product.
The γ-MnOOH reacted to the lithium source of the above reactor by hydrothermal
treatment, then formed LiMnO2. Finally, the precursor of Li1.6Mn1.6O4 was obtained by the
calcination of LiMnO2 [19].
The precursor of the lithium-ion screen needed elution before adsorption. Firstly, Li+
was stripped from the precursor by the eluent, that is, through the H–Li-ion exchange
process. A void filled with H+ was formed in the material, and the eluted void played a
role in screening and memory of Li+. The XRD of the prepared lithium-ion before and after
sieve elution is shown in Figure 5. As seen from Figure 5, the diffraction peaks appeared
at 2θ = 19.1°, 37.0°, 45.0° and 65.5°, respectively, corresponding to the (111), (311), (400)
and (440) crystal planes, and there were basically no mixed peaks. The XRD characteristic
peaks did not change significantly before and after pickling, except that the cells would
become slightly smaller. The reason for this was that during the elution, H+ was the only
ion exchanged with Li+ in Li1.6Mn1.6O4, so its basic structure did not change [20]. Addition-
ally, the XRD results could prove that our prepared Li1.6Mn1.6O4 was highly stable and
would not be destroyed by acid washing.
0 102030405060708090
0
2000
4000
6000
8000
10000
400311
Intensity/a.u.
2θ/(
o
)
Pickling before
After pickling
111
440
10,000
Figure 5. XRD paern of lithium-ion sieve before and after pickling.
The infrared spectra of the above materials before and after adsorption are shown in
Figure 6. It can be seen that the absorption peaks at 3424.7 cm−1 and 1593.1 cm−1 are related
to the vibration peaks of O-H in the adsorbed water. The absorption peak at 1112.7 cm−1 is
the stretching vibration peak of the Li–O bond. The absorption peaks at 619.1 cm−1 and 528.2
cm−1 are related to Mn(IV)-O and Mn(III)-O vibration peaks, respectively [21,22]. By com-
parison, no matter whether before the pickling, before the adsorption or after the adsorption
of Li+ ions, the infrared spectrum of Li1.6Mn1.6O4 had almost no change, indicating that the
adsorption mechanism of Li1.6Mn1.6O4 did not contain chemical bond breaking; its main
Figure 5. XRD pattern of lithium-ion sieve before and after pickling.
The infrared spectra of the above materials before and after adsorption are shown in
Figure 6. It can be seen that the absorption peaks at 3424.7 cm
−1
and 1593.1 cm
−1
are related
to the vibration peaks of O-H in the adsorbed water. The absorption peak at 1112.7 cm
−1
is
the stretching vibration peak of the Li–O bond. The absorption peaks at 619.1 cm
−1
and
528.2 cm
−1
are related to Mn(IV)-O and Mn(III)-O vibration peaks, respectively [
21
,
22
].
By comparison, no matter whether before the pickling, before the adsorption or after
the adsorption of Li
+
ions, the infrared spectrum of Li
1.6
Mn
1.6
O
4
had almost no change,
indicating that the adsorption mechanism of Li
1.6
Mn
1.6
O
4
did not contain chemical bond
breaking; its main mechanism is generally believed to be ion exchange [
23
] and surface
disproportionation [24].

Molecules 2023,28, 3737 6 of 15
Molecules 2023, 28, x FOR PEER REVIEW 6 of 16
mechanism is generally believed to be ion exchange [23] and surface disproportionation
[24].
5001000150020002500300035004000
528.2
619.1
1112.7
1593.1
3424.7
Wavenumber(cm
−1
)
After adsorption
Before adsorption
Pickling before
Figure 6. Infrared characterization of lithium-ion sieves before acid pickling and before and after
adsorption.
2.3. Lithium Recovery
2.3.1. Pretreatment of Eluent and Selection of Optimum Factors
The main metal ion components of the leaching solution were Co2+ and Li+, and the
detected contents were 5.862 g·L−1 and 0.775 g·L−1, respectively, except for a very small
amount of Fe2+, Ni2+ and Al3+. The solution was acidic; however, the acidic condition was
not conducive to the adsorption of the lithium-ion sieve. In order to eliminate the interfer-
ence of other factors, a control group was set up in the experiment, and 2 g of lithium-ion
sieve powder was taken to adsorb 100 mL 10 g·L−1 lithium solution, as shown in Figure 7,
which showed the changes of ion sieve adsorption capacity under different pH conditions.
It can be seen that when the pH of the solution was low, the concentration of H+ in the
solution was high, resulting in the ion exchange being difficult to occur. On the contrary,
the lower concentration of H+ promoted ion exchange.
678910
20
25
30
35
40
45
50
Adsorption capacity(mg/g)
pH
Effect of pH
0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0
35
40
45
50
55
60
65
70
Effect of solution concentration
Adsorption capacity(mg/g)
Solution concentration(g/L)
Figure 7. Adsorption capacity in relation to pH and solution concentration.
Figure 6.
Infrared characterization of lithium-ion sieves before acid pickling and before and after adsorption.
2.3. Lithium Recovery
2.3.1. Pretreatment of Eluent and Selection of Optimum Factors
The main metal ion components of the leaching solution were Co
2+
and Li
+
, and the
detected contents were 5.862 g
·
L
−1
and 0.775 g
·
L
−1
, respectively, except for a very small
amount of Fe
2+
, Ni
2+
and Al
3+
. The solution was acidic; however, the acidic condition
was not conducive to the adsorption of the lithium-ion sieve. In order to eliminate the
interference of other factors, a control group was set up in the experiment, and 2 g of
lithium-ion sieve powder was taken to adsorb 100 mL 10 g
·
L
−1
lithium solution, as shown
in Figure 7, which showed the changes of ion sieve adsorption capacity under different pH
conditions. It can be seen that when the pH of the solution was low, the concentration of
H
+
in the solution was high, resulting in the ion exchange being difficult to occur. On the
contrary, the lower concentration of H+promoted ion exchange.
Molecules 2023, 28, x FOR PEER REVIEW 6 of 16
mechanism is generally believed to be ion exchange [23] and surface disproportionation
[24].
5001000150020002500300035004000
528.2
619.1
1112.7
1593.1
3424.7
Wavenumber(cm
−1
)
After adsorption
Before adsorption
Pickling before
Figure 6. Infrared characterization of lithium-ion sieves before acid pickling and before and after
adsorption.
2.3. Lithium Recovery
2.3.1. Pretreatment of Eluent and Selection of Optimum Factors
The main metal ion components of the leaching solution were Co2+ and Li+, and the
detected contents were 5.862 g·L−1 and 0.775 g·L−1, respectively, except for a very small
amount of Fe2+, Ni2+ and Al3+. The solution was acidic; however, the acidic condition was
not conducive to the adsorption of the lithium-ion sieve. In order to eliminate the interfer-
ence of other factors, a control group was set up in the experiment, and 2 g of lithium-ion
sieve powder was taken to adsorb 100 mL 10 g·L−1 lithium solution, as shown in Figure 7,
which showed the changes of ion sieve adsorption capacity under different pH conditions.
It can be seen that when the pH of the solution was low, the concentration of H+ in the
solution was high, resulting in the ion exchange being difficult to occur. On the contrary,
the lower concentration of H+ promoted ion exchange.
678910
20
25
30
35
40
45
50
Adsorption capacity(mg/g)
pH
Effect of pH
0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0
35
40
45
50
55
60
65
70
Effect of solution concentration
Adsorption capacity(mg/g)
Solution concentration(g/L)
Figure 7. Adsorption capacity in relation to pH and solution concentration.
Figure 7. Adsorption capacity in relation to pH and solution concentration.
Therefore, the pH of the leaching solution was adjusted by adding an appropriate
amount of ammonia water, and the pH of the leaching solution was 9.5 when Fe(OH)
2
was completely precipitated (the residual ion concentration was less than 10
−6
mol
·
L
−1
).
Al(OH)
3
is 6, Ni(OH)
2
is 8.4 and Co(OH)
2
began to precipitate above 9 [
25
]. Therefore,
the pH was adjusted to about 9–10, and impurity ions in the leaching solution could

Molecules 2023,28, 3737 7 of 15
hardly be detected by filtering the sediment. The contents of Co
2+
and Li
+
in the leaching
solution were 5.316 g
·
L
−1
and 0.739 g
·
L
−1
, respectively. The concentration of the solution
was closely related to the adsorption capacity of the adsorbent. Figure 7showed that
with the increase in concentration, the sieve adsorption capacity of lithium-ion increased
from 19.97 mg·g−1to 34.76 mg·g−1. This was because the high concentration of Li+could
promote the H–Li ion exchange process, thus increasing the adsorption capacity. Therefore,
the leaching solution could be heated to concentrate, in order to facilitate the adsorption of
the lithium-ion sieve.
2.3.2. Lithium-Ion Sieve Adsorption
An amount of 100 mL concentrated leachate, including the concentration of Li
+
of
about 1.8 to 2 g L
−1
, was added to 2 g lithium-ion sieve adsorption. Its adsorption per-
formance is shown in Figure 8. It can be seen that the lithium-ion sieve and the leaching
solution reached the adsorption equilibrium for about 48 h. Under this condition, the maxi-
mum adsorption capacity was 38.05 mg
·
g
−1
, accounting for 98.21% of the total Li
+
content.
Furthermore, the quasi-first-order adsorption kinetics model and the quasi-second-order
adsorption kinetics model were fitted. The fitting results showed that the first-order kinetic
decision coefficient R
2
of the lithium-ion screen was 0.93542, lower than the second-order
kinetic 0.99956, indicating that the adsorption of the lithium-ion screen on the leach solution
was more consistent with chemical adsorption. Subsequently, the concentration of Co
2+
was detected. There was a certain decrease in Co
2+
concentration, about 0.29% of the total,
which was almost negligible. The reason was mainly due to the lithium ion screen having a
superior selective adsorption capacity.
Molecules 2023, 28, x FOR PEER REVIEW 7 of 16
Therefore, the pH of the leaching solution was adjusted by adding an appropriate
amount of ammonia water, and the pH of the leaching solution was 9.5 when Fe(OH)2 was
completely precipitated (the residual ion concentration was less than 10−6 mol·L−1).
Al(OH)3 is 6, Ni(OH)2 is 8.4 and Co(OH)2 began to precipitate above 9 [25]. Therefore, the
pH was adjusted to about 9–10, and impurity ions in the leaching solution could hardly
be detected by filtering the sediment. The contents of Co2+ and Li+ in the leaching solution
were 5.316 g·L−1 and 0.739 g·L−1, respectively. The concentration of the solution was closely
related to the adsorption capacity of the adsorbent. Figure 7 showed that with the increase
in concentration, the sieve adsorption capacity of lithium-ion increased from 19.97 mg⋅g−1
to 34.76 mg·g−1. This was because the high concentration of Li+ could promote the H–Li
ion exchange process, thus increasing the adsorption capacity. Therefore, the leaching so-
lution could be heated to concentrate, in order to facilitate the adsorption of the lithium-
ion sieve.
2.3.2. Lithium-Ion Sieve Adsorption
An amount of 100 mL concentrated leachate, including the concentration of Li+ of
about 1.8 to 2 g L−1, was added to 2 g lithium-ion sieve adsorption. Its adsorption perfor-
mance is shown in Figure 8. It can be seen that the lithium-ion sieve and the leaching
solution reached the adsorption equilibrium for about 48 h. Under this condition, the max-
imum adsorption capacity was 38.05 mg·g−1, accounting for 98.21% of the total Li+ content.
Furthermore, the quasi-first-order adsorption kinetics model and the quasi-second-order
adsorption kinetics model were fied. The fiing results showed that the first-order ki-
netic decision coefficient R2 of the lithium-ion screen was 0.93542, lower than the second-
order kinetic 0.99956, indicating that the adsorption of the lithium-ion screen on the leach
solution was more consistent with chemical adsorption. Subsequently, the concentration
of Co2+ was detected. There was a certain decrease in Co2+ concentration, about 0.29% of
the total, which was almost negligible. The reason was mainly due to the lithium ion
screen having a superior selective adsorption capacity.
0 5 10 15 20 25 30 35 40 45 50 55
0
5
10
15
20
25
30
35
40
Adsorption capacity
Adsorption capacity(mg/g)
t/h
-3
-2
-1
0
1
2
3
4
ln(Q
E
-Q
T
)
y=0.01431+0.02591x
R
2
=0.99956
ln(Q
E
-Q
T
)
y=2.9163-0.25705x
R
2
=0.93542
-0.2
0.0
0.2
0.4
0.6
0.8
1.0
1.2
1.4
t/Q
T
t/Q
T
Figure 8. Adsorption performance curves.
Repeated elution experiments of the lithium-ion sieve were carried out, and the re-
sults are shown in Table 2. It can be seen that the adsorption capacity of the lithium-ion
screen decreased significantly in the first three cycles at 8.4625%, 10% and 6%, and the
change slowed down to 3.82% from the fourth to the fifth time, and the final performance
of the twentieth time decreased to only 10.46% compared with the tenth time. The results
showed that the repeated adsorption performance of the lithium-ion screen was good.
Figure 8. Adsorption performance curves.
Repeated elution experiments of the lithium-ion sieve were carried out, and the results
are shown in Table 2. It can be seen that the adsorption capacity of the lithium-ion screen
decreased significantly in the first three cycles at 8.4625%, 10% and 6%, and the change
slowed down to 3.82% from the fourth to the fifth time, and the final performance of
the twentieth time decreased to only 10.46% compared with the tenth time. The results
showed that the repeated adsorption performance of the lithium-ion screen was good.
When HCl was used as the eluent, the dissolution rate of lithium was about 91% in many
cycle experiments, which could achieve a good dissolution effect.

Molecules 2023,28, 3737 8 of 15
Table 2. Results of cyclic adsorption experiment.
Number of Cycles Adsorption Capacity (mg·g−1) The Dissolution Rate of Lithium (%)
1 38.05 91.54
2 34.83 91.49
3 31.34 90.81
4 29.29 91.12
5 28.17 91.29
10 21.79 91.38
20 19.51 89.88
2.3.3. Synthesis of Lithium Chloride
The adsorbed lithium-ion sieve powder was desorbed using 0.5 mol
·
L
−1
HCl solution.
However, due to the disproportionation of Mn on the surface of the lithium-ion sieve during
the elution process, a small amount of Mn
2+
existed in the solution. The removal effect
could be achieved by dropping an appropriate amount of H
2
O
2
and diluting ammonia
water. In addition, the Co
2+
adsorbed during the adsorption process was not detected
in the solution and showed that a trace amount of Co
2+
on the adsorbent did not eluate
during the elution process. Through the above operation, the purer LiCl solution was
finally obtained.
The obtained pure LiCl solution was heated, concentrated, cooled and crystallized to
obtain some white crystals. Phase analysis was conducted by XRD, as shown in Figure 9. In
Figure 9, the diffraction peaks were sharp and basically smooth, and there were almost no
mixed peaks. The position of the diffraction peaks was consistent with that of the standard
card (PDF#74-1972). It indicated that the obtained product was crystalline compared to
pure lithium chloride crystal, and the peak at about 2
θ=
33
◦
was caused by LiCl
·
H
2
O. Its
purity reached about 93% (determined by the Li+content in solids).
Molecules 2023, 28, x FOR PEER REVIEW 9 of 16
Table 2. Results of cyclic adsorption experiment.
Number of Cycles Adsorption Capacity (mg·g−1) The Dissolution Rate of Lithium (%)
1 38.05 91.54
2 34.83 91.49
3 31.34 90.81
4 29.29 91.12
5 28.17 91.29
10 21.79 91.38
20 19.51 89.88
2.3.3. Synthesis of Lithium Chloride
The adsorbed lithium-ion sieve powder was desorbed using 0.5 mol·L−1 HCl solution.
However, due to the disproportionation of Mn on the surface of the lithium-ion sieve during
the elution process, a small amount of Mn2+ existed in the solution. The removal effect could
be achieved by dropping an appropriate amount of H2O2 and diluting ammonia water. In
addition, the Co2+ adsorbed during the adsorption process was not detected in the solution
and showed that a trace amount of Co2+ on the adsorbent did not eluate during the elution
process. Through the above operation, the purer LiCl solution was finally obtained.
The obtained pure LiCl solution was heated, concentrated, cooled and crystallized to
obtain some white crystals. Phase analysis was conducted by XRD, as shown in Figure 9.
In Figure 9, the diffraction peaks were sharp and basically smooth, and there were almost
no mixed peaks. The position of the diffraction peaks was consistent with that of the stand-
ard card (PDF#74-1972). It indicated that the obtained product was crystalline compared
to pure lithium chloride crystal, and the peak at about 2𝜃 = 33° was caused by LiCl·H2O.
Its purity reached about 93% (determined by the Li+ content in solids).
0 102030405060708090
Intensity/a.u.
2θ/(
o
)
LiCl(self-prepared)
PDF#74-1972（LiCl）
Figure 9. XRD paern of lithium chloride crystal.
2.3.4. Synthesis of Lithium Carbonate
Although the synthesis method of lithium chloride crystal was simple, the conversion
rate in the process was low, resulting in a lot of waste. Therefore, the lithium recovery
method could be designed through the following equation:
2Li+Na
CO=2Na
+Li
CO↓ (3)
According to Equation (3), only by controlling the dosage of Na2CO3 and the reaction
temperature can Li2CO3 precipitation be generated, and then the reaction suspension can
be filtered, washed and dried to obtain Li2CO3 crystals.
Figure 9. XRD pattern of lithium chloride crystal.
2.3.4. Synthesis of Lithium Carbonate
Although the synthesis method of lithium chloride crystal was simple, the conversion
rate in the process was low, resulting in a lot of waste. Therefore, the lithium recovery
method could be designed through the following equation:
2Li++Na2CO3=2Na++Li2CO3↓(3)
According to Equation (3), only by controlling the dosage of Na
2
CO
3
and the reaction
temperature can Li
2
CO
3
precipitation be generated, and then the reaction suspension can
be filtered, washed and dried to obtain Li2CO3crystals.

Molecules 2023,28, 3737 9 of 15
1. The influence of the pH;
As 0.5 mol
·
L
−1
HCl was used as eluent, the recovered lithium solution became acidic
(pH 0~1). The formation pH of Li
2
CO
3
crystal was determined by many experiments to
be about 8~9, so the solution pH could be adjusted by adding an appropriate amount of
NaOH before the reaction to promote the formation of Li
2
CO
3
crystals. Figure 10 shows
a comparison of the quality and purity of the final product when 10 mL of a 30 g
·
L
−1
Na
2
CO
3
solution was added to 100 mL of leachate at different pH values (5, 6, 7, 8, 9, 10, 11
and 12). It can be seen that as the pH of the initial solution rose, the crystal yield of Li
2
CO
3
gradually increased because a small amount of Na
2
CO
3
reacted with H
+
when the solution
was acidic. The purity of the product tended to rise at the beginning. The closer the pH was
to 9, the easier Na
2
CO
3
would react to form crystals. When the pH exceeded 9, the purity
began to decline. Therefore, the pH of the recovered lithium solution could be adjusted to
7~9, which is conducive to the generation of high quality Li2CO3crystals.
Molecules 2023, 28, x FOR PEER REVIEW 10 of 16
1. The influence of the pH;
As 0.5 mol·L−1 HCl was used as eluent, the recovered lithium solution became acidic
(pH 0~1). The formation pH of Li2CO3 crystal was determined by many experiments to be
about 8~9, so the solution pH could be adjusted by adding an appropriate amount of
NaOH before the reaction to promote the formation of Li2CO3 crystals. Figure 10 shows a
comparison of the quality and purity of the final product when 10 mL of a 30 g·L−1 Na2CO3
solution was added to 100 mL of leachate at different pH values (5, 6, 7, 8, 9, 10, 11 and
12). It can be seen that as the pH of the initial solution rose, the crystal yield of Li2CO3
gradually increased because a small amount of Na2CO3 reacted with H+ when the solution
was acidic. The purity of the product tended to rise at the beginning. The closer the pH
was to 9, the easier Na2CO3 would react to form crystals. When the pH exceeded 9, the
purity began to decline. Therefore, the pH of the recovered lithium solution could be ad-
justed to 7~9, which is conducive to the generation of high quality Li2CO3 crystals.
3456789101112
0.29
0.30
0.31
0.32
0.33
0.34
Li
2
CO
3
Quantity of product
Purity of product
pH
Quantity(g)
90
91
92
93
94
95
96
Purity(%)
Figure 10. Effect of pH on quantity and purity of output.
2. The influence of the dosage of Na2CO3;
The dosage of Na2CO3 was changed, and different volumes of 30 g·L−1 Na2CO3
solution (5, 7, 10, 15 and 20 mL) were added into 100 mL of eluent, respectively. Other
parameters were fixed unchanged, and the experiment was conducted; the results are shown
in Figure 11. As seen in Figure 11, with the increase in the dosage of Na2CO3, more crystals of
Li2CO3 can be obtained. When the dosage of Na2CO3 was greater than 15 mL, without regard
to the loss in the filtration process, the mass increase of the obtained Li2CO3 crystals was small
because the Li in the eluent had almost completely formed precipitate. However, the crystal-
lization purity of Li2CO3 tended to decrease slightly with the increase of the dosage of
Na2CO3 because a small amount of sodium carbonate was wrapped in the Li2CO3 crystal
during the crystallization process [26].
Figure 10. Effect of pH on quantity and purity of output.
2. The influence of the dosage of Na2CO3;
The dosage of Na
2
CO
3
was changed, and different volumes of 30 g
·
L
−1
Na
2
CO
3
solution (5, 7, 10, 15 and 20 mL) were added into 100 mL of eluent, respectively. Other
parameters were fixed unchanged, and the experiment was conducted; the results are
shown in Figure 11. As seen in Figure 11, with the increase in the dosage of Na
2
CO
3
, more
crystals of Li
2
CO
3
can be obtained. When the dosage of Na
2
CO
3
was greater than 15 mL,
without regard to the loss in the filtration process, the mass increase of the obtained Li
2
CO
3
crystals was small because the Li in the eluent had almost completely formed precipitate.
However, the crystallization purity of Li
2
CO
3
tended to decrease slightly with the increase
of the dosage of Na
2
CO
3
because a small amount of sodium carbonate was wrapped in the
Li2CO3crystal during the crystallization process [26].
3. The influence of reaction temperature.
Temperature was also a key factor in the formation of Li
2
CO
3
precipitation. Other
conditions remained unchanged, and the temperature was controlled at 20, 30, 40, 60
and 80
◦
C for the experiments. The results are shown in Figure 12. It can be seen from
Figure 12 that the crystal weight of Li
2
CO
3
increased with the increase in temperature. The
solubility of Li
2
CO
3
decreased with the increase in temperature. According to the study of
Tao et al. [
26
], the Li
2
CO
3
aggregates generated at low temperature were mainly flake, while
the Li
2
CO
3
aggregates generated at high temperature were mainly rod-like. Compared
with rod-like aggregates, flake aggregates were more likely to contain impurities, which
was also the reason why the product purity was lower at low temperature.

Molecules 2023,28, 3737 10 of 15
Molecules 2023, 28, x FOR PEER REVIEW 11 of 16
4 6 8 10 12 14 16 18 20 22
0.29
0.30
0.31
0.32
0.33
0.34
0.35
Li
2
CO
3
Quantity of product
Purity of product
The dosage of Na
2
CO
3
(mL)
Quantity(g)
98.2
98.4
98.6
98.8
99.0
99.2
99.4
99.6
99.8
Purity(%)
Figure 11. Effect of dosage of Na2CO3 on the quantity and purity of output.
3. The influence of reaction temperature.
Temperature was also a key factor in the formation of Li2CO3 precipitation. Other
conditions remained unchanged, and the temperature was controlled at 20, 30, 40, 60 and
80 °C for the experiments. The results are shown in Figure 12. It can be seen from Figure
12 that the crystal weight of Li2CO3 increased with the increase in temperature. The solu-
bility of Li2CO3 decreased with the increase in temperature. According to the study of Tao
et al. [26], the Li2CO3 aggregates generated at low temperature were mainly flake, while
the Li2CO3 aggregates generated at high temperature were mainly rod-like. Compared
with rod-like aggregates, flake aggregates were more likely to contain impurities, which
was also the reason why the product purity was lower at low temperature.
20 30 40 50 60 70 80
0.26
0.28
0.30
0.32
0.34
0.36
0.38
Li
2
CO
3
Quantity of product
Purity of product
Temperature(℃)
Quantity(g)
98.6
98.8
99.0
99.2
99.4
99.6
Purity(%)
Figure 12. Effect of temperature on quantity and purity of output.
For 100 mL of the treated lithium-containing solution (Li concentration is about 40~50
mg·L−1), considering the influence of the dosage of Na2CO3 and the reaction temperature,
using 15 mL Na2CO3 solution with a concentration of 30 g·L−1 and a control temperature
of 60 °C could make the production of Li2CO3 larger and the purity higher.
Figure 11. Effect of dosage of Na2CO3on the quantity and purity of output.
Molecules 2023, 28, x FOR PEER REVIEW 11 of 16
4 6 8 10 12 14 16 18 20 22
0.29
0.30
0.31
0.32
0.33
0.34
0.35
Li
2
CO
3
Quantity of product
Purity of product
The dosage of Na
2
CO
3
(mL)
Quantity(g)
98.2
98.4
98.6
98.8
99.0
99.2
99.4
99.6
99.8
Purity(%)
Figure 11. Effect of dosage of Na2CO3 on the quantity and purity of output.
3. The influence of reaction temperature.
Temperature was also a key factor in the formation of Li2CO3 precipitation. Other
conditions remained unchanged, and the temperature was controlled at 20, 30, 40, 60 and
80 °C for the experiments. The results are shown in Figure 12. It can be seen from Figure
12 that the crystal weight of Li2CO3 increased with the increase in temperature. The solu-
bility of Li2CO3 decreased with the increase in temperature. According to the study of Tao
et al. [26], the Li2CO3 aggregates generated at low temperature were mainly flake, while
the Li2CO3 aggregates generated at high temperature were mainly rod-like. Compared
with rod-like aggregates, flake aggregates were more likely to contain impurities, which
was also the reason why the product purity was lower at low temperature.
20 30 40 50 60 70 80
0.26
0.28
0.30
0.32
0.34
0.36
0.38
Li
2
CO
3
Quantity of product
Purity of product
Temperature(℃)
Quantity(g)
98.6
98.8
99.0
99.2
99.4
99.6
Purity(%)
Figure 12. Effect of temperature on quantity and purity of output.
For 100 mL of the treated lithium-containing solution (Li concentration is about 40~50
mg·L−1), considering the influence of the dosage of Na2CO3 and the reaction temperature,
using 15 mL Na2CO3 solution with a concentration of 30 g·L−1 and a control temperature
of 60 °C could make the production of Li2CO3 larger and the purity higher.
Figure 12. Effect of temperature on quantity and purity of output.
For 100 mL of the treated lithium-containing solution (Li concentration is about
40~50 mg·L−1
), considering the influence of the dosage of Na
2
CO
3
and the reaction tem-
perature, using 15 mL Na
2
CO
3
solution with a concentration of 30 g
·
L
−1
and a control
temperature of 60 ◦C could make the production of Li2CO3larger and the purity higher.
2.4. Recovery of Cobalt
Two basic routes for cobalt recovery were designed: one was to regenerate lithium
cobaltate by leaching solution after adsorption, and the other was to extract pure cobalt
chloride crystals by precipitation. The former avoided the use of a large number of chemical
reagents but added the steps of vacuum-drying and high-temperature calcination. The
latter was simpler to operate but used more chemical reagents.
2.4.1. Regeneration of Lithium Cobalt Oxide-Based Battery by Soft Chemical Method
The element concentration ratio was adjusted to 1:1 according to the concentration
of Co
2+
in the leaching solution. The corresponding lithium chloride crystals obtained in
the above steps were added to the leaching solution as lithium sources; then, the pH of the
leaching solution was adjusted to 7 with ammonia. Firstly, this was placed in a water bath

Molecules 2023,28, 3737 11 of 15
kettle at 80
◦
C, then magnetically stirred for about 2 h at 200 r
·
min
−1
until the formation
of sol. Subsequently, the sol formed a dry gel through vacuum drying for 24 h, and then
the calcination was carried out by a tubular furnace at 700
◦
C temperature to obtain the
regenerated LiCoO2crystal.
The calcination temperature had a great influence on the performance and structure
of the regenerated LiCoO
2
. In summary of the studies [
27
–
29
], when the temperature was
500
◦
C, the regenerated LiCoO
2
was mostly a spinel structure, and when it rose to between
700
◦
C and 900
◦
C, the regenerated LiCoO
2
was a hexagonal phase structure. With the
increase in temperature, the peak strength would gradually increase and the crystallinity
would increase. However, when the temperature exceeded 900
◦
C, lithium ions would
gradually come out and decompose to produce Co
3
O
4
. Therefore, this study chose the
calcination condition at 700
◦
C for 5 h to obtain regenerated LiCoO
2
particles. Phase
analysis by XRD showed that the position of the diffraction peak was consistent with that
of the standard card (PDF#74-1972), with the characteristic diffraction peaks of (003), (101)
and (104), except that there were two LiCoO
2
characteristic peaks of hexagonal layered
structures (018) and (110) [
28
,
29
], indicating that LiCoO
2
material with a better structure
was obtained using the sol–gel method (Figure 13). SEM images of the regenerated lithium
cobaltate are shown in Figure 14. Seen from Figure 14, the crystals of lithium cobaltate
were obvious.
Molecules 2023, 28, x FOR PEER REVIEW 12 of 16
2.4. Recovery of Cobalt
Two basic routes for cobalt recovery were designed: one was to regenerate lithium
cobaltate by leaching solution after adsorption, and the other was to extract pure cobalt
chloride crystals by precipitation. The former avoided the use of a large number of chem-
ical reagents but added the steps of vacuum-drying and high-temperature calcination. The
laer was simpler to operate but used more chemical reagents.
2.4.1. Regeneration of Lithium Cobalt Oxide-Based Baery by Soft Chemical Method
The element concentration ratio was adjusted to 1:1 according to the concentration of
Co2+ in the leaching solution. The corresponding lithium chloride crystals obtained in the
above steps were added to the leaching solution as lithium sources; then, the pH of the
leaching solution was adjusted to 7 with ammonia. Firstly, this was placed in a water bath
kettle at 80 °C, then magnetically stirred for about 2 h at 200 r·min−1 until the formation of
sol. Subsequently, the sol formed a dry gel through vacuum drying for 24 h, and then the
calcination was carried out by a tubular furnace at 700 °C temperature to obtain the regen-
erated LiCoO2 crystal.
The calcination temperature had a great influence on the performance and structure of
the regenerated LiCoO2. In summary of the studies [27–29], when the temperature was 500
°C, the regenerated LiCoO2 was mostly a spinel structure, and when it rose to between 700
°C and 900 °C, the regenerated LiCoO2 was a hexagonal phase structure. With the increase
in temperature, the peak strength would gradually increase and the crystallinity would
increase. However, when the temperature exceeded 900 °C, lithium ions would gradually
come out and decompose to produce Co3O4. Therefore, this study chose the calcination
condition at 700 °C for 5 h to obtain regenerated LiCoO2 particles. Phase analysis by XRD
showed that the position of the diffraction peak was consistent with that of the standard
card (PDF#74-1972), with the characteristic diffraction peaks of (003), (101) and (104), ex-
cept that there were two LiCoO2 characteristic peaks of hexagonal layered structures (018)
and (110) [28,29], indicating that LiCoO2 material with a beer structure was obtained us-
ing the sol–gel method (Figure 13). SEM images of the regenerated lithium cobaltate are
shown in Figure 14. Seen from Figure 14, the crystals of lithium cobaltate were obvious.
0 102030405060708090
110
018
104
101
Intensity/a.u.
2θ/(
o
)
Regeneration of LiCoO
2
PDF#75-0532(LiCoO
2
)
003
Figure 13. XRD paern of regenerated lithium cobaltate.
Figure 13. XRD pattern of regenerated lithium cobaltate.
Molecules 2023, 28, x FOR PEER REVIEW 13 of 16
Figure 14. SEM images of regenerated lithium cobaltate.
2.4.2. Synthesis of Cobalt Chloride
Due to the excessive amount of citric acid added in the leaching experiment, large
amounts of citric acid crystals might be generated if crystallization is directly evaporated.
Therefore, a solution with a concentration of 5 mol·L−1 NaOH was slowly dripped in and
filtered when the solution produced a large amount of blue–gray precipitate, and then a
precipitate of coarse cobalt hydroxide was obtained. Hydrochloric acid was used for dis-
solution, then filtered, evaporated and crystallized for the remaining solution. Finally, rel-
atively pure CoCl2 was obtained. As shown in Figure 15, the position of the diffraction
peak of the prepared CoCl2 was consistent with that of standard card (PDF#25-0242), in-
dicating that its crystallinity was good, and the crystal purity was about 87.9% after mul-
tiple tests.
0 102030405060708090
Intensity/a.u.
2θ/(
o
)
CoCl
2
(self-prepared)
PDF#25-0242(CoCl
2
)
Figure 15. XRD paern of cobalt crystallization.
3. Materials and Methods
3.1. Reagents and Instruments
Main reagents: lithium hydroxide monohydrate, citric acid, tartaric acid, hydrogen
peroxide (Shanghai Maclin Biochemical Technology Co., LTD., Shanghai, China), manga-
nese chloride tetrahydrate, sodium hydroxide, hydrochloric acid and ammonia (Si-
nopharm chemical reagent co., LTD ,Shanghai, China). In this study, the raw material was
the fine particles obtained by the discharge, disassembling, crushing and screening of a
waste lithium cobalt oxide-based baery whose surface binder and acetylene ash were
removed by calcination. The experimental water was deionized water.
Main instruments: UV-26001 UV-visible spectrophotometer (Shimadzu Company,
Kyoto, Japan), Nicolet 380 Fourier Transform infrared spectrometer (Thermo Fisher
Figure 14. SEM images of regenerated lithium cobaltate.
2.4.2. Synthesis of Cobalt Chloride
Due to the excessive amount of citric acid added in the leaching experiment, large
amounts of citric acid crystals might be generated if crystallization is directly evaporated.
Therefore, a solution with a concentration of 5 mol
·
L
−1
NaOH was slowly dripped in and

Molecules 2023,28, 3737 12 of 15
filtered when the solution produced a large amount of blue–gray precipitate, and then
a precipitate of coarse cobalt hydroxide was obtained. Hydrochloric acid was used for
dissolution, then filtered, evaporated and crystallized for the remaining solution. Finally,
relatively pure CoCl
2
was obtained. As shown in Figure 15, the position of the diffraction
peak of the prepared CoCl
2
was consistent with that of standard card (PDF#25-0242),
indicating that its crystallinity was good, and the crystal purity was about 87.9% after
multiple tests.
Molecules 2023, 28, x FOR PEER REVIEW 13 of 16
Figure 14. SEM images of regenerated lithium cobaltate.
2.4.2. Synthesis of Cobalt Chloride
Due to the excessive amount of citric acid added in the leaching experiment, large
amounts of citric acid crystals might be generated if crystallization is directly evaporated.
Therefore, a solution with a concentration of 5 mol·L−1 NaOH was slowly dripped in and
filtered when the solution produced a large amount of blue–gray precipitate, and then a
precipitate of coarse cobalt hydroxide was obtained. Hydrochloric acid was used for dis-
solution, then filtered, evaporated and crystallized for the remaining solution. Finally, rel-
atively pure CoCl2 was obtained. As shown in Figure 15, the position of the diffraction
peak of the prepared CoCl2 was consistent with that of standard card (PDF#25-0242), in-
dicating that its crystallinity was good, and the crystal purity was about 87.9% after mul-
tiple tests.
0 102030405060708090
Intensity/a.u.
2θ/(
o
)
CoCl
2
(self-prepared)
PDF#25-0242(CoCl
2
)
Figure 15. XRD paern of cobalt crystallization.
3. Materials and Methods
3.1. Reagents and Instruments
Main reagents: lithium hydroxide monohydrate, citric acid, tartaric acid, hydrogen
peroxide (Shanghai Maclin Biochemical Technology Co., LTD., Shanghai, China), manga-
nese chloride tetrahydrate, sodium hydroxide, hydrochloric acid and ammonia (Si-
nopharm chemical reagent co., LTD ,Shanghai, China). In this study, the raw material was
the fine particles obtained by the discharge, disassembling, crushing and screening of a
waste lithium cobalt oxide-based baery whose surface binder and acetylene ash were
removed by calcination. The experimental water was deionized water.
Main instruments: UV-26001 UV-visible spectrophotometer (Shimadzu Company,
Kyoto, Japan), Nicolet 380 Fourier Transform infrared spectrometer (Thermo Fisher
Figure 15. XRD pattern of cobalt crystallization.
3. Materials and Methods
3.1. Reagents and Instruments
Main reagents: lithium hydroxide monohydrate, citric acid, tartaric acid, hydrogen
peroxide (Shanghai Maclin Biochemical Technology Co., Ltd., Shanghai, China), manganese
chloride tetrahydrate, sodium hydroxide, hydrochloric acid and ammonia (Sinopharm
chemical reagent Co., Ltd., Shanghai, China). In this study, the raw material was the fine
particles obtained by the discharge, disassembling, crushing and screening of a waste
lithium cobalt oxide-based battery whose surface binder and acetylene ash were removed
by calcination. The experimental water was deionized water.
Main instruments: UV-26001 UV-visible spectrophotometer (Shimadzu Company, Ky-
oto, Japan), Nicolet 380 Fourier Transform infrared spectrometer (Thermo Fisher, Waltham,
MA, USA), Regulus8100 Scanning electron Microscope ( HITACHI, Ibaraki, Japan), X-ray
Photoelectron Spectrometer (Thermo Fischer, Waltham, MA, USA), Optima 8000 Induc-
tively Coupled Plasma Spectrometer (PERKINELMER, Waltham, MA, USA), Tube furnace.
3.2. Experimental Process
Leaching of valuable metals: Weighed 0.3 mol citric acid (tartaric acid) and dissolved
in 200 mL deionized water, followed by ultrasonic shock for 5 min. After that, appropriate
amount of hydrogen peroxide solution was added (tartaric acid does not require hydrogen
peroxide), weighed 2, 4, 6, 8 and 10 g of treated waste battery positive powder and added it
to the beaker. The beaker was placed in a water bath kettle at the temperature of 90
◦
C and
a rotating speed of 600 r
·
min
−1
. ICP-AES was used to determine the metal concentration in
the leaching solution, and the metal leaching rate was calculated as shown in Equation (4):
η=c×V
m×100% (4)

Molecules 2023,28, 3737 13 of 15
In the formula:
η
is metal leaching rate, %; cis the mass concentration of metal ions,
g/mL; Vis filtrate volume, mL; mis the theoretical mass of each metal in lithium cobaltate
powder, g.
Preparation of the adsorbent: Took 0.1 mol MnCl
2·
4H
2
O and dissolved it in 100 mL
deionized water, added a quantitative amount of LiOH
·
H
2
O (molar ratio of lithium to
manganese was 4:1) and mixed evenly. After mixing, slowly added appropriate amount of
30% H
2
O
2
, about 5 mL, continued to stir for 2 h, and then aged for 24 h. Transferred the
formed colloid to a 100 mL Teflon-lined high-pressure reactor by rinsing with supernatant,
then added partial supernatant to about 70% of the autoclave volume and heated at 120
◦
C
for 24 h. The material obtained was filtered, washed, dried, then ground with mortar and
calcined in Muffle furnace at 400
◦
C for 5 h to get the precursor of ion sieve. Weighed
10 g lithium-ion sieve precursor powder and placed it in 0.5 mol
·
L
−1
HCl solution, stirred
it with 200 r
·
min
−1
magnetic force for about 12 h, and then obtained lithium-ion sieve
adsorbent through filtration, washing and drying process.
Separation of metal ions: Put 2 g ion sieve powder into 100 mL leach solution, stirred
at about 20
◦
C with 200 r
·
min
−1
magnetic force for 12 h, sampled at selected time period,
measured Li
+
concentration, and drew saturation adsorption curve to study the adsorption
performance of lithium-ion sieve. Then, the solution was filtered, lithium-ion sieve powder
was collected, and the adsorbed lithium ion sieve was placed in 0.5 mol
·
L
−1
HCl solution
for desorption to obtain a higher purity lithium solution.
4. Conclusions
1.
The leaching of the cathode material of a lithium cobalt oxide-based battery with
citric acid and a hydrogen peroxide system was investigated. The leaching rates of
86.21% and 96.9% for Co and Li can be achieved at a reaction temperature of 90
◦
C,
stirring speed of 600 r
·
min
−1
and a solid–liquid ratio of 10 g
·
1 L
−1
. The determination
coefficients R
2
of Co and Li fitting curves were 0.95334 and 0.99447 according to the
chemical reaction control velocity equation, indicating that leaching was a chemical
reaction. In the tartaric acid system, the leaching rates of Co and Li were 90.34% and
92.47%. The leaching of Li in the tartaric acid system was rapid, reaching 75.29% in
the first 30 min;
2.
The prepared Li
1.6
Mn
1.6
O
4
lithium-ion screen was used to conduct the lithium sepa-
ration. The maximum adsorption capacity of the lithium screen was 38.05 mg
·
g
−1
,
and the dissolution rate of lithium was about 91%. Through elution, purification and
other steps, the adsorbed lithium was transformed into a lithium chloride solution.
There were two methods for lithium recovery: (1) Relatively pure lithium chloride
crystals can be obtained by direct concentration and crystallization, with a detected
purity up to 93%. (2) Li
2
CO
3
crystals were generated by adding Na
2
CO
3
, and the
purity of Li2CO3crystals can reach 99.59%;
3.
Two methods were used to conduct the tests for the recovery of cobalt: (1) CoCl
2
crystals were obtained by means of reprecipitation and recrystallization, and the
purity can reach 87.9%; (2) LiCoO
2
was directly regenerated using the sol–gel method
and lithium chloride was used as a lithium source. XRD characterization showed
that LiCoO
2
had good crystallinity. The two methods had different advantages and
disadvantages: the former was simpler to operate, and the latter used fewer chemical
reagents. The above two methods could provide ideas for recycling waste cathode
materials of lithium cobalt oxide-based batteries.
Author Contributions:
H.W.: Conceptualization, Methodology, Data Curation, Writing—Original
draft preparation; G.C.: Supervision, Reviewing and Editing, Funding acquisition; L.M., G.W.,
X.D. and R.C.: Resources, Formal analysis, Project administration, Participation in discussions and
suggestions. All authors have read and agreed to the published version of the manuscript.

Molecules 2023,28, 3737 14 of 15
Funding:
This research was funded by [the Anhui Provincial Key Research and Development Plan]
grant number [2022107020025], and [key projects of natural science of universities in Anhui Province]
grant number [KJ2021A0617], and [the Top-notch Talents Program of universities in Anhui Province]
grant number [2019gxbjZD24].
Institutional Review Board Statement: Not applicable.
Informed Consent Statement: Not applicable.
Data Availability Statement: Not applicable.
Conflicts of Interest:
The authors declare that they have no known competing financial interest or
personal relationships that could have appeared to influence the work reported in this paper.
Sample Availability: Samples of the compounds are available from the authors.
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