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From greenhouse gas to feedstock: formation of ammonium carbamate fromCO2 and NH3 in organic solvents and its catalytic conversion into ureaunder mild conditions

Francesco Barzagli,a Fabrizio Mani*a and Maurizio Peruzzinib

Received 12th October 2010, Accepted 8th February 2011DOI: 10.1039/c0gc00674b

The capture of carbon dioxide by ammonia in both aqueous and non-aqueous solutions wasinvestigated at atmospheric pressure and 273 K under different operating conditions. The CO2

capture is fast and efficient ranging between 78 and 99%, depending on both the NH3

concentration and the solvent nature. The precipitation of solid mixtures of ammoniumbicarbonate, ammonium carbonate and ammonium carbamate occurred in ethanol–watersolution. Selective precipitation of ammonium carbamate was achieved by reacting gaseous CO2

and NH3 in anhydrous ethanol, 1-propanol or N,N-dimethylformamide (DMF) in a flow reactorthat operates in continuous. In the second step of the process, the pure ammonium carbamate isused to produce urea with good yield (up to 54% on carbamate basis) at 393–413 K in the presenceof inexpensive Cu(II) and Zn(II) catalysts. The yield of urea depends on several factors includingthe catalyst, the reaction temperature and the reaction time. Identification and quantification ofurea in the reaction mixtures was obtained by analysis of its 13C NMR spectrum. A preliminarymechanistic interpretation of the catalytic reaction is also briefly presented and commented.

Introduction

The reduction of anthropogenic CO2 emissions is consideredone of the most urgent challenges1 and imperative strate-gies must be adopted to limit CO2 emissions by improvingenergetic efficiency, bolstering the use of alternative energysources (biomass, wind farms and photovoltaic cells), favouringthe change in fuels and adopting efficient CO2 capture andsequestration technologies. Among the last technology, theammonia scrubbing process provides the advantage of high CO2

loading capacity and absorption efficiency with no absorbentdegradation.2 However, this process suffers from serious energypenalties due to NH3 loss avoidance and regeneration and for thecost related to its separation from concentrated CO2 that mustbe compressed and sequestered. The ultimate goal for setting upa sustainable CO2 capture process must be therefore aimed atlowering the costs of the process and, even more important,at maximizing the net balance CO2(captured) - CO2(emitted),3 thatrepresents the “true” CO2 avoided.

In our laboratory we are developing a new concept of CO2

capture technology which combines the necessary CO2 abate-

aUniversity of Florence, Department of Chemistry, via della Lastruccia 3,50019, Sesto Fiorentino, Firenze, Italy. E-mail: fabrizio.mani@unifi.itbICCOM CNR, via Madonna del Piano 10, 50019, Sesto Fiorentino,Firenze, Italy

ment with the production of commercially valuable products.4

Turning carbon dioxide into a feedstock for producing usefulcommodity chemicals in mild conditions would indeed circum-vent most of the drawbacks of the energy consuming stepsof CO2 desorption, absorbent regeneration as well as of CO2

transportation and disposal in geological cavities, oceans orelsewhere.

In a previous experimental study,5 we reported that theabsorption of CO2 by 0.85–10.0 M NH3 (1.46–18.1 wt%)aqueous solutions occurs with high efficiency and load capacityproducing solutions of bicarbonate, carbonate and carbamateammonium salts in a ratio which depends on the amount ofabsorbed CO2 with respect to the concentration of free NH3

in solution, as inferred by 13C NMR spectroscopy. The highwater solubility of the ammonium salts of each species, HCO3

-,NH2CO2

- and CO32-, prevented the crystallization of any solid

compound at the end of the absorption experiments.The reaction of gaseous CO2 and NH3 in dry conditions

at ambient temperature and under atmospheric pressure isexoenthalpic and produces ammonium carbamate [see below,reaction (4)]. However, the process is scarcely suited for practicalCCS applications because the low reaction rate of CO2 and NH3

in the gas phase causes a severe loss of NH3 and/or scarceCO2 removal efficiency, accompanied by the difficult removal ofsolid ammonium carbamate from the absorbent reactor.6 On thecontrary, doing the reaction in liquid phase facilitates the process

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Fig. 1 Simplified flow diagram of the absorber-filtration cyclic configuration (A) and simplified sketch of the gaseous CO2–NH3 absorber (B).

and makes easy the separation of the solid from the solution,thus allowing an efficient recycle of the unreacted scrubbingsolution. In summary, the CO2 capture by a wet method shouldbe highly preferable.

In this paper we report our results on CO2 capture by NH3

in ethanol–water solution as well as in anhydrous ethanol, 1-propanol and N,N¢-dimethylformamide (DMF). The processesresult in the formation of solid mixtures of ammonium bicarbon-ate and ammonium carbamate or of pure ammonium carbamatethat contain all of the captured CO2. The capture of CO2 hasbeen carried out at room pressure with chilled NH3 (273 K)in order to minimize NH3 loss. Remarkably, the unreactedammonia solution, once separated by filtration from the solidcompounds, is entirely reclaimed into the absorbent reactor.

Ammonium bicarbonate is a market product with a varietyof uses. In particular, it has been used as a nitrogen fertilizer inChina for over 30 years.7 Recently, the use of both ammoniumbicarbonate and carbamate has been patented for the recovery offreshwater from seawater by forward osmosis.8 Ammonium car-bamate has been also tested as a NH3 generator for NOx abate-ment in diesel exhaust gases, to recover manganese from steel-making plant slag and to soil remediation.9 More important,ammonium carbamate is the intermediate for the production ofurea, the most used nitrogen fertilizer worldwide (more than 108

metric tons per year).10 The commercial production of urea isbased on substantially similar patented processes where carbondioxide is reacted with excess of ammonia (NH3/CO2 molarratio up to 4) at high temperature (450–500 K) and pressure(150–250 bar) to produce ammonium carbamate, which is thendehydrated to urea.

Here we report our studies about the conversion of the solidammonium carbamate, obtained in the anhydrous process ofCO2 capture by NH3, into urea. The process takes place inrelatively mild conditions, i.e. 393–413 K and £14 bar (thepressure generated by the thermal decomposition of ammoniumcarbamate). The reaction was accomplished in a sealed vesselcontained ammonium carbamate in the presence of a transitionmetal catalyst (1–1.5 wt%). Depending on both the reactionparameters (temperature and time) and the catalyst nature,

conversion of carbamate into urea ranged between 17 and 54%.Apart for the crucial use of the catalyst, the urea formation fromheating solid ammonium carbamate has been recognized for along time.11

Results and discussion

Formation of solid ammonium carbamate and ammoniumbicarbonate

Due to the low solubility of both ammonium carbamate andbicarbonate in ethanol, some CO2 capture experiments werecarried out in ethanol–water ammonia solutions (275–335 mL)using three different NH3 concentrations (1.10, 2.06 and 2.72 M;see Experimental). The absorption experiments were carriedout at 273 K using an home-built glass absorber4b immersedinto a thermostatted bath. The CO2/N2 gas mixture (12% v/v)simulating flue gas was flowing at the bottom of the absorbentapparatus through a sintered glass diffuser. During the absorp-tion experiments the slurry was circulated with a peristalticpump in a closed loop between the absorber and the filtrationunit (Fig. 1A). In the latter, the solid was continuously separatedfrom the solution that was reclaimed back to the absorber.In a different type of absorption experiments, gaseous CO2

and, separately, NH3 were continuously introduced through twosintered glass diffusers in the absorber unit (Fig. 1B) containingthe anhydrous solvent (300 mL, ethanol, 1-propanol or DMF).The gas exiting from the absorbent unit was dried and purifiedfrom NH3 before being analysed with a gas chromatograph.

The equilibria which describe the reaction of carbon dioxidewith aqueous ammonia are:12

NH3 + CO2 + H2O � NH4+ + HCO3

-, K eq(273) = 3.22 ¥ 103 (1)

NH3 + HCO3- � NH2CO2

- + H2O, K eq(273) = 7.30 (2)

NH3 + HCO3- � CO3

2- + NH4+, K eq(273) = 2.76 ¥ 10-1 (3)

If an excess of NH3 is maintained in the system, the overallreactions can be rewritten as

1268 | Green Chem., 2011, 13, 1267–1274 This journal is © The Royal Society of Chemistry 2011

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2NH3 + CO2 � NH2CO2- + NH4

+, K eq(273) = 2.35 ¥ 104 (4)

2NH3 + CO2 + H2O � CO32- + 2NH4

+, K eq(273) = 8.89 ¥ 102 (5)

The concentration of both NH2CO2- and CO3

2- decreases byincreasing CO2 absorption; meanwhile, a simultaneous increaseof bicarbonate occurs according to the equilibria (6) and (7)

CO32- + CO2 + H2O � 2HCO3

- (6)

NH2CO2- + CO2 + 2H2O � 2HCO3

- + NH4+ (7)

In general, a greater ammonia concentration increases theefficiency of CO2 removal, but lowers the loading capacityand increases the loss of NH3 from the scrubbing solution.In the NH3/H2O–C2H5OH absorption experiments, no freshNH3 was added to the absorber to replace that consumed sothat the absorption efficiency decreased with increasing of theabsorbed CO2. The absorption experiments were stopped afterca. 200 min. The average CO2 removal efficiency in the entireexperiments is comprised between 78.5% (NH3 1.10 M) and98.9% (NH3 2.72 M). The solids recovered at the end of the ab-sorption experiments were different mixtures of NH4HCO3 andNH2CO2NH4 together with a smaller amount of (NH4)2CO3.Pure NH4HCO3 or NH2CO2NH4 were never obtained at the endof each experiment. In order to estimate the composition of thesolid obtained during the absorption process, 13C NMR spectrawere recorded in D2O and compared with those of standardsolutions of NH4HCO3 and NH2CO2NH4 in the same solvent.The results of three experiments at different NH3 concentrationsare reported in Table 1.

The maximum amount of solid (6.58 g after 184 min) andthe maximum concentration of carbamate in the solid mixture(97.5% after 114 min) were obtained with the 2.06 M NH3

solution that represents the best compromise amongst severalopposite effects, as, for example, the amount of the speciesin solution that increases with NH3 concentration and thesolubility of the different species that increases with increasingwater–ethanol ratio that, in turn, increases with NH3 concen-tration. Furthermore, the amount of carbamate should increasewith NH3 concentration, but should decrease with increasingof water–ethanol ratio, according to the reverse of reaction(2). Analogous experiments were carried out using other liq-uid mixtures, such as pyridine/water, dimethylsulfoxide/water,

DMF/water, dioxane/water, 2-propanol/ethanol/water, 1-butanol/ethanol/water. In all of these experiments differentamounts and different ratios of the three ammonium salts wereobtained, without any selectivity.

In the experiments, aimed at obtaining pure NH2CO2NH4,anhydrous operating conditions were sought in order to avoidthe competitive formation of HCO3

- and CO32- [reactions (1)

and (5)]. Gaseous NH3 and, separately, CO2 in molar ratiocomprised between 0.96 and 2.20 were continuously fed atthe bottom of the absorbent reactor (Fig. 1B) containing300 mL of different organic solvents (ethanol, 1-propanol,DMF) thermostatted at 273 K. The flow rate of the NH3/CO2

mixture was the best compromise between the rate of solidcarbamate formation and the CO2 capture efficiency. Increasingthe flow rate increases the amount of carbamate at the expenseof a reduced CO2 capture efficiency and of an appreciable loss ofammonia. Decreasing the flow rate has the opposite effect. Theduration of each experiment was fixed at 8 h. Each experimentwas carried out at least in duplicate for any solvent and anyconcentration used. The efficiency of CO2 capture was alwaysgreater than 93% (with a maximum value of 98% in DMF) byrunning the reaction with a slight excess of NH3 (NH3/CO2 ratio2.10–2.20) with respect to the stoichiometry of the reaction (8)

2NH3(gas) + CO2(gas) → NH2CO2NH4(solid) (8)

The results of typical experiments carried out under differentoperating conditions are reported in Table 2 and summarised inFig. 2. The slight differences obtained with the three solventsare mainly due to the slight differences of NH3/CO2 ratiosand, presumably, to the different solubility of either gas andcarbamate.

The conversion of NH3 into solid NH2CO2NH4 rangedbetween 71.8 and 79.4% [Table 2, entries 3, 6, 9; on molarscale, according to the stoichiometry of reaction (8)]. Absorptionexperiments with ca. 1/1 NH3/CO2 ratio (0.96–1.03) gave 93.6–96.9% (Table 2, entries 1,4,7) conversion of NH3 in crystallinesolids, but the CO2 capture efficiency was reduced to 55–66%. Again, the search for the best compromise between CO2

capture efficiency and NH3 conversion to pure NH2CO2NH4

asked for a NH3/CO2 molar ratio of ca. 1.5 that resultedin high CO2 capture (85–90%) and NH3 conversion (94–98%;Table 2, entries 2, 5, 8) to crystalline NH2CO2NH4. In all

Table 1 CO2 absorption efficiency and composition of the solid mixture of ammonium carbamate, bicarbonate and carbonate obtained fromdifferent aqueous ammonia–ethanol solutions

Composition (%)f (n/n)

c/mol dm-3H2O–C2H5OHa

(v/v) tb/min CO2 absc (%)CO2/NH3

d

(n/n) me/g NH2CO2- HCO3

- CO32- M̄ g

solid/CO2h

(n/n)

1.10 0.0741 128 86.8 0.385 3.30 75.5 15.8 8.8 79.8 0.354192 78.5 0.537 5.61 48.7 42.6 8.6 80.6 0.431

2.06 0.148 114 90.5 0.181 2.88 97.5 2.2 0.3 78.1 0.340184 89.8 0.289 6.58 86.8 8.6 4.6 79.6 0.481

2.72 0.222 124 100 0.145 1.33 87.3 8.3 4.5 79.0 0.127187 98.9 0.215 3.80 83.3 10.3 6.6 79.6 0.245

a Volume ratio between aqueous ammonia and ethanol. b Absorption time. c Average absorption efficiency during the absorption time. d Averageloading during the absorption time. e Mass of the solid mixture. f Average composition, on molar scale, of the solid mixtures recovered at the end ofany experiment. g Average molar mass computed from the mixture composition. h Molar ratio between the solid mixture and the absorbed CO2.

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Table 2 CO2 absorption efficiency and yield of NH3 conversion into solid ammonium carbamate at different operating conditions

Entry solvent NH3/CO2 n/n CO2(abs) % NH2CO2NH4/NH3 % (n/n) NH2CO2NH4/CO2 % (n/n) NH2CO2NH4, m/g h-1

1 ethanol 1.03 66.4 93.6a 74.7 2.332 ethanol 1.59 89.7 94.0a 83.1 3.773 ethanol 2.17 93.1 77.0 89.7b 4.084 1-propanol 0.96 54.6 96.6a 66.9 2.345 1-propanol 1.51 88.1 98.2a 84.0 3.936 1-propanol 2.20 94.5 71.8 83.7b 3.817 DMF 0.96 63.4 96.9a 82.1 2.618 DMF 1.58 85.0 93.7a 85.9 3.759 DMF 2.10 98.0 79.4 85.2b 4.21

a NH3 is the species in defect with respect to the stoichiometric ratio. b CO2 is the species in defect with respect to the stoichiometric ratio.

Fig. 2 CO2 absorption efficiency (—) and ammonia conversion intocarbamate (%) (---) as a function of the NH3/CO2 molar ratio.

of these experiments, ammonium carbamate was obtained inquite pure form, as inferred from the comparison of 13C NMRspectra in D2O of the different experimental samples with thatof reagent grade commercial ammonium carbamate. In eachexperiment, the unreacted solution was entirely recycled to theabsorbent apparatus. Remarkably, the loss of NH3 from the moreconcentrated absorbent solution never exceeded 1.5% (on molarbasis) for the entire duration of each absorption experiment.

Conversion of ammonium carbamate into urea

The urea manufacture at industrial scale, typically an urea plantproduces 1500 tons per day, is carried out by feeding excessammonia and carbon dioxide to the synthesis reactor at 450–500 K, 150–250 bar. The process involves the intermediateformation of ammonium carbamate [reaction (8)] that is suc-cessively dehydrated to form urea

NH2CO2NH4 → NH2CONH2 + H2O (9)

While reaction (8) is fast and exothermic (DH◦ =-151 kJ mol-1), reaction (9) is slow and endothermic (DH◦ =32 kJ mol-1)13 and does not go to the completion underindustrial processing conditions. As a matter of fact, the rathersmall enthalpy value suggests that the reaction (9) should besubstantially right hand shifted at high temperature and thelimiting value of urea yield is dictated rather by the reactionrate than the equilibrium value. Moreover, we can point out that

the reaction (9) is an oversimplification of the multi-phase andmulti-component equilibria occurring at high temperature andpressure. The set of the main reactions are:

NH2CO2NH4(s) � 2NH3(g) + CO2(g)

2NH3(g) + CO2(g) � (NH2)2CO(l) + H2O(l)

NH2CO2NH4(s) + H2O(l) � NH4HCO3(s) + NH3(g)

NH2CO2NH4(s) + 2H2O(l) + CO2(g) � 2NH4HCO3(s)

Therefore, the final yield of urea is the result of the competitionbetween the conversion reactions of carbamate into urea andits decomposition reactions to, mainly, bicarbonate. The aboveequilibria proceed differently to each other at the same workingconditions. In general, the yield of the reaction is in the order 30–55% on NH3 basis (60–70% on CO2 basis) and strongly dependson reaction temperature, pressure, time and NH3/CO2 ratio.

In order to circumvent the energy penalties affecting theconventional industrial production of urea requesting highworking temperature and pressure, we looked for the possibilityto promote the dehydration of NH2CONH4 by carrying outthe reaction in the presence of a transition metal catalyst.Surprisingly, very few reports may be found in both scientificand patent literature describing the use of a catalyst to bringabout the synthesis of urea and none of them involves the twostep procedure which is industrially used.14

In keeping with our expectations, heating ammonium car-bamate contained in a sealed vessel (home-built stainless steelairtight container) for 2–3 days at 393–413 K (in one experimentup to 433 K) in the presence of different catalysts (catalystloading: 1.0–1.5%, on molar basis, Table 3) resulted in anincreased production of urea confirming the possibility todevelop a catalytic synthesis of urea at relatively low pressurebecause the maximum pressure generated by the thermaldecomposition of carbamate in the autoclave was 14 bar at413 K. At the end of the reaction, 13C NMR spectroscopy inD2O (see Experimental) showed that the solid mixture containedvariable amounts of unreacted carbamate (d = 165.84), urea(d = 162.95) and carbonate/bicarbonate mixture (d = 163.3–164.3) depending on the operating conditions. No other productwas detected in the 13C NMR spectra, in particular no trace ofbiuret (d = 158.08) was found. Presumably, dimerization of ureato biuret occurs only at temperature and pressure higher thanthose occurring in our experimental conditions. Urea could be

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Table 3 Catalysts and experimental conditions employed in the carbamate conversion into urea and yield of the reactionsa

Entry Catalyst % (mol)b T/K Time/day Ureac (% mol)

1 copper 15.3 393 3 50.32 CuCl2·2H2O 1.00 393 3 48.13 CuCl2·2H2O 1.00 413 3 53.64 CuO 1.00 393 3 53.75 CuCO3·Cu(OH)2 0.97 403 2.5 51.06 ZnCl2 2.50 403 2 32.07 ZnCl2 1.25 413 3 45.08 NiCO3·2Ni(OH)2·H2O 1.12 413 3 25.99 NiCO3·2Ni(OH)2·H2O 1.00 433 3 31.010 RuCl3 1.26 413 2.5 33.011 MnCl2 1.54 393 3 21.012 FeCl3 1.00 393 3 17.3

a Blank experiments up to 403 K and 3 days gave less than 3% yield. b The catalyst percentage is referred to the starting amount of carbamate.c Percentages are referred to the starting amount of carbamate.

recovered from the solid residue by heating the reaction mixture(333 K) till to constant weight, which removed the unconvertedNH2CO2NH4, the ammonium carbonate/bicarbonate mixtureand water as gaseous NH3, CO2 and H2O, leaving pure urea andthe catalyst. The purity of the product was checked by 13C NMRanalysis and the yield of the reaction (9) could be determinedafter subtracting the weight of the catalyst from the constantweighted residue. Both gaseous NH3 and CO2 obtained from thedecomposition of the mixture of the unreacted carbamate andof the by-product ammonium bicarbonate could be recoveredand recycled, whereas the catalyst could be separated from themelt urea and reused.

A variety of reaction promoters in different ratio with respectto carbamate, were tested and the most efficient ones arereported in Table 3. The most efficient catalysts were copper(II)and zinc(II) compounds and elemental copper. Under the sameoperating conditions, manganese, iron, chromium and cobaltcompounds gave a conversion lower than 25%. RuCl3 liesbetween the two groups of catalysts with a 33% of conversion.The other tested metal compounds gave generally less than 5%of urea. Blank experiments run without any catalyst gave lessthan 3% of urea again confirming the catalytic activity of theadded metal salt. In general, increasing the catalyst loading didnot improve the yield and in most experiments a catalyst loadingof 1–1.5% (on molar scale) with respect to carbamate gave thebest results. In contrast, 15% loading of powdered elementalcopper was necessary to obtain the same results of the Cu(II)compounds indicating that Cu(0) is a poor catalyst with respectto copper(II) salts. Increasing both decomposition temperature(Table 3, entries 2, 3 and 8, 9) and time (entries 6, 7) increasedthe conversion.

The role of the metal catalyst in the reaction mechanismresponsible for the carbamate conversion to urea (eqn (9)) is,at the present, uncertain. However, although unravelling themechanism of the catalytic reaction is out of our current interest,some hypotheses seem plausible.

If one consider that (i) either hydrated or anhydrous metalcompounds have a similar effect and (ii) the amount of thecatalyst is 1–1.5% of the water formed by the reaction (9), it isconceivable that the metal salts cannot simply shift the reactionto the right acting as dehydrating agents. Furthermore, theformation of metal(II)–urea complexes that could also favour

the conversion reaction or the capacity of the metal(II) ionsto provide a platform to bring two molecules of NH3 in closeproximity to each other thus facilitating the reaction with CO2,cannot explain the better performances of Cu(II) and Zn(II)compared to Ni(II) and Co(II), for example, as all of the fourmetal ions give ammonia15 and urea complexes.16 On theseassumptions, we may tentatively propose that the carbamateconversion to urea can occur through the following steps [M(II)stands for either Cu(II) or Zn(II)]

NH CO NH 2NH + CO [M(NH ) ]

NH CO

2 2 4heat

3 2M( )

3 62+ CO2

2

II⎯ →⎯ ⎯ →⎯⎯ ⎯ →⎯⎯

NNH + H O + [M(NH ) ] [M(NH ) ]2 2 3 42+ 2NH3

3 62+⎯ →⎯⎯

(10)

The ammonia formed by the thermal decomposition ofammonium carbamate may easily coordinate Cu2+ or Zn2+

ions forming the corresponding hexakis-amino complexes[M(NH3)6]2+ (I). Then, the reaction of CO2 with two adjacentmolecules of ammonia may well account for the metal-assistedformation of one urea molecule as a dihapto-coordinated ligandin the transient [M{k2-N,N (NH2)2CO}(NH3)4]2+ complex (II).Water is also released along this reaction step. The urea complex,once formed, may easily exchange urea with the excess ofammonia restoring the more stable hexakis-amino complex (I),possibly via the tetracoordinated [M(NH3)4]2+ complex (III).Fig. 3 illustrates this mechanistic hypothesis. If we taken forgranted this putative hypothesis, the greater efficiency of Cu(II)and Zn(II) with respect to the other tested metal ions couldbe readily explained on the basis of the peculiar coordinatingproperties of Cu(II) and Zn(II) ions that, in the presence of anexcess of several ligand molecules, including ammonia, formsix-coordinated complexes (I) containing four strongly andtwo labile coordinated ligands.15 Thus, two cis disposed NH3

molecules of either [Cu(NH3)6]2+ or [Zn(NH3)6]2+ complexesmay easily react with CO2 yielding the urea complex II whicheventually undergoes substitution reactions where the coordi-nated urea molecule is easily replaced by two ammonia ligands.The proposed reaction pathway could also explain the reducedconversion yield as the metal(II)/carbamate ratio increases andthe poor performances exhibited by other transition metal saltswith respect to Zn2+ and Cu2+ ions. Indeed, the formation of six-coordinated Cu(II) and Zn(II) complexes at high temperature

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Fig. 3 Proposed pathway of the M2+ catalysed reaction betweenCO2 and two coordinated ammonia molecules (M = Cu, Zn). Ureadisplacement from the coordination sphere by the stronger ligand NH3

may regenerate easily the active catalytic species.

requires a large excess of ammonia with respect to the metal(II)ion while Cr(III), Fe(III), Mn(II), Co(II) and Ni(II), whichgave worse catalytic performances, form very stable and inertoctahedral six-coordinated ammonia complexes and, becauseof this kinetic inertness, the ligand displacement step is hard tobe accomplished under the used working conditions.

Studies are in progress to substantiate this putative hypothesisby investigating the specific role of amino complexes in thecatalytic decomposition of ammonium carbamate to urea.Preliminary results indicate that catalytic tests carried out withisolated amino complexes [M(NH3)6]2+ may increase both thedecomposition rate of ammonium carbamate and the urea yield.

Experimental

All reagents were reagent grade. Ammonium carbamate,ammonium bicarbonate, ethanol, 1-propanol and N,N¢-dimethylformamide (Sigma-Aldrich) were used as received.Standard NH3 solution 15.2 M (Sigma-Aldrich) was used toprepare the ethanol/water/ammonia solutions. Pure CO2 andN2 used to simulate flue gas and pure NH3 were obtained byRivoira. Flow rates of N2, CO2 and NH3 were measured withgas mass flow meters (Aalborg) equipped with gas controllers(Cole Parmer). The inlet and outlet CO2 concentrations in theflue gas mixture were measured with a Varian CP-4900 gaschromatograph calibrated with a 10% v/v CO2/N2 referencegas (Rivoira).

The cyclic absorption–filtration device consisted of the ab-sorber and the filtration units that are connected to eachother by means of a peristaltic pump (Masterflex L/S) that

allows the absorbent slurry and the filtered solution to circulatecontinuously in a closed loop between the absorber and thefiltration unit (Fig. 1A). The temperature of the absorbentsolution was kept constant at 273 K by a thermostatted waterbath (Julabo model F33-MC refrigerated bath).

To mimic flue gas, a gas mixture containing 12% (v/v) CO2 inN2, was continuously fed into the absorber through a sinteredglass diffuser (16–40 mm pores) at the bottom of the absorbentsolution with a flow rate of 14 dm3 h-1. The vent gas cameout from the top of the absorber. The outlet gas was driedby flowing in turn through a condenser cooled at 268 K, aconcentrated H2SO4 solution and a gas purification tower filledwith P2O5, before being analysed with a gas chromatograph.The gas chromatograph measured the percentage of the CO2

absorbed by the ammonia solutions at intervals of 10 min.In the NH3/H2O/C2H5OH experiment, the absorber device

was a home-built glass cylinder with a diameter of 56 mm andheight 300 mm equipped with a thermometer and a combinedpH electrode and fitted with three polyethylene disks threadedon a 2 mm glass rod.4b This arrangement increases the liquidturbulence therefore providing the reaction mixture with asufficient residence time. The absorbent solution was obtainedby mixing 270 mL of ethanol and, separately, 20.0, 40.0, 60.0 mLof NH3 15.2 M, according to the different experiments. Thevolumes of water–ethanol solutions were 275, 295 and 335 mLand the respective NH3 concentrations were 1.10 M, 2.06 M and2.72 M. During all of the CO2 absorption experiments no newNH3 was added to the absorber unit. Each experiment lastedabout 3 h.

In the experiments aimed at obtaining pure carbamate, bothCO2 (12% v/v in N2) and NH3 were simultaneously introducedthrough two separate gas diffusers into 300 mL of the appro-priate solvent (ethanol, 1-propanol or DMF) contained in thethermostatted (273 K) absorber (Fig. 1B). The NH3/CO2 flowratio was in the range 0.96–2.20 and each experiment lasted 8 h.At the end of each experiment, the solid collected by the filtrationunit was washed with CO2 saturated ethanol and diethyl ether inturn before being dried at room temperature with a flow of pureCO2 to avoid the decomposition of either ammonium carbamateor bicarbonate. The conversion of ammonium carbamate intourea was carried out in a home built reactor that comprises aninternal Teflon R© container (the volume of the cylinder is 50 mL)sealed by means of an airtight Teflon R© cap. The Teflon R© reactor iscontained in stainless steel vessel with a screw cap and equippedwith a pressure gauge. The reactor is heated at the appropriatetemperature (393–413 K) by means of a silicone oil heatingbath (IKA HB4). In each experiment, the Teflon R© reactor ischarged with 8.00 g of ammonium carbamate intimately mixedwith the catalyst selected from a variety of oxides or saltsof chromium(III), manganese(II), iron(III), cobalt(II), nickel(II),copper(II), zinc(II), ruthenium(III) and from pure Cr, Fe, Cu, andZn metals.

As an example, the decomposition of ammonium carbamatein the presence of CuO is described below. The other experimentswere run similarly by replacing CuO with the appropriatecatalyst.

In a typical experiment, 8.0 g of ammonium carbamate(0.103 mol) were carefully mixed with 0.080 g of CuO (1.01 ¥10-3 mol) and heated to 393 K for three days. A small amount

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(weighted) of the solid mixture recovered from the reactor atthe end of the conversion experiment was dissolved in D2Oand analysed by 13C NMR spectroscopy. The integration of thesignals (see below) due to HCO3

-/CO32- (d = 163.15), NH2CO2

-

(d = 165.92) and (NH2)2CO (d = 162.92) allowed us to quantifythe molar ratio of the species. Additionally, the solid mixturewas heated in the air at 333 K to constant weight in order todecompose the unreacted ammonium carbamate and the by-product ammonium bicarbonate and to evaporate water. Theresidue, subtracting the mass of the starting CuO and summedto the mass of the sample used for NMR spectrum, weighted3.32 g (0.0553 mol) and was identified as pure urea (53.7% of thestarting ammonium carbamate, on molar scale) by 13C NMRspectroscopy in D2O solution (d = 162.92) upon comparisonwith that of reagent grade urea. The 13C NMR spectra wereobtained with a Bruker AvanceIII 400 spectrometer. Chemicalshifts are to high frequency relative to tetramethylsilane asexternal standard at 0.00 ppm. CH3CN was used as internalreference (CH3CN, d = 1.47). The standard pulse sequence withproton decoupling and NOE suppression was used to acquirethe 13C{1H} with the following acquisition parameters: pulseangle = 90.0◦, at = 1.36 s, dl = 0 s, data points = 65 K, ns = 500–2 ¥103. Increasing the acquisition time and/or the relaxation delay(up to 60 s) does not produce substantial changes in the relativepeak areas of the CO carbon atoms. Normally, the integrationof 13C NMR resonances does not allow reliable quantificationof species having carbon atoms in different environments, due tothe different spin–lattice T1 relaxation time.17 In the species weare dealing with, i.e. carbamate, bicarbonate, carbonate and ureathe CO carbons have a similar chemical environment, so thatthey likely exhibit similar T1. As a matter of fact, to substantiatethis assumption, we have carried out several 13C NMR spectraon standard solutions containing accurately weighted amountsof ammonium carbamate, ammonium bicarbonate and ureain different molar ratios and we have found a quantitativerelationship (maximum error 5%) between the relative peakareas of the CO 13C resonances and the known concentrations ofeach species. The quantification method is therefore empiricallyquite reliable likely reflecting similarities of the relaxation ratefor similar carbons in carbamate, bicarbonate and urea.

Conclusions

The results here reported confirm the advantage of CO2 re-moval by NH3 in both aqueous and non–aqueous solutions interms of absorption efficiency and loading capacity. In orderto circumvent the drawbacks of the energy consuming stepsassociated with NH3 regeneration, its separation from CO2 andCO2 sequestration deep underground or under oceans, we havedevised a procedure that transforms all of the captured CO2 intocommodity chemicals. In particular, the reaction of gaseous NH3

and CO2 in organic solvents (ethanol, 1-propanol and DMF)produces solid ammonium carbamate in a quite pure form andwith a good yield. The best compromise between efficiency ofCO2 capture (85–93%) and yield of solid carbamate with respectto NH3 (93–98%) is obtained with about 1.5 NH3/CO2 molarratio. The CO2 capture by the wet method here reported is to bepreferred to the reaction of CO2 and NH3 in gaseous phase due

to their low rate of reaction, to the severe loss of NH3 and/or tothe scarce CO2 removal efficiency.

Even more interesting, the solid ammonium carbamate re-covered from the CO2 capture is converted into urea undermild conditions with respect to those used in the conventionalindustrial processes for urea manufacturing. We have found that,in the presence of 1–1.5% (on molar scale) of either copper(II)or zinc(II) compounds, ammonium carbamate heated at 393–413 K in a sealed vessel without applying any external pressureproduces urea with variable yields (45–54% with respect tocarbamate) depending on the operating conditions. The entireprocedure of ammonium carbamate and urea production heredescribed offers significant advantages in terms of energy gainwith respect to the higher temperatures and pressures requiredby the conventional industrial processes still representing anefficient capture of anthropogenic CO2 flue emission. Even ifthe present rate of catalytic formation of urea is too low forindustrial application, we have proved that some catalysts canaccelerate the carbamate dehydration to urea in rather mildconditions. Studies are in progress to find more efficient catalystsand to optimise the process.

Acknowledgements

Financial support from MIUR (Rome, Italy) and Ente CRF(Florence, Italy), through FLORENCE HYDROLAB Project, isgratefully acknowledged.

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