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Organic Acids for Stimulation
Purposes: A Review
Luai Alhamad, Ahmed Alrashed, Eiman Al Munif, and Jennifer Miskimins, Colorado School of Mines
Summary
Hydrochloric acid (HCl) is the acid of choice for acidizing operations in most carbonate formations, and is the base acid that is commonly paired with hydrofluoric acid (HF) in most sandstone applications. However, high dissolving power, high corrosion rate, lack of
penetration, and sludging tendency coupled with high temperature (HT) can make HCl a poor choice. Alternatively, weaker and lesscorrosive chemicals, such as organic acids, can be used instead of HCl to avoid these issues. The objective of this paper is to provide an
intensive review on recent advancements, technology, and problems associated with organic acids. The paper focuses on formic, acetic,
citric, and lactic acids.
This review includes various laboratory evaluation tests and field cases that outline the use of organic acids for formation-damage
removal and dissolution. Rotating-disk-apparatus (RDA) results were reviewed to determine the kinetics for acid dissolution of different
minerals. Additional results were collected from solubility, corrosion, coreflooding, inductively coupled plasma, X-ray diffraction, and
scanning-electron-microscope (SEM) diffraction tests.
Because of their retardation performance, organic acids have been used along with mineral acids, mainly a formic/HCl mixture, or
as a standalone solution for HT applications. However, the main drawback of these acids is the solubility of reaction-product salts. This
challenge has been a limiting factor of using citric acid with calcium-rich formations because of the low solubility of calcium citrate.
However, the solubility of the salts associated with formic, acetic, and lactic acid can be increased when these acids are mixed with gluconic acid because of the ability of gluconate ion to chelate calcium-based precipitation. In terms of formation-failure response, organic
acids are in lower risk of causing a failure compared with HCl, specifically at deep formation treatments. Organic acids have also been
used in other applications. For instance, formic acid is used in HT operations as an intensifier to reduce the corrosion rate caused by
HCl. Formic, acetic, and lactic acids can be used to dissolve drilling-mud filter cakes. Citric acid is commonly used as an ironsequestering agent.
This paper shows organic acid advances, limitations, and applications in oil and gas operations, specifically in acidizing jobs. The
paper differentiates and closes the gap between various organic acid applications along with providing researchers an intensive guide
for present and future research.
Organic Acids
An organic acid is any organic compound with acidic characteristics that does not dissociate completely in water and donates hydrogen
ions to water molecules. Mineral acids, such as HCl, are stronger than organic acids because of their complete dissociation in water
(Robert and Crowe 2000; Burgos et al. 2004). Consequently, organic acids are called weak acids. The strength of mineral and organic
acids can be characterized by the value of the acid-dissociation constant (Ka) (Robert and Crowe 2000; Buijse et al. 2004). This constant
is used to measure the acid-ionization strength. It is common to take the logarithmic value of Ka, referred to as pKa, and to use it for
acid-strength comparisons.
High pKa values reflect weak acids, while low values reflect strong acids. For example, the pKa value of HCl is less than unity,
while the pKa value of formic acid is 3.751. The acid-dissociation process is represented by the acid chemical species where [HA] is an
acid, [A–] is a conjugate base, and [Hþ] is a hydrogen proton, as presented in Eq. 1. Organic acids cannot reach a complete dissociation
because of their conjugate base nature to reform the acid again; consequently, conjugate base stability determines the acid strength
(Robert and Crowe 2000; Buijse et al. 2004; Chang et al. 2008).
HA
A þ Hþ : . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð1Þ
Organic acids are classified according to their chemical-structure groups, with the carboxyl group being the most common. Carboxyl
organic acids mainly consist of a carboxyl group and other molecules (R-COOH). For recognition purposes, carboxylic acids are known
depending on the number of carbon atoms and associated molecular structure (Badea and Radu 2018). For example, formic acid is called
methanoic acid because it contains only one carbon atom (HCOOH), while acetic acid is known as ethanoic acid because it contains two
carbon atoms (CH3COOH). Because citric acid contains three carboxyl groups and lactic acid contains one hydroxy group, they are part
of the tricarboxylic acids and alpha hydroxy acids, respectively (Al-Khaldi et al. 2005; Al-Harbi et al. 2012; Badea and Radu 2018).
The most commonly used organic acids in the oil and gas industry are formic, acetic, citric, and lactic acids because of their ability
to meet different application specifications (Al-Khaldi et al. 2005; Chang et al. 2008; Nasr-El-Din et al. 2009). The strong reactivity
and corrosivity of HCl, coupled with more-HT applications, encouraged researchers to further study organic acids for reservoirstimulation operations. Organic acids are much weaker than HCl, which can inflict a slower reaction with formation rocks and lower
corrosion with steel components (Al-Khaldi et al. 2005; Chang et al. 2008). Among the four studied organic acids in this review, formic
acid is the most corrosive acid when it is used without a corrosion inhibitor (CI). Citric acid can cause some corrosion problems, but it
is usually used at a low concentration (de Wolf et al. 2017).
The objectives of this paper are to recognize these acids and their applications in reservoir-stimulation operations; to understand
their chemical interactions with stimulated formation, ions in solutions, and steel tubulars; and to highlight advances and limitations
associated with these acids.
C 2020 Society of Petroleum Engineers
Copyright V
This paper (SPE 199291) was accepted for presentation at the SPE International Conference and Exhibition on Formation Damage Control, Lafayette, Louisiana, USA, 19–21 February 2020,
and revised for publication. Original manuscript received for review 25 December 2019. Revised manuscript received for review 26 May 2020. Paper peer approved 14 June 2020.
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Formic Acid Overview
Formic acid (HCOOH) is a monoprotic acid that is commonly used to stimulate carbonate formations and to enhance CIs for HT applications. The chemical structure of formic acid is shown in Fig. 1.
O
H
OH
Fig. 1—Formic acid chemical structure.
In stimulation operations, formic acid has been applied widely to remove formation damage in carbonate formations and to dissolve
carbonated rocks in sandstone operations. Formic acid is considered the second-most reactive organic acid after citric acid, which
assures a strong reactivity with formation materials. Accompanied by another acid or as a standalone fluid, formic acid has secured its
place as a stimulation-fluid option (Robert and Crowe 2000; Buijse et al. 2004). However, formic acid concentration is constrained to
9 wt% because of the low solubility of its salt, calcium formate, which results from complexing with calcium ions in solution (Rabie
et al. 2015). Table 1 shows the solubility of different organic acid calcium salts in water.
Acid
Reaction Product (Salt)
Solubility (g/L)
Acetic
Calcium acetate
300 (Rabie et al. 2015)
Formic
Calcium formate
166 (Rabie et al. 2015)
Citric
Calcium citrate
0.85 (Nasr-El-Din et al. 2007)
Lactic
Calcium lactate
79 (Rabie et al. 2015)
Table 1—Solubility of different organic acid calcium salts in water.
Formic acid has also been used in corrosion-protection blends because of its ability to aid other CI packages (Brezinski 1999;
Al-Katheeri et al. 2002). Brezinski (1999) reported that formic acid could aid corrosion inhibition by forming carbon monoxide (CO)
on the metal surface that can intensify corrosion inhibition through different mechanisms. Formic acid mixed with strong acid and heat
undergoes a dehydration process to produce water and CO. Since 1940, CO has been used in corrosion protection (Heusler and
Cartledge 1961). The decomposition of formic acid occurs through two paths: decarbonylation (Eq. 2) or dehydration (Eq. 3).
HCOOH ! CO2 þ H2 ; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð2Þ
HCOOH ! CO þ H2 O: . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð3Þ
Acetic Acid Overview
Acetic acid is another monoprotic acid (CH3COOH) and is considered the second-simplest carboxylic acid after formic acid. Its chemical structure is shown in Fig. 2.
H
H
C
H
O
C
O
H
Fig. 2—Acetic acid chemical structure.
In stimulation operations, acetic acid can be paired with another acid to dissolve carbonates, can be used as an iron agent, or can be
used to remove drilling-fluid filter cakes (Ewing et al. 1983; Buijse et al. 2004; Al Moajil and Nasr-El-Din 2007). It is usually used up
to 13 wt% to avoid an excessive level of acetate ions in solution that can result in producing calcium acetate salt. Calcium acetate solubility in water is shown in Table 1.
On the other hand, other studies showed that acetic acid mixed with HCl or formic acid could not contribute or spend appropriately
because of a slow ionization step (Chatelain et al. 1976; Burgos et al. 2005). Increasing acetic acid concentration creates two challenges. The first challenge is the precipitation of the reaction product of calcium acetate. The second challenge is the slow acid reactivity as the acid concentration increases, where less acetic acid is being converted or dissociated, as described by Chatelain et al. (1976).
The acetic acid-ionization fraction is shown in Fig. 3.
Citric Acid Overview
Citric acid (C6H8O7) is a weak acid that occurs naturally from citrus fruits. It is considered a triprotic acid that has three carboxylic
(COOH) groups and one hydroxyl (OH) group. In solution, its anionic derivative is citrate (C6H8O3–
7 ), which can complex with different
ions to create a salt or an ester (Al-Khaldi et al. 2003). The chemical structure of citric acid is shown in Fig. 4.
Citric acid has been widely used in the oil and gas industry as an iron-control agent to stabilize ferric and ferrous ions from spent
acid (Smith et al. 1969). Recently, citric acid has been investigated in stimulation operations to stimulate carbonate formations as a
standalone system or to stimulate sandstone formations with HF. However, the extremely low solubility of its reaction product (calcium
citrate) casts doubts on the use of citric acid as a stimulation fluid for calcium-rich rocks (Al-Khaldi et al. 2003, 2010; Burgos et al.
2004; Al-Harbi et al. 2012, 2013; Mohamed et al. 2015). The solubility of calcium citrate is shown in Table 1.
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Fraction Converted
1.0
0.8
0.6
2 wt%
10 wt%
15 wt%
0.4
0.2
0
0
25.0
50.0
75.0
Temperature (°C)
Fig. 3—Acetic acid converted fraction per temperature for 2, 10, and 15 wt% concentrations (Chatelain et al. 1976).
O
HO
O
OH
O
OH
OH
Fig. 4—Citric acid chemical structure.
Chabarek and Martell (1959) discussed the citric acid capability of complexing with other ions such as magnesium, iron, and calcium. Citric acid can produce water-soluble compounds and water-insoluble compounds. When complexing with other ions, the behavior of citric acid and the resulting compound depends on ionic strength, pH, and temperature. It can act as a sequestering agent that
produces a water-soluble compound or as a chelating agent that produces a water-insoluble compound (Chabarek and Martell 1959;
Al-Khaldi et al. 2003).
Lactic Acid Overview
Lactic acid [CH3CH(OH)CO2H] is an a-hydroxycarboxylic acid, shown in Fig. 5, that has served stimulation operations as an ironcontrol agent and as a main component in drilling-fluid mudcake (Elkatatny and Nasr-El-Din 2012). Lactic acid produces calcium lactate salt when it reacts with carbonate, which has moderate solubility in water (Al Moajil and Nasr-El-Din 2014). Although lactic acid
is comparable with formic acid in terms of dissociation constant and dissolution of calcite, it has a limited application in stimulation
operations because of the reaction-salt solubility (Rabie et al. 2014).
O
OH
OH
Fig. 5—Lactic acid chemical structure.
Lactic acid can be used in the solid form as a precursor that hydrolyzes with water under certain temperatures to release the acid in
situ. This advantage can bring good treatment results because corrosion challenges can be alleviated because of the in-situ released
acid, where a well tubular would not be exposed to the acid. Along with that, lactic acid precursors can serve as a diverting agent in
multizone acid-stimulation treatments or hydraulic fracturing. In terms of leakoff rate, solid lactic acid reduces the risk of losing the
treatment’s fluid into the formation because the acid would be released when it is in the targeted place (Nasr-El-Din et al. 2009).
Organic-Acid-Dissociation Process
Mineral acids, such as HCl, dissociate completely in water, whereas weak acids do not. During the organic-acid-dissociation process
(Eq. 4), the dissociated state (A–) and the undissociated state (HA) coexist within the solution because of the instability state and reverse
reaction behavior of organic acids (Nasr-El-Din et al. 2007, 2009; Al-Harbi et al. 2012).
HA
A þ Hþ : . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð4Þ
There have been different studies and research conducted to understand the reasons behind the incomplete dissociation of organic
acids (Chatelain et al. 1976; Williams et al. 1979; Buijse et al. 2004; Chang et al. 2008; Al-Khaldi et al. 2010). Buijse et al. (2004) mentioned that as the solution pH increases, more Hþ is generated, which shifts Eq. 4 to the right. However, during the acid and rock reaction, carbon dioxide (CO2) gas is generated and kept in solution because of the reservoir pressure being greater than 1,000 psi. CO2
enhances the reaction’s chances to be reversed, which would keep the pH in the range of 4 to 6. At these values, organic acids are not
completely dissociated. In addition, CO2 buffers the acid solution through the generation of carbonic acid, which then buffers the solution pH to nearly 4.5 (Williams et al. 1979; Buijse et al. 2004; Chang et al. 2008). Dissolved CO2 can generate carbonic acid through
Eqs. 5 through 7.
H2 O þ CO2
H2 CO3
H2 CO3 ; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð5Þ
Hþ þ HCO
3 ; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð6Þ
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HCO
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Hþ þ CO2
3 : . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð7Þ
The acid dissociation constant Ka, shown in Eq. 8, is a constant that measures the concentration ratio of both sides of the equation. The
dissociation constant helps to measure the acid’s affinity to release a proton (Hþ). However, it is more convenient to express the dissociation constant using Eq. 9,
Ka ¼
½ A ½ H þ ; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð8Þ
½AH
pKa ¼ logðKa Þ: . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð9Þ
Strong acids have low values of pKa compared with weak acids that have larger pKa values. For example, HCl dissociates entirely in
aqueous solutions, so it has a pKa value that is less than unity, compared with 3.08 for citric acid. Table 2 shows different acid Ka and
pKa values, starting with the strongest.
Acid
HA
Ka
pKa
Hydrochloric
HCl
>1
<1
Nitric
HNO3
>1
Sulfuric
HSO–4
1:210
–2
1.92
Citric, first step
C6 H8 O7
8:4104
3.128
Hydrofluoric
HF
7:210–4
3.14
Formic
HCOOH
1:7710–4
3.75
<1
–4
Lactic
CH3 CHðOHÞCO2 H
1:3810
3.86
Acetic
CH3 COOH
1:7610–5
4.756
Table 2—Dissociation constants for different acids at 25 C.
Formic and acetic acids are monoprotic acids that dissociate by following Eqs. 10 and 11. The ionic species for both acids are
shown in Figs. 6 and 7; it can be seen that both acids require a high pH value to reach complete spending.
HCOOH ðformic acidÞ
HCOO þ Hþ pKa ¼ 3:75; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð10Þ
CH3 COOH ðacetic acidÞ
CH3 COO þ Hþ pKa ¼ 4:756: . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð11Þ
1.0
HCOO–
HCOOH
0.8
α
0.6
pH = 1.02
onset of free HCOO–
0.4
0.2
0
0
2
4
6
8
10
12
14
pH
Fig. 6—Distribution of formic acid conjugate base as a function of solution pH value (Al-Harbi et al. 2012).
1.0
CH3COO–
CH3COOH
0.8
0.6
α
pH = 2.01
onset of free CH3COO–
0.4
0.2
0
0
2
4
6
8
10
12
14
pH
Fig. 7—Distribution of acetic acid conjugate base as a function of solution pH value (Al-Harbi et al. 2012).
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Citric acid is a triprotic acid that consists of different ionic species. Ionic species distribution is dependent on solution pH and the
dissociation constant, as illustrated in Fig. 8. In addition, citric acid dissociates stepwise, as in Eqs. 12 through 15 (Chabarek and
Martell 1959; Smith and Martell 1989).
H3 C6 H4 O6 OH
H2 C6 H4 OH þ Hþ pKa1 ¼ 3:128; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð12Þ
H2 C6 H4 O6 OH
HC6 H4 O6 OH2 þ Hþ pKa2 ¼ 4:76; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð13Þ
HC6 H4 O6 OH2
C6 H4 O6 OH3 þ Hþ pKa2 ¼ 6:4; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð14Þ
C6 H4 O6 OH3
C6 H4 O6 OH4 þ Hþ pKa2 ¼ 11:6: . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð15Þ
1.0
Citric2–
0.8
Citric3–
Citric
γ
0.6
0.4
Citric4–
Citric–
0.2
0
0
2
4
6
8
10
12
14
Equilibrium pH
Fig. 8—Distribution of citric acid conjugate bases as a function of solution pH value (Al-Khaldi et al. 2003).
Lactic acid is a hydroxy carboxylic acid that dissociates partially in water to give lactate and hydrogen ions, as shown in Eq. 16.
Dissociation of lactic acid vs. solution pH is shown in Fig. 9.
LA $ Hþ þ L pKa ¼ 3:86: . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð16Þ
1.0
0.9
–
Lactate (L )
0.8
Mole Fraction
0.7
Lactic acid (HL)
0.6
pKa = 3.79
0.5
0.4
0.3
0.2
0.1
0
0
1
2
3
4
5
6
7
8
pH
Fig. 9—Distribution of lactic acid conjugate base as a function of solution pH value (Rabie et al. 2014).
Carbonate Stimulation
HCl is the most commonly used acid for carbonate-acidizing operations and has been used for many years (Buijse et al. 2004). However, the strong dissolution level, high corrosion rate, and poor etching pattern drove the industry and academia to retard HCl using
emulsified acid, gelled acid, or surfactants. Unfortunately, all these methods have their limitations (Nasr-El-Din et al. 2007). Weak
acids have been used in carbonate-stimulation operations as a choice instead of HCl because of their retardation reaction with carbonate
and as iron-control agents because of their ability to stabilize ferric and ferrous ions in solution (Elkatatny and Nasr-El-Din 2012).
Because weak acids are less reactive than HCl, many investigations have been completed to identify weak acid benefits in stimulation operations. The conventional organic acids, formic and acetic, tackled some of the main challenges, such as tubular corrosion, acid
penetration, and reaction rate. Along with that, citric acid was introduced in capsule form to generate in-situ acid states (Al-Khaldi
et al. 2003). However, there are two drawbacks to these organic acids in stimulation treatments. The first drawback is the solubility of
their reaction products. The second drawback is the reaction of those acids with carbonates because the reaction is reversible, which
prevents the acid from reaching a complete reaction stage (Nasr-El-Din et al. 2009).
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Weak acids react with carbonate through hydrogen ions that dissociate in water from organic acids. The hydrogen attack on carbonate rocks would produce calcium ions, water, and CO2 that are dissolved in solution, as can be seen in Eq. 17.
2Hþ þ CaCO3
Ca2þ þ H2 O þ CO2 : . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð17Þ
In terms of the corrosiveness of organic acid, Nasr-El-Din et al. (2012) studied the corrosion rates of carbon-steel tubulars caused by
different organic acids. In the absence of a CI (at 1,000 psi and 300 F), formic acid showed to be the top corrosive organic acid to L-80
tubulars, whereas the glutamic acid N,N-diacetic acid (GLDA) system was the least corrosive one. Fig. 10 illustrates the corrosion rates
of the different tested organic acid systems. Alhamad and Miskimins (2020) investigated the corrosion rate of 9 wt% lactic acid and
found that it can be as corrosive as formic acid, with a tendency of causing iron lactate precipitation.
6-hour Metal Loss (lbm/ft2)
1.2
1.0
0.8
0.6
0.4
0.2
20 wt%
GLDA
20 wt%
HEDTA
10 wt%
10 wt%
10 wt%
Acetic acid Citric acid Formic acid
Fig. 10—Corrosion rates of formic acid, citric acid, acetic acid, hydroxyethylene-diamine-tetra-acetic acid (HEDTA), and GLDA
(no CI) (Nasr-El-Din et al. 2012).
The following subsections show different studies conducted on formic, acetic, citric, and lactic acids for carbonate acidizing, along
with example field results.
Formic and Acetic Acids. One of the most highlighted studies for formic acid and acetic acid thermodynamic limitations was conducted by Chatelain et al. (1976). This study established correlations between unreacted organic acid fraction with temperature and initial acid concentration in reaction with limestone and dolomite samples. Both formic- and acetic-acid-converted fractions showed the
same trend when concentrations and temperature were increased, as shown in Fig. 11. Reaction limitations were justified by dissolved
CO2 gas and reaction-product salts of calcium formate and calcium acetate (Chatelain et al. 1976).
1.0
5 wt% formic
Fraction Converted
Fraction Converted
1.00
0.96
0.92
10 wt%
0.88
0.84
0.80
0
50
100 150
200
Temperature (°F)
250
300
5 wt%
0.8
0.6
2 wt% acetic
10 wt%
15 wt%
0.4
0.2
0
0
50
100 150
200
250
300
Temperature (°F)
Fig. 11—Converted fraction of (left) formic acid and (right) acetic acid vs. temperature for different initial concentrations (Chatelain
et al. 1976).
Dill and Keeney (1978) investigated the addition of formic acid to HCl to lower the corrosion rate and to increase wormhole length
in limestone stimulations. After adding the CI, corrosion tests showed that 7.5% HCl and 10% formic acid were less than the corrosion
rate of 0.05 lb/ft2, without the need for any intensifier, up to 400 F for 6 hours. Moreover, the blend was retarded with a new chemical
retarder that decreased the blend reactivity, which resulted in a weaker reaction rate. They reported that although formic acid can lower
the corrosivity of HCl, after a particular concentration, the weak acid would contribute to an increase in corrosion rate, as shown in
Fig. 12 (Dill and Keeney 1978).
Van Domelen and Jennings (1995) reported one of the first massive formic/acetic-acid-blend treatments in a limestone formation.
Their goal was to develop a new organic acid blend that would be as effective as HCl with lower reaction and corrosion rates for HT
operations. Their laboratory work found that a 9 wt% formic and 13 wt% acetic acid blend can have the same dissolving power as a
15 wt% HCl solution, but with retarded reaction and less corrosion risk. The blend was also tested with a gelling agent to retard the
reaction by building viscosity. The retardation test showed that organic acid would yield better viscosity than HCl by providing more
stability to the gelling-agent polymers. The blend was advanced for application in the Arun Formation in Indonesia, where the average
reservoir temperature is 350 F. The field responded positively in terms of total production and tubular protection as the increase of the
total production was sustained for an extended period with an average of 1.55-times production improvement (Van Domelen and
Jennings 1995).
After the successful treatment in Arun Formation in Indonesia, da Motta et al. (1998) applied a 7% formic and 5% acetic acid mixture to stimulate and remove calcium carbonate scale in the Merluza Formation offshore of Brazil. Their study was inspired by the need
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to find an alternative blend to the HCl-based blend that could assure maximum protection for CR-13 alloys. Initially, a 10% acetic acid
was used to remove calcium carbonate scale, but according to the laboratory results, a 7% formic and 5% acetic acid mixture is four
times more effective than acetic acid alone. Formic acid is far stronger than acetic acid, as can be seen in the dissociation constants presented in Table 2. In the mixture, acetic acid would buffer the solution, providing more retardation, and formic acid would increase the
blend strength. Solubility tests showed that 10% acetic acid would need 2 hours to dissolve calcium carbonate scale, while the blend
would need only 30 minutes. Field results showed that the new blend was able to increase the cumulative production from the treated
wells from 837 000 to 1.576 million m3/d (da Motta et al. 1998).
Corrosion in mg/cm2 (lbm/ft2)
58.8
(0.12)
10%
HCl
49.0
(0.10)
39.2
(0.08)
7¹⁄₂%
HCl
29.4
(0.06)
19.6
(0.04)
9.8
(0.02)
0
2
4
6
8
10
12
14
Formic Acid (by wt%)
Fig. 12—Corrosion rates vs. formic acid concentration blended with 7.5 or 10% HCl (Dill and Keeney 1978).
Fredd and Fogler (1998) studied the acetic acid reaction with calcite using an RDA. Over the pH range of 2.3 to 2.9, the dissolution
rate of acetic acid was mass-transfer limited because of the reversible reaction of the acid. This implies that the reaction is limited by
the transport of reactant and reaction products to and from the calcite surface, respectively. At a pH value greater than 3.7, they reported
that the reaction was surface limited, and under that value, the dissolution was in a transition between the two reaction characteristics
(Fredd and Fogler 1998). Fig. 13 shows the surface reaction of any organic acid and the reactant and product transportation to and from
the reaction surface.
Ca2+
HA
CO2
Diffusion
Dissociation
+
–
HA H + A
Boundary
layer
+
2+
2H + CaCO3 ↔ Ca + H2O + CO2
Fig. 13—Surface reaction of an organic acid with rock surface and the reactant and product transportation to and from the
reaction surface.
One of the main challenges for carbonate acidizing using HCl is maintaining the optimal injection rate. Because HCl is highly reactive with calcium carbonate at HTs, a face-dissolution pattern can occur at a low injection rate (Akanni and Nasr-El-Din 2016). HCl
requires a high injection rate to reach the optimal results that would exceed the maximum-allowable rate (Fredd and Fogler 1998;
Al-Duailej et al. 2013; Akanni and Nasr-El-Din 2016). Ultimately, the low reactivity of organic acids allows them to be injected at an
optimal injection rate to develop the desired wormhole pattern (Huang et al. 2000; Al-Duailej et al. 2013; Akanni and Nasr-El-Din 2016).
Huang et al. (2000) illustrated this fact by coreflooding different cores using 11.6% acetic acid at room temperature. Results showed
that acetic acid requires a lower injection rate and more pore volume (PV) to breakthrough compared with the previous 3.4% HCl
linear-modeling results. They reported that the advances of acetic acid can be more feasible when it is used at HT applications, as per
linear-modeling results. In addition, the slow reactivity of acetic acid would result in larger wormholes than those created from HCl.
However, according to an estimation driven from the radial-Darcy-flow equation, the skin factor would be lowered with wormholes as
small as the ones obtained from HCl (Huang et al. 2000).
Acetic acid was also investigated for removing drilling-fluid filter cake; HCl was not a suitable choice because of its incompatibility
issues with unconsolidated formations. Nasr-El-Din et al. (2001) reported that 10 wt% regular and emulsified acetic acids showed good
solubilities in laboratory tests. However, laboratory and field tests concluded that the emulsified acid could impose damages because of
fines migration. Both acetic acid formulas were applied in the Shu’aiba Carbonate Reservoir in Saudi Arabia, where the emulsified acid
results were less than expectations, while the regular acid showed significant productivity improvement (Nasr-El-Din et al. 2001).
Promising laboratory and field results for formic and acetic acids motivated different companies to consider these weak acids for
field applications. Nasr-El-Din et al. (2002, 2003) applied 9 wt% formic with 15 wt% HCl gelled acids to stimulate the Khuff Formation
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in Saudi Arabia. The main idea was to maintain a slow reaction using formic acid and a gelling agent along with maintaining a highly
reactive mixture that is less corrosive than 28 wt% HCl to protect CR-13 tubulars. Corrosion tests were performed on L-80 steel and
SCR-13 coupons while maintaining corrosion rates of 0.05 and 0.03 lb/ft2, respectively. The tests were conducted for 3 hours under
2,500 psi and 250 F. An addition of 15 gal/1,000 gal organic acid CI ensured that corrosion rates were less than 0.02 lb/ft2 and free of
any pitting. As was expected from the laboratory results, all field trials responded positively, with a substantial increase in gasproduction and minimum-corrosion results. Al-Mutairi et al. (2005) used the same formula of in-situ gelled acid to acid fracture multiple wells in the same formation. Effective results were obtained, similar to the previous treatments conducted by Nasr-El-Din et al.
(2002, 2003); the formic/HCl blend showed enhancement compared with other systems in terms of well responses and
corrosion outcomes.
An intensive study was performed with 13/9% acetic/formic acids by Buijse et al. (2004), where they built a model to capture formic/
acetic acid reactions with carbonates. The model was straightforward and was compared with actual laboratory results to confirm its
validity and accuracy. One of their most significant outcomes was the fact that even with 13/9% acetic/formic acid blends being comparable with 15% HCl in terms of acid strength, the incomplete spending of organic acids undermines the dissolving power to that of 8% HCl
(Buijse et al. 2004). The model demonstrated the low dissociation of acetic acid in the 13/9% blend because most of the dissolution was
contributed by formic acid, while the acetic acid contribution was insignificant. However, the acetic acid addition plays a significant role
in slowing the reaction and keeping the pH in the acidic range. The model also showed the advances of organic acids in terms of penetration depth. For example, the 13/9% blend reached up to 100 m, while 15% HCl only reached 40 m because of rapid spending. Fig. 14
shows the acid concentration as a function of penetration depth generated from the Buijse et al. (2004) model. Field application of this
blend was performed in Maraca Field in Venezuela, where the total production increased from 3,600 to 20,500 BOPD.
Acid Concentration (mol/L)
5.0
Fracture height: 20 m
Fracture width: 0.5 cm
Fracture length: 100 m
Pump rate: 15 bbl/min
Volume: 100 m3
Acid viscosity: 10 cp
Reaction order n: 0.6
Rate constant k: 0.023
4.5
4.0
3.5
3.0
15% HCI
2.5
2.0
13/9% Acetic/formic
1.5
1.0
Incomplete spending
0.5
0
0
20
40
60
80
100
Distance from Wellbore (m)
Fig. 14—Acid concentration as a function of penetration depth generated from the Buijse et al. (2004) model.
Nonetheless, Burgos et al. (2005) reported that wells treated with the 13/9% acetic/formic acids in Maraca Field showed a decline in
production after a few months. Encapsulated citric acid treatments were conducted, but poor results occurred because of unbroken polymers in the pad fluids, poor etching, and precipitation of calcium citrate salt (Burgos et al. 2004). As a result, laboratory and simulation
tests were performed on two blends using acetic and formic acids coupled with a low concentration of HCl. The addition of 5 wt% HCl
to formic acid showed production improvement in formations with permeability of 15 md or greater. The addition of HCl to acetic acid
overshadowed the acetic acid contribution because acetic acid would partially spend and dissociate in water because of its low
dissociation constant.
Conversely, formic acid has a higher dissociation constant that would add more value when coupled with HCl. The proposed mixture was 7/11% HCl/formic acid. The 7% HCl could be spent within the wellbore-zone area, and the formic acid would generate
extended conductive wormholes by evolving in the generated wormhole tips. HCl was limited to 7% to avoid sludging and severe corrosion, while formic acid was limited to 11% to avoid calcium formate precipitation. The new blend was applied in the previous wells
that were stimulated with the 13/9% acetic/formic blend. Results showed that better productivity-index values were obtained, except in
a well that was already around maximum conductivity from the previous treatment (Burgos et al. 2005). Fig. 15 shows a comparison of
productivity-index values between wells treated with the 13/9% acetic/formic acids and the 7/11% HCl/formic acids.
2.00
1.80
Organic acid
7/11 acid
Expon. (organic acid)
Expon. (7/11 acid)
1.60
PI [(B/D)/psi]
1.40
1.20
1.00
0.80
0.60
0.40
0.20
0
0
0.5
1.0
1.5
2.0
2.5
phi*h (ft)
Fig. 15—Productivity index (PI) vs. storability for wells treated with 13/9% acetic/formic and 7/11% HCl/formic acids (Burgos
et al. 2005).
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Nasr-El-Din et al. (2006) used the same HCl/formic acid blend applied in the early 2000s in the Khuff Formation in Saudi Arabia.
However, they compared the performance of in-situ gelled acid with gelled acid in the long-term production outcomes. Previously, HCl
was crosslinked to produce an in-situ gelled acid that has low viscosity in the surface, which is easier to pump, and builds viscosity
after reacting with the formation, which slows the reaction. In-situ crosslinked acids are usually used to build higher viscosity in situ
that can overcome the HT and yield a slower reaction rate. In the Khuff Formation, more extended wormholes were needed to sustain
higher conductivity from created fractures. Laboratory tests showed that crosslinked HCl alone without formic acid would not be the
best choice in terms of viscosity. The addition of formic acid would stabilize the blend viscosity. Also, temperature stability was better
for the HCl/formic blend than HCl alone for both spent-acid (where viscosity should be high) and live-acid states (where viscosity
should be low). Coreflood tests showed that gelled and in-situ gelled acids would achieve narrower and longer wormholes than HCl
alone because the acid did not spend so fast thanks to the gelling ability to retard the acid system. A total of 11 successful field trials
were conducted using the 15/9% HCl/formic acid blend, where the results showed better performance for the in-situ gelled acid over
the gelled acid because of better leakoff control caused by the high viscosity. Fig. 16 shows the cumulative production of wells treated
with both blends. Flowback samples indicated that both blends maintained the integrity of CR-13 tubing with insignificant corrosion
damage (Nasr-El-Din et al. 2006).
Cumulative Production (MMcf)
45,000
40,000
35,000
30,000
25,000
20,000
Gelled HCL/formic acid
15,000
Crosslinked HCL/formic
10,000
5,000
0
0
50
100
150
200
250
300
350
400
450
Production Time (days)
Fig. 16—Cumulative production for wells treated with gelled acid or in-situ gelled acid (Nasr-El-Din et al. 2006).
Surfactant-based acids started to gain more attention in the mid-2000s because of their superiority over polymer-based acids in
terms of possible unbroken polymer damage. Surfactant-based acids have been mainly used to divert acid treatment because viscosity is
built by the acid reaction with rocks to produce an impermeable gel. This gel can be broken by hydrocarbons after treatment (Li et al.
2011). With more demands for organic acids along with surfactant-based acids, it was essential to understand how these systems would
affect each other when combined together.
Li et al. (2011) investigated the behavior of the apparent viscosity of amidoamine oxide surfactants when organic acids were in solution, using a high-pressure/HT (HP/HT) rheometer. The addition of formic and acetic acids reduced the apparent viscosity of the spent
system regardless of the acid concentration. Nevertheless, the apparent viscosity of amidoamine oxide surfactants has an indirect correlation with organic acid concentrations because more acid strength in solution would cause more reduction in the solution’s apparent
viscosity. Although formic acid did not follow this trend, the apparent viscosity stayed the same in all concentrations. This same effect
of organic acids and formic acid was reported by Cetin and Nasr-El-Din (2017). This effect enabled the system to be used in a broader
range of applications when it is coupled with formic acid. Another finding was noted for the system when used with citric acid because
citrate ions have high stability with calcium ions in solution, which produced calcium citrate salt. The apparent system viscosity
increased using citric acid because of the suspension of calcium citrate in solution.
Li et al. (2011) concluded that organic acids affect the amidoamine oxide surfactant system by rearranging the surfactant micelles
and changing the shape of the micelles. This effect happens because of the high affinity between calcium ions and organic acids, which
left fewer calcium ions in solution to form rod-shaped micelles that build the spent-acid viscosity. To compensate for this effect, the
authors recommended increasing the surfactant loading while considering compatibility, wettability, and rheological tests at downhole
conditions (Li et al. 2011).
Rabie et al. (2012) also studied the effect of formic acid when mixed with HCl to make in-situ gelled acid. The difference between
the Rabie et al. (2012) work and previous authors is the initial HCl concentration. Lower initial concentration of HCl would let the
organic acid contribute more toward spending the acid. Higher initial HCl concentration would dissolve most of the calcium carbonate.
For example, Li et al. (2011) used 20 wt% HCl mixed with an organic acid to spend the acid with calcium carbonate, while Rabie et al.
(2012) used HCl between 0 and 5 wt%.
Coreflood and RDA tests were conducted by Rabie et al. (2012), which showed that a high concentration of formic acid made the
in-situ gelled acid behave like gelled acid, with more PVs needed to break through. This means that more formic acid in the solution
leads to more-viscous live acid because organic acids give more stability to the polymers used in the gelling agent. In the spent case,
more formic acid led to less-viscous gel compared with HCl alone because of the incomplete spending of formic acid. RDA tests
showed that the dissolution rate was associated with the formic acid concentration. When the formic acid concentration was increased,
the dissolution rate decreased. This behavior was also recorded for in-situ gelled acid based on HCl only, where less dissolution was
noticeable when the HCl concentration was increased. In their research, it was recommended not to use formic acid with HCl to generate in-situ gelled acid because the blend would not develop sufficient viscosity when spent because of low HCl concentration and
incomplete spending of formic acid (Rabie et al. 2012).
Carbonate acidizing using formic and acetic acids has become a common practice in stimulation operations, specifically in HT applications. Optimization and elimination of these acids’ drawbacks can extend their application fields. It can be concluded from the previous laboratory and field applications that reaction-product-salt limits prevent these acids from being used at greater than the maximum
concentration before calcium salts are produced. In an attempt to overcome this challenge, Mohamed et al. (2015) mixed formic or
acetic acid with gluconic acid to produce a new salt that has a higher solubility than these calcium salts. They used coreflood tests to
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investigate the optimal ratio between organic and gluconic acids based on molarity. A 15 wt% acetic acid was used alone to coreflood a
limestone core and showed white precipitation (calcium acetate) in the core inlet, as shown in Fig. 17. However, the introduction of gluconic acid at a 1:1 molar ratio of gluconic to acetic acids prevented this precipitation. The optimal ratio dependent on minimum PV to
breakthrough was found to be 1:4 of gluconic to acetic acids. A blend with 15 wt% acetic acid and 12 wt% gluconic acid was free of
precipitation. Formic acid was used up to 12.5 wt% without calcium formate precipitation when 1:7 of gluconic/formic acid molar ratio
was applied (Mohamed et al. 2015).
Fig. 17—Core-inlet face flooded with 15 wt% acetic acid that caused calcium acetate precipitation (Mohamed et al. 2015).
Safari and Fragachan (2018) compared 15% HCl and 10% acetic acid stimulation results on the rock mechanical response in carbonate formations. Unoptimized stimulation fluid can lead to counterbalance the production gain by increasing local porosity. This increase
can reduce the rock strength and can increase the risk of compaction. Safari and Fragachan (2018) used a rock-failure model to conclude that acid type, acid volume, and acid strength have the main effects on alternating the stimulation efficiency. The efficiency was
mainly investigated from the perspective of rock mechanical response.
Safari and Fragachan (2018) showed that acetic acid needed a higher pumping rate (higher volume) to reach the same wormhole
length as the one obtained with HCl. However, the formation-failure response was very high in the case of HCl, specifically at elevated
depths where the stresses are high. Fig. 18 shows a comparison of the compaction-failure profile of the treated wellbore at 250 F
obtained from stimulating with HCl and acetic acid at 9,454 ft. In terms of production, the low reactivity of acetic acid ensured better
production-enhancement sustainability that resulted in better performance than the HCl case. Using acetic acid over HCl showed a difference of 85 B/D, which can justify the surplus cost of acetic acid.
Fig. 18—Comparison of the compaction failure profile of the treated wellbore at 2508F obtained from stimulating with (left) HCl and
(right) acetic acid at 9,454 ft (Safari and Fragachan 2018).
Citric Acid. Citric acid reacts with carbonate through the hydrogen-ion attack that takes place after dissociation of the citric acid in
solution, as shown in Eqs. 18 through 20 (Al-Khaldi et al. 2003).
2Hþ þ CaCO3
Ca2þ þ H2 O þ CO2 ; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð18Þ
H2 AOH þ Ca2þ
CaH2 AOHþ ; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð19Þ
Ca2þ þ 2ðCaH2 AOHþ ÞðaqÞ
Ca3 ðAOHÞ2 ðsÞ # þ4Hþ : . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð20Þ
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As per Eqs. 18 through 20, with more dissociation of hydrogen ions from citric acid and its species, citrate ions become more available for calcium ions to start complexing. As the solution pH increases, free calcium ions tend to stabilize more with the formation of
insoluble calcium citrate salt (Al-Khaldi et al. 2003).
Blauch et al. (2003) reported the use of encapsulated citric acid technology for well-production-enhancement treatment in carbonate
formations. Their results were positive and showed 10-fold increase in reservoir pressure and larger than onefold increase in production.
Moreover, the tested technology showed better results in fracture conductivity than other conventional acidizing systems, such as 20%
HCl and emulsified acid, as can be seen in Fig. 19.
100,000
Conductivity (md-ft)
10,000
250°F, HCl 20%
1,000
250°F, emulsified acid
1 lbm/gal EA equivalent at 250°F
100
No acid, fractured
10
250°F 20/40 cp at 2 lbm/ft2
1
0
2,000
4,000
6,000
8,000
10,000
12,000
Closure Stress (psi)
Fig. 19—Fracture-conductivity-test comparison for different systems (Blauch et al. 2003).
Furthermore, they investigated the encapsulated-acid-technology mechanisms using environmental SEM (ESEM). The capsule,
shown in Fig. 20, is mainly coated with a hydrogenated-vegetable-oil material. The performance is profoundly affected by its melting
point to release the acid. Using the ESEM (heat-stage method), Blauch et al. (2003) studied the heat-release mechanisms. They found
that the melting point is discrete, and acid dispersion is controlled by manipulating the coating composition and diffusional properties
of the coating membrane. Burgos et al. (2004) added that acid release could also be affected by formation temperature and closure
stress, which can crush the capsules and release the acid. Nevertheless, two main mechanisms are followed by the encapsulated acid
technology: the diffusional osmotic exchange release and the thermal melt release of protective coating (Blauch et al. 2003). Figs. 21
and 22 show ESEM photos of the two mechanisms.
Fig. 20—Sand-size crystal that contains citric acid coated with hydrogenated-vegetable-oil material (Blauch et al. 2003).
Fig. 21—ESEM photo of acid being released through a diffusional osmotic exchange mechanism (Blauch et al. 2003).
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Fig. 22—ESEM photo of acid being released thermally (Blauch et al. 2003).
Burgos et al. (2004) reported contrary results when the encapsulated citric acid was investigated in a carbonate formation. They
tested the capsules through conducting a field trial by fracturing a formation with a regular gel pad. Then, the encapsulated citric acid
would be placed inside the fracture to stimulate the fracture faces. However, no significant production enhancement was observed.
The negative results were attributed to three different explanations. The first one was associated with the used fluid pad because it
contained a high concentration of polymers. Breaking these polymers is a common problem in hydraulic-fracturing operations, and
polymers left inside the formation were the first potential source of the low enhancement. The second explanation was the poor etching
generated from the released acid. Acid etching is vital to make functional connectivity between the formation and the wellbore that can
assist hydrocarbon flow. The last explanation was the citric-acid-reaction product with carbonate, calcium citrate, which was produced
because of the high stability between calcium and citrate ions (Al-Khaldi et al. 2003). The low solubility of the salt imposes formation
damage within the created channels that would plug hydrocarbon pathways to the wellbore (Burgos et al. 2004).
Different tests were performed to investigate the poor results and to understand the contribution of calcium citrate toward the poor
results. Limestone rocks were placed with 10% citric acid in an autoclave under 250 F and 1,200 psi. Laboratory results showed that
acid spending was low because the pH values of different samples increased from 1.7 to 2.7 after 24 hours. The solution pH curves of
the 13/9% acetic/formic acid and the 10% citric acid are plotted vs. time in Fig. 23. Calcium citrate was not noticed in the liquid samples, but white precipitation was noticeable in the examined limestone rocks. On the basis of such, Burgos et al. (2004) concluded that
using citric acid can generate severe formation damage that can undermine the use of citric acid in stimulation operations (Burgos
et al. 2004).
4.5
pH after 24 hours: 4.2
4.0
13/9% acetic/formic
3.5
pH
3.0
pH after 24 hours: 2.7
2.5
10% citric
2.0
1.5
1.0
0
100
200
300
400
500
Time (minutes)
Fig. 23—Solution pH values of 13/9% acetic/formic acid blend and 10% citric acid vs. time (Burgos et al. 2004).
Al-Khaldi et al. (2003) investigated the citric acid reaction with carbonate rocks by understanding the geochemical reaction paths
and calcium citrate precipitation. According to Al-Khaldi et al. (2003), the pH value of a solution is a significant factor for calcium
citrate precipitation. By mixing calcium carbonate with different strengths of citric acid solutions, calcium citrate precipitation was
found to start depositing between pH values of 2.7 and 3.2, as shown in Fig. 24. Iron(III) ions within solution could prevent calcium
citrate precipitation because of the high stability constant between iron(III) and citrate ions. Magnesium ions in solution did not have a
positive effect on calcium citrate precipitation (Al-Khaldi et al. 2003).
In this research, it was concluded from coreflood and fracture-conductivity tests that CO2 has a positive effect on preventing citrate
ions from being free to complex with calcium ions when pore pressure was 900 psi. Tests conducted in atmospheric conditions showed
a negative effect of citric acid on limestone and dolomite cores caused by CO2 gas being released out of solution. On the other hand,
tests conducted in in-situ conditions showed permeability and conductivity improvement without calcium citrate precipitation
(Al-Khaldi et al. 2003).
Intending to determine the citric acid diffusivity from the calcite reaction, Al-Khaldi et al. (2010) used an RDA for various citric
acid strengths. Citric acid and calcite reaction was described by three steps: reactants’ transfer to the calcite surface, reaction at the calcite surface, and transfer of reaction product away from the surface (Al-Khaldi et al. 2010). The slowest step determines the overall
reaction. The three steps are shown in Fig. 13.
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Concentration of Ca in
Solution (mg/L)
6000
4.0 wt% citric acid
5000
4000
7.6 wt% citric acid
3000
2000
2.1 wt% citric acid
5.8 wt% citric acid
1000
0
2.0
2.2
2.4
2.6
2.8
3.0
3.2
3.4
Equilibrium pH
Fig. 24—Calcium-ion concentration vs. solution pH values for different initial citric acid strengths (Al-Khaldi et al. 2003).
Calcite-Dissolution Rate (gmol/cm2/s)
The results showed that citric acid could not effectively dissolve calcium carbonate without precipitating calcium citrate. At 2.1 and
7.6 wt% of citric acid, calcium citrate can precipitate when there are 5,000 and 2,500 ppm of free calcium ions in solution, respectively.
Although the effective-diffusion coefficient of citric acid becomes greater with a higher initial acid concentration, calcium citrate
would precipitate more because of the presence of more calcium ions in the solution (Al-Khaldi et al. 2010).
RDA tests were performed at 1,000 psi and 50 C, and the reaction regime was dependent on initial acid strength. At low acid concentration, the dissolution rate was mass-transfer limited regardless of rotation speed. For high acid concentration, the dissolution rate
was mass-transfer limited up to 500 rev/min. At higher rotational speed, the dissolution rate was surface-reaction limited because of calcium citrate acting as a barrier between the acid solution and the calcite. Calcite-dissolution rates derived from different citric acid solutions vs. rotational speed are shown in Fig. 25. It was found that the dissolution rate did not increase significantly with increasing acid
strength, except between 1 to 2 wt%. This finding is cost effective because there is no need to use a high concentration of citric acid
where 2 wt% would achieve the same results of 7.5 wt% with less calcium citrate precipitation (Al-Khaldi et al. 2010).
×10–7
12
Initial citric acid concentration (wt%)
1
2
10
5
7.5
8
6
4
2
0
0
2
4
6
8
10
12
ω 0.5, S –0.5
Fig. 25—Calcite-dissolution rate for different citric acid solutions vs. rotational speed (Al-Khaldi et al. 2010).
In a goal to increase calcium citrate solubility, Mohamed et al. (2015) used gluconic acid as a chelating agent to be mixed with a
citric acid solution. Gluconic acid effects were evaluated using coreflood tests only, where the results showed that gluconic acid would
not add any value to calcium citrate solubility because of the high stability between calcium and citrate ions.
Lactic Acid. Lactic acid reaction with carbonate has been a subject of study for many researchers (Lund et al. 1973; Plummer et al.
1978; Sjöberg and Rickard 1984). This reaction can be classified into three subreactions: hydrogen-ion reaction, water reaction, and carbonic acid reaction with calcite, shown in Eqs. 21 through 23, respectively.
Hþ þ CaCO3 $ Ca2þ þ HCO
3 ; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð21Þ
H2 O þ CaCO3 $ Ca2þ þ HCO
3 þ OH ; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð22Þ
H2 CO3 þ CaCO3 $ Ca2þ þ 2HCO
3 : . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð23Þ
The hydrogen-ion reaction is represented by Eq. 21 and is the dominant reaction at a low pH value. Eq. 22 represents the waterdissolution reaction that is dominant at a high pH value. The last reaction, Eq. 23, represents the carbonic acid reaction with calcite.
It is only considered when CO2 partial pressure is higher than 0.1 atm, and the solution pH value is higher than 5. Because this study is
focusing on the reaction of lactic acid with calcite under acidic solutions with a low pH value, the hydrogen-ion reaction is the only
one considered.
Lactic acid has received less attention than other organic acids because of the solubility of its reaction product of calcium lactate.
Similar to citric acid, lactic acid is an a-hydroxycarboxylic acid, but with much higher solubility than citric acid in terms of their
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calcium salts (Nasr-El-Din et al. 2009; Rabie et al. 2015). Salt solubilities are shown in Table 1. Nasr-El-Din et al. (2007) studied lactic
acid reaction with carbonate in solid forms. The main goal was to address limitations with other conventional acid systems, including
corrosion, reaction-product precipitations, and acid-leakoff rate. The solid beads release lactic acid under exposure of heat and water,
which can be placed inside a fractured formation to add more etching onto the fracture faces.
Different laboratory tests were performed during their study to investigate the hydrolysis of the solids, solubility of limestone cores
caused by the released acid, compatibility, core etching, and corrosion tests. The results showed that solids hydrolysis is highly dependent on temperature because the solids could release the acid in less than 1 hour at a temperature of 300 F, while it needed 12 hours at
200 F. A borate gel was also tested to suspend the solids, but the final solution pH value was noted to be higher than that in the case of
using deionized water. The high pH was explained by the fact that the borate gel is a high-pH solution, and the solids would consume a
portion of the acid to increase the solution pH value. Corrosion tests showed that the released lactic acid would require a low concentration of an organic acid CI to keep the corrosion rate at less than the acceptable value. A high etching pattern was noticed when the
solids were placed inside a vessel with a core in the bottom (Nasr-El-Din et al. 2007). The solid beads are shown in Fig. 26.
Fig. 26—Solid beads of lactic acid (Nasr-El-Din et al. 2007).
Laboratory results were positive enough to apply the new solid beads in a deep gas formation in Saudi Arabia where a well was fractured using a high-pH borate gel (Nasr-El-Din et al. 2009). Solid beads were added to the fracturing fluid on the fly to carry it inside
fractures. The well was shut down for 24 hours to allow the appropriate time for the acid to be hydrolyzed. The treatment evaluation
was conducted by means of flowback-sample analysis and production rate. After 9 months of production, results were below expectations because the gas rate only improved from 18 to 19.5 MMcf/D, and the skin factor was lowered from 7.5 to 5.5. Nasr-El-Din et al.
(2009) attributed this low improvement to four reasons. First, lactic acid needed more time and more water to reach complete hydrolysis
because the flowback samples showed higher calcium- and magnesium-ion concentrations in the second day of treatment. Second, the
released lactic acid reacted with mill scale, which produced iron(II) lactate precipitation. Third, the beads were not placed entirely
inside the fracture; a portion was left in the wellbore. Fourth, lactic acid was partially consumed with the fracturing fluid to increase the
pH value (Nasr-El-Din et al. 2009).
Rabie et al. (2014) studied the reaction of lactic acid with carbonate using 1, 5, and 10 wt% concentrations and limestone. Dissolution rates of the mentioned concentrations were measured at 1,500 rev/min and reported in Table 3. A kinetic study was also reported
for 5 wt% lactic acid on limestone and dolomite cores using the RDA test. Their results showed that at low temperatures, the reaction is
limited by the surface reaction. At HTs and low rotation speeds (less than 1,500 rev/min), the reaction is limited by surface reaction and
mass transfer. However, at HTs and high rotation speeds, the reaction was limited by the surface reaction (Rabie et al. 2014). In this
study, diffusivity values for 5 wt% lactic acid were reported at 200 and 250 F. Table 4 shows the reported diffusivities compared with
other acid values.
Lactic Acid
Concentration
Dissolution Rate (gmol/cm2s)
1 wt%
1.4010–6 6 9.2910–8 at 150 F
5 wt%
4.5310–6 6 1.1010–7 at 200 F
10 wt%
1.1510–5 6 2.0810–7 at 250 F
Table 3—Dissolution rates of 1, 5, and 10 wt% lactic acid from
carbonate reaction (Rabie et al. 2014).
Rabie et al. (2014) mixed lactic acid with gluconic acid to improve the reaction of salt solubility. According to their study, the new
mixture produced a soluble salt with solubility of 400 compared with 79 mg/L for calcium lactate salt. Acid-dissolving-capacity tests
showed that a 1:1 molar ratio between lactic and gluconic acids was the optimal ratio to dissolve the maximum amount of calcium carbonate and to avoid calcium lactate precipitation. Other ratios showed either precipitation of calcium lactate or low dissolving capacity.
RDA tests were also performed on lactic acid, gluconic acid, and a mixture of both at a 1:1 ratio. Results showed that the mixturedissolution rate was comparable with that value for lactic acid. Table 5 shows the resulted dissolution rates. Coreflood tests confirmed
the risk of using lactic acid alone in high concentration because calcium lactate appeared in effluent samples when 20 wt% lactic acid
was used to coreflood a limestone core. The tests also proved the advantage of coupling lactic acid with gluconic acid because no precipitation occurred. The most optimal molar ratio in terms of PVs needed to break through was 1:1.
Alhamad (2020) extended the previous work on lactic acid and investigated its performance with limestone cores when deionized
and synthetic seawater were in use. The main results were obtained from solubility, corrosion, and coreflood tests. Calcium lactate precipitation was found to be a limiting drawback of using lactic acid at concentrations of 9 wt% or higher. However, this precipitation
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was avoided by coupling lactic acid with gluconic acid at 1:1 molar ratio. The solubility test showed that increasing lactic acid molarity
over gluconic acid molarity would be associated with an unbalanced ratio between lactate and gluconate ions within the acid solution.
Therefore, gluconate ions would have less chelation ability to completely prevent calcium lactate precipitation. Fig. 27 shows the main
results from Alhamad (2020) on the solubility test for a lactic/gluconic acid mixture.
Acid
Concentration
(M)
Temperature
( C)
Diffusivity
(cm2/s)
Reference
HCl
0.50
25
3.610–5
Kung (1998)
HCl
0.56
93.33
8.310–5
Conway et al. (1999)
HCl
0.56
121
1.110–4
Conway et al. (1999)
–5
Formic
0.50
25
3.010
Kung (1998)
Acetic
0.50
25
1.1510–5
Vitagliano and Lyons (1956)
Citric
0.25
25
6.4610–6
Muller and Stokes (1957)
Citric
0.26
40
6.810–6
Al-Khaldi et al. (2010)
–6
Al-Khaldi et al. (2010)
Citric
0.39
50
4.510
Lactic
0.56
26.67
5.9110–7
Rabie et al. (2014)
–5
Lactic
0.56
93.33
1.1010
Rabie et al. (2014)
Lactic
0.56
121
1.4710–5
Rabie et al. (2014)
Table 4—Diffusivity values of different acids.
Acid
Dissolution Rate
(gmol/cm2s)
Lactic acid
1.7673510–6
Gluconic acid
9.6110710–7
Lactic/gluconic acid at 1:1 molar ratio
1.4263110–6
Table 5—Dissolution rates of lactic acid, gluconic acid, and mixture
of both acids (Rabie et al. 2014).
Calcium Ions vs. Time
Ca Ions (ppm)
45,000
40,000
35,000
Precipitation
results of calcium
lactate deposition
A drop in calcium ions as a
reflection of calcium lactate
deposition
30,000
25,000
20,000
15,000
10,000
5,000
9 wt% lactic acid
14 wt% (1:1) lactic acid/gluconic acid
0
0
20
40
60
80
100
120
Time (minutes)
Fig. 27—Calcium-ion concentrations vs. time for the reaction of 9 wt% lactic acid and 14 wt% (1:1) of lactic and gluconic acids with
crushed core at 708F and 1,000 psi (Alhamad 2020).
Preparing the two-acid mixture of lactic and gluconic acids with synthetic seawater (sulfate concentration of 4,300 ppm) was associated with calcium sulfate precipitation. Calcium and sulfate ions were free in solution to complex together and to produce the lowsolubility compound of calcium sulfate. However, the Alhamad (2020) investigation showed that equilibrium time for calcium and
sulfate within the two-acid mixture was longer than the case of 15 wt% HCl solution. Gluconate ions within the two-acid mixture were
suspected to be the reason behind the delay in calcium sulfate precipitation. Fig. 28 shows the sulfate-ion concentrations vs. time for
the case of the two-acid mixture and HCl solutions.
Alhamad (2020) performed a sensitivity analysis on the lactic/gluconic acid mixture using a coreflood apparatus. Injection rate, test
temperature, total acid concentration, and diluted-water type were the main parameters in the sensitivity analysis. The analysis showed
the risk of a face-dissolution pattern when the two-acid mixture was pumped at a low injection rate that was less than the optimal value.
In addition, a ramified dissolution pattern was noticed when a high injection rate was used. Manipulating other parameters produced
results that followed the same behavior noticed when an HCl solution is used with slower acid-rock reactivity. Corrosion tests in
Alhamad (2020) work showed low corrosion rates of the two-acid mixture when compared with HCl solutions. However, the corrosion
damage was high enough to dissolve a considerable concentration of iron(II) ions to complex with lactate ions. This bonding caused
iron(II) lactate precipitation, which was avoided by minimum loading of a CI that was suitable for organic acids. Fig. 29 shows iron-ion
concentrations vs. time and iron(II) lactate precipitation.
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Sulfate-Ion Concentrations vs. Time
4,500
3,500
27 wt% (1 day)
27 wt% (10 days)
15 wt% HCI (1 day)
2,500
1,500
500
–500
0
40
20
60
80
100
120
Time (minutes)
Fig. 28—Sulfate-ion concentrations vs. time for the case of the two-acid mixture and HCl solutions (Alhamad 2020).
Iron-Ion Concentrations in 9 wt% Lactic Acid Solutions
Iron Ions (ppm)
25,000
Precipitation
results of iron
20,000 lactate deposition
A drop in iron ions as a
reflection of iron lactate
deposition
15,000
200°F
300°F
10,000
5,000
0
0
30
60
90
120
150
180
210
240
Time (minutes)
Fig. 29—Iron-ion concentrations in solution for 9 wt% lactic acid solution that was tested at 200 and 3008F at 1,000 psi for 4 hours
(Alhamad 2020; Alhamad and Miskimins 2020).
Sandstone Stimulation
The most common acid for sandstone-stimulation operations is mud acid, which is a blend of hydrochloric and hydrofluoric acids. The
primary acid is HF, which reacts with aluminosilicate. According to Martin (2004), sandstone stimulation is obtained by removing formation damage, such as fines, and drilling damage. Along with HF acid, HCl acid is required to keep the fluid pH low, thus blocking
byproduct precipitation and removing carbonate contents.
It is essential to understand the reaction behavior of HF/HCl or organic acid with sandstone to identify the source of formation
damage. HF-based acids follow three reaction steps when reacting with sandstone. The primary reaction releases aluminum ions
because of the acid interaction with aluminosilicate (Shuchart 1995; Gdanski 2000), as shown in Eq. 24.
þ
þ Mþ þ water: . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð24Þ
ð5þxÞHF þ M Al Si þ ð3 x þ 1ÞHþ ! HSiF5 þ AlFð3xÞ
x
Silica gel [(Si)OH4] precipitates during the secondary reaction because of the silicon (HSiF5) production in the primary reaction
(Gdanski 1999). The secondary reaction is shown in Eq. 25, where x is the average F/Al ratio and M represents cations such as Na and
K. Gdanski (1999, 2000) found that the secondary reaction goes to completion quickly at temperatures higher than 125 F. HCl or
organic acids are coupled with HF to avoid silica gel precipitation. However, HCl is limited to removing carbonate materials in claysensitive formations. In addition, HCl should not be used with all clay minerals at a temperature higher than 300 F because of destabilizing clay structures (Buijse et al. 2004; Chang et al. 2008; Zhou and Nasr-El-Din 2013).
þ
x=5 HSiF5 þ M Al Si þ ð3 x þ 1ÞHþ þ H2 O ! AlFxð3xÞ þ Mþ þ silica gel: . . . . . . . . . . . . . . . . . . . . . . . . . . . ð25Þ
A tertiary reaction is associated with an increase in the solution pH value and an increase in aluminum concentration. Hydroxyl aluminum fluoride precipitation occurs after this reaction because of the spent-acid dilution by formation brine and the reaction with formation minerals (Gdanski 1998). The tertiary reaction is shown in Eq. 26, and hydroxyl aluminum fluoride processes are shown in
Eqs. 27 and 28. The main three reactions associated with sandstone stimulation are shown in Fig. 30, which shows that formation
damage caused by secondary and tertiary reactions can plug formation pores (Ueda et al. 2019).
AlFx þ M Al Si þ Hþ ! AlFy þ Mþ þ SiO2 xH2 O; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð26Þ
AlFx þ OH ! AlFy ðOHÞ3x ; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð27Þ
AlFx þ MCO3 ! AlFy ðOHÞ3x þ M2þ þ CO2 : . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð28Þ
The success of the treatment is associated with choosing the appropriate fluid to stimulate with minimal formation damage. Spending rate, high corrosion rate, and HCl-sensitive clays such as illite were reported to be different causes of developing retarded HF acid
with organic acids instead of the conventional mud acid. Al-Dahlan et al. (2001) reported that the reaction rate of the mud acid was less
than the acceptable value and byproduct precipitation was possible. They conducted many tests and experiments to introduce organic
acid with HF acid because the mixture would have the nature of retardation, cause a low corrosion rate, and be compatible with sensitive formations. Still, the two significant limitations of organic acids can be found in sandstone stimulations, including the concentration
limit and the completion of the reaction (Nasr-El-Din et al. 2007; Chang et al. 2008).
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Primary Reactions
HF(aq) + mineral(s) + HCl(aq) → AIFx(aq) + H2SiF6(aq)
Potential precipitates: CaF2, MgF2
Secondary Reactions
H2SiF6(aq) + mineral(s) + HCl(aq) → Si(OH)4(s) + AIFx(aq)
Potential precipitates: Si(OH)4, K2SiF6, Na2SiF6
Tertiary Reactions
AIFx(aq) + mineral(s) → AIFy(aq) + Si(OH)4(s); x > y
Potential precipitates: Si(OH)4, AI(OH)3
Distance from Wellbore
Fig. 30—Formation damage caused by secondary and tertiary reactions in sandstone stimulation (Udea et al. 2019).
One of the first sandstone organic acid treatments was conducted by Abrams et al. (1983), in which they used methyl formate to generate formic acid in situ that would react with an injected ammonium fluoride to produce HF in-situ. This technique was designated to
be applied in the East Bay Area in Louisiana, USA, where mud-acid treatments were associated with clay and fines migrations. Regular
mud acid can only remove clay damage within a few inches around the wellbore. Such a retarded acid would push HF further inside the
formation to increase permeability. Field results showed that methyl formate with an ammonium fluoride mixture dissolved more suspended fines and clays, causing an increase in oil and water production over a long period of time. The production data for conventional
and organic/HF treatments are shown in Fig. 31.
Self-generating acid treatments
Average BOP Increase Per Well
60
50
40
30
Range of average result of various
conventional treatments
20
Self-generating mud-acid
treatments, first 28 wells, 1973
conventional treatments 102 wells 1973
10
0
0
1
2
3
4
5
6
7
8
9
10
11
12
Months Since Job
Fig. 31—Comparison of acid-stimulation techniques for East Bay Field (Abrams et al. 1983). BOP 5 barrels of oil per day.
Shuchart and Gdanski (1996) pointed out the limitation of using mud acid in sandstone formations that contain a considerable
amount of chlorite and zeolite. Strong acids such as HCl can destabilize the alumina layers, causing a collapse in the minerals’ structure.
A collapse in clay structure would result in producing silica gel and fines migration. However, clay problems can also occur by using
organic acids. Organic acids have a slow ionization process that causes low ion exchange with clays. Ammonium chloride has to be
pumped with organic acids to overcome this problem (Shuchart and Gdanski 1996; Yang et al. 2012). In addition, organic acids tend to
increase the F/Al ratio in solution, which is an invitation for aluminum fluoride (AlF3) precipitation. Shuchart and Gdanski (1996) used
a new organic system that can alter the aluminum fluoride equilibria, which can minimize the precipitation amount in solution. Field
trials in the Gulf of Mexico showed an increase in production rate caused by injecting the new system (Shuchart and Gdanski 1996).
Rogers et al. (1998) ran several coreflood tests on composite samples from the Ewing Bank Field to optimize a stimulation fluid.
The used mud acid (10% HCl and 1.5% HF) showed a 74% reduction in productivity on coreflood tests. X-ray diffraction and SEM
tests were used to determine the reasons behind the permeability reduction. Test results showed that HCl was very reactive to the core
samples, specifically to zeolite, which caused fines migration and silica gel precipitation. Different formulations of citric acid with HF
were proposed to achieve a proper stimulation solution. 10% citric acid and 1.5% HF with CI showed the optimal result because it successfully dissolved zeolite with minimum silica precipitation. The authors did not report any AlF3 precipitation during laboratory testing. The formulation was applied in the field, and the drawdown increased from 4.78 to 12.73 BOPD/psi (Rogers et al. 1998).
The F/Al ratio had been a challenge in treatment plans for injection wells in the Chunhua Oil Field in eastern China, where significant formation damage occurred caused by clays and fines (Wang et al. 2000). Wang et al. (2000) used a mixture of acetic acid with HF
to decrease AlF3 precipitation by altering the aluminum fluoride equilibria. Previous field treatments with regular mud acid resulted in a
slight increase of injectivity followed by a sharp decrease within 1 month because of fines migration, secondary precipitation, and
incompatibility issues. The new formulations passed all laboratory tests and were applied in a damaged well. Results showed that
the skin factor dropped from 13 to –0.21, and the injection rate increased from 0 to 100 m3/d (Wang et al. 2000).
Fluoride-based precipitation is the main challenge for organic acid use with HF in sandstone stimulation. This problem was noticed
by Yang et al. (2012) in a study where formic acid was coupled with HF to stimulate sandstone cores. The study showed that formic
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700
5000
5 wt%
NH4CI
9 wt% FA + 5 wt%
NH4CI
9 wt% FA 9 wt% FA
+1 wt% HF +5 wt%
NH4CI
5 wt% NH4CI
Concentration (mg/L)
Pressure Drop Across the Core (psi)
acid could be helpful to dissolve carbonate material in HT applications but also caused a formation damage of calcium fluoride precipitation. Fluoride-based precipitation was reduced when a low concentration of HF was used. Fig. 32 shows how a mixture of formic acid
and HF damaged a Bandera Sandstone core when HF concentration was 1 wt%.
600
500
4000
5 wt% NH4CI 9 wt% FA
+5 wt%
NH4CI
3000
9 wt% FA 9 wt% FA
+1 wt% HF +5 wt%
NH4CI
5 wt% NH4CI
Ca
Fe
2000
Mg
1000
Si
AI
400
0
2
4
6
8
10
12
14
16
18
20
0
0
2
Cumulative Volume Injected (PV)
4
6
8
10
12
14
16
18
20
Cumulative Volume Injected (PV)
Fig. 32—Coreflood result of a Bandera core treated with a mixture of 9 wt% formic acid and 1 wt% HF. On the left, the pressure profile showed an increase in the differential pressure after pumping the mixture. On the right, the calcium-ion concentration
decreased after 10 PV because of calcium fluoride precipitation (Yang et al. 2012).
Al-Harbi et al. (2012) investigated aluminum and iron(III) ions’ effects on organic acids and HF solution performance. Inductively
coupled plasma and X-ray diffraction tests were used to track elemental composition in solution and identify precipitated solids. Solutions were prepared using formic, acetic, and citric acids with different amounts of dissolved aluminum and iron in solution. The solution pH value was manipulated through the addition of sodium hydroxide. Organic acids were prepared along with HF using two ratios:
10 wt% of organic acid and 0.5 wt% of HF, or 10 wt% of organic acid and 1.5 wt% of HF.
Iron(III)-based precipitation was not an issue at low-pH solutions despite the iron concentration in solution. However, when the
solution was neutralized to a pH value of approximately 4.3, Na3FeF6 precipitation was noticeable. Aluminum-based precipitation was
dependent on the F/Al ratio. Organic acids with 1.5 wt% of HF would cause aluminum fluoride precipitation because they exceeded a
F/Al ratio of 1.9, while organic acids with 0.5 wt% of HF did not cause any precipitation (Al-Harbi et al. 2012), as shown in Fig. 33.
Al-Harbi et al. (2013) investigated the organic acid mixture with HF acid ability to overcome mud-acid challenges. The idea was to
slow down HF generation (shown in Eqs. 29 and 30) through dissolving ammonium bifluoride with organic acids. In mud acid, HCl dissociates much faster than organic acids, causing fast HF generation and inappropriate retardation.
HA ! Hþ þ A ; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð29Þ
Hþ þ NH4 HF2 ! NHþ
4 þ 2HF: . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð30Þ
Fig. 33—10 wt% formic acid with (left) 0.5 wt% HF and (right) 1.5 wt% HF (Al-Harbi et al. 2012).
The investigation was performed by testing the dissolution of kaolinite, illite, and chlorite with organic acids and HF using coreflood
tests on Berea Sandstone cores at 250 F. The investigation concluded that the initial concentration of HF plays a significant role in
terms of free fluoride ions, which feed aluminum fluoride precipitation. The formic/acetic acid mixture showed poor results in terms of
permeability improvement, while citric acid at 10 wt% showed better results only in a low HF acid concentration of 0.5 wt%; the permeability improved by 19%. However, an organic acid mixture with HF did not yield better results than mud acid because HF concentration was very limited to low strength to avoid calcium and fluoride precipitation (Al-Harbi et al. 2013).
Akaighe et al. (2016) determined HF strength generation from organic acids using a fluoride ion-selective electrode. This method
helped to evaluate the needed amount of fluoride source (in this case, ammonium fluoride) to generate 1.5% HF coupled with the tested
organic acid. The results showed that acetic acid required 3.96 wt% of ammonium fluoride to reach the desired level of HF, followed by
formic acid with 3.67 wt%. Among all tested organic acids, citric and etidronic acids required the least amount of ammonium fluoride,
with amounts of 3.38 and 3.24%, respectively. Because conventional mud acid can dissociate entirely in water, it only needs 3.02%.
However, coreflood tests, at 150 F using Berea cores, showed that citric acid was the best choice to be coupled with 1.5% HF because
citrate ions can complex with aluminum and iron ions (Akaighe et al. 2016). Besides citric acid, other organic acids did not show effective stimulation results when mixed with HF because of fluoride-based precipitation. Table 6 shows the coreflood results.
Shah et al. (2017) reported a field trial in an onshore field in India where they stimulated a sandstone reservoir with organic acid.
Two case studies corresponding to two wells were reported. The use of an HF/acetic acid mixture resulted in a 500 and 325% increase
in injectivity, respectively, in the two reported wells. These results outperformed the reported HCl/HF results for the same field.
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Treatment Fluid
Permeability
Enhancement
12 wt% acetic acid/1.5 wt% HF
–1.90%
12 wt% formic acid/1.5 wt% HF
–1.10%
15 wt% citric acid/1.5 wt% HF
3.70%
10 wt% HCl acid/1.5 wt% HF
–5.60%
Total Pages: 27
Table 6—Coreflood results from using organic acids/HF and mud
acid (Akaighe et al. 2016).
A study was performed by Ueda et al. (2019) to stimulate a volcanic tuffaceous sandstone formation in Kita-Akita Field in northern
Japan, where conventional mud acid was limited because of high analcime content. Coreflood and compatibility tests indicated that the
use of formic acid would cause similar damage to the use of HCl because of analcime-based damage. When acetic acid was coupled
with HF in 10/0.5 wt% concentrations, permeability enhancement was noticeable. The low dissociation of acetic acid reduced the acidanalcime alteration, which was reflected in a better retained permeability than the case of using formic or formic/acetic acid mixtures.
Other Applications
In addition to stimulation applications, organic acids are used for other applications, as described in the following subsections.
Iron Control. Iron control is a critical practice in reservoir-stimulation operations because of the tendency of ferric and ferrous ions to
complex with other ions in solution and produce soluble and insoluble products. Insoluble products can invade and precipitate inside
the treated formation and plug hydrocarbon pores and flow paths near the wellbore. Iron can exist in two forms, either Feþ2 or Feþ3; the
former appears in a reduction environment, and the latter appears in an oxidation environment coupled with air (Taylor et al. 1999).
The Feþ2 state is found more in reservoir conditions, specifically within sour wells. According to Crowe and Minor (1985) and Crowe
(1986), Feþ2 is found in spent acid five times more often than Feþ3 depending on the well and formation conditions. This is a preferable
ratio because Feþ3-based precipitation can occur at pH values greater than 1, while Feþ2-based precipitation occurs at pH values greater
than 6 unless hydrogen sulfide is present (Crowe and Minor 1985; Crowe 1986; Taylor et al. 1999).
Table 7 shows the most common iron-based precipitations that occur during reservoir stimulation. Ferric hydroxide can precipitate as
fast as the acid spends and pH values reach 1 (Crowe and Minor 1985; Crowe 1986; Taylor et al. 1999). Ferrous hydroxide can be precipitated at a pH value of 6 unless hydrogen sulfide is present, because the insoluble ferrous sulfide deposition would be precipitated at a pH
value of 2 (Crowe 1986; Taylor et al. 1999). Elemental sulfur can be precipitated in sour wells because of the reduction of Feþ3 to Feþ2
(Smith et al. 1969). McLeod (1984) mentioned that there is a possibility of ferrous carbonate precipitation when Feþ2 is free in solution.
Compound
Chemical Formula
Environment
Reference
Ferric hydroxide or iron(III) hydroxide
FeðOHÞ3
Greater than pH of 1
Taylor et al. (1999)
Ferrous hydroxide or iron(II) hydroxide
FeðOHÞ2
Greater than pH of 6
Smith et al. (1969)
Ferrous sulfide or iron(II) sulfide
FeS
Greater than pH of 2
–
In sour wells
Smith et al. (1969)
–
–
Sulfur
S0
Reduction
–
In sour wells
Crowe (1986)
–
–
Ferrous carbonate or iron(II) carbonate
FeðCOÞ3
Spent acid
McLeod (1984)
Table 7—Most common iron-based precipitations and depositional environment.
Each stage of reservoir-stimulation operations can prompt iron ions in solution from acid-transportation pipes and tanks to the
treated wellbore tubulars and formation (Taylor et al. 1999). An appropriate procedure is highly recommended to avoid iron contamination caused by rust destruction in acid tanks (Crowe and Minor 1985). A contaminated solution could have either Feþ2 or Feþ3 with a
higher level of Feþ3 because of the oxidation of Feþ2 from dissolved oxygen (Crowe 1986). During acid injection, mill scale and corrosion products can be easily dissolved by acid to add more iron in solution. Nonetheless, this can be protected by a precleaning stage and
pickling the tubulars to wash out these products (Smith et al. 1969; Crowe and Minor 1985; Crowe 1986; Taylor et al. 1999).
Because substantial risk can come from iron-based precipitations, researchers have investigated different agents to keep iron complexes soluble in solution (Smith and Mancillas 1969; Dill and Fredette 1983; Ewing et al. 1983; Crowe and Minor 1985; Crowe 1986;
Taylor et al. 1999; Assem et al. 2013; Zhang et al. 2016; Ramanathan and Nasr-El-Din 2019). Smith and Mancillas (1969) mixed different
chelating agents in acid solutions to produce water-soluble complexes and to stabilize iron(II) and iron(III) ions at high pH values. Dill
and Fredette (1983) studied chelating-agent performance in sour conditions in the presence of oxidizing agents. Their results indicated
that the chelating-agent performance was limited by solubility and temperature and can be affected by sour and oxidizing conditions.
During the 1980s, organic acids were frequently studied as iron-control agents to overcome challenges associated with other chelating
agents. Ewing et al. (1983) tested citric and acetic acids to chelate up to 50,000 ppm of Feþ3 ions in HCl solutions between 80 and 170 F.
Compatibility results showed that 5 gal/1,000 gal of citric acid alone could chelate ferric ions below 2,000 ppm in 7.5% HCl. However,
after the excess loading of ferric ions and citric acid are in solution, iron(III)-based and calcium citrate precipitations occurred within the
first testing hour. Adding nitrilotriacetate enhanced the citric acid performance because the two chemicals were able to chelate 5,000 ppm
of ferric ions in 20% HCl. Acetic acid was not an efficient iron-control agent, but it acted as a buffering agent to keep the solution pH
value at a low level to prevent ferric-based precipitations. At 50,000 ppm of ferric ions, 20 gal/1,000 gal of acetic acid with 30 gal/1,000 gal
of both citric acid and nitrilotriacetate were enough to keep iron(III) ions in the 20% HCl solution (Ewing et al. 1983).
Crowe and Minor (1985) and Crowe (1986) identified sources of Feþ2 and Feþ3 where the latter can be a severe problem when the
acid is spent. These papers showed the importance of understanding the treated well conditions before applying the treatment. Sour
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wells were identified as free ferric hydroxide conditions because of the presence of hydrogen sulfide. However, ferrous hydroxide precipitation would occur. Glass-bottle tests were performed to examine different iron-control agents with 15% HCl. Citric acid was the
best agent for all formation conditions, including sour conditions, whereas acetic acid was the worst. Acetic acid performed poorly
once the temperature increased, while other agents performed better with higher temperatures, as can be seen in Fig. 34.
6.0
5.5
Fe in Solution (mg/L x 1000)
5.0
Stabilizer concentration = 12 000 mg/L
Initial Fe(III) concentration = 10 000 mg/L
4.5
4.0
3.5
NTA
3.0
2.5
Citric acid
2.0
1.5
Tetrasodium EDTA
1.0
0.5
100
Acetic acid
150
200
250
300
350
400
Temperature (°F)
Fig. 34—Performance of different iron-control agents in terms of iron(III) ions in solution vs. temperature (Crowe and Minor 1985).
NTA = nitrilotriacetate; EDTA 5 ethylenediamine-tetra-acetic acid.
Taylor et al. (1999) examined ferric hydroxide precipitation in the presence and absence of iron-control agents. The new finding in
their work was the solution pH value at which ferric hydroxide started to precipitate. At a pH value of 2, ferric hydroxide started to precipitate in see-through-cell tests without the presence of any chelating agent, while it was suspected to precipitate at a pH value of 3 in
preceding research. With the addition of citric acid, ferric hydroxide did not precipitate, but calcium citrate started to precipitate once
excessive loading of citric acid was added. Nevertheless, the addition of acetic acid with low loading of citric acid showed a positive
effect on chelating Feþ3 ions in solution.
Assem et al. (2013) examined the effect of iron(III)-based precipitation on well productivity during matrix-acidizing operations. The
conducted experiments measured the pressure drop across a limestone core in a case where there were iron(III) ions in the HCl solution
vs. a case where there were no iron(III) ions present. The presence of iron(III) ions increased the pressure drop across the core, indicating formation damage, which was later confirmed by computed-tomography scan images, where most of the damage was seen on the
core’s inlet face. A higher acid-injection rate showed a reduction in the amount of iron(III) precipitates.
Organic acids have been used to dissolve iron(II) sulfide scales as standalone systems or by mixing them with other compounds.
Ramanathan and Nasr-El-Din (2019) studied the effectiveness of iron(II) sulfide dissolution using various organic acids and introduced
new mixtures that improved the dissolution efficiency. Among the tested standalone acid systems, maleic acid followed by citric acid
showed to be the best in terms of dissolution capacity (solubility of 10.6 and 7.6 g/L, respectively), whereas the lactic acid was the
worst (4.3-g/L solubility). The newly tested mixtures involved adding synergists such as potassium oxalate and potassium citrate to
Na2EDTA, where iron(II) sulfide solubility was enhanced to 22.3 and 27.7 g/L, respectively.
Zhang et al. (2016) studied the negative effect of ferric and ferrous ions on scale-inhibitor efficiency and the effect of adding EDTA
and citric acid to reverse the negative effect using a laser-nucleation-detection apparatus. Results showed that ferric hydroxide precipitation would adsorb the scale inhibitor to its surface, causing low inhibition efficiency. The addition of EDTA improved the inhibition
efficiency because of the elimination of ferric hydroxide precipitation by direct chelation. Citric acid was found to be able to control
ferric ions through a competitive adsorption mechanism where iron(III) complexes with citrate instead of scale-inhibitor molecules.
In summary, citric, acetic, and lactic acids have been tested to stabilize ferric and ferrous ions in solution for many years. Among all
organic acids, citric acid has been successfully able to chelate ferric ions under different sets of conditions because of the stable formation of iron(III) citrate soluble compounds. However, excessive loading of citric acid in solutions saturated with free calcium ions can
impose the risk of calcium citrate precipitation. Acetic acid use is limited to be paired with other chelating agents to be effective in a
wide range of temperatures. Smith et al. (1969) reported that the stability of citrate ions with iron(III) is double the stability of acetate
ions with iron(III). Lactic acid, as an iron-control agent, has received less attention because of its effectiveness at low temperatures and
the possibility of calcium lactate precipitation at high loading (Smith and Mancillas 1969). Formic acid has not been used as an ironcontrol agent because of the low stability between its conjugate base (formate) and iron ions. However, well conditions, temperature,
stimulation-acid concentration, and expected iron levels in solution all affect the appropriate combination and concentration for choosing organic acids to control iron ions (Dill and Smolarchuk 1988).
Filter-Cake Removal. Filter cake is an essential step in drilling operations to protect pay zones from particles and fines invasion
(Al Moajil and Nasr-El-Din 2013). Different viscous and weighting materials are used to build up the cake, which include calcium carbonate,
barite, ilmenite, manganese tetroxide (Mn3O4), and hematite (Al Moajil and Nasr-El-Din 2013). Still, near-wellbore damage is attributed to
mudcake particles and filtrate invasion in the drilled formations (Almond et al. 1995; Leschi et al. 2006). One of the stimulation engineer’s
tasks is to remove and clean mudcakes from the pay-zone sandface to provide a smooth passage for hydrocarbon flow toward the wellbore.
HCl, organic acids, precursors of organic acid enzymes, oxidizers, chelating agents, and solvents are used to remove water-based or
oil-based filter cakes. Dissolving and cleaning the filter cake is highly dependent on the materials used to build these cakes. For
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example, HCl cannot dissolve a filter cake that is derived from barite and cannot be used to remove hematite-based filter cake because
the reaction would release iron(III) in solution (Al Moajil and Nasr-El-Din 2013). Also, organic acids should not be used alone to
remove Mn3O4-based filter cakes because of incomplete cleanup (Al Moajil et al. 2018). Many studies have been conducted on filtercake removal using organic acids and their precursors because they address part of the challenges associated with removal fluids, such
as reaction rate, zonal coverage, and corrosion protection.
Al-Otaibi et al. (2006) examined encapsulated lactic acid in precursors to remove drilling-fluid filter cake. Acid precursors advance
over other removal systems through slow reaction, complete zonal coverage, and better corrosion protection. The precursors are triggered by downhole conditions and water contact, as in Eq. 31.
Precursors þ water ! Acid þ alcohol: . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ð31Þ
Al-Otaibi et al. (2006) used an HP/HT dynamic fluid-loss cell to evaluate lactic acid precursors for water-based filter-cake removal.
An HP/HT pendant-drop instrument was also used to measure the surface tension of the released live acid and spent acid as a function
of temperature and acid concentration. The kinetics study of the precursors showed a high correlation between the amount of acid
released and temperature. For example, a solution with 10 vol% loading of the precursors would release 10 wt% acid within less than
30 minutes at 250 F, while it needed more than 30 hours at 150 F. The prepared spent lactic acid had a low surface tension, which
accelerates spent acid flowing back to surface. HP/HT fluid loss showed that precursors would remove a significant amount of calcium
carbonate particles in the drilling-mud filter cake but would not remove clay and silica particles (Al-Otaibi et al. 2006).
Precursors of formic and acetic acids were also available and tested for the same purpose. Al Moajil and Nasr-El-Din (2007) tested a
water-based mudcake with precursors of formic, acetic, and lactic acids that released concentrations of 6, 7, and 8 wt% of acids, respectively. Results showed that formic acid was the best in terms of reactivity, required soaking time, surface tension, and tested temperatures. Acetic acid precursors performed poorly because they needed much more time to release the encapsulated acetic acid. Precursors
of lactic acid were comparable with precursors of formic acid when the time and temperature were 6.25 hours and 250 F, respectively.
Alotaibi et al. (2010) tested precursors of formic, acetic, citric, and lactic acids to dissolve formate-based mudcake. The precursors
were loaded at 10 vol% concentration with the cake in an HP/HT filter press. The results showed that precursors of lactic acid outperformed all other acids in terms of dissolving the cake for 5 hours at 300 F and lowering the surface tension.
Despite the advantage of the previous chemicals, a more-optimized chemical is needed to reach complete mudcake removal. Ineffective mudcake removal is a critical issue in maximum-reservoir-contact wells where overbalanced horizontal drilling is active. Horizontal wells are more susceptible to near-wellbore formation damage because of longer contact time with drilling fluid, openhole
completions, and long pay-zone interval (Rostami and Nasr-El-Din 2010; Elkatatny and Nasr-El-Din 2012).
Rostami and Nasr-El-Din (2010) explained a way to optimize horizontal-well cleanup. This method is achieved by including a solid
ester in the drilling-fluid recipe that can produce lactic acid once contacted with water under reservoir conditions. This solid advances
over other systems because it can dissolve calcium carbonate particles and biopolymers in the cake. HP/HT static-filter-press tests were
performed to evaluate the new idea where the solids were mixed with the drilling fluid, and a differential pressure of 300 psi was applied
to form the cake. The results showed that a higher 70% removal efficiency could be explicitly achieved at 230 F, where filtrate volume is
also minimum. SEM tests were performed to inspect the remaining cake of any calcium lactate precipitation. SEM results showed that no
calcium lactate was produced, and calcium carbonate was the dominant species. However, this new idea needed 20 hours or more to reach
high dissolving efficiency; a soaking time of 5 hours would only dissolve 4% of the cake (Rostami and Nasr-El-Din 2010).
The same concept was investigated by Elkatatny and Nasr-El-Din (2012) through conducting filter-press tests using a 1-in. core
instead of the ceramic disk to simulate a sandface. Elkatatny and Nasr-El-Din (2012) were interested in looking at the amount needed
of lactic acid solids and the calcium carbonate particle size in drilling fluid to achieve the optimal removal efficiency and the minimum
filtrate volume. The results, supported with computed-tomography scan tests, showed that 100% removal efficiency occurred at
24 hours of soaking time and 250 F testing temperature. This was achieved by using a 3:1 lactic acid solid/calcium carbonate ratio in
the drilling-fluid recipe along with stimulating the cake with 10 wt% alpha amylase enzymes to remove polymer materials.
Al Moajil et al. (2018) conducted laboratory studies to remove filter cake and organic sludge built by oil-based drilling-fluid/
mud-acid interaction. They used 5 to 6 wt% liquid formic acid and 1 to 2 wt% HCl with the addition of 10 vol% of mutual solvent. Mud
cakes were prepared using laboratory and field conditions, and the mixture showed a removal efficiency of 99% for both when tested at
165 F. Analytical tests showed that the deposited sludge from the acid-solubility test with crushed core was mainly diesel and contamination from drilling-fluid additives. Solubility and coreflood tests coupled with effluent analysis and computed-tomography scans
revealed that the mixture would dissolve the created mudcake and would cause insignificant damage to the formation (Al Moajil
et al. 2018).
As can be noticed from the previous studies, precursors of lactic and formic acids provide a variety of reaction times from which a
completion engineer can choose. Length of zonal coverage and treatment time can determine the optimal choice between the fast acid
release of formic acid or the slow acid release of lactic acid. On the basis of this, Nasrallah and Vinci (2018) successfully implemented
precursors of lactic acid to remove drilling-fluid filter cake after confirmation of slow acid release by laboratory testing. The use of
formic acid precursors in this field was associated with high brine loss in formation caused by a fast reaction. Fig. 35 shows a comparison of completion-fluid loss in different wells associated with precursors of formic or lactic acids (Nasrallah and Vinci 2018).
Although calcium carbonate is the most commonly used weighting material in mud circulation, its low density undermines its application when high-density fluid is needed (Al Moajil and Nasr-El-Din 2013). Mn3O4-based fluid is an alternative solution. Organic acids
have been tested excessively to dissolve Mn3O4-based mudcake. However, Mn3O4 reaction with acids is complex and can produce lowsolubility reaction products.
Al Moajil and Nasr-El-Din (2007) evaluated different acids to dissolve a Mn3O4-based drilling-fluid cake. HP/HT visual cells were
used to conduct solubility measurements between 5 g of the cake and 80 to 90 cm3 of the chosen acid between 70 and 284 F at 200 psi.
HP/HT filter press was used as well to examine the effectiveness of acid precursors of formic and lactic acids. Solubility tests showed
that precursors of lactic acid (11.7 and 14.9 wt% acid concentrations) dissolved more solids than precursors of formic and acetic acids.
Acids in liquid forms showed that solid solubility is highly dependent on acid-reaction product (Al Moajil and Nasr-El-Din 2013). For
example, citric acid dissolved the highest amount of cake, but when the temperature increased from 212 to 284 F, the solubility
decreased severely because of the precipitation of manganese citrate.
Analytical tests of the reacted Mn3O4 solids showed that it is coated with a polymeric material that would lower the acid’s dissolving
ability. When these coated materials were degraded using a particular enzyme, all acid solubilities increased significantly. According to
the study results, lactic and citric acids performed better than formic and acetic acids with the caution of citrate-based precipitation
(Al Moajil and Nasr-El-Din 2007).
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Total Brine Losses (bbl)
8,000
7,000
6,000
5,000
Formic acid
4,000
Lactic acid
3,000
2,000
1,000
0
1
2
3
4
5
6
Wells
Fig. 35—Comparison of completion-fluid loss in different wells associated with precursors of formic or lactic acids (Nasrallah and
Vinci 2018).
Al Moajil and Nasr-El-Din (2014) added 1 wt% HCl to 4 wt% lactic acid to increase the acid strength in removing a Mn3O4-based
filter cake. The mixture tended to perform better at low temperatures; it dissolved 85% at 250 F and 100% at 190 F. The high reactivity
of HCl might cause this at high temperatures. The authors also showed that lactic acid or HCl alone as a stimulation fluid would perform
poorly because the maximum dissolving efficiency was up to 70% (Al Moajil and Nasr-El-Din 2014).
Wagle et al. (2017) laboratory tested the use of different organic acids to remove drilling filter cakes with the presence of barite and
Mn3O4 in the water-based mud. Results showed that acetic acid was the most efficient organic acid (40.5% efficiency) in terms of filtercake removal, whereas citric acid was the least performing (25.8% efficiency). They attributed the latter acid’s weak performance to the
manganese citrate product resulting from the reaction. Fig. 36 illustrates the filter-cake-removal performance of the tested
organic acids.
Removal Efficiency (%)
50
40.5
40
32.8
30
25.8
20
10
0
15 wt% formic acid
solution
15 wt% acetic acid
solution
15 wt% citric acid
solution
Fig. 36—Drilling-filter-cake-removal efficiency of formic, acetic, and citric acids (Wagle et al. 2017).
Chendrika et al. (2017) applied 15 wt% acetic acid coupled with 6 vol% of CIs in the Madura Sea Field in Indonesia for sweet wells
with conditions lower than 305 F and 8,100 psi. Laboratory tests showed that 7% of formic acid could dissolve all cake very fast, while
7% acetic acid was slower, which can give more time for uniform spotting of the acid. The acetic acid concentration was increased to
12 and 15% to increase the acid’s strength. Field tests on two wells showed good response in terms of wellhead pressure and gas rates
(Chendrika et al. 2017). However, flowback samples and precipitation investigation were not reported.
Additional field applications were reported using formic and acetic acids. 10 wt% formic acid was used instead of HCl on an oilbased filter-cake-removal project in Lunskoye Field in the far east of Russia (Ruksanor et al. 2012). Laboratory results indicated that
formic acid coupled with mutual solvent performed better than HCl in terms of uniform removal. However, oil-wetted calcium carbonate mudcake in sandstone did not dissolve efficiently in the absence of mutual solvent. A field treatment showed a 2.5-fold increase
after a formic acid-based treatment.
A mixture of 5 wt% HCl and 5 wt% formic acid was used in Unayzah-A Formation in Saudi Arabia as a wash-acid treatment to
remove water-based mudcake. An injectivity-optimization study was performed for acidizing stimulation treatments in this sandstone
formation by Al-Mohammad et al. (2012). The treatment contained the aforementioned mixture as a preflush stage, followed by mud
acid. The study suggested a new method to evaluate a mudcake fluid dissolver where a coreflood apparatus was used along with solubility test to evaluate mudcake-removal efficiency. In the coreflood test, a water-based mudcake was built in the core-inlet face before
injecting the preflush mixture. According to the study tests, formic acid acted as a buffering agent and required extra soaking time to
remove the drilling-fluid mudcake efficiently, even when it was combined with HCl (Al-Mohammad et al. 2012). A possible improvement for the 5:5 wt% mixture can be noticed if the test temperature was increased from 186 to 250 F or higher to increase the
mixture reactivity.
Precursors of formic acid have also been implemented in low-temperature reservoirs. Davidson et al. (2012) reported that formic
acid precursors provided complete cake removal thanks to the low temperature of Captain Field in Scotland. Their study showed that
90% of the acid would be released within 2 days because the field temperature is as low as 30 C. Degouy et al. (2015) reported a successful formic acid precursor treatment in the Alkhalij Field offshore Qatar, where the temperature is approximately 53 C.
Overview
Using the reviewed papers, Table 8 was constructed to provide a comprehensive comparison between the four organic acids in the
fields of carbonate stimulation, sandstone stimulation, mixing with gluconic acid, iron control, and filter-cake removal, in addition to
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their dissociation constants and corrosion rates. Using the papers’ deduced results, future tests are recommended in the areas of lactic
acid stimulation because more field data are needed to draw a strong conclusion. Moreover, the filter-cake-removal efficiency of the
four organic acids needs to be further studied because various field applications produced different conclusions. Overall, the literature
needs to be enriched with more quality field-application data of the four organic acids to enhance our understanding.
Category
Formic Acid
Acetic Acid
Citric Acid
3.751
4.756 (weakest among
the four)
3.128 (strongest among
the four acids)
3.86
Corrosion rate of
carbon-steel tubulars
(300 F, 10 wt%)
>1 lbm/ft2 metal loss
(in 6 hours)
(Nasr-El-Din et al.
2012)
0.6 lbm/ft2 metal loss
(in 6 hours)
(Nasr-El-Din et al.
2012)
0.7 lbm/ft2 metal loss
(in 6 hours)
(Nasr-El-Din et al.
2012)
0.45 lbm/ft2 metal loss
(in 4 hours) (Alhamad
and Miskimins 2020)
Carbonate stimulation
Formic and acetic acids are usually blended
together (acetic provides retardation, formic
increases blend strength)
HT applications (250 F and higher; 170 F and
higher for surfactant-based fluids)*
Injection-rate optimization
Applied in laboratory and field (better results when
both acids are mixed together)
Encapsulated citric
acid technology
Most results were
poor (mainly
attributed to calcium
citrate precipitation)**
Applied as solid beads
Results were less than
expected (mainly
attributed to calcium
lactate precipitation)†
More field application is
needed
Sandstone stimulation
Mixed with HF for sandstone stimulation
Tested in sandstones with high chlorite and zeolite (where mud acid usually has problems)
Drawback of aluminum fluoride precipitation††
Citric acid provided the best results among the four acids but associated with citrate-based precipitation
Dissociation constant
(pKa)
Mixing with gluconic
acid
Iron control
Mixture helped to prevent precipitation or increase
salt solubility (Mohamed et al. 2015)
Did not improve
solubility (Mohamed
et al. 2015)
Mixture improved
solubility (Alhamad
2020)
Has not been used
because of low
stability between
formate and iron ions
The worst agent among
the three acids. It has
to be paired with
chelating agents at
high temperatures of
150 F or higher (Ali
and Hinkel 2000)
The best agent among
the three acids.
Excessive loadings
cause calcium citrate
precipitation
Lower performance than
citric acid. Perform
better at low
temperature of 100 F
or lower (Ali and
Hinkel 2000)
Precursors performed
well (especially with
calcium carbonate)
More effective when
used on Mn3O4 cakes
(Wagle et al. 2017)
More effective when
used on Mn3O4
cakes but associated
with citrate-based
precipitation
(Al Moajil and
Nasr-El-Din 2007,
2013)
Precursors performed
well (especially with
calcium carbonate
when paired with low
concentration of HCl)
(Al Moajil and
Nasr-El-Din 2014)
Corrosive without a CI
Limited to 9 wt% to
avoid formate-based
precipitation (unless
gluconic acid is mixed
with) (Mohamed et al.
2015)
Low stability with iron
ions
Increases F/Al ratio in
sandstone acidizing††
Limited to 13 wt% to
avoid acetate-based
precipitation (unless
gluconic acid is mixed
with) (Mohamed et al.
2015)
Low performance as an
iron-control agent
(because of
temperature limitation)
Low reactivity with
treated rock
Increases F/Al ratio in
sandstone acidizing††
Limited in most
acidizing application
because of citratebased precipitation**
Limited to 9 wt% to
avoid lactate-based
precipitation (unless
gluconic acid is used)
(Mohamed et al. 2015)
Low performance as an
iron-control agent
(because of
temperature limitation)
Filter-cake removal
Limitations
Lactic Acid
*
(Dill and Keeney 1978; Van Domelen and Jennings 1995; da Motta et al. 1998; Nasr-El-Din et al. 2001, 2002, 2003, 2006; Al-Mutairi et al. 2005; Burgos et al. 2005; Li et al.
2011; Cetin and Nasr-El-Din 2017).
(Burgos et al. 2004; Al-Khaldi et al. 2003, 2010; Mohamed et al. 2015).
†
(Nasr-El-Din et al. 2007, 2009).
††
(Shuchart and Gdanski 1996; Nasr-El-Din et al. 2007; Chang et al. 2008; Yang et al. 2012; Al-Harbi et al. 2012).
**
Table 8—Summary table of the main aspects of formic, acetic, citric, and lactic acids.
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Conclusions
This paper provides an in-depth review of the main organic acids, including formic, acetic, citric, and lactic acids, and their uses in the
industry, particularly in carbonate- and sandstone-stimulation operations. The review also sheds light on other various applications of
the organic acids, such as iron control and drilling-filter-cake removal. The following conclusions can be inferred from this review:
1. Formic acid has been broadly used for formation-damage removal in carbonate formations and dissolving carbonates in sandstone
stimulations. It is a good option in corrosion-protection blends through the formation of carbon monoxide.
2. Acetic acid can be paired with other acids to dissolve carbonates or remove drilling-fluid filter cakes. Acetic acid as a standalone
acid can be less reactive compared with other organic acids.
3. Citric acid has been used in the industry as an iron-control agent to stabilize ferric ions from spent acid. Calcium citrate solubility
is very low, which makes citric acid a poor choice for calcium-rich rock stimulation.
4. Lactic acid precursors can serve as diverting agents in multizone acid-stimulation treatments. They can be also used as a main component in drilling-fluid mudcake.
5. There are mainly two drawbacks to using organic acids in carbonate stimulation, including the incomplete reaction between the
acid and carbonate and the poor solubility of the reaction products.
6. Mixing acetic, formic, or lactic acids with gluconic acid produced a salt with higher solubility, which greatly reduced precipitation.
7. Encapsulated citric acid technology was tested in carbonate formations and the precipitation of calcium citrate was found to be a
limiting factor of using citric acid with calcium carbonate.
8. The results of lactic acid stimulation of carbonate using a solid acid form were promising in the laboratory, but when taken to the
field, the outcome was not as encouraging because of byproduct precipitations.
9. Organic acids were tested when combined with HF acid to stimulate sandstone formations because of the retardation ability of the
mixture. Field tests showed better results in terms of dissolving suspended fines and clays for the organic/HF mixture treatment
compared with the mud acid. However, one of the challenges of organic/HF treatments of sandstone formations is fluoridebased precipitations.
10. In terms of iron control, tests showed that citric acid was the best agent because of its ability to chelate ferric ions under different
conditions. Acetic acid was a poor iron-control agent, but it acted as a buffering agent to keep the solution pH value at a low level
to prevent ferric-based precipitations.
11. In terms of filter-cake removal, using precursors of lactic and formic acids showed a high efficiency, especially in calcium
carbonate-based filter cakes. However, in Mn3O4-based filter cakes, the combination of HCl and an organic acid (e.g., lactic acid)
showed to be more effective than using the organic acid alone.
Nomenclature
k ¼ rate constant
Ka ¼ dissociation constant
pKa ¼ dissociation logarithmic constant
n ¼ reaction order
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SI Metric Conversion Factors
bbl 1.589 873 101 m3
F
( F–32)/1.8 C
ft 3.048 101 m
in. 2.54 10 cm
psi 6.894 757 10 kPa
rev/min 0.104 719 76 10 rad/s
978
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