Manganese Vanadium Oxide Compounds as Cathodes for Lithium ...

Mat. Res. Soc. Symp. Proc. Vol. 658 © 2001 Materials Research Society
Manganese Vanadium Oxide Compounds As Cathodes For Lithium Batteries
J. Katana Ngala, Peter Y. Zavalij and M. Stanley Whittingham*
Institute for Materials Research and Department of Chemistry
State University of New York at Binghamton,
Binghamton, NY13902-6000, USA.
ABSTRACT
A pure form of the compound [(CH 3 ) 4 N] 0.2 Mn 0.06 V 2 O 5+δ •0.2H 2 O has been synthesized by
the hydrothermal technique. A manganese richer compound, [(CH 3 ) 4 N] 0.18 Mn 0.1 V 2 O 5+δ •0.35H 2 O,
has also been synthesized. By comparison to [(CH 3 ) 4 N] 0.17 Fe 0.1 V 2 O 5+δ •0.17H 2 O, the structure of
these compounds consists of double sheets of vanadium oxide with tetramethyl ammonium ions
in the interlayer spaces. Both compounds display reversible lithium intercalation with sharp end
points. The lithiation capacity is lowered by the increase in manganese content.
INTRODUCTION
Transition metal oxides may act as cathodes in lithium secondary batteries partly due to the
ability of the metal ions to exhibit variable oxidation states, and their highly negative free
energies of reaction with lithium ions, in contrast to other chalcogenides.
The commercially available layered LiCoO 2 has good cyclability. However, it finds limited
use such as in lap tops and cellular phones, due to its high cost. The spinel compound, LiMn 2 O 4 ,
which was found to act as an intercalation host in the early 1980s, has been developed for
commercial use as a replacement for LiCoO 2 in second-generation Li-ion batteries; the
manganese is lower cost than cobalt. Unfortunately, only 0.5 Li per Mn may be inserted and deinserted
reversibly [1]. This corresponds to a capacity of about 110 Ah/kg in contrast to 130
Ah/kg for LiCoO 2 .
Vanadium oxides have been extensively investigated as possible cathodes for lithium
batteries. A class of compounds of particular interest for potential use in batteries is that
consisting of the δ-V 2 O 5 structure. This structure type contains double vanadium oxide sheets
with the vanadyl groups on the outside of the sheets; the vanadium is in a distorted octahedral
coordination rather than the square pyramid typical of V 2 O 5 itself and the well known bronzes,
such as LiV 2 O 5 . One of the most studied vanadium oxide cathodes,V 6 O 13 , contains these sheets
and cycles lithium well [2] in secondary cells. Typical simple members of this class are
Ni 0.25 V 2 O 5 [3], [N(CH 3 ) 4 ] 0.3 Fe 0.1 V 2 O 5 and Zn 0.4 V 2 O 5 [4].
We report here on the preparation of a pure phase of [N(CH 3 ) 4 ] y Mn z V 2 O 5+δ •nH 2 O
(compound A) by hydrothermal synthesis. Earlier we reported [5] this phase as one component
of a two phase mixture. We also report on an isostructural material (compound B) with a higher
manganese content. We compare their electrochemical behavior. Hydrothermal synthesis and
other soft chemistry techniques such as sol-gel synthesis, have been widely applied for the low
temperature syntheses of metal oxides. These techniques are useful in accessing metastable metal
oxides.
GG9.16.1

EXPERIMENTAL
Compound A was hydrothermally synthesized by the treatment of V 2 O 5 (Alfa-Aesar) with
an aqueous solution of HMnO 4 , in the presence of the templating cation [(CH 3 N) 4 ]OH (Alfa-
Aesar). The reacting mole ratio was V 2 O 5 :HMnO 4 :[(CH 3 ) 4 N]OH = 1:0.15:2. Nitric acid was
added to lower the pH. HMnO 4 was obtained by ion exchanging the potassium in KMnO 4
(Fisher) using an acidic cation exchange resin. The mixture was transferred to an autoclave after
thorough stirring, placed in an oven at 165 0 C and left to react for 3 days. The pH of the reaction
mixture was initially 2.4 and 3.5 at the end of the reaction. This pH range along with the given
reacting mole ratios ensured pure product. The brown clay-like solid was filtered and rinsed with
distilled water.
Compound B was prepared by taking the above product (compound A) and adding to it
LiOH and HMnO 4 , in the mole ratio 3:1:1, respectively. The same heat treatment as above was
followed. The reaction proceeded under similar pH conditions (2.2-3.5). A dark brown clay was
isolated.
X-ray powder diffraction was performed using Cu Kα radiation on a Scintag θ-θ
diffractometer equipped with a Ge (Li) solid state detector. The data was collected from 2θ = 2 0
to 2θ = 90 0 with 0.02 0 steps and 10 sec. per step count time. The TGA data was obtained on a
Perkin Elmer model TGA 7, the FTIR on a Bruker Equinox and electron microscopy on a JEOL
8900. A mixture of the sample, carbon black and polyvinylidene fluoride as the binder, was
prepared in the weight ratio 8:1:1, respectively. Cyclopentanone was added to the mixture and
thoroughly mixed to obtain a slurry that was coated on a 1 cm 2 aluminium disk using a dropper.
After air–drying the disk was placed in an oven at about 150 0 C for effective dehydration for at
least 5 hours. Swagelok test cells were used with lithium metal, adhered to a nickel disk, as the
anode. The electrolyte was 1M LiPF 6 dissolved in a 2:1 mixture of ethylene carbonate (EC) and
dimethyl carbonate (DMC).
RESULTS AND DISCUSSION
Compound A:
EDS analysis on compound A indicated the presence of V and Mn in the ratio Mn:V =
0.06:2. This corresponds to z = 0.06 in the formula, [N(CH 3 ) 4 ] y Mn z V 2 O 5+δ·nH 2 O (standard
deviation = ±0.01). The amount of Mn obtained here is less than previously reported, probably
due to extra manganese present in the other phases of the earlier product.
Powder x-ray analysis, shown in Figure 1, indicates a single phase product in contrast to
that reported earlier [5] which had mixed phases. The pattern, which suggests low crystallinity,
was indexed using a monoclinic unit cell in the space group C2/m, with the following
parameters: a = 11.66 Å, b = 3.61 Å, c = 13.91 Å and = 108.8 0 . The low crystallinity is
confirmed by the SEM image, which shows a fuzzy layered morphology.
The powder x-ray pattern gives a repeat distance of 13.1 Å, consistent with δ-V 2 O 5 sheets
having TMA ions sandwiched in the interlayer spacing. The in-plane spacing of 3.6 Å points to
disordered organic cations as the diameter of these cations is >5Å. Unfortunately, Rietveld
GG9.16.2

efinement could not be performed as the quality of the XRD pattern was insufficient. The
structure was deduced from the similarity of its XRD pattern with that of
[(CH 3 ) 4 N] 0.17 Fe 0.1 V 2 O 5+δ·0.17H 2 O whose structure we reported earlier [4]. From this analogy, we
postulate that the structure of compound A consists of VO 6 distorted octahedra forming double
sheets of V 2 O 5 as shown in Figure 2.
300
Relative Intensity
250
200
150
100
50
0 0 1
0 0 2
2 0 -1
0 0 3
1 1 0
0 0 4
1 1 2
4 0 1
0 0 6
6 0 -3
5 1 1
6 0 3
0 0 10
0
10 20 30 40 50 60 70 80 90
2θ
Figure 1: Powder x-ray pattern for [N(CH 3 ) 4 ] y Mn z V 2 O 5+ •nH 2 O;z = 0.06.
Figure 2: The structure for [N(CH 3 ) 4 ] y Mn z V 2 O 5+ •nH 2 O.
The FTIR spectrum (Figure 3) indicates the presence of a V=O group by the peak at 1012
cm -1 . The vanadyl groups are projected away from the double layers. The band at 3421 cm -1 is
due to the presence of water, whereas those at 517 cm -1 and 754 cm -1 are attributed to V-O
vibrations in the framework. The presence of the tetramethylammonium ion is indicated by the
three bands situated at 2900 cm -1 , 1481 cm -1 and 950 cm -1 .
GG9.16.3

0.5
Transmittance
0.45
0.4
0.35
0.3
0.25
4000 3500 3000 2500 2000 1500 1000 500
Wavenumber, cm -1
Figure 3: FTIR spectrum for [N(CH 3 ) 4 ] y Mn z V 2 O 5+ •nH 2 O; z = 0.06.
Thermal analysis, performed under oxygen atmosphere, results in two weight losses as
displayed in Figure 4. The 1.8% weight loss in the range, 100 0 C-250 0 C is associated with water
of solvation between the sheets. This corresponds to n = 0.20 in the formula:
[(CH 3 ) 4 N] y Mn z V 2 O 5+δ •nH 2 O. The weight loss before 100 0 C is associated with surface water. The
7.6% weight loss over the range 250 0 C-300 0 C is due to decomposition of the organic,
corresponding to y = 0.21 in the formula: [(CH 3 ) 4 N] y Mn z V 2 O 5+δ •nH 2 O.
100
95
Under O
2
Weight %
90
85
80
100 200 300 400 500 600
Temperature, °C
Figure 4: TGA profile for [N(CH 3 ) 4 ] y Mn z V 2 O 5+ • •nH 2 O; z = 0.06.
The product obtained by subjecting the sample to a constant temperature of 200 0 C for 4
hours shows an XRD pattern that is identical to that of the original sample. This implies that
water is not crucial for the integrity of the structure. However, upon heating to 600 0 C, the
structure collapses as evidenced by the XRD pattern of the resulting product. The pattern is
consistent with a mixture of MnV 2 O 6 and V 2 O 5 .
The electrochemical data shows an open circuit potential of 3.62 V, consistent with
vanadium in the +5 oxidation state. Figure 5 displays the charge–discharge profile of 7 cycles. It
indicates initial capacity of 220 Ah/kg, corresponding to about 2 Li per formula unit. The same
capacity was reported for the earlier product [5]. We still observe the plateau at about 2.8 V,
suggesting some phase change during cycling.
GG9.16.4

5
4
Volts
3
2
1
Initial discharge
0
-50 0 50 100 150 200 250
Capacity, Ah/kg
Figure 5: Cycling of [N(CH 3 ) 4 ] y Mn z V 2 O 5+ ; z = 0.06 at 0.1 ma/cm 2 .
Compound B
EDS analysis yields the approximate ratio Mn:V = 0.1:2, corresponding to z = 0.1(1) in the
formula [(CH 3 ) 4 N] y Mn z V 2 O 5+δ .nH 2 O. On the other hand, the powder XRD pattern of this product
is identical to that of compound A, so that that the additional Mn content does not alter the δ-
phase structure.
Figure 6 gives the thermal analysis profile for this compound. The 3.0% weight loss over
the temperature 82 0 C-200 0 C corresponds to n = 0.35 in the formula [(CH 3 ) 4 N] y Mn z V 2 O 5+δ .nH 2 O,
whereas the 6.3% loss in the interval, 200 0 C-350 0 C corresponds to y = 0.18. Thus, concomitant
with increase in Mn content in the structure is a decrease in the amount of the organic cation and
an increase in water of solvation. This suggests that probably the manganese ions, as the organic
cations, are situated in the interlayer spaces in hydrated form. Thus the introduction of more
manganese ions causes the displacement of the organic cations.
100
Under O
2
Weight percent
95
90
85
80
100 200 300 400 500 600
Temperature, °C
Figure 6: Thermogravimetric profile for [N(CH 3 ) 4 ] y Mn z V 2 O 5+ .•nH 2 O; z = 0.1.
GG9.16.5

The electrochemical behavior of this compound I shown in Figure 7. It exhibits an open
circuit potential of 3.59 V, which is slightly less than that for compound A, due to the higher
manganese content which will probably reduce the overall oxidation state of the oxide matrix.
The cell was charged initially; essentially no capacity was observed during this charging process
indicating no lithium in the sample, even though the reaction medium contained much lithium.
5
4
Volts
3
2
1
Initial charge
0
-50 0 50 100 150 200 250
Capacity, Ah/kg
Figure 7: Cycling of [N(CH 3 ) 4 ] y Mn z V 2 O 5+ ; z = 0.1 at 0.1ma/cm 2 .
Figure 7 shows that the capacity of the cell, 180 Ah/kg, is significantly reduced from the
220 Ah/kg of sample A. The reason for this is not fully understood, as the 0.04 additional Mn
ions cannot account for the drop of around 0.4 Li per formula unit simply based on oxidation
state changes or site occupancy. It is possible that this loss is in part a kinetic effect, with the
additional manganese ions pinning the layers together much as just a few per cent additional
titanium ions in Li x Ti 1+y S 2 can cut the lithium diffusion coefficient by several orders of
magnitude [6].
CONCLUSIONS
The pure form of the δ-phase compound [(CH 3 ) 4 N] 0.2 Mn 0.06 V 2 O 5+δ •0.2H 2 O was
hydrothermally synthesized. A more manganese rich phase, [(CH 3 ) 4 N] 0.18 Mn 0.1 V 2 O 5+δ •0.35H 2 O
was synthesized from the former compound. Both compounds display reversible cycling as
cathodes, although the latter has lower capacity than the former.
We are currently investigating other δ-phase compounds of the series, to which the two
compounds belong, with the aim of evaluating their performance as cathodes for lithium ion
batteries. Work is also in progress to explore other types of pillars for the layers and to study the
manganese oxide analogs.
GG9.16.6

ACKNOWLEDGEMENTS
We thank the Department of Energy, Office of Transportation Technologies, through
Lawrence Berkeley Laboratory, and the National Science Foundation through Grant
DMR–9810198 for partial support of this work. We also thank Bill Blackburn for the SEM
studies.
REFERENCES
1. A. R. Armstrong, A. D. Robertson, R. Gitzendanner, and P. G. Bruce, J. Solid State Chem.,
145, 549-550(1999).
2. Ö. Bergström, H. Björk, T. Gustafsson, J. O. Thomas, J. Power Sources, 81-82(1999) 685.
3. Y. Oka, T. Yao, N. Yamamoto, J. Solid State Chem., 132 (1997) 323.
4. F. Zhang, P. Y. Zavalij and M. S. Whittingham, Mater. Res. Bull., 32 (1997) 701.
5. F. Zhang and M. S. Whittingham, Electrochem. Comm., 2 (2000) 69.
6. M. S. Whittingham, Progress in Solid State Chemistry, 12 (1978) 41.
GG9.16.7