Chelating Agents in the Oilfield

Abstract

The focus in this chapter will be on the chemical subset of chelating agents commonly used in the petroleum industry. It will start by defining the functionality known as chelation. It will then share the possible applications of these chelating agents. Such applications include utilizing them as straight acidizing fluids, iron control agents, and inorganic scale removers. The chapter will then compare the corrosivity of these molecules towards typical metals used in industry. Finally, it will touch on the environmental aspect by sharing insights on the thermal degradation profile of the used chelating agents and its impact on nature.

1. Introduction

Chelating agents are compounds having at least two groups, usually referred to as ligators, that can transfer electrons to generate coordination bonds with a central metal atom. Due to the broad definition of the term, many compounds containing two or more electron donating groups, such as oxalates, fall into this category. However, in the oil and gas industry, the term “chelating agents” is often used to refer to a particular subclass of chelating agents known as aminopolycarboxylic acids (APCA). As the name suggests, these chelating agents contain one or more amine groups which form multiple metallocyclic rings from a single molecule via the formation of numerous coordinate bonds (Figure 1). In this review, we aim to provide readers with a basic understanding of chelation chemistry and the many applications of chelating agents in the oil and gas industry.

Figure 1.

Chelating agent rings in a Ca-EDTA complex [1].

1.1 Stability constants

The equilibrium constant is an indication to whether the products of a reaction will remain stable in the solution or revert to its reactants. In the case of APCAs, this is typically referred to as the stability constant as it serves as an indication of the stability of the chelated product. The stability of a chelated product improves with a higher stability constant. Eqs. (1) and (2) can be used to determine this constant.

$$\text{M}+\text{Y}\rightleftarrows \text{MY}.$$E1

$$StabilityConstant=\log \left(\frac{\left[MY\right]}{\left[\text{M}\right]\left[\text{Y}\right]}\right).$$E2

Eq. (2) displays the stability constant in relation to the reactant and product concentrations in mol/L where M and Y react to produce the product MY. Additionally, this constant has been shown to rely on numerous factors, including the size of the ring created during chelation, the number of rings established, the basicity of the chelating agent, the type of donor atoms, and the central metal atom [2, 3]. For most chelating agents, a chelate ring size of five atoms offers the most stable structure, and any other chelation ring size would result in a loss in the structure’s stability [4]. Since it is impossible to generalize the extent to which each of these specific characteristics affects the stability constant for all chelating agents, a combination of these factors can result in variations in the stability constants. Understanding the stability constants is an important first step to make when deciding on the appropriate chelating agent to be used in the field; selecting a chelating agent with low affinity towards the type of metal ion to be complexed would result in a failed treatment.

Examples of the chelating agents used in the oil industry include: nitrilo triacetic acid (NTA), hydroxyethylimino diacetic acid (HEIDA), ethylenediamine tetraacetic acid (EDTA), hydroxyethylethylenediamine triacetic acid (HEDTA), diethylenetriamine pentaacetic acid (DTPA), L-glutamic acid N, N-diacetic acid (GLDA), methylglycine diacetic acid (MGDA), 1,4,7,10-tetraazacyclododecane-1,4,7,10-tetraacetic acid (DOTA), and trans-l,2-cyclohexylenediamine tetraacetic acid (CDTA). A table of stability constants for these chelating agents can be found in Table 1.

	Ca^2+	Mg^2+	Fe^2+	Fe^3+	Ba^2+	Sr^2+	Al^3+	Zn^2+	Zr^4+
NTA	6.4	5.4	8.8	15.8	4.8	5.0	11.4	10.6	20.8
HEIDA	4.8	3.5	6.8	11.6	2.8	3.8	7.7	8.4
EDTA	10.7	8.7	14.3	25.7	7.7	8.6	16.1	17.5	29.9
HEDTA	8.4	7.0	12.2	19.8	6.2	6.9	15.6	14.7
DTPA	10.9	9.3	16.5	28.0	8.6	9.7	18.4	18.8	36.9
GLDA	5.9	5.2	8.7	15.2	3.5	4.1	12.2	11.5
MGDA	7.0	5.8	8.1	16.5	4.9	5.2		10.9
DOTA	17.2	11.8	20.2	29.4	12.9	15.0	17.0	20.5
CDTA	13.1	11.1	18.9	30.0	8.6	10.5	19.5	19.3	29.9

Table 1.

Stability constants at 77°F (25°C) [5, 6, 7, 8].

Of these chelating agents, the most used are NTA, HEIDA, EDTA, HEDTA, DTPA, and GLDA [9, 10]. As can be seen in Figure 2, they are made up of between one and three amine groups and multiple carboxylic arms that branch off the amine. Readers may notice that several of these chelating agents have similar structural formulas, such as EDTA and HEDTA. This is due to the addition of a hydroxyl functional group in place of a carboxylate functional group to improve the water solubility of the chelating agent. While this replacement comes at the cost of its stability constant (Table 1), HEDTA does not face any solubility problems while EDTA tends to precipitate at low pH. These APCAs can be described by the general chemical formula HxY, where x depends on how many carboxylate groups are protonated in the molecule. At higher pH levels, the APCA can be deprotonated until all protons are removed. The coordination interactions between the amine/carboxylate groups of deprotonated APCA and the metal ion then serve to chelate the metal ion [14].

Figure 2.

Structural formulas of commonly used aminopolycarboxylic acids in the petroleum industry [11, 12].

1.2 Mechanism of mineral dissolution

One important application of chelating agents in the oilfield is to dissolve formation rock and scale deposits. Therefore, it is necessary to understand the method of mineral dissolution. The rock and scale deposits are found in the wellbore or in the tubulars. These minerals are often covered with oil leaving them oil wet. This necessitates several preparation stages to remove the oily layers and to generate a water wet surface that would allow for the dissolution mechanisms discussed in this section.

The method of chelating agent attack on rocks such as calcite or dolomite, happens through a mechanism called chelating agent assisted dissolution. At a high pH, solution coordination would be the primary dissolution mechanism while surface complexation would be prevalent at lower pH [15, 16, 17]. The pH of the chelating agent can be adjusted by the addition of strong acids and bases such as HCl and NaOH. The optimal pH is in the range of 3–5 for low pH applications and above 10 for high pH applications.

While solution coordination involves the physical dissolution of the metal ion before it is chelated in solution (Figure 3), surface complexation involves a different process as described by Nowack and Sigg [18] and Fredd and Fogler [19].

Figure 3.

Representation of solution coordination mechanism of EDTA chelating agent (more common at high pH) [1].

The surface complexation mechanism can be described in five steps as seen in (Figure 4).

1. Chelating agent diffusion from solution to the surface of the mineral.

2. The initial adsorbed compound is formed.

3. Transforming the complex into one that can pull away from the surface.

4. Complex is released from the mineral surface.

5. Complex diffusion into solution.

Figure 4.

Representation of surface complexation mechanism of EDTA chelating agent (more common at low pH) [1].

Among these five steps, the rate-limiting step can be either step 3 or 5 [18]. Chelating agents are frequently utilized at pH 4 in oilfield acidizing processes. This pH is adjusted by incorporating diluted HCl. It should be noted that despite the presence of H⁺ at this pH, chelating agent assisted dissolution accounts for 90% of dissolution [19] via the previously stated surface complexation process. Surface complexation refers to the creation of a chelating agent-cation complex with cations on the surface. This strengthens the bond between the chelating agent and the ion while weakening the bonds of the crystal structure surrounding the ion. Due to the protonation of oxygen atoms in the presence of H⁺ at this pH, oxide bonds in the lattice are also weakened. This allows the chelating agent to extract the atom, thus making surface complexation the major process of dissolution at low pH.

2. APCAs for mineral acidizing

2.1 Subsurface formations containing carbonate minerals

The primary objective of carbonate acidizing is to increase permeability through the creation of wormholes that bypass the formation damage close to the wellbore (Figure 5). Conventional matrix acidizing involves the use of strong acids such as HCl which result in the consumption of large volumes of acid throughout the process [20]. However, at temperatures above 200°F (93°C), HCl becomes highly corrosive and causes face dissolution instead of forming wormholes. Therefore, chelating agents can be used as an alternative acidizing fluid in these conditions.

Figure 5.

Representation of a carbonate acidizing operation yielding wormholes that bypass the damage zone and improves conductivity of the formation.

As shown earlier, chelating agents do not depend solely on hydrogen ions for their dissolution process, which is different from that of HCl. With chelating agents, chelation and hydrogen ion attack simultaneously at low pH, increasing dissolution [14, 19]. HCl and organic acids (e.g. formic and acetic acid) have a 2:1 stoichiometry, compared to 1:1 for regularly used chelating agents (acid:carbonate) [21]. The dissolution of calcium carbonate at pH between 4 and 5 by chelating agents can be expressed by Eqs. (3)–(5) [22]:

Attack with hydrogen:

$${\text{2H}}^{+}+{\text{CaCO}}_3\to {\text{Ca}}^{2+}+{\text{CO}}_2+{\text{H}}_2\text{O}.$$E3

Calcium chelation:

$${\text{Ca}}^{2+}+{\text{H}}_2\text{Y}\rightleftarrows {\text{CaY}}^{2-}+{\text{2H}}^{+}.$$E4

Mixture of both:

$${\text{H}}_2{\text{Y}}^{2-}+{\text{CaCO}}_3\rightleftarrows {\text{CaY}}^{2-}+{\text{H}}_2\text{O}+{\text{CO}}_2.$$E5

The use of chelating agents to acidize carbonate formations has been extensively studied by many authors. The first to do so were Fredd and Fogler [14], who showed that wormholes could be created by injecting chelating agents at pH ranges of 4–13 and noted the absence of sludge or asphaltene formation throughout the process. They later investigated the kinetics of carbonate dissolution using CDTA, DTPA, and EDTA at pH ranges of 3–12 [19]. In this work, they determined that the pH, type, and concentration of the chelating agent were all factors in the dissolution.

Subsequently, many authors have investigated EDTA as an alternative acidizing fluid. Due to its lower reactivity than HCl [22], EDTA was observed to produce more effective wormholes and could be used at temperatures up to 400°F (204°C) [23, 24].

Due to the poor solubility of EDTA, more soluble APCAs such as HEDTA and HEIDA were investigated by other authors. While not as effective, they were also shown to be able to produce dominant wormholes at temperatures between 250°F (121°C) and 400°F (204°C) [6, 21, 23, 25].

A more recent APCA developed was GLDA which overcomes the solubility issues with EDTA and is also biodegradable. This APCA was tested by Mahmoud et al. and was found to be effective at forming wormholes between 180°F (82°C) and 300°F (149°C) [24]. In addition, LePage et al. [26]. showed that GLDA had greater solubility and calcium dissolution capacity at 300°F (149°C) across a wide pH range.

In addition, the method of preparing chelating agents was also found to impact its performance in acidizing treatments. Chelating agent solutions are typically prepared by mixing the salt of the chelating agent with water and acidifying the resulting solution using conventional acids. Tests were done to examine the impact of using various acids (to get the chelating agent’s pH to a value of 4) on the carbonate acidizing performance [27]. The authors found that 5 wt.% GLDA adjusted with acetic acid dissolved limestone with greater efficacy than the same solution adjusted with HCl.

2.2 Subsurface formations containing sandstone minerals

Unlike carbonate acidizing, acid treatments in sandstone reservoirs are not aimed towards increasing the native matrix permeability, but rather to repair the damage that was done during drilling and completion operations. Clay, carbonates, sulfates, zeolite, feldspar, sand, and oxides are just a few of the minerals found in sandstone reservoirs. While clays are present in both carbonate and sandstone formations, acidizing in sandstone reservoirs can become complicated due to the presence of cementing materials. These cementing materials are often limestone and act to hold individual sand grains together. If they are removed during treatment, sand production can occur, resulting in damage to pipes and equipment due to erosion [28, 29]. Therefore, using a weaker stimulation fluid such as a chelating agent instead of HCl, helps to reduce the possibility of sand production.

Furthermore, when selecting chelating agents for sandstone acidizing, the type of counter-cation must also be considered. Commercial chelating agents are produced in a variety of salts with cations such as Na⁺, K⁺, or NH₄⁺. Ammonium salts of chelating agents should be utilized in hydrofluoric (HF) acid applications to avoid fluorosilicate precipitation. While this is true for common chelating agents, recent research has shown that chelating agents can be designed to overcome the problem of fluorosilicate precipitation with non-ammonium chelating agents [30, 31].

Both laboratory and field treatments have demonstrated the utility of common APCAs for stimulating sandstone. It has been shown that chelating agents can sequester cations such calcium, iron, magnesium, and aluminum, which inhibits various precipitations from occurring during the process of sandstone acidizing [32, 33, 34, 35]. In laboratory tests, sandstone cores were subjected to a variety of common chelating agent treatments at approximately 200°F (93°C) and permeability was found to have increased [36, 37]. In these experiments, EDTA was found to cause formation damage by precipitating in the core and clogging the pores due to its poor solubility. It was also discovered that although chelating agents helped to increase the permeability of the core, the HF main flush decreased the permeability due to secondary formation damage. A new chelating agent treatment combined with HF was evaluated by Mohd Jeffry et al. [38] in sandstone cores. While the formulation of the chelating agent was not shown, the authors observed a 264% increase in permeability after two such treatments were used.

Following the successful laboratory experiments, chelating agents have therefore been applied to field treatments and have been shown to be able to restore permeability to damaged formations [30, 31, 37]. However, authors have described different methods of applying chelating agent treatments. Legemah et al. [39] described a two-step process, using a low volume, high concentration formulation for the first treatment and a high volume, low concentration formulation for the second. Typically, the chelating agent treatment is applied in a single step. Therefore, before field application of a chelating agent stimulation treatment, it is necessary to conduct laboratory scale tests to determine the method of treatment as well as to identify other potential issues that can arise, especially with sandstone reservoirs.

3. Chelation of iron

The chelating properties of chelating agents can be used to tackle another major problem in the oil and gas industry, namely iron. Ferric (Fe³⁺) iron has a reputation for causing many issues during stimulation treatments by stabilizing damaging emulsions and inducing the precipitation of sludge and asphaltene [40, 41]. Iron has also been found to be incompatible with several additives used in the oil industry, including viscoelastic surfactants (VES). In addition, it is well known that iron hydroxides easily precipitate, which damages the formation by clogging the pores and lowering permeability. Iron fluoride precipitation is an additional issue when HF is used in sandstone acidizing. Taylor et al. [42, 43] studied the precipitation of iron hydroxide at room temperature, at spent HCl acid conditions, and in the absence of gases such as H₂S. They observed that ferric ions (Fe³⁺) precipitate at pH above 1, while ferrous ions (Fe²⁺) precipitate at a pH above 6 as shown in Eqs. (6) and (7):

When pH > 1:

$${\text{Fe}}^{3+}+{\text{3OH}}^{-}\to \text{Fe}{\left(\text{OH}\right)}_3.{\text{mH}}_2\text{O}.$$E6

When pH > 6:

$${\text{Fe}}^{2+}+{\text{2OH}}^{-}\to \text{Fe}{\left(\text{OH}\right)}_2.{\text{mH}}_2\text{O}.$$E7

Iron can be produced from a number of sources in the formation, such as formation rock and tubulars (Fe²⁺, Fe³⁺), formation water (Fe²⁺), and treatment fluids (Fe³⁺). The concentration of iron in fluids during an acid treatment may reach 4000 mg/L at the wellhead and may surpass 100,000 mg/L by the time it reaches the formation [44, 45]. As a result, competent iron control is needed to prevent precipitation, incompatibilities, and formation damage.

Spent acids typically have a pH between 4 and 4.5 due to the formation of carbonic acid when carbon dioxide, a product of carbonate acidizing, dissolves in water. As a result, any Fe³⁺ ions present will precipitate as Fe³⁺ hydroxide when the pH reaches this level. To prevent this, iron-control agents such as APCAs are added to chelate the ions and keep them in solution. One of the first and most used iron control agents was citric acid, as illustrated by Hall and Dill [46]. However, it loses 50% of its potency at 350°F (176°C) and can result in the precipitation of calcium citrate and asphaltene/sludge issues [47].

The effectiveness of chelating agents as iron control agents is dependent on its solubility and its iron stability constants. If the chelating agent experiences solubility issues, such as EDTA, in the acidizing fluid, then it cannot chelate iron ions. Furthermore, the stability constant of the iron-chelating agent complex is also important due to competition from hydroxide ions in the solution. Therefore, the Fe³⁺-chelating agent complex must be more stable than Fe³⁺ hydroxide to prevent the precipitation of Fe³⁺ hydroxide [43]. In general, chelating agents have greater stability constants for iron and can keep iron soluble in solutions over a wide pH range as shown in Figure 6.

Figure 6.

Chelating agent optimum pH range for iron (III) (re-graphed from [48]).

It should be noted that in sour environments (H₂S), iron control is more difficult as hydrogen sulfide is a strong reducing agent that causes precipitation through the formation of iron sulfides of low solubility and higher stability compared to the iron-chelating agent complex. As certain forms of iron sulfide can be hard to remove, it is ideal to prevent their formation instead of remediation. Consequently, hydrogen sulfide scavengers must be used in combination with chelating agents to keep iron soluble in solution in sour environments [49, 50]. More information regarding iron sulfide scale is provided later in this chapter.

4. Inorganic scale removal

4.1 Calcium and barium sulfate

Inorganic scale is a huge problem in the oilfield as it tends to form restrictions in the wellbore and the surface pipelines that ultimately result in reducing production rates. Inorganic scale precipitation results from mixing two incompatible water sources, where the first contains an abundance of metal ions such as calcium, magnesium, or barium, and the other contains anions such as bicarbonates or sulfates. Another form of inorganic scale precipitation can occur through self-precipitation initiating from one water source. This occurs due to changes in the temperature and pressure conditions which ultimately change the thermodynamic equilibrium of ions in solution causing it to form inorganic scales. Inorganic scale precipitation is seen at choke vales, bends in the surface pipelines, the top 500 ft of the well, or even downhole at the intake of electrical submersible pumps.

The two most common types of inorganic scales faced in the oilfield are carbonate and sulfate precipitations. While inorganic acids such as HCl provide a cheap and easy method to remove carbonate scales, they are often unable to dissolve sulfate scales. In this scenario, chelating agents are required [49].

In the past, EDTA was commonly used to treat calcium scales in boilers and oilfield surface pipelines [50, 51, 52]. Following that, MGDA and HEIDA were tested and proved to be effective [21, 53, 54]. It was also shown that GLDA could work for this function as well [55].

Barium sulfate (BaSO₄) is the hardest to remove due to its low solubility in water (2 mg/L) [56, 57]. Chelating agents such as EDTA, HEDTA, and DTPA were extensively studied in barium sulfate scale removal treatments [58, 59, 60]. After recognizing the difficulty of dissolving barium sulfate in the oil and gas industry, authors tested and compared several chelating agents (DTPA, DOTA, CDTA and EDTA) for removal treatments [61]. Through the testing it was noted that DTPA at pH above 12 gave the best results, even when compared to chelating agents that have higher complex stability constants with barium (Ba-DOTA > Ba-DTPA).

More recently, researchers showed that larger chelating agents such as macropa and macroquin were able to outperform DTPA and DOTA at pH 8 and pH 11 conditions [62]. This opened an area of research to compare these bigger molecules to common oilfield chelating agents. Through these studies authors were able to show that the use of macropa at pH 8 outperformed the use of EDTA, HEDTA, or DTPA. Although, at high pH (>12), a dilute solution of penta potassium DTPA showed the most prominent results in barium sulfate dissolution [57, 63, 64, 65].

It was also shown that combinations of chelating agents such as DTPA/GLDA, DTPA/EDTA, or DTPA/MGDA resulted in lower performance than DTPA alone. This is mainly because barite dissolution is inhibited by the physical presence of other chelating agent molecules [66, 67]. It was also noted that barite dissolution with DTPA increases as temperature increased, indicating a thermally activated process [63, 68, 69].

Furthermore, researchers have also shown that the performance of some chelating agents decreases with an increase in concentration due to steric hindrance of chelating agent molecules on the surface of the scale [70]. Consequently, it is recommended to apply a multi-cyclic injection technique with dilute chelating agent solution [63, 71].

Several chelating agent and dissolution catalyst formulations have been proposed for sulfate scale dissolution [13, 57, 72]. Examples of dissolution catalysts are oxalic acid, fluoride, dithionate, citrate, thiosulfate, nitriloacetate, mercaptoacetate, hydroxyacetate, amino acetate and formate brine [13, 57, 63, 73, 74, 75, 76, 77, 78, 79]. Carbonate and formate ions are commonly used today [13, 63, 80]. Halogens such as fluoride show very good synergy with EDTA but negatively affect DTPA’s ability to dissolve barium sulfate [57, 68].

Paul and Fieler [57] showed the importance of cation choice to achieve a successful scale solvent, especially when a sizable amount is added to adjust the pH. Their testing compared lithium, sodium, potassium, and cesium and pointed out that the dissolution is enhanced as the size of the cation is increased. Of these cations, potassium would be the optimum choice since cesium is difficult to obtain.

4.2 Iron sulfide scale

Another inorganic scale seen in the oilfield is sulfide-based scales. Although zinc, lead, and mercury sulfide do exist in the oilfield, iron sulfide is the most common sulfide scale. This is due to the abundance of iron sources in the oilfield. Iron sulfide is black colored, oil wet, heterogenous, and explosive scale. It can come in many different forms depending on a variety of factors including temperature, pressure, and exposure time to H₂S.

Iron sulfide scale can be dissolved by conventional acid systems if the iron to sulfur ratio is close to unity [81, 82]. However, this reaction can regenerate (H₂S) and precipitates elemental sulfur. Therefore, H₂S scavengers must be added if HCl or chelating agents are to be used [83, 84, 85]. In cases where the iron to sulfur ratios is close to 0.5, iron sulfide removal can become difficult to dissolve through conventional systems. This is a result of the iron being protected by sulfur inside the mineral structure, which prevents the iron from being leached by acids, making its dissolution difficult. Common solutions in this scenario include milling or using oxidizing agents such as chlorites/chlorine dioxide or permanganates to oxidize iron sulfide to soluble Fe³⁺ ions [86, 87]. Table 2 shows a list of common iron sulfide species.

Iron sulfide	Chemical composition	Iron:sulfur ratio
Mackinawite	Fe9S8	~1
Troilite	FeS	1
Pyrrhotite	Fe7S8	~1
Pyrite	FeS2	0.5
Marcasite	FeS2	0.5

Table 2.

Common iron sulfide scale [13].

Research has shown that EDTA at pH 8–10 was capable of dissolving almost 80 wt.% of iron sulfide scale after 20 hours at 175°F (80°C) [88]. Additionally, the use of 20 wt.% EDTA and DTPA were also determined to be capable of preventing iron sulfide formation with an inhibition efficiency of 83 wt.% and 86 wt.%, respectively [89]. Other chelating agents such as GLDA and MGDA can also be used to remove FeS scale [78]. Furthermore, the use of converting agents such as (K₂CO₃) with DTPA at pH >11 proved to be efficient in dissolving iron sulfide scales with high concentration of sulfur [90]. The converting agent converts iron sulfide to iron carbonate which can be dissolved by chelating agents. This method can dissolve over 85 wt.% of the iron sulfide field scale samples at 155°F (68°C).

5. Steel corrosion by aminopolycarboxylic acids

Considering chelating agents are frequently used as substitutes for the conventional HCl acidizing recipes at high temperatures, it becomes essential to understand the capabilities of metal corrosion by chelating agent [10]. Corrosion to the tubulars can lead to a variety of concerns regarding the well’s safety. Moreover, corrosion can cause destruction to downhole pumps and equipment resulting in reduced production rates and significant financial losses.

The metal corrosion mechanism is comparable to that of mineral surfaces [91]. Through a process known as surface complexation, ligands attach and dissolve the metal oxide layer of the tubular.

As stated by Ng et al. [92], chelating compounds corrode at pH 4 through a two-step process that involves chelating agents that promote the removal of the iron oxide layer and a redox interaction between the chelating agent and the base metal. As the base metal Fe is being oxidized to Fe²⁺, the chelating agent’s carboxylic groups are being reduced into aldehydes.

Orthophosphates, silicates, chromates, amines, aldehydes, alkaloids, nitro and nitroso, thiourea, phenols, naphthol, and chelating chemicals are a few examples of corrosion inhibitors that can be utilized to decrease corrosion rates in challenging environments. The type and extent of corrosion, length of protection, and temperature affect the corrosion inhibitor choice.

The corrosion rates of chelating agents in low pH have been investigated as part of several tests using corrosion-resistant alloys and carbon steels. If the corrosion rates are kept under 0.05 lb/ft² over 6 h (45.4 mmpy), they are generally regarded as acceptable. For chrome steel, corrosion limit is lowered to 0.03 lb/ft² over 6 hours (27.2 mmpy) due to the higher material costs [93, 94]. The corrosion rate depends on many factors such as the APCA and the type of metal corroded. Some examples from literature tests at pH 4 can be seen in Table 3 for low carbon steel tubulars, and in Table 4 for corrosion resistant alloys.

Metal	APCA	Conc, wt.%	Temperature, °F (°C)	Inhibitor, v. %	Duration, hours	Corrosion rate, lb/ft^2 (mmpy)	Source
N-80	GLDA	20	300 (149)	0	12	0.724 (328.8)	[92]
	HEDTA	20	300 (149)	0	12	0.803 (364.6)
	EDTA	20	300 (149)	0	12	0.858 (389.6)
	MGDA	20	300 (149)	0	12	0.642 (291.5)
	GLDA	20	350 (176)	0	12	0.754 (342.4)
	HEDTA	20	350 (176)	0	12	0.974 (442.3)
	EDTA	20	350 (176)	0	12	1.07 (485.9)
	MGDA	20	350 (176)	0	12	0.76 (345.1)
	HEDTA	20	350 (176)	1	12	0.0102 (4.6)
	MGDA	20	350 (176)	1	12	0.00561 (2.5)
L-80	GLDA	20	300 (149)	0	6	0.5937 (539.2)	[95]
	GLDA	20	300 (149)	0.005	6	0.0262 (23.8)
	GLDA	20	300 (149)	0.001	6	0.0394 (35.8)	[96]
	GLDA	20	300 (149)	0	6	0.1956 (177.6)
	GLDA	20	300 (149)	0.001	6	<0.05 (<45.4)	[97]
	HEDTA	20	300 (149)	0	6	0.8341 (757.5)	[95]
	HEDTA	20	300 (149)	0.005	6	0.1300 (118.1)
	HEIDA	20	300 (149)	0	6	0.6519 (592.0)	[96]
	MGDA	20	300 (149)	0	6	0.4596 (417.4)

Table 3.

Low-carbon steel (LCS) corrosion rate summary.

Metal	APCA	Conc., wt.%	Temperature, °F (°C)	Inhibitor, v. %	Duration, hours	Corrosion rate, lb/ft^2 (mmpy)	Source
13Cr	GLDA	20	300 (149)	0	6	0.0009 (0.8)	[96]
	GLDA	20	300 (149)	0	6	0.0080 (7.3)	[55]
	HEDTA	20	300 (149)	0	6	0.3310 (300.6)
	HEDTA	20	300 (149)	0	6	0.5253 (477.1)	[98]
	HEIDA	20	300 (149)	0	6	0.0590 (53.6)	[96]
	MGDA	20	300 (149)	0	6	0.0900 (81.7)
	ASDA	20	300 (149)	0	6	0.0563 (51.1)
	HEDTA	20	350 (176)	0.3	4	0.0100 (9.1)	[91]
	GLDA	20	350 (176)	0.5	6	0.0496 (45.0)	[98]
	GLDA	20	350 (176)	0	6	0.3478 (315.9)
S13Cr	GLDA	20	350 (176)	0	6	0.0187 (17.0)	[98]
22Cr	GLDA	20	300 (149)	0	6	0.0001 (0.09)	[99]

Table 4.

Corrosion resistant alloy (CRA) corrosion rate summary.

6. Degradation of aminopolycarboxylic acids

The effectiveness of chemical treatments in the oil and gas industry depends on the performance of the chemical agents. Its degradation under extreme conditions can render the treatment ineffective. In the oilfield, this can occur due to high bottomhole temperatures that cause the APCA to break down. Therefore, care must be taken when selecting the APCA for the treatment.

Thermal decomposition of chelating agents has been studied by many authors. Research on the degradation of NTA and EDTA has shown that EDTA hydrolyzes to produce HEIDA and IDA, around 500°F (260°C) and pH 9.5 [100]. HEIDA was later shown to undergo further hydrolysis to yield ethylene glycol and IDA.

N-methyliminodiacetic acid (MIDA), the first breakdown product of NTA, was converted at 560°F (293°C) into methylsarcosine, which was then converted into trimethylamine. This process was non-hydrolytic and proceeded stepwise, eventually yielding carbon monoxide, carbon dioxide, formaldehyde, and methylamines [101]. As a result of the average field temperatures being much lower, these decompositions are unlikely.

Several authors [21, 75, 76, 77, 102] investigated how GLDA decomposed thermally. When heated at 300°F (149°C) and 350°F (176°C), GLDA displayed thermal stability temperature equivalent to that of HEDTA; the breakdown products were cyclic GLDA and formic acid [26]. IDA, oxotetrahydrofuran-2 carboxylic acid, hydroxyglutaric acid, and acetic acid were identified as the GLDA products of decomposition through mass spectrometry [103].

At temperatures as high as 350°F (176°C), MGDA and HEIDA exhibited outstanding stability, showing no degradation after 6 hours [36]. But it was observed that ASDA decayed rapidly, with less than 10 wt.% of the APCA left after the same amount of time at 300°F (149°C). Since the byproducts of degradation will no longer be able to carry out the original ligand’s intended function, such as controlling iron, removing scale, or acidifying the matrix, these results necessitate the application of an APCA with sufficient thermal stability for high-temperature wells encountered in the oilfield.

The biodegradability of APCA used in the oilfield has also come under close examination because of a recent rise in environmental regulations. Chelating agents in the environment have the power to chelate metal ions, which allows them to affect metal speciation and bioavailability as well as remobilize toxic heavy metals into groundwater [104, 105, 106]. Because of the potential for groundwater contamination, this adds another degree of consideration which must go into choosing the best APCA for treatment. Depending on the amount of nitrogen atoms that exist in the structure, chelating compounds can differ in their susceptibility to biodegradation. When compared to chelating compounds with numerous nitrogen atoms, those with a single nitrogen atom tend to be more biodegradable [107, 108, 109, 110].

7. Summary

There are many challenges with applying conventional HF or HCl acids for oilfield stimulation. Chelating agent research has progressed over the previous decades to identify the best application for them in the petroleum industry as a replacement to conventional acid systems. Common chelating agents were addressed in this work, accompanied by a thorough analysis of the unique features that distinguish presented chelating agents. A summary of the structural formulas and chemical properties of some of the most common chelating agents was also provided.

Chelating agents have been utilized for oilfield stimulation either as a stand-alone treatment or in combination with traditional acids to lessen some of the downsides. They also show good performance even at high temperatures and serve to reduce the number of additives required for stimulation and scale removal treatments, such as corrosion inhibitors due to their low corrosivity. In addition, their ability to chelate iron is invaluable to prevent formation damage from iron hydroxide precipitation and sludge formation, common phenomena that can quickly result in severe formation damage. Furthermore, many chelating agents are environmentally friendly and will decompose readily when exposed to the elements and minimize environmental impact. Despite this, they are still resistant to thermal decomposition and are able to perform high temperature treatments.

Despite these benefits, it is important to exercise care when selecting the appropriate chelating agent for the application. Due to the wide variety of options, properties such as solubility, stability constants, reactivity, and price must be considered when choosing the most appropriate chelating agent. By providing a comprehensive summary of the chemistry and applications of chelating agents in the oil and gas industry, we hope readers will be able to select the best chelating agent for their treatment.

Nomenclature

APCA	aminopolycarboxylic acids
ASDA	L-aspartic acid N, N-diacetic acid
CDTA	trans-l,2-cyclohexylenediaminetetraacetic acid
CRA	corrosion resistant alloys
DOTA	1,4,7,10-tetraazacyclododecane-1,4,7,10-tetraacetic acid
DTPA	diethylenetriaminepentaacetic acid
EDTA	ethylenediaminetetraacetic acid
GLDA	L-glutamic acid N, N-diacetic acid
HCl	hydrochloric acid
HEDTA	hydroxyethyl ethylenediaminetriacetic acid
HEIDA	hydroxyethyliminodiacetic acid
HF	hydrofluoric acid
IDA	iminodiacetic acid
LCS	low carbon steel
MGDA	methylglycinediacetic acid
MIDA	N-methyliminodiacetic acid
NTA	nitrilotriacetic acid
VES	viscoelastic surfactant