formate man a6 ph and buffering

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P A G E 1S E C T I O N A 6

Section A6

pH and Buffering

NOTICE AND DISCLAIMER. The data and conclusions contained herein are based on work believed to be reliable; however, CABOT cannotand 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. CABOTDISCLAIMS 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 CABOTdoes not accept and hereby disclaims liability for, any damages whatsoever in connection with the use of or reliance on this information or anyproduct to which it relates.

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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 .............................................................5

A6.2.1 How the carbonate / bicarbonate buffer works ............................. 5

A6.2.2 Buffer protection against CO 2 influx ............................................... 6

A6.2.3 Buffer protection against H 2 S influx ................................................6

A6.3 Buffering requirement and buffer capacity ..........................................7

A6.3.1 Buffer capacity ............................................................................... 7

A6.3.2 Buffer concentration ....................................................................... 7

A6.3.3 Buffer requirement for field use ......................................................8

A6.3.4 Determining buffer concentration and capacity ................................ 8

A6.3.5 Maintaining buffer concentration and capacity ................................. 10

References ..................................................................................................... 10

CHEMICAL AND PHYSICAL PROPERTIES

F O R M A T E T E C H N I C A L M A N U A LC A B O T S P E C I A L T Y F L U I D S

V E R S I O N 1 – 1 2 / 0 7 P A G E 1S E C T I O N A 6



C A B O T S P E C I A L T Y F L U I D S

P A G E 2 V E R S I O N 1 – 1 2 / 0 7

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

S E C T I O N A 6

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 + ]:

[ ]+−= H pH log

−

−

−

−

−

−

− −

−

−

(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:

( )+−=H

a pH log

−

−

−

−

−

−

− −

−

(2)

where +H a

−

−

−

−

− −

is the activity.

Commonly used high-density oilfield brines (CaCl 2 ,CaBr

2 , 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 therefore have a natural buffering capacity at

low pH:

OHHCOOHOHHCOO aK

23 + ⎯→←+− +

−

−

−

−

−

−

−

−

(3)

pK a = 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. Figure 1 shows how the pH of unbuffered formate

brine changes with the addition of a strong acid.

Table 1 ‘Theoretical’ formate / formic acid molar ratio

as a function of pH

pH Approx. formate / formic acid molar ratio6.75 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.

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 offormic acid

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




V E R S I O N 1 – 1 2 / 0 7

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


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 a

carbonate / bicarbonate buffer. Buffers are

designed to resist changes in fluid pH and cancope with large influxes of acid gas.

A6.1.2 Measuring pH in formate brines

pH is a measure of the hydrogen ion activity of a

solution. Because pH is dependent on activity, a

property that cannot be measured easily or

predicted theoretically, it is difficult to determine an

accurate value for the pH of very concentrated

solutions, such as high-density brines. Consequently,

any pH values that are measured in neat formate

brines are not a true measure of hydrogen ion

activity and should only be used in a relative sense.

pH can be measured in formate brines by potentio-

metric measurements or with pH paper. In both

cases, dilution of the fluid with nine parts (vol/vol)

deionized water is recommended in order to get

more accurate readings. Please notice that this

method is not recommended in concentrated

halide brines because interference with the divalent

cations (Ca 2 + and Zn

2 + ) causes unpredictable pH

changes with dilution [2][3].

Some examples of pH measurements with

glass electrode and pH papers in buffered and

unbuffered formate brines as a function of dilution

are shown in Figure 2 and Figure 3. The pH papers

were BDF pH indicator sticks. These sticks

indicated a pH of 14 in all of the undiluted buffered

formate brines. As these brines were all buffered

with both carbonate and bicarbonate (6.25 ppb

K 2 CO

3 and 3.75 ppb KHCO

3 ), we know that the real

pH in these systems should be close to the pKa2 of

this buffer system (i.e. 10.2). This is very much in

agreement with the pH levels that are measured in

the diluted brines. It is evident that measuring pH

directly in undiluted formate brine with pH paper

can result in errors of up to four pH units. The data

also shows that measuring pH in undiluted formatebrines with a glass pH electrode can lead to errors

of up to one pH unit in the alkaline direction.




P A G E 4 V E R S I O N 1 – 1 2 / 0 7


S E C T I O N A 6

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

KFo 1.56 s.g.unbuffered (glass electrode)

KFo 1.56 s.g. unbuffered (pH paper)

KFo 1.561 s.g. buffered (glass electrode)KFo 1.561 s.g. buffered (pH paper)

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

CsFo 2.199 s.g. buffered (glass electrode)

CsFo 2.199 s.g. buffered (pH paper)

CsKFo 2.0 s.g. buffered (glass electrode)

CsKFo 2.0 s.g. buffered (pH paper)

pH measurements in buffered and unbuffered potassium formate brine

pH measurement in buffered CsFo and CsKFo brines

Dilution factor

Dilution factor

pH

pH

Figure 3 Effect of dilution when measuring pH in buffered 2.0 s.g. / 16.7 ppg CsKFo brine and buffered 2.2 s.g. / 18.3 ppg CsFo brine.

Figure 2 Effect of dilution when measuring pH in buffered 2.0 s.g. / 16.7 ppg CsKFo brine and buffered 2.2 s.g. / 18.3 ppg CsFo brine.




P A G E 6 V E R S I O N 1 – 1 2 / 0 7


S E C T I O N A 6

A6.2.2 Buffer protection against

CO 2 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:

−

−

( ( ) aqCO gCO 2 2 ⎯→←

−

−

−

−

− −

) (6)

−

−

( ) ( ) aqCOHOH aqCO 3 2 2 2 ⎯→←+

−

−

−

−

− −

(7)

−

−

( ) ( ) ( ) aqH aqHCO aqCOH aK +−

+ ⎯ ⎯→← 33 21

−

−

−

−

− −

(8)

1 a pK

−

−

= 6.35

Depending on the original pH of the receiving brine

system, the dissolved CO 2 remains in the brine as

either carbonic acid (H 2 CO

3 ) or bicarbonate (HCO

3 - ),

according to reaction 8. This is demonstrated in

Figure 5. As more CO 2 gas enters into the brine,

the carbonic acid concentration builds up and the

pH drops and allows unbuffered brines to acidify.

The three different brine systems in Figure 5 willreact 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 the 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 CO 2 , further lowers the pH. The CO

2 largely

converts to carbonic acid, 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 the 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 .

• 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). If there is any chance of an acid

gas influx, the use of unbuffered formate brines is

not recommended.

A6.2.3 Buffer protection against

H 2 S influx

Influx of CO 2 into a wellbore is often accompaniedby hydrogen sulfide (H

2 S ). H

2 S is a weak acid with a

−

−

1 a pK of around 7. Unless the buffer 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 H 2 S is converted to HS - in a

buffered formate brine does not mean that the gas

is scavenged and made permanently safe. If the

buffer was to be overwhelmed by an excessive

influx of CO 2 / H

2 S , then H

2 S gas would come back

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 = pK a2 = 10.2 (upper buffer level) and then at

pH = pK a1 = 6.35 (lower buffer level). In the case that the added acid is carbonic acid (from CO 2 influx), the pH can never drop muchlower than pK a1

= 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




V E R S I O N 1 – 1 2 / 0 7



out of solution when pH dropped to below around

7.0. CO 2 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 2 S without

first checking if it is contaminated withHS - . If there

is any concern about H 2 S -related corrosion then

H 2 S scavenger should be added (see Section B6

‘Compatibility with metals’ and Section B5

‘Compatibility with additives’).

A6.3 Buffering requirement andbuffer capacity

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 the

buffer capacity and the total buffer concentration

need to be monitored and maintained.

A6.3.1 Buffer capacity

Buffer capacity is defined as the moles of acid orbase necessary to change the pH of one liter of

solution by one unit.

In alkaline brines that are buffered with a carbonate /

bicarbonate buffer, the following equilibrium exists:

−

−

−+−

⎯ ⎯→←+ 3

2

3 2 HCOHCO aK

−

(9)

2 a pK = 10.2

In the field, buffer capacity is typically lost by exposure

to influx of acid gas. As acid gas initially enters the

brine, the carbonate (CO 3 2- ) is gradually converted

to bicarbonate (HCO 3 - ), whilst the pH remains at

around the upper buffer level (pH =

−

2 a pK

−

−

−

−

−

= 10.2).

When all the 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 the buffer

capacity at the upper buffering level is zero. Any

further influx of acid gas can now easily lower the

pH down to the lower buffer level (

−

−

1 a pK

−

−

−

−

= 6.35)

provided by the bicarbonate. (See Figure 5.)

It is important to notice that whilst the pH of a

buffered formate brine is a function of the ratio of

concentrations of carbonate and bicarbonate, the

buffer capacity depends upon the actual carbonate

concentration.

Whenever the buffer capacity in the field is getting

low, it can be increased by a simple addition of OH -

(NaOH or KOH ):

−

−

OHCOOHHCO 2

2

33 + ⎯→←+ −−−

−

−

−

−

(10)

In other words, it is unnecessary to add new

carbonate buffer in order to regain buffer capacity.

A6.3.2 Buffer concentration

The total buffer concentration in a brine that is buff-

ered with carbonate / bicarbonate is defined as the

combined concentration of carbonate (CO 3 2- ),

bicarbonate (HCO 3 - ), and carbonic acid (H

2 CO

3 ).

If the total buffer concentration has been removed

from the fluid in other reactions, new buffer will

need to be added to the fluid.

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

1 0

1 1

1 2

0 5 0 1 0 0 1 5 0 2 0 0 2 5 0 3 0 0 3 5 0 4 0 0 4 5 0 5 0 0

BBL gas influx / BB L b u ff er ed fo r m a te b r i n e ( 2% CO2, 21°C / 70° F , 1 at m )

pH

Increasing time of CO 2 influx

pH in various brine systems as a function of CO 2 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 ( H 2CO

3 ),

which promotes corrosion




P A G E 8 V E R S I O N 1 – 1 2 / 0 7


S E C T I O N A 6

There are three ways in which the buffer concentration

in a formate brine can be altered during field use:

1. Inux of acid gas (CO 2 ) increases buffer

concentration:

−

−

−−

⎯→ ⎯ ++ 3 2 2

2

3 2HCOOHCOCO

−

−

−

−

− −

−

−

−

(11)

Influx of CO 2 converts carbonate from the buffer to

bicarbonate. The buffer capacity therefore drops

whilst the total buffer concentration (carbonate +

bicarbonate + carbonic acid) increases by the

amount of CO 2 entering the brine.

2. Inux of multivalent cations decreases

buffer concentration:

−

−

)( )( )( 3

2 2

3 sCaCO aqCa aqCO ↓ ⎯→ ⎯ + +−

−

−

−

−

− −

−

−

(12)

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, and new buffer should be

added to the fluid.

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 reaction is reversible and the

establishment of equilibrium in closed HPHT well

systems usually limits formate decomposition to a

few percent in typical formate brine formulations.

Formate brines that are heavily buffered with

carbonate / bicarbonate equilibrate faster and

decompose less.

A6.3.3 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 exposed to well conditions for a long time,

a high level of buffering would be appropriate. In

applications where well exposure times are short,

and in applications where no acid gas is expected,

a smaller buffer concentration will do.

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 the

addition of bicarbonate does not contribute to the

initial buffer capacity it will contribute positively to

the final pH of the brine if the buffer is overwhelmed

by acid gas influx, and it will contribute positively to the fluid’s thermal stability.

It has been shown that the pH of buffered formate

brine is dependent on the ratio between carbonate

and bicarbonate according to the following equation:

−

−

−

−

−

−

)exp( ] [

] [

3

2

3 pHB AHCO

CO×=

−

−

(13)

where

−

−

−

−

−

−

−

111069164.9 − = A

−

−

−

−

−

−

−

−

−

29762135. 2=B

−

−

Figure 6 Relationship between the carbonate to bicarbonate ratio and the pH in a formate brine. The pH is measured with a calibrated

glass electrode in a variety of laboratory and field brines, diluted with nine parts water.

0

1

2

3

4

5

6

7

8

9

10

11

7 7.5 8 8.5 9 9.5 10 10.5 11pH

C a r b o n a t e / b i c a r b o n a t e m o l a r r a t i o

exp( ] [

] [

3

2

3

HCO

CO ) pH×=−

−

·11

1069164.9 −

29762135. 2




V E R S I O N 1 – 1 2 / 0 7


P A G E 1 1S E C T I O N A 6

A6.3.5 Maintaining buffer concentration

and capacity

In order to get the full benefit of the carbonate /

bicarbonate buffer in the formate brine, both the

buffer concentration and the buffer capacity should

be maintained during field use.

Buffer capacity (carbonate + hydroxide) can easily

be determined by the standard API phenolphthalein

titration. Some rough level of bicarbonate (over-

whelmed buffer) can be determined by a correlation

between the phenolphthalein titration endpoint and

measured pH, as shown previously.

In most field applications, it has been found that

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 with KOH .

A potential disadvantage with this method, however,

is that it allows the concentration of bicarbonate tobuild up. Excessive concentrations of bicarbonate

are known to cause rheology and fluid loss problems

in water-based muds [6].A good indication that the

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.: “CO 2 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 # SPE86502, 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] 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.

formate man a6 ph and buffering

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