SPE 152870

Dual Chelant Mechanism for the Deployment of Scale Inhibitors in Controlled Solubility/Precipitation Treatments M. J. Todd & A. J. Savin Clariant Oil Services, K. S. Sorbie, Heriot-Watt University

Copyright 2012, Society of Petroleum Engineers This paper was prepared for presentation at the SPE International Conference and Exhibition on Oilfield Scale held in Aberdeen, UK, 30–31 May 2012. This paper was selected for presentation by an SPE program committee following review of information contained in an abstract submitted by the author(s). Contents of the paper have not been reviewed by the Society of Petroleum Engineers and are subject to correction by the author(s). The material does not necessarily reflect any position of the Society of Petroleum Engineers, its officers, or members. Electronic reproduction, distribution, or storage of any part of this paper without the written consent of the Society of Petroleum Engineers is prohibited. Permission to reproduce in print is restricted to an abstract of not more than 300 words; illustrations may not be copied. The abstract must contain conspicuous acknowledgment of SPE copyright.

Abstract

In scale inhibitor squeeze treatments, precipitation of the inhibitor within the formation can lead to extended squeeze lifetimes.

However, such processes also have the potential to cause formation damage unless they are carefully designed and controlled.

The formation of a partially soluble inhibitor/metal complex within a reservoir is the objective for almost all precipitation

squeeze packages. However, historically there are numerous ways this is achieved almost all of which require a limited

operational window to be deployed successfully. In this paper, we describe the development of a novel dual chelant system

which provides a method for controlling both the “wanted” and “unwanted” precipitation of the scale inhibitor package within

the formation. The highly tunable nature of the system allows for ease of pumping at more extreme conditions (higher and

low temperatures, calcium levels etc.) than have previously been possible. By using the dual chelant mechanism described in

this paper, a package can be tuned to precipitate within a certain time frame both at low and high temperatures in brines with

varying degrees of salinity and hardness. The scale inhibitor (SI) itself is a chelant or ligand for divalent ions present (mainly

Ca2+) and this is denoted L2 and the second chelant, L1, is added to the system at certain design concentrations, as explained in

the paper.

In many situations, the high divalent metal ion content of a produced brine, or formation water can limit the successful

pumping of a scale inhibitor due to high levels of calcium, for example. Under these conditions the dual chelant mechanism

can also be deployed to prevent scale inhibitor phase separation.

This paper discloses the theory of how the dual chelant mechanism works using computer modeling and the subsequent

confirmation of the simulations by laboratory testing. The importance of the pKa of the SI (L2) and the added chelant, L1, and

the relative metal binding constant interactions between L1/ L2 and Ca2+ are explained and investigated. The comparison of the

dual chelant mechanism versus conventional packages is demonstrated by core flood experiments. The dual chelant

mechanism gives a clear improvement in squeeze lifetime and controllability and provides a platform for the development of

many types of controlled solubility scale inhibitor treatment.

Introduction

The deposition of inorganic solids within oil production equipment is a severe problem that if left untreated can

seriously impair or in some cases stop oil production altogether. Inorganic scale can be produced from two main scenarios.

Firstly the formation of carbonate scale from changes in pressure and pH of the formation water as it is produced through the

facility with the oil and gas. Secondly, sulphate scale production can be induced by the mixing of incompatible waters. When

sea water is used as reservoir pressure support from injector wells, the sulphate rich sea water mixes with the resident

formation water which typically contain Ba2+ and Sr2+ ions, resulting in supersaturated solutions producing BaSO4 and SrSO4

scales. This is not generally regarded as a problem if these scales deposit within the reservoir, in fact this can be advantageous

as the resultant SO42- stripping reduces the supersaturation and the likelihood of scale formation in more sensitive areas.

However if the Ba2+ and Sr2+ rich formation water and SO42- rich sea water mix in the near well bore area resulting in a high

supersaturation (high scaling tendency) then the scale formation can severely hinder oil production blocking perforations,

constricting pipelines and impeding the function of safety critical devices such as down hole safety valves (DHSV). In these

2 SPE 152870

cases the scale formation is damaging to the producer well and must be removed. However, workovers to remove scale

mechanically are expensive and time intensive. The industry has long focused on prevention of scale formation by applying a

range of scale inhibitors such as phosphonates and various polymeric species.

Scale inhibitors are attractive due to their sub-stoichiometric mode of action as nucleation inhibitors or their inclusion into the

growing crystal structure and the resulting distortion known as crystal growth modifiers. This differs from traditional chelating

agents which need to be applied in a 1:1 molar ratio. This provides obvious benefits when considering large application

volumes. Scale inhibitors generally function in the region of 1-10 mg/L depending on the scaling severity and the conditions of

application. The exact concentration required to prevent scale is determined in laboratory experiments as the minimum

inhibitor concentration (MIC) and refined during field application.

Scale inhibitors are generally applied in two ways.

1) Continual injection

2) Squeezed into the near wellbore formation

Continual injection gives the best control and is the least wasteful as the concentration of the scale inhibitor applied can be

varied as the conditions change. However there must be a suitable chemical injection capillary in the region affected by, or

upstream of the scaling region. This is ideal if a suitable injection string exists but is an expensive option to retro-fit. Also

injection strings/capillaries can only reach down to certain points above the perforations in the well and it is almost impossible

to continually inject scale inhibitor into the near well bore formation.

Squeezing (forcing under pressure) the inhibitor into the reservoir formation allows the inhibitor to be placed in the near

wellbore area, this is generally followed by an over-flush to displace the inhibitor far enough into the higher temperature

region. The well is usually shut-in for some period (4-24 hours depending on the treatment design) to allow the inhibitor to be

retained through adsorption or precipitation onto the rock matrix within the pores Sorbie et al. (2005). When the well is

brought back online the produced water causes desorption, or dissolution of the retained inhibitor protecting the near well bore

and production equipment. These squeeze applications are expensive due to deferred oil production, as a result it is important

that the treatment last as long as possible. It is also imperative the formation is not damaged in any way to impede oil

production.

Precipitation vs. Adsorption Squeezing

Adsorption squeezes are the main stay of downhole scale prevention in the oil industry, but in some cases they do not deliver

the desired protection lifetime due to a number of factors e.g. low temperature, clean sandstones thus less high surface area

clays to act as adsorption sites or high porosity, etc. There have been numerous techniques developed to increase squeeze

lifetime. One of the most widely used is the application of precipitation squeeze treatments. In precipitation treatments the

injected scale inhibitor reacts with divalent cations (usually Ca2+, although others have been used) that are either present in the

formation or injected with the inhibitor in the squeeze package. Essentially the scale inhibitor is acting as a chelating agent for

the divalent cations, the efficacy of which is pH dependent. The result is a divalent metal salt of the scale inhibitor which is

only slightly soluble in the produced fluid. The advantage is that a relatively large mass of scale inhibitor is stored in the pore

space as a solid. The main mechanism for delivering the scale inhibitor to protect the facility is then dissolution after which

adsorption may proceed. The rate of inhibitor release is then governed by the inhibitor complex solubility Cs which depends on

temperature, [M2+], pH and inhibitor concentration Malandriono et al (1995). Malandriono et al (1995) modeled core flood

data and proposed a number of mechanisms for inhibitor retention and release from a reservoir rock matrix due to SI

precipitation mechanisms.

Factors Controlling Scale Inhibitor Precipitation

Several laboratory based mechanistic studies have been conducted to determine the optimum conditions for extension

of squeeze life through precipitation of scale inhibitors, Tantayakom (2004), Tantayakom (2005), Browning (1995), Kan

(1994), Rabaioli (1995). Tantayakom (2004, 2005) investigated the effects of divalent cation concentration, pH and ionic

strength of two phosphonate (ATMP and DTPMP) precipitating solutions at room temperature and the resultant solids. It was

found that in both cases the increase in pH resulted in a higher divalent cation (M+) to scale inhibitor (SI) ratio of the solids. As

this ratio of M+:SI of the precipitates increased the solubility decreased in both cases. The inclusion of Mg2+ into the reactions

had dramatic affects. Increasing the Mg:Ca ratio in the solution increased the solubility of the solids, e.g. the ksp for Mg:ATMP

1:1 was 0.049 M2 whereas the solids with Ca:ATMP 1:1 had a ksp of 0.017 M2 three times lower Tantayakom et al. (2004) .

This is an important factor when considering that the injection fluid may be sea water. It is interesting to note that even at pH

10 the highest Ca:DTPMP ratio reported was 3.82, whereas stoichiometrically one would expect a 5:1 ratio based on the

SPE 152870 3

ionizable protons (pKa) available on a DTPMP molecule. Browning (1995) conducted a similar study with the simple

phosphonate scale inhibitor HEDP. Similar observations to Tantayakom et al were made however it was noted the pH

alteration and hence altering the M+:SI ratio resulted in very different crystal morphologies which in turn had an effect on the

dissolution rate. Kan (1994) observed similar findings for DTPMP noting that the precipitates with a higher crystallinity

exhibited slower dissolution rates than the amorphous solids hence a longer squeeze lifetime would be expected. Kan (1994)

related the laboratory studies to field application and hypothesized that the precipitates form during a squeeze shut-in would be

of this highly crystalline nature. Rabaioli (1995) conducted a study but using a polyacrylate scale inhibitor. The same general

rules were found for the polymer as observed for the phosphonates. Rabaioli (1995) constructed full phase diagrams with

temperature which showed the changes in solubility with temperature. It was noted that polydispersity of the polymer will

affect the way it precipitates in the reservoir with higher molecular weights being the least Ca2+ tolerant and least soluble. This

will in turn affect the scale inhibition properties of the returning fluids.

Kinetics were not studied in any of the above studies but it is hypothesized that the precipitates which appear the most

attractive in terms of low dissolution rate are probably the fastest to form and hence the most difficult to use. The above

studies highlight the importance of the following factors when considering a treatment package.

Choice of scale inhibitor

pH of precipitating solution

Ca2+ : scale inhibitor ratio

The resulting supersaturation

The temperature of deployment

This paper focuses on exploiting the above findings in a safe application of the technology. Precipitation technology is still

less widely used than adsorption treatments because it carries the inherent risk of formation damage if the treatment is not

correctly placed. In these cases the treatment itself causes scale to form, blocking perforations or face plugging the formation.

Thus controlled placement of such a treatment is essential to a successful squeeze.

Deployment Methods

The simplest method for inducing scale inhibitor precipitation is to inject an acidic scale inhibitor into a reservoir (normally

carbonate) where Ca2+ is present and is released in the presence of the inhibitor as described in the patent by Smith (1972).

After dissolution and CO2 release, the pH of the injection fluid increases causing deprotonation of the scale inhibitor and

formation of the scale inhibitor precipitate. However, there is limited control with regards to placement with such a treatment

and formation damage is possible. If the reservoir is of mixed mineralogy (e.g. sandstone and carbonate) introduction of an

acidic scale inhibitor could dissolve the matrix cementing carbonate materials resulting in formation damage. In the cases

where the well to be treated does not have sufficient carbonate mineralogy the Ca2+ must be provided by other means and one

such treatment relies on the ion exchange capacity of resident clay minerals within the reservoir as described in the patent by

Lawson (1983) . However such a method may induce clay swelling and or migration resulting in significant formation

damage. The main way of providing the [Ca2+] required for scale inhibitor precipitation is to either over-flush with a Ca2+ rich

fluid or provide the Ca2+ in the inhibitor slug and control the precipitation reaction in another manner Carlberg (1987, 1989),

Hen (1992), Olson (1992), Bourne (1997), Todd (2010), McRae (2004) and Lynn (2003).

A popular method is to provide an acidic scale inhibitor solution with a high [Ca2+], such that at elevated pH the scale inhibitor

calcium salt will form. The inclusion of a thermally activated pH modifier has been used to cause the deprotonation of the

scale inhibitor after a time delay in the elevated temperature of the reservoir as described by Hen (1992). The thermally

activated modifier is usually a urea based component which produces NH3 under thermal hydrolysis as shown in equation 1.

NH2CONH2 + H2O 2NH3 + CO2 (1)

This type of treatment has been widely applied in the North Sea as described by Olson (1992), Bourne (1997) and Todd

(2010). However, the main limitation of this type of treatment is that the thermal degradation of urea is the rate limiting step. It

can only be applied at temperatures greater than approximately 85oC with an upper limit of approximately 130oC. Below 85oC

the hydrolysis reaction does not take place or is so slow that the shut-in times required would be too long to make the

treatment economically viable. As the temperature increases, the hydrolysis rate increases such that at any temperature much

higher than 130oC the reaction rate is so rapid that premature precipitation is very likely. In some cases, this kind of treatment

can be applied at elevated temperatures but the pumping capabilities have to be substantial as indicated in Todd (2010). At

temperatures lower than 85oC, the hydrolysis of urea has to be activated by other means e.g. enzymatic control has been

reported to have worked very well by McRae (2004). However, in this case the rate limiting step is still the hydrolysis of urea

and hence there is not any additional control element to this technique. Other techniques do exist for low temperature

precipitation for example Lynn (2003) describes transporting the components separately in an emulsion form and, once placed

4 SPE 152870

in the reservoir, the emulsion breaks and the components mix. This is similar to over-flushing with a Ca2+ rich solution and

relying on adequate mixing in the reservoir which is unlikely to give the same degree of precipitation as other methods.

All of the above applications rely on a similar method for forming the precipitate, i.e. mixing Ca2+ and the desired scale

inhibitor at elevated pH to form the SI_Ca salt. The subtlety comes in providing a) the correct mixing and b) providing a

controlled pH increase in such a way as to maintain full control of the process.

Carlberg (1987, 1989) describes a different approach to designing precipitation squeezes which he refers to as operating

through “dual chelant” mechanism. The method described addresses point a) above as the Ca2+ and scale inhibitor solutions

are pumped together, and point b) does not apply as the scale inhibitor solution is pre-adjusted to a pH at which precipitation

would take place. Carlberg introduces a secondary chelant, the first being the scale inhibitor, which controls the [Ca2+]

available to the scale inhibitor during the mix and in turn is completely soluble at all [Ca2+]. The key to the Carlberg method is

that the secondary chelate used (NTA_Ca(aq)) has a lower stability constant (calcium binding constant) (Equation 2) than the

scale inhibitor chelate (SI_Ca(aq)). If this system is left for sufficient time, it is thermodynamically inevitable that the

precipitate will form at any temperature (SI_Ca(s)) since the calcium will repartition onto the SI from the added chelant. Thus,

this system relies of the kinetics of calcium redistribution from the added chelant to the SI and when this occurs to a suitable

degree the SI_Ca complex will precipitate. Carlberg et al describe experiments and application in wells at near room

temperature (26oC). In other words, the added chelant in Carlberg’s approach, NTA, is effectively a delaying agent and as such

this mechanism is using a kinetic control. As the temperature is increased the reaction time will increase which will

significantly limit the application range of this type of dual chelant system.

K = [M] [L]

[ML]

(2)

Where:

K = Stability constant,

[ML] = The concentration of the chelate formed from metal cations and either the secondary chelant or the scale inhibitor.

[M] = The concentration of the metal cations.

[L] = The concentration of either the secondary chelant or the scale inhibitor.

The central idea of this new technology is to use a secondary competitive chelant (denoted L2) with a calcium stability

constant equal or higher than that of the scale inhibitor (L1). The secondary chelant (L1) modifies and controls both the

availability and solubility of the working complex in such a manner that at low temperatures the precipitate will never form. At

increased temperatures the solubility of the SI_Ca complex is regulated in order to provide a controlled reaction. This dual

chelant system is described through the use of mathematical modeling and the supporting laboratory findings as described

below Experimental Table 1 Sea Water Composition

Species Concentration (mg/L)

Sodium 10890

Potassium 460

Magnesium 1368

Calcium 428

Strontium 7

Chloride 19766

Bicarbonate 140

Sulphate 2960

Control Study

In order to assess the level of precipitation control from the dual chelant system (solution 1) it was compared to a

system with no secondary competitive chelant (solution 2) and a urea hydrolysis package (solution 3), where the secondary

chelant is referred to as L1. In all cases the concentration of DTPMP and Ca2+ were kept constant at 0.025M and 0.028M,

respectively. The various solution compositions as given in Table 2. All pH adjustments were made with 50% KOH solution.

SPE 152870 5

Table 2 Composition of solutions for control study

Solution Number [DTPMP]=[ L2] (moles) [L1] (moles) [Ca2+

] (moles) Initial pH

1 0.025 0.014 0.028 6.2

2 0.025 0 0.028 6.2

3 (Urea) 0.025 0 0.028 2

The neat solutions were monitored at room temperature for 2 hours, the solutions were then mixed 16% v/v in sea water (Table

1) and monitored at room temperature for a further 2 hours. They were then placed in an oven at 52oC and monitored for 24

hours. The existence and extent of any precipitate was noted visually as a haze or a precipitate.

Sensitivity Study

In order to determine the sensitivity of the model a series of experiments were conducted. In all solutions the pH,

[DTPMP] and [Ca2+] were kept constant, while the [L1] was varied. Four concentrations of L1 were chosen; 0.008M, 0.014M,

0.023M and 0.031M. The solutions were made according to Table 2 and all pH adjustments were made with 50% KOH

solution.

Table 3 Composition of Solutions for the sensitivity study

Solution Number [DTPMP] [Ca2+

] [L1] Initial pH

1 0.025 0.028 0.008 6.2

2 0.025 0.028 0.014 6.2

3 0.025 0.028 0.023 6.2

4 0.025 0.028 0.031 6.2

For each solution 6 duplicates were mixed 16% v/v in sea water and monitored for 5 minutes.

The following procedure was then carried out:

- One duplicate from each solution was left at room temperature and the remaining 5 were placed in an oven at 30°C.

- The samples were monitored for haze and precipitation over 24 hours. If the samples did not show any haze or precipitation

after 24 hours the temperature was increased by 10°C.

- If samples still did not show any haze or precipitation after 24 hours the temperature was again increased by a further 10°C

and monitored for 24 hours. This process was repeated until a precipitate was seen in each sample, or the temperature had

reached a maximum of 120oC.

- Once the samples for a solution had precipitated at a given temperature one of the duplicates was removed and the precipitate

filtered and weighed on a four decimal point balance.

- The remaining duplicates remained in the oven and the temperature increased. At each temperature increase another

duplicate was removed and again the precipitate weighed. This process was repeated until all the duplicates had been

depleted.

The resulting mass vs. temperature plots were compared to those generated in the model presented below.

Results and Discussion Control Study

The purpose of the control experiment was to show that the precipitation of the DTPMP salt (L2_Ca(s)) can be

thermodynamically controlled by the addition of the secondary competitive chelant L1. The solutions 1 to 3 were monitored at

room temperature for 2 hours. Solution 2 started to phase separate immediately, the resulting solutions are presented below in

Figure 1. It can be seen that there is a large mass of solid precipitate in solution 2. However, solutions 1 and 3 remain clear

with no phase transition observed. This shows that the addition of the competitive chelant L1 stops the formation of the

DTPMP-calcium salt. Solution 3 remained clear as there was not sufficient heat to hydrolyse the urea. A comparative study

was performed where solutions 1 and 3 were systematically checked at room temperature for several months with no sign of

phase separation.

6 SPE 152870

Figure 1 Solutions 1-3 after 2 hours at room temperature.

Solutions 1 and 3 were diluted 16% v/v in sea water and placed in an oven at 52oC and monitored for 24 hours. This

temperature was chosen as it is an example of an actual field application temperature at which traditional adsorption squeeze

applications had provided poor squeeze lifetimes. Solution 1 showed the initial stages of phase separation after 4 hours at

52oC, whereas solution 3 remained transparent with no indication of phase transition throughout the test period. Figure 2

shows solutions 1 and 3 after thermally aging at 52oC for 24 hours. 24 hours was chosen as it represents the longest,

economically viable shut-in time allowed by most operators.

Figure 2 Solutions 1and 3 diluted 16% v/v in laboratory sea water and thermally aged at 52

oC for 24 hours.

As expected, there was not sufficient thermal energy at 52oC to cause hydrolysis of the urea and in turn initiate the

DTPMP-calcium solid formation. However, solution 1 shows a considerable mass of precipitate. The indefinite prevention of

precipitation provided by the addition of the secondary chelant at room temperature has been observed. The ability to provide

phase separation at low temperatures that are not possible using conventional controlled precipitation techniques illustrate the

advantages of such a technology. These results look very promising, provided the addition of L1 does not lead to increased

formation damage over the comparative scale inhibitor adsorption packages.

Sensitivity Study

A series of experiments was carried out to investigate the effect of the secondary chelant concentration [L1] has on the

phase boundaries and hence the separation parameters of the DTPMP calcium salt (L2_Ca(s)). Figure 3 shows the mass (m) of

precipitate from each package with increasing temperature at a variety of L1 concentrations, varying from 1.22 to 4.88

mMoles/L. In all these experiments the L2:Ca (where L2 is DTPMP) ratio was kept constant as was the initial pH of the

treatment package. At the lowest concentration (1.22 mMoles of L1), precipitate was observed at room temperature of mass, m

= 22332 mg/L. As the temperature is increased to 52oC the mass of precipitate increases to m = 46210 mg/L of treatment

solution. As the temperature is increased further the mass of precipitate remains relatively constant up to 80oC.

SPE 152870 7

Figure 3 Mass of DTPMP-calcium precipitate with increasing temperature at various secondary chelant (L1) concentrations

The concentration of L1 was increased to 2.32 mMoles/L. With this increase in L1 concentration no precipitate was observed at

room temperature. However as the temperature is increased to 30oC, m =7676 mg/L of precipitate was observed, which was

significantly lower than the onset of precipitate observed at an L1 concentration of 1.22 mMoles/L. This observation indicates

a reduction in mass with increased L1 concentration. As the temperature is elevated to 40 then 52oC a steady increase in the

mass of DTPMP precipitate is observed, from m = 11546 to m= 17284 mg/L, respectively. As the temperature is increased

further to 60oC, the mass of precipitate measured was m = 31720 mg/L showing a deviation from the previously linear

behavior. At 70 and 80oC the masses recorded were m = 42802 and m = 46432 mg/L indicating a reduction in the mass of

precipitate with further increases in temperature. The final mass measured for an L1 concentration of 2.32 mMoles/L is

analogous to that measured for 1.22 mMoles/L of L1 at the same temperature, m = 46432 mg/L and m = 48018 mg/L,

respectively.

At an L1 concentration of 3.66 mMoles/L, no precipitate was observed between room temperature and 60oC. As the

temperature was increased to 70oC a mass of, m = 13490 mg/L of precipitate was observed. This is significantly lower than

that observed for the lower concentrations of L1. At 80oC, there is a rapid increase in the mass of solids measured, m = 35946

mg/L which is considerably lower than the mass measured at the same temperature for lower L1 concentrations. The

temperature increase to 90oC resulted in a mass of m = 49886 mg/L precipitates. This is analogous to the final masses

measured at 80oC for 1.22 and 2.32 mMoles of L1. It would appear that at concentrations ranging from 1.22 to 3.66 mMoles/L

of L1 the onset of precipitation is shifted to higher temperatures, however the final mass appears analogous, i.e. m = 46000-

50000 mg/L.

The final concentration of L1 tested in this study was 4.88 mMoles/L. There was no precipitate observed between room

temperature and 70oC. As the temperature was increased to 80oC, the mass of precipitate was m = 24354 mg/L. This is a 10oC

increase in the onset temperature compared to an L1 concentration of 3.66 mMoles/L. At 90oC, the mass measured was m =

31982, and 100oC resulted in m = 30100 mg/L, 110oC produced m = 27182 mg/L and 120oC resulted in m = 31448 mg/L. The

variation between the mass measured from 90 to 120oC is believed to be within experimental error, although this has not been

quantified. The maximum mass of precipitate utilising an L1 concentration of 4.88 mMoles/L is considerably less than those

observed at lower concentrations, approximately m = 20000 mg/L.

0

10000

20000

30000

40000

50000

60000

0 10 20 30 40 50 60 70 80 90 100 110 120

Mas

s o

f P

reci

pit

ate

(m

g/L)

Temperature (°C)

1.22 mMoles/L L1

2.32 mMoles/L L1

3.66 mMoles/L L1

4.88 mMoles/L L1

8 SPE 152870

In general the following observations can be made regarding the thermodynamic control offered by the dual chelant system:

The onset temperature for precipitation increases with increasing L1 concentration.

The mass of precipitate increases with increasing temperature.

For concentrations of L1 between 1.22 and 3.66 mMoles/L the maximum mass of precipitate appears to converge

between m = 46000 and 50000 mg/L

At an L1 concentration of 4.88 mMoles/L the mass of precipitate is considerably less at all temperatures measured, m

~20000mg/L.

The results presented in Figure 1 show that the onset temperature and mass of precipitate can be controlled

thermodynamically. The addition of the secondary chelant, L1 prevents the precipitation of the DTPMP-calcium salt at low

temperatures. This study did not look at the kinetics, i.e. the times to precipitation. However it should be noted that because it

is thermodynamically impossible for the formation of precipitation at lower temperatures with the addition of certain L1

concentrations this shows a different mode of action that previously described by Carlberg. This distinction is now

demonstrated further using equilibrium modeling of the dual chelant system below.

The Model for the Dual Chelant System

In the Carlberg (1987, 1989) case a second chelant or ligand (L1) was deployed whose binding with calcium was

weaker than the main scale inhibitor (L2) which is also a calcium chelant. The calcium rich brine and L1 chelant were mixed in

advance and this mixture was blended with scale inhibitor (L2) and immediately injected into the formation. Because the

binding constant of the scale inhibitor was higher than of the added chelant, the calcium would gradually (i.e. kinetically)

repartition from the chelant to the scale inhibitor. Once sufficient calcium had redistributed, then the Ca_ L2 (Ca_SI) complex

would precipitate because of its reduced solubility in the presence of calcium. Thus, the Carlberg system relied on the kinetics

of Ca repartitioning for the system to operate. The system described in the experimental results above and by the theory below

does not rely on the kinetics for its operation.

For this new system, we deploy two chelants as follows:

(i) L1 = X which has a higher binding constant with Ca compared with the main inhibitor L2 (DETPMP) (Kb~ 1010) but this

complex L1_Ca is always very soluble (at all temperatures). In fact, the 2 chelant system reported here can work if the L1 – Ca

binding constant is comparable or somewhat lower than that of the SI since the system can be controlled by varying the

concentrations of chelant L1;

(ii) L2 (DETPMP or any other precipitating SI) as previously (Kb ~ 1010) but this complex L2_Ca has a varying solubility with

temperature – see below.

To understand and model how this system works, we must write out the competing binding reactions between the 2 chelants

(L1 = X and L2 = DETPMP) and calcium along with the coupled precipitation reaction as follows, where the xi denote the

molar concentrations (activities) at equilibrium:

2

1 (aq) ( ) 1 ( )

2 1 4

_

aq aqL Ca L Ca

x x x

Binding const 1K (3)

2

2 (aq) ( ) 2 ( )

3 1 5

_

aq aqL Ca L Ca

x x x

Binding const 2K (4)

2 ( ) 2 ( )

5 6

_ _

aq sL Ca L Ca

x x

(5)

Solubility of complex = 2sC which varies with temperature

Therefore, at equilibrium, there are six species as follows (the first 5 in the aqueous phase and the last one as a precipitated

solid): 1 (aq)L , 2

( )aqCa , 1 ( )_ aqL Ca , 2 (aq)L , 2 ( )_ aqL Ca and 2 ( )_ sL Ca . Therefore, there are 6 unknowns which we must

find in an equilibrium calculation for this system. We suppose that we start with the following initial concentrations of

1 (aq)L , 2

( )aqCa and 2 (aq)L in the system (no complexes formed):

SPE 152870 9

Initial concentrations: 10[ ]initialCa x , 1 (aq) 20[ ]initialL x and 2 (aq) 30[ ]initialL x

To solve the above equations (for 1 2 3 4 5 6, , , , , x x x x x x ), we proceed as usual in equilibrium systems and use the chemical

equilibria and mass balances as follows:

Equilibria: 11 2

1

[ _ ]

[ ][ ]

L CaK

L Ca 4

1

2 1.

xK

x x (6)

22 2

2

[ _ ]

[ ][ ]

L CaK

L Ca 5

2

3 1.

xK

x x (7)

Mass balances: Ca: 10 1 4 5 6x x x x x (8)

L1: 20 2 4x x x (9)

L2: 30 3 5 6x x x x (10)

NB: Although it appears as if there are 6 unknowns in the above 5 equations, there are actually only 5 unknowns because we

can always eliminate either 5x or 6x as follows:

Either

(a) there is no precipitate and hence 6x = 0 and we just have 5 unknowns ( 1 2 3 4 5, , , , x x x x x )

or

(b) there is a precipitate and it forms until 5 2sx C and so 5x is now known and set to this solubility value and we just have

5 unknowns ( 1 2 3 4 6, , , , x x x x x ).

The details of how these equations are solved are not important but this has been done by the authors in order to give a

numerical demonstration of how the 2 chelant Ca_ L1_ L2 system works as shown below.

Numerical Calculations for the Two Chelant Case (The Ca_ L1_L2 System): The system of equations listed above in Eqs. 6 -

10 are now solved for some example cases that demonstrate how the system works. Firstly, we need to define the solubility

curve of the 2 ( )_ sL Ca complex, which is a function of temperature, T, and we denote this by, Cs2 = Cs2 (T). Here, we

assume a simple model for Cs2 (T) as shown in Figure 4 for demonstration purposes.

Figure 4 Plot of solubility, CS2 (T) (M) vs. temperature, T (

oC) for the SI_Ca (L2_Ca )complex.

0

0.01

0.02

0.03

0.04

20 30 40 50 60 70 80 90 100

Solubility of the L2_Ca Complex (M)

T (oC) -->

Solubility(Cs2)

10 SPE 152870

Three examples are presented which differ only in the ordering of the magnitude of the calcium binding constants of the

chelant (L1) and the scale inhibitor (L2), i.e. K1 and K2.

Example 1: Takes the initial conditions of [Ca] = 0.045M, [SI] = [L2] = 0.04M and equal binding constants, K1 = K2 = 1E10.

This is close to the calcium binding constant for DETPMP. We assume a fixed solubility of the SI_Ca (L2_Ca) complex

corresponding to T = 70oC in Figure 4; i.e. CS2 = 0.01M. The concentration of L1 is then varied as the amount of precipitate

(M) is calculated by solving the above equilibrium equations.

The results in Figure 5 show that the appearance of precipitate can be controlled such that if we add [L2] = 0.15M, then no

precipitate will appear at 70oC for the conditions of Example 1. However, if the system heats up to above this temperature,

then precipitate will start to appear. This is best viewed as point (b) on the schematic Figure 8, as discussed below.

Figure 5 Plot of amount (M) of the L2_Ca complex precipitating vs. the concentration of [L1] at a fixed solubility level of complex (Cs2 = 0.01M) corresponding to, T = 70

oC; Example 1.

Example 2: Again, the initial conditions are [Ca] = 0.045M, [SI] = [L2] = 0.04M, but the binding constants are as follows, K1

=1E10, and K2 = 1E9. That is, the added chelant (L1) has a higher binding constant than that of the scale inhibitor (L2). The

solubility of the SI_Ca (L2_Ca) complex is again as given in Figure 4 as a function of temperature, T. For these conditions,

the amount of precipitated L2_Ca complex (M) vs. T is given in Figure 6 for a range of concentrations of the added chelant,

[L1] = 0 to 0.1M. When the binding constant of the added chelant (L1) is higher than that of the SI (L2), then by adding a

relatively small amount of L1, then it can control the amount of precipitated complex very well. Note that if sufficient chelant,

the amount of precipitated complex can be controlled and indeed completely suppressed (at [L1] = 0.1 M in this case; see

Figure 6). Note also that the theoretical mass of precipitate vs. T predictions in Figure 6 are qualitatively very similar to the

experimental results in Figure 3. Indeed, the experiments were designed using the model and this qualitative behavior is

actually a prediction of the model. This gives us high confidence that the mechanistic model described above (Eqs. 6 to 10)

capture the main features of this dual chelant system.

Example 3: In this example, the initial conditions are again [Ca] = 0.045M, [SI] = [L2] = 0.04M. However, the calcium

binding constant of the added chelant (L1), K1 =1E9, which is lower than that of the scale inhibitor, K2 = 1E10. The solubility

of the SI_Ca (L2_Ca) complex is again as given in Figure 4 as a function of temperature, T. For these conditions, the amount

of precipitated L2_Ca complex (M) vs. T is given in Figure 7 for a range of concentrations of the added chelant, [L1] = 0 to

0.85M. In this case, the fact that the binding constant of the added chelant (L1) is lower than that of the SI (L2), means that we

have to add a relatively larger amount of L1 in order to control the amount of precipitated complex. However, the same degree

of control of precipitated amount is still possible but using a sufficiently higher concentration of added chelant L1.

These examples illustrate how the new chelant system works and shows how this differs from previously proposed dual

chelant systems.

0.00E+00

1.00E-02

2.00E-02

3.00E-02

0 0.05 0.1 0.15 0.2

Am

ou

nt

of

L2_

Ca

Pre

cip

itat

e(M

)

Chelant conc., [L1] -->

Amount of L2_Ca (SI_Ca) precipitate (M) as a function of the chelant conc. [L1][Ca] = 0.045M; [SI] = [L2] = 0.04M; K1 = K2 = 1E10; Fixed solubility, Cs2 = 0.01M

SPE 152870 11

Figure 6. Plot of the amount (M) of the L2_Ca complex precipitated vs. T for a range of concentrations of [L1] for the solubility Cs2(T) in Figure 1; Example 2

Figure 7. Plot of the amount (M) of the L2_Ca complex precipitated vs. T for a range of concentrations of [L1] for the solubility Cs2(T) in Figure 1; Example 3.

How the Different Dual Chelant Method Work: Here we use the schematic solubility (Cs2) vs. temperature plot in Figure 8 to

explain how the dual chelant methods of Carlberg and this paper work mechanistically. The mixed chelant process described

by Carlberg (1989) utilises a secondary chelant where K1<K2 and due to the relative concentrations used the precipitation of

L2_Ca(s) at any application temperature is inevitable. This is illustrated in Figure 8 as point (a), as the condition of the mixed

chelant system is such that it lies to right of the phase boundary denoted as Cs2. The useful deployment reaction comes from

the kinetic control, or the relatively slow repartitioning of Ca2+ from L1 (the added chelant) over to the scale inhibitor, L2. Once

this reaches the point of super saturation the L2_Ca(s) is formed. It is proposed that as the temperature increases the reaction

rate will increase as the relative level of supersaturation increases. However, this latter factor is not studied here and will be

dependent on several field specific parameters including the temperature gradient during deployment and the rate of

temperature increase.

The dual chelant system described in this paper (point (b) on the schematic of Figure 8) relies on L1 having a K1≥K2 such that

the L2_Ca complex is held at a concentration below its solubility at the storage temperature (T ~ 70oC for point (b) in this

example in Figure 8). When the temperature is then increased by placement into the hotter reservoir region, point (b) then

moves to the right as shown in Figure 8 and a controlled precipitation of L2_Ca complex occurs. As the L2_Ca complex

0.00E+00

5.00E-03

1.00E-02

1.50E-02

2.00E-02

2.50E-02

3.00E-02

3.50E-02

25 50 75 100

Mas

s Si

_Ca

Pp

t(M

)

[L1] =0

[L1] = 0.01M

[L1] = 0.02M

[L1] = 0.03M

[L1] = 0.055M

[L1] = 0.10M

Mass precipitate SI_Ca (Molar) vs. T (oC) for various [L1][SI]= [L2] = 0.04M; [Ca] = 0.045M; K1=1.0E10; K2=1.0E9

T (oC) -->

0.00E+00

5.00E-03

1.00E-02

1.50E-02

2.00E-02

2.50E-02

3.00E-02

3.50E-02

25 50 75 100

Mas

s Si

_Ca

Pp

t(M

)

[L1] =0

[L1] = 0.01M

[L1] = 0.02M

[L1] = 0.04M

[L1] = 0.055M

[L1] = 0.10M

[L1] = 0.30M

[L1] = 0.50M

[L1] = 0.85M

Mass precipitate SI_Ca (Molar) vs. T (oC) for various [L1][SI] = [L2] = 0.04M; [Ca] = 0.045M; K1=1.0E9; K2=1.0E10

T (oC) -->

12 SPE 152870

precipitates, then the entire system (Eq. 6 – 10) re-equilibrates and more Ca2+ repartitions. The same competitive reaction is

demonstrated in the control study and illustrated in Figure 2. Solution 2 was stable under storage temperature but when applied

at 52oC the phase boundary Cs2 was crossed causing supersaturation and forming L2_Ca(s).

Figure 8 Hypothetical schematic illustrating where specific treatment packages lie in relation to the solubility of the L2_Ca complex Cs2. Where the Carlberg (1987, 1989) case is presented at (a), Example 1 where K1=K2 is presented at (b). The effect of altering [L1] is illustrated by the two headed arrow (c).

The alteration of precipitation on-set temperature is modelled in examples 2 and 3 and illustrated in Figures 6 and 7

respectively. Example 2 describes a dual chelant system where K1>K2 analogous to that observed in the sensitivity study

(Figure 3), i.e. at L2: L1 ratios ≥1 there is sufficient precipitation control with the onset temperature increasing with increasing

[L1], i.e. to prevent precipitation at room temperature and to cross the phase boundary (Cs2) at elevated temperature [L1] is half

that of [L2]. However, if the Carlberg case (example 3) is considered, the [L1] required for precipitation control is such that the

L2: L1 <<1, to prevent precipitation at room temperature and to cross the phase boundary (Cs2) at elevated temperature [L1] is

2.5 times that of [L2]. In both examples 2 and 3 described in the model a level of control over the solubility limit Cs2 is

introduced by altering the concentration of the competitive chelant [L1]. The same result is observed qualitatively for the

sensitivity study displayed in Figure 3, the exact concentrations are different as the exact values of K1 and K2 are not known

for laboratory study. However as the results are analogous, i.e. L2: L1 ≥1 there is sufficient precipitation control with low

concentrations of L1, i.e. to prevent precipitation at room temperature and to cross the phase boundary (Cs2) at elevated

temperature [L1] is 0.56 times that of [L2] which is analogous to that proposed in the model.

The relationship between L1 and L2 is illustrated by the double headed arrow at (c) in Figure 8. This shows that Cs2 is

essentially shifted by increasing or decreasing the [L1] as the [Ca] available to L2 is varied.

It appears that there is a useful reaction kinetic introduced by the competitive reaction between L1 and L2 as the time to

precipitation observed in the control study for solution 1 was 4 hours. Considering the application temperature is low (52oC)

this provides significant placement control and hence greater prevention of formation damage compared to other systems for

low temperatures such as the Carlberg (1987, 1989) mixed chelant system, or to a greater extent systems which rely on Ca rich

overflushes. The exact effects of L1 on reaction kinetics are currently being studied at the time of writing and will be discussed

in a future paper. This is a step improvement compared to previous precipitation control mechanisms such as urea hydrolysis

as the temperature window for application is not as limited.

0

1000

2000

3000

4000

5000

6000

0 10 20 30 40 50 60 70 80 90 100 110 120 130 140 150

[Ca]

(a.

u.)

Temperature (°C)

Cs2

(b)

(a)

(c)

2 Phases

1 Phase

SPE 152870 13

Conclusions

A dual chelant system has been developed where a secondary competitive chelating agent (L1) controls the [Ca]

available to the primary chelant (the scale inhibitor- L2). Laboratory experiments have shown that this competition results in a

stable solution that without the secondary chelant present would precipitate immediately. Furthermore, at low application

temperatures there is a controlled phase separation such that even at low pump rates the package could be placed the desired

distance from the well-bore preventing formation damage induced by formation face plugging. This system can be tuned by

careful blending of the [L2], [Ca] and [L1] to perform at conditions (such as high temperature) where conventional systems

have had limited success.

An equilibrium model of the system has been developed which demonstrates the operational principles of the novel dual

chelant mechanism proposed in this work. This mechanism is principally utilised when the calcium binding constant of the

added chelant (L1) is higher than that of the SI (chelant L2). Under these conditions (K1 >>K), then a small amount of added

chelant can be used to control the precipitation envelope very effectively.

However, if the calcium binding constant of L1 is comparable with or somewhat lower than that of the SI (L2) (K1 <K2), then

the system can still be made to work quite effectively but higher concentrations of L1 are required.

This novel dual chelant approach is not reliant on the kinetic of calcium repartitioning between a weaker chelant (L1) and the

SI (L2) as described by Carlberg (1987, 1989).

Acknowledgements The authors would like to thank Clariant Oil Services for granting permission to publish this paper.

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