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 CO

2 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 To check if a newer version of this section exists please visit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

F O R M A T E T E C H N I C A L M A N U A L

SECTION A6PH AND BUFFERING

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 5 –10/14

C A B O T

C H E M I C A L A N D P H Y S I C A L P R O P E R T I E S

http://www.formatebrines.com/manual

P A G E 2 V E R S I O N 5 – 1 0 / 1 4


S E C T I O N A 6

C A B O T

pH

Addition of strong acid

pH behavior of unbuffered formate brines

0

2

4

6

8

10

12

14

pKa = 3.75 50% formate 50% formic acid

No formic acid

Traces of formic acid

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+]:

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha

OHHCOOHOHHCOO aK 23 +→←+−

−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK

323 COHHHCO ←+ +−

1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←

( ) ( )aqCOHOHaqCO 3222 →←+

( ) ( ) ( )aqHaqHCOaqCOH +− +332

−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322

3 sCaCOaqCaaqCO ↓→+ +−

( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=

[ ]−+−−= 3logloglog2

HCOPKpH CO−OH −2

3CO −

3HCO [ ]−2

3CO [ ]−

3HCO 42SOH [ ]−OH [CO

[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02

3 LmolPOHCO f×=+ −−

[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(1)

In more concentrated solutions, the behavior of the ions in the solution depends not on their concentrations, but on activities. Thus, in reality, a more precise definition is:1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(2)where

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

is the activity.

Commonly used high-density oilfield brines (CaCl2,

CaBr2, and ZnBr

2) have a naturally acidic pH. Attempts

to raise the pH to alkaline levels in these halide-based brines can result in precipitation of insoluble calcium or zinc salts, e.g. Ca(OH)

2, Zn(OH)

2.

Formate salts dissolved in water exhibit a naturally alkaline pH (8 – 10). The pH of the formate brines can be adjusted to almost any level with common acids and bases without causing the precipitation of insoluble salts. The pH of fluids based on formate brines can therefore be safely adjusted to the level that delivers the optimal performance.

The formate ion is a buffer in itself, and formate brines have a natural buffer capacity at pH = 3.75:

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(3)pKa = 3.75

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 approximately ten, as shown in Table 1. This means that in concentrated cesium formate brine with a formate concentration of around 10 mol/L and a pH of around 10 – 10.5, the concentration of formic acid is less than 0.000001 mol/L. Figure 1 shows how pH of unbuffered formate brine changes with addition of a strong acid.

Table 1 ‘Theoretical’ formate / formic acid molar ratio as a function of pH.

pH Approx. formate / formic acid molar ratio

10.75 10,000,000

9.75 1,000,000

8.75 100,000

7.75 10,000

6.75 1 000

5.75 100

4.75 10

3.75 1

2.75 0.1

1.75 0.01

0.75 0.001

The pKa value in formate brines has been shown to increase with temperature [1]. In very concentrated brines, pH (and thereby pKa) are poorly defined.

Figure 1 Graph shows how the pH of unbuffered formate brine changes with the addition of a strong acid.

S E C T I O N A : C H E M I C A L A N D P H Y S I C A L P R O P E R T I E S

V E R S I O N 5 – 1 0 / 1 4 P A G E 3S E C T I O N A 6

C A B O T

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 high-density 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.

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

P A G E 4 V E R S I O N 5 – 1 0 / 1 4


S E C T I O N A 6

C A B O T

6.0

7.0

8.0

9.0

10.0

11 .0

12.0

13.0

14.0

15.0

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16

10.0

11 .0

12.0

13.0

14.0

15.0

0 1 2 3 4 5 6 7 8 9 1 0 1 1 1 2 1 3 1 4 1 5 1 6

KFo 1.56 g/cm3 unbuffered (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)

CsFo 2.2 g/cm3 buffered (glass electrode)

CsFo 2.2 g/cm3 buffered (pH paper)

CsKFo 2.0 g/cm3 buffered (glass electrode)

CsKFo 2.0 g/cm3 buffered (pH paper)

pH measurements in buffered and unbuffered potassium formate brine

pH measurements in buffered cesium formate brine and buffered cesium / potassium formate brine

Dilution factor

Dilution factor

pH

pH

Figure 3 Effect of dilution when measuring pH in buffered 2.2 g/cm3 / 18.3 lb/gal cesium formate brine and buffered 2.0 g/cm3 / 16.7 lb/gal cesium / potassium formate brine.

Figure 2 Effect of dilution when measuring pH in buffered and unbuffered 1.56 g/cm3 / 13.0 lb/gal potassium formate brine.



C A B O T

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 can be made in formate brines both by pH electrode (potentiometric measurements) and by use of pH paper, although the pH electrode method is more accurate. Due to the significant difference between pH measured in neat formate brines and in diluted formate brines, it is important to always record whether the measurement was made in diluted or neat brine.

If pH is measured in neat brine, one would clearly also have to report the method of measurement (pH paper vs. glass electrode) and the brine type and concentration, otherwise the measured pH value is meaningless.

A6.2 pH buffering of formate brines with carbonate / bicarbonate buffer

Formate brines used in oilfield applications should be buffered by addition of potassium or sodium carbonate and potassium or sodium bicarbonate. The main purpose of this buffer is to provide an alkaline pH and to prevent the pH from fluctuating as a consequence of acid or base influxes into the brine.

Maintaining an alkaline pH environment in formate brines with carbonate / bicarbonate buffer is important for the following reasons:

• Alkaline pH helps control corrosion (see Section B6)• Presence of carbonate / bicarbonate provides special

protection against CO2 corrosion (see Section B6)

• Alkaline pH helps lower formate decomposition rates (see Section A13)

• Presence of carbonate / bicarbonate helps limit amount of formate decomposition (see Section A13)

• Alkaline pH helps stabilize polymers and other additives (see Section B5)

• Presence of carbonate lowers the risk of H2S gas

release (see Section A6.2.3 on next page)• Presence of carbonate improves well control by

sequestering influx of CO2 (see Section A6.2.2 on

this page)

The main reason for loss of pH control in oilfield fluids is the influx of acid gases such as CO

2 and H

2S. These are both

weak acids with a pKa higher than the pKa of formic acid.

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–).

The carbonate / bicarbonate buffer system provides strong buffering at two different pH levels:

• Upper buffering level at pH = 10.2

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(4)

where

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

= 10.2

At pH = 10.2 (

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

) the buffered solution contains the same amount of carbonate (

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

) and bicarbonate (

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

).

• Lower buffering level at pH = 6.35

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(5)

where

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

= 6.35

At pH = 6.35 (

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

) the buffered solution contains the same amount of bicarbonate (

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

) and carbonic acid (

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

).

The exact levels of

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

and

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

will vary somewhat with temperature and pressure.

Figure 4 demonstrates how a pure carbonate buffer works in water when a strong acid is added. The carbonate reacts with added acid until all carbonate is consumed. As long as there is still carbonate left in solution, pH remains high around the ‘higher buffer level’ (10.2±1). As soon as the carbonate is consumed, pH drops rather quickly down to the ‘lower buffer level’ where it remains as long as bicarbonate is available to react with the added acid for conversion to carbonic acid. In order for pH to drop below this second buffer level, an acid needs to be added that is stronger than the carbonic acid, which is formed. As any CO

2 gas influx into

the buffered solution dissolves and converts to carbonic acid, a CO

2 influx is therefore not capable of pulling the

pH much below this second buffer level.

A6.2.2 Buffer protection against CO2 influx

The major cause of acidification of conventional completion brines is influx or diffusion of carbon dioxide gas (CO

2) into the wellbore from the surrounding rock

formations:

P A G E 6 V E R S I O N 5 – 1 0 / 1 4


S E C T I O N A 6

C A B O T

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(6)

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(7)

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(8)

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

= 6.35

Depending on the original pH of the receiving brine system, dissolved CO

2 remains in the brine as either

carbonic acid (H2CO

3) in equilibrium with dissolved CO

2

gas or bicarbonate (HCO3–), according to reaction 8. This

is demonstrated in Figure 5. As more CO2 gas enters into

the brine, carbonic acid concentration builds up and pH drops and allows unbuffered brines to acidify.

The three different brine systems in Figure 5 react in the following ways to a CO

2 influx:

• Conventional divalent halide brines can not be buffered with carbonate / bicarbonate because the corresponding metal carbonate (CaCO

3, ZnCO

3)

precipitates out of solution resulting in formation of solids in the clear packer / completion fluid. These divalent brines have a naturally low pH (2 – 6), and the influx of CO

2, dependent on the partial pressure of

CO2, further lowers the pH. The CO

2 partly converts to

carbonic acid (Equation 7), which is very corrosive.

• Buffered formate brines are capable of buffering large amounts of CO

2. Unless the influx is unusually

large, the brine maintains a pH around the upper buffer level (pH = 10.2), which is high enough to prevent

carbonic acid being present in the brine. With a large influx of CO

2, the pH drops down to the lower buffering

level (pH = 6.35) where it stabilizes. Measurements of pH in formate brines exposed to various amounts of CO

2 have confirmed that pH never drops below 6 – 6.5.

This pH is still close to neutral, meaning that this brine system cannot be ‘acidified’ to any great extent by exposure to CO

2. However, carbonic acid and a small

amount of formic acid are also present.• Unbuffered formate brines: The pH of these brine systems responds in a similar fashion to halide brines when exposed to CO

2 gas. However, they do have a

higher initial pH, and the pH drop will be limited as the formate brine is a buffer in itself (pKa = 3.75). At such low pH a significant amount of corrosive formic acid is present in the fluid. If there is any chance of an acid gas influx, the use of unbuffered formate brines is highly discouraged.

A6.2.3 Buffer protection against H2S influx

Influx of CO2 into a wellbore is often accompanied by

hydrogen sulfide (H2S). H

2S is a weak acid with a

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

of around 7. This means that at a pH of 7, equal amounts of hydrogen sulfide (H

2S) and hydrogen bisulfide

(HS–) will be present in the brine. At higher pH, more HS– will be present and at lower pH more H

2S will exist.

Therefore, unless the carbonate buffer in the formate brine is overwhelmed by large influxes of CO

2, the

carbonate buffer traps and retains this toxic gas in its less harmful form, namely bisulfide, HS–.

The fact that any H2S is converted to HS– in buffered

formate brines does not mean that the gas is

Figure 4 The pH in water buffered with carbonate as a function of added acid (H+). The x-axis shows the fraction of the buffer that is consumed by the added acid. As can be seen, carbonate buffers twice, first at pH = pKa2

= 10.2 (upper buffer level) and then at pH = pKa1

= 6.35 (lower buffer level). If the added acid is carbonic acid (from CO2 influx), the pH

can never drop much lower than pKa1 = 6.35.

pH behavior of carbonate / bicarbonate buffer when adding strong acid

Addition of strong acid

Fraction of buffer consumed

3

4

5

6

7

8

9

10

11

12

0 0.2 0.4 0.6 0.8 1 1 .2 1 .4 1.6 1 .8 2

pH

pKa2

pKa1



C A B O T

scavenged and made permanently safe. If the buffer was to be overwhelmed by an excessive influx of CO

2 /

H2S, then H

2S gas would come back out of solution

when pH dropped to below around 7.0. CO2 gas would

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 muds, and never lower the pH or let the buffer deplete in a formate mud or brine that has been exposed to H

2S

without first checking if it is contaminated with HS–. If there is any concern about H

2S-related corrosion then H

2S

scavenger should be added (see Section B6, Compatibility with Metals and Section B5, Compatibility with Additives).

A6.3 Buffer addition and maintainance

Whenever formate brine is used in the field, it is important to maintain the ability of the buffer to resist acid influxes. In order to do this, both buffer capacity and total buffer concentration need to be monitored and maintained.

A6.3.1 Buffer capacity

In buffered formate brine, it is the carbonate component of the buffer that provides buffering at the alkaline pH of 10.2. Bicarbonate is mainly added in order to balance alkalinity of the carbonate as pH is a function of the carbonate-to-bicarbonate ratio. The carbonate concentration alone is therefore the true measure of the brine’s buffer capacity.

The scientifically correct definition of buffer capacity is: “The number of moles of acid or base necessary to change the pH of one liter of solution with one unit”.

As can be seen from the graph in Figure 5, the decrease in pH of one unit is not really a good measure of how much carbonate buffer is present in the brine, and therefore the true capacity of this buffer. Cabot therefore uses the actual carbonate concentration as a measure of the capacity of the buffer rather than the scientifically defined ‘buffer capacity’.

In alkaline brines that are buffered with carbonate / bicarbonate buffer, the following equilibrium exists between carbonate and bicarbonate:

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(9)

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

= 10.2

In the field, carbonate is typically lost by exposure to influx of acid gas. As acid gas initially enters the brine, carbonate (CO

32–) is gradually converted to

bicarbonate (HCO3

–), whilst pH remains at around the upper buffer level (pH =

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

= 10.2). When all carbonate is converted, the buffer loses its ability to maintain pH. The carbonate component of the buffer system is now referred to as ‘overwhelmed’ or ‘swamped’ and has no more capacity to buffer pH at the upper buffering level. Any further influx of acid gas can now easily lower pH down to the lower buffer level (

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

= 6.35) provided by the bicarbonate. (See Figure 5.)

It is important to notice that whilst pH of a buffered formate brine is a function of the ratio of concentrations of carbonate and bicarbonate, the capacity of the buffer to maintain pH around 10 – 10.5 depends upon the actual carbonate concentration.

Figure 5 pH as a function of CO2 influx in a typical halide brine, an unbuffered formate brine, and a buffered formate brine.

4

5

6

7

8

9

10

11

12

0 50 100 150 200 250 300 350 400 450 500BBL gas influx / BBL buffered formate brine (2% CO2 21°C / 70°F, 1 atm)

pH

Increasing time of CO2 influx

pH in various brine systems as a function of CO2 influx volume

Buffered formate brine

Unbuffered formate brine

Calcium bromide brine

pH>6.35: CO

2 mainly converted to

bicarbonate (HCO3-),

which does not promote corrosion

pH<6.35: CO

2 mainly converted to

carbonic acid (H2CO

3),

which promotes corrosion

P A G E 8 V E R S I O N 5 – 1 0 / 1 4


S E C T I O N A 6

C A B O T

A6.3.2 Total buffer concentration

The total buffer concentration in a brine that is buffered with carbonate / bicarbonate is defined as the combined concentration of carbonate (CO

32–),

bicarbonate (HCO3

–), carbonic acid (H2CO

3), and any

carbon dioxide (CO2) gas dissolved in the brine. If the

total buffer concentration has been removed from the fluid in other reactions, new buffer will need to be added to the fluid.

There are three ways in which the buffer concentration in a formate brine can be altered during field use:

1. Influx of acid gas (CO2) increases buffer

concentration:

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(10)

Influx of CO2 converts carbonate from the buffer to

bicarbonate. The carbonate concentration (buffer capacity) therefore drops whilst the total buffer concentration (carbonate + bicarbonate + carbonic acid + dissolved CO

2) increases by the amount of CO

2

entering the brine.

2. Influx of multivalent cations decreases buffer concentration:

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(11)

An influx of multivalent cations consumes the buffer by precipitating out insoluble calcium carbonate.

The total amount of carbonate / bicarbonate buffer available decreases by the amount of carbonate that is precipitated.

3. Formate decomposition increases buffer concentration Small amounts of soluble carbonate and bicarbonate can form as a result of formate decomposition if the brine is exposed to high temperature for an extended period of time (See Section A13, Thermal Stability). The dominant decarboxylation reaction is reversible, and the establish ment of equilibrium in closed HPHT well systems usually limits formate decomposition to a few percent in typical formate brine formulations.

A6.3.3 Determining buffer concentration and capacity

For standard water-based mud filtrates, API RP 13B-1 [5] recommends that carbonate and bicarbonate content are measured by pH titrations. Alkalinity in the form of

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

content,

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

content, and

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

content is determined by the combination of a phenolphthalein titration to an endpoint of pH = 8.2, and a methyl orange titration to an endpoint of pH = 3.1.

In formate brines, the determination of the methyl orange titration endpoint is complicated by the formate / formic acid equilibrium that is present at pH = 3.75 (

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

formic acid). This is illustrated in Figure 6. As can be seen, the formate / formic acid equilibrium starts to establish at a pH significantly higher than the methyl orange endpoint. The fact that only one of the two standard titration endpoints can be determined in a

Figure 6 Titration curves for buffered water and buffered potassium formate. Both fluids contain the same amount of added carbonate / bicarbonate buffer (17.8 kg/m3 / 6.25 lbs/bbl K

2CO

3 and 10.7 kg/m3 / 3.75 lbs/bbl KHCO

3). The

phenolphthalein and methyl orange endpoints from the standard API RP 13B-1 alkalinity titrations are shown. No methyl orange endpoint can be detected in the buffered formate brine due to the formate / formic acid equilibrium starting to establish at a higher pH.

Titration curves

2

3

4

5

6

7

8

9

10

11

12

13

-0.3 -0.2 -0.1 0.0 0.1 0.2 0.3 0.4 0.5

pH

Added H+ [mol/L]Added OH- [mol/L]

Phenolphthalein endpointMethyl orange endpointWater + standard bufferKFo + standard buffer



C A B O T

formate brine, means that the standard API alkalinity test method is unsuitable for determining carbonate and bicarbonate concentrations in this brine system.

The consequences of this are:- The phenolphthalein endpoint can be used to

determine the combined hydroxide and carbonate concentration [

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

]+ [

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

].- The methyl orange endpoint cannot be used to

determine the total buffer concentration (including

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

). Even if it is possible to detect an endpoint in a very diluted formate brine, the use of this endpoint would give erratic calculations.

Another method is therefore required for measuring the bicarbonate concentration. One way of doing this is by using a Garrett Gas Train (GGT). The GGT determines the total amount of carbonate and bicarbonate. The method is time consuming and therefore not very popular on the rig, and does not differentiate between the different buffer components.

Carbonate and bicarbonate concentrations have been measured by Cabot using a dual-titration method. Recently a simpler and more accurate method, based on simple pH measurement and phenolphthalein titration, has been developed.

A simple field or laboratory method for determining buffer concentrationLaboratory testing of formate brines with known additions of carbonate and bicarbonate has shown that pH of buffered formate brines is dependent on the carbonate-to-bicarbonate ratio [6], [7]. The following relationship, R, has been found between the carbonate and bicarbonate molar ratio and brine pH:

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA ×R =

HCO

CO×=−

−

(12)

where

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

and [CO32–‐] and [HCO

3–] are the molar concentrations of

carbonate and bicarbonate. This relationship is shown in Figure 7 and is valid for pH measured with a glass electrode in formate brine diluted with nine parts deionized water. This relationship can be used to determine buffer concentration and buffer capacity of buffered formate brines in the field. This means that both carbonate and bicarbonate concentrations can be determined just by measuring pH and performing the standard phenolphthalein titration. The method is as follows:

1. Prepare a sample consisting of 5 mL fluid sample (brine or mud filtrate) and 45 mL deionized water.

2. Measure pH of this sample with a calibrated glass electrode.

3. Perform a titration to pH = 8.2 with 0.02N HCl or H

2SO

4 and report the phenolphthalein endpoint p

f

as the volume, V (mL), of titrant required per mL of fluid sample:

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(13) Depending on pH, four situations exist:1. pH > 11.1

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(14)

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(15) Unless large amounts of [

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

] have been added to this fluid, one can assume that most of this alkalinity is from carbonate.

2. pH = 11.1

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(16)

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(17)

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(18)

3. 9.0 < pH < 11.1

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(19)

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(20) From the carbonate / bicarbonate pH relationship

in Equation 12 and Figure 6, determine the carbonate / bicarbonate ratio, R. Calculate the bicarbonate concentration as:

[HCO3

–] = [CO3

2–] (mol/L) / R (21) 4. pH < = 9.0

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

(22) [ ] insignificant2

3 =−CO insignificant

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

difficult to determine pH needs to be adjusted to above 9.0 with

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

18

19

20

21

22

23

24

25

17

16

26

27

[ ]+−= HpH log

( )+−=H

apH log +Ha


−+

+

− →←+ 3

2

32 HCOHCO apK

→1apK

2apK

apK


1apK

−2

3CO −3HCO

−3HCO 32 COH

( ( )aqCOgCO 22 →←



−+− →←+ 3

2

3 HCOHCO

OHCOOHHCO 2

2

33 +→←+−−−

−− →++ 322

2

3 2HCOOHCOCO

)()()( 322


( ) ( )aqHHCOOHgCO K +−+→←+ 322

[ ] [ ]

2

3

COP

HCOHK

−+ ×=



3CO −

3HCO [ ]−2

3CO [ ]−


[ ]fPLmgCO ×=− 1200)/(2

3

fP×02.0

[ ]fPmkgCO ×=− 2.1)/( 32

3

[ ]fPCO ×=− 42.0)(ppb2

3[ ]−2

3CO

)

[ ] [ ] )/(02.02


[ ] 03 =−

HCO

[ ] 0=−OH

[ ] 03 =−

HCO

[ ] 0==−OH

[ ] )/(02.02

3 LmolPCO f

)/5(mLpf

×=−

[ ]]CO

)/()exp

[

(3

3

2

Lmol

Vol

pHBAHCO ..=

−−

3.894 • 10-10=A

2.193=B

[ ] 0=−OH

28 [ ] 02

3 =−CO

−OH −2

3CO −

3HCO apK

[ ]−OH [ ]−2

3CO [ ]−

3HCOfP 42SOH

)exp(] (mol/L)[

] (mol/L)[

3

23 pHBA x

HCO

CO×=

−

−

[HCO3

-] = [CO3

2-]/R (mol/L)

→← 1aK

2aK

before the bicarbonate concentration can be determined.

The equivalent amounts of sodium or potassium carbonate and bicarbonate can easily be calculated from these molar concentrations. This simple field method is explained in more detail in Section C2.

A6.3.4 Buffer requirement for field use

The recommended buffer concentration required in formate brines depends on the application. The amount of time the brine will be in contact with the reservoir fluids, and the expected level of acid gas influx, are important factors. In well suspension and packer applications where the formate brine may be

P A G E 1 0 V E R S I O N 5 – 1 0 / 1 4


S E C T I O N A 6

C A B O T

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.

Adding only soluble carbonate to the brine provides a high level of buffer capacity, but the pH might become

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-to-bicarbonate ratio. (See Equation 12 and Figure 7).

When determining buffer levels for field applications, one needs to consider that some oilfield formate

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.

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 [CO

32–] to [HCO

3–] (see figure 6 above),

b) equivalent potassium carbonate (K2CO

3) to potassium bicarbonate (KHCO

3) ratio (wt/wt), and c) equivalent sodium

carbonate (Na2CO

3) to sodium bicarbonate (NaHCO

3) ratio (wt/wt).

0

2

4

6

8

10

12

14

16

18

9 9.5 10 10.5 11pH

H2ONaFoKFoKFo (non-analytical)CsKFo[CO3

2-]/[HCO3-] (mol/mol)

R =

[C

O3

2- ]/

[HC

O3

- ]

0

2

4

6

8

10

12

14

16

18

9 9.5 10 10.5 11pH

R =

[C

O3

2- ]/

[HC

O3

- ]

[CO32-]/[HCO3

-] (mol/mol)K2O3/KHCO3 (wt/wt)Na2CO3/NaHCO3 (wt/wt)


V E R S I O N 5 – 1 0 / 1 4 P A G E 1 1S E C T I O N A 6

C A B O T

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 CO

32– 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.

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 fluid-loss 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 High-Density Cesium Formate Brine: The Kvitebjørn Experience, 2004 – 2006,” SPE 105733, Amsterdam, February 2007.

section a6 ph and buffering

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