Water Formed Scale - oilprocessing.netoilprocessing.net/data/documents/Water-Formed-Scale.pdf · 1.1.7 The mole A mole is defined as the amount of substance that contains as many

Books are available here

1 http://www.oilprocessing.net/oil/

Water Formed Scale

Oil Field Scale, formation, Prediction, and Inhibition.

Prepared by

Yasser Kassem

https://www.amazon.com/Yasser-Kassem/e/B07KXLM5RZ?ref=dbs_p_pbk_r00_abau_000000

http://www.oilprocessing.net/oil/



Preface

I would like to present a simple book in organic scales, and water formed scale.

The book is dedicated to oil and gas engineers and technicians dealing with oil

and gas produced water, water injection and water treatment facilities.

Best Regards

Yasser Kassem

Oct. 2nd, 2019


Discussion board


Free parts of fundamentals of oil and gas processing book

http://oilprocessing.net/data/documents/oil-and-gas-processing-fundamentals.pdf

اإلنجليزي للكتاب ترجمة وهو والغاز البترول معالجة عن بالعربية كتاب ألول مجاني تحميل

http://oilprocessing.net/data/documents/ogparabic.pdf

احفظ) الرابط على وهو الماوس يمين من تختار أن الممكن ومن, ذلك بعد وتحفظه الملف تفتح ممكن

الكمبيوتر أو توب الالب على مباشرة بتنزيله وتقوم... ( ك الهدف أو الملف .












Contents Chapter 1 ......................................................................................................................................... 5

Fundamentals of Chemistry and water Analysis ............................................................................. 5

1.1 Elements and Compounds ..................................................................................................... 5

1.1.1 Pure Substances, Compounds, Elements ....................................................................... 5

1.1.2 Atom: .............................................................................................................................. 5

1.1.3 Atomic Number .............................................................................................................. 6

1.1.4 Atomic Mass Number ..................................................................................................... 6

1.1.5 Atomic Weight ................................................................................................................ 7

1.1.6 Molecules and Molecular Weights ................................................................................. 8

1.1.7 The mole ....................................................................................................................... 10

1.1.8 Ions and Valence .......................................................................................................... 10

1.1.9 Radicals ......................................................................................................................... 11

1.1.10 Equivalent Weights ..................................................................................................... 11

1.1.11 Using Equivalent weights ........................................................................................... 12

Chapter 2 ....................................................................................................................................... 14

Water Formed Scale ...................................................................................................................... 14

2.1 Introduction ......................................................................................................................... 14

2.2 Scale Types .......................................................................................................................... 14

2.2.1 Calcium Carbonate ....................................................................................................... 15

2.2.2 Calcium Sulfate ............................................................................................................. 18

2.2.3 Barium Sulfate .............................................................................................................. 20

2.2.4 Strontium Sulfate.......................................................................................................... 23

2.2.5 Iron Compounds ........................................................................................................... 25

2.3 Predicting Scale Formation .................................................................................................. 27

2.3.1 The value of Solubility calculations .............................................................................. 27

2.3.2 The Basics for Solubility Calculations ........................................................................... 27

2.3.3 Solubility Product Principle .......................................................................................... 28

2.3.4 Saturation Ratio ............................................................................................................ 29

2.3.5 Calcium Carbonate Scaling Calculations ....................................................................... 29

2.3.6 Sulfate Scaling Calculations .......................................................................................... 34

2.3.7 Barium Sulfate Solubility Calculations .......................................................................... 35





2.3.8 Strontium Sulfate Solubility Calculations ..................................................................... 36

2.4 Example Scale Calculations.................................................................................................. 37

Chapter 3 ....................................................................................................................................... 45

Scale problems & Inhibition .......................................................................................................... 45

3.1 Introduction ......................................................................................................................... 45

3.2 Scale mechanisms ................................................................................................................ 45

3.3 Causes of Scales ................................................................................................................... 45

3.3.1 Change in temperature ................................................................................................ 45

3.3.2 Mixing Incompatible water .......................................................................................... 52

3.3.3 Loss of carbon dioxide (increase in pH) ........................................................................ 53

3.4 Field and Lab Examination of Scales .................................................................................... 54

3.5 Qualitative and Quantitative Tests ...................................................................................... 55

Chapter 4 ....................................................................................................................................... 57

Scale Inhibition .............................................................................................................................. 57

4.1 Theory of Scale Inhibition .................................................................................................... 57

4.1.1 Introduction .................................................................................................................. 57

4.2 Complexation ...................................................................................................................... 57

4.3 Adsorption ........................................................................................................................... 59

4.3.1 Threshold Inhibition ..................................................................................................... 64

4.4 Laboratory Evaluation of Inhibitors ..................................................................................... 65

4.4.1 Precipitation Tests ........................................................................................................ 66





Chapter 1

Fundamentals of Chemistry and water Analysis

1.1 Elements and Compounds

Matter is the “stuff” of which the universe is composed. Matter has two characteristics: it has

mass and it occupies space. Matter is found in three physical “states” (solid, liquid, and gas).

1.1.1 Pure Substances, Compounds, Elements Substance: a particular kind of matter with a definite fixed composition. Substances cannot be

broken down into simpler substances by chemical reactions are called elements. Most

substances are chemical combinations of elements. These are called compounds. Compounds

are made of elements and can be broken down into elements. Properties of the compound are

not usually related to the properties of the elements that compose it. Compounds have the same

chemical composition at all times.

Matter

Mixtures

1- Variable composition

2- May be separated by physical methods

3- Different mixtures have different

properties

4- Example: air

Pure Substances

1- Fixed composition

2- Cannot be separated into simpler substances

3- Can only be changed by chemical methods

4- Properties do not vary

Homogeneous

mixture

Heterogeneous

mixture

Element

Cannot be decomposed

into simpler substance,

i.e., (oxygen- Nitrogen -

Sodium)

Compounds

Can be decomposed

into simpler

substance, i.e., (CO2-

Water)

Table. 1.1. Mixture, substances differences.

1.1.2 Atom: Atoms are the smallest unit particle of an element.

An atom is composed of a positively-charged nucleus orbited by one or more negatively-charged

particles called electrons. A simplified schematic representation of this arrangement is illustrated

in Figure 1.1. The nucleus is the core of an atom. It has a positive charge because it usually

consists of two particles, the neutron and the proton (hydrogen is the exception with only a proton

in the nucleus). The neutrons are electrically neutral, and the protons are electrically positive. A

nucleus with one proton has a charge of +1 (or simply 1), and a nucleus with two protons has a

+2 charge. Together the neutrons and protons give the nucleus its mass, but the proton alone

gives the nucleus its positive charge.

The particles that orbit the nucleus are electrons. They are very small, with a mass only 1/1835

the mass of a proton or neutron. Each electron is negatively charged, and the charge of one

electron is equal in magnitude (but opposite in sign) to the charge of one proton. The number of

electrons orbiting a nucleus is exactly equal to the number of protons contained in that nucleus.

The equal and opposite charges cancel each other, and the atom as a whole is neutral.





The electrons are bound in the atom by electrostatic attraction. The atom remains neutral unless

some external force causes a change in the number of electrons.

Because the nucleus is composed of neutrons and protons that are about 1835 times heavier

than an electron, the nucleus contains practically all the mass of the atom, but constitutes a very

small fraction of the volume. Although electrons are individually very small, the space in which

they orbit the nucleus constitutes the largest part of the atomic volume.

Some of the properties of the atom and its component parts are summarized in Table 1.2. The

masses listed in Table 1.2 are measured in atomic mass units (amu), which is a relative scale in

which the mass of a proton is about 1.0.

Figure. 1.1.Schematic of a Simple Atom (Helium)

Table. 1.2. Properties of the Atom and its Fundamental Particles.

1.1.3 Atomic Number

The number of protons in the nucleus plays such an important role in identifying the atom that it is

given a special name, the atomic number. The symbol Z is often used for atomic number (or

number of protons). Hydrogen has an atomic number of 1. The atomic number is also equal to

the number of electrons.

1.1.4 Atomic Mass Number

The sum of the total number of protons, Z, and the total number of neutrons, N, is called the

atomic mass number. The symbol is A. Not all atoms of the same element have the same atomic

mass number, because, although the Z is the same, the N and thus the A are different. Atoms of

the same element with different atomic mass numbers are called isotopes.





1.1.5 Atomic Weight

In Table 1.2, the masses of atomic particles are given in atomic mass units (amu). These units

represent a relative scale in which the mass of the isotope carbon-12 is used as the standard and

all others are related to it. Specifically, 1 amu is defined as 1/12 the mass of the carbon-12 atom.

Since the mass of a proton or a neutron is approximately 1 amu, the mass of a particular atom

will be approximately equal to its atomic mass number, Z.

The atomic weight of an element is generally more useful than isotopic masses. The atomic

weight of an element is defined as the weighted average of the masses of all of its natural

occurring isotopes. On this scale, the atomic weight of hydrogen (H) is 1.0 amu, that of helium

(He) is 2.0 amu, and that of magnesium is 24.3 amu. This tells us that the He atom is twice the

mass of H atoms while the Mg atoms are about 24 times heavier than H atoms.

Table. 1.3. gives the atomic weight of selected elements.

Table 1.3. Symbols and atomic weight of selected elements.





Figure 1.2. Illustrates Hydrogen, Helium , and Magnesium atoms.

Figure. 1.3. Periodic table of elements.

Figure. 1.2. Hydrogen, Helium , and Magnesium atoms.

1.1.6 Molecules and Molecular Weights

Atoms combine to form molecules, which can be defined as the smallest particle of a compound

which can exist. Thus, atoms have the same relation to elements as molecules have to

compounds.

A molecule of a compound is formed by the union of two or more atoms of the elements of which

the compound is composed. The molecular weight (or formula weight) is the relative mass of a

single molecule compared to the mass of carbon-12 atom. Molecular weight is found by adding

the atomic weights of the elements which form the molecule. An example of calculation is given in

table. 1.4.

Table .1.4. Molecular weight calculation.

Molecular weights, like atomic weights, are also relative weights. A molecular weight of 18 means

that one molecule of water is 18/12 = 1.5 times as heavy as one atom of carbon.

Another example: The chemical formula of sulfuric acid is H2SO4, which means that it contains 2

atoms of hydrogen, 1 atom of of sulfur and 4 atoms of oxygen. Therefore, its molecular weight is:

(2x1)+32+(4x16) = 98

Thus, a mole of H2SO4 contains 98 units of mass, of which two are hydrogen, 32 are sulfur, and

64 are oxygen.





Figure.1.3. Periodic table.





1.1.7 The mole A mole is defined as the amount of substance that contains as many elementary entities (atoms,

molecules, or other particles) as there are atoms in 12 grams of pure carbon-12 isotope. “the

number equal to the number of carbon atoms in 12 grams of carbon.”

One mol of something consists of 6.022 x 1023 units of that substance.

According to the definition, one mole refers to 6.02 x 1023 “elementary entities” whose identities

must be specified. We could have a mole of atoms or a mole of molecules (or a mole of ions,

electrons or any particles).

The mass of one mole of atoms of a pure element is numerically equal to the atomic weight of

that element (in amu’s), expressed in grams. For example, 1 mole of elemental sodium contains

6.02 x 1023 sodium atoms and has a mass of 23.0 grams.

Similarly, one mole of ions contains 6.02 x 1023 ions, and this number of ions has a mass equal to

the atomic weight or formula mass expressed in grams. Hence, 1 mole of Mg+2 contains 6.02 x

1023 magnesium ions, has a mass of 24.3 grams.

Another example: If you obtain a sample of pure gold, which has an atomic weight of 197 amu,

and weight 197 grams, you will have one mole of gold, which contains 6.02 x 1023 atoms of gold.

Thus, one atom of gold weights:

197/(6.02 x 1023) = 3.27 x 10-23 g

When dealing with gases, we have to be careful, since some gases exist as atoms while others

exist as molecules containing two atoms. Oxygen commonly exists as O2 molecules, so one mole

of oxygen has the mass of 32.0 grams contains 6.02 x 1023 molecules. This concept is further

illustrated in table. 1.5.

Unless otherwise stated, the mass contained in a mole is expressed in grams. However, moles

can also be expressed in other units, such as pounds, tons, or kilograms.

Table. 1.5. Mass of one mole of atoms of some common elements.

Using carbon as an example:

1 mole C = 12 g

1 lb-mole C = 12 lb. = 5448 g = 5448/12 = 454 g-mol

1 ton-mole C = 12 tons. = 24000 lb = 2400/12 = 2000 lb-mole or 908000 g-mole

1.1.8 Ions and Valence As previously stated, atoms are made up of protons, neutrons, and electrons. Each atom has an

equal number of protons and electrons, so their charge balance out and the net charge is zero. If,

however, an atom should loss or gain electrons, an imbalance in charge will result since the

number of protons in nucleus stays constant. Therefore, if an atoms gains electrons it will have a





net negative charge. If it loses electrons it will have a net positive charge. Whenever this

happens, the atom is no longer called an atom. It becomes an ion which may be defined as an

atom or group of atoms containing an electrical charge. A positively charged ion is called a cation

while a negatively charged ion is called an anion.

The amount of charge is called the valence and is a measure of the element’s chemical

combining power. When hydrogen ionizes it loses its electron and has a net positive charge or

valence of +1.

H – e- → H+

Calcium ionizes by losing 2 electron and thus the calcium ion has a valence of +2.

Ca – 2e- → Ca2+

Chlorine ionizes by gaining 1 electron and chlorine ion has a valence of -1.

Cl + e- → Cl-

The formula of hydrochloric acid, HCl, tells us that one atom of hydrogen combines with one atom

of chlorine. The sum of the valence must be zero since compounds have a net charge of zero.

H+ + Cl- → = HCl ……. (+1) + (-1) = 0

Similarly for salts

Ca2 + 2Cl- → CaCl2 ……. (+2) + [2 x( -1)] = 0

1.1.9 Radicals A radical is a group of atoms found in a certain compounds which reacts as a unit, i.e., as if it

were a single atom or ion. The radicals in the compounds following are enclosed in parentheses.

H2SO4 --- CaCO3 ---- BaSO4

If these compounds were ionized we would find that the radicals behaves as polyatomic ions:

SO4-2 (sulfate ion) ------------------ CO3

-2 (Carbonate ion)

1.1.10 Equivalent Weights When elements combine to form a given compound they do so in a fixed and invariable ratio by

weight. This ration can be predicted by means of equivalent weight. For and element or ion:

Equivalent weight = Atomic weight / Valence

The valence of an element is the charge it exhibits when it is ionized.

Some elements, such as iron, have more than one equivalent weight because it can exist in more

than one valence state. For example, iron has two different equivalent weights because it can

exist as either ferrous ions (Fe+2) or ferric ion (Fe+3).

The equivalent weight of the compound is:

Equivalent weight = Molecular weight / Net positive valence

A compound has no charge. The net positive valence of a compound is the total number of

positive charges that would result if the compound were ionized in aqueous solution.





1.1.11 Using Equivalent weights One of the main uses of equivalent weight is to permit us to easy determine the quantity of

reactants and products in a chemical reaction.

If two elements, A and B will combine to form compound C, they will do so on an equivalent

basis. One equivalent of A will combine with one equivalent of B to form one equivalent of C.

For example:

1 Equivalent Ca+2 + 1 Equivalent CO3-2 = 1 Equivalent CaCO3

(20 ) + (30 ) = ( 50)

Similarly for ferric chloride:

Fe+3 + 3Cl- = FeCl3

Weight reacting = 56 + 3 (35.5) = 162.5

Equivalent weight = 18.7 + 35.5 = 54.2

No of equivalents = 3 + 3 = 3

We see that one atomic weight of irons combined with three atomic weights of chlorine to give

one mole of FeCl3. More important, however, 3 equivalents of iron have combined with 3

equivalents of chlorides to give 3 equivalents of ferric chloride. Equivalents always combine on

one-to-one basis.

Some examples of equivalent weights are given in tables 1.5, 1.6, and 1.7.





Table. 1.5 .Equivalent weight of some elements and compounds.







Chapter 2

Water Formed Scale

2.1 Introduction Solubility is defined as the limiting amount of a solute which can be dissolved in a solvent under a

given set of physical conditions.

The chemical species of interest to us are present in aqueous solutions as ions. Certain

combinations of these ions are compounds which have very little solubility in water. The water

has a limited capacity for maintaining these compounds in solution, and once this capacity, or

solubility is exceeded, the compound precipitate from solutions as solids. Therefore, precipitation

of solid materials which may form scale can occur if both of the following conditions are satisfied:

1- The water contains ions which are capable of forming compounds of limited solubility.

2- There is a change in physical conditions or water composition which lowers the solubility

below the concentration presents.

Solid precipitates may either stay in suspension in water, or they may form a coherent scale on

the surface such as a pipe wall. Formation plugging may occur by filtration of suspended particles

from the water. Alternatively, a solid scale may form on the formation face. Either is undesirable.

The difficulty of removal varies with the type of plugging which has occurred.

Scale formation frequently restricts flow through injection and flow lines, and tubing strings. It

cause pump wear or plugging and creates additional rod load when formed on sucker rods. Fire

tubes in all types of heaters fail prematurely when scale formation results in overheating.

Corrosion is often more severe under scale deposit.

Water formed scales are responsible for many production problems and their effective control

should be one of the primary objectives of any efficient water handling operation.

2.2 Scale Types Few types of many scales are found in oilfield waters. These scales are listed in table ….

Table. 2.1. Common oilfield scale.





2.2.1 Calcium Carbonate Calcium carbonate scale can be formed by the combination of calcium and other carbonate or

bicarbonate ions as follows:

Ca+2 + CO3-2 → CaCO3 ↓ Eq. 2.1

Ca+2 + 2(HCO3-) → CaCO3 ↓ + CO2 + H2O Eq. 2.2

2.2.1.1 Effect of CO2 Partial Pressure The presence of CO2 increases the solubility CaCO3 in water. When carbon dioxide dissolves in

water, it forms carbonic acid, which ionizes according to the following series of equations:

CO2 + H2O ↔ H2CO3 Eq. 2.3

H2CO3 ↔ H+ + HCO3- Eq. 2.4

HCO3- ↔ H+ + CO3

-2 Eq. 2.5

Only a small percentage of the bicarbonate ions dissociate at the pH values found in most

injection water to form H+ + CO3-2 , as shown in figure 2.1.

Bicarbonate ions are usually greater than carbonate ions present under normal circumstances.

Therefore, it is thought that equation (2.2), is the more accurate expression for the precipitation of

calcium carbonate.

As the concentration of CO2 in solution is increased, the reaction shifts to the left, resulting in less

CaCO3 precipitation. The water also becomes more acidic (the pH decreases) with the addition of

CO2 to the water.

The amount of CO2 that will dissolve in water is proportional to the partial pressure of CO2 in the

gas in contact with the water.

Partial Pressure of CO2 = (Mole Fraction of CO2 in Gas) x (Total Pressure) Eq. 2.6

Mole Fraction of CO2 in Gas = %CO2 in Gas /100 Eq. 2.7

Hence, if a two-phase (gas + water) system is operating at 100 psia and the associated gas

contains 10 mol % CO2 , the partial pressure of CO2 in the gas is:

(0.1)x(100) = 10 psia.

So, if either the system pressure or the percentage of CO2 in the gas were to increase, the

amount of CO2 dissolved in the water also would increase.

Figure 2.2 illustrates the effect of CO2 partial pressure on the pH of water containing little or co

dissolved minerals. This data should not be applied to brines since the presence of dissolved

minerals changes the relationship between pH and the amount of dissolved CO2. The effect of

CO2 pressure on the solubility of CaCO3 in pure water is shown in figure 2.3.

This data illustrated that CaCO3 solubility increases with increase CO2 partial pressures. The

effect becomes less pronounced as the temperature increases.

The reverse is also true. It is one of the major cause of CaCO3 scale precipitation. At any point in

the system where a pressure drop is taken, the partial pressure of CO2 in the gas phase

decreases, CO2 comes out of solution, and the pH of the water rises. This shifts reaction 2.2 To

the right and may cause CaCO3 precipitation.





Figure 2.1. Ionization of carbonic acid as a function of pH.

Figure 2.2 Effect of CO2 partial pressure on the pH value of water.





Figure 2.3. Effect of CO2 partial pressure on calcium carbonate solubility.

2.2.1.2 Effect of Total Pressure The solubility of calcium carbonate is a two phase system increases with increase pressure for

two reasons:

1- Increased pressure increases the partial pressure of CO2 and increases the solubility of

CaCO3 in water as previously explained.

2- Increased pressure also increases the solubility due to the thermodynamic consideration.

Pressure drops are one of the primary causes of calcium carbonate scale deposition in

production systems. In addition to decreasing the solubility of CaCO3 due to the loss of CO2 and

the thermodynamic pressure effect, pressure drops across chokes and valves induce turbulence

in the water which helps to overcome supersaturation effects and initiate precipitation.

In single phase (all water) systems, such as water injection system, increased pressure increases

the solubility of calcium carbonate solely due to thermodynamic considerations.

2.2.1.3 Effect of Temperature Contrary to the behavior of most materials, calcium carbonate becomes less soluble as

temperature increases- The hotter the water gets, the more likely CaCO3 scale will form.

Hence, a water which is non-scaling at the surface may result in scale formation in an injection

well if the downhole temperature is sufficiently high. This is also the reason that CaCO3 scale is

often found on the fire-tubes of heating equipment.

The solubility of CaCO3 scale in pure water at 1 atmospheric CO2 partial pressure as a function of

temperature os shown n figure 2.4.

Methods for calculating the temperature at which CaCO3 scale may be anticipated are given in a

later section.





Figure 2.4 Calcium carbonate solubility.

2.2.1.4 Effect of Dissolved Salts Calcium carbonate solubility increases as the salt content of the water increases. For instance,

adding 200,000 mg/L NaCl to distillate water increases the CaCO3 solubility from 100 mg/L to 250

mg/L.

Actually the higher the total dissolved solids (not counting calcium or carbonate ions), the greater

the solubility of CaCO3 in the water and lower the scaling tendency up to a maximum of about

200,000 mg/L.

2.2.1.5 Summary of Factors affecting CaCO3 scale formation In summary the, the likelihood of forming calcium carbonate scale:

Increases with temperature

Increases with partial pressure of CO2 decreases.

Increases as pH increases.

Increases as total dissolved salts decreases.

Increases as total pressure decreases.

2.2.2 Calcium Sulfate The precipitation of calcium sulfate from water results from the reaction:

Ca+2 + SO4-2 → CaSO4 ↓

2.2.2.1 Forms of Calcium Sulfate

Many calcium sulfate deposits found in the oilfield are gypsum. According to Oddo and Tomson,

the most likely scale to form from brines will be gypsum at temperatures less than 1760F (800C).

Between 1760F (800C) and 2500F (1210C), any of the three types of calcium sulfate may form,

with gypsum being more likely at the low end of the temperature range and anhydrite more likely

at the high end. Hemi-hydrate is commonly found in this temperature range in non-turbulent

systems with high ionic strengths. Above 2500F (1210C), any calcium sulfate formed will almost

certainly be anhydrite.





2.2.2.2 Effect of Temperature Gypsum solubility in pure water increases with temperature up to about 1000F (380C), then

decreases with temperature as shown in figure 2.5.

This is a quite different from the temperature-solubility behavior of CaCO3. First, gypsum is

considerably more soluble than CaCO3 in the normal temperature range of interest. Second, the

maximum in the gypsum curve tells us that an increase in temperature could either increase or

decrease the solubility of gypsum depending on which part of the temperature curve we’re

concerned with. This is decidedly different from CaCO3 where an increase in temperature always

decreases the solubility.

Note that above about 1000F (380C), anhydrite becomes less soluble than gypsum, so it could

reasonably be expected that anhydrite might be the preferred form of CaSO4 in deeper, hotter

wells.

Actually, the temperature at which scale changes form from gypsum to anhydrite or hemi-hydrate

is a function of many factors, including pressure, dissolved salt content, flow conditions, and the

speed at which different forms of CaSO4 can precipitate from solution.

Predicting which form of calcium sulfate will precipitate under given set of conditions is very

difficult. Even though anhydrite would be expected above 1000F (380C), in preference to gypsum

due to its lower solubility, gypsum may be found at temperature up to 2120F (1000C).

It is often difficult to precipitate anhydrite directly from solution. However, with the passage of

time, gypsum can dehydrate to form anhydrite.

Above 2120F (1000C), anhydrite will precipitate directly in a stirred or flowing stream. If the system

is quiescent, hemi-hydrate solubility becomes limiting. Conversion to anhydrite could be expected

with time.

Figure 2.5 Solubility of calcium sulfate in pure water.





2.2.2.3 Effect of Dissolved Salts The presence of NaCl or dissolved salts other than calcium or sulfate ions increases the solubility

of gypsum or anhydrite just as it does for calcium carbonate up to a salt concentration of about

150,000 mg/L. Further increases in salt content decreases CaSO4 solubility. (Refer to figure 2.6)

2.2.2.4 Effect of Pressure Increase pressure increases the solubility of all forms of calcium sulfate due to the

thermodynamic considerations.

Pressure drops are one of the primary causes of calcium sulfate scale deposition in production

systems. In addition decreasing the solubility due to the thermodynamic considerations, pressure

drops across chokes and valves induce turbulence in the water which helps to overcome

supersaturation effects and initiate precipitation.

The effect of pressure and temperature on anhydrite solubility is shown in figure 2.7. Note that

pressure effect decreases as temperature increases.

2.2.2.5 Effect of pH pH has a little or no effect on the solubility of calcium sulfate.

Figure 2.6 Solubility of gypsum in NaCl brines.

2.2.3 Barium Sulfate Barium Sulfate is the least soluble of the scales we have discussed thus far.

Ba+2 + SO4-2 → BaSO4 ↓

Table 2.2 compares the solubility of the three scales mentioned thus far, in distilled water at 770F

(250C)

Scale Solubility (mg/L)

Gypsum 2080

Calcium Carbonate 15

Barium Sulfate 2.3

Table 2.2 Comparative solubilities at 250C.





The extreme insolubility of BaSO4 makes it very likely that scaling will occur if both Ba+2 and

SO4-2 ions are present in water.

Most barium sulfate scales also contain some strontium sulfate.

Figure 2.7. Effect of pressure and temperature on CaSO4 (anhydrite) solubility.

2.2.3.1 Effect of Temperature Barium sulfate solubility increases with temperature up to to 2120F (1000C). The solubility in

distilled water increases from 2.3 mg/L at to 770F (250C) to 3.9 mg/L at to 2030F (950C) as shown

in figure 3.8. The increase is fairly substantial percentage wise, but barium sulfate is still quite

insoluble even at this higher temperature.

Above to 2120F (1000C).the solubility decreases with temperature in water with TDS (Total

dissolved salts) less than about 50,000 ppm. The solubility in higher salinity waters shows

normal solubility behavior and increases with temperature as shown in figure 3.9.

Because of the increase in solubility over normal temperature ranges, barium sulfate usually

presents no downhole scaling problems in injection well if it is non-scaling at surface conditions. It

is more commonly a problem in producing or water supply wells.

2.2.3.2 Effect of Dissolved Salts The solubility of barium sulfate in water is increased by foreign dissolved salts just as the in the

case of calcium carbonate and calcium sulfate. The addition of 100,000 mg/L of NaCl to distilled

water increases the solubility of BaSO4 from 2.3 mg/L to 30 mg/L at 770F (250C). Maintaining

100,000 mg/L of NaCl and increase the temperature to 2030F (950C), will increase BaSO4

solubility to about 65 mg/L.





Figures 2.8 and 2.9 illustrates the effect of salinity on BaSO4 solubility over a wide range of

temperatures. The data in figure 2.8 is plotted as a function of molar ionic strength rather than

sodium chloride concentration. An ionic strength of 1.0 is equivalent to a sodium chloride

concentration approximately 60,000 mg/L. Ionic strength is defined later in this chapter with a

calculation example.

As a rule of thumb, you can assume that BaSO4 solubility will double as the temperature is raised

from 770F (250C) to 2030F (950C) regardless the dissolved salt concentrations. The effect of

dissolved salts is much more pronounced, as mentioned above in the addition of 100,000 mg/L

NaCl example.

Figure 2.8 Barium sulfate solubility in NaCl solutions.





Figure 2.9 Effect of temperature on Barium sulfate solubility at atmospheric pressure.

2.2.3.3 Effect of Pressure Increased pressure increases the solubility of barium sulfate due to the thermodynamic

considerations.

Pressure drops are one of the primary causes of barium sulfate scale deposition in production

systems. In addition decreasing the solubility due to thermodynamic considerations, pressure



Figure 2.10 shows the effect of pressure and temperature on barium sulfate solubility in pure

water.

2.2.3.4 Effect of pH pH has a little or no effect on the solubility of barium sulfate.

2.2.4 Strontium Sulfate Strontium Sulfate is considerably more soluble than barium sulfate, with a solubility of 129 mg/L

in pure water at a temperature of 770F (250C).

Sr+2 + SO4-2 → SrSO4 ↓

2.2.4.1 Effect of Temperature Strontium Sulfate solubility decreases with temperature. In pure water the solubility decreases to

68 mg/L at 2570F (1250C).

2.2.4.2 Effect of Dissolved Solids The solubility of SrSO4 in water increases as the NaCl content of the water increases up to a

maximum of approximately 175,000 mg/L. Further increases in salinity result in decreasing

solubility. In brines containing calcium or magnesium, the apparent solubility of SrSO4 is greater

than a NaCl brine of equivalent ionic strength. Refer to figure 2.11.





Figure 2.10 Effect of pressure and temperature on barium sulfate solubility in pure water.

Figure 2.11. Strontium sulfate solubility in NaCl solution.

2.2.4.3 Effect of Pressure

Strontium sulfate solubility in NaCl brines slightly increases with pressure.

Pressure drops are one of the primary causes of strontium sulfate scale deposition in production

systems. In addition decreasing the solubility due to thermodynamic considerations, pressure







Until the advent of seawater injection in the middle east, pure SrSO4 scales was seldom

observed and was not considered as a major problem in water injection operations. However,

serious SrSO4 scale problems have occurred in producing wells in number of middle east fields

after breakthrough of seawater, due to mixing of sulfate-bearing seawater and the strontium in

formation waters in the producing wellbores.

In a majority of cases, however, strontium co-precipitate with barium to form (Ba,Sr) SO4 scale.

2.2.4.4 Effect of pH pH has a little or no effect on the solubility of strontium sulfate.

2.2.5 Iron Compounds Iron ions present in water may be either naturally present in the water or the result of corrosion.

Formation waters normally contain only a few mg/L of natural iron and values as high as 100

mg/L are rare. Higher iron contents are invariably the result of corrosion. Precipitated iron

compounds are a common cause of deposit formation and injection well plugging, as well as

being indicative of serious corrosion problem.

Corrosion is usually the result of CO2, H2S, or oxygen dissolved in the water. Most of the scales

containing iron are corrosion products. However, iron compounds can also form by precipitation

of natural formation iron even if corrosion is relatively mild.

Carbon dioxide can react with iron to form iron carbonate scale. Whether or not scale actually

forms will depend on the pH of the system. Scale formation is much more likely above pH 7.

Figure 2.12 illustrates the solubility of iron carbonate in fresh water.

Figure 2.12 Iron carbonate solubility diagram.





Hydrogen sulfide will form iron sulfide as a corrosion product which is quite insoluble and usually

forms a thin, adherent scale. Suspended iron sulfide is the cause of “black water”.

The iron sulfide diagram in figure 2.13 illustrates the concentration of Fe2+ (ferrous ion) which will

stay in solution at various pH values and H2S concentration in fresh water. Fe3+ (ferric ion) is

seldom found at pH values above 3.0.

Oxygen combines to form several compounds. Ferrous hydroxide, Fe(OH)2, ferric hydroxide

Fe(OH)3, and ferric oxide Fe2O3, are common scales resulting from contact with air. For example,

when air-free water containing dissolved ferrous iron ions and bicarbonate ions is contacted by

air, ferric hydroxide can be formed.

2Fe+2 + 4HCO3- + H2O + ½ O2 → 2Fe(OH)3 ↓ + 4CO2

Ferrous iron (Fe+2) is oxidized in the presence of air to give ferric iron (Fe+3), and ferric hydroxide

results. This is practically insoluble above pH 4. As shown in figure 2.14, if oxygen can be

excluded, 100 ppm of Fe(OH)2 “ferrous hydroxide” will still in solution at pH 8.5.

“Red water” is the result of suspended particles of Fe2O3, another product of oxygen and iron.

Iron compounds can also result from the action of certain bacteria (gallionella ferrginea) which

live in water in the presence of air. These bacteria take Fe+2 ions from the water and deposit ferric

hydroxide.

In summary, the chemistry of iron compounds is much more complex than that of previously

discussed compounds. This is due primarily to the fact that iron commonly exists in two oxidation

states in water, Fe+2 (ferrous) and Fe+3 (ferric). These two ions form compounds with the same

anions that possess very different solubilities. It is difficult to quantitatively predict the behavior of

iron compounds. It is far more important to prevent their formation.

Figure 2.13 Iron sulfide stability diagram.





Figure 2.14 Iron hydroxide stability diagram.

2.3 Predicting Scale Formation

2.3.1 The value of Solubility calculations Solubility calculations or scaling indexes may be used to predict the formation of certain types of

scales. The value obtained from these calculation procedures should be taken merely as

guidelines. They indicate the degree of “scaling tendency”, or the likelihood of scale formation.

Simplifying assumptions have been made in the derivation of each equation; solubility in naturally

occurring waters is a complex phenomenon.

It should be emphasized that if scale formation is indicated by calculation, it serves as an alarm. If

you are looking at possible water source, you should avoid those which show scaling tendencies

or make a provision for treatment in your planning. Similarly, one should avoid mixing waters

which would result in a composite analysis which exhibits scaling tendencies under system

conditions.

A calculated scaling tendency in an existing system should focus attention on the fact that scale

formation is likely and you should begin inspecting the system for signs of scale formation

immediately. The following section is a summary of some of the published scale prediction

equations, which are used for oil field waters.

2.3.2 The Basics for Solubility Calculations All scale prediction methods are based on laboratory measurements of the solubility of a specific

compound at equilibrium conditions. Normally, solubilities are measures in synthesized water

over a range of temperature and atmospheric pressure.





2.3.3 Solubility Product Principle When a sparingly soluble salt is added to water, cations and anions from the crystal lattice of the

solid pass into solution until the solution becomes saturated. In the saturated solution the

equilibrium exists between the ions in solution and ions present in the solid crystal lattice.

Using barium sulfate (BaSO4) as an example:

BaSO4 → Ba+2 + SO4-2

Solid Solution

At a given temperature and pressure the product of the activities of the ions in the saturated

solution is constant and is called the “Thermodynamic solubility product constant”, KSP.

KSP = a Ba+2 x a SO4 -2 Eq. 2-8

Where

a Ba+2 = Barium ion activity

a SO4 -2 = Sulfate ion activity

The activity of an ion is defined as the product of the ion concentration and the activity coefficient.

a Ba+2 = C Ba+2 x Y Ba+2 Eq. 2.9

a SO4 -2 = C SO4 -2 x Y SO4 -2 Eq. 2.10

Where

C Ba+2 = Ba+2 Concentration, moles/L

C SO4 -2 = SO4 -2

Concentrations, moles/L

Y = activity coefficient, where it is a function of temperature, pressure, and ionic strength.

At saturation, the resulting solubility product is defined as the ion product constant, or “conditional

solubility product constant, Kc” and is defined in equation 2.11.

Kc = KSP / ( Y Ba+2 x Y SO4 -2 ) = (C Ba+2)( C SO4 -2 ) at saturation Eq. 2.11

Where

Kc = conditional solubility product constant, molar units.

Based on this principle, it is possible to evaluate a solution with respect to the possibility of the

precipitation of a given salt at a given temperature and pressure by comparing the value of ion

product constant at those conditions with the product of the measured concentrations of the ions

in the solution.

If we have a solution which contains a given amount of dissolved BaSO4, and the measured

concentrations of Ba+2 and SO4-2 are CBa+2 and CSO4 -2 , the possibilities are as follows:

1. The solution is saturated with BaSO4 (Solubility limit)

(C Ba+2)( C SO4 -2 ) = Kc Eq. 2.12

2. The solution is under saturated with BaSO4. . Precipitation cannot occur

(C Ba+2)( C SO4 -2 ) < Kc Eq. 2.13





3. The solution is supersaturated with BaSO4. Precipitation can occur

(C Ba+2)( C SO4 -2 ) > Kc Eq. 2.14

Thus, precipitation can occur only in the last case.

2.3.4 Saturation Ratio Continuing with the example of BaSO4 dissolved in water, the saturation ratio (sometimes called

the supersaturation ratio) is defined as the ratio of iron product to the ion product constant.

SR = IP/Kc = (C Ba+2)( C SO4 -2 ) / Kc Eq. 2.15

Where

IP = Ion product

SR = Saturation ratio

(C Ba+2) and (C SO4 -2 ) = Measured concentration of Ba+2 and SO4-2 in solution.

Thus, it is also possible to express the conditions necessary for precipitation in the term of the

saturation ratio:

1. SR = 1 The solution is saturated with BaSO4.

2. SR < 1 The solution is undersaturated with BaSO4.

Precipitation cannot occur

3. SR > 1 The solution is supersaturated with BaSO4.

Precipitation can occur

2.3.5 Calcium Carbonate Scaling Calculations

2.3.5.1 Calcium Carbonate Saturation Indexes It is common to express the degree of supersaturation, and hence the likelihood of precipitation of

CaCO3 from the solution in terms of the saturation index, which is defined as follows:

Saturation Index = log10 (IP/Kc ) Eq. 2.16

Although there is agreement as to the definition of the saturation index “stability index”, equations

uses two abbreviations (SI) and (Is).

The saturation index, which we will refer to as SI, is a measure of the degree of supersaturation,

and thus the driving force available to cause precipitation. The larger the value of SI, the greater

the likelihood that scale will occur. It does not predict the amount of scale which will precipitate.

Figure 2.15 illustrates the relationship between the supersaturation ratio and the saturation index.

2.3.5.2 Langelier Saturation Index This is a well known index was developed to predict whether a fresh saturated with dissolved

oxygen would form calcium carbonate scale or be corrosive.

The Langelier saturation index is calculated from the following empirical equation:

SI = pH - pHs Eq. 2.17

where

SI = Stability Index





pH = Actual pH of water

pHs = pH at which water would be saturated with CaCO3.

pHs = pCa + pAlk M + Ct Eq. 2.18

pCa = log [1 / (Moles Ca+2 /liter)] Eq. 2.19

pAlk M = log 1 / (equivalent M alkalinity / liter) Eq. 2.20

M alkalinity = total alkalinity = CO3-2 + HCO3

- , equivalent / L Eq. 2.21

where

Ct = a constant which is a function of the total dissolved solids and temperature.

As previously explained, if SI > 0, precipitation of CaCO3 is indicated. A negative value of SI

indicates that the water is corrosive if dissolved oxygen is present. This index indicates the

tendency of ware to precipitate calcium carbonate, but it does not indicate the amount of

precipitate.

Figure 2.15 Saturation index (SI) versus saturation ratio (SR).

2.3.5.3 Ryznar Stability Index Ryznar developed an empirical equation for calculating the stability index of fresh water at

atmospheric pressure.

SI = 2pHs – pH Eq. 2.22

Ryznar’s stability index value always positive. They can be interpreted as follows:

1- Stability index < 6.5 indicates CaCO3 scale formation. The smaller the index, the larger

the amount of scale formed.

2- Stability index > 6.5 indicates corrosion if dissolved oxygen is present. The larger the

index, the more sever the anticipated corrosion.





2.3.5.4 Stiff and Davis Method Stiff and Davis empirically extended the Langelier method to apply to oilfield brines. Their

equation is as follows:

SI = pH - pHs Eq. 2.23

where

SI = Stability Index

pH = Actual pH of water

pHs = pH at which water would be saturated with CaCO3.

pHs = K + pCa + pAlk Eq. 2.24

SI = pH - K - pCa - pAlk Eq. 2.25

Where

SI = Stability index. If SI is negative, the water is undersaturated with CaCO3 and scale formation

is unlikely. If SI is positive, scale is likely to form.

K = A constant which is a function of salinity, composition and water temperature. Values of K are

obtained from a graphical correlation with ionic strength and temperature of water. (Figure 2.16)

pCa = log [1 / (Moles Ca+2 /liter)] Eq. 2.26

pAlk M = log 1 / (equivalent M alkalinity / liter) Eq. 2.27

M alkalinity = total alkalinity = CO3-2 + HCO3

- , mg/ L Eq. 2.28

u =0.5 (C1Z12+ C2Z2

2+ C3Z32 +…..+ CnZn

2) Eq. 2.29

where

C = Concentration of the ion in moles/liter.

Z = Valence of the ion.

In order to calculate SI we must know the temperature, pH, and the HCO3- and CO3

-2

concentrations. In addition, a complete water analysis is necessary to enable calculation of ionic

strength.

It is essential that values of pH, HCO3- and CO3

-2 be measured in the field immediately after

sampling, since these parameters change very quickly once the sample is removed from a

pressurized system. Valid calculations cannot be made from laboratory analyses.

Unfortunately, even field measurements of pH values will not suffice when attempting to apply

this method to downhole conditions in production and injection wells. The pH must be calculated.

Equations are given in the following section which enable estimation of pH values at elevated

pressures and temperatures.

Values of K as a function of ionic strength in Figure 2.16, the curves are based on experimental

data in the following range:

Molar ionic strength: 0 – 3.6

Temperature: 32, 86, and 1220F (0,30 and 500C)

Pressure: 1 atmosphere(14.7 psig , 101.3 KPa).

All curves outside of this data range were extrapolated.

Figure 2.17is a chart for determination of pCA and pAlk .





The results of calculations may be summarized as follows:

If SI is negative, the water is undersaturated with CaCO3 and scale formation is unlikely.

If SI is positive, the water is supersaturated with CaCO3 scale is likely to form.

If SI = 0, the water is saturated with CaCO3.

An example calculation is given later in this chapter.

2.3.5.5 Oddo and Tomson Method The equation developed by Oddo and Tomson enable the calculation of the saturation index. Is,

and considers the effect of total pressure as well as varying CO2 partial pressure. They also

developed equations which permits the calculation of pH.

The equation here were published 1994.

Any System (Gas Phase Present or Absent) where the pH is Known

Is = log [(Ca+2)(HCO3-)]+pH–2.76+9.88 x10-3 T + 0.61x10-6 T2 –3.03x10-5 P– 2.348 u0.5

+0.77 u Eq. 2.30

Where

Ca+2 = Calcium ion concentration, moles/L

HCO3- = Bicarbonate ion concentration, moles/L

T = Temperature, 0F

P = Total absolute pressure, psia

u = Molar ionic strength, moles/L

Gas Phase Absent

These equations are applicable in water injection system and in production systems where the

system pressure is greater than the bubble point pressure of the fluids.

1- Determine Caq, the amount of CO2 dissolved in the water. This can be determined directly

by on-site titration, or it can be calculated using equation 3.31.

Caq = log PCO2 - 2.212 – 6.51 x 10-3 T + 10.19 x 10-6 T2 – 1.29 x 10-5 P – 0.77 u0.5

– 0.059 u Eq. 2.31

2- Calculate Is or the pH as desired from the following two equations

Is = log [(Ca+2)(HCO3-) / Caq ]+3.63 +8.68 x10-3 T+8.55x10-6 T2 –6.56x10-5 P– 3.42 u0.5

+1.373 u Eq. 2.32

pH = log [(HCO3-) / Caq] + 6.39 – 1.198 x 10-3 T + 7.94 x 10-6 T2 – 3.53 x 10-5 P – 1.067

u0.5 +0.599 u Eq. 2.33

3- It is also possible to calculate the change in Is or pH in a system without a gas phase

using the following equations:

ΔIs = 8.68 x 10-3 ΔT + 8.55 x 10-6 Δ(T2) - 6.56 10-5 ΔP Eq. 2.34

ΔpH = 1.198 x 10-3 ΔT + 7.94 x 10-6 Δ(T2) – 3.53 10-5 ΔP Eq. 2.35





Gas Phase is Present and the pH is Unknown

1- Calculate fg, the fugacity coefficient of CO2 gas.

Eq. 2-36

2- Calculate yg, the mole fraction of CO2 in the gas phase at the specified T and P. Given

that yt is the mole fraction of CO2 in the gas at the surface.

Eq. 2.37

where

BWPD = Barrels of water per day

BOPD = Barrels of oil per day

MMscfd = Million standard cubic feet per day.

3- Calculate the ionic strength.

Eq. 2.38

Where all ions concentrations are in mg/L and are obtained from a water analysis.

4- Calculate Is or the pH as desired.

Eq. 2.39

Eq. 2.40

This method is said to be valid over the following data range:

Molar ionic strength: 0-4.0

Temperature: 32-392 0F (0-2000C)

Pressure 0-20,000 psig





2.3.5.6 Estimation of the amount of CaCO3 scale formed It is possible to estimate the maximum amount of scale which could form, assuming the system is

at equilibrium.

If a solution is supersaturated with a salt (such as CaCO3, CASO4, BaSO4 or SrSO4), precipitation

can be expected.

Stiff and Davis

Maximum amount of scale can be precipitated in pounds per thousand barrels (PTB) =

Eq. 2.41

Where

G = Ca+2 + HCO3-, moles/L

X = Ca+2 - HCO3-, moles/L

Table. 2.3. Calcium carbonate scaling severity

Oddo and Thompson

Eq. 2.42

Eq. 2.43

2.3.6 Sulfate Scaling Calculations Traditional approach to Sulfate Scale Solubilities

Solubility values for CaSO4, BaSO4 or SrSO4, can be calculated using the following equation,

providing values of conditional solubility products, Kc, are known for each compound:

Eq. 2.44

The actual concentration of CaSO4 in solution is equal to the smaller of the Ca+2 or SO4-2

concentrations (expressed in meq/liter) in the water of interest, since the smaller concentration

controls the amount of calcium sulfate which can be formed.

The calculated calcium sulfate solubility, S (meq/liter), is compared with the actual concentration

to determine if scale formation is likely.

This formula can be used to calculate the solubility of any divalent salt such as CaSO4, BaSO4 or

SrSO4.





Table. 2.4. Interpretation of sulfate scale calculations

2.3.6.1 Calcium Sulfate (Gypsum) Solubility Calculations The data measured by Skillman, McDonals and Stiff has widely used to estimate the solubility of

gypsum in oil field brines. They measured ion product constant in simulated oilfield brines over

the following range:

Temperature: 50,95,122 and 176 0F (10,35,50 and 80 0C)

Ionic strength: 0-6.0 moles/L

Pressure: 1 atmosphere (101.3 KPa)

The following procedure is recommended to assess the possibility of gypsum precipitation from a

given brine:

1- Calculate the molar ionic strength using equation 2.38.

2- Obtain appropriate value of KC for the temperature of interest from figure.2.18.

3- Determine the excess common ion concentration, X, in moles/liter. This is simply the

difference between the calcium concentration and the sulfate concentration.

4- Calculate the solubility of gypsum in meq/liter by solving equation 2.44.

5- Calculate the actual concentration of gypsum in the water, which is equal to the smaller

of the Ca+2 or SO4-2 concentration expressed in meq/liter.

6- Compare the calculated solubility with the actual concentration to determine if

precipitation of gypsum is likely.

An Example calculation is given in this chapter.

2.3.7 Barium Sulfate Solubility Calculations It is possible to estimate the solubility of barium sulfate for waters which contain predominately

sodium and chloride ions and very little magnesium or calcium ions using the solubility data

measured by Templeton. Molar conditional solubility product constant calculated from his data

are presented in figure 2.19 over the following range:

Temperature: 77,95,122,149,176 and 203 0F (25,35,50,65,80 and 95 0C)







The solubility can be calculated using equation 2.44.

, and the probability of BaSO4 precipitation evaluated using the same procedure as previously

outlined for gypsum.

Because BaSO4 has such limited solubility, the appearance of Ba+2 and SO4-2 ions in any water

indicates a danger of scale formation.

2.3.8 Strontium Sulfate Solubility Calculations The solubility of strontium sulfate in sodium chloride solutions can be calculated in the same

manner as the solubilities of the other sulfate scales.

The following equation is based on Jacques and Bourland data and can be used to estimate

values of Kc.

log Kc = X/R Eq. 3.45

where

Kc = Conditional solubility product constant, molar units

X = 1/T

Eq. 2.46

The units are:

T = 0K = 0C +273

P = Total pressure, psig

u = Ionic strength, moles/L

The coefficients of the equation are:

The equation applies over the following range:

Temperature: 100- 300 0F (38- 149 0C)


Pressure: 100-3000 psig

Kc values measure by Fletcher, French and Collins for strontium sulfate is given in figure 2.20.

Their data was measured over the following range:

Temperature: 50,75,122 and 156 0F (10,35,50, and 69 0C)

Ionic strength: 0.1-5.25 moles/L






2.4 Example Scale Calculations Ionic Strength, CaCO3 Scaling Index and CaSO4 Solubility

Temperature = 60 0C ; pH =7.04

Calcium Carbonate Scaling Index Calculation

K = 2.24 from figure 2.16.

pCa =1.67 from figure 2.17.

pAlk = 2.05 from figure 2.17.

SI = pH – (K+pCa+pAlk) = 7.04 – (2.24 + 1.67 + 2.05) = +1.08

SI > 0, so CaCO3 scale is likely.

Calcium Sulfate Solubility Calculation

Kc = 9.2 X 10-4 from figure 2.18.

4 Kc = 36.8 x 10-4

Actual CaSO4 Concentration = 36.3 meq/L

S > Actual, so CaSO4 scale is unlikely.





Figure 2.16. Values for Stiff and Davis K for CaCO3 Scale calculation.





Figure 2.17. Conversion of mg/L Calcium and Alkalinity into pCa and pAlk.





Figure 2.18. Calcium sulfate (Gypsum) Conditional Solubility Product Constants.





Figure 2.18. (Continued) Calcium sulfate (Gypsum) Conditional Solubility Product Constants.





Figure 2.19. Barium sulfate Conditional Solubility Product Constants.





Figure 2.20. Strontium sulfate Conditional Solubility Product Constants.





Figure 2.20. (Continued) Strontium sulfate Conditional Solubility Product Constants.





Chapter 3

Scale problems & Inhibition

3.1 Introduction

Wells producing water are likely to develop deposits of inorganic scales. Scales can and do coat

perforations, casing, production tubular, valves, pumps, and downhole completion equipment,

such as safety equipment and gas lift mandrels. If allowed to proceed, this scaling will limit

production, eventually requiring abandonment of the well.

Technology is available for removing scale from tubing, flow-lines, valves, and surface

equipment, restoring at least some of the lost production level. Technology also exists for

preventing the occurrence or reoccurrence of the scale with conventional inhibitor technology

3.2 Scale mechanisms

As brine, oil, and/or gas proceed from the formation to the surface, pressure and temperature

change and certain dissolved salts can precipitate. This is called “self-scaling.” If a brine is

injected into the formation to maintain pressure and sweep the oil to the producing wells, there

will eventually be a commingling with the formation water. Additional salts may precipitate in the

formation or in the wellbore (scale from “incompatible waters”). Many of these scaling processes

can and do occur simultaneously. Scales tend to be mixtures. For example, strontium sulfate is

frequently found precipitated together with barium sulfate.

3.3 Causes of Scales Many types of scales form in petroleum producing operations. This module discusses the most

common and most disabling types: the water-formed deposits. Other types of scales, such as

those formed by corrosion or bacteria, are covered in other courses. Water-formed deposits are

seldom pure compounds. They are usually mixtures of compounds in a mixed-crystal lattice or as

layers of different compounds. As a starting point, we will discuss the scale compounds barium

sulfate (BaSO4), calcium sulfate (CaSO4), calcium carbonate (CaCO3), and strontium sulfate

(SrSO4).

Water-formed scales can form for a number of reasons. Defining the cause of the scale is

essential to selection of the best remedy. Sometimes, a simple change in the production system

is enough to prevent scale formation. In other cases, chemical inhibitors must be added. The

common causes of water-formed deposits are

• Change in temperature

• Mixing of incompatible waters

• Loss of carbon dioxide (increase in pH)

• Drop in pressure

3.3.1 Change in temperature Temperature change is a contributing factor in the deposition of barium sulfate and, in some

cases, calcium carbonate. Deep gas and condensate wells in the Gulf of Mexico and parts of





Texas are plagued by the deposition of barium sulfate in lower reaches of the wells and in the

subsurface safety valve (SSSV). These wells produce very little water, maybe only 10 barrels

per day. The problem is attributed to the drop in temperature when gas expands in the tubing. If

formation water is saturated with respect to barium sulfate, then some deposition can be

expected as the temperature drops (Figure 3.1).

Figure 3.1 Effect of temperature of the solubility of BaSO4 in water.

The effect of temperature on the solubility of calcium carbonate is inverse to that shown for

barium sulfate (Figure 3.2).

The solubility of calcium carbonate decreases with increasing temperature. The data in Figure 3.2

were obtained in a closed container in the presence of carbon dioxide. The decrease in solubility

with temperature is greater if carbon dioxide is released from the system. The reasons for this will

be discussed later. From Figure 3.2, one would presume that injecting a water saturated with

CaCO3 into an injection well would cause scale formation

because of the increased temperature as the water descends into the well. This seldom results in

problems for the injection well because the well bore and formation near the injection well are

soon cooled by the injected water. Furthermore, calcium carbonate precipitates slowly because

its solutions easily exceed saturation. Note that the solubilities of CaCO3 are about

1,000 times those of BaSO4. In producing wells, the temperature decreases as water rises in the

tubing, causing an increase in the solubility of CaCO3. At the same time, carbon dioxide is

released from the water due to a decrease in pressure as the water rises in the production tubing.

This causes a decrease in the solubility of CaCO3. Generally, the loss of CO2 predominates and

CaCO3 precipitates in the upper reaches of the tubing. Because CaCO3 precipitates slowly, it also

leaves scale on surface equipment. Fortunately, computer programs are available that give the

equilibrium conditions for CaCO3 solutions at any temperature and pressure. With this

information, the scaling tendency can be predicted.





Calcium carbonate scaling in produced waters also results when a water incompatible with the

formation water is injected. Complexities arise in predicting incompatibility. Although a water may

be compatible as injected, it can become incompatible through interactions with minerals in the

oil-producing formation. This is especially true of calcareous formations.

Figure 3.2 Effect of temperature of the solubility of CaCo3 in water.

Calcium sulfate solubility increases with temperature below 860F (300C), but the solubility

decreases with temperature above 860F (Figure 3.3).

Calcium sulfate exists in three forms: gypsum (CaSO4 • 2H2O), which contains two water

molecules; hemihydrate (or, subhydrate, CaSO4 • 1/2H2O); and anhydrite (CaSO4). Only gypsum

and anhydrite occur in scales and in petroleum reservoirs. Anhydrite absorbs water very strongly

at temperatures below 1220F (500C), making it a good drying agent. The

crystalline form of the anhydrite and resulting gypsum are exactly the same except for the

presence of water in the latter. Plaster of Paris is made by slurrying anhydrite in water.

Anhydrite absorbs the excess water and converts to a dry, cohesive mass of gypsum. Calcium

sulfate scales sometimes occur in producing wells even though only formation water is produced.

The reasons for this are not clear. It is presumed, however, to be caused by the dissolution of

anhydrite in the formation, which precipitates as gypsum when the temperature of the water is

lowered. Calcium sulfate precipitates as anhydrite if the temperature is greater than 1220F (50 0C). Certain salts cause anhydrite to precipitate at temperatures as low as 1040F (400C). An

injected water need contain neither calcium ions or sulfate ions to be incompatible with the





formation water if the petroleum reservoir contains gypsum or anhydrite deposits. Cool injected

water can dissolve large quantities of anhydrite then precipitate it as the water is heated by the

reservoir. This problem occurs extensively in West

Texas. There plugging of the formation near the well bore is common with resulting losses in

production.

Figure 3.3 Effect of temperature of the solubility of CaSO4 in water.

Strontium sulfate (SrSO4) is the rarest of the water-formed scales to be discussed here, but it is a

problem in some areas. The solubility of SrSO4 in water (114 mg/l at 7 0F) is less than CaSO4

(2080 mg/l) but greater than BaSO4 (2.3 mg/l) under comparable conditions. The change in

solubility of SrSO4 with temperature is the same as gypsum in that solubility increases with

temperature up to 1040F (400C) but decreases at higher temperatures (Figure 3.4).

SrSO4 scales result from the injection of water that is incompatible with the formation water.

An example would be injecting a water, such as sea water, that is high in sulfate into a formation

where the water contains high levels of strontium.

Higher pressures increase the solubility of all the scales discussed above. The main reason for

the increase in solubility is that the volume of a system composed of scale and water decreases

when the scale dissolves. This can be expressed simply as shown in Figure 3.5, which illustrates

the effect of pressure on the solubility of anhydrite (CaSO4). A secondary factor in increasing the

solubility is that water is compressed to a volume decrease of about 2 percent on the application

of 500 bars hydrostatic pressure. This decrease in volume increases the salt concentration per

unit volume, and higher salt concentrations increase the solubility of scales. The effect of

pressure on solubility decreases with increasing temperature and total dissolved solids (TDS) of

the water. In the case of CaCO3, higher partial pressures of carbon dioxide (CO2) increase the

solubility. This effect of CO2 applies only to CaCO3 and not to BaSO4, CaSO4, or SrSO4.





Figure 3.4 Effect of temperature of the solubility of SrSO4 in water.

Figure 3.5 Effect of pressure on the solubility of CaSO4 (Anhydrite) in water.





Carbon dioxide increases the solubility of CaCO3 as shown in Figure 3.6. The reverse condition

also applies. A loss of CO2 from solution results in the precipitation of CaCO3.

The loss of CO2 can occur as the result of increasing the temperature or dropping the pressure.

In some cases, CaCO3 scales in producing systems have been controlled by preventing release

of CO2 from the system. Because of the effect of the loss of CO2, CaCO3 scaling can result from

the production of produced water only. It can also result from injection of an incompatible water.

Sea water is almost saturated with CaCO3.

The solubility of BaSO4 with increasing salt concentration (in molality) is shown in Figure 3.7.

The increase in solubility is greater at low salt concentrations. The solubility of CaSO4 goes

through a maximum at about 3.0 m (molality) salt concentration, presumably because the added

salt ties up sufficient water molecules to encourage the precipitation of anhydrite.

The solubility of CaCO3 also goes through a maximum with salt concentration at about 2 m

(molality) because the solubility of CO2 decreases with increasing salt concentration (Figure 3.8).

Figure. 3.6. Effect of pressure on the solubility of Calcium carbonate.





Figure 3.7. Effect of salt concentration on the solubility of barium sulfate

Figure. 3.8. Effect of Salt Concentration on the Solubility of Calcium Carbonate at 25°C





3.3.2 Mixing Incompatible water Incompatible waters are those that produce a precipitate when mixed. Incompatible waters can

be waters produced from different zones and then commingled in the producing systems, or they

can be a formation water and an injected water. Most often, the injected water contains an ion

that reacts with an ion in the formation water to cause precipitation and scaling when the two

waters mix in a producing well. Examples of incompatible waters are given in table 3.1.

Table 3.1 .Incompatible ions.

The illustration in Figure 23 gives only the ions that take part in the scaling reactions. Actual

waters contain other ions that affect the precipitation reactions. Two incompatible waters and their

produced mixture are given in

Table 3.2. Production of a mixture of these waters resulted in severe scaling of the producing well

by CaSO4. Notice that the formation water is high in calcium and the injected water is high in

sulfate, resulting in the precipitation of CaSO4. The injected water in table 3.2 is much lower in the

concentrations of most ions than is the formation water.

Table 3.2 .Composition of incompatible water and the mixture of the two.





Therefore, we could conclude that injected water was breaking through to the producing well if

the salt content of the produced water decreased. Also, we could accomplish the same thing by

monitoring the concentration of only one ion; for example, Mg++ in the produced water.

When the Mg++ concentration decreased, it indicates that injected water is breaking through. The

composition of the produced water in terms of the fractions of injected water and formation water

can be calculated from the concentration of Mg++ in each water. For example, the fraction of

injected water in the produced water as shown in Figure 3.2 is given by:

3.3.3 Loss of carbon dioxide (increase in pH) The loss of carbon dioxide is usually due to the decrease of pressure which decrease the

solubility of CO2 and results is an increase of pH.

The pH of a produced water generally has little effect on scaling by BaSO4, CaSO4, or SrSO4, but

a change in pH by loss of CO2 in a producing well has a large effect on scaling by CaCO3.

The efficiency of chemical scale inhibitors for all the scales is affected by pH. The high pressures

in many oil producing reservoirs cause large concentrations of CO2 to go into solution, resulting in

a drop in pH (Figure 3.9).

Figure. 3.9. Effect of CO2 Partial pressure on pH.

Calcium carbonate reacts with hydrogen ions in water, causing the pH of the water to rise.

Therefore, injecting a water into a limestone reservoir generally results in an increase in pH of the

injected water.





3.4 Field and Lab Examination of Scales The first step in controlling a scale problem is to identify the scale with certainty. This is difficult to

do visually because the various scales can occur in different forms and colors (Figures 3.10 to

3.12).

Figure. 3-10. Calcium carbonate and calcium sulfate scale.

Figure. 3-11 Strontium sulfate and Barium sulfate scale





Table.3.3 Properties of common scale.

Mixed scales are common. Calcium carbonate mingled with barium sulfate is shown in

Figure 3.12. If two scales have a common ion, such as CaSO4 and BaSO4, the least soluble

compound forms first. If a water is capable of forming both CaSO4 and BaSO4, then BaSO4 is

found in the bottom of a producing well. If there is enough SO4-2

left to form CaSO4, it will be

found in shallower depths of the well.

Figure. 3-12. mixed calcium carbonate/barium sulfate scale

3.5 Qualitative and Quantitative Tests Identification of a scale in the field by simple physical or chemical tests requires considerable

experience with various types of scales unless the scale is calcium carbonate. Calcium carbonate

dissolves in a weak acid such as vinegar with the liberation of carbon dioxide. If a scale "fizzes" in

vinegar (acetic acid) and dissolves, then it is calcium carbonate. The other scales require closer

scrutiny as shown in Figure 3.13. The tests given in Figure 38 are not fool-proof but provide an

educated guess of the identity of a scale. Chemical verification is still required if there is doubt or

if the scale is complex.





Figure 3.13. Simple Field Tests for Identifying Scales

Discussion board






احفظ) الرابط على وهو الماوس يمين من تختار أن الممكن ومن, ذلك بعد وتحفظه الملف تفتح ممكن














Chapter 4

Scale Inhibition

4.1 Theory of Scale Inhibition

4.1.1 Introduction There are two major chemical methods for preventing, or inhibiting, scale formation. These are

removal of the offending cation (Ba+2, Ca+2, Sr+2) or anion (SO4-2, CO3

-2) from solution or addition

of a chemical that prevents crystal growth. The removal of offending ions is accomplished by ion

exchange, precipitation, or complexation. Each of these techniques requires the addition of a

stoichiometric quantity of chemical. Because large amounts of added chemicals are required,

removal of the offending ion is practiced only where the concentration of offending ions is very

small, such as in a boiler feed water.

Preventing crystal growth requires only the addition of a very low concentration of chemical,

maybe only about 2 ppm of an effective chemical. For treating large amounts of water that

contain high concentrations of scale-forming ions, preventing crystal growth is the preferred

method.

Removal of ions

Anion exchange

Cation exchange

Precipitation

Complexation

Acidification (CaCO3 only)

Prevent crystal growth

Adsorbed chemicals

In this section, Theory of Scale Inhibition, we will discuss complexation and adsorption.

Complexation is rarely used for preventing scale in produced waters. We include it here because

it is useful for small amounts of water or low concentrations of cations.

Adsorption scale inhibitors are good complexing agents for cations. Understanding this property

is important in the formulation and use of adsorption scale inhibitors. Adsorption scale inhibitors

are the most important to us in preventing scale in oil field produced waters and most of our time

will be spent on this topic.

4.2 Complexation Recall that complexing agents, or chelating agents, are useful for dissolving scale. They react

with cations to form a new ion that is more soluble.

M++ + X ↔ X M++ Eq. 4.1

Where: M++ is a cation such as Ca++, Ba++, or Sr++

X is a chelating agent such as EDTA or DTPA

Complexing agents can lower the concentration of a cation so its concentration no longer

exceeds the solubility product constant. They can therefore, be used to prevent scale and to





dissolve scale. Complexing agents are commonly added to cooling water and boiler waters for

these purposes.

Complexing agents must be added in a stoichiometric quantity relative to the amount of cation to

be complexed. Most complexing agents such as EDTA and DTPA cost about $1.00/lb so they are

not economical in many cases. Another disadvantage to complexing agents is that they complex

more than one cation. In other words, they are generally not specific for only the cation selected

for complexing. The ability of a complexing agent to complex a given cation is expressed by the

stability constant, K, as previously explained. Using the reaction given in Eq. 4.1, the stability

constant can be expressed as shown in Eq. 4.2. The stability constant is the ratio of complexed

cation to uncomplexed cation and is, therefore, an expression of the strength of the complex.

Eq. 4.1

The stability constants of EDTA for several cations are given in Figure 4. If EDTA is added to a

mixture of the cations given in table 4.1, the Ca++ ions will be complexed first because they form

the strongest complex (log K = 10.59) with EDTA. However, if Ba++ ions are to be complexed,

then sufficient EDTA must be added to complex all of the Ca++ and Mg++ ions first. The amount

of EDTA required might be many times that required to complex only the Ba++. The amounts of

EDTA required to complex Mg++, Ca++, and Ba++, based on mole-for mole addition, are given in

table 4.2.

Table.4.1. Some Stability Constants for EDTA Complexes.

Table.4.2. Amount of Na3EDTA Required to Complex Various Cations.

Suppose a water contains 1 ppm each of Mg++, Ca++, and Ba++ ions. It would require 9.6 ppm

of Na3EDTA to complex the Ca++ but 25.2 ppm to complex the Mg++ because sufficient EDTA to

complex the Ca++ must also be added. Complexing the Ba++ would require 28.0 ppm of

Na3EDTA. Almost all complexing agents behave in this manner. It would be advisable to measure

the amount required experimentally if a mixture of cations is present.





4.3 Adsorption Adsorption is the interaction of a substance with a solid surface. Absorption is the taking up of

one substance by another throughout its bulk. This section is titled adsorption because it is

believed that surface adsorption is the way that some scale inhibitors are able to prevent scale at

concentrations of only 1 ppm or so.

About 80 years ago, it was discovered that some inorganic phosphorous compounds, called

metaphosphates, could prevent CaCO3 scaling in irrigation systems when only 2 ppm were

added to the water. This small amount of metaphosphate could maintain a supersaturation of

over 200 ppm of CaCO3 in solution. A low concentration of metaphosphate could not only prevent

CaCO3 scaling, but it also caused previously-formed CaCO3 deposits to disappear.

Clearly, this was not a stoichiometric phenomenon and so it was called a threshold treatment

because it maintained the dissolved CaCO3 on the threshold of precipitation. The term threshold

treatment today means the use of substoichiometric quantities of scale inhibitor, specifically,

concentrations in the range of 20 ppm, or less. With modern threshold chemicals, we often

observe effective scale inhibition with only 0.5 ppm of inhibitor.

Metaphosphates belong to a family of inorganic phosphates called condensed phosphates,

dehydrated phosphates, or polyphosphates. They are inexpensive and still used in municipal

waters and cooling waters. The preparation of a polyphosphate is written in the following

reactions. The monomolecular phosphoric acid, called orthophosphoric acid, is heated with an

alkali to expel water and form sodium hexametaphosphate, an excellent scale inhibitor. The

orthophosphates have no scale-inhibiting properties at all. In practice, it was found that solutions

of sodium hexametaphosphate rapidly lost their ability to inhibit scale. The reason is that the

hexametaphosphate reacts with water, or hydrolyzes, and reverts back to the inactive

orthophosphate. This loss of scale-inhibiting activity is called reversion.

Formation 6H3PO4 + 3Na2O → Na6(PO3)6 + 9H2O Reversion Na6(PO3)6 + 6H2O → 6NaH2PO4

At one time, polyphosphates were widely used in oil fields to inhibit scale. They were found to be

effective against CaCO3, CaSO4, and BaSO4. Because of the reversion problem, solutions of

polyphosphates could not be depended on for long life when squeezed into formations. The rate

of reversion increases with temperature and acidity so the squeezed solutions rapidly lost

effectiveness under downhole conditions (Table 4.3).

Table 4.3. Reversion of Inorganic Polyphosphates at 25 °C





To overcome the reversion problem, controlled-solubility, or glassy, polyphosphates were

developed. These contained cations other than sodium to limit their solubility. The glassy

phosphates in pellet form were fractured into formations with propping agents. Although scale

inhibition for up to 4 years was reported, penetration of the pellets and life of the treatment were

very uncertain.

Another disadvantage of the inorganic phosphates was overtreatment. If too much phosphate

was used and reversion occurred, calcium orthophosphate, Ca3(PO4)2, precipitated from solution.

The solubility of calcium phosphate is only 20 mg/l, thus, it is a more offensive scale than is

CaCO3. The glassy, controlled-solubility, phosphates offered a way to avoid overtreatment.

Baskets of the compound were hung from the tubing of producing wells, placed in source water

tanks, and put in flowing streams. The glassy phosphates slowly dissolved to provide inhibition for

a period of time after which the basket was refilled.

Even though their cost was low, disadvantages of the polyphosphates caused them to give way

to the more expensive, but more reliable, organic compounds used today. The use of inorganic

polyphosphates is almost unknown in oil fields today but much of our application technology and

knowledge of the mechanisms of scale inhibition were developed for the polyphosphates.

Several mechanisms have been proposed whereby threshold scale inhibition occurs. Generally,

nucleation is not prevented by scale inhibitors, but the scale nuclei remain very small – in the

order of only a micron (10-6 m) in diameter. Restriction of crystal growth by adsorption on one or

more crystal faces appears to be the dominant mechanism for scale inhibition by threshold

inhibitors. If a CaCO3 crystal is placed in a metastable saturated solution of CaCO3, the crystal

will grow. However, if the crystal is first immersed in a solution of threshold inhibitor and then

placed in the saturated solution, it will not grow.

Crystals nucleated in the presence of threshold inhibitors have a different form than those

nucleated under normal conditions. In extreme cases, where some scale formation occurs in

inhibited waters, the scale is softer and bulkier and more easily sloughed off than scale formed in

uninhibited waters.

Surfaces treated with scale inhibitor, and then exposed to scale-forming water, have a delayed

tendency to scale. Supposedly, the inhibitor adsorbs on the solid surface and prevents adherence

of scale nuclei.

Scale inhibitors are also good deflocculants. In other words, they cause agglomerations of

crystals to disperse. Scale inhibitor molecules contain ionic groups. These ionic groups impart the

same electrical charge on solids to which they adsorb. The like charges repel one another,

causing the solids to disperse. Very large polymers of the same inhibitors can adsorb on more

than one particle, binding them together. The large, or high molecular weight polymers, are used

as flocculants to aid settling and filtration of solids.

In summation, threshold scale inhibitors are able to adsorb on solid surfaces. As a result of this

adsorption, they are able to prevent crystal growth, change crystal forms, decrease adherence,

and disperse crystals. The result of these effects is that these compounds prevent problems from

scale formation even when present in very low concentrations.

The organic phosphorous, or organophosphorus, scale inhibitors have almost entirely displaced

the inorganic phosphorous compounds from use in the oil field. Other organic compounds

including phosphoric acid esters, polyacrylates, and polyacrylamides are also widely used. The

main reasons for the shift to the organic compounds are that they are more stable and they can

be “tailored” to be very effective under extreme conditions.





One of the earliest organophosphorus scale inhibitors was aminotrimethylenephosphonate

(ATMP), which was marketed as Dequest 2000 by Monsanto Chemical Co. (Figure 4.1). The dry

pentasodium salt is called Dequest 2005DN and a 40% solution of the pentasodium salt is

Dequest 2006.

The acid form of ATMP is a strong acid. A 1% solution has a pH of less than 2. It must therefore

be handled in acid-resistant equipment. A 1% solution of the pentasodium salt solution has a pH

of 10 to 11 and is thus less hazardous to handle. The two forms are equal in all respects in their

scale-inhibiting properties when diluted in process water. Properties of the two products are given

in table 4.4.

Figure 4.1. Structure of ATMP and Its Pentasodium Salt

Table 4.4. Properties of Commercial ATMP





The compatibility of phosphoric acids and their sodium salts with a wide variety of solvents makes

them readily formulated for a variety of purposes. The phosphonic acids are available in a variety

of molecular weights and copolymers with carboxylic acids. An example of a compound having a

higher molecular weight than ATMP is ethylenediaminetetramethylenephosphonic acid (EDTMP)

shown in Figure 4.2. Any number of the group shown in brackets can be added.

Figure 4.2. Ethylenediaminetetramethylenephosphonic Acid (EDTMP)

Phosphonates resist hydrolysis much better than do the inorganic condensed phosphorous

compounds. ATMP is stable for 90 years in water at pH 7 and 185 °F. A comparison of the

stability of ATMP with sodium pyrophosphate is given in table 4.5.

Table 4.5. Hydrolysis of Dequest and Sodium Pyrophosphate at pH 7.0

The hydrolysis rate of organic phosphonates is constant over all pHs. The rate for condensed

phosphates increases as the pH decreases in water solutions. Organic phosphates rival the

phosphonates in scale-inhibiting properties but are less stable to hydrolysis. They are, however,

more stable than the inorganic phosphorous polymers. The phosphates cost less than

phosphonates and should be considered for applications where stability is not a problem.

Structures of some phosphate groups are given in Figure 4.3. Note that in a phosphate, the

phosphorous is connected to carbon through an oxygen atom while, in a phosphonate, the

phosphorous is connected directly to carbon.





Figure 4.3. Structures of Organic Phosphate Esters

Other modern scale-inhibiting chemicals are the polyacrylic acids and polyacrylamides (Figure

4.5). These compounds are effective scale inhibitors when having molecular weights less than

about 20,000. Compounds having molecular weights higher than 20,000 are good flocculants.

Figure 4.5. Structures of Some Scale-inhibiting Polyacrylates

Molecular weight is an important variable in the scale-inhibiting properties of an organic polymer.

For AMP, the monomer is the best inhibitor for BaSO4 and CaCO3 but the dimer is best for

gypsum (Table 4.6). Commercial formulations seldom consist of active ingredients having a

single molecular weight. Usually, a spectrum of molecular weights will be present.

Table 4.6. Effect of Molecular Weight on Scale Inhibition by AMP





The effective concentrations of inhibitor found in Figure 16 are for the given conditions only.

Generally, the amount of inhibitor required increases with the degree of supersaturation (Table

4.7), decreases at lower temperatures, and increases with pH (Table 4.8).

Table 4.7. Effect of Degree of Supersaturation of CaSO4 on the Concentration of AMP Required for Scale

Inhibition

Table 4.8. Effect of Temperature and pH on the Inhibition of CaSO4 and CaCO3 Precipitation by AMP

4.3.1 Threshold Inhibition As stated earlier, threshold inhibition is effective scale inhibition at concentrations of inhibitor less

than the stoichiometric concentration required to complex the offending ions. In actual practice,

the effective concentration of a good scale inhibitor for a particular system is about 5 ppm, or

less. Systems with high pH, high suspended solids, high supersaturation, or high temperatures

might require a higher concentration of scale inhibitor. Generally. the addition

of higher-than-required concentrations of the organic phosphonates, acrylates, and acrylamides

are only wasteful and do no other harm. Table 4.9 is an example of overtreatment with partially

hydrolyzed polyacrylamide (PHPA). One hundred per cent inhibition of CaCO3 precipitation was

obtained from 5 to 20 ppm of PHPA. The excess of inhibitor over 5 ppm is wasted.

Excess concentrations of inorganic phosphorous compounds and organic phosphate esters can

be harmful if the compounds hydrolyze to form orthophosphate. High concentrations of

orthophosphate can precipitate Ca3(PO4)2, thus contributing to scale problems.

Proprietary scale inhibitor formulations usually contain about 20% by weight of the active scale-

inhibiting compound. Frequently, a recommended scale inhibitor dosage is given in terms of the





inhibitor as formulated. For example, a 5 ppm dose of inhibitor formulation would be only 1 ppm

of the inhibitor compound. If any confusion exists about the meaning of a recommended dosage,

you should clarify the units before proceeding with the treatment.

If you think the dose refers to pure compound when it actually refers to 20% formulated

compound, then you will overtreat by a factor of five. In the reverse situation, you would

undertreat by a factor of five.

Table 4.9. Inhibition of CaCO3 Precipitation by Partially Hydrolyzed Polyacrylamide

4.4 Laboratory Evaluation of Inhibitors Scale inhibitor testing is imperative if a large-scale treatment program is anticipated. Cost

effective treatment, of course, is the major goal, but other considerations must be made also.

There are probably some 2000 scale inhibitor formulations on the market, many of which will not

be suitable for the particular system to be treated. Even though we have seen that organic

polymers of phosphonates, phosphates, and acrylates are good scale inhibitors, the purity,

formulation, and average molecular weight of the inhibitor will vary among suppliers.

Some of the factors that require testing for a particular system are listed in Figure 4.10.

Composition of the scale is important because some scale inhibitors are effective against some

scale compounds but are not effective against other scale compounds. Many chemical scale

inhibitors are effective when the degree of supersaturation is low but most are ineffective when

the degree of supersaturation is high. The temperature of the operating system must be

duplicated, if possible. All scale inhibitors have an upper temperature limit above which they are

not effective. The selected scale inhibitor must be compatible in not decreasing the function nor

causing precipitation of other chemicals in the system, such as corrosion inhibitors, biocides,

oxygen scavengers, or emulsion breakers. The scale inhibitor must form clear solutions with the

water in which it is to be used. Downstream effects, such as foaming and oil-water separation,

must not be aggravated by the scale inhibitor. Use of a particular scale inhibitor will require that

large quantities of the concentrated formulations be shipped and stored in bulk or drums. The

effects of local weather on the chemical should be determined to avoid freezing or overheating. A

frozen drum of chemical is useless until thawed. Formulations that contain volatile solvents have

been known to rupture drums when stored in intense sunlight. Finally, the only way to calculate

the cost for using a particular chemical is to determine the concentration at which it is effective.

Most of the tests required are self-explanatory. We will devote most of our time on tests to find

effective scale inhibitors and their effective concentration. These tests are continuously useful

because, even if an effective compound is found, new, more effective formulations might become





available. Occasional review of the scale inhibitor program and revision as necessary can result

in great savings in the chemical treatment program.

• Composition of the scale

• Degree of supersaturation

• Temperature

• Compatibility with other chemicals

• Compatibility with the water

• Downstream effects

• Weather

• Cost

Table 4.10. Factors to Consider in Selecting a Scale Inhibitor

There is no definitive simple test for finding an effective scale inhibitor. The test should simulate

the system with which we are concerned as closely as possible. The best test is in the system

itself, but inhibitors should be screened first in a rapid, inexpensive, simple test to select the best

chemicals for the purpose. Broadly, the tests we will consider are divided into static and dynamic

tests. The latter more closely simulate operating conditions but are more expensive and time-

consuming to run.

4.4.1 Precipitation Tests The simplest scale test is to mix two incompatible waters containing scale inhibitor and measure

the amount of precipitate that forms. The effectiveness of the inhibitor can be expressed as the

per cent of precipitate that was prevented (in the next equation) or the per cent of scale forming

compound retained in solution.

In essence, a precipitation test consists of mixing two volumes of water, one of which contains the

cation and the other the anion of the scale compound. The mixture is capped and stored at

temperature for a predetermined period of time, usually 24 hours, after which the precipitate is

filtered out and weighed. A test is run without scale inhibitor to determine “S” for the equation in

Figure 21. For inhibitor tests, the inhibitor can be put in either water or both, whichever more

closely simulates the actual system. In testing for a squeeze job, the inhibitor would be put in both

waters.

E = [S - F] x 100 / S where: E = Inhibitor effectiveness, % S = Weight of precipitate formed in absence of inhibitor F = Weight of precipitate formed in presence of inhibitor

The volumes of water that are mixed should simulate those encountered in actual practice. Each

water should have a total dissolved solids content and pH approximating actual conditions.

Several difficulties might be encountered in designing a precipitation test for a particular purpose.

Ideally, we would use the actual waters that we want to inhibit, but this is not always possible. If

two waters are mixing downhole, we can obtain only a mixture of the waters. Also, produced

waters are seldom stable and reliable in behavior because of the presence of hydrocarbons,





reduced substances such as iron, unstable compounds, and dissolved gases. It is also likely that

two actual incompatible waters might not form a precipitate when mixed. This is particularly true if

the scale compound is one that readily

supersaturates, such as gypsum or strontium sulfate.

The most reliable and consistent data are obtained by using synthetic waters that simulate the

known compositions of the actual waters. Precipitation can be insured by adjusting the

supersaturation or by seeding the mixture of waters.

The offending ions can be incorporated in the test waters using the compounds listed in Table

4.11. When injected water is produced, a range of compositions and supersaturations are going

to be encountered in production because of the variable ratio of produced water to injected water

that is produced. In this case, it is best to run several tests covering the range of water

compositions that are expected to produce scaling conditions. Remember that precipitation can

be forced in the tests by adding more of the scaling compounds listed in Table 4.11.

Table 4.11. Compounds Added to Test Waters to Simulate Precipitates

For inhibitor additions to the test solutions, it is convenient to make a 1% (wt) stock solution of the

inhibitor. One ml of 1% solution contains 10 mg of inhibitor. So if you wanted to make 1 liter of

solution containing 5 ppm (5 mg/l) of inhibitor, you would add 0.5 ml of the stock solution.

Amounts of 1% stock solution to add for 50 ml of test solution are given below. Remember to add

enough inhibitor for the final volume of test solution. If two 25- ml volumes of water are to be

mixed for a test, then the final volume for which inhibitor concentration is calculated is 50 ml.

Table 4.12. Volume of 1% (Wt) Inhibitor Solution Required for Test Solutions.





Discussion board






احفظ) الرابط على وهو الماوس يمين من تختار أن الممكن ومن, ذلك دبع وتحفظه الملف تفتح ممكن












Documents

Water Formed Scale - oilprocessing.netoilprocessing.net/data/documents/Water-Formed-Scale.pdf · 1.1.7 The mole A mole is defined as the amount of substance that contains as many