A Template-Free, Ultra-Adsorbing, High Surface Area Carbonate Nanostructure
High Surface Area Carbonate
Nanostructure. PLoS ONE 8(7): e68486. doi:10.1371/journal.pone.0068486
A Template-Free, Ultra-Adsorbing, High Surface Area Carbonate Nanostructure
Johan Forsgren 0 1
Sara Frykstrand 0 1
Kathryn Grandfield 0 1
Albert Mihranyan 0 1
Maria Strmme 0 1
Richard G. Haverkamp, Massey University, New Zealand
0 Current address: Department of Materials Science and Engineering, McMaster University , Hamilton , Canada
1 1 Division for Nanotechnology and Functional Materials, Department of Engineering Sciences, The A ngstro m Laboratory, Uppsala University , Uppsala , Sweden , 2 Division for Applied Materials Science, Department of Engineering Sciences, The A ngstro m Laboratory, Uppsala University , Uppsala , Sweden
We report the template-free, low-temperature synthesis of a stable, amorphous, and anhydrous magnesium carbonate nanostructure with pore sizes below 6 nm and a specific surface area of , 800 m2 g21, substantially surpassing the surface area of all previously described alkali earth metal carbonates. The moisture sorption of the novel nanostructure is featured by a unique set of properties including an adsorption capacity ,50% larger than that of the hygroscopic zeolite-Y at low relative humidities and with the ability to retain more than 75% of the adsorbed water when the humidity is decreased from 95% to 5% at room temperature. These properties can be regenerated by heat treatment at temperatures below 100uC.The structure is foreseen to become useful in applications such as humidity control, as industrial adsorbents and filters, in drug delivery and catalysis.
. These authors contributed equally to this work.
Nanotechnology is starting to influence most scientific areas and
this key enabling technology is foreseen to significantly impact all
materials science dependent industries during the coming decades
. The interest in high surface area nanostructured materials
from 1990 onwards has increased exponentially for all classes of
porous materials and at the beginning of 2013, according to the
ISI Web of Knowledge, there were in total about 60,500 records
on zeolites, 20,500 records for mesoporous silica and 13,100
records on metal organic framework (MOF) materials, whereas
before 1990 these numbers were insignificant. The most common
way to produce high surface area materials with
micro-mesoporous structures, i.e., pores with diameters below 50 nm, is by using
soft templates and building around them a more rigid structure
after which the template is eluted with a solvent or burnt away to
produce the rigid porous material.
In the current work we will show that it is possible, at low
temperatures and without the use of templates, to synthesize a
unique high surface area nanostructure with a well-defined
poresize distribution of sub 6 nm pores of a widely used, non-toxic and
GRAS (generally-recognised-as-safe)-listed material that is already
included in the FDA Inactive Ingredients Database ; viz.
Magnesium is the eighth most abundant element in the earths
crust and essential to most living species. It can form several
structures of hydrated carbonates such as nesquehonite
(MgCO3?3H2O), and lansfordite (MgCO3?5H2O), a number of
basic carbonates such as hydromagnesite (4
MgCO3?Mg(OH)2?4 H2O), and dypingite (4 MgCO3?Mg(OH)2?5 H2O), as
well as the anhydrous and rarely encountered magnesite (MgCO3)
. In contrast to other alkali earth metal carbonates, chemists
have found anhydrous magnesium carbonate difficult to produce,
particularly at low temperatures. Above 100uC, magnesite
(crystalline MgCO3) can be obtained from Mg(HCO3)2 solutions
by precipitation. However, at lower temperatures, hydrated
magnesium carbonates tend to form, giving rise to what has been
referred to as the magnesite problem .Yet, not only chemists have
been intrigued by magnesium carbonates. Although abundant in
nature, where crystalline forms exist as traces in most geological
structures, pure magnesium carbonate is seldom found on its own
in larger deposits, a fact that has puzzled geologist for more than a
In 1908, Neuberg and Rewald tried to synthesise magnesite in
alcohol suspensions of MgO . However, it was concluded that
MgCO3 cannot be obtained by passing CO2 gas through such
suspensions due to the more likely formation of magnesium
dimethyl carbonate (Mg(OCOOCH3)2). Subsequent studies by
Kurov in 1961  and Buzagh in 1926  only reiterated the
assumption that MgO preferentially forms complex dimethyl
carbonates when reacted with CO2 in methanol. A further
overview of early works is provided in detail in Text S1 and Figure
Yet, by changing the synthesis conditions in comparison to what
has been described earlier, we here report the successful formation
of a magnesium carbonate, hereafter referred to as Upsalite, in a
reaction between MgO, methanol and CO2 resulting in an
anhydrous, micro-mesoporous and large surface area structure.
We further show that the moisture sorption of the material is
featured by a unique set of properties including an adsorption
capacity ,50% larger than that of the hygroscopic zeolite-Y at low
relative humidities and with the ability to retain more than 75% of
the adsorbed water when the humidity is decreased from 95% to
5% at room temperature. The humid material is easily
regenerated to regain its moisture sorption characteristic upon storage at
Results and Discussion
The synthesis is carried out well below 100uC, while previously
reported amorphous structures of magnesium carbonate have
been formed at higher temperatures by thermal decomposition of
hydrated magnesium carbonates  or of a double salt of
magnesium ammonium carbonate . In the current work, CO2
is not bubbled through the methanolic suspension, instead the
reaction vessel is pressurised with CO2 to moderate relative
pressures (13 bar). Initially, the temperature is kept at 50uC in
order to facilitate a reaction between MgO and methanol, and
after , 3 h the temperature is decreased to room temperature.
This results in formation of a rigid gel in the reaction vessel after ,
4 days. When dried in air at 70uC, the gel solidifies and collapses
into a white and coarse powder that is primarily X-ray amorphous
with traces of unreacted and crystalline MgO, see XRD pattern in
Figure 1a. The sharp peaks at 2 h equal to 43u and 62u originate
from the unreacted MgO , while the halo peak between 2 h
values of 25u and 40u is indicative of at least one amorphous phase.
Raman spectroscopy reveals that the powder indeed is
composed of a carbonate (Figure 1b), where the band at
,1100 cm21 corresponds to vibration of the carbonate group
. Moreover, a broad halo, or the so-called Boson peak, with a
maximum at ,100 cm21, is further witness to the amorphous
character of the powder .
When examined with Fourier transform infrared spectroscopy
(FTIR, Figure 1c) the material displayed absorption bands at
,1440 cm21, ,1100 cm21 and ,850 cm21, which all
correspond to the carbonate group . No water of crystallisation is
visible in this spectrum [13,19]. The anhydrous character of the
bulk material is further confirmed by Thermal Gravimetric
Analysis (TGA) (see Figure S2).
X-ray photoelectron spectroscopy (XPS) confirms the
anhydrous nature of the MgCO3. Energy resolved spectra were
recorded for the Mg2p and O1s peaks (Figure 1d,e) which were
found to be positioned at 52.1 eV and 533.5 eV, respectively,
which is indicative of MgCO3 . Further, the O1s peak does not
contain any components for crystal water which, expectedly,
would have appeared at 533533.5 eV . The shoulder seen at
535.6 eV is located between the binding energies for liquid water
(539 eV) and ice (533 eV)  and is, therefore, representative of
surface adsorbed water as previously described for adsorbed water
on carbon fibres [23,24]. The shoulder seen at 531.0 eV shows the
presence of MgO in the powder. No presence of Mg(OH)2 was
observed in the bulk, which would have resulted in a peak at
532.4 eV .
Having proved the formation of amorphous anhydrous MgCO3
by XRD, Raman, FTIR, and XPS, we postulate the following
simplified route of synthesis, based on the presence of
HOMgOCH3 as an intermediate (as confirmed with FTIR in Figure S3)
and the necessity of heat treatment in the last synthesis step:
The presence of a hydroxyl group in the vicinity of the methoxy
group in HOMgOCOOCH3 makes the compound labile and will
therefore favour internal transition to a solvate MgCO3?CH3OH
and loss of methanol upon heating to 70uC:
?MgCO3 : CH3OH
As it is evident from the above cascade scheme, the postulated
reaction of MgCO3 formation goes through several steps, some of
which are equilibrium reactions, namely 1a and 1b. A schematic
description of the reaction steps is presented in Figure 2.
In order to analyse the pore structure and water sorption
capacity of Upsalite, N2 and H2O vapour sorption analyses were
performed. Figure 3a shows the N2 sorption isotherm for Upsalite,
which exhibits a typical Type 1 shape according to the IUPAC
classification . The SSA of 800 m2 g21 for the material
(Table 1) was derived from such isotherms according to the
Brunauer-Emmet-Teller (BET) equation , substantially
surpassing the SSA of all previously described alkali earth metal
carbonates, where crystalline forms of magnesium carbonates
typically have a SSA of 418 m2 g21 . This high SSA places
Upsalite in the exclusive class of high surface area nanomaterials
including mesoporous silica, zeolites, MOFs, and carbon
Figure 3c displays the H2O vapour sorption isotherm for
Upsalite and, based on the large amount of H2O adsorbed at low
RHs, it is evident that the material is highly hydrophilic . The
limited desorption of moisture from the material when the vapour
pressure is reduced from 95% is further proof of the strong
interaction between water molecules and the material. It should,
however, be noted that no signs of hydrate formation in the
material are seen using XRD after the sorption isotherm is
completed, and that the sorption isotherm can be repeated with
undistinguishable results after heat treatment at moderate
temperatures (95uC) under vacuum. This contrasts to the
regeneration of moisture sorption properties of, e.g., Zeolites
typically requiring heat treatments at temperatures between 150uC
Further, both the N2 and the H2O vapour sorption isotherms
were analysed in order to establish the microporous properties of
the material according to the Dubinin-Astakhov (D-A) model ,
see Table 1. The hydrophilic nature of the material is further
reflected in the greater characteristic energy for adsorption of H2O
compared to N2. The discrepancy in the limiting micropore
volume (w0) in which the value obtained from the N2 sorption
isotherm is to be regarded as the true value and modal
FRTaImRasnpescpteructmru.mT h.eThtehrbeaenvdisoibblesearvbesdoraptti,on11b0a0ndcsm(21144s0tecmms2fr1o,m110vi0brcamti2on1 aonfdth8e5C0Oc3mg2r1o)uapreanaldl dthuee htoalovibcaranttieornesdoaftt1h0e0CcOm3 2g1roisupa.Bdo)soanndpeea)kX.PcS)
Mg2p and O1s peaks. The Mg2p peak at 52.1 eV and the O1s peak at 533.5 eV stem from MgCO3, the O1s peak at 531.0 eV from MgO and the O1s peak
at 535.6 eV from surface adsorbed water. The solid lines represent the measured spectrum. The coloured lines are calculated using the CasaXPS
software and represent the fitted curves (obtained using Gaussian-Lorentzian functions) and the subtracted background (obtained using a Shirley
equivalent pore size obtained from the two sorption isotherms is
most likely due to site-specific interaction between the H2O species
and the material, not only in the micropores but also on the
exterior of the material and in pores larger than 2 nm .
Figure 3b shows the incremental and cumulative pore volume
obtained through density functional theory (DFT) calculations on
the N2 sorption isotherm. From these calculations it infers that
about 98% of the pore volume is inherent to pores with a diameter
smaller than 6 nm, while the remaining pore volume is made up of
pores with a broad size distribution between 8 and 80 nm centred
at ,16 nm.
When examined with scanning electron microscopy (SEM),
these pores in the larger size range are clearly visible (Figure 4ab).
Furthermore, the highly porous nature of the material is evident
from the scanning transmission electron microscopy (STEM)
tomography work available in the Supporting Information (see
Video S1 and S2 and Figure S4; S4.1). Such three-dimensional
reconstructions allow for visualisation of the internal pore structure
of the material via a series of slices through the volume.
Measurements from these slices, which essentially represent
cross-sections, confirm that pore widths are consistently 16 nm,
while pore heights vary between 8 and 50 nm. However, these
larger pore networks that are visible with SEM and STEM are not
Total pore volume[b] (cm3/g)
w0, limiting micropore volume[c] (cm3/g)
Equivalent surface area in micropores[c] (m2/g)
Characteristic energy of adsorption[c] (kJ/mol)
Modal equivalent pore width[c] (nm)
Correlation coefficient of fit[c]
[a]According to the BET equation
[b]Single point adsorption at P/P0<1
[c]According to the Dubinin-Astakhov equation
noted throughout the entire material, which is consistent with the
limited contribution from these pores to the total pore volume as
determined by DFT. In the transmission electron microscopy
(TEM) image in Figure 4c the smaller pores, dominating the
contribution to the total pore volume sensed by nitrogen sorption,
can be distinguished. An enlargement of pores was found to take
place under the electron beam where the sample was unstable for
long periods of time. This enlargement is most likely due to
remaining organic groups leaving the sample. A representative
image recorded after a longer period (, 1 min) under the electron
beam is shown in Figure S4 (S4.2).
The water sorption capacity of the material is interesting from
an industrial and technological point of view and it is, hence,
compared to three commercially available desiccants, namely
fumed silica (SSA: 196 m2 g21), hydromagnesite (SSA: 38 m2 g21)
and the microporous Zeolite Y (SSA: 600 m2 g21, silica/alumina
ratio 5.2:1), see Figure 3c. For comparison, all samples were
degassed at 95uC under vacuum for 10 h prior to analysis. The
H2O vapour adsorption isotherm for Upsalite displays similarities
with the hydrophilic zeolite at very low RHs (,1%) and shows an
even higher adsorption capacity compared to the zeolite at RHs
between 1 and 60%. This behaviour contrasts largely to that of the
other two non-porous materials, i.e. fumed silica and
hydromagnesite, which mainly adsorb H2O at RH .60%.
Amorphous magnesium carbonates produced by high
temperature thermal decomposition of hydromagnesite, nesquehonite, or
magnesium ammonium carbonate double salt have previously
been reported to be unstable upon hydration [15,30]. In
particular, instability of the hydromagnesite decomposed material
was evident by a weakening of the carbonate bond . Such
weakening was observed by a shift towards lower temperatures, as
well as by the development of a shoulder and a split into two or
more peaks, of the carbonate decomposition peak located above
350uC in differential TGA (dTGA) spectra . In this respect,
Upsalite appears to remain stable upon hydration. After 11 weeks
of storage at RT and 100% RH no peak split, shoulders or
movement of the carbonate decomposition peak towards lower
temperatures is observed in dTGA spectra (see Figure S5). In fact,
the decomposition peak for the carbonate bond is shifted towards
higher temperatures as compared to the as-synthesized material
(Figure S2 and S5), indicating a strengthening of the carbonate
Figure 3. Sorption isotherms and DFT-based pore size distribution for Upsalite. a) N2 sorption isotherm at 77 K. b) Incremental pore
volume (violet) and cumulative pore volume (blue) obtained from N2 sorption isotherm. c) Moisture sorption isotherm at room temperature for
Upsalite (blue), Mg5(CO3)4(OH)2?4H2O (green), Aerosil (red) and Zeolite Y (black). The arrows indicate the direction of the pressure change.
We report herein the template-free formation of a stable,
amorphous magnesium carbonate nanostructure formed at low
temperature in a methanol solution of MgO with CO2. The
obtained magnesium carbonate is featured with a unique structure
of pores almost exclusively in the sub 6 nm size range and
extraordinarily high surface area, which has never been reported
before, neither for natural nor synthetic magnesium carbonates.
The material described in this work is further featured with
extraordinary water sorption properties that can be regenerated at
temperatures below 100uC. The material is foreseen to find its use
in a number of applications including humidity control and
delivery systems for therapeutic or volatile agents.
Materials and Methods
In the current work 4 g magnesium MgO powder was placed in
a glass bottle together with 60 ml methanol and a stirring magnet.
The solution was put under 3 bar CO2 pressure and heated to
50uC. After approximately 4 hours the mixture was allowed to
cool to RT and the carbon dioxide pressure was lowered to 1 bar,
and the reaction continued until a gel had formed. When a gel was
obtained, the carbon dioxide pressure was removed and the gel
was allowed to solidify and dry at ,70uC during 3 days. A
schematic description of the synthesis is found in Fig. 2.
X-ray diffraction. XRD analysis was performed with a
Siemens/Bruker D5000 instrument using Cu-Ka radiation.
Samples were ground and put on a silicon sample holder with
zero background prior to analysis. The instrument was set to
operate at 45 kV and 40 mA.
Raman spectroscopy. A Reinshaw Ramanscope was used
for the Raman studies. The Raman instrument was calibrated
with a silicon wafer using the band at 521 cm21 prior to the
studies. A 524 nm argon-ion laser with a 10 mm spot size was used
Fourier transform infrared spectroscopy. The FTIR
studies were performed on a Bruker IFS 66v/S spectrometer
using an Attenuated Total Reflectance (ATR) sample holder from
SENSIR. 50 scans were signal-averaged in each spectrum and the
resolution was 4 cm21. Before the measurement a background
scan was recorded and thereafter subtracted from the spectrum for
X-ray photoelectron spectroscopy. The XPS experiments
were conducted on a Phi Quantum 200 Scanning ESCA
microprobe instrument. Prior to analysis, the samples were sputter
cleaned using argon ions for 10 min at 200 V to remove surface
adsorbed contaminations. A full spectrum was recorded together
with energy resolved spectra for Mg2p and O1s. During the
acquisition, an electron beam of 20 mA was used together with
argon ions to neutralise the non-conducting sample. The peak
fittings were made with CasaXPS software, the curves were fitted
using Gaussian-Lorentzian functions and the background was
subtracted using a Shirley function. The spectra were calibrated
against the O1s peak for magnesium oxide (531.0 eV) instead of
the C1s peak at 285.0 eV for adventitious carbon, which otherwise
is commonly used as a reference. However, in the case of MgO,
the binding energy for adventitious carbon is not reliable as
reference since hydrocarbons interact with magnesium oxide in a
way that shifts the C1s peak randomly making it unsuitable as a
reference. Therefore, the O1s peak for MgO (531.0 eV) is instead
proposed to be used as an internal reference . The presence of
magnesium oxide in the samples was confirmed by XRD analysis
prior to the XPS study.
N2 sorption analysis. Gas sorption measurements were
carried out with N2 at 77 K using an ASAP 2020 from
Micromeritics. The samples were degassed at 95uC under vacuum
for 10 h prior to analysis. The SSA was determined by applying
the BET equation  to the relative pressure range of 0.050.30
for the adsorption branch of the isotherm. The D-A equation was
employed on the appropriate pressure region for adsorption in
micropores. The BET and D-A calculations were performed with
the ASAP 2020 V3.04 software from Micromeritics delivered
together with the analysis equipment. The pore size distribution
was determined using DFT analysis carried out with the DFT Plus
software from Micrometrics using the model for N2 at 77K for
slitshape geometry with low regularisation (l = 0.005). The standard
deviation of the DFT fit was 2.037 cm3/g.
Scanning electron microscopy. For the SEM analyses, a
Leo 1550 instrument from Zeiss equipped with an in-lens detector
was used. Prior to the studies, the samples were cooled with liquid
N2, crushed and put on a stub holder with double-sided carbon
tape. As a last step prior to analysis the sample was sputter coated
with a thin layer of gold/palladium.
Thermal gravimetric analysis. TGA analysis was carried
out under a flow of air in an inert alumina cup with sample sizes of
approximately 15 mg. The samples were heated from RT to
700uC with a heating ramp of 10uC min21 using a
Thermogravimetric analyser from Mettler Toledo, model TGA/SDTA851e.
Water vapour sorption. An ASAP 2020 instrument from
Micromeritics was used for the water sorption studies. Prior to
analysis the samples were degassed at 95uC under vacuum for
Figure 4. Electron microscopy images of Upsalite. a) SEM micrograph of Upsalite. Scale bar, 1 mm. b) Higher magnification SEM of a region in
a) clearly showing the textural porosity of the material. Scale bar, 200 nm. c) Representative TEM image of Upsalite showing contrast consistent with
a porous material. The image is recorded with under-focused conditions to enhance the contrast from the pores. Scale bar, 50 nm.
10 h. The D-A equation was employed on the appropriate
pressure region for adsorption in micropores. The affinity
coefficient (b) for water was set to 0.2 in the D-A calculations,
which has been shown to be an appropriate value for analysis of
polar surfaces .
Scanning transmission electron microscopy and electron
tomography. Scanning transmission electron microscopy
(STEM) samples were prepared by dispersing the powder in
ethanol and placing 20 ml on a QuantifoilH TEM grid.
Experiments were performed on an FEI Tecnai F20 (FEI Company, The
Netherlands) operated at 200 kV. Images were recorded on a
high-angle annular dark-field detector (HAADF). The Dual-Axis
Tomography Holder Model 2040 (Fischione Instruments, PA,
USA) was used in a linear tilt scheme to acquire a single-axis
tiltseries with image acquisition increments of 2u. Automated
focusing, image shifting, and acquisition of HAADF STEM
images over an angular range of 662u were achieved using the
Explore3D software (FEI Company, The Netherlands). The 3D
reconstructions were computed using a simultaneous iterative
reconstruction technique, with 20 iterations, in Inspect3D (FEI
Company, The Netherlands). Models for 3D visualisation were
created in Amira Resolve RT FEI (Visage Imaging Inc., USA).
Transmission Electron Microscopy (TEM). HRTEM
images were taken with a JEOL-3010 microscope, operating at
300 kV (Cs 0.6 mm, resolution 1.7 A). Images were recorded
using a CCD camera (model Keen View, SIS analysis, size
102461024, pixel size 23.5623.5 mm) at 30 000100 0006
magnification using low-dose conditions on as-crushed samples.
Figure S1 X-ray pattern of the material obtained when
water was deliberately added to the synthesis. All peaks in
the pattern corresponds to nesquehonite, Mg(HCO3)(OH)?2 H2O,
(PDF# 00-020-0669). No signs of residual MgO can be detected
in the pattern.
Figure S3 FTIR spectrum for the in-situ sample
collected from the reaction vessel after 3 hours of reaction,
together with a reference sample.
Figure S4 STEM Tomography and TEM. Internal pore
structure from electron tomography. Slices through two
electron tomographic reconstructions of the Upsalite material.
Essentially representing an internal cross-section, pore size
measurements from these volumetric slices indicate pore sizes
with one dimension measuring on average 16 nm, and the other
varying between 50 nm and 8 nm (S4.1). TEM image showing
Upsalite sample after a long period (,1 min) under the electron
Overview of earlier work.
We thank Prof. Peter Lazor for assistance with interpretation of the Raman
spectra, Dr. Alfonso E. Garcia-Bennett for recording the TEM images and
Dr. Martin Sjodin for initial discussions regarding the reaction mechanism.
Conceived and designed the experiments: JF AM MS. Performed the
experiments: SF KG JF. Analyzed the data: JF AM SF MS KG.
Contributed reagents/materials/analysis tools: JF AM SF MS KG. Wrote
the paper: JF AM SF MS KG.
1. Fadeel B , Kasemo B , Malmsten M , Strmme M ( 2010 ) Nanomedicine: Reshaping clinical practice . J. Intern. Med . 267 : 2 - 8 .
2. European Commission ( 2012 ) A European strategy for Key Enabling Technologies - A bridge to growth and jobs . Brussels: COM 341 final.
3. Mihranyan A , Ferraz N , Strmme M ( 2012 ) Current status and future prospects of nanotechnology in cosmetics . Prog. Mater. Sci . 57 : 875 - 910 .
4. Varshney HM , Shailender M ( 2012 ) Nanotechnology; Current Status In Pharmaceutical Science: A Review . Int . J. Therap. Appl . 6 : 14 - 24 .
5. Truitt B ( 2009 ) Magnesium Carbonate . In Handbook of Pharmaceutical Excipients , 6th ed.; Rowe R , Sheskey P , Quinn M , Eds.; Pharmaceutical Press: London; 397 - 400 .
6. Deelman JC ( 2011 ) In Low Temperature Formation of Dolomite and Magnesite [Online] version 2.3 ed.; Compact Disc Publications: Eindhoven , 2011 . Available: www.jcdeelman.demon.nl. Accessed 2012 Apr 18 .
7. Pohl W ( 1990 ) Genesis of Magnesite Deposits - Models and Trends. Int. J. Earth. Sci . 79 : 291 - 299 .
8. Neuberg C , Rewald B ( 1908 ) Ueber Kolloide Und Gelatinose Verbindungen Der Erdalkalien . Colloid Polymer Sci . 2 : 354 - 357 .
9. Kurov VI ( 1961 ) Alkyl Carbonate Salts .5. Methyl Carbonate Salts of Bivalent Metals . Russ. J. Gen. Chem . 31 : 9 - 11 .
10. Buzagh A ( 1926 ) Ueber Kolloide Losungen Der Erdalkalikarbonate . Kolloid Z. 38 : 222 - 226 .
11. Khan N , Dollimore D , Alexander K , Wilburn FW ( 2001 ) The Origin of the Exothermic Peak in the Thermal Decomposition of Basic Magnesium Carbonate . Thermochim. Acta 367-368 : 321 - 333 .
12. Hashimoto H , Tomizawa T , Mitomo M ( 1968 ) Exothermic Process in the Differential Thermal Analysis Curves of Basic Magnesium Carbonate . Kogyo Kagaku Zasshi 71 : 480 - 484 .
13. Sawada Y , Uematsu K , Mizutani N , Kato M ( 1978 ) Thermal Decomposition of Hydromagnesite 4 MgCO3NMg(OH)2N4 H2O under Different Partial Pressures of Carbon-Dioxide . Thermochim. Acta 27 : 45 - 59 .
14. Choudhary VR , Pataskar SG , Gunjikar VG , Zope GB ( 1994 ) Influence of Preparation Conditions of Basic Magnesium Carbonate on Its ThermalAnalysis . Thermochim. Acta 232 : 95 - 110 .
15. Dell RM , Weller SW ( 1959 ) The Thermal Decomposition of Nesquehonite MgCO3N3 H2O and Magnesium Ammonium Carbonate MgCO3N(NH4)2- CO3N4 H2O . Trans Faraday Soc 55 : 2203 - 2220 .
16. Clark CB ( 1946 ) X-Ray Diffraction Data for Compounds in the System CaOMgO-SiO2 . J. Am . Ceram. Soc. 29 : 25 - 30 .
17. Rutt HN , Nicola JH ( 1974 ) Raman Spectra of Carbonates of Calcite Structure . J. Phys. C Solid State Phys . 7 : 4522 - 4528 .
18. Gouadec G , Colomban P ( 2007 ) Raman Spectroscopy of Nanomaterials: How Spectra Relate to Disorder, Particle Size and Mechanical Properties . Prog. Cryst. Growth Charact. Mater . 53 : 1 - 56 .
19. Raade G ( 1970 ) Dypingite, a New Hydrous Basic Carbonate of Magnesium, from Norway . Am. Mineral . 55 : 1457 - 1465 .
20. Santamaria M , Di Quarto F , Zanna S , Marcus P ( 2007 ) Initial Surface Film on Magnesium Metal: A Characterization by X-Ray Photoelectron Spectroscopy (XPS) and Photocurrent Spectroscopy (Pcs) . Electrochim. Acta 53 : 1314 - 1324 .
21. Splinter SJ , Mcintyre NS , Lennard WN , Griffiths K , Palumbo G ( 1993 ) An Aes and Xps Study of the Initial Oxidation of Polycrystalline Magnesium with Water-Vapor at Room-Temperature . Surf. Sci. 292 : 130 - 144 .
22. Siegbahn K , Nordling C , Johansson G , Hedman J , Heden PF , et al.( 1969 ) Esca Applied to Free Molecules. North-Holland Publishing Company: Amsterdam.
23. Xie YM , Sherwood PMA ( 1990 ) X-Ray Photoelectron Spectroscopic Studies of Carbon-Fiber Surfaces .11. Differences in the Surface-Chemistry and Bulk Structure of Different Carbon-Fibers Based on Poly(Acrylonitrile) and Pitch and Comparison with Various Graphite Samples . Chem. Mater. 2 , 293 - 299 .
24. Gardner SD , Singamsetty CSK , Booth GL , He GR , Pittman CU ( 1995 ) Surface Characterization of Carbon-Fibers Using Angle-Resolved Xps and Iss . Carbon 33 : 587 - 595 .
25. Sing KSW , Everett DH , Haul RAW , Moscou L , Pierotti RA , et al.( 1985 ) Reporting Physisorption Data for Gas Solid Systems with Special Reference to the Determination of Surface-Area and Porosity ( Recommendations 1984 ). Pure Appl. Chem . 57 : 603 - 619 .
26. Brunauer S , Emmett PH , Teller E ( 1938 ) Adsorption of Gases in Multimolecular Layers . JACS 60 : 309 - 319 .
27. Pires J , Pinto ML , Carvalho A , de Carvalho MB ( 2003 ) Assessment of Hydrophobic-Hydrophilic Properties of Microporous Materials from Water Adsorption Isotherms . Adsorption 9: 303 - 309 .
28. Dubinin MM , Astakhov VA ( 1971 ) Description of Adsorption Equilibria of Vapors on Zeolites over Wide Ranges of Temperature and Pressure . In Molecular Sieve Zeolites-Ii, American Chemical Society: Washington, D C . 102 : 69 - 85 .
29. Verhoeven L , Lodewyckx P ( 2001 ) In Comparison of Dubinin-Radushkevich Micropore Volumes Obtained from N2, CO2 and H2O-Adsorption Isotherms , Carbon 2001 proceedings, Lexington (KY, USA), American Carbon Society: Conference proceeding avalibe from : http://acs.omnibooksonline.com/data/ papers/2001_10.2.pdf.
30. Botha A , Strydom CA ( 2003 ) Dta and FT-IR Analysis of the Rehydration of Basic Magnesium Carbonate . J. Therm. Anal. Calorim . 71 : 987 - 995 .
31. Ardizzone S , Bianchi CL , Fadoni M , Vercelli B ( 1997 ) Magnesium Salts and Oxide: An Xps Overview . Appl. Surf. Sci . 119 : 253 - 259 .
32. Dastgheib SA , Karanfil T ( 2005 ) The Effect of the Physical and Chemical Characteristics of Activated Carbons on the Adsorption Energy and Affinity Coefficient of Dubinin Equation . J. Colloid Interface Sci . 292 : 312 - 321 .