Chapter 5. Significance of Some Inorganic Constituents and Physical Properties of Oilfield Waters

Chapter 5.

SIGNIFICANCE OF SOME INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES OF OILFIELD WATERS

In general, the concentrations of the constituents in various natural solids of reservoir rocks must be considered along with the amounts that are found in associated oilfield waters. Some possible chemical reactions between host rock and reservoir water may deplete or enrich the concentration of the constituents in oilfield waters. Another important factor is the solubility of a constituent. The ionic potential, determined by dividing the ionic radius by the valence, influences the solubility of elements. For example, elements with low ionic potential are more likely t o remain in true ionic solution. Elements commonly found in oilfield waters have the following ionic potentials: sodium, 0.95; calcium, 0.50; magnesium, 0.33; chlorine, 1.81; bromine, 1.95; and iodine, 2.16. Apparently the cation (magnesium) and the anion (chlorine) would be the most likely to remain in true ionic solution; however, several other variables occur during diagenesis which lead to depletion or enrichment of constituents in waters.

Lithium Lithium is the lightest alkali metal; it has a distinctly smaller radius, 0.60 8,than the other alkalies and is the smallest of all singly charged cations. It is one of the less abundant elements, and its abundance in the earth’s crust is about 6.5 x wt.% (Fleischer, 1962). Here again, it is an exception because in general, the lighter elements tend to be more abundant than the heavier elements. It is lithophilic in that it tends t o be associated with the silicate phase in rocks (Ahrens, 1965); however, because of its small size, it supposedly cannot replace the abundant alkali metals in mica. It and the other alkali metals exist in a uniform positive one state of oxidation and are inherently ionic. Their chemical behavior depends almost entirely upon electron loss, and their chemistry is simpler than that of any of the other metallic elements (Moeller, 1954). Lithium is potentially toxic to plants (Hem, 1970), yet it is regularly found in plant ashes, which indicates that it normally is present in soil waters (Goldschmidt, 1958). Coal ashes of Neurode, Silesia, contained up to 198 ppm lithium, whereas soils in northeast Scotland contain 30-5,000 ppm. The content of lithium in sediments ranges up to 6 ppm in quartzites and sandstones, up to 15 ppm in calcareous rocks, and up t o 120 ppm in clays and shales.

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

134 TABLE 5.1

Properties of the alkali metals Property

Lithium

Sodium

Potassium

Rubidium

Cesium

Atomic number Nonhydrated radius (A) Hydrated radius (A) Outer electronic configuration Atomic weight Ionization potential (V)

3

11

19

37

55

0.60

0.95

1.33

1.48

3.82

3.58

3.31

-

1s' 2s' 6.939 5.390

2s22p6 3s' 22.990 5.138

3s' 3p6 4s' 39.102

-

4s24p6 5s' 85.47

4.339

1.69

4.176

5s' 5p6 6s' 132.905 3.893

TABLE 5.11 Five relative concentration changes of some dissolved ions during evaporation of sea water and brine* Constituents

Concentrations (mg/l) Sea water

Lithium Sodium Potassium Rubidium Magnesium Calcium Strontium Boron Chloride Bromide Iodide

CaSO4

0.2 2 11,000 98,000 350 3,600 0.1 1 1,300 13,000 400 1,700 60 7 5 40 19,000 178,000 65 600 0.05 2

NaCl

MgS04

11 12 140,000 70,000 23,000 37,000 6 8 74,000 80,000 100 10 1 10 300 310 275,000 277,000 4,000 4,300 5 7

KCI

27 13,000 26,000 14 130,000 0 0 750 360,000 8,600 8

MgC12

34 12,000 1,200 10 153,000 0 0 850 425,000 10,000 8

*Approximate mg/l. Columns headed sea water, CaS04, etc., represent stages in sea water evaporation. For example, sea water contains 0.2 mg/l of lithium, after calcium sulfate has precipitated the residual brine contains about 2 mg/l of lithium, after sodium chloride has precipitated the residual brine contains about 11 mg/l of lithium, the residual brine contains about 12 mg/l of lithium after magnesium sulfate precipitates, 27 mg/l of lithium after potassium chloride precipitates, and 34 mg/l of lithium after magnesium chloride precipitates.

LITHIUM

135

a,

The hydrated radius of lithium is 3.82 as shown in Table 5.1 (Moeller, 1954). The ionic potential is 0.60, and the polarization is 1.67. The polarization is quite high and is a measure of its replacing power in an exchange system. Apparently it can replace strontium, calcium, and magnesium since their polarizations are 1.77, 2.02, and 3.08, respectively. Some surface waters of the volcanic sodium chloride type are enriched in lithium (White, 1957). Lithium from Searles Lake brine is recovered as Li2NaP04 (Brasted, 1957). The content of lithium in oilfield waters is usually less than 10 mg/l but in some Smackover formation waters from east Texas, concentrations up t o 500 mg/l are present. When a brine containing lithium goes through an evaporite sequence, lithium is one of the elements whose concentration does not decrease, as illustrated in Table 5.11, in the liquid phase as various minerals precipitate (Collins, 1970). Fig. 5.1 illustrates the enrichment of lithium as compared t o an evaporite sequence in some subsurface brines from Tertiary, Cretaceous, and Jurassic age sediments. Fig. 5.2 illustrates a similar enrichment for some brines taken from Pennsylvanian and Mississippian age sediments (Collins, 1969a). Possibly lithium was liberated and potassium was depleted by exchange reactions with clay minerals, degradation of lithium containing minerals, or simply a leaching of minerals, primarily silicates, which contain lithium. Lithium substitutes in the structure of several common minerals and forms few minerals of its own. If the minerals in which it has substituted should degrade or break down with depth, the lithium might be resolubilized, thus increasing its concentration in the aqueous phase. White et al. (1963) postulated that because the lithium concentration in magmatic waters is related to volcanic

LITHIUM, mgll

Fig. 5.1. Comparison of the lithium concentrations in some Tertiary (T),Cretaceous (C), and Jurassic (J) age formation waters from Louisiana with an evaporating sea water.

136

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

u 10

20

3

LITHIUM, mg/l Fig. 5.2. Comparison of the lithium concentrations in some Mississippian (M) and Pennsylvanian (P) age formation waters from Oklahoma with an evaporating sea water.

emanations, the increase in the lithium content of deeper waters might be related to the same cause.

Sodium The most abundant member of the alkali-metal group is sodium, ranking number 6 with respect t o all the metallic elements. The radius of the sodium ion is 0.95 A, and its geochemistry is controlled to some extent by calcium because of the similarity of their ionic radii. Its abundance in the earth's crust is about 2.8 wt.% (Fleischer, 1962). Table 5.1 shows that its outer electronic configuration is 2s' 2p6 3s' , with a first ionization potential of 5.138 V, indicating that its single outer electron is less firmly held than in the lithium atom with a first ionization potential of 5.390 V. The ionization potential is a measure of the chemical reactivity - the lower the potential, the greater the reactivity. Table 5.1 (Moeller, 1954) also illustrates some qf its other properties. According t o Ahrens (1965),sodium is lithophilic, and many distinctly lithophile elements have valence electrons outside a closed shell of eight electrons. The ionic radius decreases as the charge on the cation increases. Sodium does readily participate in solid solution relationships because its radius is small, making replacement of cations with 30% larger radii difficult. The amounts of sodium in argillaceous sediments and marine shales are about 1,000ppm and 1,300ppm, respectively (Goldschmidt, 1958).

SODIUM

137

Sodium in solution tends to stay in solution; it does not readily precipitate with an anion, and it is less easily adsorbed by clay minerals than are cesium, rubidium, potassium, lithium, barium, and magnesium. The major source of sodium in sea water can be attributed t o the weathering of rocks. Some sodium probably was derived through volcanic activity. The ocean and evaporite sediments contain the bulk of the sodium. Igneous rocks contain appreciably more sodium than sedimentary rocks with the exception of evaporites. Sea water contains about 11,000 mg/l of sodium, as illustrated in Table 5.11. The concentration of sodium increases in brine as it evaporates, t o about 140,000 mg/l, when halite precipitates. Most oilfield waters contain more sodium than any other cation, and most oilfield waters are believed to be of marine origin. Fig.5.3 is a log-log plot of the chloride concentration versus sodium of some subsurface brines taken from sediments of Tertiary, Cretaceous, and Jurassic age. The straight line is a plot of chloride versus sodium concentrations for some evaporite waters, and indicates the enrichment of sodium ions until halite (NaC1) precipitates - at a chloride concentration of about 140,000 mg/l (compared t o that of normal sea water, 19,000 mg/l). The plot of the concentrations of sodium versus chloride for these subsurface brines falls very near the normal evaporite curve, indicating that the concentration mechanism may be related to an evaporite process (Collins, 1970). Fig. 5.4 is a similar plot for some subsurface brines taken from sediments of Pennsylvanian and Mississippian age (Collins, 1969a). Several of these samples are somewhat depleted in sodium which indicates that

SODIUM,

g/l

Fig. 5.3. Sodium versus chloride concentrations for some formation waters taken from Tertiary (T), Cretaceous (C), and Jurassic (J) zge sediments and compared to evaporating sea water.

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

138

500F --/ 200t-

0

I-

/

- Nor ma1 evaoor it e associated b i n e

/"

SODIUM, mg/l Fig. 5.4. Sodium versus chloride concentrations for some formation waters taken from Pennsylvanian (P) and Mississippian (M) age formation sediments and compared t o evaporating sea water.

diagenetic processes, such as ion-exchange or ultra-filtration reactions involving clays and/or carbonates, may operate to deplete the sodium concentration in waters in older sediments. Potassium The second most abundant member of the alkali-metal group is potassium; its abundance in the crust of the earth is about 2.55 wt.% (Fleischer, 1962). Like the other alkali metals, it is lithophilic, and with its large ionic radius (see Table 5.1) it participates in forming solid solutions and forms its own minerals, such as feldspar and mica. The potassium feldspars are resistant to leaching by water, which may account for the low potassium concentrations in many natural waters. Clay minerals readily adsorb potassium, and in illite it is incorporated into the crystal structure in such a manner that it cannot be removed by ion-exchange reactions (Lyon and Buckman, 1960). Potassium is less easily hydrated than sodium, and is more easily adsorbed by colloids; therefore, it is retained in sediments and soils in greater abundance than sodium. It is an essential element t o plants and animals. According to Gol&chmidt (1958),potassium in pulverized potassium feldspars is absolutely unavailable t o plants. The concentrations of potassium in carbonates, sandstones, and shales is about 2,700, 10,700,and 26,600 ppm, respectively (Mason, 1966). Potas-

139

POTASSIUM .lvv

200

~~

-

- 100 -

1 POTASSIUM,

I I I IIll

g/ I

Fig. 5.5. Potassium versus chloride concentrations for some formation waters taken from Tertiary (T), Cretaceous (C), and Jurassic (J) age sediments and compared to evaporating sea water.

sium concentrates primarily in hydrolysates (clay minerals), such as illite and glauconite, and in evaporites. Table 5.11 illustrates how the concentration of potassium in the aqueous phase increases until sylvite (KC1) precipitates. The concentration of potassium in some subsurface brines usually is depleted with respect to an evaporite-associated sea water. Fig.5.5 illustrates the relation of potassium in some subsurface brines taken from sediments of Terti500 -

,'

-

-Nmal

evaporite curve

-

5--

m

POTASSIUM, mg/l Fig. 5.6. Comparison of the potassium concentrations in some Pennsylvanian (P) and Mississippian (M) age formation waters from Oklahoma with an evaporating sea water.

140

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

ary, Cretaceous, and Jurassic ages to an evaporite-associated sea water (Collins, 1970). Fig.5.6 illustrates the same relation for some subsurface brines taken from Pennsylvanian and Mississippian age sediments (Collins, 1969a). The depletion of potassium in subsurface brines might be caused by its uptake by clays. For example, montmorillonite-type clay minerals systematically change to illite with increasing depth of burial, due to thermal diagenesis; and, as a result of this transformation, they lose interlayer (bound) water (Burst, 1969). This change appears t o begin at a temperature above 90°C. (This freed interlayer water can be readily expelled, and its movement probably is important in the first migration stage of hydrocarbons.) Laboratory experiments at elevated temperatures and pressures indicate that montmorillonite loses its interlayer water and transforms into illite in the presence of potassium-enriched water (Khitarov and Pugin, 1966). The structural variations of the expandable minerals in clays apparently are also influenced by the potassium content of the associated waters.

Rubidium Rubidium, like the other alkali metals is lithophilic, and its abundance in wt.%, which is greater than that of the earth’s crust is about 3.0 x lithium (Fleischer, 1962). It tends to be removed from solution more readily than lithium, primarily because of its ability to replace potassium in mineral structures. Table 5.11 indicates that it precipitates from an evaporite along with sylvite to a greater extent than lithium, and it has a high chemical is only about 10% larger than the reactivity. The radius of its ion, 1.48 potassium ion, so it can be accommodated into the same crystal lattices. Because of this, it forms no minerals of its own. Rubidium and cesium occur sympathetically in nature; that is, both are commonly found in amazonite, vorobyevite, and beryl (Goldschmidt, 1958). Rubidium is a member of series NH4-K-Rb-Cs, and members of this series are more similar in their chemical and physical properties than are the members of any other group, with the exception of the halogens. Rubidium concentrates in the late crystallates, particularly those of granitic derivation, and it has a greater tendency t o be adsorbed by clays than has potassium. It is removed from igneous rocks by water leaching and then adsorbed by hydrolysate sediments and soils. Shales contain about 250 ppm of rubidium; deep-sea red clays, about 400 ppm; and some glauconites, about 500 ppm (Goldschmidt, 1958).Sea water contains about 0.12 mg/l of rubidium; subsurface brines contain up t o 4 mg/l. Higher concentrations of rubidium probably can be found in brines associated with rocks containing potassium minerals, such as microcline feldspars, or lepidolite mica.

a,

141

CESIUM

Cesium Cesium is the heaviest alkali metal and also the rarest, with an abundance of about 7 x wt.% in the earth’s crust (Fleischer, 1962). It has an ionic radius of 1.69 8,which is distinctly larger than potassium, and it cannot replace potassium in minerals as easily as rubidium; probably because of this, it forms its own minerals. It is leached from igneous and metamorphic rocks by water during weathering, and is adsorbed by hydrolysate sediments and soils more readily than rubidium or potassium. Its low ionization potential indicates that it has the greatest chemical reactivity of the alkali metals. Cesium and rubidium were discovered in 1860 by Robert Bunsen by use of spectral analysis, a method which he and Kirchhoff invented. Cesium concentrates primarily like rubidium, in marine argillaceous sediments. Some shales contain about 15 ppm; deep-sea red clays, 20 ppm; and glauconite, 15 ppm of cesium (Goldschmidt, 1958). Sea water contains 5x mg/l of cesium, and some subsurface brines contain up to 1mg/l. Beryllium Beryllium is a member of the alkaline earth group in the periodic chart of the elements, but few of its properties are similar t o the more abundant members, such as magnesium, calcium, and strontium. Beryllium, like lithium, is a light element with an atomic weight of 9.012 (Table 5.111; see also Moeller, 1954), and like lithium, it is an exception t o the rule that light elements are more abundant than heavy elements. The earth’s crust contains wt.% of beryllium (Fleischer, 1962). about 6 x In sedimentary rocks, beryllium is restricted primarily to hydrolysates and especially to bauxites enriched in aluminum (Goldschmidt, 1958). Shales contain about 6 ppm, and some coal ashes contain up to 8,000 ppm, although generally only about 4 ppm. The concentration of beryllium in sea TABLE 5.111 Properties of the alkaline earth metals Property

Beryllium Magnesium

Atomic number 4 Ionic radius (A 1 0.31 Outer electronic configuration 1s’ 2s’ Atomic weight 9.012 Ionization potential (V) 9.320

12 0.65 2s’ 2p6 3s’ 24.31 7.644

Calcium

Strontium

Barium

20

38

56

0.99 3s2 3p6 4s2 40.08 6.111

1.13 4s’ 4p6 5s2 87.62 5.692

1.35 5s2 5p66s’ 137.34 5.210

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

142

water is about 5 x lo-' mg/l, and some subsurface brines contain 0.02-4.2 mg/l. Since beryllium is highly toxic, waters containing it should be handled with caution.

Magnesium One of the more abundant members of the alkaline earth group of metals, magnesium makes up about 2.1 wt.% (Fleischer, 1962) of the crust of the earth. Magnesium is dissolved during chemical weathering, mainly as the chloride and sulfate. Ferromagnesian minerals in igneous rocks and magnesium carbonate in carbonate rocks are generally considered t o be the principal sources of magnesium in natural waters. Carbon dioxide plays an important role in the dissolution of magnesium from silicate and carbonate minerals. Waters associated with either granite or siliceous sand may contain less than 5 mg/l of magnesium, whereas those associated with either dolomite or limestone may contain over 2,000 mg/l of magnesium. Elements commonly found in oilfield waters have the following ionic potentials: sodium, 0.95; calcium, 0.50; magnesium, 0.33; chlorine, 1.81; bromine, 1.95; and iodine, 2.16. Apparently the cation (magnesium) and the anion (chlorine) would be the most likely to remain in true ionic solution; however, several other variables occur during diagenesis which lead to depletion of magnesium in waters. Depletion of magnesium in some waters probably is a result of the replacement reaction t o form dolomite, CaMg(C0, ) 2 . Whole mountain masses are made of dolomite, which is formed by the regular substitution in the calcite

2oo

t

J

C

/

?$

Normal evaporite curve

'so0

500

rpoo

2,000

5ooO

MAGNESIUM, mg I I

lop00

2 0 m ,

5Q(

Fig. 5.7. Comparison of the magnesium concentrations in some Tertiary (T), Cretaceous (C), and Jurassic (J) age formation waters from Louisiana with an evaporating sea water.

CALCIUM

c

143

Normal evaporite curve

500

M M P

r 20 10

1,000

I 0,000 lO0,OoO MAGNESIUM, mg/l

Fig. 5.8. Comparison of the magnesium concentrations of some Pennsylvanian (P) and Mississippian (M) age formation waters from Oklahoma with an evaporating sea water.

crystal lattice of alternate ions of calcium and magnesium. The large differences in the ionic radii of Ca (0.99A) and Mg (0.65A) are the reason for this diadochy. Magnesium ions in aqueous solution have a large attraction for water molecules and probably are surrounded by six water molecules in octahedral arrangement. This may account for the paucity of magnesium in soils, because the small cation becomes large by hydration. Sodium has a similar reaction, but potassium, which does not, is readily adsorbed by soil colloids. Shales, sandstones, and carbonates contain 15,000,7,000,and 47,000 ppm of magnesium, respectively (Mason, 1966). Subsurface brines contain from less than 100 mg/l t o more than 30,000 mg/l; however, many subsurface brines are depleted in magnesium if compared to a sea water evaporite sequence, (Table 5.11). Sea water contains about 1,300 mg/l. Fig. 5.7 is a plot of chloride versus magnesium for some subsurface brines taken from Tertiary, Cretaceous, and Jurassic age sediments. The position of the normal evaporite curve indicates that all of these waters were depleted in magnesium with respect to this curve (Collins, 1970). Fig. 5.8 is a plot showing similar depletion of some subsurface brines taken from some sediments of Pennsylvanian and Mississippian age.

Calcium The abundance of calcium in the crust of the earth is about 3.55 wt.% (Fleischer, 1962),making it the most abundant of the alkaline earth metals,

144

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

but only in the crust; in the earth as a whole, magnesium is much more abundant. Calcium is dissolved as bicarbonate as a result of chemical weathering of calcium-bearing minerals. Waters associated with limestone, dolomite, gypsum, or gypsiferous shale usually contain an abundance of calcium, but waters associated with granite or silicious sand may contain less than 10 mg/l of calcium. Slight changes in the pH of waters containing calcium bicarbonate will cause calcium carbonate to precipitate, and calcium carbonate is one of the most common deposits found in plugged oilfield lines, equipment, and reservoirs. Precipitation of calcium carbonate in the sea is the prime mode of the origin of limestone. The solubility of calcium carbonate in sea water increases with salinity and increasing partial pressure of carbon dioxide, but it decreases with increasing pH, calcium content, and temperature. The solubility of calcium sulfate decreases with increasing temperature. Shales, sandstones, and carbonate rocks contain about 22,100, 39,100, and 302,300ppm of calcium, respectively (Mason, 1966).Sea water contains 400 mg/l and subsurface brines often contain 2,000-3,000 mg/l, with some as high as 30,000 mg/l. Fig. 5.9 is a plot of chloride versus calcium concentrations for some subsurface waters taken from Tertiary, Cretaceous, and Jurassic age sediments. The amount of calcium in these waters increases with increasing salinity, and the waters from the older sediments appear to contain more calcium. Fig. 5.10 is a similar plot for some subsurface brines taken from sediments of Pennsylvanian and Mississippian age. These samples all appear to be enriched in calcium relative t o the evaporite curve, and the concentration of calcium appears to increase with increasing salinity.

200

-

\ 0

Normal evaporite curve

-

100-

500

1 , m

2

1

p

I 1 I I I111 5poo

lop00

29ooo

CALCIUM, mg/l

Fig. 5.9. Comparison of the calcium concentrations of some Tertiary (T), Cretaceous (C), and Jurassic (J) age formation waters from Louisiana with an evaporating sea water.

STRONTIUM

145

&-\-

Normal evaporite curve

M

P P

Ii 201

/

M M 1

00 CALCIUM, mg/ I Fig. 5.10. Comparison of the calcium concentrations of some Pennsylvanian (P) and MEsissippian (M) age formation waters from Oklahoma with an evaporating sea water.

Strontium Strontium, a minor element compared t o calcium and magnesium comprises about 0.03 wt.% of the earth's crust (Fleischer, 1962). Table 5.111 illustrates some of its properties, and it resembles calcium chemically. Strontium has a tendency to work upward during fractional crystallizaticn because of its relatively large radius (Goldschmidt, 1958).It occurs abundantly with potassium in volcanic rocks, alkali rocks, and pegmatites. Dissolved strontium results from water leaching of rocks, and it has been postulated that the strontium in petroleum-associated waters also may be a byproduct of the organic decay processes which originally formed petroleum. Strontium is only a microconstituent in most terrestrial animals, but several species of marine animals contain considerable quantities of strontium in their skeletons (Odum, 1951). Table 5.11 indicates that strontium may reach a concentration of 60 mg/l during sea-water evaporation, and then most of it precipitates with calcium sulfate. The amount of sulfate in the water influences the amount of strontium that remains in solution. Data by Sillhn and Martell (1964)indicate that if the sulfate activity in a water is 100 mg/l, the strontium activity can be about 28 mg/l. Davis and Collins (1971)studied the solubility of strontium sulfate in strong electrolyte solutions and found that 958 mg/l of strontium is soluble in a synthetic brine solution of ionic strength 3.05,

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

146

containing ions of sodium, calcium, magnesium, potassium, chloride, bromide, and iodide. Calcium chloride concentration apparently has a very pronounced effect upon the solubility of strontium sulfate. Celestite and strontianite occur commonly in sediments. Carbonate sediments contain up t o 1,200 ppm of strontium; dolomites, usually less than

“““I

I C -Cretaceous J -Jurassic

C

J

c cc 2,000 C

cC

T

20

1 0

50

I IIII

100

I

2 a

I IIL

STRONTIUM, mgll

Fig. 5.11. Comparison of the strontium concentrations of some Tertiary (T), Cretaceous (C), and Jurassic (J) age formation waters from Louisiana with an evaporating sea water. 500

-

-

-

200 -

\

P

Ill

I

50

100 ZOO STRONTIUM, mg/l

500

1,000

Fig. 5.12. Comparison of the strontium concentrations of some Pennsylvanian (P) and Mississippian (M) age formation waters from Oklahoma with an evaporating sea water.

BARIUM

147

170 ppm; and secondary gypsum, up t o 1,100 ppm (Goldschmidt, 1958). Sea water contains about 8 mg/l of strontium, but subsurface brines contain up to 3,500 mg/l. Fig. 5.11 is a plot of chloride versus strontium content for some subsurface brines taken from some Tertiary, Cretaceous, and Jurassic age sediments. Most of these samples were enriched in strontium compared to the evaporite-associated water, and it is possible that a mechanism similar to dolomitization could cause the enrichment. In comparison t o calcium, the strontium appears to be increasingly accumulated; for example, only five samples (from Tertiary sediments) fell within the normal evaporite curve. Fig. 5.12 is a similar plot for some subsurface brines showing similar results taken from sediments of Mississippian and Pennsylvanian age. Barium Barium, like strontium, is a minor element, comprising 0.04 wt.%, of the earth’s crust; it is more concentrated in igneous rocks and less concentrated in sedimentary rocks (Fleischer, 1962). It, like the other alkaline earth metals, is predominantly lithophile. Table 5.111 illustrates some of the properties of barium; its ionic radius, 1.35 A, permits it t o replace potassium, but usually not calcium and even less commonly magnesium. Barium forms more of its own minerals than does strontium. Barium is readily absorbed by colloids, like potassium, and is therefore retained by soils or precipitated with hydrolysates; it is also concentrated in deep-sea manganese nodules (Hem, 1970). Barium dissolves as bicarbonate, chroride, or sulfate during weathering processes, and migrates in aqueous solutions as these compounds. The solubility of barium sulfate increases when hydrochloric acid or chlorides of the alkali or other alkaline earth metals are present in solution. The properties of barium are similar t o those of strontium. Both precipitate through loss of carbon dioxide from a bicarbonate-bearing solution, or as sulfates by the action of sulfuric acid, sulfides, or sulfates. Strontium, however, is less likely t o be absorbed by clays than barium, because its ionic radius is smaller and its ionic potential is higher. Encrustation deposits taken from plugged pipes of waterflood systems for secondary recovery of oil, where barium is present, usually contain barium, calcium, strontium, iron, and traces of other metals. Barium may cause problems in waterflood systems by reacting with the chromate-type oxygencorrosion inhibitors, forming water-insoluble barium chromate. The amount of barium found in sandstones, shales, and carbonates is about 180, 450, and 90 ppm, respectively (Goldschmidt, 1958).Sea water contains about 0.03 mg/l, and subsurface brines may contain more than 100 mg/l; however, many brines contain less than 10 mg/l. Davis and Collins (1971)found that 59 mg/l of barium sulfate is soluble in a synthetic brine with an innic strength of 3.0487, containing sodium, calcium, magnesium,

TABLE 5.IV Properties of aluminum. copper. iron, lead, manganese, and zinc property

Aluminum

Copper

Iron

Lead

Manganese

Zinc

Atomic number Ionic radius (A)

13 0.50

26 0.76(+2) 0.64(+3) . .

82 1.20(+2) 0.84(+4) . .

25 0.80(+2) 0.46c+7 1..

30 0.74

Outer electronic configuration Atomic weight Ionization potential (V)

27 0.96(+1) 0.691+21 . ,

2s22p63s23p1 26.98 5.984

3s23p63d'04s' 63.54 1.723

3s23p63d64s' 55.54 1.165

3s' 3p6 3d54s' 54.938 1.168

3s2'3p63d" 4s' 65.37 9.391

4d'05s'5p64f'5d106s'6p' 207.19 7.415

MANGANESE

149

potassium, chloride, bromide, and iodide ions. Many analyses performed by wet chemical methods indicate rather high concentrations of barium in some subsurface brines. Some of these high results probably should be attributed to strontium plus barium rather than barium only, because satisfactory separation of the two in wet chemical methods is very difficult to accomplish. Manganese Manganese is a member of the VII B group of elements and is well known for its multiplicity of oxidation states. Essentially it is cationic, and the Mn+4 oxidation state usually is found in sediments. Its (+2) ionic radius is 0.80 8,while the ferric iron radius is 0.76 8 (see Table 5.IV); reasonable amounts of interchange in crystal lattices between these two ions are possible. The abundance of manganese is about 0.1 wt.% of the earth’s crust (Fleischer, 1962). Manganese is present in many oilfield brines because it is readily dissolved by waters containing carbon dioxide and sulfate. Except for titanium, manganese is the most abundant trace element in igneous rocks. Nearly all mineral groups of petrological importance contain manganese. During weathering, manganese is dissolved mainly as the bicarbonate. Decomposition of the bicarbonate leads to the formation of M d 4 compounds. In a reducing type of environment Mn+ compounds migrate in aqueous solutions. Mn+ compounds are less mobile, and Mn+4 compounds precipitate from aqueous solutions. In general, manganese remains in solution at a low redox potential and precipitates at a high redox potential. According to Goldberg (1963), manganese oxide nodules on the ocean bottom occur in both shallow water and deep water environments. He attributes these deposits to slow oxidation of dissolved manganese in areas where the waters contact an oxide surface. In most subsurface brines, the manganese is in the reduced form (Mn+*)because the redox potential is low and in subsurface brines probably would be the pH is less than 7.0. Any suspended with particulate matter or complexed by organic compounds, rather than in ionic solution. Shales and carbonates contain about 850 ppm and 1,100 ppm, respectively, of manganese (Mason, 1966). Sea water contains about 0.002 mg/l, and many subsurface brines contain 1.0 to 6.0 mg/l of manganese. Iron Iron is a member of the VIII group of elements and is predominantly siderophile. However, because it has an affinity for sulfur, it is also thiophile; and because it commonly enters into silicate minerals, it is lithophile as well. It is an ubiquitous element, with an abundance of about 5 wt.% of the earth’s crust (Fleischer, 1962).

150

INORGANIC CONSTITUENTS A N D PHYSICAL PROPERTIES

Iron, cobalt, and nickel possess atomic radii that differ only about 2% or less, so that the crystal chemistry of the three are related. The divalent ions of nickel, magnesium, cobalt, and iron have similar ionic radii; consequently, their chemistries in the sequence of isomorphous crystallization of mixtures are similar. The trivalent ions of iron and cobalt are similar in size, but the high oxidation potential of cobalt prevents much replacement (Goldschmidt, 1958). The solubility of iron compounds in ground waters is a function of the type of iron compound involved, the amounts and types of other ions in solution, the pH, and the Eh. According t o Larson and King (1954), 100 ppm of ferrous iron can stay in solution at pH 8 and pH 7; the theoretical maximum is about 10,000 ppm. The effects of many other ions, plus temperature and pressure differentials, such as those common to oilfield waters, have not been thoroughly studied. When a ground water in which ferrous iron is dissolved contacts the atmosphere, the following reaction can occur: 2Fe2++ 4HCO3- + H20 + 1 / 2 0

2

+ 2Fe(OH), + 4C02

Sandstone contains iron oxide, iron carbonate, and iron hydroxide, and shales and carbonate rocks contain oxides, carbonates, and sulfides of iron. Oilfield waters with characteristic low redox potentials dissolve some iron from the surrounding rock. The iron occurs in such waters at two levels of oxidation, ferrous or ferric. Knowledge of the amount and type of iron compounds in oilfield waters is used to estimate the amount of corrosion that is occurring in the production system, and t o determine the type of treatment required if the water is t o be used for waterflooding. This knowledge also enables determination of the Eh of the in situ water, because the Eh can be calculated from the Fe+2 and Fe+ values. Shales, sandstones, and carbonates contain about 47,200, 9,800, and 3,800 ppm, respectively, of iron (Mason, 1966). Sea water contains about 0.01 mg/l, and subsurface brines contain from traces to over 1,000 mg/l of iron. Copper Copper is a member of the VIII group of elements, and it is characteristically thiophile; the largest concentrations of it are found in various sulfur compounds. The earth’s crust contains about 0.01 wt.% of copper (Fleischer, 1962). Its compounds are dissolved easily during weathering, if the pH of the solution is less than 4.5. Many of the water-soluble copper compounds are salts of organic acids such as acetic, citric, and naphthenic. Much of the copper that is dissolved is precipitated afterward as sulfide. Traces of copper remain in the oceans, but its content is kept low because of the adsorption on, or combination with, marine organisms. Miholic (1947) presented an age

ZINC

151

division for mineral waters based on the presence of heavy metals in waters associated with joints and faults caused by tectonic movements of different geological ages. He placed copper as the predominant heavy metal in the Caledonian Group of the Orogenic Epoch (post-Silurian). Biochemical processes are known to be responsible for enriching a deposit in metals such as uranium, copper, and vanadium; therefore, this classification is restricted to waters of igneous origin. Most shales and carbonates contain about 45 and 4 ppm, respectively, of copper, with sandstones containing less than 1 ppm (Mason, 1966). Sea water contains about 0.003 mg/l, and most subsurface brines analyzed in this laboratory contained from less than 0.5 mg/l up to about 3 mg/l. The solubility of copper generally decreases with decreasing redox potential and increases with increasing redox potential if reduced sulfur is present. Most subsurface oilfield brines have relatively low redox potentials. zinc Zinc is a member of the I1 B group of elements and is predominantly thiophile. Its abundance in the crust of the earth is about 0.013 wt.% (Fleischer, 1962). Its geochemistry results from the similarity of its divalent ionic radius and the radii of Mg+’, Ni+?, Co+’, Fe+’, and Mn+’ (Goldschmidt, 1958). Zinc is dissolved readily as sulfate or chloride from acid rocks, such as granite, during weathering. Conversely, zinc is not dissolved easily from limestone with which it is deposited. Most alkaline waters do not extract zinc; however, a solution of NH,, NH,NO,, and NaC10, can extract and hold small quantities of zinc; the more acidic the water, the greater the amount of zinc extracted. Zinc is precipitated as the sulfide, oxide, carbonate, or silicate. Traces of zinc are found in sea water, but eventually zinc is deposited in carbonated sediments or in bottom muds or sapropels as sulfide. Shales, sandstones, and carbonates contain about 95, 16, and 20 ppm, respectively, of zinc (Mason, 1966). Sea water contains about 0.01 mg/l, and subsurface brines contain traces to more than 500 mg/l of zinc. Mercury Mercury is a member of the I1 B group of elements, which also includes zinc and cadmium. It is relatively abundant for a heavy element, but still must be considered scarce, with an abundance of about 4 x lo-’ wt.% of the crust of the earth (Fleischer, 1962). Most commercial deposits of mercury are of hydrothermal origin and are related to magmatic rocks; the commercial ore is cinnabar, HgS, or the liquid metal itself (Goldschmidt, 1958). Mercury is predominantly thiophile, and its geochemistry is controlled by the fact that it is volatile, with a boiling point of 357”C, and can be reduced to the metal by ferrous iron. Therefore, in a magmatic environment

152

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

the temperature and the redox potential control its occurrence. It is transported in hot springs (White et al., 1963). Shales, sandstones, and carbonates contain about 0.4, 0.03, and 0.04 ppm, respectively, of mercury. Sea water contains 3 x lo-’ mg/l, and subsurface oilfield brines contain 0-0.15 mg/l. The samples containing 0.15 mg/l of mercury were found in relatively dilute brines taken from the Cymric and the Rio Bravo oilfields in California. Free mercury is found in the oils produced from these fields, and the ages of the producing formations range from Eocene t o Pleistocene. The mercury content of natural waters has been used t o locate cinnabar deposits (Dall’Aglio, 1968). The amounts of mercury in waters appear t o increase with increasing bicarbonate concentration. Karasik et al. (1965) found that saline waters containing 200,000 mg/l of chloride contain very small amounts of mercury, which suggests that anionic complexes such as HgC14-* may not be important transporters of mercury. Brackish waters containing up t o 3,000 mg/l dissolved solids, up to 400 mg/l of bicarbonate, and the iodide ion sometimes contain up to 10 ppb of mercury, while stronger brines contain <0.1 ppb of mercury, which suggests that mercury may be transported as Hg14-* in brackish waters.

Lead Lead is a member of the IV A group of elements; it is ubiquitous in the earth, but its abundance in the crust is only about 0.002 wt.% (Fleischer, 1962). It is extracted from its minerals during weathering and migrates in the form of soluble-stable compounds. It is particularly soluble in acetic and other acids. Because the bicarbonate form is more soluble than the carbonate, lead can be transported as the bicarbonate. Most of the lead is precipitated from waters before they reach the sea. Hemley (1953) studied lead sulfide solubility related to ore deposition from saline waters. He concluded that lead-complex concentrations increase with increasing concentrations of bivalent sulfur and decrease at pH values above 7. The solubility of lead is limited primarily by the solubility restrictions of its sulfide and sulfate in reducing and oxidizing systems. How its solubility is influenced by many other ions, such as those found in a brine, has not been sufficiently studied. Shales, sandstones, and carbonates contain about 20, 7, and 9 ppm of lead, respectively. Sea water contains about 0.003 mg/l, and subsurface brines contain trace amounts to more than 100 mg/l of lead. Cadmium Cadmium is a member of the I1 B group of elements and may be considered one of the rarer elements; its abundance is about 3 x lo-’ wt.% of the earth’s crust (Fleischer, 1962). It is strongly thiophile, but its chemistry

BORON

153

differs from that of zinc in that it will precipitate from a strong acid solution, whereas zinc will not. There are few independent cadmium minerals, and its distribution is mainly that of a “guest” atom or ion in minerals. It frequently is present in lead-zinc deposits and occurs in solid solution in hypogene sulfides. A main carrier of cadmium is sphalerite, and oxidation of sphalerite or other sulfides containing cadmium will release the soluble cadmium sulfate. Shales and carbonates contain about 0.3 and 0.035 ppm of cadmium, respectively, and sandstones contain less than 0.01 ppm (Mason, 1966). Sea water contains about 0.0001 mg/l, and the subsurface oilfield brines may contain from 0 to about 0.001 mg/l of cadmium. Subsurface brines of the sulfate type in contact with lead-zinc deposits probably contain higher concentrations of cadmium. Boron Boron is a member of the I11 A group of elements, and it is an oxyphile and lithophile element. Its abundance in the crust of the earth is about 0.001 wt.% (Fleischer, 1962). It has small atomic and ionic radii. Knowledge of the presence of boron compounds in oilfield waters is important for several reasons. Boron is useful in identifying the sources of brines intrusive t o oil wells, or in fresh-water lakes or streams. In concentrations exceeding 100 mg/l, it affects electric log deflections. Boron is present in oilfield brines as boric acid, inorganic borates, and organic borates. When it is present as undissociated boric acid, it is an important buffer mechanism, being second only to the carbonate system. It may be precipitated as the relatively insoluble calcium and magnesium borates. Kazmina (1951) calculated the borate-chloride coefficient of some Russian oilfield waters. With a plot of the borate-chloride coefficient in logarithmic coordinates as a function of chloride content, he distinguished genetic groups of natural waters found in oil-bearing regions. Mitgarts (1956) studied the significance of boron and other elements in petroleum prospecting. In general, boron, together with bromine and iodine, is always associated with waters accompanying petroleum. Like chlorine, it can be considered an element of marine origin. The solubility of most boron compounds, the hydrolytic cleavage of boron salts, and their ability t o be occluded and coprecipitated with other compounds account for the extensive migration of boron. Soluble-complex boron compounds in brines and connate waters probably are there as a result of the decay of the same plants and animals that were the source of petroleum. Shales, sandstones, and carbonates contain about 100, 35, and 20 ppm, respectively, of boron. Sea water contains about 4.8 mg/l, and subsurface oilfield water contains from trace amounts to more than 100 mg/l. Fig. 5.13 is a plot of chloride versus boron concentrations of some oilfield brines taken from some sediments of Tertiary, Cretaceous, and Jurassic age. The plot

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

154

---I200

Normal evaporite curve,

/

%

T T

a 0

" I

50

J

"k/~ /

30

c

J

T

C "

T c

BORON,mg/I

Fig. 5.13.Comparison of the boron concentrations of some Tertiary (T), Cretaceous (C), and Jurassic (J) age formation waters from Louisiana with an evaporating sea water.

I

P

BORON, m g / l

Fig. 5.14. Comparison of the boron concentrations of some waters from Pennsylvanian (P) and Mississippian (M) age sediments with an evaporating sea water.

indicates that the majority of these brines are enriched in boron relative to a normal evaporite-formed brine, and that the samples that were depleted in boron may have contained dissolved halite. Fig. 5.14 is a similar plot for some samples taken from some Pennsylvanian and Mississippian age sediments. Boron is one of the elements whose concentration in the aqueous phase increases as a brine is evaporated, as illustrated in Table 5.11.

ALUMINUM

155

Aluminum Aluminum is the third most abundant element in the earth’s crust, but its concentration in natural waters usually is less than 1 mg/l. The ionic radius of trivalent aluminum is 0.57 (Goldschmidt, 1958),and it usually behaves as a cation when 6-coordinated with oxygen compounds. However, when 4-coordinated, it usually acts like the central atom of an anion. The 4-coordination usually, but not exclusively, is associated with minerals formed at high temperatures, but the 6-coordination is associated with minerals formed at low temperatures, which includes most sediments in the petroleum environment. The clay minerals illite, kaolinite, and montmorillonite often contain about 13.5, 21, and 11%aluminum, respectively. Quartzites, sandstones, limestones, and shales contain about 0.7, 3.0, 0.6, and 10% aluminum, respectively. During weathering silica will leach out and leave aluminum hydroxide behind (Pirsson and Knopf, 1947), and sedimentation processes leave only about 0.4 mg/l aluminum in sea water. According t o Hem (1970), the cation A P 3 predominates in solutions with a pH of 4.0 or less. Above pH 4.5, polymerization gives rise to an aluminum species with a gibbsite (aluminum hydroxide) structural pattern. Above pH 7.0, the dissolved form is the anion A1 (OH),-. The pH of the water is the main control of the amount of alumium that is likely to be present in natural waters. A water with a pH less than 4.0 may contain 1%or more of aluminum; for example, waters associated with acid mine drainage. Oilfield waters contain trace amounts t o more than 100 mg/l of aluminum.

a

A 1ha 1inity Alkalinity is defined as the capacity of a solution to neutralize an acid, usually t o a pH of 4.5. A solution with a neutral pH of 7.0 may have a considerable amount of alkalinity; therefore, alkalinity is a capacity function, in contrast to pH, which is an intensity function. The alkalinity-pH ranges originally coincided with methyl orange and phenolphthalein color end points. The potentiometric titration produces more accurate alkalinity results, and it utilizes an end point where the most abrupt pH change occurs while specific increments of a standard acid are added. Alkalinity usually is caused by the presence of bicarbonate, carbonate, or hydroxyl ions in a water; however, the weak acids such as silicic, phosphoric, and boric can contribute titratable alkalinity species. Carbon dioxide, which is dissolved in circulating waters as bicarbonate or carbonate as a result of the carbon cycle, is the prime source of alkalinity in shallow ground waters. However, in deep subsurface brines, additional carbon dioxide probably is dissolved as a result of diagenesis of inorganic and organic compounds. Most oilfield waters contain no hydroxyl ions, and most of them contain

156

INORGANIC CONSTITUENTS A N D PHYSICAL PROPERTIES

no carbonate ions, but they do contain bicarbonate ions. Some oilfield waters in the Rocky Mountain area are alkaline and contain both primary and secondary alkalinity, where primary alkalinity is that associated with the alkali metals and secondary alkalinity is that associated with the alkaline earth metals. For example, the Green River formation waters that are in or near trona beds may contain more than 20,000 mg/l of carbonate and 5,000 mg/l of bicarbonate. Most oilfield waters from other areas contain from about 100 to 2,000 mg/l of bicarbonate.

Acidity The basis of acidity is the solvated hydrogen ion H 3 0 + ,which is found in nature. Volcanic emanations produce HF, HCl, and H2SO4, probably formed by reactions between water and constituents associated with the magma. Waters associated with peat may contain organic acids, rain waters may contain carbonic acid, and waters associated with reducing conditions and anaerobic bacteria may contain H2 S. Acidity, as contrasted to alkalinity, is the capacity of a solution to neutralize a base, usually from below pH 4.5 t o pH 7.0. Most oilfield brines normally do not contain acidity. New wells or reworked wells often are acidified or "acidized" with a strong mineral acid or a combination of mineral and organic acids. This treatment causes the produced water to contain a certain amount of acidity until all of the acid is neutralized or diluted. Because of the large quantities of acids used in some treatments, it m a y take 6 months or more for the water produced from a treated well t o return to normal. Organic acids and organic acid salts commonly are found in oilfield waters, and the concentration ranges from trace amounts to more than 3,000 mg/l. Silica Silicon is the second most abundant element in the earth's crust, which contains about 27 wt.% of it (Fleischer, 1962). It always occurs in a combined form. Most of the silicon compounds involve structures with oxygen, and there are about a thousand silicate minerals in the earth's crust; however, those which are predominant are relatively few in number. The solubility of silica in water is a function of temperature, pressure, pH, and other ions in solution. Most silica in natural water probably is in the form of monomolecular silicic acid, H4 Si04 or Si(OH)4. Collins (196913) studied the solubility of a serpentine in solutions of calcium chloride and sodium chloride a t temperatures from 30" to 200°C and pressures from 176 t o 1,055 kg/cm2. The solubility calculated as silicon molarity in solution increased with increasing concentrations of sodium chloride, increasing pressure, and increasing temperature up to about 125°C. Between about 125" and 2OO0C, the solubility decreased with increasing temperature. The solu-

157

AMMONIUM NITROGEN

bility of silicon in the presence of calcium chloride solutions decreased with increasing Concentration of calcium chloride and with increasing temperature above about 100°C. The solubility increased with increasing pressure and with increasing temperature between 30' and about 100°C. Subsurface petroleum-associated brines usually contain less than 30 mg/l of dissolved silica; Rittenhouse et al. (1969) report that their silica content ranges from about 1 to 500 ppm as silicon, and that some low salinity waters contained a higher median content of silica than more saline waters in other areas. Ammonium nitrogen Ammonium contains nitrogen in the N-3 oxidation state, a reduced form. Nitrogen can occur in all of its states of oxidation, ranging from -3 t o +5. Oxidation of the reduced forms produces nitrogen gas, N 2 , and other nitrogen species up to nitrate, NO3-. Ammonia, NH3, forms during the anaerobic decay of organic nitrogenous material. The petroleum genetic environment produces ammonia, which transforms to ammonium, NH4, in many petroleum-associated waters because the redox potential is too low to oxidize the ammonia to nitrate. The ammonium ion is too weak t o be successfully titrated; however, Collins et al. (1969), developed a technique using formaldehyde, whereby a produced strong acid can be titrated. TABLE 5.V Ammonium content of 10 subsurface brine samples Sample

State

Formation

1 2 3 4 5 6 7 8 9 10

Utah Utah Utah Okla. Okla. Utah Utah Okla. Okla. Okla.

Navaho Green River Lower Green River Morrow Rue Uinta Surface, Green River Green River Hunton Oswego Chester

30-143 0 852-1,719 71 914-1,737 91 2,713-2,715 0 1,958-1,963 116 717-1,111 2,069 5 68 1-1,5 5 6 2,696-2,77 1 434 1,928-1,951 233 2,408-2.437 23

The NH4N content of several oilfield brines was determined and a wide variation in concentration was found. Table 5.V illustrates the amounts of NH4N found in 10 samples taken from subsurface rocks in Oklahoma and Utah.

158

INORGANIC CONSTITUENTS A N D PHYSICAL PROPERTIES

Phosphorus

Phosphorus occurs in the earth’s crust almost exclusively as the ion, and a large percentage of it is contained in the apatite group of minerals, which primarily are related to igneous rocks. The crust of the earth contains about 0.12 wt.% of phosphorus. It is a member of the V A group of elements with oxidation states ranging from -3 to +5. In contrast t o nitrogen, phosphoric acid and phosphates are not oxidizing agents. The phosphorus species present in most natural waters probably is the phosphate anion, and it usually is reported as an equivalent amount of the orthophosphate ion (PO4 )-3, the final dissociation product of phosphoric acid, H3PO4. This dissociation occurs in four steps, giving four possible phosphate forms: H3PO4, H2PO4-, HP04-2, and P 0 4 - 3 . In the alkalinity titration, any HP04-2 is converted to H2P04- and appears as bicarbonate. Shales, sandstones, and carbonates contain about 700, 170, and 400 ppm, respectively, of phosphorus. Sea water contains about 0.07 mg/l. A detailed study of the content of phosphorus in subsurface brines has not been made, but of the few that have been analyzed, most have contained less than 1 mg/l. Arsenic Arsenic is a member of the V A group of elements and probably occurs in nature mainly in the form of arsenides and sulfarsenides; it rarely occurs in its elemental form. It is comparatively rare, and the earth’s crust contains about 0.0005 wt.% of it (Fleischer, 1962). In an acidic environment, the oxidized ion, A s O ~ - ~is, mobile, and mineral arsenates tend t o be solubilized. The arsenates usually are formed in oxidation zones in contact with atmosphere and free oxygen, and arsenic will precipitate with ferric iron hydroxide. Glauconitic sediments have been found which contain up t o 70 ppm of arsenic (Goldschmidt, 1958). Subsurface oilfield brines may contain arsenic as HAs02- or H2As04, depending upon the Eh and pH. A low Eh may favor the HAs02- form. Shales, sandstones, and carbonates contain about 13, 1, and 1 ppm, respectively, of arsenic (Mason, 1966). Sea water contains about 0.003 mg/l and subsurface oilfield brines contain from 0 to 10 mg/l. Compounds containing arsenic sometimes are used in corrosion inhibitors; therefore, information concerning well treatments should be obtained before assuming that any arsenic found occurs naturally. Oxygen Oxygen is the most abundant element in the earth’s crust, which contains about 49 wt.% of it (Fleischer, 1962). I t is capable of existing in many types of combinations, and even though it is highly active, it occurs extensively in

SULFUR

159

the free form. Most combined oxygen is ionic; however, it forms a covalent molecule with hydrogen, namely water. It also forms complex oxy-salts with various metals. The oxygen content of rocks decreases with depth. The solubility of oxygen in water is primarily a function of temperature and pressure, and surface waters at ambient conditions may contain 7.63 mg/l a t 3OoC and 11.33 mg/l at 10°C (Hem, 1970). The amounts of dissolved oxygen in subsurface petroleum-associated waters is usually low, and in most in situ conditions it is undetectable because of the low redox potential of the environment. It can cause corrosion problems in the well pipes, but in most cases it is atmospheric oxygen that mixes with the produced brine during production operations that causes oxygen corrosion. Sulfur Sulfur is a member of the VI A group of elements and is widely dispersed in sedimentary and igneous rocks as metallic sulfides. The crust of the earth contains about 0.05 wt.% of sulfur (Fleischer, 1962). Free sulfur often is related to volcanic activity and can be deposited directly as a sublimate. Many commercial deposits, however, are associated with sedimentary gypsum, and probably result from biogenic activity such as that. of anaerobic bacteria. Large deposits of sulfur are found in caprocks of anhydrite overlying some salt domes. Hydrogen sulfide, often found in oilfield waters, is formed by anaerobic bacteria. One such species of bacteria is the Desulphouibrio,which obtains its oxygen from sulfate ions, causing them to be reduced to hydrogen sulfide. complexed ) with Sulfur in surface water usually occurs in the form ( S 6 oxygen as the sulfate anion S04-2. As previously mentioned, the conversion of oxidized sulfur t o a reduced form commonly involves a biogenic process, and such a reduction may not occur unless these bacteria are present. The Eh of subsurface oilfield brines usually is somewhat reducing, and the sulfur species in such environments can include hydrogen sulfide (H2 S), sulfite and thionates (S406-’). Detailed studies of the sulfur species in subsurface brines have not been made, and it is likely that other forms of sulfur are present in some brines. The temperature, pressure, Eh, pH, and other constituents in solution all influence the types of dissolved sulfur that occur in oilfield brines. Shales, sandstones, and carbonates contain about 2,400, 240, and 1,200 ppm, respectively, of sulfur (Mason, 1966). Sea water contains 900 mg/l of sulfur as sulfate, and subsurface oilfield brines contain from none up to several thousand milligrams per liter. The amount of sulfate in the brine is influenced by bacterial activity and by how much calcium, strontium, and barium is present. If these three cations are present in relatively high concentrations, the amount of sulfate present will be low. However, some brines containing high concentrations of magnesium and low concentrations of the other alkaline earth metals may contain high concentrations of sulfate.

160

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

Selenium Selenium is a member of the VI A group of elements and occurs in -2,0, +4,and +6 valence states, respectively. It is a scarce element with an abundance of about 9 x 10" wt.% of the crust of the earth (Fleischer, 1962).

Large areas of North America are underlain by seleniferous rocks and soils. These seleniferous rocks are of sedimentary origin and range in age from Late Paleozoic to Holocene. Selenium is the only known element that can be absorbed by plants in sufficient amounts to make them lethal when eaten by animals (Trelease, 1945).Fig. 5.15 illustrates the distribution of seleniferous vegetation. Sandstones, shales, and carbonates contain about 0.6, 0.05,and 0.08 ppm, respectively, of selenium. Sea water contains about 0.004 mg/l of selenium. A few subsurface oilfield brines from areas where selenium is present in soils were analyzed at this laboratory, but no selenium was detected in the brines analyzed. Most brines are present in a petroleum environment under reducing conditions, and in such an environment, selenium likely is reduced to the element and precipitated. However, in areas where outcrop water flows through petroleum-bearing formations, it is possible that selenium in the form of t)e anion Se03-2 may be present.

1 Fig. 5.15. Distribution of seleniferous vegetation in the United States.

FLUORINE

161

Fluorine Fluorine is a member of the VII A group of elements and is the most electronegative of all the elements. Its ionic radius is 1.33 A, which is about the size of OH- and O-*;therefore, it enters a variety of minerals. The earth’s crust contains about 0.03 wt.% of fluorine (Fleischer, 1962). In solutions, fluorine usually forms the fluoride F- ion; at a low pH, the HFo form might occur. It also can form strong complexes with aluminum, beryllium, and ferric iron. Fluorine occurs in several minerals, but the only common industrial source is fluorspar (CaF2). It occurs as HF or SiF in volcanic emanations, and even as the free element in (stinkfluss) “stinking fluorspar” of Wolssendorf, Bavaria. The solubility of calcium fluoride (fluorite) in water at 25OC is about 8.7 ppm of fluoride (Aumeras, 1927);this solubility could be affected by other dissolved constituents. Sodium fluoride is very soluble, and magnesium fluoride is more soluble than calcium fluoride; therefore, a petroleumassociated water that is deficient in calcium and has been in contact with rocks containing fluoride minerals will contain appreciable quantities of fluoride. Shales, sandstones, and carbonates contain about 740, 270 and 330 ppm of fluorine, respectively (Mason, 1966). Sea water contains about 1.3 mg/l, and natural waters with a dissolved solids concentration of less than 1,000 mg/l usually contain less than 1 mg/l of fluoride. However, concentrations up t o 50 mg/l have been reported (Hem, 1970). Not many subsurface petroleum-associated brines have been analyzed for fluoride, but a few are known t o contain up t o 5 mg/l. Chlorine Chlorine is a member of the VII A group of elements and is the most important member of the group with respect to water. The crust of the earth contains about 0.19 wt.% of chlorine (Fleischer, 1962);some estimates place the fluorine abundance above the chlorine abundance. Volcanic activity produces the gas hydrogen chloride and sometimes chlorine, but much less frequently. The caliche evaporite deposits in Chile contain the perchlorate ion C104-; however, the mechanism by which it formed is not clear. Several minerals contain the chloride ion. The chloride ion does not form low-solubility salts. I t is not easily adsorbed on clays or other mineral surfaces. It is not significant in oxidation and reduction reactions, and it forms no important solute complexes. Chloride is very mobile in the hydrosphere, yet it is relatively scarce in the earth’s crust. It is the predominant anion in sea water and in most petroleum-associated waters. It is found in all natural waters, and its average concentration in rainwater is about 3 mg/l (Hem, 1970). Chloride salts are

162

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

very soluble; therefore, chloride is usually not removed from solution except during freezing or evaporation processes and in hyperfiltration, as water moves through some types of clay beds (White, 1965). Shales, sandstones, and carbonates contain about 180, 10, and 150 ppm, respectively, of chloride. Sea water contains about 19,000 mg/l of chloride, the principal anion in the sea. The chloride content of the hydrosphere is much greater than can be accounted for by weathering of rocks, and it has been postulated that the primordial atmosphere may have been rich in chlorine compounds. The volcanic emission of chlorine gases appears a more plausible explanation, however. Oilfield brines usually contain relatively high concentrations of chloride; in some brines the concentration may be 200,000 mg/l or more. Chloride usually is the predominant anion in oilfield brines. Table 5.1 illustrates how its concentration can increase in an evaporite-associated brine. Evaporation probably is the only geochemical process which appreciably affects the chloride content of the seas. Bromine Bromine is a member of the VII A group of elements and it behaves somewhat similarly t o chlorine. The crust of the earth contains about 0.0005 wt.% of bromine (Fleischer, 1962). It usually occurs as the ion bromide Br-, and it does not form its own minerals when sea water evaporates (Valyashko, 1956). It forms an isomorphous admixture with chloride in the solid phases. The order of crystallization (see Table 5.1) is halite (NaCl), sylvite (KCl), carnallite (MgC12-KC1*6H20), and/or kainite (MgS04 *KC1=3H2 0), and at the eutectic point, bischofite (MgCl? -6H2 0). Each of these chlorides entrains bromide in the solid phase. This distribution accounts for the relative enrichment of bromide in the liquid phase because with each crystallization more bromide is left in solution than is entrained in the solid phase. Mun and Bazilevich (1962) reported that, in fresh-water lakes, bromide accumulates in the muds, that its concentration is proportional t o the organic-matter concentration in the sediments, and that it is not influenced by the pelitic fraction. In muds of salt lakes, the higher the bromide concentrations in the brine, the higher it is in the muds. In general, the bromide content in the pore solutions increased with depth, but the bromide content in muds decreased with depth, owing to more complete decomposition of organic bromine compounds. Bromide is two t o three times more concentrated in carnallite than in sylvite and five to ten times more than in halite (Myagkov and Burmistrov, 1964). Apparently, concentration and dilution are responsible for the complex distribution of bromide in rocks of a carnallite zone. The determining factor in the replacement of chloride by bromide is the mineral composition rather than the bromide concentration in the brine.

BROMINE

163

Braitsch and Herrmann (1963) found that the absolute bromide content of rocks can be used to determine primary and secondary paragenesis. Distribution of bromide between solution and crystals of halite, sylvite, carnallite, and bischofite, and the effects of other ions plus temperatures between 25" and 83OC, confirm this. This method was also applied to determine the temperature of primary potash deposits. An investigation of the bromide/ sodium chloride relation in salt deposits revealed that bromide can be used to determine the stratigraphy of evaporite-salt deposits (Baar, 1963). Derivation of theoretical profiles of bromide thickness versus salt thickness indicated that, with constant inflow, evaporation, and reflux, the thickness profiles were all monotomic logarithmic functions. The irregular and high bromide concentrations of some salt deposits were attributed to inflow of bromide-rich bitterns from an adjacent potash basin (Holser, 1966). Shales, sandstones, and carbonates contain about 4,1,and 6 ppm, respectively, of bromide (Mason, 1966). Sea water contains about 65 mg/l of bromide, and subsurface petroleum-associated brines contain from less than 50 to more than 6,000 mg/l of bromide. Fig. 5.16 illustrates the bromide concentration plotted versus the chloride concentration for some subsurface brines taken from Tertiary, Cretaceous, and Jurassic age sediments. This plot indicates that the waters from these Tertiary age sediments are depleted in bromide relative to a normal evaporite brine, whereas those from the Cretaceous and Jurassic age sediments are enriched'in bromide.

C

BROMIDE, mg / I

Fig. 5.16. Comparison of the bromide concentrations in some formation waters from Tertiary (T), Cretaceous (C), and Jurassic (J) age sediments from Louisiana with an evaporating sea water.

INORGANIC CONSTITUENTS A N D PHYSICAL PROPERTIES

164

Normal evaporite curve

P PI

M

300

1,000 BROMIDE, m g l l

3000

I

10

Fig. 5.17. Comparison of the bromide concentrations in some formation waters from Pennsylvanian (P) and Mississippian (M) age sediments from Oklahoma with an evaporating sea water.

Fig. 5.17 is a similar plot for some brines taken from some Pennsylvanian and Mississippian age sediments. The bromide concentrations in these brines do not appear t o be significantly different. Brines containing 1,500 to 8,000 mg/l of bromide, with calcium and magnesium chloride as the major constituents, are formed by evaporation of sea water and associated sedimentation rather than by dissolution of salts. Increase in temperature causes a phase shift in the solid and brine phases, resulting in an increase of bromide in solution. Iodine Iodine is a member of the VII A group of elements, and of the four members discussed in this chapter, it is the least abundant, since it comprises only about 3 x lo-' wt.% of the earth's crust (Fleischer, 1962). I t forms three minerals of its own; namely, iodoargyrite (AgI), iodoembolite [Ag(Cl,Br,I)], and miersite [(Ag,Cu)I]. Marine plants, such as kelp and plankton algae, concentrate iodine. The distribution of iodide in marine and oceanic silts and interstitial waters indicates that near-shore ocean Sediments contain more iodide than deep-sea sediments. Red clays and calcareous sediments contain less iodide than organic-bearing argillaceous sediments. The iodide concentration in the marine and oceanic sediments decreases with depth, but the iodide concen-

IODINE

165

tration in the interstitial waters increases with depth (Shishkina and Pavolva, 1965). The iodide in bottom water layers and in the interstitial water of muds in some Japanese lakes was found t o be selectively captured by flocculated iron, manganese, and aluminum hydroxides which sank to anaerobic layers (Sugawara et al., 1956). Reduction of the hydroxides releases iodide to the bottom waters. However, the release of iodide is incomplete, and the flocculates reach the bottom muds where the Eh is even more negative, resulting in high accumulation of iodide in interstitial water of muds. The primary source of organic matter in marine and oceanic basins is photosynthesis by plankton algae. Algae are directly or indirectly the food resource of all the remaining life in the basins, and the proliferation rate differential and the types of feeding organisms influence the sediment deposition rate as well as the amount of iodide and bromide in the sediment (Bordovskii, 1965). Shales, sandstones, and carbonates contain about 2.2, 1.7, and 1.2 ppm, respectively, of iodide (Mason, 1966). Sea water contains about 0.05 mg/l, and most subsurface petroleum-associated brines contain less than 10 mg/l; however, some have been found to contain up t o 1,400 mg/l. Fig. 5.18 is a plot of the chloride concentrations versus the iodide concentrations for some brines taken from some Pennsylvanian and Mississippian age sediments. Iodide is tremendously enriched in all of these brines compared to the normal evaporite-associated brine. Some mechanisms such as leaching or solubilization of iodine, iodate, or iodide compounds, ion filtration, anion exchange, and desorption had t o occur, t o account for this enrichment of iodide in the aqueous phase. A similar plot for some waters taken from Tertiary, Cretaceous, and Jurassic age sediments gave similar results except that these particular brines were not as heavily enriched in iodide. The iodide concentration of some subsurface waters is dependent on the proximity of argillaceous deposits containing organic matter, rather than on dissolved mineralization. Gas may play an important part in the accumulation of iodide in subsurface waters. Some gas structures are bounded by iodide-rich waters, and the iodide content is depleted at a distance from the gas structure (Ovchinnikov, 1960). Studies of some reservoirs, Holocene to Miocene in age, in lagoonal sedimentary basins of thick sediments with wide areal extent, indicate that a genetic relation exists between iodide in the formation waters and the accompanying natural gas (Marsden and Kawai, 1965). Possibly the high concentrations of iodide are the result of concentration by algae and other marine organisms from ancient sea waters; their remains became part of the sediments, and later the iodide was solubilized. However, because the iodide usually is strongly incorporated in the sediment, such sediments must contain large quantities of iodide, and other mechanisms must operate to solub i k e the iodide in associated waters.

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

166 1

,

1,200

-

I#OOO

cn

500

2 sI

200

\

w'

O

5

/1

~ w mevaporite i

/

loot

50E t

20

CUM

/

/

0

3

'

, ,

/-

0

/

M

vs

P

P

MM M M F P

M MhP

IODIDE, mg/l Fig. 5.18. Comparison of the iodide concentrations of some formation waters from Pennsylvanian (P) and Mississippian (M) age sediments from Oklahoma with an evaporating sea water.

Theoretically, only iodate is thermodynamically stable in sea water (SillCn, 1961). The exact form of iodine in oilfield brines has not been investigated. These forms probably will vary with the salinity, Eh, and other factors. Sugawara and Terada (1957)established that both iodide and iodate are present in comparable amounts in sea water. Biologists found that iodine-concentrating algae ultilize only the iodate form (Shaw, 1962). Significance of some physical properties

Redox potential The redox potential often is abbreviated as Eh, and may also be referred to as oxidation potential, oxidation-reduction potential, or pE. It is expressed in volts, and at equilibrium it is related to the proportions of oxidized and reduced species present. Standard equations of chemical thermodynamics express the relationships. Eo is the standard potential of a redox system when unit activities of participating substances are present under standard conditions. Eo is related to standard free energy change in a reaction by the equation:

where n is the number of unit negative charges (electrons) shown in the

PHYSICAL PROPERTIES

167

redox reaction and f is the Faraday constant in units that give a potential in volts (94,484 absolute coulombs). Standard free energy values are given in texts such as that of Latimer (1952). When the system is not under standard conditions, the redox potential is expressed by the Nernst equation: Eh = E o

T (oxidized species) +R log nf (reduced species)

where R is the gas content (1.987 cal. degree mole), and T is the temperature in degrees Kelvin. Geochemical literature and biochemical literature, such as that of Pourbaix (1950), present increasing positive potential values to represent increasing oxidizing systems and decreasing potential values to represent reducing systems. The sign of Eh used in this manner is opposite to standard American practice in electrochemistry. Zobell (1946) established basic procedures for measuring the Eh of geologic-related materials. The Zobell solution containing 0.003M potassium ferrocyanide and 0.003M potassium ferricyanide in a 0.1M potassium chloride solution has an Eh of 0.428 V at 25OC. Minor temperature variations can be calculated using the equation: Eh = 0.428+).0022 ( t - 25) where t = temperature of the sample in degrees Celsius. Garrels and Christ (1965) describe procedures for determining Eh equilibria of mineral substances. Particularly useful are the procedures described for constructing diagrams showing fields of stability for various mineral substances as functions of pH and Eh. Fig. 5.19 is an Eh/pH diagram. Such stability field diagrams might be constructed for the substances comprising petroleum and should be of considerable help in understanding the mechanisms of origin, accumulation, and chemical stability of petroleum. Unfortunately, this approach does not yield simple results because most oxidation reactions involving hydrocarbons and other petroleum constituents are not reversible in the usual sense. Furthermore, thermodynamic data are available for only a small fraction of the large number of reactions and products that are possible. Attempts t o obtain useful results from Eh measurements in natural media involve numerous difficulties. In a natural medium, such as petroleumassociated water, there are many variables, none of which is controlled, which individually or collectively may have little or great influence on Eh measurements made on the water. Many chemical substances, such as ferric or ferrous ions, various organic oxidation-reduction systems, sulfides, and sulfates, may be present in the water in large or small amounts. Even controlled systems in the laboratory often produce unaccounted-for variances. In the field, the lack of knowledge of actual participating species may seriously impair proper interpretation of Eh readings. Eh measurements made

168

INORGANIC CONSTITUENTS A N D PHYSICAL PROPERTIES

1,200 -

H2O

000 600

<

I

I

400-

z

W

1

200

-

Acid Fa++

/Modern

I

I I

sea water

Alkaline

on poorly poised media (media with poor Eh stability), such as some oilfield waters, involve additional uncertainties. Response of electrodes in such solutions is sluggish, and electrodes are easily contaminated with trace amounts of substances which will produce invalid readings. In the natural environment, reactions occur that involve protons and electrons, such as: FeS04 + 2H20

* S04-2

+ FeO-OH + 3H+ + e-

Such reactions depend upon both the pH and the Eh of the system, and the equilibrium line of such reactions is Eh = E o - 59 a/n pH mV, where a is the number of protons. Knowledge of the redox potential is useful in studies of how compounds such as uranium (Naumor, 1959), iron, sulfur (Hem, 1960), and other minerals (Cloke, 1966; Pirson, 1968) are transported in aqueous systems. The solubility of some elements and compounds is dependent upon the redox potential and the pH of their environment. The Eh/pH diagram shown in

PHYSICAL PROPERTIES

169

Fig. 5.19 can be used to predict that ferrous ions are the more common form of dissolved iron and that ferric ions will precipitate in an oxidizing environment if the pH is above 1. Similar diagrams can be drawn for other constituents. Some water associated with petroleum is “connate” water, and Fig. 5.19 indicates that such water has a negative Eh; this has been proven in various field studies (Buckley et al., 1958). The Eh of some petroleum-associated waters in the Anadarko Basin ranged from -270 mV to -300 mV (Collins, 1969a). Knowledge of the Eh is useful in determining how t o treat a water before it is injected into a subsurface formation (Ostroff, 1965). For example, the Eh of the water will be oxidizing 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 then reinjected. In such a situation, the Eh value is useful in determining how much iron will stay in solution and not deposit in the well bore. Organisms that consume oxygen cause a lowering of the redox potential. In buried sediments, it is the aerobic bacteria that attract organic constituents which remove the free oxygen from the interstitial water. Sediments laid down in a shoreline environment will differ in degree of oxidation as compared t o those laid down in a deep-sea environment (Pirson, 1968). For example, the Eh of the shoreline sediments may range from -50 t o 0 mV, but the Eh of deep-sea sediments may range from -150 to -100 mV. The aerobic bacteria die when the free oxygen is totally consumed; the anaerobic bacteria attack the sulfate ion, which is the second most important anion in the sea water. During this attack, the sulfate reduces t o sulfite and then t o sulfide; the Eh drops to -600 mV; H2S is liberated, and CaC03 precipitates as the pH rises above 8.5 (Dapples, 1959).

The term pH means the logarithm (base 1 0 ) of the reciprocal of the hydrogen-ion concentration, and the pH of pure water at 25OC is 7.0, which means that there is lo-’ ‘mole per liter of H+ in solution. When other constituents are solubilized by water, the pH probably will change because the chemical equilibrium shifts as new ions combine with H+ or OH-. The presence of slightly dissociated acids or bases will tend to buffer the solution, and the addition of H+ or OH- will shift the pH only a small amount until the acids or bases are changed to salts. The pH of oilfield waters usually is controlled by the carbon dioxidebicarbonate system. Because the solubility of carbon dioxide 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 used for water identification or correlation purposes, but it will indicate possible scale-forming or corrosion tendencies of a water. The

170

INORGANIC CONSTITUENTS A N D PHYSICAL PROPERTIES

pH also may indicate the presence of drilling-mud filtrate or well-treatment chemicals. A detailed study indicates that virtually no environment exists on or near the earth’s surface where the pH/Eh conditions are incompatible with organic life (Baas Becking et al., 1960). Because COz is the main byproduct of organic oxidation and the building material of plant and much bacterial life, it must be expected t o play a dominant role. It dissolves in HzO, producing the bicarbonate ion and a free hydrogen ion. The concentration of the hydrogen ion is 1 x moles per liter (pH 7) at 25OC in pure water, but when saturated with COz, it rises t o 1 x lo-’ moles per liter (pH 5). The equilibrium conditions of carbon dioxide, carbonic acid, and the bicarbonate ion are: Hz 0 + COz

* HzCO, * HC03-

+ H+

* 2H+ + C03-’

and the pH of each equilibrium in ocean water is pH 5, pH 6.3, and pH 10.3. This reaction moves to the right with increasing temperature in a closed system. In the presence of organic constituents, the equilibria are modified, and the pH range can extend from 2 to 12.

Fig. 5.20. Changes in pH as a result of the addition of carbonate ions to distilled water and water solutions containing sodium and chloride ions.

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. Addition of the carbonate ion to sodium chloride solutions will raise the pH, as illustrated in Fig. 5.20. If calcium were present, calcium carbonate would precipitate. The reason why 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.

PHYSICAL PROPERTIES

171

Ionic radii Table 5.VI contains data concerning the radii of the nonhydrated ions, the hydrated ions, the ionic potential, and the polarization. The size of the ions is of interest concerning the mobilities or the relative transport coefficient of a given ion through a clayshale membrane system or the replacement coefficient in a clay ion-exchange system. The ionic potential is of interest because elements with low ionic potentials are the most likely to remain in true solution. The polarization, which is equivalent to the valency divided by the ion radius, is of interest because the larger the polarization, the lower the replacing power in an exchange system (Collins, 1970). The ionic potentials of the constituents involved in the diagenesis are important (Hem, 1960). Those that stay in true ionic solution to rather high pH levels include Na+, K+, Mg+’, Fe+’, Mn+’, Ca+’, Sr+*,and Ba+’ ;they are the soluble cations, and their ionic potentials range from 0.3 to 1.3, where the ionic potential is the ratio between the ionic charge and the ionic radius. Constituents that are precipitated by hydrolysis are those with ionic potentials of 3-12 and include such ions as A P 3 , Fe+3, S P 4 , and M r P 4 . Constituents which form soluble complex ions and usually go into true ionic solution include B+3, C 4 ,N + 5 , P+’, S 6 ,and Mn+’ ;their ionic potentials are over 12. In general, the hydroxides of the soluble cations possess ionic bonds; therefore, they are soluble. The hydrolysates, or those ions precipitated by hydrolysis from hydroxyl bonds, and the soluble complex ions both have hydrogen bonds.

TABLE 5.VI Radii, valence, ionic potentials, and polarization Constituents

Nonhydrated radius (A)

Valence

Hydrated Ionic radius (A) potential

Polarization

Lithium Sodium Potassium Calcium Magnesium Strontium Barium Boron Chloride Bromide Iodide Sulfate

0.60 0.95 1.33 0.99 0.65 1.13 1.35 0.23 1.81 1.95 2.16 2.90

+1 +1 +1 +2 +2 +2 +2 +3 -1 -1 -1 -2

3.82 3.58 3.31 4.12 4.28 4.12 4.04

1.67 1.05 .75 2.02 3.08 1.77 1.48

-

3.32 3.30 3.31 3.79

0.60 0.95 1.33 0.50 0.33 0.57 0.68 0.08 1.81 1.95 2.16 1.45

-

0.55 0.51 0.46 0.69

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

172

Density Equations that were developed for sea water can be applied to oilfield waters to obtain approximate values for engineering studies. The density values (ao) at O°C and atmospheric pressure are related to the chlorinity (CZ) as follows:

-0.069 + 1.4708 CZ - 0.00157 C12 + 0.0000398 C13 = UO where CZ = chlorinity (see Table 3.111). The density is very dependent upon temperature:

where D = a complex function of ao, and temperature and D values can be obtained from Knudsen’s Hydrographic Tables (Knudsen, 1901).

Vapor pressure The relative lowering of the vapor pressure of oilfield water can be calculated with the following equation: Ap/po = 0.538 x

S

where p o = the vapor pressure of distilled water at the same temperature, and S = the salinity (see Table 3.111) (Kellog and Company, 1956, 1966,

1968).

Boiling point A first approximation of the boiling point elevation can be calculated from:

At = 0.0158S where S = the salinity. Freezing point is :

An empirical equation which can be used t o estimate the freezing points

t

= 4.0086 - 0.064633 ((TO)

- 0.0001055 ( 0 0 ) ~

See “Density” for an explanation of terms.

PHYSICAL PROPERTIES

173

Viscosity The viscosity will increase with decreasing temperature and with increasing salinity. The viscosities of sodium chloride solutions of the same ionic strength can be used to estimate oilfield water vicosities. Osmotic pressure

A relationship between osmotic pressure ( P o ) and the depression of the freezing point at 0°C is (in atmospheres):

The osmotic pressure at other temperatures can be estimated (Kellog and Company, 1956,1966,1968):

Po

x

(1+ 0.00367t)

Specific heat The values for the specific heat, c p , of oilfield waters can be approximated from the values of an equivalent sodium chloride solution. Thermal conductivity The thermal conductivity coefficient, A, can be calculated from thermal capacities because the ratio of thermal conductivities of two materials is the same as that of the thermal capacities of equal volumes. The values for X at various temperatures are available in a “Saline Water Conversion Technical Data Book” (Kellog and Company, 1956,1966,1968).

Surface tension The surface tension of an oilfield water increases with decreasing temperature and with increasing salinity. An empirical formula which can be used t o calculate it is:

75.64 - 0.144t + 0.0399 Cl = surface tension (dynes/cm*) where t = temperature in Celcius, and CZ = the chlorinity (see Table 3.111; Kellog and Company, 1956,1966,1968).

174

INORGANIC CONSTITUENTS AND PHYSICAL PROPERTIES

References Ahrens, L.H., 1965. Distribution o f the Elements in Our Planet. McGraw-Hill, New York, N.Y., 110 pp. Aumeras, M., 1927. Studies of ionic equilibria, 11. Equilibrium of calcium fluoride in dilute hydrochloric. J. Chem. Phys., 24:548-571. Baar, C.A., 1963. How to use the bromine test to determine the stratigraphical position in rock salt series. Neues Jahrb. Mineral. Geol. Palaontol., Monatsh., 7( 1):145-153. Baas Becking, L.G.M., Kaplan, I.R. and Moore, D., 1960. Limits of the natural environment in terms of pH and oxidation-reduction potentials. J. Geol., 68:243-284. Bordovskii, O.K., 1965. Source of organic matter in marine basins. Mar. Geol., 3( 1/2):5-31. Braitsch, 0. and Herrmann, A.G., 1963. Zur Geochemie des Broms in Salinaren Sedimenten, Teil I. Experimentelle Bestimmung der Br-Verteilung in verschiedenen naturlichen Salzsystemen. Geochim. Cosmochim. Acta, 27 :361-391. Brasted, R.C., 1957. The halogens. In: M.C. Sneed, J.L. Lewis and R.C. Brasted (Editors), Comprehensive Inorganic Chemistry, 3. Van Nostrand, New York, N.Y., 250 pp. Buckley, S.E., Hocott, C.R. and Taggart, Jr., M.S., 1958. Distribution of dissolved hydrocarbons in subsurface waters. In: L.G. Weeks (Editor), Habitat of Oil. American Association of Petroleum Geologists, Tulsa, Okla., pp.850-882. Burst, J.F., 1969. Diagenesis of Gulf Coast clayey sediments and its possible relation to petroleum migration. Bull. A m . Assoc. Pet. Geol., 53:73-93. Cloke, P.L., 1966. The geochemical application of E h - p H diagrams. J. Geolog. Ed., 14:140-1 48. Collins, A.G., 1969a. Chemistry of some Anadarko Basin brines containing high concentrations of iodide. Chem. Geol., 4:169-187. Collins, A.G., 1969b. Solubilities of some silicate minerals in saline waters. U.S. Off. Saline Water Res. Dev. Progr. Rep., No. 472, 27 pp. Collins, A.G., 1970. Geochemistry of some petroleum-associated waters from Louisiana. US.Bur. Min. Rep. Invest., No.7326, 31 pp. Collins, A.G., Castagno, J.L. and Marcy, V.M., 1969. Potentiometric determination of ammonium nitrogen in oilfield brines. Environ. Sci. Technol., 3:274-275. Dall’Aglio, M., 1968. The abundance of mercury in 300 natural water samples from Tuscany and Latium (Central Italy). In: L.H. Ahrens (Editor), Origin and Distribution o f the Elements. Pergamon Press, Oxford, p.1065-1081. Dapples, E.C., 1959. The behavior of silica in diagenesis. In: H.A. Ireland (Editor), Silica in Sediments - SOC.Econ. Paleontol. Mineral., Spec. Publ., No.7, pp.36-55. Davis, J.W. and Collins, A.G., 1971. Solubility of barium and strontium sulfates in strong electrolyte solutions. Environ. Sci. Technol., 5:1039-1043. Fleischer, M., 1962. Recent estimates of the abundances of the elements in the earth’s crust. U.S. Geol. Surv. Circ., No.285, 9 pp. Garrels, R.M. and Christ, C.L., 1965. Solutions, Minerals, and Equilibria. Harper and Row, New York, N.Y., 450 pp. Goldberg, E.D., 1963. The oceans as a chemical system. In: M.N. Hill (Editor), Composition of Sea Water: Compamtive and Descriptive Oceanography. Interscience, New York, N.Y., 2:3-25. Goldschmidt, V.M., 1958. Geochemistry. Oxford University Press, London, 730 pp. Hem, J.D., 1960. Some chemical relationships among sulfur species and dissolved iron : chemistry of iron in natural water. U.S.Geol. Surv., Water Supply Paper, No, 145942, pp.57-73. Hem, J.D., 1970. Study and interpretation of the chemical characteristics of natural waters. U.S. Geol. Suru., Water Supply Paper, No. 1473, 363 pp. Hemley, J.J., 1953. A study of lead sulfide solubility and its relation to ore deposition. Econ. Geol., 48:113-137.

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Holser, W.T., 1966. Bromide geochemistry of salt rocks. In: Jon L. Rau (Editor), 2nd Symposium on Salt. The Northern Ohio Geological Society, Cleveland, Ohio, 1:248-275. Karasik, M.A., Goncharov, Yu. and Vasilevskaya, A.Ye., 1965. Mercury in waters and brines of the Permian salt deposits of Donbas. Geochem. Int., 2:82-87. Kazmina, T.I., 1951. The boron-chloride coefficient in oilfield waters. Dokl. Akad. Nauk S.S.S.R . , 77 :301. Kellog, M.W. and Company, 1956, 1966, 1968. Saline Water Conversion Technical Data Book and Supplements 1 and 2. (Catalog No. 1, 1.7712: Sa3, Supp.1, Supp.2, Supt. of Documents), U.S. Government Printing Office, Washington, D.C. Khitarov, N.I. and Pugin, V.A., 1966. Behavior of montmorillonite under elevated temperatures and pressures. Geochem. Int., 3:621-626. Knudsen, M., 1901. Hydrographic Tables. G.E.C. Gadd, Copenhagen, 6 3 pp. Larson, T.E. and King, R.M., 1954. Corrosion by water at low flow velocity. J. A m . Water Works Assoc., 46 :1-9. Latimer, W.M., 1952. Oxidation Potentials. Prentice-Hall, New York, N.Y., 2nd ed., 352 PP. Lyon, T.L. and Buckman, H.D., 1960. The Nature and Properties o f Soils. Macmillan, C., New York, N.Y., 567 pp. Marsden, S.S. and Kawai, K., 1965. “SuiyBsei-Ten’nengasu” - a special type of Japanese natural gas deposit. Bull. Am. Assoc. Petrol. Geol., 49:286-295. Mason, B., 1966. Principles of Geochemistry. John Wiley and Sons, New York, N.Y., 329 PP. Miholic, S., 1947. Ore deposits and geologic age. Econ. Geol., 42:713. Mitgarts, B.B., 1956. The composition of subterraneous waters in its value for petroleum prospecting, according to the data from the Fergana region. Vses. Nauch. Issled., Geol. Inst., Vopr. Neftepoiskovi Gidrogeol., 18:58-93 (in Russian). Moeller, T., 1954. Inorganic Chemistry. John Wiley and Sons, New York, N.Y., 966 pp. Mun, A.I. and Bazilevich, Z.A., 1962. Distribution of bromine in lacustrine bottom muds. Geochemistry, 2:199-205. Myagkov, V.F. and Burmistrov, D.F., 1964. Distribution of bromine in the carnallite beds of the Verkhnekamsk deposit. Geochem. Int., 1:701-703. Naumor, G.B., 1959. Transportation of uranium in hydrothermal solution as a carbonate. Geochemistry, 1:5-20. Odum, H.T., 1951. Notes on the strontium content of sea water, celestite radiolaria, and strontianite snail shells. Science, 114:211. Ostroff, A.G., 1965. Introduction t o Oilfield Technology. Prentice-Hall, Englewood Cliffs, N.J., 412 pp. Ovchinnikov, N.V., 1960. Patterns in the alteration of the chemical composition of subsurface waters of the Azqv-Kuban Trough and the distribution of iodine and bromine therein. Izv. Vyssh. Uchebn., Zaved., Geol. Razvedka, 3:134-138. Pirson, S.J., 1968. Redox log interprets reservoir potential. Oil Gas J., 66:69-75. Pirsson, L.V. and Knopf, A., 1947. Rocks and Rock Minerals. John Wiley and Sons, New York, N.Y. 3rd ed., 179 pp. Pourbaix, M.J.N., 1950. Thermodynamics o f Dilute Aqueous Solutions. Longmans, Green and Co., New York, N.Y., 151 pp. Rittenhouse, G., Fulton, Jr., R.B., Grabowski, R.J. and Bernard, J.L., 1969. Minor elements in oilfield waters. Chem. Geol., 4:189-209. Shaw, T.I., 1962. Physiology and biochemistry of algae. In: R.A. Lewin (Editor), Halogens. Academic Press, New York, N.Y., pp. 247-253. Shishkina, O.V. and Pavolva, G.A., 1965. Iodine distribution in marine and oceanic muds and in their pore fluids. Geochem. Int., 2:559-565. Sillbn, L.G., 1961. Physical chemistry of sea waters. A m . Assoc. Adv. Sci. Publ., No.67, pp. 549-581.

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Sill&, L.G. and Martell, A.E., 1964. Stability constants of metal-ion complexes. Chem. Soc. Lond., Spec. Publ., No.17, 754 pp. Sugawara, K. and Terada, K., 1957. Iodine distribution in the western Pacific Ocean. J. Earth Sci., 5:81-102. Sugawara, K., Koyama, T. and Terada, K., 1956. Co-precipitation of iodide ions by some metallic oxides with special reference to iodide accumulation in bottom water layers and interstitial waters of muds in some Japanese lakes. UNESCO/NSRIC/65,presented at Int. Congr. Theor. Appl. Limnol., Helsinki, Finland, 1956, 11 pp. Trelease, S.F., 1945. Selenium in soils, plants, and animals. Soil Sci., 160:125-131. Valyashko, M.G., 1956. Geochemistry of bromine in processes of halogenesis and use of bromine content as a genetic and prospecting criterion. Geochemistry, 6:33-49. White, D.E., 1957. Magmatic, connate and metamorphic waters. Bull., Geol. Soc. A m . , 68 :1669. White, D.E., 1965. Saline waters of sedimentary rocks. In: A. Young and J.E. Galley (Editors), Fluids in Subsurface Environments - A m . Assoc. Pet. Geol., Mem. 4 , pp.342-366. White, D.E., Hem, J.D. and Waring, G.A., 1963. Data of geochemistry. U.S. Geol. Surv. Prof. Paper, No.440-F, pp. FI-F67. Zobell, C.E., 1946. Studies on the redox potential of marine sediments. Bull. A m . Assoc. Pet. Geol., 30:477-513.