Received: 22 November 2019
| Revised: 16 December 2019 | Accepted: 17 December 2019
DOI: 10.1002/cey2.29
REVIEW
Recycling of mixed cathode lithium‐ion batteries for
electric vehicles: Current status and future outlook
Tyler Or
| Storm W. D. Gourley | Karthikeyan Kaliyappan | Aiping Yu |
Zhongwei Chen
Department of Chemical Engineering,
University of Waterloo, Waterloo,
Ontario, Canada
Correspondence
Zhongwei Chen, Department of Chemical
Engineering, University of Waterloo,
200 University Avenue West, Waterloo,
ON N2L 3G1, Canada.
Email: zhwchen@uwaterloo.ca
Funding information
Natural Sciences and Engineering
Research Council of Canada
Abstract
Worldwide trends in mobile electrification, largely driven by the popularity of electric
vehicles (EVs) will skyrocket demands for lithium‐ion battery (LIB) production. As
such, up to four million metric tons of LIB waste from EV battery packs could be
generated from 2015 to 2040. LIB recycling directly addresses concerns over long‐
term economic strains due to the uneven geographic distribution of resources
(especially for Co and Li) and environmental issues associated with both landfilling
and raw material extraction. However, LIB recycling infrastructure has not been
widely adopted, and current facilities are mostly focused on Co recovery for
economic gains. This incentive will decline due to shifting market trends from
LiCoO2 toward cobalt‐deficient and mixed‐metal cathodes (eg, LiNi1/3Mn1/3Co1/3O2).
Thus, this review covers recycling strategies to recover metals in mixed‐metal LIB
cathodes and comingled scrap comprising different chemistries. As such,
hydrometallurgical processes can meet this criterion, while also requiring a low
environmental footprint and energy consumption compared to pyrometallurgy.
Following pretreatment to separate the cathode from other battery components, the
active material is dissolved entirely by reductive acid leaching. A complex leachate is
generated, comprising cathode metals (Li+, Ni2+, Mn2+, and Co2+) and impurities
(Fe3+, Al3+, and Cu2+) from the current collectors and battery casing, which can be
separated and purified using a series of selective precipitation and/or solvent
extraction steps. Alternatively, the cathode can be resynthesized directly from the
leachate.
KEYWORDS
acid leaching, comingled LIB scrap, hydrometallurgy, NMC, selective precipitation, solvent
extraction
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INTRODUCTION
Lithium‐ion batteries (LIBs) have dominated the secondary
energy storage market due to their unmatched combination
of energy density (150‐200 Wh/kg, normalized by device
mass), power output (>300 W/kg), and cycle stability
(~2000 cycles) coupled with lower costs due to the
increasing global production capacity.1 Large‐scale
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© 2020 The Authors. Carbon Energy published by Wenzhou University and John Wiley & Sons Australia, Ltd.
Carbon Energy. 2020;1–38.
wileyonlinelibrary.com/journal/cey2
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demands for LIB production have recently been driven by
the popularity of electric vehicles (EVs). Moreover, LIB
technology is expected to play an important role in
stationary energy storage systems that require high power
output, enabling energy harvesting from intermittent
natural sources (ie, wind, solar, and geothermal).2
The major components of LIBs are the negative and
positive electrodes, electrolyte, and separator (Figure 1).
The negative and positive electrodes correspond to the
anode and cathode, respectively, during discharge and are
often referred to as such. The separator is an electrically
insulating membrane (eg, polypropylene) that prevents an
electrical short‐circuit due to contact between the electrodes
but is permeable to ion diffusion. The electrolyte is typically
a lithium salt (eg, LiPF6 and LiClO4) dissolved in a mixture
of ethylene carbonate and either dimethyl carbonate,
diethyl carbonate, or ethyl methyl carbonate. This organic
solvent is chosen due to its high electrochemical stability,
allowing the battery to operate on a higher voltage range.
The operating principle for LIBs relies on the intercalation
and deintercalation of Li+ between the electrodes—during
charging, the positive electrode serves as a “source” of Li+
ions, where a power source is applied to the battery to
oxidize the transition metal oxide, which causes the release
of a Li+ ions into the electrolyte (ie, deintercalation) and
simultaneously releases electrons into the external circuit
(Figure 1). The electrons combine with intercalated Li+ at
the graphite‐based negative electrode. During discharge, the
reverse reaction occurs spontaneously, where both electrons and Li+ ions are simultaneously released from the
negative electrode, and the current resulting from released
electrons can power a load. A variety of compositions for
the positive electrode (cathode) active materials (ceramic
FIGURE 1
Major components and operating mechanism of
LIBs. LIB, lithium‐ion battery [Color figure can be viewed at
wileyonlinelibrary.com]
OR ET AL.
intercalation host) are commercially available, with the
most common being LiCoO2 (LCO), LiFePO4 (LFP),
LiMn2O4 (LMO), LiNi1/3Mn1/3Co1/3O2 (NMC‐111, abbreviated as NMC), and LiNi0.8Co0.15Al0.05O2 (NCA). LIBs in
the market are referenced according to the cathode
composition, as this dictates the battery performance. The
various cathode materials have advantages and trade‐offs
with respect to energy density, power capabilities, cost,
toxicity, safety, and stability, which are summarized in
Table 1. To fabricate the electrode, the active materials are
mixed with conductive additives (carbon black) and a
polymer binder such as poly(vinylidene fluoride) (PVDF) to
form aggregates (>80 wt% active material) bound to metal
current collectors (Cu and Al).
1.1
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Emerging market for EVs
Global EV sales (including all‐electric and plug‐in hybrid
EVs) have exponentially increased over the past decade
(Figure 2A). Projections from academic institutions and
consulting firms present unanimously positive outlooks for
the EV market growth. The International Energy Agency
estimates based on current and expected policies that global
EV sales could reach four million in 2020 and 21.5 million
by 2030, corresponding to an approximately 24% yearly
sales growth and a stock value of $13 and $130 million,
respectively.6 EV sales are primarily driven by government
policies, such as tax breaks for EV purchases and
developments in charging infrastructure. Additionally,
decreasing prices for LIB packs due to expansions in
production capacity and steady improvements in the
driving range for EVs are significant factors to increasing
EV sales. In 2018, approximately 5.1 million EVs were on
the road globally, doubling the amount from the previous
year, with the largest markets located in China, Europe, and
the United States, respectively.7 LIBs will remain the
secondary energy storage technology of choice for EVs in
the foreseeable future. Older technologies such as lead‐acid
(Pb‐acid) and nickel‐metal hydride (Ni‐MH) batteries have
low energy density comparably, which limits their EV
application to low cost and limited driving range vehicles
(eg, e‐bikes, scooters, and some hybrid vehicles) and
regenerative brake charging. On the other hand, although
emerging technologies such as lithium‐sulfur and metal‐air
batteries have ultrahigh‐energy storage potential, they
currently demonstrate poor cycle life and power output
and thus will likely enter the market in the distant future.8
The rising popularity of EVs continues to escalate
demands for LIB production exponentially. In 2017,
EVs dominated the energy output of LIBs, more than
doubling the usage from portable electronics
(Figure 2C).9 This is projected to further increase
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OR ET AL.
T A B L E 1 Comparison of most common commercially available LIBs
Cathode material
Nominal
voltage, V
Typical gravimetric
capacity, mAh/g
Typical volumetric
capacity, mAh/cm3
LiCoO2
3.6
145
550
Co is toxic and expensive.
Typical use: portable electronic
devices
Cost: medium
Lifetime: medium
LiMn2O4
4.0
120
496
Mn abundant and environmentally
friendly.
Typical use: power tools and e‐bikes
Cost: low
Lifetime: low
LiNi1/3Mn1/3Co1/3O2
3.7
170
600
Designed to reduce Co
Typical use: portable electronic
devices and EVs
Cost: high
Lifetime: high
LiFePO4
3.3
165
589
Fe abundant and environmentally
benign, high thermal stability
Typical use: power tools and e‐bikes
Cost: medium
Lifetime: high
LiNi0.8Co0.15Al0.05O2
3.7
200
700
Highest specific energy density, used
in Panasonic batteries for Tesla EVs
Cost: high
Lifetime: medium
Li4TiS2
1.9
210
697
Highly stable and safe
Cost: very high
Lifetime: very high
Comments
Note: Gravimetric and volumetric capacity normalized based on cathode active materials.3-5
Abbreviations: EV, electric vehicle; LIB, lithium‐ion batteries.
(+31.6%) by 2025. In addition, EV demand will produce a
shift in the market share for the various types of LIBs
(ie, cathode chemistry; Figure 2D). While LCO batteries
are still prevalent in consumer electronics (eg, mobile
phones and laptops) as the technology is mature and
reliable, they are not ideal for EV applications due to
their relatively short cycle life, safety issues from thermal
runaway reactions (exothermic release of oxygen leading
to fire and explosion), and the high raw material cost of
Co (Figure 2B).3,10 The original 2008 Tesla Roadster
incorporated LCO batteries, but since then, nearly all EV
manufacturers have used NMC type or NCA batteries for
commercial vehicles.3 NCA batteries are notable for
possessing the highest energy density among commercial
LIBs (Table 1) and are incorporated into modern Tesla
EV battery packs, but other EV manufacturers have
chosen NMC cells due to their higher cycle life.11 LFP
batteries are appealing in general due to their high cycle
life, thermal stability, and composition of abundant and
environmentally benign materials. However, their high
cost relative to their energy density is limiting their
application in EVs.12 For example, LFP battery packs
manufactured by A123 were implemented in the 2014
Chevrolet Spark, but in subsequent models, they were
replaced with a Ni‐rich composition (ie, NMC or NCA
type) to reduce the battery pack mass and expand the
driving range. However, it should be noted that LFP
technology currently generates interest in China due to
the scarcity of Ni and Co resources in the region, and has
been developed and adopted by major EV manufacturers,
such as the BYD Company and the Wanxiang Group
Corporation that acquired A123 in 2013.13 Based on all of
these considerations, NMC type LIBs are projected to
dominate the global battery market share in 2025
(Figure 2C,D).
1.2
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Incentives for LIB recycling
The growth of the EV market and LIB production
imposes demand for infrastructure and strategies to
handle LIB waste and potentially recover precious metals
from the cathode. According to a material flow analysis
study conducted by Richa et al,14 anywhere from 0.33 to
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OR ET AL.
FIGURE 2
A, Global new EV sales to date. B, Price (represented in bar size) of valuable metals used in LIBs in 2018 from the US
Geological Survey. Projection of LIB market share based on (C) application and (D) cathode composition. The projection assumes that Tesla,
Inc remains the only major EV manufacture to adopt NCA cells in 2025. EV, electric vehicle; LIB, lithium‐ion battery [Color figure can be
viewed at wileyonlinelibrary.com]
4 million metric tons of LIBs from EV battery packs could
be generated from 2015 to 2040 based on their
conservative to extreme estimates, respectively. The
estimation accounts for different projections on EV sales,
the lifespan of LIB cathodes (8‐10 years on average), and
the number of cells per EV battery pack. Their baseline
projection of 1.3 million metric tons of LIB waste from
EVs yields a commodity value of approximately three
billion USD assuming 100% collection and recovery of all
metals, including aluminum, copper, nickel, cobalt, iron,
and steel. LIB recycling is also critical toward preserving
precious resources. The Cobalt Institute estimates that
approximately 50% of cobalt produced worldwide is
currently used in secondary batteries, with the vast
majority in LIBs and a minor amount in Ni‐MH
batteries.15 Moreover, cobalt is an important component
in integrated circuits, semiconductors, magnetic recording devices (eg, hard disks), and various high‐strength
alloys for applications ranging from space vehicles to
prosthetic and dental applications. It has also generated
significant research interest as a catalyst in a wide range
of applications.16-19 However, cobalt is considered a
critical resource as approximately 60% of the worldwide
mine production in 2018 originated from the Dominican
Republic of Congo, where the geopolitical instability and
unethical labor practices are well documented.15,20,21
The US Geological Survey estimates that 39% of all
lithium produced is used in primary and secondary
lithium‐based batteries.22 The consensus regarding
lithium availability suggests that although the exact
quantity of recoverable global lithium reserves is
uncertain, they should be able to meet long‐term
projected demands (up to 2100).21,23,24 However, the
uneven geographic distribution of the reserves can cause
price spikes for raw lithium materials. For instance,
cumulative lithium demands in China may exceed the
country’s lithium reserves around 2028.25 In addition,
global demands for lithium could surpass production by
2050.26,27 These concerns have motivated research
toward sodium‐ion batteries (SIBs) as they possess a
similar operating mechanism, cathode chemistry, and
theoretical energy density to LIBs. However, SIBs are far
from commercial realization due to their poor cycle life.
It is evident that LIB recycling is the most direct solution
to mitigate strain on precious raw material reserves.
In addition to material savings, LIB recycling has
positive impacts toward energy consumption and the
environmental protection. Li, Ni, Co, and Al production
requires high energy to be extracted from virgin resources
and causes the release of significant quantities of greenhouse gases (GHG) from transportation and smelting
processes.28 Although Al, Cu, and Fe are not currently
perceived as valuable components of LIBs, recycling can
save 95%, 85%, and 74%, respectively, of the total energy
required to obtain them through ore extraction, while also
preventing a substantial amount of harmful CO2 and SO2
emission.29,30 Peters et al31 compiled a review of life cycle
analysis studies in the literature and concluded that
producing 1 kWh of storage capacity on an average LIB
(across all cathode chemistries and geometries) resulted in a
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OR ET AL.
cumulative energy demand of 328 kWh and 110 kg of CO2
emissions. Recycling has the potential to significantly
reduce these emissions, especially if pyrometallurgical
methods are avoided.32 In addition, most LIBs are currently
disposed in municipal landfills unless restricted by regional
policies. Environmental hazards occur when water enters
the landfill and leaches toxic metals from LIBs. The issue is
exacerbated by the fact that anaerobic microorganisms in
the landfill produce acids that can corrode the battery
casing over time. However, legislation can be put in place to
mitigate these issues and drive the development of LIB
recycling infrastructure. The European Union Battery
Directive (2006) prohibits the landfilling of LIBs and sets
a target “recycling efficiency” of 50% by weight, although
the legislation was initially written with Zn‐based batteries
in mind.33,34 Various provincial/state legislations around
the world have also outlawed the landfilling of LIBs.
1.3 | Current and future LIB recycling
infrastructure
LIBs implemented in the first generation of EVs have likely
reached their end of life starting approximately 2015, which
presents an urgent need to proactively develop waste
management systems. Table 2 summarizes the recycling
method of current facilities to the best of knowledge. As
seen in Table 2, most industrial facilities use pyrometallurgy,
where spent LIBs are fed directly into a furnace and smelted
at high temperatures (>1000°C). In this process, the plastic,
electrolyte, and carbonaceous components are decomposed/
eliminated and valuable metals (Co, Ni, Cu, etc) are
collected as molten metals and alloys. Hazardous emissions
must be removed with flue cleaning systems to meet
stringent environmental regulations. This is the traditional
method of industrial‐scale e‐waste recycling adopted from
the mining industry due to its simplicity and high
productivity and process capacity.35 However, this process
typically does not recover Li, which is lost as slag with other
refractory oxides and gases. In addition, the recovered alloys
require further refinement through hydrometallurgical
routes.36 Pyrometallurgical recycling can potentially result
in a net increase in GHG emissions and energy consumption
compared to raw material extraction, due to the incineration
of nonvaluable battery components.37
Most existing LIB recycling plants primarily focus on
recovering cobalt. This is because (a) recycling plants
must focus on recovering older LIBs (mostly LCO) that
T A B L E 2 Summary of process in major LIB recycling facilities worldwide to the best of knowledge
Recycling plant
Input
Summary of process
Umicore
NiMH, LIBs
Large‐scale batteries dismantled in dedicated disassembly line. Battery scrap
is melted down (without prior pretreatment) with addition of carbonaceous
materials to reduce the metals. Alloy of valuable metals (Co, Ni, and Cu)
collected whereas Li, Mn, Fe, and Al oxides are disposed as slag. Valuable
metals recovered in hydrometallurgical process— acid leaching, followed
by precipitation as salts (impurities also removed by precipitation).38
Retriev Technologies
(Toxco)
Primary and secondary
Li‐batteries
Batteries chilled in liquid nitrogen (−196°C) before shedding to nullify
violent reactions associated with lithium metal short‐circuiting. Li+/Li
metal removed by dissolution in water. Metals separated by sieving/
filtering, then recovered though a hydrometallurgical process. Li recovered
as Li2CO3.38-40
Recupyl
LIBs
Batteries shredded using a rotary shearing machine under inert atmosphere
(Ar and CO2) to avoid violent Li reactions. Li/Li+ removed by dissolution in
water. Metal oxides are separated as a fine powder and recovered though a
hydrometallurgical process. Li recovered as Li2CO3 or Li3PO4.41
Duesenfeld (LithoRec)
LIBs
Shredding under inert atmosphere and heating to evaporate electrolyte.
Electrolyte recovered by condensation and CO2‐assisted extraction.
Magnetic and sieving separation to remove battery casing and separator.
Heating at 400°C‐600°C to separate active material from current collectors.
Recover active material metals through hydrometallurgical process.42
Inmetco
NiCd, NiMH, LIBs
Batteries smelted to form alloy containing Co, Ni, and Fe. Other metals (eg,
Li components) disposed as slag.39
Sony‐Sumitomo
LIBs
Batteries incinerated at 1000°C, where organics, fluorides, and lithium are
removed as fly ash. Co recovered from metal residue though
hydrometallurgical process.38
Abbreviation: LIB, lithium‐ion battery.
6
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are now reaching the end of their lifespan and (b) cobalt
is by far the most valuable metal in LIBs (Figure 2B) and
is the primary driver for profitability in recycling. Wang
et al43 estimated in 2016 that based on state‐of‐the‐art
recycling efficiencies reported in the literature, one
metric ton of LCO batteries can yield $8900 in value
while an equivalent mass of LMO yields merely $890. For
LMO and LFP batteries, the Cu current collector foil is
the most valuable component. Thus, LCO must comprise
at least 21% of the total LIB scrap in order for current
recycling plants to be profitable. Proposed strategies to
address this, such as battery labeling and sorting based on
cathode chemistry, do not address the need for recycling
infrastructure to be robust enough to handle metal
recovery beyond Co. Taken together, it is evident that
future recycling infrastructure cannot primarily focus on
recovering Co to maximize profits, especially given the
market trends for LIB cathode chemistries driven by the
EV market (Figure 2). In the near future, NMC‐111 will
likely be substituted with LiNi0.6Mn0.2Co0.2O2 (NMC‐622)
and LiNi0.8Mn0.1Co0.1O2 (NMC‐811), which comprise
even smaller quantities of Co.21 Thus, recycling plants
should handle diverse mixed‐type cathodes and comingled scraps containing various cathode chemistries
with high efficiency. In addition, although Li recovery is
not currently a priority as Li2CO3 is perceived as cheap
and abundantly available, this is unsustainable given the
long‐term projected demands for LIBs. Ideally, recycling
plants should also use processes that have lower energy
input and environmental imprint.
This work aims to review LIB recycling developments
in the literature in the context of meeting demands for
the growing EV market. Practical challenges associated
with the collection and disassembly of EV battery packs
are also discussed. With respect to metal recovery, a
particular emphasis is placed on hydrometallurgical
processes, as it is robust in handling various metal
compositions. Compared to pyrometallurgy, hydrometallurgy requires low energy input, produces low quantities
of toxic emissions, and ideally can recover all valuable
metals in LIB scrap at high purity. In this route, LIB cells
are pretreated to separate the cathode from other
components in the battery (current collectors, separator,
electrolyte, casing, etc) where Cu and Al can be collected.
The cathode is then leached in acid to dissolve Li, Ni, Mn,
and Co. The leachate can then be treated with a series of
selective precipitation and/or solvent extraction steps to
either recover the metals as salts or subsequently
resynthesize the cathode directly from the leachate. The
various approaches and optimized protocols reported in
the literature to leach and recover mixed‐type cathodes
and/or comingled LIB scrap (generating a complex
leachate) are highlighted and discussed herein.
OR ET AL.
2 | PRETREATMENT OF LIB
C EL L S
The first stages of LIB recycling involve a pretreatment
step to separate the cathode active materials from the
battery casing, separator, current collector, electrolyte,
carbonaceous additives, and connections. The approaches utilized in the literature should be divided into
lab‐scale and large/industrial‐scale methods. Lab‐scale
approaches result in excellent cathode active material
separation and efficiency and are often used in experiments that focus on leaching and/or subsequent metal
recovery. On the other hand, industrial‐scale approaches
have high process capacity and throughput, but metal
separation is less refined.
2.1
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Lab‐scale pretreatment
In a typical lab‐scale approach, the LIB cell is first
discharged by soaking the cell in a saturated brine solution
for approximately 24 hours. This is necessary to reduce risks
of violent short‐circuiting and exothermic reactions of
lithium deposits in the anode with oxygen and water,
leading to ignition of the highly flammable organic solvent.
The casing is then disassembled manually to retrieve the
cathode and Al current collector sheet. Separation of the
cathode active material from the current collector, binding
agent, and conductive additive is achieved either by (a)
high‐temperature calcination (350°C‐600°C) to decompose
the organic binder, additives, and electrolyte and liberate
the active material as a fine powder; (b) dissolution of
PVDF in N‐methyl‐2‐pyrrolidone (NMP) assisted using heat
and/or sonication, followed by drying and filtration44; and
(c) use of a strong base to dissolve the Al current collector
(forming NaAlO2 [aq]).45,46 It should be noted that NMP
dissolution is not applicable to electrodes that use a poly
(tetrafluoroethylene)‐based binder, as it is nonpolar,
whereas PVDF possesses alternating –CH2– and –CF2–
units. Furthermore, as this process does not remove all
PVDF and conductive carbon from the active material,
NMP dissolution may be followed by calcination to degrade
the organic components.
There is ongoing research to reduce the negative
environmental impacts of these techniques and improve
their compatibility with large‐scale processes. For instance, due to the volatility and high cost of NMP,
alternative reagents to dissolve PVDF have been investigated, including ionic liquids and molten salts (AlCl3‐
NaCl), which can also be reused.47,48 Additionally, to
avoid the release of highly toxic hydrogen fluoride (HF)
vapor when PVDF is calcined at approximately 500°C,
PVDF can be reacted with CaO to reduce its
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OR ET AL.
decomposition temperature to 300°C and thus avoid HF
liberation.49
2.2
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Industrial‐scale pretreatment
Table 2 shows that pretreatment in industrial recycling
facilities typically involves mechanically crushing LIBs
followed by scrap separation through sieving. This is
necessary as lab protocols involving the manual
disassembly and separation of battery components do
not address the requirements for scaling up. Crushing
is performed under cryogenic cooling or inert atmosphere, as this lowers the reactivity of Li/Li+ and thus
reduces risks of explosions, fires, and toxic gas
emissions (eg, organic solvent, fluorophosphates, and
HF vapor liberation) when penetrating LIB cells. A
potentially cheaper and lower energy alternative
process would involve discharging spent LIBs in salt
solution before crushing. This is promising due to its
simplicity, effectiveness, and also having the additional
benefit of absorbing the heat evolved during discharge.38 Li et al50 quantified the discharge kinetics
of 18 650 cells in various brine concentrations (5, 10,
and 20 wt% NaCl), concluding that 10 wt% NaCl was
the most effective with a 72% voltage drop (from fully
charged state) after approximately 6 hours of soaking.
However, it should be noted that metal contaminants
(especially Al and Fe) were leached in solution, mostly
originating from corrosion on the metal casing. To
address this, Shaw‐Stewart et al assessed LIB discharge
in various salt compositions, demonstrating that halide
(containing Cl−, Br−, or I−) and acidic (pH < 4, eg,
NaHSO4−) salts cause rapid corrosion of the battery
casing, and highly alkaline (pH > 12) salts can perforate the casing.51 This leads to hazards from leakage
that can generate HF acid if LiPF6 is present (LiPF6 +
H2O → LiF + POF3 + 2HF). With this in mind, they
identified NaNO2 solution as a promising discharge
electrolyte with low corrosion rate.
After crushing, the cell scrap is often separated
gravimetrically using a series of vibrating sieves. In
general, battery casing materials (Al‐Fe‐Mn alloy),
plastics, and current collectors are separated as coarse
particles while fine particles comprise mostly of electrode
materials.52,53 This separation occurs because the battery
casing and metal foils are more ductile and difficult to
crush while the electrode active materials exist originally
as fine powders. A magnetic separator may also be used
to remove metal casing pieces (high Fe content) from the
cathode powder.54 Wang et al43 processed cryogenically
cooled LIB cells in a mechanical shedder (granulator)
and sequentially sieved the scraps to separate them into
five different size fractions. When shredding LCO
batteries, they obtained an 82 and 68 wt% Co composition
among metals in the ultrafine (<0.5 mm mesh size) and
fine (0.5‐1 mm) fractions, respectively, suggesting promise as a crude separation method. The study also
processed NMC‐type LIBs (Li1.05(Ni4/9Mn4/9Co1/9)O2),
obtaining 92 wt% active material (50% Ni, 22% Mn, 20%
Co) in the ultrafine fraction. It should be noted that the
cryogenic cooling step has a positive impact on electrode
separation and yield. The cooling crystallizes PVDF,
which induces crack formation throughout the electrode
material, whereas the yield and tension strength of the Al
current collector increases.55 As a result, the current
collector remains somewhat intact after shearing, while
the electrode powder detaches efficiently and can be
collected in high purity after sieving. Assuming the
cathode active material is the recovery target (Co and Ni‐
rich batteries), the clear disadvantage of this technique is
the loss of active material in the coarser fractions that are
mixed with metals (Cu, Al, Fe) from the battery casing
and current collectors. There is a trade‐off between yield
and purity when selecting the sieve mesh size for the fine
fraction.53 Furthermore, the fine fraction will contain a
mixture of cathode metals and graphite coated with
PVDF and carbon black.56 Removal of the carbonaceous
component by heat treatment and/or separation may be
required to improve metal recovery efficiency. This
calcination is often performed at less than 600°C, as
high temperatures (900°C) could generate a molten Al
(impurity) coating on the active material and affect
subsequent leaching efficiencies.54 He et al57 used
flotation to separate LiCoO2 and graphite, which relies
on their difference in wettability. As the carbon additive
coating on the metals masks this property, the coating
was first degraded using Fenton’s reagent (FeSO4 +
H2O2). Wang et al58 performed a similar process and
achieve good metal recovery via flotation after degrading
the carbon coating at 450°C for 15 minutes. Furthermore,
if Cu is the recovery target (valuable component in LFP
and LMO batteries), the coarse fractions will also contain
Al, Fe, and plastics, thus requiring further treatment,
such as the removal of plastics through electrostatic
separation. It should also be noted that crushing LIBs
releases volatile organic compounds from the electrolyte
(especially dimethyl carbonate and tert‐alkylbenzenes),
thus requiring extensive ventilation and air purification
equipment.50 To address this, crushing in the presence of
water flow (ie, wet crushing) has been proposed.52
However, wet crushing can potentially increase operating
costs and lower the purity of cathode active materials
collected in the fine fractions.
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2.3
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Electrolyte and graphite recovery
In both lab and industrial‐scale pretreatment approaches
described here, the electrolyte and other organic constituents are usually decomposed and/or discarded.
Although the economic justification for recycling these
components is questionable, recovery approaches have
been explored in the literature. Vaporized electrolyte
solvent evolved from crushing and heating can be
recovered by condensation, while the liquid organic
solvent can potentially be dissolved in water and
subsequently recovered by distillation, which is viable
due to its high boiling point (242°C‐248°C) relative to
water.59 To recover electrolyte immobilized in the pores
of the electrode material and separator, liquid or supercritical CO2 extraction can be applied on the shredded
material.60,61
In lab‐scale LIB pretreatment, graphite is recovered by
first manually separating the anode component. Separation of graphite from the Cu current collector can then be
achieved through NMP dissolution, acid dissolution, or
mechanical scraping.62 As discussed previously, LIBs
pretreated by crushing and sieving contain electrode
materials in the fine fractions—separation of carbons and
metals can be achieved through flotation or magnetic
separation.57,58,63,64 The main challenge in graphite
recycling is the removal of electrolyte, solid‐electrolyte‐
interface (SEI) layer, and other carbonaceous components (binder and conductive additive) for purification.
Thus, high‐temperature treatment is likely required to
remove carbons and SEI‐layer components, while the
electrolyte can be removed through evaporation or CO2‐
assisted extraction.65
3 | D I S A S S E M BL Y O F E V
B A TTERY P A CKS
One challenge, particularly with respect to processing EV
battery packs, is that lack of automation infrastructure
and design standardization to feasibly disassemble the
packs. EV battery packs can be relatively complex, but
generally comprise of LIB cells (cylindrical, prismatic, or
pouch geometry) assembled into modules with heating/
cooling components (eg, tabs, pipes, or plates with
refrigerant solution). The battery pack also contains the
battery management unit (safety electronics to prevent
over‐charge/discharge) and other control electronics
housed in an insulated casing.42 EV manufacturers likely
employ different battery pack designs. For instance, the
Tesla Model S 85 kWh battery modules contain 444 (6
series, 74 parallel) cylindrical NCA cells (18 650, 3.6 V,
3.2 Ah). Connecting 16 modules in series yields a
nominal voltage of 345.6 V and 7104 cells in total
(Figure 3A). On the other hand, the Audi Q5 plug‐in
hybrid EV possesses four modules that each contain 18
prismatic or pouch NMC cells (5 Ah, 3.7 V) all connected
in series, yielding a total of 72 cells and 266 V
(Figure 3B).66 Establishing automation in the disassembly of EV battery packs has the potential to reduce
processing costs and reduce safety concerns due to the
high operating voltages of the packs. Although research
on optimizing the disassembly of complex products is
well established,67 little work has been dedicated
specifically toward EV battery packs.68 In general,
disassembly lines have a low degree of automation due
to (a) heterogeneous designs from various manufacturers;
(b) modern design complexities (eg, miniaturization and
close‐packing); and (c) environmental factors that add
variability (eg, physical and corrosion damage). Herrmann et al69 proposed and conducted a methodology to
assess the automation potential of disassembling an EV
battery pack. The authors assessed 15 steps of the
disassembly process based on the technical difficulty
(eg, welded connections more difficult to disassemble
compared to screws) and necessity (eg, economic and
safety benefits). They concluded that some aspects are
well suited for automation—for instance, dismantling of
the modules and connections must be done manually,
while the individual cells could be autonomously
extracted using a prototype gripper system where the
jaw contact plates can simultaneously monitor the cell
state of health through its potential and internal
resistance.70 Wegener et al66 manually disassembled an
Audi Q5 battery pack to assess practical issues with
automated disassembly.71 Some complexities include the
various screw types inserted at different orientations and
cables that are difficult to access. They proposed a hybrid
automated system where a robotic manipulator could
work alongside a human worker to remove screws and
bolts and liberate the modules (Figure 3C).72 These
findings clearly demonstrate that it is difficult to realize
automation with current EV battery pack designs, and
design standardization (eg, reduction in number of joints)
is required to allow ease of access to the LIB cells and
separation of electronic components.73 Achieving this
will likely require political pressure due to the proprietary nature of EV pack designs.
4
|
ACID LEACHING
To recover valuable metals in high purity, hydrometallurgical processes are often performed on the pretreated
battery scraps. In this approach, the separated cathode
metals are leached as ions typically in acid solution.
OR ET AL.
| 9
F I G U R E 3 A, Schematic of Tesla 85 kWh battery pack comprised of 16 modules and 7104 total cylindrical cells. B, Schematic of Audi
PHEV 5 kWh battery pack comprised of four modules and 72 total prismatic or pouch cells. C, Overview of proposed EV battery pack
recycling steps based on current automation potential. EV, electric vehicle [Color figure can be viewed at wileyonlinelibrary.com]
Upon leaching, valuable metals are recovered, either
through extraction and purification or direct resynthesis
of the cathode material as reviewed in Sections 6 and 7,
respectively. Hydrometallurgical recovery is promising
for large‐scale applications due to its low energy
requirement, which mostly comes from heating a
reaction vessel solution. In addition, it is robust and
suitable for mixed cathode compositions, as ideally all
metal types can be leached in solution and selectively
recovered in high purity and efficiency/yield. Although
selective recovery can potentially be arduous, development and optimization in this area can eliminate the
need to sort LIBs by cathode chemistry for recycling
facilities.
The leaching agent comprises an acid solution
(inorganic/mineral or organic) often coupled with a
reducing agent (eg, H2O2, NaHSO3, glucose, citric acid,
etc.). Li+ is easily leached in solution with its efficiency
strongly correlated with the acidic strength (ie, displacement of Li+ with H+ to induce solubilization). However,
cathode transition metals have low solubility as they exist
in +3/+4 valence states in discharged cathodes and are
difficult to leach due to the strong M–O bonds. The
reductant aids leaching by reducing the metals toward a
divalent state (M2+), which is critical toward achieving
high leaching efficiencies (>95%), especially for Co and
Mn.74 H2O2 is the most commonly used reductant as it
possesses a low oxidation potential, low reagent cost, and
produces harmless byproducts. Six key parameters affect
leaching kinetics and yield: temperature, reaction time,
acid concentration, reductant concentration, liquid/solid
ratio (inverse of pulp density), and stirring/agitation.
Ideally, they should be minimized to reduce the material
cost and required energy input. When developing novel
leaching systems, these parameters are first optimized by
assessing each as the bottleneck toward achieving high
leaching efficiency (Figure 4). In general, higher reaction
temperatures improve leaching kinetics and efficiency, as
it increases the frequency of reactant collisions and
provides the activation energy for the endothermic
reaction. However, temperatures that are too high can
reduce efficiency due to the vaporization and/or decomposition of acids and reductants (especially H2O2).
Similarly, high acid concentrations promote leaching,
but concentrations that are too high could restrict the
diffusion of leached products. The liquid/solid ratio is the
most important parameter to minimize in industrial‐scale
leaching as it significantly affects the operating cost. In
some work, leaching is coupled with ultrasonic cavitation, which forms rapid bubble formation and collapse in
10
|
OR ET AL.
solution. This generates high‐impact collisions and localized
heating, enhancing reactivity between the solid and solvent,
and thus improving leaching kinetics and efficiency.75
Table 3 summarizes the optimized leaching parameters and
efficiencies from various approaches in the literature.
4.1
|
Inorganic acid leaching
Inorganic acids are commonly used due to their
effectiveness and low reagent cost. Compared to HNO3,
H2SO4, and H3PO4, HCl leaching requires a lower
concentration threshold to achieve high leaching rates
and efficiency.105 This is ascribed to (a) its stronger
acidity (ie, higher dissociation constant); (b) the presence
of corrosive Cl− anions that possess a lower pitting
potential compared to sulfate, nitrate, and phosphate
anions; and (c) Cl− can serve as a reductant for
M3+/M4+.80,86 As seen in Table 3, high leaching
efficiencies can be achieved for HCl without the
supplementation of a reducing agent.106-108 However,
coupling HCl with a reducing agent can lower the
concentration threshold of HCl required to reach high
efficiency.109,110 A key disadvantage of using HCl is the
evolution of toxic Cl2 vapor from the oxidation of Cl− as
shown in Equation (1), which requires specialized
equipment to handle. In addition, the Cl− anions
generate significant corrosion concerns for stainless steel
facility equipment.
6LiNi1/3Mn1/3Co1/3 O2 (s) + 24HCl(aq)
→ 2NiCl2 (aq) + 2MnCl2 (aq) + 2CoCl2 (aq)
+ 6LiCl(aq) + 3Cl2 (g) + 12H2 O(l).
(1)
On the other hand, the sulfuric acid leaching products
are more benign as shown in Equation (2). Meshram
et al74 demonstrated that when using H2SO4 without a
reducing agent, maximum leaching efficiencies of 50.2%
and 66.2% were obtained for Mn and Co respectively, due
to the presence of lowly soluble Co3+ and Mn4+ in spent
batteries. However, when coupled with a reducing agent,
Table 3 shows that excellent leaching efficiencies (>95%)
can be achieved for all metals.
6LiNi1/3Mn1/3Co1/3 O2 (s) + 9H2 SO4 (aq) + H2 O2 (l)
→ 2NiSO4 (aq)+ 2MnSO4 (aq) + 2CoSO4 (aq)
+ 3Li2 SO4 (aq) + 2O2 (g) + 10H2 O(1).
(2)
Phosphoric acid leaching has also generated interest
due to its mild acidity and low corrosivity compared to
HCl and H2SO4.91,111 Zhuang et al leached LiNi0.5Mn0.3Co0.2O2 using a mixture of H3PO4 and citric acid at
relatively mild conditions (0.6M total acid concentration,
liquid/solid ratio 50 mL/g, and 90°C for 0.5 hours) as
depicted in Equation (3). Citric acid (C6H8O7) served as
both a leaching and reducing agent, resulting in high
leaching efficiencies (>91%) for all metals.
Typical optimization process of acid leaching parameters: A, Acid concentration (L‐tartaric acid, C4H6O6), (B) reductant
concentration (H2O2), (C) pulp density, and (D) temperature and leaching time. Adapted with permission: Copyright 2017, American
Chemical Society76 [Color figure can be viewed at wileyonlinelibrary.com]
FIGURE 4
| 11
OR ET AL.
T A B L E 3 Summary of approaches for acid leaching of LIB cathodes comprising Li, Ni, Mn, and Co
Cathode material(s)
Optimized conditions
Leaching efficiency
Reference
Mixture of LCO, LMO, and NMC active
material powder
• 4M HCl
• Liquid/solid ratio 50 mL/g
• 80°C for 1 h
99% Li, 99.8% Ni, 99.8% Mn,
99.5% Co
Wang et al77
NMC, LFP, and other cathode types from
spent commercial batteries
• 6M HCl and H2O2/M2+ = 2
(mole)
• Liquid/solid ratio 8 mL/g
• 60°C for 2 h
>95% Ni, Mn, and Co
Li et al78
Spent commercial LIBs of various
compositions
• 1.75M HCl
• 50°C for 2 h
• Liquid/solid ratio 5 mL/g
99.2% Li, 99% Mn, 98% Co
Barik et al79
NCA
• 4M HCl
• Liquid/solid ratio 20 mL/g
• 90°C for 18 h
100% Li, Ni, Co, and Al
Joulie et al80
Spent commercial LIBs of various
compositions
• 4M HCl
• Liquid/solid ratio 50 mL/g
• 80°C for 1 h
99% Li, 94.5% Ni, 93.4% Mn,
98.7% Co
Dhiman and
Gupta81
Commercial LIBs of various compositions
• 1M H2SO4 (no reductant)
• Liquid/solid ratio 20 mL/g
• 95°C for 4 h
94.3% Li, 96.3% Ni, 50.2%
Mn, 66.2% Co
Meshram et al74
Mixture of NMC, LCO, LMO, and LFP active
material powder
• 4M H2SO4 and 30 wt% H2O2
~100% for Li, Mn, Ni, Co; Fe
partially dissolved
Zou et al82
NMC and LCO active material powder
• 1M H2SO4 and 1 vol% H2O2
• Liquid/solid ratio 25 mL/g
• 40°C for 1 h
~99.7% for Li, Ni, Mn, Co
He et al83
Mixture of NMC, LCO, LMO, and LFP
• 2M H2SO4 and 2 vol% H2O2
• Liquid/solid ratio 20 mL/g
• 80°C for 1 h
Unspecified
Chen et al84
Mixture of NMC, LCO, LMO, and LFP
• 2.5M H2SO4, H2SO4/H2O2 = 5
(vol/vol)
• Liquid/solid ratio 33 mL/g
• 60°C for 1 h
99% Li, 99.1% Ni, 99.1% Mn,
99.2% Co
Zheng et al85
NMC cathode scrap
• 2.5M H2SO4 and 0.8M NH4Cl
• Liquid/solid ratio 10 mL/g
• 80°C for 1 h
99.11% Li, 97.49% Ni, 97.34%
Mn, 97.55% Co
Lv et al86
Spent commercial LIBs of various
compositions
• 1M H2SO4 + 0.075M NaHSO3
• Liquid/solid ratio 50 mL/g
• 95°C for 4 h
96.7% Li, 96.4% Ni, 87.9%
Mn, 91.6% Co
Meshram et al87
Thermomechanochemical pretreatment:
NMC active material ball milled with lignite
(20 wt% dosage) and roasted at 350°C for 3 h
• H2SO4 (1.15 eq)
• Liquid/solid ratio 3.5 mL/g
• 55°C for 2.5 h
98% Ni, 98% Mn, 96% Co
Zhang et al88
Thermomechanochemical pretreatment:
NMC active material ball milled with carbon
black (10 wt% dosage) and roasted at 550°C
for 0.5 h under Ar
• 4M H2SO4
• Liquid/solid ratio 10 mL/g
• 90°C for 30 min
99.56% Ni, 99.9% Mn,
99.87% Co
Liu et al89
Spent commercial LIBs of various
compositions
• 2M H2SO4 and 4 vol% H2O2
• Liquid/solid ratio 20 mL/g
• 50°C for 120 min
>98% Li, Ni, Mn, and Co
Sattar et al90
Spent commercial LiNi0.5Mn0.3Co0.2O2
• 0.2M H3PO4 acid and 0.4M
citric acid
• Liquid/solid ratio 50 mL/g
• 90°C for 30 min
100% Li, 93.38% Ni, 91.63%
Co, 92% Mn
Zhuang et al91
Inorganic acids
(Continues)
12
|
OR ET AL.
T A B L E 3 (Continued)
Cathode material(s)
Optimized conditions
Leaching efficiency
Reference
Mechanochemical pretreatment: LIB active
material ball milled with Fe powder (1:1
weight ratio)
• 1M HNO3
• Liquid/solid ratio 333 mL/g
• 25°C for 2 h
77.15% Li, 99.9% Ni, 100%
Mn, 91.25% Co
Guan et al92
Crushed commercial LIBs of various
compositions
• HNO3
• 35 mmol HNO3 per g of scrap
• 70°C for 5 h
80% Co, 85% Li
Peng et al93
Spent commercial LCO, LMO, and NMC
18 650 cylindrical and pouch cells
• 0.5M citric acid and 1.5 vol%
H2O2
• Liquid/solid ratio 50 mL/g
• 90°C for 1 h
96.02% Li, 97.79% Ni, 99.76%
Mn, 99.20% Co
Li et al94
Spent commercial NMC
• 2M citric acid and 2 vol% H2O2
• Liquid/solid ratio 30 mL/g
• 80°C for 1.5 h
99% Li, 97% Ni, 94% Mn,
95% Co
Chen et al95
Spent commercial LIBs of various
compositions
• 1.5M citric acid and 2 vol%
H2O2
• Liquid/solid ratio 50 mL/g
• 95°C for 30 min
97% Li, 99% Ni, 95% Co
Musariri et al96
Spent commercial NMC
• 1M D,L‐malic acid and 6 vol%
H2O2
• 50°C for 30 min
Unspecified
Yao et al97
Spent commercial NMC
• 1.2M D,L‐malic acid and 1.5 vol
% H2O2
• Liquid/solid ratio 25 mL/g
• 80°C for 30 min
98.9% Li, 95.1% Ni, 96.4%
Mn, 94.3% Co
Sun et al98
Spent commercial NMC
• 2M maleic acid and 4 vol%
H2O2
• Liquid/solid ratio 50 mL/g
• 70°C for 1 h
99.45% Li, 98.58% Ni, 98.16%
Mn, 98.45% Co
Li et al99
Spent commercial LiNi0.5Mn0.3Co0.2O2 and
LCO
• 2M L‐tartaric acid and 4 vol%
H2O2
• Liquid/solid ratio 59 mL/g
• 70°C for 30 min
99.07% Li, 99.31% Ni, 99.31%
Mn, 98.64% Co
He et al76
Spent commercial NMC
• 1.5M lactic acid and 0.5 vol%
H2O2
• Liquid/solid ratio 20 mL/g
• 70°C for 20 min
97.7% Li, 98.2% Ni, 98.9%
Mn, 98.4% Co
Li et al100
NMC cathode and Al current collector cut to
10 × 10 mm pieces
• 15 vol% trifluoroacetic acid
• Liquid/solid ratio 8 mL/g
• 40°C for 3 h
68.9 wt% Li, 35.01 wt% Ni,
35.18 wt% Mn, 35.39 wt%
Co, 9.31 wt% Al in leachate
(Rest of metals in cathode
powder recovered after
calcining residual carbons)
Zhang et al101
NMC cathode and Al current collector cut to
10 × 10 mm pieces
• 3M trichloroacetic acid and
4 vol% H2O2
• Liquid/solid ratio 20 mL/g
• 60°C for 30 min
99.7% Li, 93% Ni, 89.8% Mn,
91.8% Co
Zhang et al102
NMC cathode and Al current collector cut to
10 × 10 mm pieces
• 3.5M acetic acid and 4 vol%
H2O2
• Liquid/solid ratio 25 mL/g
• 60°C for 1 h
99.97% Li, 92.67% Ni, 96.32%
Mn, 93.62% Co
Gao et al103
Organic acids
(Continues)
| 13
OR ET AL.
T A B L E 3 (Continued)
Cathode material(s)
Optimized conditions
Leaching efficiency
Reference
Spent commercial NMC
• 1M acetic acid and 6 vol% H2O2
• Liquid/solid ratio 50 mL/g
• 70°C for 1 h
98.83% Li, 97.93% Ni, 97.74%
Mn, 97.85% Co
Li et al99
Spent commercial NMC
• Citrus juice (contains citric and
maleic acid for leaching;
ascorbic acid and flavonoids as
reducing agent)
• Liquid/solid ratio 20 mL/g (solid
mass includes Cu and Al foil)
• 90°C for 30 min
100% Li, 98% Ni, 99% Mn,
94% Co
Pant and
Dolker104
Abbreviation: LIB, lithium‐ion battery.
10LiNi 0.5Mn 0.3Co0.2 O2 (s) + 22H3 PO4 (aq) + C6 H8 O7 (s)
→ 10Li+(aq)
+
+
5Ni2+(aq)
22H2 PO−4 (aq)
+
3Mn2+(aq)
+
+ 6CO2 (g) + 15H2 O(1).
2Co2+(aq)
(3)
Although few studies have specifically addressed the
leaching and recovery of NCA cathode metals, the overall
methods and challenges should be similar to other
mixed‐type compositions. Joulié et al80 achieved approximately 100% leaching efficiency of Li, Ni, Co, and Al
using 4M HCl and liquid/solid ratio 20 mL/g at 90°C for
18 hours. They also assessed leaching using HNO3 and
H2SO4 in the absence of a reducing agent, which
expectedly resulted in substantially lower efficiencies
for Ni, Co, and Al. In addition, little work has been
dedicated toward LFP cathodes due to the abundance
and low cost of Fe. However, Fe is difficult to leach due to
the high stability in the Fe–P–O bonds, thus requiring
high acid concentrations (2.5‐6M).82,112,113 Li et al114 took
advantage of this property by using dilute H2SO4 to
selectively leach Li and subsequently recover Li and
FePO4 separately. Using the leaching conditions described in Table 3, LFP cathodes introduced in a
comingled scrap will be partially leached, comprising
Li+ and small quantities of Fe3+.82
organic acids is directly correlated with their acidic strength
(pKa), especially for Li+.117 Although their acidity is
significantly lower than strong inorganic acids, organic
acids are potent leaching agents as they can (a) serve as
moderate reducing agents for multivalent transition metals
and (b) stabilize the dissolution of metallic ions by forming
chelation complexes.117,118 For instance, ascorbic acid
possesses an enediol group that can release two equivalents
of H+ (pKa1 = 4.17 and pKa2 = 11.6) and has functional
groups (hydroxyl, alkoxide, and ester) with an affinity
toward metal cations. It is also a potent reducing agent
(E0 = −0.28 V), as it possesses an oxidizable reductone
group that can resonance‐stabilize radical formation on the
alkoxide groups (Figure 5A).119 Thus in the absence of an
additional reducing agent, organic acids can outperform
inorganic acids at leaching transition metals when the
leachate pH is identical.120 Shih et al121 assessed Co
leaching using 2M H2SO4 and 1.25M citric acid
(E0 = −0.18 V), achieving efficiencies of 29% and 75%,
respectively with all other parameters held identical.
Adding a mild amount of H2O2 reductant worked
synergistically to elevate the citric acid leaching efficiency
to approximately 99%, which is a common approach in the
literature. The leaching of NMC using a generic monoprotic
organic acid (HOA) is depicted in Equation (4):
6LiNi1/3Mn1/3Co1/3 O2 (s) + 18HOA(aq) + 3H2 O2 (l)
4.2
|
Organic acid leaching
Leaching with organic acids has drawn increasing interest
in the literature due to their biodegradability, lower acidity,
and minimal corrosion.115 On the other hand, inorganic
acids may be corrosive and release toxic/harmful gases (eg,
Cl2, NOx, SOx), and require additional wastewater processing costs (eg, neutralization and large volumes of water) to
avoid secondary pollution.116 The leaching efficiency of
→ 2Ni(OA)2 (aq) + 2Mn(OA)2 (aq)
+ 2Co(OA)2 (aq) + 6LiOA(aq) + 3O2 (g)
+ 12H2 O(l).
(4)
From Table 3, it is evident that high efficiencies can be
achieved with a variety of organic acids coupled with a
reducing agent. Among them, citric acid is a highly potent
leaching agent due to its triprotic nature, comprising three
14
|
OR ET AL.
carboxylic acid groups (pKa1 = 3.08, pKa2 = 4.74, and
pKa3 = 5.40), which can contribute to acidity and the
formation of strong chelation complexes. This explains its
higher efficacy compared to other compounds such as
maleic acid (diprotic, pKa1 = 1.94, and pKa2 = 6.22) and
acetic acid (monoprotic, pKa = 4.76).122,123 A comparison of
the chelation complexes that can be formed from these
compounds is shown in Figure 5C.124 While oxalic acid
possesses potent acidity (diprotic, pKa1 = 1.23, and pKa2 =
4.19), it exhibits poor leaching efficiency for Ni, Mn, and Co
due to the formation of a lowly soluble metal oxalate
(MC2O4) precipitate. 125 Oxalic acid leaching is typically
studied on LCO cathodes, as Li+ can easily be leached while
the oxalate counterion can serve as a precipitation agent to
recover Co.126 For mixed cathodes, oxalic acid has
generated interest as a selective Li+ leaching agent. 127
It is worth mentioning that due to the low acidity and
corrosivity of organic acids, they exhibit minimal
dissolution of Al foil (<10 wt% dissolution).102,103,128
Thus, separation of the cathode materials from the Al
current collector before leaching is unnecessary, and at
the same time the Al foil can be recovered at high purity
(Table 4).
5 | LEACHING DEVELOPMENTS
A N D UN C O N V E N T I O N A L
APPROACHES
5.1
|
Biohydrometallurgy
Biohydrometallurgy relates to the use of microorganisms
(ie, bacteria and fungi) that secrete various organic acids
(from the metabolic Krebs cycle) for metal leaching. The
main advantages of this method include the ecofriendly
composition and lowered material cost from substituting
the use of chemically synthesized acids.135 However, this
process requires long incubation/leaching times to
achieve good leaching efficiency (1‐2 weeks), high
liquid/solid ratio, and lengthy preparation steps to
culture the microorganism.136 Key parameters that must
be monitored and optimized in the bioreactor to
maximize organic acid production include the type and
concentration of carbohydrates (energy source for microbes) in the culture medium, pH, temperature, degree
of aeration, presence of toxic trace elements, and the
microbial population density.137,138
Bahaloo‐Horeh et al139 assessed the leaching of LIB
scrap comprised of Li, Ni, Mn, Co, Al, and Cu using
Aspergillus niger fungi, which is known to secrete citric,
gluconic, malic, and oxalic acid. They developed a
numerical model to predict the production of these acids
as a function of sucrose concentration, inoculum size,
and initial pH. Sucrose was found to be the most
important input parameter and was positively correlated
with the production of malic, citric, and gluconic acid,
while also negatively correlated with oxalic acid, which is
desired as oxalic acid produces precipitates with Co, Ni,
and Cu and thus lead to poor leaching efficiency. Under
optimized conditions, a leaching efficiency of 100% Cu
and Li, 77% Mn, and 75% Al was achieved at 50 mL/g
liquid/solid ratio, while 64% Co and 54% Ni were leached
at 100 mL/g liquid/solid ratio.
Xin et al explored the bioleaching of NMC, LMO, and
LFP cathodes using a combination of sulfur‐oxidizing
(Acidithiobacillus thiooxidans) and Fe2+‐oxidizing
A, Oxidation and deprotonation of ascorbic acid. B, Potential coordination complexes of M2+ cations with carboxylate
groups. C, Potential chelation complexes among acetic, maleic, and citric acid with monodentate and bridging bidentate coordination
depicted [Color figure can be viewed at wileyonlinelibrary.com]
FIGURE 5
HCl
HCl
HCl + H2O2
HCl
H2SO4 + H2O2
H2SO4 + NaHSO3
Leachate comprising Al3+, Cu2+,
Li+, Mn2+, and Co2+
Leachate comprising Fe2+/3+,
Cu2+, Li+, Ni2+, Mn2+, and Co2+
Leachate comprising Al3+, Li+,
Ni2+, and Co2+
2265 Li+, 75 Ni2+, 18 700 Mn2+,
32 730 Co2+
1000 Li+, 2000 Ni2+, 2000 Mn2+,
6560 Co2+
Leaching conditions
Leachate comprising Li+, Ni2+,
Mn2+, and Co2+
Selective precipitation
Leachate composition, mg/L
Purity: 90.25 wt% Co, 96.36 wt% Ni
1. Oxidative precipitation of Co2+ using NaClO (1
eq) as Co2O3 at pH = 3
2. Precipitation of Ni2+ as Ni(OH)2 at pH = 11
1. CoC2O4 precipitate formed using oxalic acid at
pH = 1.5 and 50°C for 2 h
2. Precipitate Mn2+ as MnCO3 at pH = 7.5
3. Precipitate Ni2+ as NiCO3 at pH = 9.0
4. Concentrate leachate (x2) and adjust pH to 14 to
precipitate Li2CO3
Joulie et al80
Li et al78
Barik et al79
Wang et al77
Reference
Efficiency: 98.94% Co, 89% Ni, 92%
Mn
Purity: 95.91% Co (coprecipitated
with 3.81% Ni and 0.28% Mn), 98%
Li
(Continues)
Meshram et al87
Efficiency: 94% Mn, 91% Ni, 95% Co, Nayl et al129
90% Li
Partial separation
Efficiency: 99% Cu
1. 99% of Cu2+ precipitated by displacement
reaction with Fe powder (1.5 eq)
2. Remaining Fe2+ in solution precipitated at
pH = 4 and 90°C.
3. Remaining Ni2+, Mn2+, and Co2+ in solution can
be used for direct resynthesis
1. Precipitate Mn2+ as MnCO3 at pH = 7.5
2. Similarly, adjust pH to 9.0 to form NiCO3
precipitate after 1 h
3. Form Co(OH)2 precipitate at pH = 11‐12 after 2 h
4. Precipitate remaining Li+ as Li2CO3
Efficiency: 97.7% Mn coprecipitated
with 26.9% Co
Efficiency: 80% Li, >99% Ni
Purity: 96.97% Li, 97.43% Ni, 98.23%
Mn, 96.94% Co
Purity and efficiency of recovery
1. Mn selectively precipitated using NaOCl (1.5 eq)
at pH = 1.5 for 30 min
2. Remove trace amounts of Al and Cu at pH = 4.5
and 5.5, respectively,
3. Remaining coprecipitated using Na2CO3
1. Adjust to pH = 2
2. Selectively oxidize Mn2+ with KMnO4 (0.5 eq) at
40°C
3. Adjust pH using NH3(aq) to form [Ni
(NH3)6]2+(aq) complex. Form red precipitate by
adding DMG (two eq) at pH = 9
4. Adjust pH to 0 with HCl, then to pH = 11 with
NaOH to break [Co(NH3)6]2+(aq) complex and
form Co(OH)2(s)
5. Remaining Li+ in leachate precipitated as Li2CO3
by adding saturated Na2CO3 at 100°C
Optimized protocol
T A B L E 4 Summary of approaches to separate and recover metals from a complex leachate by selective precipitation, selective oxidation, and/or solvent extraction
OR ET AL.
| 15
Leaching conditions
Efficiency: 98.6% Co, 99.9% Mn, and Dhiman and Gupta81
99.6% Li
1. Fe3+ precipitated at pH = 3.5 and 95°C for 2 h
2. Oxidize Mn2+ with (NH4)2S2O8 at pH = 4
(Continues)
|
11 Al3+, 31 Fe3+, 14 Cu2+, 570 Li+, HCl
385 Ni2+, 3148 Mn2+, 2880 Co2+
Tsakiridid and
Agatzini133
Partial separation
Efficiency: 99.9% Co and 99.7% Ni
coextracted with 0.2% Mn
H2SO4
Leachate comprising Li+, Ni2+,
Co2+, Mn2+, and Mg2+
1. Coextract Co2+ and Ni2+ using 20% Cyanex 301
at O/A = 1, pH = 2, and 50°C
2. Strip Co2+ and Ni2+ using 5M HCl at O/A = 2
and 50°C
H2SO4 + H2O2
Leachate comprising Fe3+, Al3+,
Cu2+, Li+, Ni2+, Mn2+, and Co2+
Partial separation
Joo et al131
2+
Efficiency: 93% Ni coextracted with
0.15% Mn, 0.23% Co, and 0.19% Li
Shih et al121
Chen et al84
Chen et al130
Reference
Partial separation
Joo et al132
Efficiency: 98.3% Mn coextracted
with 4.11% Co, 1.06% Ni, and 0.25%
Li
H2SO4 + H2O2
Leachate comprising Fe3+, Al3+,
Cu2+, Li+, Ni2+, Mn2+, and Co2+
Partial separation
Efficiency: 99.5% Mn coprecipitated
with 8% Ni and 3% Co
99.6% Co, 98.7% Ni, and 95.4% Cu
from solvent extraction
89% Li, 98% Ni, 97% Mn, 98% Co
Purity and efficiency of recovery
1. Fe3+, Al3+, and Cu2+ precipitated at pH = 4.8
2. Mn2+ selectively extracted using 0.43 M D2EHPA
and 0.7M Versatic 10 acid at O/A = 1
3. Impurities in organic phase removed using 0.1M
EDTA at O/A = 4
4. Mn2+ stripped using 0.5M H2SO4 at O/A = 2
1. Mn2+ selectively precipitated using KMnO4 at
pH = 3 for 1 h
2. Ni2+, Co2+, and Cu2+ separated from Li+ using
D2EHPA solvent extraction at O/A = 1.5, pH = 6,
and 45 min
1. Fe3+, Al3+, and Cu2+ precipitated at pH = 4.8
2. Ni2+ selectively extracted using 0.23M LIX 84‐I
and 1.41M Versatic 10 acid at O/A = 1, pH = 5,
and 25°C
H2SO4 + H2O2
Leachate comprising Al3+, Cu2+,
Li+, Ni2+, Mn2+, and Co2+
1. Ni2+ precipitated using DMG (two eq) at pH = 6,
then recovered as NiCl2 by HCl dissolution.
2. Co2+ precipitated as CoC2O4 using (NH4)2C2O4
(1.2 eq) at pH = 6 and 55°C
3. Mn2+ solvent extracted using D2EHPA then
stripped with H2SO4
4. Remaining Li+ precipitated as Li3PO4
1. Precipitate Fe3+ as Fe2O3 by adjusting pH using
NaOH
2. Selectively extract Cu2+ using commercial agent
(Mextral 5640H)
3. Oxidative precipitation of Mn2+ using KMnO4
4. Selectively extract Co2+ using commercial agent
(nickel loaded Mextral 272P)
5. Remaining Ni2+ and Li+ precipitated as Ni(OH)2
and Li3PO4, respectively
Citric acid
Optimized protocol
H2SO4 + H2O2
1960 Fe3+, 1780 Cu2+, 1490 Li+,
4290 Ni2+, 5680 Mn2+, 7180 Co2+
1248 Li+, 4930 Ni2+, 5269 Mn2+,
5837 Co2+
Solvent extraction and combined processes
Leachate composition, mg/L
T A B L E 4 (Continued)
16
OR ET AL.
3+
Efficiency: 99% Li, >99% Ni, >98%
Mn, 99.9% Co
Purity (wt%): 98.7% Li, ~100% Ni,
99.7% Mn, 99.9% Co
1. Mn2+ selectively oxidized using KMnO4 (1.2 eq)
at pH = 2.5 and 80°C
2. Ni2+ precipitated using DMG (2 eq) at pH = 5
and 80°C
3. Two‐step precipitation of Co2+ using Cyanex 272
(50% saponified, 0.64M O/A = 1, pH = 5) and
stripped with H2SO4
4. Li+ precipitated as Li2CO3 (1.2 eq Na2CO3) at
pH = 12 and 90°C
3790 Li, 9660 Ni, 9560 Mn, 10 090
Co
Note: Conditions at room temperature (~25°C) unless specified otherwise.
H2SO4 + H2O2
Citric acid + H2O2
Co2+ recovery only
Efficiency: 98% Co2+
Purity: 99% Co2+
2930 Li+, 5861 Ni2+, 4097 Mn2+,
6319 Co2+
1. Fe3+ precipitated at pH = 3‐3.5 and 95°C for 2 h
2. Mn2+ selectively oxidized using (NH4)2S2O8 (1.8
eq) at pH = 4 and 70°C
3. Cu2+ precipitated at pH = 5.5
4. Co2+ extracted with 25 wt% P507 at pH = 3.5, O/
A = 1.5, and stripped with H2SO4
5. Co2+ precipitated using (NH4)2C2O4 (pH = 1.5,
1.15 eq)
Purity and efficiency of recovery
3600 Fe3+, 1800 Cu2+, 2500 Li+,
H2SO4 + H2O2
500 Ni2+, 1800 Mn2+, 20 600 Co2+
3. Precipitate Cu and Al at pH = 5.5
4. Solvent extract Co2+ using 0.2M Cyphyos IL 102,
O/A = 1, followed by stripping using 0.05M HCl
at O/A = 1
5. Adjust pH to 9 using NH3 and precipitate Ni2+
with DMG at 50°C‐60°C
6. Precipitate remaining Li+ as Li2CO3 at pH = 11
and 100°C
2+
Optimized protocol
Efficiency: 89% Li+, 98% Ni2+, 95%
Mn2+, >95% Co2+
Purity: 99.07% Li+, 98.46% Ni2+,
98.47% Co2+
Leaching conditions
1. Ni2+ selectively precipitated with DMG (two eq,
pH = 8), then stripped with and recovered as
hydroxide precipitate
2. Co2+ selectively precipitated using (NH4)2C2O4
(1.2 eq, pH = 6). Coprecipitated Mn2+ scrubbed
with dilute oxalic acid.
3. Mn2+ extracted with Na‐D2EHPA (20 vol%,
pH = 4, O/A = 2) and stripped with H2SO4
4. Remaining Li+ precipitated as Li3PO4
Leachate composition, mg/L
T A B L E 4 (Continued)
Sattar et al90
Chen et al134
Chen and Zhou95
Reference
OR ET AL.
| 17
18
|
OR ET AL.
(Leptospirillum ferriphilum) chemotrophic bacteria with a
food source of elemental sulfur and pyrite respectively.140
L. ferriphilum can produce H2SO4 from the biooxidation
of elemental sulfur, while A. thiooxidans catalyzes the
release of Fe2+ from pyrite, which can be used to reduce
Mn4+ and Co3+ (from Fe2+/Fe3+ redox) and is essential
to achieve high leaching efficiencies.141 After optimizing
the pH to enhance microbial growth, more than 95%
leaching efficiency for Li, Ni, Mn, and Co was obtained at
100 mL/g liquid/solid ratio.
5.2
|
Mechanochemical pretreatment
Mechanochemical methods have been explored as a
pretreatment step before leaching. This involves the
cogrinding of cathode active material with reducing agents
to break and deform the crystal structure, decrease particle
size, and generate products with decreased activation
energy and improved reactivity for leaching.142 As a result,
the cathode metals can be leached at mild concentrations
and low liquid/solid ratios. In some work, the mechanochemical reaction produces a soluble Li product that can
be selectively leached. This is interesting, as in the
leaching techniques reviewed thus far, all metals in
mixed‐type cathodes are dissolved. Recovering high‐
purity Li (>99.5%, battery grade) that is appropriate for
cathode synthesis from a complex leachate comprising
Al3+, Fe3+, Cu2+, Li+, Ni2+, Mn2+, Co2+, and other metals
can be challenging and time‐consuming. Thus, selective Li
leaching can prioritize its recovery at high purity. Yang
et al143 demonstrated a mechanochemical activation route
to selectively de‐lithiate a LiNi0.5Mn0.3Co0.2O2 cathode.
Here, the active material was ball milled with hydrated
Na2S, generating LiOH, Na2SO3, and Ni0.5Mn0.3Co0.2(OH)2
as products. The selective breakage of the Li–O bond in
the cathode is ascribed to its relatively low bond
dissociation energy. Thus, Li could be selectively leached
(95.1% efficiency) in deionized water at room temperature
due to the low solubility of Ni0.5Mn0.3Co0.2(OH)2, and
subsequently recovered as Li2CO3 precipitate at 99.96 wt%
purity. Li2CO3 and Ni0.5Mn0.3Co0.2(OH)2 could be used as
reagents to resynthesize the cathode material as reviewed
in Section 7.
Zhang and Hu et al88 demonstrated a thermomechanochemical approach to selectively extract Li.144 In this
approach the cathode active material (NMC and LCO)
was ball milled with lignite carbon (20 wt%) and
subsequently calcined at 650°C for 3 hours. This process
reduced NMC and LCO and converted it mostly into
Li2CO3, Ni, Co, and MnO. To selectively leach Li, Li2CO3
was converted into a soluble compound (LiHCO3)
through a carbonation reaction. The dissolved Li+ was
then filtered and recovered by heating the solution at
100°C to regenerate the Li2CO3 precipitate. They subsequently demonstrated how Ni, Co, and MnO can easily
be leached in H2SO4 where more than 96% efficiency was
achieved at a low liquid/solid ratio (3.5 mL/g) and
without requiring a reducing agent, as the metals already
exist at low valence states. Through a similar concept, Liu
et al89 ball‐milled NMC with carbon black (10 wt%
dosage) and calcined the powder at 550°C for 0.5 hours
under Ar flow, producing Li2CO3, Ni, Co, NiO, and MnO.
Due to the low solubility of Li2CO3, 93% of Li was
selectively leached with water at a relatively high liquid/
solid ratio of 30 mL/g at 25°C for 1.5 hours. The other
metals were then leached with 4M H2SO4 at 10 mL/g,
90°C, and 0.5 hours, producing efficiencies of more than
99.5%. Alternatively, Guan et al92 used iron powder as the
reducing agent that was cogrinded with the cathode
active materials. X‐ray photoelectron spectroscopy characterization confirmed that the mechanochemical reaction reduced Co3+ to Co2+ and oxidized Fe to Fe3+. As a
result, leaching efficiencies of 77.15% Li, 91.25% Co, 100%
Mn, and 99.9% Ni were achieved using 1M HNO3 at
333 mL/g, 25°C, and 2 hours, whereas without Fe
reduction, only 39% Li, 38.67% Ni, 33.19% Mn, and
20.43% Co efficiency could be obtained under identical
leaching parameters. However, a disadvantage of this
approach is the large quantity of Fe impurity introduced
into the leachate.
5.3
|
Ammoniacal leaching
Although leaching with basic ammoniacal solutions is
known to be less efficient than acid leachants, thus
requiring higher concentrations, leaching times, and
liquid/solid ratios, this approach has generated interest
due to the observed selectivity toward Li, Ni, Co, and Cu
in addition to ammonia/ammonium being a more
environmentally friendly reagent. The ammoniacal solution generally comprises ammonia as the leaching agent,
an ammonium compound as the buffering agent, and a
sulfite or thiosulfate compound as the reducing agent.
Selective leaching is driven by the formation of stable
coordination complexes between ammonia and metallic
ions, particularly Li+, Ni2+, Co2+, and Cu2+, whereas Fe,
Mn, and Al typically remain precipitated as hydroxides.
145
Although leaching with ammonia alone is thermodynamically favorable, the reaction has poor kinetics, and
thus minimal leaching is observed for Li, Ni, and Co even
at high concentrations.146-148 However, Cu can easily be
leached due to the high formation constant of the Cu
(NH3)42+ complex.147 The ammonium compound lowers
the pH of the leachate and generates a buffer solution
| 19
OR ET AL.
(NH3/NH4+ buffer range: 8.25‐10.25) that is necessary to
form and stabilize the metal‐ammonia complexes.
Similar to acid leaching, the reducing agent lowers the
metal valence state and energy barrier of dissolution and
is essential to reach high leaching efficiencies.
Zheng et al149 assessed the leaching of cathode active
material powder in ammoniacal solution, where under
optimized conditions of 4M NH3, 1.5M (NH4)2SO4
buffering agent, 0.5M Na2SO3 reducing agent, liquid/
solid ratio 100 mL/g, and 80°C for 5 hours, leaching
efficiencies of 95.3% Li, 89.8% Ni, 80.7% Co, and merely
4.3% Mn were obtained. Using a sulfite reducing agent
decreased the amount of Mn leached, as this generated
(NH4)2Mn(SO3)2∙H2O as a precipitation product. As a
result, selectivity against Mn was improved—the leachate
purity comprised 98.6% Li+, Ni2+, and Co2+ and only
1.4% Mn2+. Several studies have noted that high leaching
efficiencies (>95%) for all metals can be obtained using a
two‐step process where the solid residue is filtered and
leached again using identical conditions.149,150 Meng
et al150 proposed that the second step is required due to
the sluggish leaching kinetics caused by the formation of
(NH4)2Mn(SO3)2∙H2O/Mn3O4 precipitate coated on the
active material particles. To avoid the use of sodium
sulfite and subsequent formation of Mn precipitate,
Wang et al151 pretreated LiNi0.5Mn0.3Co0.2O2 through a
thermomechanochemical reduction reaction with graphite, similar to the approaches described previously. The
reduced products were then easily leached with ammoniacal solution (NH3, NH4HCO3, and H2O2) at high
efficiencies (81.2% Li, 96.4% Ni, and 96.3% Co). A key
disadvantage of ammoniacal leaching is that the formation of stable metal‐ammonia complexes can impede the
separation and recovery of metals through selective
precipitation and solvent extraction as described in the
following sections.
6 | SELECTIVE M ETAL
RECOVERY FROM LEACHATE
From a leachate comprised mostly of Li+, Ni2+, Mn2+,
Co2+, Fe3+, Al3+, and Cu2+, the metals can be selectively
recovered through a series of precipitation and/or solvent
extraction steps. When spent LIBs are pretreated by
crushing and sieving, various quantities of Fe3+, Al3+,
and Cu2+ are present in the leachate as impurities from
the current collectors and battery casing. Thus, they are
first removed by coprecipitation or coextraction. While
Al3+ and Fe3+ are sparingly soluble in aqueous solution,
Cu2+ may cause coprecipitation/coextraction complications with Ni2+, Mn2+, and Co2+. On the other hand,
LIBs pretreated by lab‐scale approaches often produce
little to no detectable impurities, which simplifies the
subsequent recovery. It should be noted that Fe in the
leachate is likely to present as Fe3+, as Fe2+ can be
oxidized by H2O2 in an acidic solution or other high‐
valence metals (M3+/M4+).152
6.1
|
Selective precipitation
Valuable metals can be recovered from the leachate
through a series of pH adjustments to form metal
hydroxide precipitates, reactions with anionic weak bases
to form sparsely soluble products (eg, carbonates,
phosphates, and oxalates), and/or selective oxidation
reactions to form metal‐oxide precipitates. Recovered
metallic salts may then be thermally treated to produce
crystalline metal oxides (eg, Co3O4), which can be used as
reagents for solid‐state reactions.153,154 In general,
regarding the precipitation of salts, the leachate temperature and pH are important factors to optimize their
solubility. The solubility of salts possessing a basic
counter‐anion is inversely proportional to pH as seen in
Equation (5) using the example of an M(II) carbonate.
MCO3 (s) + H2 O(l) ⇌ M2+(aq) + CO32−(aq)
+ H2 O(l) ⇌ M2+(aq)
+ HCO−3 (aq) + OH−(aq).
(5)
While an increase in the leachate pH will improve the
overall recovery efficiency, it may cause coprecipitation of
different metals. It is worth mentioning that at significantly
high pH, soluble anionic species can form (eg, Cu(OH)3−,
Al(OH)63−). In practice, pH adjustment of the leachate is
performed by adding concentrated NaOH. Use of a weak
base such as NH3 (NH4OH) requires large volumes to alter
the pH and may form stable soluble complexes with
metallic ions (eg, excess NH3 forms [M(NH3)6]2+ complex).
Temperature also has significant impacts on solubility
thermodynamics. While precipitation is usually an
entropically favorable process due to the net release of
water in the metallic ion solvation shell, the dissolution
of many salts is an endothermic process. Thus, some salts
require temperature optimization (eg, CoC2O4) to achieve
high recovery efficiency. Metal carbonates and hydroxides are notable examples where its dissolution is
exothermic, and thus its solubility is typically inversely
proportional to temperature.155-157 Figure 6 highlights the
decrease in solubility of common metal hydroxides and
carbonates in LIB leachates as a function of pH and
temperature.131 The graphs were generated based on the
KSP (25°C) of the compounds and standard enthalpy of
dissolution.158 While this can be useful to predict the
20
|
OR ET AL.
FIGURE 6
Solubility of common metal hydroxides in LIB leachate as a function of (A) pH and (B) temperature. C and D, Trend in
solubility of metal carbonates used in selective precipitation. LIB, lithium‐ion battery [Color figure can be viewed at wileyonlinelibrary.com]
relative order of precipitation in a mixed‐metal leachate,
the accuracy is affected by notable factors, such as how
different anions in the leachate can produce various
interactions with the metallic cations, concentration‐
dependent ion interactions, assumptions for the activity
coefficient, and the possibility of forming different
precipitate compositions (eg, Fe3+ precipitated as jarosite
in sulfate media).134,152 Due to the relatively large KSP of
Cu(OH)2, it has the highest risk of coprecipitation with
Ni(OH)2 and Co(OH)2. However, Figure 6A indicates that
at pH = 6.5 and 25°C, Cu(OH)2 theoretically has very low
solubility such that [Cu2+] = 3 mg/L (ppm) while
[Ni2+] = 32 000 mg/L. To assess the suitability of this
approach, Gratz et al159 dissolved a mixture of Ni2+
(13.54 g/L), Mn2+ (12.26 g/L), Co2+ (12.32 g/L), and Cu2+
(989.5 mg/L) in H2SO4/H2O2 leachate media. When the
pH was raised from 1 to 6.5, the degree of precipitation
observed was 4%, 1%, 4%, and 96.4%, respectively.
However, in the absence of Cu2+ and with all other
conditions identical, negligible precipitation of Ni2+,
Mn2+, or Co2+ was observed. This indicates that
selectively precipitating Cu2+ (along with Al3+ and
Fe3+) by pH adjustment is effective when its concentration is low relative to other metals. Cu(OH)2
induces coprecipitation of Co and Ni likely through the
formation of Cu–Co and Cu–Ni complexes that are not
accounted for in Figure 6A.152,160,161 To illustrate this,
Suzuki et al161 dissolved equivalent molar concentrations
of Al3+ (54 mg/L), Cu2+ (127.1 mg/L), and Co2+
(117.9 mg/L) in sulfate media and observed 100% Al3+
and Cu2+ precipitation at pH = 7 along with 50% of Co2+
coprecipitated, which deviated significantly from modeled precipitation curves.
From a sulfuric acid leachate of spent commercial
LIBs pretreated by crushing and sieving, Kang et al
reported an approach to remove large quantities of
Al3+ (1800 mg/L), Cu2+ (782.7 mg/L), and Fe3+
(159.5 mg/L) first by adjusting the leachate pH to 6.5
using NaOH and CaCO3, resulting in significant
coprecipitation among all metals. The precipitate was
washed with water to redissolve valuable metals that
were returned to the leachate, resulting in more than
99% removal of impurity ions albeit with notable losses
in valuable metals (−6.7% Co, −14.9% Mn, −19.4% Ni,
and −1.6% Li).162 On the other hand, under similar
pretreatment and leaching conditions, Joo et al131,132
precipitated (~100%) small quantities of Al3+ (830 mg/
L), Cu2+ (5.6 mg/L), and Fe3+ (10.4 mg/L) simply by
adjusting the leachate pH to 4.8 with concentrated
NaOH, with negligible coprecipitation of Ni, Mn, and
Co. These studies highlight the importance of the
pretreatment process to effectively separate impurity
metals from valuable components.
Due to the similar solubility of Ni(OH)2, Mn(OH)2,
and Co(OH)2 (Figure 6), it is challenging to achieve
separation from pH and temperature adjustment alone,
| 21
OR ET AL.
so other precipitates with improved selectively are
formed. For instance, dimethylglyoxime (DMG,
C4H8N2O2) is an analytical chelation reagent with high
selectivity toward Ni2+, forming a 2:1 DMG:Ni2+ complex
and a red precipitate in basic conditions as depicted in
Equation (6)163:
Ni2+(aq) + 2C4 H8 N2 O2 (aq) + 2OH−(aq)
→ Ni‐(C4 H7 N2 O2)2 (s) + 2H2 O(l).
(6)
Ni2+ can then be stripped as NiCl2 using concentrated HCl and the recovered DMG (white powder)
reused. Chen and Zhou selectively precipitated Ni2+
from a citric acid leachate at pH = 8 and DMG/Ni2+
(molar ratio) = 2.95 Minimal coprecipitation with Li+,
Mn2+, or Co2+ (<1%) was observed, possibility due to
the stable complexes already formed with citric acid/
citrate. Similarly, Sattar et al90 separated large and
similar concentrations Ni2+ (9.66 g/L) and Co2+
(10.09 g/L) in sulfuric acid leachate by DMG precipitation (~100% Ni2+ purity) using pH = 5 and DMG/
Ni2+ = 2. DMG can also precipitate Ni2+ in NH3
solution, where Ni2+ and other metals exist as a soluble
[Ni(NH3)6]2+ complex, which can potentially be
exploited to further enhance its selectivity.77
Oxalate ions are known to be excellent ligands for M2+
ions due to their dibasic characteristics. They are highly
selective toward Co2+ as precipitation can be achieved
rapidly at a low pH range (1‐1.5).134 It is worth
mentioning that CoC2O4 is a commonly used precursor
to generate Co3O4 catalysts or metallic Co powder by
decomposition. In most work, the pH and C2O42−:Co2+
ratio (slight excess desired) is optimized to maximize
Co2+ recovery efficiency while avoiding Mn2+ and Ni2+
coprecipitation. From a sulfuric acid leachate Meshram
et al,87 recovered Co2+ first via CoC2O4 precipitation at
pH = 1.5 and 50°C, resulting in an approximately 99%
recovery efficiency and a 95.91% purity (coprecipitated
with 3.81% Ni2+ and 0.28% Mn2+). As carbonate
precipitation is more selective than hydroxide precipitation in separating Ni2+ and Mn2+ (Figure 6C), the
leachate pH was adjusted to 7.5 and mixed with
concentrated Na2CO3 to generate MnCO3 (92% efficiency), and finally to pH = 9.0 to precipitate NiCO3
(89% efficiency). However, the metal purity of the
carbonate precipitates was not explicitly addressed.
Similarly, Nayl et al129 treated sulfuric acid leachate with
Na2CO3 to form MnCO3 (94% efficiency) and NiCO3 (91%
efficiency) at pH = 7.5 and 9.0, respectively, although the
purity was not specified.
Alternatively, ions can be selectively oxidized to form
metal‐oxide precipitates with high purity. The oxidation
of Mn2+ using permanganate (MnO4−) is depicted in
Equation (7):
3Mn2+(aq) + 2MnO−4 (aq) + 2H2 O(l) → 5MnO2 (s)
+ 4H+(aq)E0 = 0.45V.
(7)
MnO4− has a high reduction potential, making it
favorable for oxidative precipitation of Mn2+ into MnO2,
Mn3O4, or MnOOH (s).164 This reaction is thermodynamically selective over Co2+ and Ni2+ as seen in Equations
(8) and (9). 165
3Co2+(aq) + 2MnO−4 (aq) + 2H2 O(l)
⇌ 3CoO2 (s) + 2MnO2 (s) + 4H+(aq)
E0 = 0.01V,
(8)
3Ni2+(aq) + 2MnO−4 (aq) + 2H2 O(l)
⇌ 3NiO2 (s) + 2MnO2 (s) + 4H+(aq)
E0 = −0.02V.
(9)
From these equations, it is evident that the leachate
pH influences the thermodynamic favorability of the
reactions. A higher pH range (3‐4) enhances Mn
precipitation efficiency, but may introduce Co and Ni
coprecipitation, thus requiring optimization. Similarly,
ammonium persulfate ((NH4)2S2O8) has been demonstrated as a selective oxidant for Mn2+ (Equation (10)),
where the pH and S2O82−:Mn2+ ratio were optimized to
minimize coprecipitation with Co, while Ni coprecipitation was not observed.134
Mn2+(aq) + (NH 4)2 S2 O8 (aq) + 2H2 O(l)
→ MnO2 (s) + (NH 4)2 SO4 (aq) + H2 SO4 (aq)
+ 2H+(aq)E0 = 0.89V.
(10)
The oxidative precipitation approach demonstrates
good selectivity with the main drawback being the high
reagent cost of the oxidants.166 It is not ideal in HCl
leaching media, as the Cl− ions could participate in a
competing reaction by reducing the oxidants or high‐
valence metals, evolving Cl2(g) and precipitates in the
process (Equation (11)).167
2MnO−4 (aq) + 8H+(aq) + Cl−(aq) → 2MnO2 (s)
+ 3Cl2 (g) + 4H2 O(l).
(11)
Finally, Li+ in the leachate is often precipitated last as
it is highly soluble and largely unaffected by pH
adjustments. Li+ is reacted with saturated Na2CO3 or
22
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OR ET AL.
Na3PO4 to form Li2CO3 or Li3PO4 precipitate, respectively, at high pH (~11). As Li2CO3 is relatively soluble
(Figure 6C), precipitation is often aided by concentrating
the Li+ in the leachate and using elevated temperatures
(~100°C).157 The high solubility of Li+ can be exploited to
selectively recover high‐purity (battery grade) Li2CO3,
which is ideal as a cathode synthesis reagent. After
leaching, Gao et al adjusted the pH of the solution to 11
to form a Ni‐Mn‐Co hydroxide coprecipitate.103,128
Following filtration, the remaining Li+ in the leachate
was precipitated as Li2CO3 with purity more than 99.9%.
6.2
|
Solvent extraction
Solvent extraction relates to the selective transfer of
metallic ions from aqueous to organic phase aided by
complexation agents (extractants). Extraction compounds
comprise of an active component that forms coordination
complexes with metallic ions and a modifier component
that aids their miscibility in the nonpolar organic phase
(eg, kerosene or toluene). The metal is “stripped” from
the organic phase in strong acid, where the extractant is
recycled and the metal can then be physically recovered
from aqueous solution through precipitation or reduction. This approach is suitable for industrial‐scale
applications, as the main operating cost is the consumption of reagents for pH adjustments used to optimize the
selectivity and recovery efficiency. As solvent extraction
relies on the formation of metal‐chelate complexes,
the counter‐anion in the acidic leachate can influence
the extraction properties. Notably, Cl− anions are
stronger inner‐sphere ligands compared to SO42− anions,
and as such the metal‐anion complex will influence
the desirable choice of extractant.161,168 For instance,
when considering the pH extraction isotherm of
2‐ethylhexylphosphonic acid mono‐2‐ethylhexyl ester
(commercial name PC88A), Mn is extracted before Co
in sulfate media, while the opposite occurs in chloride
media.169 Dhiman and Gupta81 compared the extraction
of Co2+ using Cyphos IL 102 in HCl, HNO3, and H2SO4
media, showing that extraction efficiency increased
drastically with respect to (HCl), whereas extraction in
H2SO4 and HNO3 was poor at all concentrations. This
indicates that formation of a Co2+‐anionic‐chloro species
is critical toward the extraction mechanism of Cyphos IL
102. To the best of knowledge, only a few works have
tested solvent extraction in organic acid media, which
can form metal chelates (Figure 5).95,130 Although the
work demonstrated good extraction efficiency of Mn2+ in
citric acid media using di(2‐ethylhexyl) phosphoric acid
(D2EHPA), the role of various organic ligands on
extraction behavior has not been delineated. For instance,
the strong interactions between Co2+ and citric acid could
influence the choice of extractant. This section aims to
review solvent extraction theory and strategies to separate a
complex leachate comprised of common metals in LIB
scrap (Mn, Co, Ni, Li, Fe, Cu Al). All studies reviewed here
have employed H2SO4/H2O2 as the lixiviant.
Bis(2,4,4‐trimethylpentyl) phosphinic acid (Cyanex
272) is a common commercial extractant that traditionally has been used to separate Ni and Co from
leached laterite ore.170,171 Metal extraction is based on
a cation exchange mechanism and thus is often
saponified (Equation (12)) to improve efficiency, as
the Na+ counterion is more easily displaced by the
metal.
Na+(aq) + ½(HA)2 (org) → NaA(org) + H+(aq).
(12)
HA and NaA refer to the “acid‐form” and “sodium‐
form” extractant respectively. (HA)2 indicates that Cyanex
272 forms a hydrogen‐bonded dimeric structure in nonpolar
solvents. The extraction of M2+ and M+ (ie, Li+) is depicted
in Equations (13) and (14), respectively172,173:
M2+(aq) + A−(org) + 2(HA)2 (org) ⇌ MA2n3HA(org)
+ H+(aq),
(13)
M+(aq) + A−(org) + (HA)2 (org) ⇌ MA ⋯ 2HA(org).
(14)
The equations show a 2:1 reactant ratio for (MA)2:M2+
and 1:1 for (MA)2:M+, indicating a stronger chelating
affinity for M2+ over Li+. For M2+ extraction, the
corresponding extraction equilibrium constant (Ke) is:
Ke =
[MA2 ∙3HA][H+]
.
[M 2+][A− ][(HA)2]2
(15)
[MA ∙ 3HA]
2
)
Defining the distribution coefficient (KD = [M
2+]
as the distribution of the metal between the organic and
aqueous phase (ie, extraction efficiency), Ke can be
represented in Equation (16):
Ke =
KD [H+]
.
−
[A ][ ( HA) 2]2
(16)
By taking the logarithm of Equation (16) and
rearranging, the M2+ extraction equation is reexpressed
in Equaion (17):
logKD = logK e + log[A− ] + 2log[(HA)2] + pH. (17)
| 23
OR ET AL.
Through a similar derivation, the extraction of M+ is
expressed in Equation (18):
logKD = logK e + log[A− ] + log[(HA)2].
(18)
These equations model the relationship among
metal extraction efficiency, pH, and the concentration
of extractant. In general, extraction efficiency (KD )
increases with respect to pH and extractant concentration (for a single extractant). Equation (18) indicates that Li + extraction is constant with respect to
pH, which makes the separation of Li+ from M 2+
using Cyanex 272 a facile method. Separation of
different M2+ cations can be achieved by optimizing
the pH and [(HA)2 ]. Figure 7 shows experimental
results from Cyanex 272 extraction that confirm the
pH‐independent extraction of Li+ (Figure 7C) and the
2:1 chelation of (HA) 2:M 2+ (Figure 7D).
The equilibrium pH at which 50% of a metal can be
isothermally extracted is known as the pH50. Selectivity
of one metal (M1) over the other (M2) as a function of pH
is indicated by their difference in pH50 (∆pH50(M1 −M2) )
and more generally by the separation factor (ß) as defined
in Equations (19) and (20), respectively174:
∆pH50(M1 −M2) = pH50(M1) − pH50(M2).
β=
KD (M1)
KD (M2)
.
(19)
(20)
As a general guideline, ∆pH50(M1 −M2) should be more
than one to achieve good separation. Equations (21) and
A, Selective extraction of Cu2+, Fe3+, and Al3+ with Ancorga M5460 by pH optimization. B, Cyanex 272 extraction pH isotherm.
C and D, Analysis of Cyanex 272 extraction mechanism based on Equation (24). Adapted with permission: Copyright 2015, Elsevier172
FIGURE 7
24
|
OR ET AL.
(22 represent a more general expression for a divalent
metal in sulfate and chloride media, respectively,
assuming the anions form stable ligands with the metal.
The equations also assume that the extractant forms a
dimer in an organic solvent, which is the case for
phosphoric, phosphonic, phosphinic, and carboxylic
acids.121,175
M2+(aq) + xSO2−
4 (aq) + mH2 A2 (org)
(org) + nH+(aq), (21)
⇌ M(SO4 ) x (A2H2‐n)(2‐2x‐n)+
m
M2+(aq) + xCl−(aq) + mH2 A2 (org)
(org) + nH+(aq),
⇌ M(Cl) x (A2H2‐n)(2‐x‐n)+
m
Ke =
KD [H+]n
,
[SO4 2− /Cl−]x [H2 A2]m
(22)
(23)
logKD = logK e + m log[H2 A2]
+ x log[SO4 2− /Cl−] + n pH.
(24)
The general extraction equation is shown in Equation
(24) where the coefficients m, x, and n can be determined
empirically to analyze the extraction mechanism as shown in
Figure 7.
Nayl et al172 selectively purified Li, Ni, Mn, and Co from
an H2SO4/H2O2 leachate comprised of Li+, Ni2+, Mn2+,
Co2+, Fe3+, Al3+, and Cu2+ using a series of solvent
extraction and selective precipitation steps. Ancorga M5460
(commercial extractant comprised of 5‐nonsalicylaldoxime as
the active substance known to have an affinity toward Cu2+
and 2,4,4‐trimethyl‐1,3‐pentanediol diisobutyrate as the
modifier component) was first used to remove impurity ions
(Cu2+, Fe3+, and Al3+; Figure 7A). Using an oil/water
volumetric ratio (O/A) of unity, 99.2% of Cu2+ was extracted
at pH = 1 while 94% Fe3+ and 95.6% Al3+ was extracted at
pH = 2.0 to 2.2. The remaining ions in the aqueous solution
were then treated with saponified Cyanex 272 (0.04M;
pH = 5; O/A = 1) to extract 81% Co2+ and 92.5% Mn2+ along
with 5.94% Ni2+ and 10% Li+ (Figure 7B). The Ni2+ and Li+
impurities in the organic phase were removed by precipitation using dilute Na2CO3 and returned to the mother leach
liquor. Co2+ and Mn2+ were then stripped from the
extractant using H2SO4, where Mn2+ was precipitated as
MnCO3 (99.7% purity) at pH = 7.5 by adding saturated
Na2CO3, while the remaining Co2+ was precipitated as Co
(OH)2 (>99% purity) by increasing the pH to 11 with NaOH.
From the mother leach liquor, saturated Na2CO3 was added
to recover NiCO3 (99.4% purity) at pH = 9 and Li2CO3 (99.6%
purity) at pH = 11 to 12. This work demonstrates the
feasibility of recovering valuable metals in high efficiency
and purity from a complex LIB leach liquor.
Besides Cyanex 272, other commonly used extractants
such as D2EHPA show selectivity toward M2+ metals over
Li+, making this approach ideal for Li recovery.121
However, the individual separation of Co2+, Mn2+, and
Ni2+ can be challenging when using solvent extraction
alone, for instance, the |ΔpH50 (Co–Mn)| is small for PC88A,
D2EHPA, or Cyanex 272, resulting in significant coextraction.169,176,177 An approach to address this is to use a
mixture of different extractants to improve selectivity. The
extractants can form interactions described either as a
mixed complex formation or reverse micelle aggregation,
resulting in synergistic or antagonistic extraction efficiency
for certain metals.178 Pranolo et al179 used a mixture of
Ionoquest 801 and Ancorga M5640 to selectively remove
LIB leachate impurities. Ionoquest 801 alone demonstrated
an extraction preference order of Fe3+ >> Al3+ > Cu2+ >
Co2+, Ni2+, Li+ as a function of pH with unavoidable
coextraction of Cu2+ and Co2+. The addition of Ancorga
M5640 downshifted the pH extraction isotherm of Cu2+
substantially (ΔpH50 (Cu) = 3.45), allowing for selective
removal of all impurity ions at pH = 4 to 4.5. Joo et al131
used a mixture of 2‐hydroxy‐5‐nonacetophenone oxime
(LIX 84‐I) and Versatic 10 acid to selectively extract Ni2+
from LIB scrap leached with H2SO4. They first adjusted the
leachate pH to 4.8 to precipitate Fe3+, Al3+, and Cu2+
impurities, with no detectable coprecipitation of Ni2+, Co2+,
and Mn2+. Compared to using LIX 84‐I alone, the addition
of Versatic 10 downshifted the pH extraction isotherm for
Ni2+ but upshifted the isotherm for Co2+ in a
concentration‐dependent manner. The mixture of both
extractants clearly produced a synergistic enhancement on
the extraction of Ni2+ (ie, KD (Mix) − (KD (A) + KD (B)) > 0 for
arbitrary extractants A and B). This resulted in good
selectivity for Ni2+ (ΔpH50 (Co–Ni) = 1.89 and ΔpH50
(Mn–Ni) = 2.16) under optimized conditions (0.23M LIX 84‐
I, 1.41M Versatic 10 acid, O/A = 1, pH = 5, and 25°C). The
research group also explored the use of D2EHPA and
Versatic 10 acid to separate Mn2+ from the same leachate
composition.132 Versatic 10 acid appeared to disrupt the
extraction mechanism between D2EHPA and Co2+/Mn2+
in a concentration‐dependent manner. Using 0.43M
D2EHPA, 0.7M Versatic 10 acid at O/A = 1, the excellent
separation between Co2+ and Mn2+ was achieved (ΔpH50
(Co–Mn) = 5.5). The ßMn/Co calculated was 33.97 compared to
merely 14.3 when using D2EHPA alone.
7 | CATHODE RESYNTHESIS
D I RE C T L Y F R O M L E AC H AT E
From the prior discussion, it is clear that separating and
recovering cathode metals in high purity from a complex
leachate can be challenging and time‐consuming,
| 25
OR ET AL.
requiring many steps, such as pH adjustment, filtering,
washing, and concentration adjustments. Thus, researchers have recently proposed resynthesizing mixed cathodes directly from the leachate.
The first step involves removing impurity metals
(mostly Cu2+, Al3+, and Fe3+), which will be prevalent
in the leachate if the cells are pretreated by crushing and
sieving. As discussed in Section 6, this can be achieved by
precipitation through pH adjustment or solvent extraction. After impurity removal, the leachate concentration
is measured and adjusted by adding the appropriate
reagents to achieve the desired stoichiometric ratio for
the cathode material. The cathode is then synthesized
through a sol‐gel or coprecipitation method.
It should be noted that trace amounts of impurity
metals can be beneficial toward electrochemical performance as a dopant. Although Al is electrochemically
inert, there is a notable amount of literature demonstrating how it enhances the capacity retention, thermal
stability, and rate performance of LIB cathodes by
stabilizing the lattice structure.180-186 In fact, Al doping
is employed in commercial LIB cathodes and necessary to
reach state‐of‐the‐art performance benchmarks.94 Similarly, Cu doping can improve cycle stability at the
expense of initial discharge capacity, although this has
not been explored to the extent of Al.160,187,188 Due to the
higher solubility of Cu2+/Cu(OH)2 (Figure 6), it may be
the most prevalent impurity/dopant metal. To study the
effects of Fe impurity, Park et al189 synthesized
LiNi1/3Mn1/3Co1/3FexO2 at x = 0.0005, 0.0025, and 0.01,
demonstrating that at high dopant quantity (x = 0.01),
cation mixing among Fe3+/Ni2+/Li+ was observed,
resulting in poor rate performance and decreased
discharge capacity. However, in lower dopant quantities,
Fe could improve rate performance due to an expansion
of the lattice parameters, leading to facile Li+ (de)
intercalation. Fe doping can also improve the cycle
stability of NMC, likely by suppressing active material
dissolution.189,190 Weng et al45 studied Li(Ni1/3Mn1/3
Co1/3)1‐xMgxO2 and noted that Mg doping could improve
cycle stability at the expense of the initial discharge
capacity as Mg is electrochemically inert; however, Mg
quantities up to x = 0.01 had no significant impact on
electrochemical performance. To the best of knowledge,
the combined effects of these dopants have not been
reported, which may be relevant in resynthesized
cathodes.
7.1
|
Coprecipitation resynthesis
The mixed hydroxide coprecipitation method is one of the
most common synthesis approaches for layered transition
metal oxides that generally involves two steps: (a)
precipitation of metals as a hydroxide precursor at high
pH (eg, Ni1/3Mn1/3Co1/3(OH)2 for NMC cathode) and (b)
solid‐state reaction with a stoichiometric equivalent (5%‐
10% excess) of LiOH or Li2CO3.191 The first step is typically
performed under inert atmosphere (N2) to avoid oxidation
and formation of metal oxides (especially for Mn). NH3 is
also added as a chelating agent to ensure that the metal
ions are well dispersed to balance the rates of nucleation
and crystal growth, which helps generate a homogeneous
particle size and cation distribution, leading to improved
electrochemical performance.192 The feasibility of this
process was first demonstrated by Zou et al82 who used
H2SO4 and H2O2 to leach a mixed cathode scrap. Fe3+
impurity in the leachate was removed by precipitation at
pH = 3, and the concentration ratio of Ni:Mn:Co was
adjusted to unity by adding appropriate quantities of metal
sulfate reagents. The pH was subsequently adjusted to 11
to form and collect the Ni1/3Mn1/3Co1/3(OH)2 coprecipitate, and the remaining in Li+ in the leachate was
precipitated as Li2CO3. Finally, Ni1/3Mn1/3Co1/3(OH)2 and
Li2CO3 were cogrinded and subsequently sintered at
900°C for 15 hours, generating phase‐pure NMC with
good electrochemical performance. As seen in Table 5, this
general approach is versatile and can be used to synthesize
a variety of mixed cathode compositions simply by
adjusting the leachate to the desired metal ratio.
To address stability issues associated with the oxidation of Mn(OH)2 precipitate toward MnOOH and/or
MnO2 which affects the uniformity and reproducibility of
the particle size distribution and electrochemical performance, He et al194 regenerated NMC using the carbonate
coprecipitation method. Here, the cathode was leached
with H2SO4 and H2O2, and after adjusting the Ni:Mn:Co
ratio, NH3 and Na2CO3 were added to the leachate. By
maintaining the leachate at pH = 7.5 and 60°C for
12 hours, the metals were precipitated as Ni1/3
Mn1/3Co1/3CO3, which is more stable than the hydroxide
precursor in aqueous solution.194,203-205 The precursor
morphology generated was spherical with a uniform size
distribution, which was well maintained throughout
subsequent calcination steps and regeneration of NMC.
Oxalate has also been demonstrated as an effective
coprecipitation agent to achieve a similar effect.125,206
Along these lines, Zhang et al127 demonstrated a unique
proof‐of‐concept approach where pure NMC powder was
leached with oxalic acid. As mentioned previously, oxalic
acid generates MC2O4 (M = Ni, Co, and Mn) precipitate,
and thus Li+ was selectively leached, while the powder
suspension comprised MC2O4 and unreacted NMC. NMC
was then regenerated through a solid‐state reaction with
the powder and a stoichiometric equivalent of Li2CO3.
This approach directly regenerates NMC cathodes by
H2SO4 + H2O2
H2SO4 + H2O2
H2SO4 + H2O2
H2SO4 + H2O2
H2SO4 + H2O2
NMC
NMC
NMC
NMC
Leaching conditions
NMC
Coprecipitation synthesis
Cathode resynthesized
Gratz et al159
Chen et al193
He et al194
Zheng et al85
• Unspecified
• Voltage range unspecified
• 158, 155, 149, 140, 133, 125, 133,
79 mAh/g at 0.1, 0.2, 0.5, 1, 2, 3, 5,
10 C, respectively
• ~100% capacity retention after
100 cycles at 0.5 C
• 2.7‐4.3 V vs Li/Li+
• 163.5, 135.1, and 112.6 mAh/g at
0.1, 1, and 5 C, respectively
• 94% capacity retention after 50
cycles at 1 C
• 2.7‐4.3 V vs Li/Li+
• 150.6, 148.8, 141.6, 132.9, and
120.5 mAh/g at 0.1, 0.2, 0.5, 1,
and 2 C, respectively
• 97% capacity retention after 100
cycles at 0.2 C
• Performance similar to NMC
synthesized with fresh reagents
1. Remove Fe3+, Al3+, and Cu2+ at pH = 6.47
2. Adjust Ni:Mn:Co molar ratio to unity
3. Hydroxide coprecipitation at pH = 11
4. Grind dry powder with Li2CO3, then sinter at
900°C for 15 h
1. Remove Cu2+, Fe3+, and Al3+ by pH adjustment
2. Adjust Ni:Mn:Co ratio to unity
3. Coprecipitate by adding NaOH and NH3, followed
by drying
4. Grind with Li2CO3 (5% excess, precalcine at
450°C for 5 h, then sinter at 900°C for 14 h
1. Adjust Ni:Mn:Co ratio in leachate with metal
sulfate precursors
2. Add NH3 and Na2CO3 (1:1 with leachate
concentration), maintain pH at 7.5 for 12 h
3. Form NMC carbonate precursor at 60°C
4. Calcine precursor at 500°C for 5 h, then grind
with Li2CO3 (6% excess)
5. Precalcine at 500°C for 5 h, then sinter at 900°C
for 12 h
1. Remove Cu2+, Fe3+, and Al3+ by precipitation
and solvent extraction
2. Adjust Ni:Mn:Co ratio to unity
3. Corecipitate at pH 10.2 using NaOH and NH3
(NaOH/NH3 = 5 [vol/vol])
4. Precipitate remaining Li+ as Li2CO3
5. Grind coprecipitate Li2CO3 (5% excess),
precalcine at 350°C for 4 h under Ar, then sinter
at 750°C for 10 h
|
(Continues)
Zou et al82
Reference
• 2.5‐4.6 V vs Li/Li+
• 130.2 mAh/g at 46.6 mA/g
• 82.4% after 50 cycles at 46.6 mA/g
Electrochemical performance
1. Remove Fe3+ at pH > 3
(a) Adjust Ni:Mn:Co molar ratio to unity with
metal sulfate reagents
(b) Coprecipitate as metal hydroxides at pH > 11
(c) Remaining Li+ precipitated as Li2CO3 at
40°C
2. Cathode resynthesized using conventional solid‐
state reaction with Li2CO3, sintering at 900°C for
15 h
Resynthesis protocol
T A B L E 5 Summary of approaches to resynthesize cathode from a complex leachate by the coprecipitation or sol‐gel method
26
OR ET AL.
H2SO4
H2SO4 + H2O2
LiNi0.5Mn0.3Co0.2O2
NMC, LiNi0.5Mn0.3Co0.2O2, and
LiNi0.6Mn0.2Co0.2O2
H2SO4 + H2O2
H2SO4 + H2O2
NMC
NMC, LiNi0.5Mn0.3Co0.2O2, and
LiNi0.8Mn0.1Co0.1O2
H2SO4 + H2O2
Leaching conditions
0.2Li2MnO3 ∙ 0.8LiNi1/3Mn1/3Co1/3O2
(Mn‐rich NMC)
Cathode resynthesized
T A B L E 5 (Continued)
3+
2+
1. Adjust to desired to Ni:Mn:Co molar ratio
2. Coprecipitate with NaOH and NH3 at pH 10.5
3. Grind dry powder with Li2CO3 (5% excess),
precalcine at 500°C for 5 h, then sinter at 900°C
(1:1:1, 20 h), 850°C (5:3:2, 15 h), or 750°C
(8:1:1, 20 h, under O2)
1. Adjust to desired to Ni:Mn:Co molar ratio
2. Coprecipitate with NaOH and NH3
3. Grind dry powder with Li2CO3 (3% excess),
precalcine at 500°C for 5 h, then sinter at 900°C
(1:1:1), 850°C (5:3:2), or 800°C (6:2:2) for 12 h
1. Adjust Ni:Mn:Co molar ratio to 5:3:2
2. Coprecipitate with NaOH and NH3 at pH 10.7‐
10.8
3. Grind dry powder with Li2CO3 (5% excess),
precalcine at 500°C for 5 h, then sinter at 850°C
for 15 h
1. Remove impurity ions as described above
2. Ni2+, Mn2+, and Co2+ solvent extracted using 60%
saponified D2EHPA (pH = 3.5, O/A = 1, 6 min,
and three stages)
3. Precipitate Li+ remaining in raffinate as Li2CO3
4. Strip Ni2+, Mn2+, and Co2+ with H2SO4
5. Coprecipitate by adding NaOH and NH3
(pH = 10.5)
6. Grind dry powder with Li2CO3 (5% excess),
precalcine at 450°C for 5 h, then sinter at 900°C
for 20 h
1. Precipitate Al , Fe , and Cu at pH = 4.8, then
purify further by solvent extraction (10%
D2EHPA, O/A = 0.5, and pH = 2)
2. Adjust Ni:Mn:Co molar ratio to 4:7:4
3. Coprecipitate by adding NaOH and NH3
(pH = 10.5)
4. Grind dry powder with LiOH (5% excess),
precalcine at 500°C for 5 h, then sinter at 850°C
for 20 h
3+
Resynthesis protocol
Yang et al195
Yang et al196
Liu et al197
Sa et al198
Yang et al199
• 2‐4.6 V vs Li/Li
• 248.3, 196.4, and 167 mAh/g at
0.1, 0.5, and 1 C, respectively
• 88, 85, and 80% capacity retention
after 50 cycles at 0.1, 0.5, and 1 C,
respectively
• 2.7‐4.3 V vs Li/Li+
• 150, 145, 130, and 100 mAh/g at
0.5, 1, 2, and 5 C, respectively
• 94, 92.8, and 88% capacity
retention after 100 cycles at 0.5, 1,
and 2 C, respectively
• 2.5‐4.3 V vs Li/Li+
• 174.1, 167.4, 161.2, and
155.2 mAh/g at 0.2, 0.5, 1, and
2 C, respectively
• 93.8% capacity retention after 50
cycles at 0.2 C
• 2.7‐4.3 V vs Li/Li+
• >155 mAh/g at 0.1 C
• >80% capacity retention after 100
cycles at 0.5 C
• 2.7‐4.3 V vs Li/Li+
• >168.3 mAh/g at 0.1 C
• >86% capacity retention after 50
cycles at 1 C
• Performance similar to that
synthesized with fresh reagents
(Continues)
Reference
+
Electrochemical performance
OR ET AL.
| 27
As described above
1. Adjust molar ratio by adding metal nitrate
precursors (2% Li excess)
2. Adjust pH to 8.0 using NH3(aq)
3. Stir at 80°C to obtain gel precursor,
followed by precalcination at 400°C and
sintering at 850°C for 8 h
Acetic acid + H2O2
Maleic acid + H2O2
D,L‐Malic acid + H2O2
Citric acid + H2O2
NMC
NMC
NMC
NMC
1. Adjust molar ratio by adding metal nitrate
precursors (2% Li excess)
2. Adjust pH to 8.0 using NH3(aq)
3. Stir at 80°C to obtain gel precursor,
followed by precalcination at 350°C for 2 h
and sintering at 750°C for 12 h
As described above
As described above
Citric acid + H2O2
NMC
1. Adjust molar ratio by adding metal acetate
reagents (5% Li excess)
2. Adjust pH to 7.0 using NH3(aq)
3. Stir at 80°C to obtain gel precursor,
followed by precalcination at 450°C for
4‐5 h, fine grinding, and sintering at 900°C
for 12 h
Resynthesis protocol
D,L‐Lactic acid + H2O2
Leaching conditions
NMC
Sol‐gel synthesis
Cathode resynthesized
T A B L E 5 (Continued)
Li et al94
Li et al99
Li et al99
Yao et al97
Yao et al200
• 2.8‐4.3 V vs Li/Li+
• 149.8 mAh/g at 0.2 C
• 93.9% capacity retention after 160 cycles
at 0.2 C
• 2.8‐4.3 V vs Li/Li+
• 115, 109.8, 97.7, and 87.7 mAh/g at 0.2,
1, 2, and 5 C, respectively
• 85% capacity retention after 150 cycles
at 0.2 C, 83% after 100 cycles at 1 C
• 2.8‐4.3 V vs Li/Li+
• 151.6, 148.4, 133.6, and 120.2 mAh/g at
0.2, 1, 2, and 5 C, respectively
• 84% capacity retention after 150 cycles
at 0.2 C, 87% after 100 cycles at 1 C
• 2.75‐4.25 V vs Li/Li+
• 147.2 mAh/g at 0.5 C
• 94.4% capacity retention after 100 cycles
at 0.5 C
• 2.75‐4.25 V vs Li/Li+
• 154.2, 147, 140, 136.7, and 108.6 mAh/g
at 0.2, 1, 2, 3, and 5 C, respectively
• 93% capacity retention after 50 cycles at
1C
(Continues)
Li et al100
Reference
• 2.8‐4.3 V vs Li/Li+
• 151.6, 145.7, 138.3, 129.7, and
120.6 mAh/g at 0.2, 0.5, 1, 2, and 5 C,
respectively
• 95.2% capacity retention after 70 cycles
at 0.2 C, 96% after 100 cycles at 0.5 C
Electrochemical performance
28
|
OR ET AL.
Oxalic acid
Acetic acid + H2O2
Separation of electrode
from Al foil using
trifluoroacetic acid
NMC
NMC
NMC
Abbreviation: PTFE, polyvinylidene fluoride.
Ascorbic acid
Leaching conditions
Li1.2Ni0.13Mn0.54Co0.13O2
Other approaches
Cathode resynthesized
T A B L E 5 (Continued)
1. Determine concentration of metals and
collect cathode powder
2. Remove residual PTFE from dry powder by
calcination
3. Remove Al impurities by NaOH dissolution
4. Collect powder and adjust molar ratio by
adding metal nitrate precursors to powder
and grind with Li2CO3 (10% excess)
5. Precalcine at 450°C for 5 h then sinter at
900°C for 20 h
1. Adjust Li:Ni:Mn:Co molar ratio to 3.2:1:1:1
with metal acetate reagents
2. Generate NMC precursor by spray drying
pyrolysis at 600°C
3. Calcine precursor at 800°C for 6 h
1. Selectively leach NMC powder with oxalic
acid (0.6 M, 50 mL/g, and 70°C for 10 min)
2. Grind dry powder with Li2CO3 and sinter at
900°C for 14 h
1. Adjust molar ratio by adding metal
(including Li) acetate precursors
2. Metals precipitated as oxalates using oxalic
acid
3. Hydrothermal reaction at 200°C for 8 h
4. Precalcination of dried powder at 450°C for
5 h, followed by sintering at 900°C for 12 h
Resynthesis protocol
Li et al201
Zhang et al127
Zheng et al202
Zhang et al101
• 2.8‐4.3 V vs Li/Li+
• 168 mAh/g at 0.2 C
• 91.5% capacity retention after 100 cycles
at 0.2 C
• 2.6‐4.3 V vs Li/Li+
• 157.1, 150.4, 140.5, 122.1, and
88.6 mAh/g at 0.2, 0.5, 1, 2, and 5 C,
respectively
• 95% capacity retention after 100 cycles
at 0.2 C
• 2.8‐4.5 V vs Li/Li+
• 155.4 mAh/g at 0.1 C
• 83% capacity retention after 30 cycles at
0.1 C
Reference
• 2‐4.8 V vs Li/Li+
• 237.8 mAh/g at 0.5 C
• 77.1% capacity retention after 50 cycles
at 0.5 C
Electrochemical performance
OR ET AL.
| 29
30
|
OR ET AL.
re‐lithiation, and thus is not appropriate to handle
comingled LIB scrap comprising different chemistries.
7.2
|
Sol‐gel resynthesis
In the sol‐gel synthesis method for intercalation electrodes, metal ions are homogeneously dispersed in aqueous
solution with the aid of chelating/gelling agents (often
citric acid).207 Weak bases such as ammonia or acetate
are added to stabilize the pH and enhance metal cation
binding to the chelate. Subsequent water evaporation
forms a dense sol precursor to immobilize the metal ions,
which is then calcined to decompose organics and induce
crystallinity. The sol‐gel resynthesis approach is interesting when predated by organic acid leaching, as the
leachant can simultaneously serve as the chelating agent.
Li et al demonstrated this “grave to cradle” approach by
first optimizing the leaching of LIB cathodes with citric
acid and H2O2 to achieve efficiencies more than 95% for
Li, Ni, Co, and Mn.94 The metal and citric acid
concentrations were then adjusted by adding the appropriate reagents and the gel precursor was generated by
evaporating water. The gel was then precalcined at 450°C
to liberate organics, and subsequently sintered at 900°C
to regenerate NMC. The regenerated NMC exhibited a
slightly lower initial discharge capacity compared to
NMC synthesized identically using lab‐grade reagents,
although the cycle stability and rate performance were
significantly improved. This was ascribed to the regenerated NMC possessing Al dopant due to the likely
presence of Al3+ impurity in the leachate. Using this
approach, the research group also regenerated NMC
using lactic acid as the leachant and chelating agent,
resulting in similar electrochemical performance. The
good electrochemical performance suggests that lactic
acid serves as a potent chelating agent, which is
necessary to generate uniform nucleation sites and small
crystallite sizes during calcination.100 In subsequent
work, a similar resynthesis protocol was used to compare
the performance of acetic and maleic acid.99 Although
both acids were optimized to reach high leaching
efficiencies for all metals (>97%), the electrochemical
performance of regenerated NMC from maleic acid was
significantly improved compared to acetic acid with
respect to initial discharge capacity, cycle stability, and
rate capability. The performance difference was ascribed
to the stronger chelating ability of maleic acid compared
to acetic acid as illustrated in Figure 5. As a result, the
metal cations were more uniformly dispersed with maleic
acid chelation, and following calcination, the XRD
pattern was phase pure, whereas acetic acid chelation
resulted in spinel‐structured impurities.
8 | C O N C L US I O N S AN D
PERS PECT IVES
As the world trends toward mobile electrification, LIBs
are the energy storage technology of choice over the next
half‐century. As such, ramping LIB production demands
will strain resources for precious metals and cause
environmental concerns from the waste generated, all
of which can be addressed through LIB recycling. While
the recycling of Pb‐acid and Ni‐based batteries is a
mature process, it benefits from having a consistent and
relatively simple chemical composition. On the other
hand, LIB technology emerged recently (1991) and
currently comprises various cathode chemistries, adding
complexity to the process. Current commercial recycling
facilities are private ventures largely focused on recovering precious and high‐value Co. However, driven by
technological developments and the rising popularity
EVs, the LIB cathode chemistry market share is shifting
rapidly toward cobalt‐deficient and mixed‐metal compositions. Recycling facilities must adapt to handle mixed‐
type cathodes and comingled LIB scrap comprising
diverse chemistries. Rather than focusing on the metal
commodity value, recycling incentives must consider the
energy savings and environmental benefits from reducing
landfilling, toxic emissions, and reliance on raw material
extraction. Taken together, the widespread realization of
LIB recycling will require legislation and political
pressure, likely in the form of economic incentives (eg,
refundable deposits with LIB purchases), public education, landfill disposal regulations, and defined responsibilities on the collection and disposal of LIBs for
consumers, retailers, and EV and battery manufacturers.
Figure 8 summarizes the various routes explored in
the literature to recycle mixed cathode LIBs. As reviewed
in this work, hydrometallurgical processes can meet the
ideal recycling criteria and potentially recover Li, Ni, Mn,
and Co at high efficiency and purity. However, the
pretreatment of EV battery packs is a major bottleneck
for productivity, as the degree of automation in
disassembly lines is currently limited. Automation
infrastructure requires pack design standardization
among manufacturers, particularly in reducing the
quantity and obscurity of connections to liberate the
LIB cells from other electronic components. Subsequently, the separation of LIB cell components by
crushing and sieving is rapid and automated, where Cu
can be collected in the coarse fractions while active
materials are concentrated in the fine fractions. However,
this generates losses in yield and introduces impurity
metals (Fe, Al, and Cu) in the fine fractions. It is critical
to minimize impurities at this stage to simplify the
leachate purification process. In addition, the treatment
OR ET AL.
| 31
FIGURE 8
Overview of
hydrometallurgical LIB recycling
approaches for diverse cathode
chemistries. LIB, lithium‐ion
battery [Color figure can be viewed at
wileyonlinelibrary.com]
of toxic emissions and recovery of electrolyte and
graphite from crushed LIBs has not been addressed to a
great extent.
Many studies have presented high reductive acid
leaching efficiencies for all LIB cathode metals besides
Fe. Among the various inorganic and organic acids
tested, H2SO4 is currently the most promising for large‐
scale hydrometallurgical plants due to its low cost,
effectiveness (low liquid/solid ratio required), and
familiarity from ore refining processes. Most studies on
metal recovery from the leachate (ie, solvent extraction
and selective precipitation) are performed in sulfate
media. On the other hand, HCl generates aggressive Cl−
ions that cause constant corrosion concerns for commercial plants, while organic acids are currently expensive
for large‐scale processes. In future developments, organic
acids may be interesting as an environmentally friendly
lixiviant that reduces the need for wastewater and gas
treatment systems.
Although lab‐scale tests have demonstrated the
feasibility of removing impurity ions, followed by
separating and recovering a complex leachate of Li+,
Ni2+, Mn2+, and Co2+ at high purity and efficiency, the
main concern is that many steps are required and
upscaling the process will generate high capital costs.
Facilities that process ore concentrate are focused on
recovering only 1 to 2 metals, which may involve a single
leaching and solvent extraction step.208 Here, the main
equipment required are the leaching reactor, mixer‐
settler units for extraction and stripping, and auxiliary
solvent reservoirs. Nevertheless, the feasibility of LIB
recycling must be assessed through pilot plant tests. The
direct resynthesis of LIB cathodes from the leachate can
decrease the number of steps required. In this case, it is
especially important to control the concentration of
impurities in the leachate; however, a trace amount of
Al is necessary to produce commercial‐grade performance. Various mixed‐cathode compositions have been
regenerated by coprecipitation or sol‐gel synthesis
demonstrating good electrochemical performance—
comparisons between these approaches should be cautioned against due to interlaboratory variability in the
32
|
OR ET AL.
synthesis protocol (eg, sintering temperature), electrode
fabrication method, and impurity concentrations in the
leachate.
ACKNOWLEDGMEN TS
The authors gratefully acknowledge the financial support
from the Natural Sciences and Engineering Research
Council of Canada (NSERC) and the University of
Waterloo. This work was financially supported by the
111 Project (no. D17007). Karthikeyan Kaliyappan
acknowledges the financial support from Henan Normal
University, China for this work. Tyler Or was supported
through the NSERC Canada Graduate Scholarships—
Master’s Program.
CONFLICT OF INTERESTS
The authors declare that there are no conflict of interests.
ORCID
Zhongwei Chen
http://orcid.org/0000-0003-3463-5509
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AUTHOR BIOGRAPHIES
Tyler Or received his Bachelor’s degree
in Integrated Science and Chemical
Biology from McMaster University in
2018. He is currently pursuing a PhD in
Chemical Engineering under the supervision of Professor Zhongwei Chen at
the University of Waterloo. His research is focused
on the development of materials and coatings for
lithium and sodium‐ion batteries.
Storm Gourley received his Bachelor’s
degree in Nanotechnology Engineering
from the University of Waterloo in 2018.
He is currently pursuing a Master’s in
Chemical Engineering under the supervision of Professor Zhongwei Chen at
the University of Waterloo. His research focuses on
the development and commercialization of novel
materials for next‐generation lithium‐ion batteries.
Dr Karthikeyan Kaliyappan received
his Bachelor of Technology (B.Tech,
2005) and Master of Technology
(M.Tech, 2007) in Electrochemical Engineering from Central Electrochemical
Research Institute, India. He earned his
PhD from Chonnam National University, South
Korea in Advanced Chemicals and Engineering
(2013). He is currently working as Research
Scientist at CWZE Power Inc, Waterloo, Canada.
His current fields of interests are designing high
energy density materials for energy storage devices
including lithium‐ion and sodium‐ion batteries,
and next‐generation metal‐ion capacitors. He is
also mastered in developing metal‐oxide coatings
and composites using atomic layer deposition.
Dr Aiping Yu is a Professor at the
University of Waterloo. Her research
interests focus on the development,
processing, and functionalization of nanostructured carbon materials, along
with their application as electrode materials in high performance supercapacitors. She
has published over 150 refereed journal papers,
three book chapters, and one book. These publications have received over 13 000 citations. She holds
seven patents and provisional patents for nanomaterials and device development, and two of them
have been licensed to industry.
Dr Zhongwei Chen is the Canada
Research Chair Professor (Tier 1) in
Advanced Materials for Clean Energy
at the University of Waterloo, a Fellow
of the Canadian Academy of Engineering and Vice President of the International Academy of Electrochemical Energy Science
(IAOEES). His research interests involve the
development of advanced energy materials and
electrodes for fuel cells, metal‐air batteries, and
lithium‐ion batteries. He has published two books,
nine book chapters, and more than 280 peer
reviewed journal articles with over 28 000 citations
with an h‐index of 81. He is also listed as an
inventor with over 20 US/international patents
licensed to companies internationally.
How to cite this article: Or T, Gourley SW,
Kaliyappan K, Yu A, Chen Z. Recycling of mixed
cathode lithium‐ion batteries for electric vehicles:
Current status and future outlook. Carbon Energy.
2020;1–38. https://doi.org/10.1002/cey2.29