FORM ATE
TE CH N I CA L
M A N UA L
C A B O T
CHEMICAL AND PHYSICAL PROPERTIES
SECTION A6
PH AND BUFFERING
A6.1
pH of formate brines ....................................................................................................................2
A6.1.1 Controlling pH in formate brines .................................................................................3
A6.1.2 Measuring pH in formate brines .................................................................................3
A6.2
pH buffering of formate brines with carbonate / bicarbonate buffer ........................ 5
A6.2.1 How the carbonate / bicarbonate buffer works ...................................................5
A6.2.2 Buffer protection against CO2 influx ..........................................................................5
A6.2.3 Buffer protection against H2S influx .........................................................................6
A6.3 Buffer addition and maintainance ..........................................................................................7
A6.3.1 Buffer capacity .................................................................................................................7
A6.3.2 Total buffer concentration ............................................................................................8
A6.3.3 Determining buffer concentration and capacity .........................................................8
A6.3.4 Buffer requirement for field use .................................................................................9
A6.3.5 Maintaining buffer concentration and capacity ..........................................................11
References ...................................................................................................................................................11
The Formate Technical Manual is continually updated.
To check if a newer version of this section exists please visit
formatebrines.com/manual
NOTICE AND DISCLAIMER. The data and conclusions contained herein are based on work believed to be reliable; however, CABOT cannot and does not guarantee that similar results
and/or conclusions will be obtained by others. This information is provided as a convenience and for informational purposes only. No guarantee or warranty as to this information, or
any product to which it relates, is given or implied. CABOT DISCLAIMS ALL WARRANTIES EXPRESS OR IMPLIED, INCLUDING MERCHANTABILITY OR FITNESS FOR A PARTICULAR PURPOSE AS TO
(i) SUCH INFORMATION, (ii) ANY PRODUCT OR (iii) INTELLECTUAL PROPERTY INFRINGEMENT. In no event is CABOT responsible for, and CABOT does not accept and hereby disclaims liability for,
any damages whatsoever in connection with the use of or reliance on this information or any product to which it relates.
© 2013 Cabot Corporation, MA, USA. All rights reserved. CABOT is a registered trademark of Cabot Corporation.
VERSION 4 – 09/13
FORMATE TECHNICAL MANUAL
C AB O T
A6.1 pH of formate brines
pH is a measure of the acidity or alkalinity of a
solution, numerically equal to 7 for neutral solutions,
increasing with increasing alkalinity and decreasing
with increasing acidity. For dilute solutions, pH can
be defined as the negative logarithm base 10 of the
hydrogen concentration in the solution [H +]:
The pH of formate brine can be decreased to 3.75 by
adding a strong acid, but the brine will resist further pH
change until all the formate ions have been converted
to formic acid ions.
At a pH of 3.75, the formic acid and formate anions
will exist in a 1:1 molar ratio. When the pH of a formate
brine is raised or lowered one unit from this value the
ratio of formate to formic acid will change by a factor of
pH = − log [H +] (1)
1
approximately ten, as shown in Table 1. This means that
2
aH + the behavior of the ions
pHmore
= − log
(aH + )
In
concentrated
solutions,
in concentrated cesium formate brine with a formate
in the −solution+ depends
not on their concentrations, but
concentration of around 10 mol/L and a pH of around
Ka
pKa
3
HCOO + H 3 O+Thus,
←in
→reality,
HCOOHa +more
H2 O precise definition
on
is:
10 – 10.5, the concentration of formic acid is less than
pH activities.
= − log [H ]
1
0.000001 mol/L. Figure 1 shows how pH of unbuffered
log [H +]
2
aH +
2 − formate brine
pH 2=− − log+ (aHpK
−
−
(2)
+ )a CO3
pKa
HCO 3 changes with addition of a strong acid.
CO3 + H ←→ HCO 3
4
a
where H + is the activity.
log (aH + )
Ka
+
pK
pKa
3
−
HCOO− − + H
←
a → HCOOH + H2 O
+ 3O
formate / formic acid molar ratio
pKa
HCO3 Table 1H‘Theoretical’
5
2 CO 3
K a HCO3 + H ←→ H2 CO3
+
pK
+ H 3 O ←→ Commonly
HCOOH + H2used
O
high-density
oilfield brines (CaCl2,
as a function of pH.
a
CaBr
pK ) have a− naturally acidic pH. Attempts
2−
−
2 − , and ZnBr
CO3
pKa
HCO
CO3 2+ H +←a 2→ HCO 3
4
pH 3
Approx. formate / formic acid molar ratio
6
CO−2raise
(g )←
→
CO2to(aq
)
to
the pH
alkaline
levels in these2 −halide-based
−
+ pKa
CO3
pKa
HCO 3
H ←→ HCO 3 −
pK
10.75
10,000,000
−
a
+
brines
result
of insoluble calcium
→inHprecipitation
pKa or HCO3
HCO3 +can
H ←
5
H2 CO 3
2 CO3
pK
9.75
1,000,000
−
zinc salts, e.g. Ca(OH)2, Zn(OH)
.
pKa 2
+ H + ←a → H2 CO3
HCO3
H2 CO 3
2
2
1
1
2
2
2
2
1
1
1
8.75
100,000
Formate
saltsCO
dissolved
in water exhibit a naturally alkaline
6
CO
(
g
)
←
→
(
aq
)
7.75
10,000
2
2
CO2 (aq ) + H2O ←
→ H2CO3 (aq)
7
←
→ CO2 (aq ) pH (8 – 10). The pH of the formate brines can be adjusted
6.75
1 000
to almost any level with common acids and bases without
5.75
100
Ka1
− of insoluble
causing
The pH of fluids
H2 CO3 (aqthe
) ←precipitation
→ HCO3 (aq
) + H +(aqsalts.
)
8
4.75
10
based on formate brines can therefore be safely adjusted to
3.75
1
− ) ++H2O K
− 3 (aq )
CO2 (2aq
←
H2CO
79
a→
2
the
delivers
the
CO3level
+ Hthat
←

→
HCO
3 optimal performance.
2.75
0.1
) + H2O ←
→ H2CO3 (aq)
pH = − log [H +]
1
1.75
0.01
Kion
−
The
in itself,
and formate brines
+
a 1 is a buffer
H2 COformate
− (aq ) ←
8
− → HCO23− (aq ) + H (aq )
3
0.75
0.001
HCO
+ natural
OH+ ←
→ CO + H2O at pH = 3.75:
K10
have
−
a2
1
aH +
=3)−a+log
aH +)) buffer3 capacity
aq ) ←
→ HCO3pH
(aq
H ((aq
Ka2
2−
−
+
The pKa value in formate brines has been shown to
9
CO3 + H ←
→ HCO3
Ka
+
−
K a3
−
increase with temperature [1]. In very concentrated
+
2
pK
(3)
+ H O ←→ HCOOH −+ H2 O 2−
a
H ←
11→ HCO3HCOO
CO 3 + CO 2 3+ H 2O 
→ 2HCO3
brines, pH (and thereby pKa) are poorly defined.
pKa = 3.75
−
2−
10
HCO23− + OH − ←
→
2 + CO 3 + H2O
pKa2( aq ) 
CO
( aq ) + Ca
→
12
2−
−
− ↓ CaCO3 (s)
− 32 − + H +←
CO3
pKa2
HCO 3
→ HCO 3
4 CO 32CO
+ OH − ←
→
+3H2O
2−
[
]
(mol/L)
CO
3
−
K
+
= A −x exp(
B × pH )brines
pK←
13
CO 2 (−g
− ) + +H 2O
a→ HCO3 + H− (aq)
pH behavior
of unbuffered
formate
pK3−a]1 (mol/L)
[ HCO
HCO
H 2←
H2 CO
HCO3
H2 CO 3
CO 32 3++CO
+
H 2O1→
→
2HCO
115
3
3
−
CO 2 + H 2O 
→ 2HCO3 +14
−
2−
] )
CO3 [(Haq])×+[HCO
Ca 23+( aq
→ ↓ CaCO3 (s)
12
14
K = 12
26
+
CO
(
g
)
←
→
CO
(
aq
)
PCO32 (s)2
aq ) + Ca ( aq ) 
→
2 ↓ CaCO
[CO32 −] (mol/L)
−
K
=− A x exp(2B− × pH ) −
13
CO2 ( g ) + H2O ←→
HCO[3[ + 2H−]+ (mol/L)
(aq
− )
2−
−
[CO 3 ] [HCO3 ]
[OH − ] [CO
OH [ HCO
H2 SO4
pH =− − log
HCO
15 K
CO33− ] (mol/L)
HCO 3
10 K − log PCO2 + logCO
3 ]
3
=
A
x
exp(
B
×
pH
)
) + H2O ←→
HCO3 + H + (aq)
No formic acid
[ HCO 3− ] (mol/L)
[H +]8× [HCO 3 − ]
Traces of
14−
K =
+
0. 02pH
×P
] × [HCO16
]
CO2 (aq ) +f6HP2COO2 ←
→ H2formic
CO3 (aqacid
)
73
pK a = 3.75
2−
−
−
PCO 2 15
[formate
[OH − ] [CO
]
OH − CO32 − HCO 3−50%
2 −3 ]
H SO
CO
pH
=
2 −− log K− − log PCO + log [HCO 3 ]
[CO
] ( mg /[HCO
17
[CO3 ]+ [OH− ] = 0. 02 ×2Pf− ( mol / L2) −
L )− =3 1200 × P2 f 4
3
2−
− 50% formic
−
[CO 3 ] [HCO3 ]
[OH ] [CO
OH − CO3
H2 SO4 acid
log K − log PCO2 + log[HCO 43 ] K
HCO 3
a1
H2 CO3−(aq ) ←
→ HCO3 (aq ) + H +(aq)
8
2−
[
]
[
]
CO
(
kg
/
m3) = 1. 2 × Pf
HCO
=
0
18
3
16
0. 02 ×3 Pf2
K2
2−
−
9
CO3 + H + ←
a→
HCO3
Pf
2−
2−
] ( mg / L ) = 1200 × Pf
17
[CO3− ]+ [0OH −] = 0. 02 × Pf ( mol / L )
CO23−] (ppb
[[CO
[CO32−]
)= 0. 42× Pf
[
]
OH
=
0
19
3
2
−
−
Addition
of
strong
acid
[
]
+ [OH ] = 0. 02 × Pf ( mol −/−L ) −
CO3 ( mg / L ) = 1200 × Pf 2 −
2−
[CO3 ] ( kg / m3) = 1. 2 × Pf
10
HCO
→
CO 3 + H2O
[HCO33-+]] =OH
0 ←
18
2[HCO
=
[CO
]/R
(mol/L)
20
3
3
2−
3
Figure 1 Graph shows how the pH of unbuffered
with the addition of a strong acid.
[CO3 ] ( formate
]= 0
kg / m ) =brine
1. 2 × changes
Pf
1
0
11
19
21
2−
−
CO
[OH3 −]+=− CO
3
[HCO3P]0A=G02 E+ H2 2O →S2HCO
ECTION
2−
A6
2- 2 +
CO
+ Ca
)
→ ↓ CaCO3 (s)
[HCO
] =)[CO
]/R( aq
(mol/L)
12
20
3 − 3( aq
3
22
[
OH
]
=
0
2= [CO3 ]/R (mol/L)
−
K
13
CO2 (2g− ) + H2O ←→
HCO + H + (aq)
[CO3 ]− = 0. 02 × Pf ( mol / L 3)
23
[CO32−] (ppb )= 0. 42× Pf
[CO ] (ppb )= 0. 42× Pf
[CO32−]
pf = Vol (mL) /5
2−
3
[CO32 −] (mol/L)
= A x exp( B × pH )
[ HCO − ] (mol/L)
[CO32−]
VER S IO N
4
–
09 / 13
SECTION A: CHEMICAL AND PHYSICAL PROPERTIES
A6.1.1 Controlling pH in formate brines
There are two means of controlling pH in formate brines:
• Addition of hydroxide, in the form of NaOH or KOH. This
method can be used to increase pH in unbuffered
brines or increase buffer capacity in buffered brines.
However, the OH– ion is not a buffer and in unbuffered
formate brines, pH will drop immediately when the
brine is contacted by acid gases. Relying on OH–
addition to maintain pH of a formate fluid is therefore
not advised in applications where the formate will be
exposed to influxes of acid gases from the reservoir.
• Buffering the formate brine with carbonate /
bicarbonate. Unlike the heavy bromide brines based
on the divalent calcium and zinc ions, formate brines
are fully compatible with carbonate / bicarbonate
buffer. Buffers are designed to resist changes in fluid
pH and can cope with large influxes of acid gas.
A6.1.2 Measuring pH in formate brines
pH is a measure of the hydrogen ion (H+) activity of
a solution. Hydrogen ion activity coefficients cannot
be measured experimentally. In diluted solutions,
the H+ activity is not very different from the actual H+
concentration and pH can therefore be measured quite
accurately. In more concentrated solutions, however,
where the H+ activity deviates significantly from the
H+ concentration, the true pH cannot be determined.
High-density formate brines are some of the most
concentrated aqueous solutions that exist (see Section
A3, Water Activity and Colligative Properties), and the
H+ activity varies significantly from H+ concentration
in these brines. Any attempt to measure pH in these
brines will therefore result in a misleading value.
Although pH cannot be measured accurately in highdensity oilfield brines, it is still important for users to
know something about the acidity of these fluids. For
halide brines it has been found that measuring pH
directly on the neat brine, and only using the results in a
relative sense, is the best method [2], [3]. The main use
of pH readings in formate brines is to gain knowledge
about the state of the buffer. For buffered formate
brines, Cabot recommends diluting the fluid with about
nine parts (vol/vol) deionized water in order to obtain
the most meaningful pH measurement for determining
buffer condition.
A buffered formate brine or fluid should be diluted
with about nine parts (vol/vol)deionized water
before measuring pH.
V ERSION
4
–
0 9/ 13
C A B O T
The reasons for this recommendation are listed below
and are illustrated in Figure 2 and Figure 3, which show
examples of pH measurements made with glass electrode
and pH papers (BDF pH indicator sticks) in buffered and
unbuffered formate brines as a function of dilution [4].
The benefits of measuring the pH of formate brines
after dilution are:
1.Consistency – The measured pH of a solution should
be independent of the measuring method. Figures
2 and 3 show that for two different methods of
measuring pH, i.e. glass electrode and pH paper, both
give similar results in dilute buffered formate brines,
although they differ by up to 3 pH units in concentrated
buffered formate brines. This means at least one
of these methods gives erroneous pH readings in
buffered concentrated formate brines. In unbuffered
concentrated formate brines, the difference between
the two methods is not so significant.
2.Robustness – When measuring pH directly on the
neat brine, the formate concentration in the brine
has a large effect on the apparent pH (Figure 2 and
Figure 3). Therefore, when such a method is used in
the field, one would not have any feel for what this
pH value means without knowing the concentration
of the brine. When the dilution method is used, this
variable is removed, and the measured pH value
becomes a direct indicator of the buffer’s condition.
3. Accuracy of buffer component analysis – Traditional
methods for measuring carbonate and bicarbonate
concentrations in formate brines are complicated
or require special equipment. Cabot has developed
a new analytical method that only requires users
to make two simple measurements: pH and
phenolphthalein endpoint determination (see
Section A6.3.3). This method, however, only works
if the dilution method is used.
4. Meaningful and useful pH values – When pH is
measured after dilution, realistic pH values for buffers
and pH indicators can be measured. For example, in a
diluted formate brine, carbonate / bicarbonate buffer
buffers at pH = pKa 10.2. pH indicators also change
color at correct pH value. In undiluted brines, buffer
and indicator pH levels are too high and inconsistent.
It is important to still keep in mind that diluting brine
with nine parts of water does NOT provide a true pH
measurement because it still does not give a true
measure of the hydrogen ion activity in the original
brine. However, it provides a consistent measure that is
SECTION A6
PAGE 3
FORMATE TECHNICAL MANUAL
C AB O T
pH measurements in buffered and unbuffered potassium formate brine
15.0
KFo 1.56 g/cm3 buffered (glass electrode)
KFo 1.56 g/cm3 unbuffered (pH paper)
KFo 1.56 g/cm3 buffered (glass electrode)
KFo 1.56 g/cm3 buffered (pH paper)
14.0
13.0
12.0
11 .0
pH
10.0
9.0
8.0
7.0
6.0
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
Dilution factor
Figure 2 Effect of dilution when measuring pH in buffered and unbuffered 1.56 g/cm3 / 13.0 lb/gal potassium formate brines.
pH measurement in buffered CsFo and CsKFo brines
15.0
KFo 2.2 g/cm3 buffered (glass electrode)
KFo 2.2 g/cm3 unbuffered (pH paper)
KFo 2.2 g/cm3 buffered (glass electrode)
KFo 2.2 g/cm3 buffered (pH paper)
14.0
13.0
pH
12.0
11 .0
10.0
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
Dilution factor
Figure 3 Effect of dilution when measuring pH in buffered 2.0 g/cm3 / 16.7 lb/gal CsKFo brine and buffered 2.2 g/cm3 /
18.3 lb/gal CsFo brine.
PAGE 4
SECTION A6
VER S IO N
4
–
09 / 13
SECTION A: CHEMICAL AND PHYSICAL PROPERTIES
C A B O T
A6.2.1 How the carbonate / bicarbonate buffer works
A buffered solution is defined as a solution that
resists a change in its pH when hydrogen ions (H+) or
hydroxide ions (OH–) are added. The ability to resist
changes in pH comes about by the buffer’s ability to
consume hydrogen
ions (H+) and / or hydroxide ions (OH–).
pH = − log [H +]
independent of the measuring method, tells something
about the buffer composition in the fluid, is independent
of the formate brine type and concentration, and is
robust enough for fluid engineers to use in the field.
When the dilution method is used, pH measurements
1
can be made in formate brines both by pH electrode
(potentiometric measurements) and
by use of pH
The carbonate / bicarbonate buffer system provides
+
2
aH +
pH = − log (aH + )
]
= − log [Hmethod
paper, although1 the pHpH
electrode
is more
strong buffering
at two different
pH levels:
Ka
+
−
accurate. Due to the significant difference
between
pH
pKa
3
HCOO
+
H
O
←

→
HCOOH
+
H
+
2
3
2O
aH +
−pH
log=(−aHlog
+ ) [H ]
1 pH =brines
measured in neat formate
and
in diluted formate
• Upper buffering level at pH = 10.2
Ka
brines, it is important
always record
whether
the
+
pKa
3 [H +]2to
HCOO −pH
+ H=3−Olog
←(
H2[HO+a] +
aH1 +→
) HCOOH
pK
2−
pH = −+log
−
2−
pH = − log
1measurement
was made in diluted or neat brine.
(4)
CO3
pKa
HC
CO3 + H +←a → HCO 3
4H
K apH = − log (a + )
+
−
2
a
+
pK
3
pK
+aaHH +3 O ←
→ HCOOHH + H2 O
H
2
+a a
−
pH = − log (aH + ) 2 − HCOO
pK
−
]10.2
pH =3−−+log
H=
15
→COH22−CO3 HCO −
pKa
HCO
H + [←
HCO3
H
If pH is measured
one
would
clearly
where
pK
CO3brine,
+ H +←

→ HCO
4 in neat
3
3
3
a
Ka
+
−
pK
3
HCOO
+
H
O
←

→
HCOOH
+
H
O
a
K
3
2
+
−
also
to report
of measurement
pKa2
3 have
HCOO
+ H 3 Othe
←method
a→
HCOOH
Oa
pK+ H2pK
pHpH
= −=log
(a + ) ) the−buffered
2− a +
−
−
2+[−H +] a+
=3−− +CO
log
15 electrode)
→brine
pK
HCO
H3and
←
H2 →
CO
H2HCO
(pH paper vs. glass
the←
type
At
10.2
solution
+
H

HCO 3
4pH
3 CO3
a (H pKa HCO
3
3HCO 3 contains
6
CO
(
g
)
←
→
CO
(
aq
)
pK
2−
−
2−
a
+
2
2
and concentration, otherwise the measured
pH
the
same
amount
(CO3 ) andHCO 3−
pKa
K aof carbonate
CO3 + H ←→ HCO
4
+
−
3
pKa
3
HCOO
−
− + H2 O
+ +pKa
2 + pK
2 − + H 3 O ←
−→ HCOOH
− (a
2−
a
apH = −pH
log
)
+
+H
[
]

→
pK
=
−
log
H
HCO
+
←
H
CO
HCO
5
H2−CO 3
1
H
H
value
isCO
meaningless.
bicarbonate
( HCO
CO3
4
3
a 3 ).
2 − 3 + pK
pKaa
3 + H ←→ HCO 3 3
+

→
pK
HCO
+
H
←
H
CO
HCO
5
H
CO
6
CO2 ( g ) ←
→ 1CO2 (aq )pH = − log
3
a
3 [H ]
2
3
2
3
K
+
−
− 3 + pK
pK−
+ =H 3−Olog
←(aa+→
HCOOH + H2 OpK
2 aHCOO
a
pK
)
2−
−
2− a +
+
a
→ HpH
HCO3 + H
←
CO
HCO
5
H
CO
H
•CO
Lower
level
at pH = 6.35
H
3 buffering
2
3
2
3
CO3
pKa
H
4a
+
3 + H ←→ HCO 3
A6.2 pH1 buffering
brines
with
+−
→
) log
2←
pH =6−of
logformate
a
pH
=CO
(
a
+
+ ) CO (aq )
2 ([aq
]
= −6CO
log
H
1[HCO]2 ( g−)pH
(
g
)
←
→
H
H 7 2
CO2 (aq ) + H2O ←
→ H2CO3 (aq)
K
+
a 2
carbonate /3bicarbonate
buffer
pKpK
HCOO
+H
O ←− → HCOOH + H2 O
−
−
a a
2−
+ pKa 3
→COH22−CO3 HCO − (5)
pKa
H + ←
HCO3
5 K a HCO3 +pK
H
←
→ HCO
4 pH = CO
+←
3
3 )+
3 a −+ + H O +
3
aH O
log
(aH
pK
3
HCOO
→
HCOOH
+
6
+)
CO22( g ) ←
→ CO2−(aq
2
[
]
aH +
1 = − logpH
pH
(a=H H+−) log
a
3H
2
H applications
Formate brines7used inCO
oilfield
should
be
(
aq
)
+
H
O
←
→
H
CO
(
aq
)
2
2 pK
2
3
Ka
−
−
pKHa COcarbonate
H2 CO3 (aq
→HCO
HCO3− (aq2)−+HH +CO
(aq) −
−
8
+asodium
Kor
buffered by
where
=
6.35
−potassium
→HCOOH
pK)a←
+ H2++−←
5 HCOO
− 2→ HCO
3++H 3OK a
pK
2
3HCO 3
+
Ha2→
←

4ofHCO
pK
3 addition
+ 3H3CO
a pK 3 CO3
a
pH
=
−
log
(
a
)
a
HCOO
+
H
O
←

→
HCOOH
+
H
O
3 O3 ←
2
+
+
6
CO
(
g
)
←
→
3 −CO2 (aq
2aq
)
+
H
O
←
→
H
CO
)
H CO2 (aqa)
H
pK
7
2−
−
−
2
a
2
2
3
+
2
and potassium or sodium
bicarbonate.
main
CO2 (aqK)4+ H2OThe
←
→
H
CO
(
aq
)
K
7
CO3
pK
2
−
−
HCO 3
CO
+
H
←

→
HCO
+
a
2
3
3
+
a + pKa −3
9
CO
+
H
←


→
HCO
−
− a
3
3
K
H
CO
(
aq
)
←

→
HCO
(
aq
)
+
H
(
aq
)
8
+
−
purpose of this buffer
an
alkaline
and
pH = 6.35
solution
contains
the same
→
HCO3 + H
H2 CO
HCOpK
5 is2 to 3provide
H2 CO
3 pH
3
3
3←
HCOO
+ 3H O ←a→AtHCOOH
+ H(2pK
O a2 )− the buffered
a−
pK
2 −CO (+g ) ←
a CO 2(−aq−) + pK
− a + 3pKaK−a
−
−
→
+
2 − carbonic
− (H CO ).
2a consequence
to
from
of
amount
of
bicarbonate
(
)
and
acid
CO
pK
HCO
+fluctuating
H2 2H−←

→Kas
H
CO
(
aq
)
←

→
HCO
(
aq
)
+
H
(
aq
)
8HHCO

→
pK
+
H
←
H
CO
HCO
5
CO24(the
aq )pH
+6HCO
→
)HCO
7 prevent
3
CO
pK
2
3
3
3
3
a
HCO
+
←

→
HCO
4
3
a
3
2
3
2
3
2O 3←
2CO+3 (aqCO
3
3
3 +
3
a
Ka −
−
9
CO3the
+HHbrine.
←
a→
HCO
8into
3 HCO3 (aq ) + H (aq )
−
2−
acid or base influxes
−
2 CO3 (aq ) ←→
K
2
−
−
+
a
10
HCO−3 + OH ←
→− CO 3 + H2O
pKa
−
pK
pK
9 (+aq2CO
+H ←


2−
−
a → HCO3 pK
− CO
−
CO
)←
→
)− +3a →
HCO63 + H + 5←
HCO
5
H CO 3HCO
The
exact
and
vary
2 ( g→
CO3somewhat
H +←
4H
3of H
pK
a ) + H levels
HCO
+3H2 CO
←
H2
CO
2 CO
H2 CO 3 HCO 3
3
a will
CO32 (aq
→
3
a 3 (2aq )pK
−
7 3→ HCO
+ +3 K a
K9
2O ←
2CO
− 2−
a
CO(aq+) +6
H H←in

HCO
CO
(g )←
→ CO2 (aq ) with temperature and pressure.
H2 COan
) ←→
(
aq
)→
8
3brines
Maintaining
alkaline
pHHCO
environment
formate
3 (aq
− 3 3
2 −2
10
−
HCO3 + OH − ←
→ CO 3 + H−2O − + −pKa
2− 2−
−
→ CO
HCO
+
H+ OH
←
HCO
HCO3
5
H2 CO 3
10
HCO
+H
with carbonate
is important
3for
23 CO
3 2OCO + H O 
→ 2HCO3pKa
11←→
K a CO2 (aq )buffer
2−
−+ H O ←
3 +
2
2
H2CO3 (aq
)3 the
7/+bicarbonate
2 (aq→
K
−
6
CO
(
g
)
←
→
CO
)
9
+
CO
+
H
←


→
HCO
a
6
2
2
CO
(
g
)
←
→
CO
(
aq
)
3
3
H
CO
(
aq
)
←

→
HCO
(
aq
)
+
H
(
aq
)
following reasons:
Figure
8
2
22 −
2
3 4 demonstrates
3 how a pure carbonate buffer
10 2 − HCO3 −+ OH − ←
→ CO 3 −+ H2O
2−
+
− )water
2−
CO
( aq
+ Ca 2when
( aq ) 
→
↓ CaCO
12
works
in
a
strong
acid3 (iss)added. The
3
CO
+
CO
+
H
O

→
2HCO
11
CO
+
CO
+
H
O

→
2HCO
11
→
H2CO
Ka
−
3 3 3 (aq
2)
2
7 3 CO2 (2aq )K+a26H2O ←
CO
)←
→
CO
(aq ) CO32 − +3H + ←
− 2 (g
+
9

→
HCO
2
3
−
−
2
• Alkaline
pH helps
control
corrosion
(see
Section
carbonate
reacts with added
acid until all carbonate
H2 CO
→
HCO
) +B6)
H (aq→
) H CO
8 −←
[CO32 −] (mol/L)
3 (aq
3 ((aq
CO
10
−
HCO3 + OH
→
CO
+2)H←
O
7 2+
K
+
2+
2 aq )2+− H2O ←
2
3 (aq )
2− 3
13
−2
−
CO
(
g
)
+
H
O
←
→
HCO
+
H
(
aq
)
CO
(
aq
)
+
Ca
(
aq
)

→
↓
CaCO
(
s
)
12
3
3
2
2
3
• Presence of12
carbonate
special
is consumed. As long as there is still carbonate[ HCO
left − ] (mol/L) = A x exp( B ×
( aq
(+provides
aqH) O

→
↓ CaCO
CO+)3++Ca

→
2HCO33(s)
11 CO3/2 bicarbonate
KCO
−
3
a 2 K a2 −
−
2
−
+
] (mol/L)
9 CO 8
←(

HCO
2 − pH remains[ CO
protection
COCOcorrosion
(see
Section
B6)
solution,
around= the
‘higher
3
3 +
3 )HCO
−[−
K
HHCO
)2(←

→
(aq
(aq
aq
) in
+ ] −(mol/L)
2 −high
→
H→
CO
aq
7 against
A x exp(
B × pHbuffer
)
3 aq
3( g
13
2 (aq2) + H7
2O2 ←
3)(+
CO
aq
H
O
→
(
)
CO
) +H)HK2+CO
OH←
→
HCO
+CO
H+OH
3 (aq)←
− O
−←
K
10
+
2
2
3
HCO
+
→
CO
+
H
2
2
3
2
−
2
+
− = A
−
3 [x
2] (mol/L)
− → HCO + H (aqa)
2 − 13
exp(
Bcarbonate
× pH )
HCO
( gCO
) +2HCO
H Odecomposition
3
) + Ca H2(CO
aq33)(aq

→
↓ CaCO
) 3[[)HCO
lower
formate
rates
level’
(10.2±1).
is consumed,
H
122OCO
2→
3 ]soon as the
) ←
→ HCO
+ ]H×−+[](HCO
aq)As
83←
CO 3pH+ helps
CO 2 + H
11• Alkaline
3 2( aq
3 (s
3K (aq
(mol/L)
Ka
14
2−
−
=
+
3
−
2
9
CO
+
H
←


→
HCO
+
−
(see Section A13)
pH
drops
rather
quickly
down to the ‘lower buffer level’
3 −
3 )[+
P
H
O
←
→
H
CO
(
aq
)
[
]
(mol/L)
−
− 2 (aq
2CO
7
]
[
]
CO
×
H
HCO
CO
3
2
2
3
2 3
K a ( g←
K+
2 −K
−
2 − 10
+ − K a −+
+CO
→
CO
+
K++2aHCO
14
=O
= A x−exp( B × pH )is available
+OH
132HCO
+CaCO
H− 32]−O
←
+
H (aq
(aq
− )it remains

(aq
HH
(→
aq
)3 
9
←
HCO
− as long as bicarbonate
CO38 of
( aqcarbonate
) +HCa
→
↓)HHCO
(CO
s3→
) +←
12
• Presence
/3)[bicarbonate
helps
limit
] ×2 [→
H
HCO
))+3where
H +(3=2aq
2 CO3 (aq
8←
H
HCO
PCO
[CO 32 −]
CO
+− CO

→ 2HCO3 3 − ]
] (mol/L)
113→
2 3CO
3 (3aq 3
OH − CO32 −to HCO 3−
pH
log2[K+HCO
−Hlog
15
2O
3 PCO + log [HCO
14
K =
−
2
2
−
−
−
2
−
−
−
amount of formate2 decomposition
A13)
with3 the
[COKK3− log
] (mol/L)
[COto
]carbonic
[OH − ] [C
[HCO3 ]
OH acid
[HCO
]+ added
H2 SO4
pH
PCO react
+ log
15 Section
CO3for conversion
HCO 3
K aP−−CO+2 OH+(see
− K +
3
K3a2=− +− log
− 2=− A x exp(
2− −−←
−a
+− CO
10
HCO
→
H) ←
O
B
9
CO
+
H
←


→
HCO
+ 8 (aq
2×
+pH )
13
CO
(
g
)
+
H
O
←
→
HCO
+
H
)
3
2
H
CO
(
aq
→
HCO
(
aq
)
+
H
(
aq
)
9
3
3
−
CO
+
H
←


→
HCO
−
−
2
2
2
3
[
]
[
]
2
−
−
• Alkaline pH helps
stabilize
polymers
and
other
acid.
drop
below
this second
buffer [OH − ] [CO
×3 HCO
2 −3
H−Hlog
3
(In
aqorder
)2+− Cafor
(pH
aq−)to
→
CaCO
] (mol/L)
[ − ]12
−3−
2CO
3 (s) ]
2HCO
11
[CO↓
] [HCO
[HCO
H2 SO4
pH3= +−KCO
log
P→
+3 log
15
CO
HCO
3 COOH +3H O
2O 
3 −3 HCO
14 CO
=2 K+10
CO HCO
3
3
3
→
3 + OH ←
3 level,2 an 3acid needs
2−
additives (see Section B5)
to
be
added
that
is
stronger
than
the
16
0
.
02
×
P
PCO16
f
K
[CO
] (mol/L)
2
−
−
+
a
2
3
0. 02
× P←
−
K
+
9
+
− 2−
CO
+
H

→
HCO
f 
2
+
= A x exp( B ×
3
3
13
CO
(
g
)
+
H
O
←
→
HCO
+
H
(
aq
)
[
]
[
]
−
−
2
2
−
−
−
2
−
−
−
• Presence of
lowers
risk
H
S
gas
carbonic
acid,
which
is
formed.
As
any
CO
gas
influx
× HCO
H12carbonate
−
− of
2aq
− into
3CO
−aq ) + the
(
Ca
(
)

→
↓
CaCO
(
s
)
2
2
3
−
−
2
2
2
−
[
]
[OH − ] 2 −[CO
3
3
2HCO
11153 + OH10CO
[HCO
(mol/L)
[HCO
H32]SO
CO 3
pH
log
P→
+ log
HCO 3
HCO
←
→
CO2 3+K −H++2log
HO2O
2OH
−
14
=
K 10
− CO3
3 =+−CO
3 +3H ]O
4
CO
3 ] [ HCO
HCO
OH
←
→
2−
2 − CO 317
[CO3 ] ( m
−
3
2
[
]
[
]
CO
+
OH
=
0
.
02
×
P
(
mol
/
L
)
−
−
2
−
2
release (see16Section
solution fdissolves and converts
[CO3 ] ( mgto
3buffered
17 CO [+COCO
[OH
.→
02 × Pthe
( mol
/ L ) ] (mol/L)
/ Lcarbonic
) = 1200 × Pf
PCO 2 0A6.2.3
. 02 × Pfon next
f 2HCO
3 ]++
[CO
H 2]O= 0
11 Kpage)
3 −
33
+2
2H− O ←
+
−
=
A
x
exp(
B
×
pH
)
2
+
2
−
−
13
−
2
−
−
−
CO
(
g
)
+
→
HCO
+
H
(
aq
)
−
• Presence
carbonate
improves
well
by
aH
CO− 2][HCO
influx
therefore
not capable
−
−↓ CaCO (sacid,
aq10
) +3Ca
( aq3)OH

)2 − +3[H[HCO
− ][×
] 3 ]is[HCO
−
[ 2 −of
]pulling
[) = the
]control
pH = of
− log
K −12log P2CO
+COlog
15
HCO
3 [(2HCO
3CO
(mol/L)
HCO
+→
OH
←
→
3
3 ] 2 − H2 SO4 [CO OH
3
] ( kg
[CO32 −] ( kg
/ m3CO
1. 2 × Pf
HCO
0CO
[K3HCO
18
14
= 3 2]O= 30CO
18
3
2 − [CO 2 −]+ [OH −] = 0.202
2−−[3
2+
[CO
] (buffer
×A6.2.2
Pf ( mol
/ L)3)+] =Ca
− 2HCO
−pH much
mg
/
L
)
=
1200
×
P
sequestering
of CO
(see
Section
on
below
level.
− ) this second
2s
0
.
02
×
P
3 2 2+
CO 16
H
O

→
3
f
11 17 influx
CO
(
aq
(
aq
)

→
↓
CaCO
(
12
P
f
3 + CO
2
3
CO
+
CO
+
H
O

→
2HCO
[
]
(mol/L)
11
3
CO32 3
CO
3
2
3
−
K2
+
+ ( g ) + H− O ←
= A x2 −exp(
13
this page)
2−
− B × pH )
2−
−
−
−
[H−CO
]2× [HCO 3 2] → HCO3 + H15 (aq) pH = − log[ KHCO
−
] (mol/L)
] (mol/L)
[HCO
]] 3) = 1.22−OH
− log
PCO +[CO
log[2CO
CO3
HCO 3 [ 2−[]CO 3 ]
2
−
2
−
−
3( kg32/ −m
2
+
−
−
]
K
14
=
K
+−
3protection
[
]
×x=Pexp(
HCO
=
0
−P2 +](=mol
2[×OH
18
2
−19
[
A6.2.2
Buffer
against
CO
[CO=2/3 Linflux
]A)(ppb
17
B.42
[
]
[
]
f )= 0
3
CO
+
OH
=
0
.
02
/
L
)
3
CO
( mg
1200
×× PpH
×f Pf)
0
CO3 2−
CO
(
aq
)
+
Ca
(
aq
)
→
↓
CaCO
(
s
)
13
12
(3g+) f+CO
H22O+) 
←
→
HCO
(
aq
)
16
0. 02 × Pf
− +(H
H
O

→
2HCO
− 3
11 3 ( aqCO
)CO
+2 Ca
(3aq
→
↓
CaCO
s
)
3
12 P3CO CO
3
[
]
2
3
CO3 (pp
[OH ]3= 0
19
[ HCO 3 ] (mol/L)
2
2−
− fluids is the
The main reason for loss of pH control
in× [oilfield
of conventional
2−
2−
−
−
2 −cause [of− acidification
− major
[The
] (mol/L)
−
[KHK−−+]log
] 2[HCO
CO
2HCO
2 −[
3+
3
]
(mol/L)
3
CO
[
]
− log+
[
]
[
]
=
[CO
]/R
(mol/L)
2
−
20
2 − 15
]
OH
[
]
H
SO
OH
CO
CO
HCO
pH
=
−
log
P
HCO
CO
HCO
− 14
−
3
2
+
[
]
−
3
3
K
× pH
) P3= A(xkgexp(
+
4Pf
12. 2)× dioxide
3
3 CO
3/ mB) ×=pH
=)S.0HCO
3
=←→
133 gases
[=
]2-(=mg
g. 02
)as
+×13
H
H3O[ (←
17influx of [acid
]+ [OHCO
] =18
CO
PK2[fOHCO
(and
mol
L12
/3is
Lexp(
)influx
= B1200
3 /]
+aq
(aq
+ Ca

→
↓
CaCO
(sA) x(mol/L)
such
CO
HCO
These
completion
carbon
gCO)3++CO
Hare
→
HCO
+ H)[[HCO
(. aq
2 (0
2− or×diffusion
3 brines
] ×))[HCO
163 − ]3( aq
0
023−3)×]-](mol/L)
PCO
20
2(
2 Hboth
2 P
HCO
f [CO3 3 ]/R[ HCO
[3−CO
] (ppbf )= 0. 42of
[CO32−]
] (mol/L)
× Pf
[OH −] = 02
19
CO 2
3
14
=
K
−
2
weak acids with
a pKa higher than the pKa of formic acid.P
gas (CO−2) into
the
from
the
rock
]
(mol/L)
2 − surrounding
−
−
2 − wellbore
− [ CO
−
2
−
−
3
− [ + CO
K
− CO
] [HCO
] 2[−CO
OH
P [(HCO
+ log
[HCO3 ] = 0 15 [ +] [-pH = −− ]log
) = 13. 2 ×/ LP3[)fCO
20O ←
= A3x]exp( H
B 2×SO
pH4 ) [OH
− ] ()kg
18
3 ]
3 / mHCO
13 K −21log
]H
H[OH
=[HCO
3(aq
−2− 3
+ COCO
[CO332 −]++CO
22 g )−3+
2 17 → HCO
formations:
2− [CO3 ] ( m
H. 02××HCO
−3
2
−
−
− ] = 0. 02 × Pf− ( mol
[
]
[
]
2
−
−
]
[
(mol/L)
16
0
P
×
H
HCO
[HCO
]
=
[CO
]/R
(mol/L)
HCO
20
3
[
[
]
f ] =15
CO
)
=
0
.
42
×
P
[
OH
0
CO
3 ] (ppb
[
]
3
3
14
K =19
−
[
]
OH
[
]
H2 SO4
CO
HCO
pH
=
−
log
K
−
log
P
+
log
HCO
CO
HCO
3
f
3
14
K =
3
3
3
pf = Vol (3mL) /5
21 CO [HCO3 ] = 03
PCO
22 PCO 2[OH −] = 0
−
2
−
2− 2
+
−
[CO32 −] ( kg
] = 02− −
18 3 − ]− [HCO
2−−
[/H
− L )] × [HCO
[2CO
] ([mg− /−L] )[= 1200
22−− × P
[OH
+.−02
.+02log
×2-P]/R
− − ](mL )[/5
−
=fCO0[CO
f ( mol
3 ]K
= Vol
20
[CO
]×3 PHCO
×3-]]P=pH
] (ppb
]] 4 [pfHCO
OH[HCO CO− 3−]2 3−[CO3HCO
[HCO
]
H
SO
CO
HCO
= [−CO
log
log
15−] = 17
)
=
0
.
42
[OH
0 pH16
CO
19
[
[OH − ]
f [OH
14
=−[CO
K K(mol/L)
2
]
OH
H
SO
3log
2− P
3
3
CO
3
= 23
−3 log
+
log
15−0[HCO
CO
f
3
2
4
CO 22
3 ]= 0
3
3
3
3
[OH
21
[HCO3− ] = 0 16
3 ] =P0. 02 × Pf ( mol / L )
0
.
02
×
P
CO
2
PAGE 5
f
V E R S I O N 4 18– 0 9 /[HCO
1 3 ] = 02 −
I( kg
O N2/−A
63) = 1. 2 × P
[SCOE C2 −T]CO
−
−
2−m
f
3 [CO ]+ [OH ] = 0. 02 × P ( mol / L )
] ( mg
/ L−) = 1200
[HCO3-] = [CO3172-]/R (mol/L)
20
[CO 32×−]Pf [HCO3−] [COH22−SO
3
3
==0Vol
OH/ −L )3 [CO
]× P) /5
(mL
pH = −f log K −23
log
PCO
+[CO
log−2]p−[=fHCO
15
HCO
] (pp
4
3
OH
−
3
3
19
]
0
.
02
(
mol
−
[
CO
]
2−
3
f
−
22
[
OH
]
=
0
−
3
21
[
]
HCO
=
0
[CO32 −] ( mg / L ) = 1200 × P3f
16
0. 02 × Pf 16
] = 0. 02 × Pf ( mol
+ [OH
/ L/)L )
3 17
( mol
24Pf[CO3 []HCO
3 ]=
2−
− 0. 02 ×
3
.
.
A exp (B pH )
[CO3 ] ( kg / m ) = 1. 2 × Pf
[HCO3 ] = 0
18
2−
2−
−2 −
pf =3[2-CO
[HCO3-] = [CO
]/R ((mol/L)
20-10
Vol
−/5 ] (ppb )= 0. 42
]
2[
−CO
× P[fCO 2 −] ( kg[CO
] =]−[=0]OH
OH
00−.]02
××PPf 2((−25
mol
// LL−))A−]==3.894
23
]mL2[)CO
3
10
[
]
=
0
3OH
17 3−]19
[
]
[
/ m3 3)×=P1. 2 × Pf
HCO
0
+
=
.
02
mol
CO
(
mg
/ L ) = 1200
18
−
21
[HCO
= 0[CO22
3
[CO 2×−]P(fmg
17
0. 02
3/ L ) = 1200
[16
]3=×0P. 02 × P ( mol / L )
3
CO f ]+ [OH
3
2
2
1
2
1
2
1
2
1
2
2
2
1
2
1
21
1
1
2
1
2
1
2
1
2
1
1
2
1
2
2
2
1
2
1
1
2
1
2
2
1
2
1
2
1
2
1
2
1
1
2
1
1
2
1
1
2
1
2
1
2
1
1
2
1
2
1
2
2
2
2
1
2
2
2
2
2
2
2
2
2
2
2
2−
•
2
2−
2
FORMATE TECHNICAL MANUAL
C AB O T
pH behavior of carbonate / bicarbonate buffer when adding strong acid
12
11
pK a2
10
9
8
pH
1
1
2
2
1
3
3
2
4
43
5
5
4
6
65
O
aq)
O3 (s)
q)
6
7
7
8
7
8
9
9
8
10
109
11
10
11
12
12
11
13
13
12
14
13
14
15
15
14
16
15
16
17
17
16
18
18
17
19
18
19
20 −
OH
20
19
21
20
21
22
22
21
23
7
pK a1
pH = −6log [H +]
pH = − log [H +]
5
aH +
pH = − log (aH + )
4 log (a )
aH +
pH = −
+
H
[H +]+ K a
pH = −3− log
pKa
HCOO
+
H
O
←

→
HCOOH
+ H O 0.6
3
K
0
0.2
0.4 2
1
1.2
1.4
1.6
1.8
2
pK0.8
HCOO − + H 3 O +←a→ HCOOH + H2 O
a
Fraction
of buffer consumed
aH +
pH = − log (aH + )
pKa2
2−
−
−
2−
CO32 −
pKa2
HCO 3−
CO32 − +− H +←
→KHCO
3−
a
pK+←
of strong
acid
a2 
+H O
pKAddition
HCOO
+
→
HCOOH
+
H
O
a
CO
pK
HCO
CO
+
H
←

→
HCO
3
2
3
3
3
3
a2
pKa1
−
−
+
→
pK
HCO
H +The
←
H
CO
HCO
H
CO
Figure
pH
in
water
buffered
with
carbonate
as
a
function
of
added
acid
3
a
3 − +4
2
3
2
3 (H ). The x-axis shows the fraction of
pK
1
−
a1
→
pK
HCO2 3− + H++ ←
H
CO
HCO
H
CO
3
a
2
3
2
3
pK
2−
−
−
1
the
that
by the added acid. As
be seen,
CO3 carbonate
pKcan
HCO 3buffers twice, first at pH = pKa2 = 10.2 (upper
CO3 buffer
+H ←
ais2→consumed
HCO 3
a2
buffer level) and then at pH = pKa1 = 6.35 (lower buffer level). If the added acid is carbonic acid (from CO2 influx), the pH
CO2 ( g−) ←
→
COa2 1(aq )
−
+ pK
can
drop
lower
than pKa1 = 6.35.

→
pKa1
HCO(3gnever
H
←
HCO3
H2 CO 3
3
CO
)+←
→
CO much
(aq H
) 2 CO
2
2
carbonic acid being present in the brine. With a large
influx of CO2, the pH drops down to the lower buffering
pKa
level (pH = 6.35) where it stabilizes. Measurements of
CO2 (aq ) + H2O ←
→ H2CO3 (aq)
(7)
pH in formate brines exposed to various amounts of
CO2 (aq ) + H2O ←
→ H2CO3 (aq)
2−
CO2 have confirmed that pH never drops below 6 – 6.5.
−
COK3a
pKa
HCO
3
−
+
H2 CO3 (aq ) ←
→
HCO
(
aq
)
+
H
(
aq
)
(8)
This pH is still close to neutral, meaning that this brine
3
− (aq )
a→ H CO
COCO
(aq ) + H2OK
←
2 (3aq
H
) + H +(aq)
−→ HCO
2 2 3 (aq ) ←
3
system cannot be ‘acidified’ to any great extent by
=
6.35
pK
HCO
H
CO
K a3
−2
a2 −
3
CO32 −+ H + ←

exposure to CO2. However, carbonic acid and a small
K a→ HCO3 −
+
CO3 + H ←
→ HCO3 −
athe original pH of +the receiving brine
Depending
onK
amount of formic acid are also present.
H2 CO3 (aq ) ←
→ HCO3 (aq ) + H (aq)
system,
dissolved
CO22 −remains in the brine as either
• Unbuffered formate brines: The pH of these brine
−
−
HCO23−−+ OH+ ←
CO 32 −− + H2O
K→
a(H CO
− 
carbonic
) 3in+equilibrium
with dissolved CO2
systems responds in a similar fashion to halide brines
CO3 ++ HOHacid
←

→2 HCO
3
HCO
←
→
CO
H
O
3
3
2
gas or bicarbonate (HCO3–), according to reaction 8. This
when exposed to CO2 gas. However, they do have a
− more CO gas enters into
2−
is demonstrated
in→
Figure
5. As
higher initial pH, and the pH drop will be limited as the
CO
+
CO
+
H
O

2HCO
2
2−
32 −−
2
−
HCO
OH2carbonic
←
→
CO
+concentration
H2O3 −
3++CO
3 2HCO
the
acid
builds up and pH
formate brine is a buffer in itself (pKa = 3.75). At such
CO
→
3 brine,
2 + H 2O 
3
2−
2+
drops
and allows
unbuffered
brines
to acidify.
low
significant
= − log [H +] amount of corrosive formic acid
1 pH apH
CO
( aq
)
→ ↓ CaCO
32 −( aq ) + Ca
3 (s)
CO32 −( aq ) + Ca 2 +( aq ) 
→ ↓ CaCO
(
s
)
−
2
is present in the fluid. If there is any chance of an acid
3
−
[CO3 ] (mol/L)
CO + CO 2 + H 2OK 
→ 2HCO
−
2 B × pH
aH +
=) − log (aH + )
2−
= A xgas
exp(
pH the
CO
g ) + Hdifferent
→brine
HCO3systems
+ H3 + (aq)in Figure 5 react
The23 (three
influx,
use of unbuffered formate
brines is
[
(mol/L)
−]in
CO
2O ←
3
−
K
]
[
(mol/L)
+
HCO
=
A
x
exp(
B
×
pH
)
3
CO22(following
g ) + H 2O ←
→toHCO
(aq)
−
Ka
+
−
3 +H
−
the
ways
a
CO
influx:
highly
discouraged.
2+
]
[
(mol/L)
pKa
3
HCOO
+
H
O
←

→
HCOOH
+
H
O
HCO
2
3
2
CO3 ( aq ) + Ca ( aq ) 
→ ↓ CaCO3 (s)
3
[H +] × [HCO 3 − ]
2−
[CO ] (mol/L)
= [H +] × [HCO 3 −divalent
K• Conventional
−
against
H2S influx
K]
brines
can not be3 −
pK
= A xA6.2.3
exp( B ×Buffer
pH2 −) protection
−
HCO2O2 ←→
HCO3halide
+ H + (aq
)
KCO=2 ( g ) + P
pKa
CO3 + H +←a → HCO 3
4
[ HCO 3 ] (mol/L)
buffered with
because
the
Influx
into− a wellbore is often
P carbonate / bicarbonate
−
− accompanied by
2
[CO 32 −of
] CO
[
]
[
[
]
OH − CO32 − HCO 3−
H
SO
OH
CO
HCO
pH = − log KCO−2 log PCO + log[HCO 3 − ]
pK
2
4
3
−
a
−+ (H
−
−

→ HH22SCO
corresponding
metal
, ZnCO
) 2 − HCO − hydrogen
sulfide
S).
is3 a[OH
weak
5 2 −] HCO
+
− P + carbonate
[CO
] acid
[COwith a pKa of
3 + H ]←
[HCO
]
pH = [−Hlog
− log
log[HCO 3 (CaCO
3 OH
3 CO3
2 H2 SO
] × [KHCO
4
3
3
3
3 ] CO
=
K
precipitates out of solution resulting in formation of
around 7. This means that at a pH of 7, equal amounts
PCO
solids
fluid.
These
of6hydrogen
sulfide
(H2S)) and hydrogen
bisulfide
0
. 02 × Pinf the2 clear packer / completion
− CO2 ( g ) ←
−
→
CO2 (aq
2−
−
[CO–3)2 will
] be
[OHAt− ]higher
[COpH, more
[HCO
OH–−6), and
[HCO 3 −low
] pH (2
H2the
SO4 brine.
= ×−Plog
K − log PCO +alog
CO3
HCO 3 (HS
3 ]
0pH
. 02
divalent
naturally
present
in
f brines have
2−
2−
−
[CO
] ( mg
[CO3influx
]+ [OHof
] =CO0.202
× Pf ( mol / L )on the partial pressure of
the
, dependent
HS32– −will
be/ Lpresent
) = 1200and
× Pf at lower pH more H2S will exist.
2−
−
[
] ( mg /unless
[CO
]+ [OH ]lowers
CO3, further
= 0. 02 ×the
Pf ( mol
/
L
)
CO
L ) = 1200
× carbonate
Pf
3
pH. The CO2 partly converts to
Therefore,
the
buffer in the formate
2− 2 −
2−
3
. 02
×]P=f 0
[CO[03HCO
] (mol/L)
[
]
CO
(
kg
/
m
)
=
1
.
2
×
P
f
3
3
carbonic
acid
7),)which is very corrosive.
brine
by
large
influxes
of CO2, the
= A(Equation
x exp( B × pH
2 − is overwhelmed
−
CO
(
aq
)
+
H
O
←
→
H
CO
(
aq
)
3
7
−
2
[CO3 ] ( kg / m
[HCO
) = 1.22 × Pf 2 3
] (mol/L)
[ HCO
3 ]= 0
3
2−
carbonate
buffer
traps
and retains this toxic gas in its
[CO
]
[CO32 −]+ [OH −] = 0. 02 × Pf ( mol / L )
(
mg
/
L
)
=
1200
×
P
3
f
• Buffered formate brines are capable of buffering
less harmful form, Knamely
bisulfide,
HS–.
−
a
[CO328−2]−(ppbH)2=CO03 .(42
[CO)3+2−H] +(aq)
aq )×←P
→
HCO
[OH −] =− 0
3 (aq
f
2
−
2
−
−
[
]
large
of CO2. Unless the influx is unusually
]0= 0
CO3 ] (ppb
( kg /)=m0). 42
= 1×. 2P× Pf
HCO] =3amounts
[CO
[CO3 ]
[[OH
3
Kf a
2−
−
+
large,
The
fact
that
any
H
S
is
converted
to HS– in buffered
9
CO
+
H
←


→
HCO
- the brine
2- maintains a pH around the upper buffer
3
3
2
2−
−
[HCO
2 − ] = [CO− ]/R (mol/L)
−
3 HCO 3
[CO 3 ] [HCO3 ]
[
]
[
H
SO
OH
CO
CO
22
4
3
3
level3(pH
= 10.2),
formate brines does not mean that
the gas is
[HCO
] = [CO
]/Rwhich
(mol/L)is high enough to prevent
3
[CO32−] (ppb )= 0− . 42×− Pf
[CO32−]
[OH −] = 0
2−
CO2 ( g ) ←
→ CO2 (aq )
2
1
1
2
(6)
1
2
1
2
2
CO3
1
HCO3
1
2
2
1
2
[HCO3P-−] A= G0E 62[HCO
SECTION A6
[HCO33 ]] == 0[CO3 ]/R (mol/L)
−
[OH ] = 0
[CO32 −] ( mg / L ) = 1200 × Pf
[OH −] =− 0
2− ]= 0
HCO
3
2− /L)
[[CO
3 2 −] = 0. 02 × [P
f ( mol
] ( kg / m3) = 1. 2 × P
CO
−
10
HCO3 + OH ←
→ CO 3 + H2O
11
CO 3 + CO 2 + H 2O 
→ 2HCO3
12
CO3 ( aq ) + Ca 2 +( aq ) 
→ ↓ CaCO3 (s)
VER S IO N
pf = Vol (mL) /5
pf = Vol (mL) /5
2−
4
–
09 / 13
−
2−
2−
2−
2
2
−
−
HCO 3
H2 CO 3
SECTION A: CHEMICAL AND PHYSICAL PROPERTIES
C A B O T
pH in various brine systems as a function of CO2 influx volume
12
11
Buffered formate brine
10
pH>6.35:
CO2 mainly converted to
bicarbonate (HCO3-),
which does not promote
corrosion
9
pH 8
Unbuffered formate brine
7
6
pH<6.35:
CO2 mainly converted to
carbonic acid (H2CO3),
which promotes corrosion
5
4
Calcium bromide brine
0
50
100
150
200
250
300
350
BBL gas influx / BBL buffered formate brine
] / 70°F, 1 atm)
pH = −(2%
logCO[2H +21°C
400
450
500
1
Increasing
2 time of
pHCO=2−influx
log (aH + )
aH +
Figure 5 pH as a function of CO2 influx in a typical halide brine, an unbuffered
formate
K brine, and a buffered formate brine.
pKa
3
HCOO − + H 3 O +←a→ HCOOH + H2 O
pK
−
2−
CO3 + H +←a2→ HCO 3
4
pK
− be+ seen
As
can
afrom
→ H2the
pKa
HCO
CO3graph in Figure 5, the decrease
3 +H ←
in pH of one unit is not really a good measure of
how much carbonate buffer is present in the brine,
CO2 (therefore
g )←
→ COthe
) capacity of this buffer. Cabot
and
true
2 (aq
therefore uses the actual carbonate concentration as
a measure of the capacity of the buffer rather than the
scientifically defined ‘buffer capacity’.
scavenged and made permanently safe. If the buffer 5
was to be overwhelmed by an excessive influx of CO2 /
H2S, then H2S gas would come back out of solution
when pH dropped to below around 7.0. CO2 gas would 6
first be present in equilibrium with the bicarbonate in
the brine at a lower pH (6.35). It is therefore important
to remove any HS– contamination
from used field
+
= − log
1 neverpH
muds, and
lower
the[HpH] or let the buffer deplete
7
in a formate mud or brine that has been exposed to H2S
2
aH + –
pH = − log (aH + )
without first checking if it is contaminated with HS . If
K
+
−
there is any
corrosion
H2S 8
3 concern
HCOOabout
+ H 3HO2S-related
←a→ HCOOH
+ Hthen
2O
scavenger should be added (see Section B6, Compatibility
with Metals and Section B5, Compatibility with Additives). 9
pK
−
2−
CO3 + H +←a → HCO 3
4
+
1
1
1
2−
CO3
−
HCO3
H
H
CO2 (aq ) + H2O ←
→ H2CO3 (aq)
In alkaline
brines that
are buffered with carbonate /
bicarbonate buffer, the following equilibrium exists
Ka
− bicarbonate:
+
and
H CO3 (aq )carbonate
←
→ HCO
pKbetween
3 (aq ) + H (aq )
a 2
K
2−
−
CO3 + H + ←
a→ HCO3
(9)
2−
−
pKa = 10.2 CO3
HCO 3
1
2
2
2
pH = − log [H ]
pK
−
A6.3 Buffer
addition
and
10
amaintainance
→ H2 CO3 2
HCO3 + H + ←
5
pH = − log (a
pKa2
−
2−
− CO
HCO
+ OH HCO
←
→
H2O
pKthe
3 H
In
is2+CO
typically
lost by exposure
3
a 3 field,
3
aH + carbonate
H
Whenever formate brine is used in the field, it is important
to influx of acid gas. As acid gas initially enters the
K
−
3 acidHCOO
+ H 3 O +←a→
HCOOH
+ H2 O
to maintain the ability of the buffer to resist
influxes.
brine,
(CO32–pK
) ais gradually
converted to
−
2 − carbonate
CO
+ CO 2 + H 2O 
→
2HCO3
11
6
CO
(
g
)
←
→
CO
(
aq
)
–
+ 3
2
2
In order to do this, both buffer capacity and total1 buffer
(HCO3 ), whilst pH remains at around the
]
pH = − log [Hbicarbonate
pK
2−
−
−2 −
2−
2+
concentration need to be monitored and 4maintained.
upper
buffer
(pH
= pKa↓ =CaCO
10.2).
all
CO(3sWhen
HCOcarbonate
CO3 + H +←
a → HCO
3
CO3 ( aq
) + Calevel
)
→
)
3
2
a(H aq
pH12= − log (aH + )3
+
is converted, the buffer loses its ability
to maintain
pH.
pKa
−
−
[CO32 −] (mol/L)
→ H2 CO
pK
HCO3 + H + ←
HCO3
5
H2 CO
− a +
K
3 Ka
3
A6.3.1 Buffer capacity
= A x exp( B ×
carbonate
component
the
buffer system
is −now
13 − + H3The
CO
g )→+ HCOOH
H2O ←→
H pK
(aq
3
HCOO
O +←
+ H2 O HCO3 +of
a )
2 (
[ HCO 3 ] (mol/L)
In buffered
brine,
the
referred to as ‘overwhelmed’ or ‘swamped’ and has no
CO2 (aq
) + Hit2Ois←
→carbonate
H2CO3 (aq)component
7 formate
6 alkaline
CO2 ( g ) ←
→2 −CO2 (aq
+
− buffer pH at the upper buffering
of the buffer that provides buffering at the
level.
pK capacity
2−
−
+ )more
] ×3[−HCO 3to
]
CO3
pK
HCO
CO3 + H ←a →[HHCO
4
14
=
K
pH of 10.2. Bicarbonate is mainly added in order to
Any further influx of acid gas can anow easily
lower pH3
PCO 2
Ka
−
+
pK
−
−
H2 CO
(aqcarbonate
) ←
→ HCO
(aqis
) +aHfunction
(aq
balance 8alkalinity
of 3the
as3 pH
down
the
by3
a → to
H2 CO
HCOprovided
5 ) HCO3 + H + ←
H2 2−CO
3 lower buffer level ( −pKa = 6.35)
−
−3
[CO 32 −]
OH
[
]
pH
=
−
log
K
−
log
P
+
log
HCO
15
CO3
HCO 3
CO
3
of the carbonate-to-bicarbonate
ratio.
The
carbonate
the
bicarbonate.
(See
Figure
5.)
K
2−
−
+
9
CO3 +isHtherefore
←
a→ HCO
CO2 (aq
H2O ←
→ H2CO3 (aq)
concentration
alone
the 3true7measure
of)+CO
6
→ CO2 (aq )
2 (g )←
the brine’s buffer capacity.
It0.is
to notice that whilst pH of a buffered
16
02important
× Pf
Ka
−
+
formate
brine
is a function of the ratio of
−
2−
−
H
CO
(
aq
)
←

→
HCO
(
aq
)
+
H
(
aq
)
8
2
3
3
10
HCO3 + OH ←
→ CO 3 + H2O
2−
−
The scientifically correct definition of buffer capacity
concentrations
of carbonate
[CO32 −] ( m
17
[
]
[
]
CO
+
OH
=
0
. 02
× Pf ( mol / L )and bicarbonate, the
K
2−
−
9
CO3 + H + ←
a→ HCO3 3
is: “The number of moles of acid or base necessary
toCO2 (aq )+ H2Ocapacity
the buffer to maintain pH around 10 – 10.5
←
→ H2COof
7
3 (aq )
−
2−
[CO32 −] ( kg
[HCO3−] = 0upon the actual carbonate concentration.
change11the pH of
liter of solution
unit”. 18
depends
COone
→with
2HCOone
3 + CO 2 + H 2O 
3
1
+
)
−
1
2
2
1
1
2
2
1
1
1
2
2
1
2
10
12
V ERSION
13
4
2−
3
−
2−
HCO3 + OH − ←
→ CO 3K a + H2O −
1
H2 CO3 (aq ) ←
→ HCO3 (aq ) + H +(aq)
8
CO ( aq ) + Ca ( aq ) 
→ ↓ CaCO3 (s)
–
2+
0 9/ 13
11−
K2 −a2 −] = 0− −
[
OH
2−
+
9 2 −+ COCO19

HCO
]→(mol/L)
CO
3 + H [←
CO
→
3 2HCO3 3
3
2 + H 2O 
SECTION A6
K
CO2 ( g ) + H2O ←→
HCO3 + H + (aq)
12
[H +] × [HCO 3 − ]
= A x exp( B × pH )
−
2] (mol/L)
[ HCO
[HCO
]
=
[CO
]/R (mol/L)
3
20
2+
3
CO ( aq ) + Ca− ( aq−) 
→3 ↓ CaCO
3 (s)
2−
10
2−
3
HCO3 + OH ←
→ CO 3 + H2O
K
−
+
[CO32 −] (mol/L)
PAGE 7
[CO32−] (pp
C AB O T
13
12
11
10
9
8
pH 7
6
5
4
3
2
4
pK
−
2−
CO3 + H +←a2→ HCO 3
5
pK
−
HCO3 + H + ←a1→ H2 CO3
2−
7 7
8 8
−+
+
a1
− K a−1
a 1K
H82 3CO
←

→
aq
)H++(H3aq
aq
H)2K
CO
)HCO
←

→
((aq
H2 CO
(aq
)←
→
HCO
) +HCO
) ))+ H (aq)
3 (aq
3 (
3 (aq
3 (aq
6
CO2 ( g ) ←
→ CO2 (aq )
7
10 10
CO2 (aq ) + H2O ←
→ H2CO3 (aq)
MANUAL
K a−2
2 2K
− a2 +
−
2 − 2 − + + Ka
Titration
9 + HCO
←

+→
H HCO
←3 −

9 9curves
COCO

→
HCO
3 → HCO3
3
3 3+ H ←
−
Phenolphthalein
endpoint
2−
−
2−
−2 −
−
− −
10
HCO
+ OH
←
CO
+endpoint
HCO
O + H2O
HCO
+ OH
←
→
Methyl
HCO
→
COorange
3→
3 +3OH ←
3 3+ H2O 2 3
Water + standard buffer
KFo + standard buffer
2−
−
2
−
2−
CO
+ CO

CO+ H+ OCO
+ 2HCO
H 22HCO
O 
→
2HCO3
CO11
→
11 11
3 CO
2→
3 +
2 +2H32O 2
3 3
Ka1
−
+
H2 CO3 (aq ) ←→ HCO3 (aq ) + H (aq)
22−+ 2 +
2+
2−2−
CO
( aqCO
) + Ca
)
→
↓) CaCO
(↓s)CaCO3 (s)
( aq()aq
+ →
Ca
(↓aq

→
CaCO
12 12 CO12
3 (3aq ) + Ca3 ( aq ) 
3 (s3)
Ka2
2−
−
+
CO3 + H ←
→ HCO3
2− 2−
] (mol/L)
[CO32 −] (mol/L)
[CO[3CO
]3(mol/L)
−
+
− K− + +
K K
x exp(
A×xpH
= A= xAexp(
B=×BpH
)e
CO
(
g
)
+
H
O
←
→
HCO
+
H
(
aq
)
CO
(
g
)
+
H
O
←
→
HCO
+
H
(
aq
)
13 13
CO13
(
g
)
+
H
O
←
→
HCO
+
H
(
aq
)
− −
2
3
2 2
2 22
3 3
] (mol/L)
[ HCO
[ HCO 3− ] (mol/L)
] (mol/L)
[ HCO
3 3
8
9
−
-0.2
11
CO + CO 2 + H 2O 
→ 2HCO315 15
+
−
−
2−
HCO3 + OH − ←
→ CO 3 + H2O
-0.1
0.0
0.1
14 14
10
-0.3
−
CO3
pKa2
HCO 3
(
aq
)
+
H
O
←
→
H
CO
(
aq
)
CO
(
aq
)
+
H
O
←
→
H
CO
(aq)
CO2CO
(7aq
)
+
H
O
←
→
H
CO
(
aq
)
2
2 22
F O 2R2 M2 3A T3 E− T 2E C 3H N I C A L
pKa1
HCO3
H2 CO 3
[−H]3+−] ]× [HCO 3 − ]
× [HCO
[H +[]H×+[]HCO
K= = 0.2K = 3
0.3
K 14
PCO 2
PCO P
CO 2
2
0.4
0.5
1
pH = − log [H ]
2
aH + 2 −
pH = − log (aH + )
CO3 ( aq ) + Ca 2 +( aq ) 
→ ↓ CaCO3 (s)
12
Figure− 6 Titration
curves
for
buffered
water and buffered potassium
formate.
Both
fluids
2− ×
Ka
0. 02
16
0. 02
Pf contain the same amount
+
16 16 0. 02
× P×f P[fCO
] (mol/L)
pK
HCOO
+ Hcarbonate
HCOOH + H2 Obuffer (17.8 kg/m
3
3 / 3.75
−
K a 3 / 6.25
+
3 O ←→
= Albs/bbl
x exp( BKHCO
× pH )3). The
of
added
/­bicarbonate
13
CO2 ( g ) + H2O ←→
HCO3 + Hlbs/bbl
(aq) K2CO3 and 10.7 kg/m
−
]
[
(mol/L)
HCO
2
−
2
−
−
2− 2−
−
3
2 − 13B-1 −alkalinity
phenolphthalein and methyl orange endpoints from the standard
API
RP
titrations
are
shown.
[CO/
17
17
[
]
[
]
[
]
[
]
CO
+
OH
=
0
.
02
×
P
(
mol
/
L
)
CO
+
OH
=
0
.
02
×
P
(
mol
/ L ) No methyl
[CO[CO
] ] ( mg
17
[CO3 ]3+ [OH ] = 30. 02 × Pf ( mol
/L) f
f
3 3( mg / L )
2 −due to the formate
orange
endpoint
the
buffered
formate
brine
/
formic
acid
equilibrium
starting
to
−
2−
+
−
+ pKa 2 can be− detected in
[H ] × [HCO 3 ] pKa
CO3
HCO
CO3 + H ←→ HCO 3
−
2− 2−
− 3−
2
[CO
14
K =
[HCO
/3
HCO3 ] = 0
[CO[CO
] ] ( kg
18
[HCO
establish at a higher pH.
18 18
3 ]0=[0
3 3( kg / m
3 ]=
P
CO
pK
2
−
−
pK
HCO3 + H + ←a1→ H2 CO3
HCO−
H2 CO
2−
−
− 3
[CO 32 −] [HCO3−]
[OH − ] [CO
OH
H2 SO4
pH = − log K − log PCO2 +a1log[HCO 3 3]
15
CO3
HCO 3
A6.3.2 Total buffer concentration
−
2− 2−
−
−
total
[CO)3=
[CO[CO
] ] (ppb
[OH
= 0 [OH ]of
= 0carbonate / bicarbonate buffer
19
] = ]0amount
19 19 The[OH
3 3(ppb )= 0
The total buffer concentration in a brine that is
available decreases by the amount of carbonate that
CO2 ( g ) ←
→ CO2 (aq )
2- -] =[HCO
2- 2--]]/R
16 / bicarbonate
0. 02 × Pf is defined as
[HCO
(mol/L)
= [CO
]/R (mol/L)
buffered with carbonate
20
[HCO
] = [CO[CO
]/R (mol/L)
2020 is precipitated.
3
3 3
3 33
2–
the combined concentration of carbonate
(CO3 ),
pH = − log [H +]
[CO32 −] ( mg
17
[CO32 −(H
]+ [CO
OH −]),=and
0. 02any
× Pf ( mol / L )
/ L ) = 1200 × Pf
bicarbonate (HCO3–), carbonic
acid
3. Formate− decomposition
increases
buffer
2
3
−
−
21
[
]
21
[
]
HCO
=
0
HCO
=
0
21
[
]
HCO
=
0
3
3
carbon
dioxide
in
concentration
3
− the brine. If the
a
pH
= − log
(aH + ) (CO2) gas dissolved
[CO32 −] ( kg / m3) = 1. 2 × Pf
18 H + [HCO3 ] = 0
total
buffer
Small amounts
of− soluble carbonate
and bicarbonate
= Vol
=) Vol
CO2 (aq
) + H2Oconcentration
←
→ H2CO3 (aqhas
) been removed from
p)f/5
/5 (mL) /5
pf =pfVol
(mL(mL
−
−
K
+
−
22
[OH
= 0 [OH ] = 0
2222 can[OH
] = 0]as
pKneed
HCOOfluid
+ Hin
←a→
HCOOH + new
H2 O buffer will
a
the
reactions,
to be
form
a result of formate decomposition if
3 Oother
2−
− exposed
2−
added to theKfluid. 19
the brine
temperature
2 − 2is
] (ppb
[CO32−]
0.an
42×extended
Pf
[+OH −] = 0
[CO
] = [0CO
−
. 02 ×] =Pfto0( mol
)( mol
.high
02
/ L ) )=for
a1
3
/ L×/) PL[f CO
2323 [CO23
3 ]3= 0. 02 ×3 Pf ( mol
H2 CO
pK
2−
−
−3 (aq ) + H (aq)
2 − 3 (aq
a → HCO
+ )←
period of time (See Section A13, Thermal Stability). The
CO3
pKa
HCO
CO3 + H ←→ HCO 3
3
2There
ways in
the buffer
dominant decarboxylation reaction is reversible, and
] = [COconcentration
]/R (mol/L)
20
K
2 − are
− which[HCO
+ three
a
3
3
2
] [CO3 2 − ]
pK
[CO[CO
− + H+ ←
−
a
CO


→
HCO
2 − ]3 2 −
3
− − ment3−of
→ Hcan
pKuse:
HCO
+ H ←
HCO3 2424Hthe
in a33formate
brine
establish­
equilibrium
well
a
2 CO3 be altered during field
2 CO
[HCO
( mol
/ Lclosed
) ( mol / HPHT
L)
24
[3HCO
/ .L )in
3 ] =[HCO
3 ] = . . ( mol
3 ]=
.
.
.
(
A exp
ApH
pH )
(B BpH
)exp) (B decomposition
A exp
systems usually
limits
formate
to a few
-10
-10
•
[HCO3−buffer
]= 0
•A =
−
− ) increases
2 21
A 3.894
=in3.894
103.894
10brine
1.
Influx
of−acid
gas
typical
formate
formulations.
A =25
10•-10
2525 percent
2 H2O
HCO
OH
←
→
CO(CO
CO
←
→
CO
3 )+
3 +
2 (g
2 (aq )
concentration:
pf = Vol (mL) /5
−
B 2.193
=Determining
2.193
B = 2.193
22
B =26
[OH ] = 0
2626 A6.3.3
buffer concentration and capacity
−
2−
−
− −]
[
[
]
(10)
For standard
CO 3 + CO 2 + H 2O 
→ 2HCO3
OH = 0 water-based mud filtrates, API RP 13B-1 [5]
[CO32 −] = 0. 02 × Pf ( mol / L ) 27 27 [OH27] = 0 OH = 0
23
recommends
that2 −carbonate and bicarbonate content
2− 2−
CO2 (2aq
→
[CO
− ) + H2O ←
2 + H2CO3 (aq )
] ] = [0CO3 ] = 0
2828 [CO28
3 3= 0
CO
+ Caconverts
( aq ) 
→
↓ CaCOfrom
Influx
of )CO
carbonate
are measured
by pH titrations. Alkalinity in the form
3 ( aq
3 (s) the buffer to
2
[CO3 2 − ] 2 −
2− − −
−
−2− 2−
− −
−
pK a is
OH
OH
CO
CO
bicarbonate. TheK carbonate
concentration
(buffer
of
content,
content,
and HCO
[
]
(mol/L)
pK apK
OH
CO
CO
HCO
3HCO
3 a content
[
]
mol
L
(
/
)
HCO
=
24
3 3
3
3 3
−
+
3
K
−
=
A
x
exp(
B
×
pH
)
+
.
.
a
CO
(
g
)
+
H
O
←
→
HCO
+
H
(
aq
)
(
− )
exp
A
B
pH
H
CO
(
aq
)
←

→
HCO
(
aq
)
+
H
(
aq
)
2
2
3
capacity)
therefore3 drops whilst the total buffer
determined
by the
combination
2
3
2− of
−
[ HCO 3 ] (mol/L)
− −a phenolphthalein
− −]
[]HCO
] H SO
[OH
]2−] 2−] [CO
[CO[−CO
[HCO
] ] [HCO
[OH[OH
]
P
P
3 f H SO
H2 SO4
3 3
3 3
-10
2 2orange
4 f4
concentration
titration to an endpoint
of pH3 = 8.2,
and Paf methyl
A = 3.894 •+10carbonic
K a (carbonate
2−
− 25 + bicarbonate
+
CO3 + H+ ←
→ HCO
−
3
[
]
[
]
×
H
HCO
3
acid + dissolved CO2) increases by the amount of CO2
titration to an endpoint of pH = 3.1.
K =
B = 2.193
entering the
PCO 2 brine. 26
−
2−
−
2−
−
− −
HCO
+ log
H2O[HCO [OH
formate
] = 0 OH − CO32 − HCO 3− In[CO
3 +
] [brines,
[OH − ] of the
[CO methyl
27
HCO3 ] theHdetermination
pH =3 +−OH
log K←→
− logCO
PCO
2 SO4
3 ]
3
2
2. Influx of multivalent cations decreases
buffer
orange titration endpoint is complicated by the formate /
2−
[CO3 ] = 0
28
concentration:
formic acid equilibrium that is present at pH = 3.75
−
2−
CO 3 + CO 2 + H 2O 
→ 2HCO3
2−
−
0. 02 × Pf
( pK a formic acid). This is illustrated in Figure 6. As can
OH −
CO3
HCO3
2−
CO3 ( aq ) + Ca 2 +( aq ) 
→ ↓ CaCO3 (s) − (11)
be seen, the formate / formic acid equilibrium starts to
2
−
−
[CO
]
[HCO3 ] [CO 2 P−] ( mgH/ LSO
[CO32 −]+ [OH −] = 0. 02 × Pf ( mol / L ) [OH ]
1200
× Pf
2− 3
2 )
3 f
establish
at
a4=pH
significantly
higher than the methyl
[
]
(mol/L)
CO
3
−
K
+
= A x exp( B × pH )
CO
( g ) + H2O ←→
HCO
+
H
(aq)
−
2
3
An influx
of multivalent cations consumes
the
buffer
orange
endpoint.
The fact that only one of the two
] (mol/L)
[ HCO
2−
−
3
3
[CO3 ] ( kg / m ) = 1. 2 × Pf
[HCO3 ] = 0
by precipitating
out insoluble calcium carbonate.
standard titration endpoints can be determined in a
+
−
3
4
5
6
1
2
7
3
8
4
95
106
11
7
12
138
9
14
10
15
11
16
12
17
13
18
14
19
15
20
2−
−
3
Added OH- [mol/L]
−− −
2 − −2 −
− −
2 − 2 −−
−2 −−
[CO
[C
[PHCO
OH HCO
] COCO
[CO
] ] [HC
pH
PCOKlog
log
=− −log
−+ log
log
HCO
15
CO
HCO
[HCO
pH
= −=log
KpH
−Klog
Plog
3 ][HCOOH
CO 23+ ]
3 OH
CO 2 +
3 33
3 3
33 3
2
Added
H−+log
[mol/L]
2
2
1
1
1
2
K =
[H ] × [HCO 3 ]
PAP
GE 8
CO 2
−
[OH ] = 0
SECTION A6
pH = − log K − log PCO2 + log[HCO
[HCO3 ] = [CO3 ]/R (mol/L)
-
2-
−
3
]
OH
−
2−
3
CO
HCO
−
3
[CO32−] (ppb
− )= 0. 42 × Pf
[OH − ]
[HCO3 ]
H2 SO4
[CO 32 −]
VER S IO N
[CO32−]
[
CO
4
–
09 / 13
8
1
pKa1
HCO3 + H ←3→ 3H2 CO3 CO 2
H2 CO
3 (aq
3 H2 2
32→ HCO33 (aq ) + H −(aq )
) ←pK
8 HCO
2 CO
2−
Kaaa12−1 18[ − 2 −] [HCO3+ ] =− 0
− Ka2
2 −−3
2 − ++
K
−−
−
9 PCO +CO
+ H +←

→5
[OH − ] HCO
[
H2 SO4 pK
9f =OH
log
+(aq
H
←

→
15 − −pH = − log K − log
CO
HCO
HCO
)
←
) + [HHCO
(aq
p8HCO
(Vol
3 [HCO

→HCO
←
HCO
p3f = HCO
)
Vol
mL
/5
3 ]
3 ])
32)(3mL
3−3(aq
/5
3
3
2 CO
3
a1
2 CO
3
K a1 CO 3 −
KKa2→
−
+
22 22 [OH ][OH
= 0]= 0
2 − + H + ←
−
H
CO
(
aq
)
←

→
HCO
8
a
9
CO

→
HCO
2
2
3
3 (aq
3
3
K a1→ HCO3 −
COCO
H )←

6
CO2 ( g ) ←
→ CO2 (aq )p =
32 − +CO
KH
H

→
HCO
(aq ) +3 −H +(aq)
89
CO
+
→3−2HCO
119
+2←
a22O
(mL) /5
2 32 −3+(aq
Vol
6
−
CO
(
g
)
←
→
CO
(
aq
)
f
CO
+
H
←


→
HCO
−
2O N A : C H E 2M
2− I C2A
3
[OH ] = 0
K2
−
S E C T I23
L02
D PPH (Ymol
P R O −P E R T −I E S
2−
19 32 − [OH ] = 0 9 C A CO
B O2T− + H + ←
−
[CO3[CO
] =3 00−.].02
/S LI C) /A LL )HCO
= A0××.N02
23 16
a→
HCO3
→10
H2O(2g3−−)+←
f 10
PPff (×mol
HCO
OH
←
3
K→
3 + OH ←
6CO 3 +CO
→
CO
) 32 −− + H2O
+
a222+(aqCO
2
−(+aq
− Ca
9
H
←


→
HCO
CO
)
+
(
aq
)

→
↓
CaCO
(
s
)
12
− )OH
23 − (aqO)
CO
+ H2−O←
←
→CO
H CO
3
7
10
HCO
→
23(aq
3 +
2O
10
HCO
→20
CO23322−−++3 H
H[HCO
[CO32 −] = 0. 02 × Pf ( mol / L )
3− + OH ←
2
] = [CO32-]/R (mol/L)2 −
2−
−
−
3
[
]
[
CO
17
]
[
]
[
]
CO
+
OH
=
0
.
02
×
P
(
mol
/
L
)
CO
(
mg
/
L
)
=
1200 × Pf [CO− 3 ] (mol/L)
[
CO
]
10
HCO
+
OH
←
→
CO
+
H
O
2
−
f
−2
3
33
−
K
+
3
− 3
3 2−
−
−
A x2 exp
−
2 −− ) + H O ←
CO
→
HCO
L ) 3/2 L−+) CO 2 + H 2O 13
( mol( mol
/ CO
24 24 [HCO[HCO
−
HCO
←
→=CO
23( g
2+ H O
3 + H (aq) 10
3 ]=3 ]=

→
2HCO
11
CO
+
CO

→
2HCO
11
3 + OH
3 +H
.
.
] (mol/L)
[ HCO
.
.
−
−
2
−
3
2
2
3
(
)
K
(
)
CO
(
aq
)
+
H
O
←
→
H
CO
(
aq
)
−
exp
A
B
pH
3
7
+
exp
A
B
pH
a
−
−
2
2
−
1
−
2
2
2
3
10
3
HCO
+
OH
←
→
CO
+
H
O
brine,
means
the
API alkalinity
1. Prepare
consisting
mL
[CO3 2 − ]7 formate
H
←
H )−(5=aq
−of
23−+(aq
3 2HCO
CO
COa)2sample
+ H2O→
→
11
[HCO
]((aq
CO
kg2) /+3m
1.)2 ×fluid
Pf sample
HCO
= 0standard
CO/2 L(aq
) + H18
→[that
H
)
2 CO
3 3
3 3 2HCO
33](aq
2O ←
2CO
CO
→
118
[HCO3−] = .
( mol
-10
322 −−+ CO 2 + H 2O 
3−
•
2−
[45
HCO
=0
• determining
2+
2 +− 21
carbonate
(brine
or
filtrate)
and
mL
water.
= A3.894
10for
25) 25 isAunsuitable
+ mud
3 ] deionized
= 3.894
10-10 12
CO
(
aq
)
+
Ca
(
aq
)

→
↓
CaCO
(
s
)
A exp (B . pH )test method
(
aq
)
+
Ca
(
aq
)

→
↓
CaCO
(
s
)
[
]
[
]
CO
+
CO
+
H
O

→
2HCO
12
11
×322O
H) + H
HCO
3
CO23(22aq
→
H2CO− 3 (aq)3
7
3
K←
2−
a2223+
−−
149 2. Measure
K =
CO
H )+pH
←


HCO
and bicarbonateKconcentrations
in+ this brine system.
of
this
sample
a
calibrated
glass
CO
→ 2H
11
2→
+( aq
2 − ↓with
aq
+
Ca
)

CaCO
(
s
)
12
−
− 2 + H 2O 
322 −−(+
3[→
a1
3 + 2CO
-10
3
3
•
−
−
]
(mol/L)
CO
(
aq
)
+
Ca
(
aq
)

→
↓
CaCO
(
s
)
12
CO
[
] (mol/L)
P
A = 3.894 10
H226
CO3 26
(aq ) ←

→
HCO
(
aq
)
+
H
(
aq
)
8
3
3
B K=
2.193
3
22
B
=
2.193
− 2 −++ CO CO
2− 2HCO
] = 0 = A x exp( B ×[CO
−
K11
CO
+ 2H 22O+ K[
→
+
a1
− [OH
3=+ 0. 42 ×
pH3223−−2−)]−](mol/L) = A x exp
] (ppb
electrode.
+332H((gaq
()aq
H2 CO3 (aq )19
←
→ HCO
(aq) CO2 ( g ) + H2O ←→
CO
)
P
[OH 3−3] =(aq
8
0 ) + H13
CO
13 HCOCO
+ 2Ca
( aq
)3HCO

→
↓
CaCO
(
s
)
+) 2)H
O ←→
+
H
(
aq
)
−
12
[
3CO
f
−
2
−
−
CO
−
− 3 ]](mol/L)
] (mol/L)
2+
[[2HCO
(mol/L)
K P +[−HCO
[exp
CO
OHHCl
]3) 12
CO
pH =( g−)log
KO−K a←
log
log3=[−−HCO
15 3. Perform
CO
HCO
3 H ++ (
+3aq
− this− are:
323−3 ) + Ca
COor
(3=aq
→
K CO
1 →
K a−]2OH
13
2−
consequences
titration
toHCO
pH
−
3 ( aq
]= 0
= AA) xx
exp3
=of0HCO
B = 2.193
HCO
(aq
←
←
HCO
)8.2
+HH2with
aq
22 (−3−g
2O
3 +
27
8
27 + [OH
13
2→
+→
CO
) ++)a)H
H
HCO
(−(aq
)0()0.02N
9 The CO
[HCO
−(aq
2

2 CO
3→
−]](mol/L)
(mol/L)
−2Ca
CO
2 ( aq
3− [+
K a[→
]
CO
=
.
02
×
P
(
mol
/
L
)
23
32 − + H + ←
3− CO
+
(
aq
)

↓
CaCO
s
)
12
3
23
2
K
3
f
+
]
[
(mol/L)
10
HCO
→
CO
+
H
O
HCO
3 3 + OH ←
3
[HCO
]
=
[CO
]/R
(mol/L)
=
A
x
exp(
9 - The
CO3phenolphthalein
+H ←

→
HCO
20
3
2
3
+
−
13
CO
(
g
)
+
H
O
←
→
HCO
+
H
(
aq
)
can
be used to[H ] × [HCO 3 ]
H2SO
endpoint pf −
+
− the phenolphthalein
3
3
2 − 2 −endpoint
2 4
2
3
[ and
] [report
−
K
(mol/L)
] =3 0] = 30
K 23 − ]
[OH −] = 0
− H +× HCO
−
28 28 [CO3[CO
13
CO[2[HCO
( gof
)2 −3+]] (mol/L)
H
→
HCO3
14 and Kcarbonate
=
CO
+ ←
14
K 32=the
2O ←
9
CO
+HHvolume,
a→
3
determine the combined
hydroxide
as
V
(mL),
+ required per mL
[
+] × [HCO
−K]HCO
3 of −titrant
3
=
A
x
exp(
13
CO
←3→
PCO
] HCO3 + H (aq)
16
. 02( g
×[H)P+f] ×HP[CO
HCO
[CO3 2 − ] [ HCO − ] (mol/L)
2O
14
2− 2−
K0
−
−
−
−
2
−
+
−
14
K ==22 −[sample:
[CO32 −] = 0
concentration
[−OH −CO
]+322[−CO
].
fluid
pK a pK2a
10
HCO − + OH −−OH
←
→
+− 3HCO
HCOHCO
3
[CO
−)
−] ×P
2[HCO
− − 3 −] = −
HCO
HCO
−
( mol
2O 3
3
24
3O ]
3
2 − /L
−
− 2 −]
CO
+
+
H
→
2HCO
11
[CO
2
P
OH
[
]
.
− H2 SO[4CO
10
pH
=
−
log
K
−
log
P
+
log
HCO
15
CO
HCO
HCO33 + OH21
←
→ CO
+
H
O
3
2
2
3
[
]
HCO
=
0
K− −2 log PCO + log[3HCO 3 ] 3A expOH
15CO
CO
HCO
14
KpH==2 −−3log CO
(B−3. pH
3 ]
3 3
2 cannot be used to
[)H[3+2HCO
]−× [HCO
]− [CO 2 −3
3 3
- The−methyl orange
endpoint
−
2−
2−
−
−
17
2−
[
]
[
]
−
−
CO
+
OH
=
0
.
02
×
P
(
mol
/
L
)
−
+
−
−
−
2
P
−
[[CO
2−
−
f log [HCO − ] 14
3
OH −K = CO
pH ==3 [+
− OH
log
KKHCO
log
15
] 3 ] [HCO[HCO
] 3 ] P P H 10
pK[aOH []OH ] [CO3[CO
OH
] × [CO
CO3
HCO3
2−
CO 3+
3 ]
3
3 ]P
←
→
CO
+ H AO(13)
3
SO
CO
log
−/5
log
15H SO HCO
CO332 − P HCO
HCO33−
•
=23−(−Haq
2 + P25
f
(mL
CO + log [2HCO
3−
3
f 214
determine
the total buffer
concentration
(including
3.894
10-10OH−
KppH
−
f =Vol
CO 2
) +K)Ca
( aq
)+
→
↓ =CaCO
12 42 4 CO
−
2−
22
[
3= −
3 (s)
[
OH
]
=
0
OH
[
]
CO
pH
log
−
log
P
log
HCO
15
CO
HCO
2
−
−
2
−
−
−
CO
3
P
3]
3
3
CO
+
CO
+
H
O

→
2HCO
11
CO
[
−
−
2
3
2
2
3
[
]
[
]
[
]
itHis
possible to16
detect an
endpoint
[OH ]
HCO = 0
18
CO3 11
HCO
−
3 [
Pf if
pH[=CO−2log
K − log PCO +COlog
15
0. 02
× Pf
CO 33 + ).
COEven
→
16
0. 023× Pf 2 K 26
] (mol/L)
2 SO4 2HCO3
2 +H
2O 
−
B −=+ 2.193
2 − brine, the use of this
OH − CO332− − HCO 3=− A x [exp(
])
CO 3
pH
−)Plog
KO
−←
log
PCO situations
+HCO
log[−HCO
15
in a very
diluted
formate
Depending
on
pH,
four
exist:
2=−
13
3aq
CO
(
g
+
H
→
+
H
(
16
0
.
02
×
2−
[
]
2+
CO
=
0
.
02
×
P
(
mol
/
L
)
23
2
2
3
CO
→ 2HCO3
1116
3→ ↓ CaCOf (s)
0. 02
×− CO
Pff 2 + H 2O 
[ HCO ] (mol/L)
3 +
CO 2 −( aq ) + Ca 2 + ( aq ) 
12
−
endpoint
[CO32 −] ( mg /3L ) = 1200 × Pf[CO23−2 −
17 33 (s) [CO32 −]+ [OH2−−] = 0. 02
( mol
/>
L3−2)×11.1
CO33 ( aqwould
) + Ca give
( aq )erratic

→ ↓calculations.
CaCO
12
17 × Pf1.pH
]+Pf[OH −] = 0. 0227
× Pf ( mol[OH
/ L ) ]= 0
16
0[CO
. 02
2]−
2
−
[[CO
−
[CO32 − ] (mol/L)
19 [[OH
2 −]2=
+
−]+0
3 2−
0. 02 × Pf
17
−] = 20
××) P
// LLCaCO
)) (14)
−
K
2−
f ((mol
3 H
3
+[HCO
Ca
aq
s16)
[CO
12
] 28
17 = ACO
]−B+P+]×[[)OH
])= 03.(.−02
x[CO
exp(
pH
CO
OH
02
P→
mol↓
CO
×
CO3)2 − ]
13
3 3 ([aq
3 (0
CO2 ( g ) + H2O ←→
HCO− 3− + H ++ ([aq
f
2
−
− [ CO3 −] (mol/L)
3
[
]
3
CO
=
16
0
.
02
×
K
2
−
14
= 2 −3]B+] =×[fOH
[CO3 ] ( kg /2m
exp(
pH
[ L0) the] (mol/L)
[HCO
13 Another
therefore
required
(15)
CO2 ( gmethod
) + H24
←→
HCO
+ H . (18
aq)for. measuring
HCO
0 −])=2-0. 02 × Pf ( mol / L ) 3
[HCO
( mol
17
18 = A xK[CO
− ) = 1. 2 × Pf [CO3 ]
3 ]/=HCO
2O is
3 ]3 =
3 -−
2−
2 − ] (mol/L)
[
[ HCO 33− ] (mol/L)
−
P
CO
[HCO
]
=
[CO
]/R
(mol/L)
(B pH )
20
CO
exp
A
3
2 −]
2 − [CO ]+ [OH −] = 0. 02 ×[P
K
23
[[HCO
(3mo
33 −] =
18
=−pK
A a[xCO
fexp(
3 HCO
bicarbonate
concentration.
One way of doing this is
Unless
of3−[+OH
have
]
13
)]+=large
H002O
→
HCO
H +−(]aq
) 17 been
CO
CO3 added
HCO
− amounts
18 CO
−2 − 3
2 (2g
322 −−
− [
3−]
[CO
17
[CO
[OH
]−=←
−+log
0. 02
× P+f log
( mol
/L) − ]
] (mol/L)
HCO
[CO
[H++] × 25
[HCO 3 −− ] A = 3.894 • 10-10
OH
[
3 −
pH
=
K
log
P
HCO
3 3]
15
CO
HCO
3
[
[
]
CO
3
CO
3
3
HCO
=
0
18
to this
most of22−−this −
3
3 fluid, one can assume that
14 by using
K = [Ha Garrett
] × [HCO 3Gas
] Train (GGT). The GGT determines
−
−
−
2−
[OH ] 18 [[CO
HCO3 )]==[HCO
CO33 ]][(ppb
00. 423×] Pf [CO
[OH ] =The
0
14 the Ktotal
= amount
[OH −] =− +0 is from
19
Pf 322−][−C
PCO 2 of carbonate and19bicarbonate.
[CO
−[H] =
[[OH
] ×0[HCO 3 − −]carbonate.
CO32−] ]
18
21 − alkalinity
HCO
P26
[
B = 2.193
−]33
CO 2
[
=
0
19
2
−
−
2−
−
−
[CO33 ]
[OH
19
[=
[OH ] [CO
]
OHvery CO32 −2- 14
[HCOtherefore
H SO
CO] =2 −0] [HCO
pH = is
− log
log PCO + logand
15 method
timeK −consuming
not
HCO 3− K0
3− ]
2- 3−
16
.[02
[HCO
[CO
]/R19
(mol/L)
[OH − ] [CO p = Vol (mL) /5 [CO32−] (
[CO
]/R
OH −3-] =CO
[HCO 320
]
H22SO44
CO−−]3×=3-]P0]=f P
pH = − log K − log PCO
+ −log
15
HCO
20
CO[HCO
[[HCO
OH
23
3 ] (mol/L)
33
3
f
[
]
OH
=
0
−
27
popular on the rig, and does not differentiate between
22 2. pH
[OH=]11.1
=- 0= [CO 2-2-]/R (mol/L)
2
−
] = 02 − HCO −
[HCO
19 OH −[OH CO
20
[CO
22
− 3−
=−233−−-]]log
K−−33 log
[HCO
= [CO
]/RP(mol/L)
15
20 pH
CO + log [HCO 3 ]
3(16)
3
[[CO
](
2−
[[HCO
]
OH
=
the different buffer
components.
19
-]+0[OH ] =
2- 0. 02 × P ( mol / L ) 17
3
[
CO
CO
[
]
f
32 − ] = [CO ]/R (mol/L)
3
CO3 = 0
20
3
16
0. 02 × P 28
. 023 × Pf ( mol / L )
23
(17)
[HCO3-] = [CO32-]/R (mol/L)
20
3 ]−]==00
21 2 −
[HCO3−] −= 0
16
0. 02 × Pff
21 [[CO
HCO
−
23-−−
] = [CO ]/R (mol/L)
OH concentrations
CO3
HCO3 been pK a20
[CO32 −]
[[.HCO
Carbonate
and bicarbonate
have
0[HCO
(18)
18
33 −]] = 0 3
21
2×
−
3]P
16
02
21
[
]
HCO
=
0
[
17
=
[CO322 −−]+ [OH −−] = 0. 02 × Pf ( mol / L )
f
p
(
)
CO
(
mg
/
L
)
=
1200
×
P
Vol
mL
/5
=
3
p
(
)
Vol mL /5
32 − −
f
using
a dual-titration
[CO
22 2−] method.
[OH −[]HCO
= 0 −]
17 measured
[CO3 ]+by
[OHCabot
] = 0. 02
Pf (−]mol
/ L ) [CO
= 31200
×f Pff
22
2− ]
OH3 −]3]=−(]0mg
21
[[HCO
= 0 / L [)CO
[×OH
ppf ==−Vol
((mL
Pf
3
3
H2 SO
−2 − ] =
[
mol
L
(
/
)
HCO
24
Vol
mL)) /5
/5 [ 2 −]
4
−
21
[
]
HCO
3
2
−
22
−
−
[
OH
]
=
0
3
f
and more accurate method, based
</Am
11.1
3 =0
17
[OH
].exp
CO
+(0kg
=3)0=. 02
CO32−
22 3.9.0
[HCO
OH332−<
]−]−=−pH
[[CO
(1B. 2.×pH
[HCO3−a] =simpler
×PfP()fmol / L )
0
18 Recently
=
2−
p
(
) /5 [CO
Vol
mL
2
21
[
]
=
0
](
−
f
[[CO
OH3 ]]3=]=](0=kg
0 0/. 02
19
[CO3 ] = 0. 02 × Pf 22
[HCO3 ]pH
( mol / L[)CO
23
m
)
=
1
.
2
×
P
= 0measurement and phenolphthalein
3
18 on simple
×
P
(
mol
/
L
)
23
OH
(19)
f
32 −
-10 f
•
−
2
−
A
=
3.894
10
25
−
2
−
22
[OHp ]==Vol
0 (mL ) /5 [CO ]
CO
23
3 ]]=
= 00..02
18
titration, has been developed.
(20)
CO
02 2-×× PPff ((mol
mol // LL ))
23 [[[[HCO
f
3
22
OH23−−2]3−=-]0= [CO
[HCO
]/R (mol/L)
20
2−
−
3] = 0. 023× P ( mol / L )
[
CO
23
[
CO
]
[
]
[
]
3
f
CO
From
the
carbonate
2−
(ppb
)= 0[CO
. 423 2×− ]Pf / bicarbonate
[OH −] = 0
CO32− pH relationship
19
2−
B 3=2− 2.193
−
326
−
[CO3 ] = 0. 02 × Pf ( mol / L )
[HCO3 ] = .
. 422×− ]Pf ( mol /[LCO
[OH ] =field
0
mol
([CO
/2]L−3(ppb
)] = )= 0[CO
24 for determining
19 A simple
[HCO
) 3 ] 23 the carbonate
24.
3Equation
or laboratory method
and
/
.12
−
].Figure
[[CO
02
×[CO
Pf3(3B(2 −mol
/)L( mol
)6, determine
A exp (23
B
pH ) in
exp
pH
−3−] ]−]==0.A
L
/
)
HCO
24
2[CO32−] (
OH
=
0
27
[bicarbonate
] =33−0] = . [ratio,
OH
19
[HCO
( mol / L )the bicarbonate
24
[HCOconcentration
] = [CO32- ]/R (mol/L)
20 buffer
CO3(B2 − ]..R.
Calculate
)
3exp
A
pH
.
−] = 0
-10
•
-10
21
[
HCO
•
(
)
[HCO3 ] = [CO3 ]/R (mol/L)
exp B pH ( mol / L )
A 10
A = 3.894 10 24
25
20
3
= 3.894
25
[AHCO
[CO 2 − ]
3− ] =
−
.exp
. pH )
[ACO=323.894
Laboratory
testing of formate brines with known
concentration
-10as:
28
( 2(mol/L)
- ] = 0A •• 10
2[HCO
]= ( . ) 3 .
(m
[CO
24
25
-10 3B
−]
3 Vol
[HCO
]
=
[CO
]/R
20
−
=
p
A
=
3.894
10
25
mL
3
−3 ] =
f
A exp/5(B pH )
[B[OH
( mol /−L )
HCO
24
additions− of carbonate and bicarbonate
B =shown
2.193
26 has
22
]
-10 2 − .
=
2.193
3= 0
26
•.
−
AOH= 3.894
25
pK a
exp
A 10
CO (B pH ) HCO
21
[HCO
=0
–
3− ]buffered
B
] = [CO32–3] (mol/L) / R 3 (21)
pH of
formate brines is dependent
on the 26
A = 3.894 • 10-10
25
21 that[HCO
B == 2−2.193
2.193
26 [HCO
3
3 ]= 0
[OH −] = 0
-10 2−
−= 0
[
]
27
•
OH
27
−
−−] = 0. 02
A
=
3.894
10
25
[
=
CO
×
P
(
mol
/
L
)
23
p
(
)
B
=
2.193
Vol
mL
/5
[
]
[
]
carbonate-to-bicarbonate
ratio [6], [7]. The following
26 [OH
OH3−−3]]=] =0 0 CO3f
HCO3
f
21
[[[HCO
P
H2 SO4
22
27
[OH −] = 0
) /5
2 − pf = Vol (mL
2] −= 0
B = 2.193
26 f
27
[COcarbonate
]= 0
22 relationship,
28
[OH −] = 0 R, has been found between
[[BOH
the
4.pH
<
CO
=
0
28
−3 =]9.0
3
26
OH=− ]22.193
=− 0
27
p = Vol (mL) /5
2 −] = 0
[
2
−
CO
28
[
CO
]
bicarbonate
molar ratio and brine pH:
[[OH
222
−28
−
[OHf −] =(22)
23−2 −
CO]33−−−2=−−]0= 0
0
−
27
23 and[CO
pK
32 −] = 0. 02 × Pf ( mol / L )
OH
CO
HCO
pK
OH
CO
HCO
a
[
]
3
3
mol
L
(
/
)
HCO
=
24 [[OH
a
3−
[CO3 ] = 0. 02 × Pf ( mol / L )
23
27
CO3−−2]−3=]]==0insignificant
0A .exp322(−B− . pH
28
CO
insignificant
2
−
)
−
2−
pK
OH
32 −
CO3 2−
HCO3 −
[CO3 ] = 0
2−
[CO3 ] (mol/L)
pK
28aa
−]−−=
3f2 to
3− ]
0. 02CO
/ LHCO
) [HCO
[CO323
]
[[OH
[OH −]
HCO
[×CO
] determine
−( mol
[CO
[CO
3 −23]
OH
R=
= 2A− ]× exp( B × pH ) (12)
P-10P
]]=difficult
Pf
3−
H2 SO4
0 • 10
28
[
CO
f 32− H2 SO4HCO
pK
−−
OH
CO
A
=
3.894
25
3
−
a
3
3
3
2−
−
2−]
−]
HCO 3 ]] =(mol/L)[CO 2 − ]
[[to
[[OH
−]
CO
HCO
[[HCO
( mol / L )
pH
to[[be
above
9.0PPwith
24
3 ]
H
]
]
CO3
HC
OHneeds
CO233adjusted
HCO
f
2 SO
4
HOH
− 3−
−−
2−−
f
2 SO4
[HCO33−] = A ..exp 3(B .. pH ) ( mol / L )
24
pK
OH
CO
HCO
[
]
[
]
[
]
OH
CO
HCO
[
CO
a
]
3
3
where
before
concentration
B = 2.193
Pfcan Hbe
3 2−
3
26
A exp (B pH )
2−
− the bicarbonate
3
2 SO−4
[CO3 ]
[
[OH ]
[HCO−3 ] = .
( mol / L )
24
A = 3.894 • 10-10
25
determined.
[CO(3B2−]. pH ) [HCO3−]
OH−]]= 0 A exp
25 A = 3.894 • 10-10
Pf
H2 SO4
[[OH
27
•
B = 2.193
A = 3.894
10-10
25
26
2−
] = 0 amounts of sodium or potassium
B = 2.193
equivalent
26
28 The[CO
3
–
OH −]2–= 0
27 and[[CO
can− easily be calculated
B = −2.193and bicarbonate
26 carbonate
2−
[OH −3] =‐]0and [HCO3 ] are the molar concentrations of
pK a
OH
27
CO3
HCO3
2
−
carbonate
and
bicarbonate.
This
relationship
is
shown
from
these
molar
concentrations.
This simple
field
−
[
]
CO32 − = 0
28
[
]
OH
=
0
27
2
−
−
[CO3 7] and
= 0 is valid for pH measured with a glass
28 in Figure
[CO3 ]in more[HCO
] in Section
[OH −]is explained
method
detail
C2.SO
P
3
H
f
2
4
2−
−
pK
OH −
CO 2brine
HCO with
[CO32 −] = 0
electrode
diluted
nine
−
−
pK aaparts deionized 28
OH − in formate
CO33
HCO33
A6.3.4Buffer
requirement
for− field use
water.
used−] to determine buffer
−
2−
[CO322−−] can be[HCO
[OHThis
] relationship
pK
OH −
CO3
HCO3
Pf
3−
H2 SO4
−
[
]
[
]
[
]
OH
CO
HCO
concentration and buffer
capacity3 of buffered
The recommended buffer concentrationa required
Pf formate
3
H2 SO4
2−
−
−
[COdepends
]
[on
]
[OH ] brines
HCO
brines in the field. This means that both carbonate
in formate
the
Pf
3
3 application.
HThe
2 SO4
and bicarbonate concentrations can be determined
amount of time the brine will be in contact with the
just by measuring pH and performing the standard
reservoir fluids, and the expected level of acid gas
phenolphthalein titration. The method is as follows:
influx, are important factors. In well suspension and
packer applications where the formate brine may be
5
[HCO3−] = 0
2
2
2
2
2
2
2
2
2
2
2
2
2
V ERSION
4
–
0 9/ 13
SECTION A6
PAGE 9
FORMATE TECHNICAL MANUAL
C AB O T
18
H2O
NaFo
KFo
KFo (non-analytical)
CsKFo
[CO32-]/[HCO3-] (mol/mol)
16
R = [CO32-]/[HCO3-]
14
12
10
8
6
4
2
0
9
9.5
10
pH
10.5
11
Figure 7 Relationship between the carbonate-to-bicarbonate molar ratio (R) and pH in buffered formate brines.
The relationship was developed from a range of formate brines and deionized water with known amounts of
buffer (carbonate + bicarbonate) added. The carbonate content was measured by titration to endpoint of 8.2,
i.e. phenolphthalein endpoint, after dilution with 9 parts deionized water, and pH was measured with a calibrated
glass electrode, again after dilution with 9 parts deionized water.
18
[CO32-]/[HCO3-] (mol/mol)
K2O3/KHCO3 (wt/wt)
Na2CO3/NaHCO3 (wt/wt)
16
R = [CO32-]/[HCO3-]
14
12
10
8
6
4
2
0
9
9.5
10
pH
10.5
11
Figure 8 Relationship between the carbonate-to-bicarbonate ratio (R) and pH for formate brines buffered with carbonate
and bicarbonate. The carbonate-to-bicarbonate ratio is given as a) molar ratio of [CO32–] to [HCO3–] (see figure 6 above),
b) equivalent potassium carbonate (K2CO3) to potassium bicarbonate (KHCO3) ratio (wt/wt), and c) equivalent sodium
carbonate (Na2CO3) to sodium bicarbonate (NaHCO3) ratio (wt/wt).
exposed to well conditions for a long time, a high level
of buffering is appropriate. In applications where well
exposure times are short, and in applications where no
acid gas is expected, a smaller buffer concentration
will do. In drilling fluids, which can be monitored and
conditioned at the surface, less buffer is required.
higher than wanted. This problem can be solved by
including some bicarbonate. Although addition of
bicarbonate does not contribute to buffering at high
pH (pKa= 10.2), it contributes to balancing alkalinity of
the carbonate as pH is a function of the carbonate-tobicarbonate ratio. (See Equation 12 and Figure 7).
Adding only soluble carbonate to the brine provides a
high level of buffer capacity, but the pH might become
When determining buffer levels for field applications,
one needs to consider that some oilfield formate
PAGE 10
SECTION A6
VER S IO N
4
–
09 / 13
SECTION A: CHEMICAL AND PHYSICAL PROPERTIES
brines come preloaded with buffer and some do not.
Cesium formate brine from Cabot delivered by the
manufacturing plant typically has a pH of about
10.2 – 10.4 and contains some 0.07 mol/L CO32– and
0.03 mol/L of HCO3–. This corresponds to equivalent
potassium carbonate and potassium bicarbonate
buffer levels of about 10 kg/m3 / 3.4 lbs/bbl and 3 kg/m3 /
1.0 lbs/bbl respectively. Depending on the application,
Cabot might add more buffer to the brine before it is
shipped to the field. Potassium formate brine as delivered
from the suppliers normally contains none or lower
amounts of buffer, typically up to 2.5 kg/m3 / 1 lbs/bbl
of potassium carbonate or bicarbonate, and pH can vary
considerably. Potassium and sodium formate supplied
in solid form (powder) typically contain large amounts of
carbonate, which has been added as anti-caking agent.
Such material, when dissolved in water, often exhibits a
high pH and does not normally require further buffering.
Depending on planned use, pH might need to be brought
down before such material is used in the field.
A buffer level of about 17 to 34 kg/m3 / 6 to 12 lbs/bbl
of sodium and / or potassium carbonate / bicarbonate
is recommended for most formate brine applications.
Cabot normally buffers formate brines to a pH of around
10.0 – 10.5 (measured with 1:10 dilution with deionized
water). The amounts of carbonate and bicarbonate
required to achieve this depend on the carbonate and
bicarbonate levels already in the brine. The graph in
Figure 8 shows expected pH for various sodium and
potassium carbonate / bicarbonate additions. However,
first one always needs to consider the amount of buffer
already in the brine.
A pH of 10 – 10.5 and a buffer level of
about 17 to 34 kg/m3 / 6 to 12 lbs/bbl sodium or
potassium carbonate / bicarbonate is ideal for
most formate brine applications. The use of cesium
carbonate / bicarbonate might be beneficial in
certain high-density cesium formate single-salt
formulations to achieve higher density.
A6.3.5 Maintaining buffer concentration and capacity
In order to get the full benefit of the carbonate /
bicarbonate buffer in formate brine, both the carbonate
concentration and total buffer concentration should
be maintained during field use. These concentrations
can easily be determined by the simple field method
described in Section A6.3.4 above.
V ERSION
4
–
0 9/ 13
C A B O T
In most field applications, the most practical way to
control pH and maintain buffer capacity is by adding
carbonate. This method has the advantage that the
consequences of over-treatment are not as severe
as those from KOH. A potential disadvantage with this
method, however, is that it allows the concentration
of bicarbonate to build up. Excessive concentrations
of bicarbonate are known to cause rheology and fluidloss problems in water-based muds. This has also
been experienced in formate-based muds [8]. A good
indication that the total buffer concentration is getting
low and addition of carbonate is required is that pH drops
quickly after it has been adjusted upwards with KOH.
References
[1] Leth-Olsen, H.: “CO2 Corrosion of Steels in Formate
Brines for Well Applications”, 2004 NACE, paper
# 04357, New Orleans, USA, March 2004.
[2] Prasek, B.B. et al: “A New Industry Standard for
Determining the pH in Oilfield Completion Brines,”
Paper # SPE 86502, Lafayette, LA, February 2004.
[3] Javora P.H. et al: “A New Technical Standard for
Testing of Heavy Brines”, paper # SPE 98398,
Lafayette, LA, February 2006.
[4] “Dilution factors for accurate measurement
of formate brine pH”, Cabot laboratory report
# LR-050, April 2004.
[5] API RP 13B-1: “Standard Procedures for Field Testing
Water-Based Drilling Fluids”.
[6] “Potassium Formate Titration curve using KCOOH”,
report # LR-289, Cabot Operations and Technical
Support Laboratory, Aberdeen, UK, February 2009.
[7] “Calibration titration for buffer determinations”,
report # LR-294, Cabot Operations and Technical
Support Laboratory, Aberdeen, UK, March 2009.
[8] Berg, P.C., et al.: “Drilling, Completion, and Openhole
Formation Evaluation of High-Angle Wells in HighDensity Cesium Formate Brine: The Kvitebjørn
Experience, 2004 – 2006,” SPE 105733, Amsterdam,
February 2007.
SECTION A6
PAGE 11