Spe 54721 Ms. Agentes Quelantes

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  Copyright 1999, Society of Petroleum Engineers, Inc.This paper was prepared for presentation at the 1999 SPE European Formation DamageConference to held in The Hague, The Netherlands, 31 May–01 June 1999.This paper was selected for presentation by an SPE Program Committee following review ofinformation contained in an abstract submitted by the author(s). Contents of the paper, aspresented, have not been reviewed by the Society of Petroleum Engineers and are subject tocorrection by the author(s). The material, as presented, does not necessarily reflect anyposition of the Society of Petroleum Engineers, its officers, or members. Papers presented atSPE meetings are subject to publication review by Editorial Committees of the Society ofPetroleum Engineers. Electronic reproduction, distribution, or storage of any part of this paperfor commercial purposes without the written consent of the Society of Petroleum Engineers isprohibited. Permission to reproduce in print is restricted to an abstract of not more than 300words; illustrations may not be copied. The abstract must contain conspicuousacknowledgment of where and by whom the paper was presented. Write Librarian, SPE, P.O.Box 833836, Richardson, TX 75083-3836, U.S.A., fax 01-972-952-9435. Abstract Metal chelating materials have been used as iron controlagents in acidizing treatments for many years. There areseveral disadvantages associated with use of the variouscommercially available products. These disadvantages includelow solubility and high toxicity. Furthermore, recent evidencecasts considerable doubt on the effectiveness of commonchelating agents for iron control in sour environments.A relatively simple procedure will be presented that allowspreparation of pure metal (calcium and iron) chelantcomplexes. These complexes serve as models for the in-depthevaluation of various chelating agents over a broad range of temperatures and fluid compositions and have beencharacterized with respect to solubility and thermal stability inspent acid fluids. The experimental results show loss of solubility of calcium complexes as well as thermaldecomposition of ferrous complexes at relatively lowtemperatures. Implications of these results will be discussed.The problems are exacerbated in sour environments whereiron sulfide precipitation occurs at a very low pH. Newevidence will be presented that shows iron sulfideprecipitation will occur below the pH range where chelatingagents are effective. The precipitation of iron sulfide has beenfound to be dependant upon formation basicity and NOT theoverall acidity of the bulk fluid as commonly believed. Theresult implies that simple equilibrium calculations may notreliably predict the chemistry that will occur downholeandmust be used with caution.The conclusion drawn from these results is that the onlyeffective method of preventing precipitation of iron sulfideduring sour well acid treatments is to remove H 2 S from thefluid with sulfide scavenger products. The elimination of largequantities (6-36 kg/m 3 , or 50-300lbs./Mgal) of chelatingagents will have significant cost savings with regards to fluidcost. The evidence provided in this paper shows that theseproducts are not effective and even worse, may precipitate inthe formation and cause severe damage Introduction Chelating agents (or sequestering agents) are commonly usedin well stimulation treatments 1,2 . The primary function of thesematerials is to ‘tie up’ or chemically bind hydrated metal ionsthus changing the ions’ reactivity. Effective chelation canprevent unwanted and potentially damaging effects caused bymetal ions by reacting with components in the fluids (live orreturn). For example, crosslinking and breaker properties canbe changed, thus causing rheological perturbations in fluidcharacteristics and clean-up properties.Probably the most common use of chelating agents is toprevent precipitation of metal compounds in high pH fluids.Iron(III) or ferric ion, readily undergoes hydrolysis in aqueoussolutions when the fluid pH becomes >2.5. This hydrolysisreaction will produce iron(III) hydroxide, which precipitatesduring HCl spending during an acid stimulation treatment.Inclusion of chelating agents in the stimulation fluid is quiteeffective in preventing the hydrolysis reaction and hence theprecipitation of the metal species.Thus, chelating agents appear to be both beneficial tostimulation fluids and chemically straightforward in function.However, this is not necessarily true under all conditions. Oneparticular environment of concern, and the topic of thismanuscript, is that of systems containing hydrogen sulfide(H 2 S). These systems contain only Fe(II) species. Fe(III), astrong oxidant, is chemically incompatible with H 2 S, a strongreducing agent. These two species react, forming Fe(II) andelemental sulfur.Answers to the following questions were sought:1. At what point does FeS begin to precipitate from spent acidfluids? Past experimental work is flawed and several differentvalues can be calculated from literature data.2. What is the theoretical prediction of chelate complexstability at pH values encountered in well stimulation SPE 54721Chelating Agents in Sour Well Acidizing: Methodology or Mythology Michael M. Brezinski, Halliburton Energy Services  2M. BREZINSKISPE 54721 treatments? The usual formation constants are for neutralconditions (pH7) and are far higher than return fluids willreach.3. What are the properties of both calcium and Fe(II) chelatecomplexes? In particular, will calcium and iron complexesremain in solution or precipitate (e.g. citrates have lowsolubility)? If formed, do the complexes possess enoughthermal stability to survive downhole conditions? If thecomplexes are not thermally stable, it will be likely that themetal ions will be released negating the effect of the additives.The answer to these questions requires preparation of thecomplexes.4. Can the Fe(II) complexes actually protect the metal centerfrom hydrogen sulfide and thus prevent formation of FeS inthe acidity range found for return stimulation fluids?Obviously, if the complexes react with hydrogen sulfide, theuse of chelating agents will be of no value. FeS Deposition Pathway FeS precipitation from stimulation fluids has beendocumented 3,4 . The accepted value marking the onset of thispotentially damaging reaction is pH=1.9. However, there areseveral values available for the K SP  of FeS as well as the Ka’sof H 2 S. Crowe predicts the precipitation to occur at pH=1.9from a solution containing 1000 PPM of both H 2 S and Fe (II)ions 3 . The published ‘demonstration’ of the validity of thecalculation is flawed in that the concentration of Fe(II) wasunknown (from the dissolution of a “small quantity of steelwool”). Also, the solution was spent with calcium carbonatechips to a final pH of 5 before the addition of H 2 S whichcaused FeS to form. Dill, et. al. used Fe(III) in hydrogensulfide systems and therefore probably had little H 2 Sremaining (after oxidation to elemental sulfur by Fe(III))toform FeS in their experiments 4 .The K SP  and equilibrium constants, which were obtainedfrom other sources, predict that much higher pH values maybe tolerated before the onset of FeS precipitation. Forexample, pH = 3.2 and 3.75 are calculated from data takenfrom the Handbook of Analytical Chemistry 5 and AdvancedInorganic Chemistry 6  respectively. The following experimentswere designed to probe this apparent discrepancy.These experiments were carried out anaerobically to avoidoxidation of both Fe (II) and H 2 S. 20% calcium chloride waterwas used to simulate spent 15% HCl. This fluid wasrigorously degassed using freeze-pump-thaw cycles on avacuum line. The iron concentration was 5,000 PPM, the pHof the fluid was adjusted with dilute HCl, and H 2 S wasintroduced using a gas dispersion frit. Table 2 shows theresults of the tests. There was no evidence of FeS formation atpH values from 0 to 3. Only when the acidity had decreased toa value of 4-5, did the fluid darken and form black solids.Because these experiments gave values higher than theaccepted range, another experiment was carried out thatincluded the only possible missing reactant – the formation.Table 3 shows the results obtained from this experiment.Fluids at pH values of 0 and 1 evolved CO 2  as acid spendingon the carbonate was substantial. The sample with a pH valueof two provided the key. The marble chip slowly turned black after immersion into the reaction fluid. SEM analysis showedthe surface of the chip to be covered with microcrystallineFeS. The bulk fluid remained clear and colorless. This resultimplies a surface-assisted deprotonation of H 2 S followed byprecipitation of FeS. Formation basicity appears to be theimportant driver of the precipitation reaction at these low pHvalues. One could have clear, clean return fluids but still havedeposited large quantities of FeS in the formation. Carbonateis a very basic anion. A saturated aqueous solution of sodiumcarbonate possesses a pH of 12! This phenomenon may beimportant in other areas as well (e.g., fracturing fluid –matrixcontact). Thus, normal solution equilibrium calculationsshould be used with caution in predicting downhole chemistry. Theoretical Chelate Complex Stability The following calculations show that it is not obviouscomplexes of EDTA and NTA will possess adequate stabilityto exist in the presence of hydrogen sulfide.The relative stabilities of chelated metal ions can easily becompared using the appropriate formation constant: log K f  .For example, the log K f   values for Fe(III) with NTA andEDTA are log 15.9 and log 25.1 respectively. The values forFe(II) and these common materials are much lower being log8.8 and log 14.3 respectively. Still, these log K f   values arehigh enough to give one the feeling that the metal ions will becomplexed and rendered 'inactive' with respect to currentlyaccepted damage mechanisms. However, it must beremembered that these log K ′ f 's are calculated for the fullydeprotonated chelating agent; EDTA only achieves 98%deprotonation when a pH of 12 is reached. Most of the speciesin solution still contain unionized protons. This means thateven if there is a strong interaction of the chelant with a metalion, there will still be uncomplexed coordination sitesavailable for the very reactions we are trying to stop. Figure 1shows a plot of the fractional composition of EDTA vs pH.The curves clearly show that in the pH range of return fluids,the chelating agent is mostly in the completely unionized,mono-ionized, and di-ionized form. These species will notprotect all the coordination sites of the metal center neccessaryto prevent attack by hydrogen sulfide.More realistic information with regard to the stability of the complex at lower pH values must be obtained fromconditional formation constants, K ′ f , where K ′ f   = α n K f  . The α n is calculated by the method of Schwarzenbach 7  Thiscalculation gives the relative concentrations of the partiallydeprotonated chelant species as a function of pH. The stabilityconstants become much lower in the acidity range commonlyencountered in spent acid fluids. For example, the stabilityconstants for Fe(III) with ethylenediaminetetraacetic acid(EDTA) and nitrilotriacetic acid (NTA) at pH3 become log14.5 and log 6.46 respectively. The same comparison withFe(II) shows a decrease in stability to log 3.73 and log 0.55.  SPE 54721CHALATING AGENTS IN SOUR WELL ACIDIZING: METHODOLOGY OR MYTHOLOGY3 Log K f   and Log K ′ f valuesionEDTAEDTA(pH3)NTANTA(pH3)Fe(III)25.114.515.96.46Fe(II)14.33.738.80.55This constitutes a reduction in stability of ca. 8 to 10 ordersof magnitude. Fe(II) and NTA gives a value <log 1! Thissuggests that NTA, under these conditions, will not protectFe(II) from undergoing deleterious reactions in the fluidsystem. The above suggestion will be verified by experiment( vide supra ). Complexes of NTA and EDTA Preparation of the complexes of interest is essential for thepurpose of determining important properties such as solubilityand thermal stability under realistic conditions. The commonway of preparing complexed ions usually involves dissolutionof a metal salt and chelating agent at low pH followed byaddition of base (caustic) to increase the pH to 7 or greater.Isolation is difficult because of water removal and saltcontamination.A simple method to produce pure, isolable complexes forstudy has been found. The procedure involves using IcelandicSpar crystals, an aqueous medium, and a chelating agent. Forexample, excess large single crystals of Icelandic Spar areimmersed in a measured quantity of water. The chelating agentis then added in small aliquots. Carbon dioxide will slowlyevolve as the active protons of the chelating agent spend onthe carbonate. After gas evolution ceases, the procedure isrepeated. At some point the solubility product of the calciumcomplex will be exceeded and precipitation will occur. Theprecipitated solid can then be filtered and excess Spar crystalsremoved. Washing the solid with alcohol and acetone and thenair drying (ca. 100°C) completes the workup. This techniqueproduces very pure complexes as well as allowing thesolubility of the complexes to be calculated. Also, thisprocedure readily accommodates other fluids such as 20%calcium chloride water to simulate spent 15% HCl.The calcium complex of NTA was prepared using theabove method. Analysis of the solid showed the molecularformula to be: Ca(HNTA) ã 2H 2 O. This material showed asolubility of 240 and 72.5 lbs./Mgal in water and 20% CaCl 2 respectively. Next the solubility of the complex wasdetermined at elevated temperature. A saturated solution of thecomplex in 20% CaCl 2  was heated in a pressure cell equippedwith a glass observation port. When the temperature reached204°F, massive precipitation occurred. The amount of precipitate was not quantitatively determined. A qualitativeestimate would be that most of the complex was released fromsolution at this temperature.The EDTA analog was prepared using a similar procedure.This complex showed excellent solubility (ca. 500 lbs. / Mgal), far exceeding the solubility of free EDTA that willdissolve in the treatment fluid. This salt also lost solubility atelevated temperastures (240°F).Fe(II) complexes of NTA and EDTA were prepared byspiking the initial 20% CaCl 2  /Icelandic Spar reaction mixturewith a known quantity of Fe(II) prepared from anhydrousFeCl 2  and deoxygenated 20% CaCl 2  water. In the case of theNTA, only when the quantity of NTA added to the systemexceeded BOTH the amount of Fe(II) and Ca ion did anyprecipitation occur. The precipitate was identified as thecalcium complex. Clearly, the Fe(II) complex formedpreferentially. Surprisingly, the EDTA complex showed lesssolubility than the corresponding calcium complex asdetermined by the precipitation point. Iron depletion studiesshowed the onset of complex precipitation occurred when 68lbs./Mgal of EDTA had been consumed. Therefore,Fe(II)EDTA solubility is quite low. 68 lbs./Mgal should beconsidered to be the maximum concentration of EDTA instimulation designs for acidizing. As predicted, thesesolutions were extremely oxygen sensitive, instantly turningblood red with introduction of air. Thermal Stability of EDTA complexes – can they survive? The thermal stability of ferrous EDTA in 20% CaCl 2  waterwas explored. A cloudy solution of the complex (slightlygreen due to oxidation) was heated under 500-psi nitrogenpressure. The test solution remained unchanged until thetemperature had reached 250°F. At this point, the solutionbecame both clear and colorless. The latter indicates thatreduction of trace quantities of Fe(III) had occurred. After aperiod of 6 hours, the reaction vessel was cooled to ambienttemperature and slowly depressurized. Vigorous gas evolutionoccurred during the latter stages of depressurization. Controlexperiments using nitrogen and CO 2 atmospheres in theabsence of the iron complex did not show this behavior. Thisresult highly suggests that the complex underwent thermaldecomposition, generating carbon dioxide in solution.A similar experiment was carried out using the calciumsalt of EDTA. After two hours at 250°F, the vessel was cooledand depressurized. Again, vigorous gas evolution occurredduring the latter stages of the depressurization procedure.Thus, both iron and calcium EDTA complexes appear todecompose under these conditions. This result puts an upperlimit to the successful use of this chelating agent somewherebelow 240°F.Therefore, solubilities and thermal stabilities for thecomplexes are not encouraging with respect to staying insolution or successfully stopping hydolysis reactions (hence,precipitation). Reactions of Fe(II) Complexes with H 2 S at pH3 The above preparation now allows a definitive determinationto be made regarding the complex stability in the presence of hydrogen sulfide.Solutions of Fe(II) NTA and Fe(II)EDTA in 20% calciumchloride water were prepared as described above. The pH of the solutions was adjusted to a value of 3 using dilute HCl anda pH meter. Marble chips were added followed by addition of anhydrous hydrogen sulfide via a gas dispersion frit. Within  4M. BREZINSKISPE 54721 10 minutes, both samples had formed copious quantities of black precipitate, which was analyzed as FeS. A white film of organic material was found floating on the top of the samples.This material was not analyzed but surely must be the fullyprotonated chelating agent formed by a proton transfer. I.e.,Fe(HNTA) + H 2 S →  FeS + H 3 NTA(1)These crucial experiments show that neither chelating agentwill be effective in preventing FeS deposition under theseconditions (spent acid). The NTA Precipitation Cycle –an example The following example serves to show how the experiments inthis document may come together and allow a system to bebetter understood.NTA is frequently used for iron control in sour wellenvironments. This particular chelating agent showstremendous solubility in strong acid fluids (>300 lbs./Mgal)Also, the ferrous complex is more soluble than the calciumcomplex. These properties appear to give the material anadvantage over EDTA. However, evidence presented in thispaper shows NTA (or EDTA for that matter) cannot‘encapsulate’ enough coordination sites on Fe(II) at fluid pHvalues below 2 to stop FeS formation. Also shown has beenthe loss of solubility of the calcium complex at ca. 200°F. Theincreased solubility of NTA over EDTA is explained byprotonation forming the conjugate acid of the compound. Thisacid has a pKa of <1 and will deprotonate when the acidstrength is reduced to a range of about 2% to 0.5% residualHCl. At this point, free NTA only possesses meager solubility(20 lbs. /Mgal at 65°F, 60 lbs./Mgal at 160°F, and 150lbs./Mgal at 200°F so most of the material will precipitatefrom the fluid. Next, NTA will spend on the formationforming a calcium salt. If the temperature is greater than about200°F, this complex will again form a precipitate. If on theother hand there is Fe(II) in the system, the Fe(II)NTAcomplex will form. However, the iron complex will react withH 2 S and form FeS and free NTA, which will return to thecycle. The end result will be (for a well over 200°F): all FeSwill precipitate and all NTA will precipitate as a calciumcomplex. Experimental Precautions Aqueous Fe(II) is exceedingly oxygen sensitive and must behandled under rigorously anaerobic conditions. Schlenkware,vacuum lines, and either a nitrogen or argon atmosphere areessential to obtain meaningful results. FeCl 2 ã 4H 2 O is moreoften than not, significantly contaminated with Fe(III). Thiscontamination is readily observed upon inspection of thecrystalline solid or aqueous solution of the salt. The partiallyoxidized condition will result in a green tint to the solid orsolution. Pure aqueous Fe(II) is colorless – even atconcentrations of >50,000 PPM. Also, the ion becomes astronger reducing agent upon chelation. This is the result of pushing more electrons into an already oxidatively unstablemetal center. This causes Fe(II) complexes to be even moreunstable than the hydrated ion. Hence, bubbling air through asolution of Fe(II) EDTA causes an immediate color changefrom colorless to deep red. Indicating formation of Fe(III)EDTA. Conclusions 1.   The FeS damage mechanism allows FeS precipitation tooccur at significantly lower pH values than predictedusing simple equilibrium calculations. Consideration of the formation as an active participant in all downholechemistry may give better insight into actual fluidbehavior in well stimulation treatments.2.   Chelating agents appear to be unable to preventprecipitation of FeS under realistic downhole conditions.3.   Calcium complexes of certain chelating agents exhibitreverse solubility properties with temperature and mayprecipitate at relatively low temperatures.4.   The tested complexes of chelating agents decompose attemperatures as low as 250°F setting a low uppertemperature limit for use.5.   From these results, it appears that the only courseavailable to prevent FeS formation is to remove the otherproduct from the fluid –namely H 2 S with the use of anefficient hydrogen sulfide scavenger. References 1.   Crowe, C.W., Maddin, C.M.: U.S. Patent No. 4,574,050 (1987).2.   Dill, W.R., Walker, M.L., Ford, W.G.F.: U.S. Patent No.4,679,631 (1987).3.   Crowe, C.W.: U.S. Patent No. 4,633,949 (1987).4.   Dill, W.R., Walker, M.L.: U.S. Patent No. 4,888,121 (1989).5.   Handbook of Analytical Chemistry, Meites, L. ed., 1 st  ed., NewYork, McGraw-Hill, 1963.6.   Cotton, F.A., Wilkinson, G.: “Advanced Inorganic Chemistry”,3 rd  ed., New York, Interscience, 1972. 90   Schwarzenbach, G. and Biedermann, W.:  Helv. Chim. Acta . 31 ,331(1948). Acknowledgement The author would like to Halliburton management and thank Marten Buijse of Halliburton European Research Centre forhis help in preparing this paper. TABLE 1   L   * 4 5   K a Values vs FeS PrecipitationPoint TGH 0 5  TGH 07  TGH 0 8 -C 3  ; Z 32 /: 8 Z 32 /: 3 Z 32 /9 -C 4  3 Z 32 /34 3 Z 32 /37 3 Z 32 /39 R*   XCNWG ,< 30; 504 5098 , ECNEWNCVGF   XCNWG   HQT   RTGEKRKVCVKQP   QH   (G5
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