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Produced water properties
Understanding the physical properties of the formation water that will be produced along with the oil or gas is important to a proper assessment of reserves volumes, producibility, economics, and surface facilities. As a rule, it is best to have reliable laboratory measurements of the physical properties of oilfield waters. If laboratory measurements are not available, correlations may have to be used. For example, McCain has published some of the most widely used correlations for the physical properties of oilfield waters.
This page discusses the resistivity, surface (interfacial) tension, viscosity, pH, and redox potential (Eh) of produced water. Additional properties are discussed on the pages linked below.
The resistivity of formation water is a measure of the resistance offered by the water to electrical current. It can be measured directly or calculated. The direct measurement method is essentially the electrical resistance through a l-m2 cross-sectional area of 1 m3 of formation water. Formation water resistivity, Rwg, is expressed in units of Ω-m. When resistivity of formation water is used in electric-log interpretation, the value is adjusted to formation temperature.
Surface (interfacial) tension
Surface tension (interfacial tension or IFT) is a measure of the attractive force acting at a boundary between two phases. If the phase boundary separates a liquid and a gas or a liquid and a solid, the attractive force at the boundary usually is called surface tension; however, the attractive force at the interface between two liquids is called interfacial tension (IFT). The lower the IFT, the smaller the droplet of the internal phase. At very low values of IFT, oil and water become miscible and behave as a single phase. IFT is an important factor in enhanced recovery. Also, the IFT determines the ease of separation of oil from water, because it determines the size of the oil or water droplets, depending on which phase is internal.
Most chemicals added during the course of drilling or production have a major effect on the IFT of the produced water and the hydrocarbons. Indeed, certain corrosion inhibitors added to the three-phase production stream can lower the produced water IFT enough (<1 to 5 dyne/cm) to cause the droplet size of the entrained oil to be small enough that no injection well plugging is observed, even at high oil carryover (percent levels) in the reinjected produced water. In attempting to separate the oil from the three-phase production stream, the addition of emulsion breakers changes the IFT and promotes the agglomeration of small droplets into larger ones that separate quickly. Formulating, selecting, testing, and troubleshooting emulsion breakers is the focus of an enormous amount of the effort devoted to the impacts of producing water with hydrocarbons.
Surface tension is measured in the laboratory by a tensiometer, by the drop method, or by a variety of other methods. The laboratory measurements traditionally have been difficult and done only by specialized facilities. Computerized commercial pendant-drop and falling-drop tensiometers are now available for use by chemists in more general field R&D laboratories. IFT is a critical property of produced water, but is rarely measured because of the analytical difficulties. This new technology promises to improve the ability to troubleshoot problems by directly measuring IFT instead of trial-and-error testing.
The viscosity of formation water, μw, is a function of:
- Dissolved solids
In general, brine viscosity increases with:
- Increasing pressure
- Increasing salinity
- Decreasing temperature
Dissolved gas in the formation water at reservoir conditions generally results in a negligible effect on water viscosity. There is little information on the actual numerical effect of dissolved gas on water viscosity.
Gas-in-the-water phase behaves entirely differently than gas in hydrocarbons (personal communication with J.C. Melrose, Mobil R&D Corp., Dallas, 1985). In water, the presence of the gas actually causes the water molecules to interact with each other more strongly, thus increasing the rigidity and viscosity of the water. However, this effect is very small and has not been measured to date. In the physical chemistry literature, there is an enormous amount of indirect evidence to support this concept.
For the best estimation of the viscosity of water, refer to Kestin et al. Their correlating equations involve 32 parameters for calculating the numerical effect of pressure, temperature, and concentration of aqueous NaCl solutions on the dynamic and kinematic viscosity of water. The 28 tables generated from the correlating equations cover :
- Temperature range from 20 to 150°C
- Pressure range from 0. 1 to 35 mPa
- Concentration range from 0 to 6 molal.
Figs. 1 through 3 may be used to approximate water viscosity.
McCain reported that this correlation is within 5% of graphical correlations for temperatures between 100 and 400°F, and salinities to 26 wt%.
Water viscosity at reservoir pressure can be obtained from
McCain reported that Eq. 4 fits the data to within 4% for pressures below 10,000 psia and temperatures in the range from 86 to 167°F. The fit is within 7% for pressures between 10,000 and 15,000 psia.
Figs. 1 through 3 show the effects of pressure, temperature, and NaCl content on the viscosity of water. They may be used when the primary contaminant is sodium chloride. Alternatively, the viscosities are calculated and reported by computer chemistry models at the particular temperature, pressure, and gas compositions present in the facilities.
Some engineers assume that reservoir brine viscosity is equal to that of distilled water at atmospheric pressure and reservoir temperature. In this case, it is assumed that the viscosity of brine is essentially independent of pressure (a valid premise for the pressure ranges usually encountered). In some high-temperature/high-pressure reservoirs recently developed, this assumption breaks down. In those cases, experimental measurements under the relevant temperature/pressure conditions are recommended over attempting to extrapolate the distilled-water viscosities or even the computer models.
Water (H2O) reversibly dissociates into hydrogen ions and hydroxide ions, which is described by the equilibrium constant for this chemical reaction, Keq (H2O) or simply Kw. The acidity or hydrogen ion activity of aqueous solutions controls many of its properties and is commonly expressed as the pH.
is the water dissociation reaction.
where aH+ is known as the activity of hydrogen ion in solution. The hydrogen ion activity is related to the concentration of hydrogen ions [H+] by means of the activity coefficient, γH+, giving
Solutions are known as neutral when the pH = 7, because at that point hydrogen ions and hydroxide ions are present in equal amounts, aH+ = aOH− = 10−7 M. When hydrogen ions predominate, the pH falls below 7, and the solution is described as being acidic. In the opposite case in which hydroxide ions outnumber the hydrogen ions, the pH climbs above 7, and the solution is known as basic or alkaline. The pH is commonly accurately measured with an electrode and meter, while field determinations also may use pH paper strips or colorimetric methods.
When the water is very pure and contains little dissolved salts, the value of γH+ approaches 1.0. The activity and concentration of hydrogen ion are essentially the same, so that the pH definition simplifies to
However, in what must seem to be nature’s perversity, produced water from oil reservoirs usually contains large amounts of dissolved salts. The value of γH+ is < 1.0 as a result, so the more simple form of the pH definition is not correct. Careful, direct pH measurement is the best approach for accurate pH determination, although some of the most sophisticated computer models give reasonable predictions at moderate conditions of brine concentration, temperature, and pressure.
The pH of oilfield waters usually is controlled by the CO2/bicarbonate system. Because the solubility of CO2 is directly proportional to temperature and pressure, the pH measurement should be made in the field if a close-to-natural-conditions value is desired. The pH of the water is not very useful for water identification or correlation purposes, but it does indicate possible scale-forming or corrosion tendencies of a particular water. The pH also may indicate the presence of drilling-mud filtrate or well-treatment chemicals. Organic acids, such as acetic acid, also can control the pH. The following is a typical reaction.
The pH of concentrated brines usually is less than 7.0, and the pH will rise during laboratory storage, indicating that the pH of the water in the reservoir probably is appreciably lower than many published values. In pure water or brines with little buffering capacity, like seawater, the addition of gas containing CO2 at high pressure can depress the pH to less than 2.9, making the water very reactive. This water will dissolve and corrode steel with great rapidity or, if in the reservoir, will dissolve minerals either wholly or partly. This can lead to formation damage and dramatically reduce injection and production because the newly dissolved species reprecipitate as the pressure drops at the producer well.
Addition of the carbonate ion to sodium chloride solutions will raise the pH. If enough calcium is present, calcium carbonate precipitates. The reason the pH of most oilfield waters rises during storage in the laboratory is because of the formation of carbonate ions as a result of bicarbonate decomposition caused by evolution of dissolved CO2 gas. An important consideration of CO2 gas evolution/dissolution is that it is not anything close to instantaneous; a fact that has been underappreciated by many, with very expensive and confusing consequences.
In pure water, the CO2 equilibrium takes on the order of tens of minutes (≈20 min.) to adjust to a change in CO2 pressure and for the pH to stabilize to a new level. However, with large amounts of bicarbonate in an oilfield water, the adjustment is even slower, while the buffering action of the bicarbonate itself will limit how much the pH will eventually change.
Organic acids play an extremely important role in the water chemistry. Because the volatile fatty acids, such as formic, acetic, propionic, and butyric acids, are quite commonly found in the waters, they can control the water chemistry to a large degree, especially the CO2/bicarbonate system. From a historical standpoint, this is important because analytical difficulties prohibited obtaining organic acid compositional data. Thus, much of the confusing behavior that workers observed in scale-deposition predictions based on the analysis of inorganic species compared with the actual field results turns out to be explainable once the organic acids are considered.
Historically, the typical analytical procedure for bicarbonate, the alkalinity titration, also happens to titrate the organic acids because they have Ka values similar to that of bicarbonate. Thus, scale predictions that used those bicarbonate values are somewhat incorrect, with the degree of error depending on the amounts of organic acids that were included with the bicarbonate value. Oilfield waters sometimes will have an unusual odor, which often comes from rather high concentrations of these organic acids. The formic, acetic, propionic, and butyric acids will not precipitate scale under most conditions, but certainly do buffer the water system effectively. Also, they seem to slow down the approach to CO2 equilibrium as well, so that water samples containing several hundred ppm of organic acids will not change their pH significantly when stored for several days. This also means that the dissolved CO2 in the produced water remains high even after the pressure has been reduced in a separator. It still can remain corrosive, even though it would not normally be expected to be very corrosive. One procedure to correct the bicarbonate analysis for the volatile fatty organic acid concentration is to measure the organic acid content by an independent technique, such as IC or CE, calculate the equivalent amounts of the acids, and then subtract those equivalents from the apparent bicarbonate concentration as measured by the alkalinity titration.
Naphthenic acids can precipitate and form scales, in contrast to the volatile fatty acids. Calcium naphthenate scale deposits have been identified recently in several fields that produce high-acid-number crude oils; however, the concentration of naphthenic acids in water is limited by their higher molecular weights and high oil solubility.
The redox potential (often abbreviated as Eh) may be referred to as oxidation potential, oxidation/reduction potential, or pE. It is expressed in volts or millivolts (mV), and, at equilibrium, it is related to the proportions of oxidized and reduced species present. Standard equations of chemical thermodynamics express the relationships.
The Nernst equation expresses the relationship between concentrations of oxidation-reduction couples. For example, a common redox couple involves the dissolved iron species Fe(II) and Fe(III), which can be described thermodynamically as
Eh is usually measured with a platinum electrode against a different reference, such as Ag/AgCl or saturated calomel reference electrodes. Knowledge of the redox potential is useful in studies of how compounds such as uranium, iron, sulfur, and other minerals are transported in aqueous systems. The solubility of some elements and compounds depends on the redox potential and the pH of their environment.
Some water associated with petroleum is interstitial (connate) water and has a negative Eh, which has been proved in various field studies. Knowledge of the Eh is useful in determining how to treat a water before it is reinjected into a subsurface formation. For example, the Eh of the water will oxidize if the water is open to the atmosphere, but, if it is kept in a closed system in an oil production operation, the Eh should not change appreciably as it is brought to the surface and reinjected. In such a situation, the Eh value is useful in determining how much iron will stay in solution and not deposit in the wellbore.
Organisms that consume oxygen lower the Eh. In buried sediments, it is the aerobic bacteria that attract organic constituents that remove the free oxygen from the interstitial water. Sediments laid down in a shoreline environment will differ in degree of oxidation compared with those laid down in a deepwater environment. For example, the Eh of the shoreline sediments may range from −50 to 0 mV, but the Eh of deepwater sediments may range from −150 to −l00 mV.
Aerobic bacteria die when the free oxygen is totally consumed; anaerobic bacteria attack the sulfate ion, which is the second most important anion in the seawater. During this attack, the sulfate reduces to sulfide, the Eh drops to negative potentials (approximately −600 mV), and H2S is liberated. This process is known as reservoir souring and is a major concern to engineers working on fields undergoing waterflood with injected seawater or other sulfate-containing injectant. Most waterfloods have eventually gone sour. Hydrogen sulfide generation causes problems from a health and safety standpoint because it is so poisonous. H2S also causes rapid, nearly instantaneous, failure of steel because of sulfide stress-corrosion cracking, unless the steel has been specified for "sour service." Besides the presence of sulfate ions, dissolved organic acids play a role in feeding the SRB. Predicting and mitigating reservoir souring is an active area of research. SRB and other bacteria often cause a different, much slower type of pitting corrosion on steel, known as microbially induced corrosion (MIC). MIC is commonly seen in low-flow piping areas, under deposits of solids or sludges, or in vessels and tanks.
|E||=||the voltage of the system vs. the standard hydrogen electrode|
|Eo||=||the voltage of the oxidation reaction at standard conditions (1 mole/liter, 298 K, 1 atmosphere pressure)|
|n||=||the number of electrons transferred in the reaction|
|R||=||the ideal gas constant|
|S||=||salinity in wt%|
|μw||=||viscosity of water|
- McCain, W.D. Jr.: McCain, W.D. Jr. 1990. The Properties of Petroleum Fluids, second edition. Tulsa, Oklahoma: PennWell Books.
- McCain Jr., W.D. 1991. Reservoir-Fluid Property Correlations-State of the Art (includes associated papers 23583 and 23594 ). SPE Res Eng 6 (2): 266-272. SPE-18571-PA. http://dx.doi.org/10.2118/18571-PA Cite error: Invalid
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- Collins, A.G. 1975. Geochemistry of Oilfield Waters. New York: Elsevier Scientific Publishing Co.
- Rogers, P.S.Z. and Pitzer, K.S. 1982. Volumetric Properties of Aqueous Sodium Chloride Solutions. J. Phys. Chem. Ref. Data 11 (1): 15–81. http://dx.doi.org/10.1063/1.555660
- Amyx, J.W, Bass, C.M. Jr., and Whiting, R.L. 1960. Petroleum Reservoir Engineering. New York City: McGraw-Hill Book Co. Inc.
- Kestin, J., Khalifa, H.E., and Corrcia, R.J. 1981. Tables of the Dynamic and Kinematic Viscosity of Aqueous NaCl Solutions in the Temperature Range 20-150°C and the Pressure Range 0 1-35 Mpa. J. Phys. Chem. Ref. Data 10 (1): 71.
- Zemaitis, J.F. Jr. et al. 1986. Handbook of Aqueous Electrolyte Thermodynamics. New York: American Inst. of Chemical Engineers, New York.
- Carothers, W.W. 1976. Aliphatic Acid Anions and Stable Carbon Isotopes of Oil Field Waters in the San Joaquin Valley, California. MS thesis, San Jose State University, San Jose, California.
- Baugh, T.D., Wolf, N.O., Mediaas, H., Vindstad, J.E., and Grande, K. 2004. Characterization of a Calcium Naphthenate Deposit—The ARN Acid Discovery. Petroleum Chemistry Division Preprints 2004, Am. Chem. Society Annual Meeting 49 (3): 274.
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