Matrix Acidizing Using Hcl and Other Aicds

SPE 116601 Matrix Acidizing of Carbonate Reservoirs Using Organic Acids and Mixture of HCl and Organic Acids F.F. Chang, SPE, Schlumberger; H.A. Nasr-El-Din, SPE, Texas A&M University; and T. Lindvig and X.W. Qiu, Schlumberger Copyright 2008, Society of Petroleum Engineers This paper was prepared for presentation at the 2008 SPE Annual Technical Conference and Exhibition held in Denver, Colorado, USA, 21–24 September 2008. This paper was selected for presentation by an SPE program committee following review of information contained in an abstract submitted by the author(s). Contents of the paper have not been reviewed by the Society of Petroleum Engineers and are subject to correction by the author(s). The material does not necessarily reflect any position of the Society of Petroleum Engineers, its officers, or members. Electronic reproduction, distribution, or storage of any part of this paper without the written consent of the Society of Petroleum Engineers is prohibited. Permission to reproduce in print is restricted to an abstract of not more than 300 words; illustrations may not be copied. The abstract must contain conspicuous acknowledgment of SPE copyright.

Abstract Hydrochloric acid is the most commonly used acid for carbonate acidizing due to its low cost and high dissolving power. However, there are two major drawbacks associated with using concentrated HCl solutions in deep wells. The first is its high reaction rate with carbonate rocks, which limits acid penetration in the formation. The second is its corrosivity to well tubulars. Hence organic acids become viable material for matrix acidizing to alleviate these two problems. Though organic acids provide the benefit of retardation and low corrosivity, their low dissolving capacity may still limit the wormhole penetration leading to insufficient stimulation of the formation. Therefore, opportunity exists to mix HCl with an organic acid to achieve productivity enhancement by optimizing the wormhole penetration and profile. Organic acids that are utilized in stimulating carbonate formations include formic, acetic, and more recently, citric and lactic. Selecting a suitable organic acid for a specific acidizing treatment is more difficult due to complex thermodynamic equilibrium and reaction kinetics. The reactions between organic acids and carbonate are less understood than those of HCl with carbonate rocks. Organic acid/carbonate systems are complicated because of the presence of CO2, organic ligands, and potential precipitation of the reaction products; the organic salts of calcium and magnesium. Therefore, more testing and modeling are needed to better understand these reactions. This paper discusses the required information to properly design an organic acid or HCl plus organic acid treatment. In additional to reaction kinetics, data such as carbonate dissolving capacity at reservoir temperature and pressure, solubility of reaction products, and the effect of HCl to organic acid ratio are needed to better design field treatments. Recommendations are given on what and how laboratory evaluation should be carried out to obtain this information. Introduction Oil and gas companies are developing carbonate reservoirs of deeper and deeper depths in order to meet the demand of increasing worldwide energy consumption. Enhancing productivity from these reservoirs poses a challenge in stimulation fluids due to the increase in bottom hole temperature. The rapid reaction rate between HCl and carbonate limits the penetration of HCl into the formation, especially at low pumping rates. The reaction of HCl often needs to be retarded by gelling,1 emulsifying,2 or adding viscoelastic surfactants.3 In addition to the high reaction rate, HCl is very corrosive to well tubulars. Expensive corrosion inhibitors can protect the tubulars at high temperatures only for a short period of time. These drawbacks make organic acids, such as formic and acetic, potentially attractive for stimulating high temperature wells. Organic acids have been used in well stimulation because of their low corrosivity4 and lower reaction rate with the rock. However, they have the following limitations: (1) they cannot be used at high acid concentrations. This is because of the limited solubility of their calcium salts. For instance, acetic and formic acids are typically used at concentrations less than 13 and 9 wt%, respectively to avoid precipitation of calcium acetate and calcium formate,5 (2) Organic acids have a low dissociation constant. They normally do not react to their full acid capacity because of the release of CO2 from carbonate dissolution, (3) the degree of hydrogen ion generation decreases with increasing temperature,6,7 and (4) the cost of organic acid is significantly higher than that of HCl for equivalent mass of rock dissolved. Mixing HCl and an organic acid can provide several benefits than using either acid alone.8-10 Organic acids, such as formic acid, can be used as a corrosion inhibitor intensifier for high temperature applications. More interestingly, organic acids can be used to enhance acid penetration in the formation. When an organic acid is mixed with HCl, it does not

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dissociate to generate hydrogen ions due to its low dissociation constant. Therefore, the organic acid is preserved until HCl is nearly spent. As the hydrogen ions from HCl are depleted, the carboxylic groups start to dissociate, resulting in further carbonate rock dissolution from the tip of the acid front and hence increasing acid penetration into the formation. However, there have been claims that when dissolving carbonate with an HCl-organic acid system, the conversion of organic acid is further reduced. For example, in a modeling study by Buijse et al.11 it was illustrated that in a 15 wt% HCl/10 wt% acetic acid system, only 24% of the 10 wt% acetic acid was spent. This result was believed to be caused by the large amounts of CO2 generated from the reaction of HCl with carbonates, and the reversibility of organic (HA) acid reacting with carbonates, Eq. 1.7 CaCO3 + 2 HA

CaA2+ H2O + CO2

(1)

In Eq. 1, increasing CO2 concentration drives the reaction toward the left, therefore less CaCO3 is dissolved. There is very limited experimental data to prove the validity of the models described by William et al.,7 and Buijse et al.11 Carbon dioxide also affects the reactions of organic acids via the formation of carbonic acid. Under reservoir pressure, the CO2 in most cases is dissolved in the aqueous phase and generates carbonic acid, Eqs. 2-4. Carbonic acid buffers the pH value to nearly 4.5: H2O + CO2

H2CO3

(2)

H2CO3

H+ + HCO3-

(3)

HCO3-

H+ + CO32-

(4)

Figs. 1(a) and 1(b) show the formic acid and acetic acid dissociation as a function of pH. It can be seen that at pH 4.5, formic and acetic acids are not fully dissociated. Acid-carbonate reaction kinetics has been studied by several research groups.12-16 The reaction rate and mass transfer coefficients between various acids and calcite and dolomite are well documented. However, the equilibrium chemistry of the acid-carbonate rock system is thought to be simple, therefore it has not been emphasized. For HCl, the acid spending is to completion. Therefore, its total rock dissolving capacity is straight forward. It can be calculated stoichiometrically by rewriting Eq. 1 as: CaCO3 + 2 HCl

CaCl2+ H2O + CO2

(5)

The acid reaction with calcite can be written as: CaCO3 + 2 H+

Ca2+ + H2O + CO2

(6)

Unlike HCl, organic acids and limestone reaction reaches equilibrium before the acid is completely spent. The total acid dissolving capacity requires an accurate account of the thermodynamic parameters of a system of reactions including CO2 and carbonic acid equilibrium. The system can be further complicated by the ion association of calcium ions and the anions of the organic acid, and potential precipitation of the calcium salt with the anion, A-, of the organic acid, Eqs. 7 and 8: Ca2+ + ACa2+ + 2A-

CaA+ CaA2

(7) (8)

The equilibrium constants of these reactions, especially at reservoir temperature and pressure, are not available for all reactions in the system. Therefore, it is difficult to properly model spending of organic acids. A model based on Gibbs free energy is required to study the thermodynamic equilibrium among the species in solution. The objective of this paper is to investigate the equilibrium of the system of reactions involved in the limestone dissolution by organic acids and HCl-organic acid solutions. The effect of reservoir pressure on CO2 evolved due to calcite dissolution by acid is investigated. Commercial chemical analysis software is used to establish the equilibrium of ionic species in the reactions under reservoir temperature and pressure. The software has been applied for chemical analyses in various industries. It has extensive database of thermodynamic parameters and been widely validated for its accuracy and functionality against a large number of experimental data sets. This approach can help select proper experimental parameters and design acid formulation for carbonate acidizing treatments. It complements the kinetic studies of acid-rock reaction rate.

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Equilibrium Modeling In an acid-carbonate system, the electrolyte concentrations are at high levels, and the solution cannot be treated as an idea one. An ideal solution means that the concentration approaches zero or the mole fraction approaches unity.17 Therefore, when calculating thermodynamic equilibrium, the activity of each species has to be used. The chemical analysis model utilized in this study uses the Helgeson model18 to represent the variation of species standard Gibbs free energies with temperature and pressure and a Pitzer-like activity coefficient model19 to account for deviations from thermodynamic ideality. The advantage of a Gibbs energy free energy equilibrium framework is that reaction constants are not required. The software considers all the relevant species and manifests their equilibrium according to their free energy functions. Four acid formulations were tested in the simulation to investigate their interactions with calcite rock. These acid formulations were: (1) 10 wt% acetic acid, (2) 15 wt% HCl + 10 wt% acetic acid, (3) 9 wt% formic acid, and (4) 15 wt% HCl + 9 wt% formic acid. The simulation considered a reaction vessel that contained 100 kg of the acid, calcite was sequentially added into the vessel at 1 kg increment to react with the acid at 150oF and 1,000 psi. The distribution of each chemical species was calculated at each step based on chemical equilibrium. Fig. 2 shows the case of 10 wt% acetic acid reacting with calcite. As CH3COOH reacted with CaCO3, the CO2 concentration in the aqueous increased and Ca2+ was released. After 6.2 kg of CaCO3 was dissolved, the system reached equilibrium. At the end of the reaction, 74% of the 10 wt% acetic acid was spent. This number was high compared to the 54% obtained from the model predictions by Buijse et al.11 and the 50% obtained from the experimental results of Chatelain et al.6 At 1,000 psi system pressure and 150oF temperature, the initial pH of 10 wt% acetic acid is 2.32. The pH at the end of the reaction was 4.48 due to dissolved CO2 in the system. It was notable that in the system the free Ca2+ concentration was significantly lower than the stoichiometry according to reaction (6). And the concentrations of acetate ions (CH3COO-), calcium acetate (Ca(CH3COO)2), and calcium monoacetate (Ca(CH3COO)+) were in the same order of magnitude and not negligible. This indicated that Ca2+ ions and the acetate ions were in ionic association and the reactions (7) and (8) could not be ignored.20,21 This was the reason that the model predicted higher dissociation of acetic acid than previous publication by Buijse et al.11 and Chatelain et al.6 However, Bombardieri and Martin22 showed that at 150oF and 300 psi 65% of the 10 wt% acetic acid reacted with calcium carbonate in 7 hours while the dissolution was still progressing. Though the test by Bombardieri and Martin22 was at lower pressure level, data showed that at 300 psi, the CO2 dissolved in aqueous solution was sufficient to buffer the pH to 4.3, which was not much different from the level at 1,000 psi. Hence the reduction of acid spending due to CO2 should be limited. On the other hand, increasing organic acid spending due to Ca2+ may be significant. When reactions (7) and (8) were dominant, the organic acid dissociation could be further driven to the right to increase the amount of hydrogen ion generation as follows: HA

H+ + A-

(9)

When the free organic anion (A-) was reduced due to its association with Ca2+ by reactions (7) and (8), the equilibrium of reaction (9) shifted toward the right and more H+ was produced. This was demonstrated by a simple experiment of pH measurement. The pH of a 100 cm3 solution of 10 wt% acetic acid was first measured at ambient condition. The pH of the solution was continuously monitored while CaCl2 was slowly added to the solution. Fig. 3 shows the pH of the 10 wt% acetic acid solution decreased as more CaCl2 was added to the solution, indicating the acetic acid further dissociated and more H+ was generated. Buijse’s model11 did not account for the effect of ionic association between Ca2+ and A- therefore it may be underestimating spending of the organic acid. The data presented by Chatelain et al.6 was obtained from experiments. It is difficult to argue against the validity of their results without conducting additional experiments. Nonetheless, it is worth noting that the experimental procedure used by Chatelain et al.6 to measure acid concentration was by first flushing the acid through the carbonate rock, then reducing the pressure of the acid effluent to allow CO2 to escape. The effluent was then titrated with NaOH solutions. There was still dissolved CO2 in the solution so that carbonic acid contributed to the residual acidity. No material balance was conducted to compare the acid spending with rock weight loss, nor pH measurement to demonstrate the complete vaporization of CO2. Therefore, it is the belief of the authors that more experimental work is required to properly define the capacity of organic acids under downhole conditions. The next simulation was related to calcite dissolution by a mixture of 15 wt% HCl and 10 wt% acetic acid. It is well known that when an organic acid is mixed with HCl, the organic acid will not dissociate due to high hydrogen ion concentration provided by HCl. It is when HCl is nearly completely spent that the organic acid starts to dissociate. This is beneficial to the acid penetration depth because the strength of the organic acid is preserved. It behaves as if the fresh organic acid were injected at the tip of HCl reaction front inside the formation. Another benefit of the mixed acid system is that the reaction rate is reduced.23 However, there have been claims7,11 that the organic acid spending percentage will be further reduced when it is mixed with HCl due to large amounts of CO2 generated by the HCl-CaCO3 reaction. Fig. 4 shows the chemical species distribution as the acid mixture dissolved calcite. The concentration of CH3COOH remained relatively constant when fresh HCl was reacting with the rock. After the 15 wt% HCl was fully spent dissolving about 20 kg of rock, the acetic acid started to react and the acid concentration declined sharply. Totally 28.1 kg of CaCO3 was dissolved by the 100 kg of 15 wt% HCl-10 wt% acetic acid mixture. The chemical analysis shown in Fig. 4 indicates that the effect of CO2 was less pronounced, the pH at the end of reaction was 4.29, which was not much lower than the 10 wt% acetic acid case. However, the Ca2+ plays a key role in the reaction. As the Ca2+ increased due to HCl dissolving the

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calcite, a large amount of calcium monoacetate was formed at the expense of acetate ions (CH3COO-) and calcium acetate (Ca(CH3COO)2). The CH3COO- and Ca(CH3COO)2 concentrations were reduced significantly as shown in Fig. 4. This means reaction (7) was predominant due to high Ca2+ concentration, which droves the acetic acid dissociation toward the right. At equilibrium, the acetic acid spending percentage reached 90%. There has not been published experimental data concerning the equilibrium during carbonate rock dissolution by HCl and organic mixtures. Research is underway to validate the simulation and to gain further understanding in this area. Fig. 5 shows the case of 9 wt% formic acid reacting with calcite. Generally, the trend was similar to that of acetic acid (Fig. 2). At 1,000 psi system pressure and 150oF, the initial pH of 9 wt% formic acid was 1.74. As the acid reacted with calcite, the pH and dissolved CO2 concentration increased. The CO2 in solution reached its solubility limit of 2.86% at 1,000 psi when pH increased to 3.45 after 7 kg of calcite was dissolved. Formic acid continued to dissociate until 97% of the formic acid was spent and a total of 9.7 kg of calcite was dissolved when the system reached equilibrium. An additional 1% of CO2 was generated in the form of gas. The pH at the end of the reaction was 4.48. The 97% spending was again higher than 92% from the experimental result by Chatelain et al.6 and 85% from the model calculation by Buijse et al.11 It was noticed that there was no free Ca2+ in the solution. All the Ca2+ were present in the form of calcium monoformate Ca(HCOO)+ and calcium formate Ca(HCOO)2. There was no precipitation as the two calcium salts were in the aqueous phase. The concentrations of free calcium ions (Ca2+), formate ions (HCOO-), calcium formate (Ca(HCOO)2), and calcium monoformate (Ca(HCOO)+) also explained that reactions (7) and (8) were prevalent. The final case studied was calcite dissolution by 15 wt% HCl and 9 wt% formic acid mixture. Fig. 6 shows the chemical species distribution as the acid mixture dissolved calcite. Similar to the HCl and acetic acid mixture, the concentration of HCOOH remained relatively constant until HCl reacted to near completion when 20 kg of calcite was dissolved. The formic acid concentration then started to decline sharply as it dissociated into formate and hydrogen ions to react with the rock. At the end of the reaction, the total CaCO3 dissolved was 30.4 kg by the 100 kg of 15 wt% HCl-9 wt% formic acid mixture. The pH at the end of reaction was 4.33, slightly lower than that by 9 wt% formic acid alone. The large amount of free Ca2+ ions from HCl-calcite reaction was consumed by formate ions to form calcium monoformate by ionic association. Hence, the free formate ions (HCOO-) and calcium formate (Ca(HCOO)2) was depleted from the system. Reaction (7) dominated and it drove the formic acid dissociation toward the right. At equilibrium, the formic acid spending reached 99%. It has been known that organic acid spending is significantly impacted by the CO2 in solution. A common belief6,11 in the oil and gas industry is that at 1,000 psi the CO2 generated by the acid-carbonate reaction will be dissolved in the solution. Therefore 1,000 psi has been the standard system pressure used in experiments related to carbonate dissolution. A simulation was conducted to illustrate the effect of system pressure on organic acid spending and CO2 solubility. The acid formulation for the simulation run was 15 wt% HCl and 9 wt% formic acid mixture. The acid reacted with calcite at 150oF and pressure ranging from 1 to 6,000 psi. Fig. 7 shows the pH of the system at equilibrium and the total CO2 in aqueous and vapor phases. It can be seen that a large fraction of the CO2 generated by the acid-calcite reaction was in the vapor phase below 3,000 psi. The amount of dissolved CO2 increased dramatically when the pressure was above 3,000 psi until all the vapor phase was dissolved in the aqueous solution when the pressure was above 5,200 psi. This demonstrated that 1,000 psi was not sufficient to keep all the produced CO2 in solution. This has been pointed out by a few researchers in the past.24,25 Interestingly however, the equilibrium pH remained relatively constant as long as the pressure was above 200 psi and below 3,000 psi, especially between 1,000 and 3,000 psi. This means the acid spending due to the CO2 effect was small. Therefore, from acid dissolving power aspect, application of 1,000 psi system pressure should produce reasonable results. However, when conducting kinetics experiments, CO2 vapor significantly affects the reaction rate due to the forced convection effect to promote mixing and mass transfer in the reactive boundary layer. Therefore, the reaction rate of acid on carbonates can be overestimated if insufficient system pressure is applied.24 Conclusions The chemical analysis software based on the Gibbs free energy is a very useful tool to gain insight into organic acid reactions with carbonate rocks. The objective of this paper is to highlight that further research, especially in experimentation, is needed to understand the role of organic acids in acidizing deep wells. Based on the work discussed in this paper, the following conclusions can be drawn: 1. 2. 3.

When conducting carbonate acidizing experiments, 1,000 psi system pressure may be sufficient for testing acid dissolving capacity, but a much higher pressure is required for kinetics studies. The ionic association between Ca2+ and organic anions can significantly increase organic acid dissociation. Mixing HCl with an organic acid (acetic or formic) provides benefits not only in reducing corrosivity, but also in acid penetration depth due to delayed dissociation of organic acid in the presence of live HCl.

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References 1. Nasr-El-Din H.A., Al-Mohammad, A.M., Al-Aamri, A.M., and Al-Fuwaires, O.: “Reaction Kinetics of Gelled Acids with Calcite,” paper SPE 103979, accepted for publication in SPEPO, 2008. 2. Navarrete, R.C., Holms, B.A., McConnell, S.B. and Linton, D.E.: “Emulsified Acid Enhances Well Production in HighTemperature Carbonate Formations,” paper SPE 50612 presented at the 1998 SPE European Petroleum Conference, The Hague, The Netherlands, October 20 - 22. 3. Al-Mohammad, A.M., Nasr-El-Din H.A., Al-Aamri, A.M., and Al-Fuwaires, O.: “Reaction of Calcite with SurfactantBased Acids,” paper SPE 102838 presented at the 2006 SPE Annual Technical Conference and Exhibition, San Antonio, TX, 24-27 September. 4. Harris, F.N.: “Applications of Acetic Acid to Well Completions, Stimulations and Reconditioning,” JPT 13(7) (1961) 637-639. 5. Robert, J.A. and Crowe, C.W.: “Carbonate Acidizing Design,” Reservoir Stimulation, Economides, M.J. and Nolte, K.G. 3rd Edition, John Wiley & Sons Inc. (2000) 17-11. 6. Chatelain, J.C. Silberberg, I.H., and Schechter, R.S.: “Thermodynamic Limitations in Organic-Acid/Carbonate Systems,” SPEJ 16(4) (1976) 189-195. 7. Williams, B.B., Gidley, J.L., and Schechter, R.S.: “Acidizing Fundamentals,” SPE Monograph Volume 6, 1979. 8. Dill, W.R. and Keeney, B.R.: “Optimizing HCl-Formic Acid Mixtures for High Temperature Stimulation,” paper SPE 7567 presented at the 1978 Annual Fall Technical Conference and Exhibition of the SPE-AIME, Houston, TX, Oct. 1-3. 9. Taylor, K.C., Al-Katheeri, M.I., and Nasr-El-Din, H.A.: “Development and Field Application of a New Measurement Technique for Organic Acid Additives in Stimulation Fluids,” SPEJ 10(2) (2005) 152-160. 10. Katheeri, M.I., Nasr-El-Din, H.A., Taylor, K.C., and Grainees, A.H.: “Determination and Fate of Formic Acid in High Temperature Acid Stimulation Fluids,” paper SPE 73749 presented at the 2002 SPE International Symposium on Formation Damage Control, Lafayette, LA, Feb. 20-21. 11. Buijse, M., deBoer, P., Breukel, B., Klos, M., and Burgos, G.: “Organic Acids in Carbonate Acidizing,” SPEPF 19(3) (2004) 128-134. 12. Fredd, C.N. and Fogler, H.S.: “The Kinetics of Calcite Dissolution in Acetic Acid Solution,” Chem. Eng. Sci. 53(22) (1998) 3863-3874. 13. de Rozieres, J., Chang, F.F. and Sullivan, R.B.: “Measuring Diffusion Coefficients in Acid Fracturing Fluids and Their Application to Gelled and Emulsified Acids,” paper SPE 28552 presented at the 1994 Annual SPE Conference, New Orleans, LA, Sep. 25-28. 14. Nierode, D.E. and Williams, B.B.: “Characteristics of Acid Reaction in Limestone Formations,” SPEJ 11 (1971) 406418. 15. Roberts, L.D. and Guin, J.A.: “The Effect of Surface Kinetics in Fracture Acidizing,” SPEJ 14 (1974) 385-395. 16. Nasr-El-Din, H.A., Al-Mohammad, A.M., Al-Aamri, A.D., Al-Fahad, M.A. and Chang, F.F.: “Quantitative Analysis of Reaction Rate Retardation in Surfactant-Based Acids,” paper SPE 107451, accepted for publication in SPEPO, 2008. 17. Laitinen, H.A. and Harris, W.E.: “Chemical Analysis,” 2nd edition, 1975, McGraw-Hill Book Company. 18. Shock, E.L., Helgeson, H.C., and Sverjensky, D.A.: “Calculation of the Thermodynamic Properties of Aqueous Species at High Pressures and Temperatures: Standard Partial Molal Properties of Inorganic Neutral Species,” Geochim. Cosmochim. Acta 53 (1989) 2157-2183. 19. Pitzer, K.S.: “Thermodynamics of Electrolytes. I. Theoretical Basis and General Equation,” J. Phys. Chem. 77 (1973) 268-277. 20. Loos, D., Pasel, C., Luckas, M., Schmidt, K.G., and Herbell, J.D.: “Experimental Investigation and Modelling of the Solubility of Calcite and Gypsum in Aqueous Systems at Higher Ionic Strength,” Fluid Phase Equilibrium 219 (2004) 219-229. 21. Nancollas, G.H.: “Thermodynamics of Ion Association, Part II – Alkaline-Earth Acetates and Formates,” J. Chem. Soc. (1956) 744. 22. Bombardieri, C.C. and Martin, T.H.: “Acid Treating Process,” US Patent 3,251,415, May 17, 1966. 23. Dill, W.R.: “Reaction Time of Hydrochloric-Acetic Acid Solutions on Limestone,” paper SPE 00211 presented at the 16th Southwest Regional Meeting of the American Chemical Society, Oklahoma City, OK, December 1-3, 1960. 24. Mumallah, N.A.: “Factors Influencing the Reaction Rate of Hydrochloric Acid and Carbonate Rock,” paper SPE 21036 presented at the 1991 SPE International Symposium on Oilfield Chemistry, Anaheim, CA, Feb. 20-22. 25. Crowe, C.W., McGowan, G.R., and Baranet, S.E.: “Investigation of Retarded Acids Provides Better Understanding of Their Effectiveness and Potential Benefits,” SPEPE 5(2) (1990) 166-170.

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1.00 HCOOH

HCOO-

Fraction of Species

0.80

0.60

0.40

0.20

0.00 0.0

1.0

2.0

3.0

4.0

5.0

6.0

7.0

pH Fig. 1a: The fraction of species from the dissociation of formic acid at ambient conditions. 1.00 CH3COO-

CH3COOH

Fraction of Species

0.80

0.60

0.40

0.20

0.00 0.0

1.0

2.0

3.0

4.0

5.0

6.0

pH Fig. 1b: The fraction of species from the dissociation of acetic acid at ambient conditions.

7.0

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7

1.8

5

1.6 4

1.4 CH3COOH Ca2+ CH3COOCa(CH3COO)+

3

Ca(CH3COO)2

1

CO2

pH

Concentration (mol/l)

1.2

pH

0.8 2 0.6

0.4

1

0.2

0

0 0

1

2

3

4

5

6

7

8

9

10

Calcite Dissolved by 100 kg of Acid (kg)

Fig. 2: Equilibrium concentrations of main species when 10 wt% acetic acid reacted with calcite at 150oF and 1,000 psi.

2.2 2

pH

1.8 1.6 1.4 1.2 1 0

5

10

15

20

25

30

Weight of CaCl2 Added, g

Fig. 3: Effect of calcium ion concentration on the pH of 100 cm3 of 10 wt% acetic acid at ambient conditions.

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SPE 116601

1.8

5

1.6 4

CH3COOH Ca2+ CH3COO-

1.2

3

Ca(HCOO)+ Ca(CH3COO)2

1

CO2

pH

Concentration (mol/l)

1.4

2

pH

0.8 0.6

1

0.4 0 0.2 0

-1 0

5

10

15

20

25

30

35

Calcite Dissolved by 100 kg of Acid (kg) Fig. 4: Equilibrium concentrations of main species when 15 wt% HCl and 10 wt% acetic acid reacted with calcite at 150oF and 1,000 psi. 5

2 1.8

4

1.4

HCOOH Ca2+

1.2

3

HCOOCa(HCOO)+ Ca(HCOO)2 CO2 pH

1 0.8

pH

Concentration (mol/l)

1.6

2

0.6 0.4

1

0.2 0

0 0

2

4

6

8

10

12

Calcite Dissolved by 100 kg of Acid (kg) Fig. 5: Equilibrium concentrations of main species when 9 wt% formic acid reacted with calcite at 150oF and 1,000 psi.

SPE 116601

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5

2 1.8 HCOOH Ca2+ HCOOCa(HCOO)+ Ca(HCOO)2 CO2 pH

1.4 1.2

4

3

2

1

pH

Concentration (mol/l)

1.6

0.8 1

0.6 0.4

0

0.2 0

-1

0

5

10

15

20

25

30

35

Calcite Dissolved by 100 kg of Acid (kg)

Fig. 6: Equilibrium concentrations of main species when 15 wt% HCl and 9 wt% formic acid reacted with calcite at 150oF and 1,000 psi. 5

350

300

250 3

200 pH CO2 (aq) CO2 vapor

2

150

CO2 vapor (mole)

CO2 (aq) (mol/l) and pH

4

100 1 50

0 0

1000

2000

3000

4000

5000

0 6000

Pressure (psi) Fig. 7: pH value and CO2 phase distribution as a function of pressure when 15 wt% HCl + 9 wt% formic acid reacted with calcite at 150oF.