A simple colourimetric method for determining seawater alkalinity using bromophenol

blue.

Vikashni Nand1 and Michael J Ellwood1

1 Research School of Earth Sciences, Australian National University, Canberra, ACT 2601,

Australia.

Running Head:

Seawater alkalinity using bromophenol blue

Key words:

Seawater alkalinity, seawater pH, dissolved inorganic carbon, bromophenol blue.

Corresponding author:

Vikashni Nand

Research School of Earth Sciences

The Australian National University

Building 142 Mills Road Acton

Canberra, ACT 2601 Australia

vikashni.nand@anu.edu.au

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Abstract

The development of small portable USB-spectrophotometer systems makes monitoring

alkalinity and pH possible in the field and remote locations. Here we present a method

utilising purified Bromophenol Blue (BPB) as an end-point indicator for making simple one-

point alkalinity measurements with spectrophotometric detection. The approach utilises

purified BPB dye whose absorbance characteristics have been determined over a range of

temperatures and salinities. The end-point pH for titrated samples was determined using the

BPB absorbance ratio (R(t) = 25 A590/A436) for the acid and base forms via the following

equation: pH=pK a+ log [ (R [25]−e1)(e2−R [25] . e3) ] where e1, = 0.00533, e2 = 2.232, e3 = 0.0319. A pKa

of 3.513 was determined for the dissociation of the second proton from the BPB dye. The

temperature (t) dependence of R can be expressed using the following relationship:

R(25)=R(t ) [1+(0.006774 ± 0.000009)(25−t)]. The dependence of the pKa on salinity (S) was

weak and can be expressed as p Ka (S )=p Ka (35)+[(0.00174 ± 0.00008) (35−S ) ]. Application of

the method for determining the alkalinity of in-house and certified standards typically

produced an uncertainty of ± 1.5 μmol kg−1 for purified BPB dye. When the impure BPB dye

was used as an end-point indicator the uncertainty for alkalinity measured was slightly higher

at approximately ± 3-4 μmol kg−1. Hence, if high precision alkalinity measurements are not

required (≥ 4 µmol kg-1) then utilisation of the unpurified BPB maybe suitable. We also

compared the use of BPB to two other dyes: bromocresol purple (BCP) and bromocresol

green (BCG). The utilisation of all three dyes for end-point determination produced

comparable results with an overall precision of ± 4 µmol kg-1. The one-point titration method

using BPB was utilised at a remote field location, One Tree Island, Australia, and was found

to be suitable for producing accurate and precise alkalinity data in a timely manner;

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approximately 10-15 samples can be determined per hour. When combined with seawater pH

measurements, the one-point titration method allows the full marine carbonate system to be

fully constrained without the need for high-tech spectrophotometric equipment and

comprehensive laboratory facilities.

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Introduction

Comprehensive studies of global oceanic inorganic carbon cycling are essential to understand

the oceanic uptake, transport, and storage of anthropogenic carbon dioxide. High-precision

carbonate system measurements are required to monitor the changes in seawater carbonate

chemistry. The characterization of the seawater carbonate system requires the measurement

of at least two key system parameters, e.g. total dissolved inorganic carbon (DIC), alkalinity,

carbon dioxide fugacity, and pH (Millero et al., 1993a){Millero, 1993 #154@@author-

year;Yao, 2007 #228}. Alkalinity and DIC are the preferred parameters for field

measurements and form the basis of analytical assessments of oceanic CO2 cycling and ocean

acidification. Although DIC and alkalinity are preferred parameters for constraining the

marine carbonate system, their measurement requires high-tech equipment, e.g. expensive

spectrophotometers with temperature control, to obtain high-precision data. Typically,

alkalinity measurements require sophisticated titration systems and high-quality pH

electrodes, which can be finicky to use, i.e. requiring an understanding of the Nernstian

behaviour of the electrode (Millero et al., 1993b), and with sometimes short-life spans.

More recently seawater pH has become a common parameter to measure and to help

constrain the uptake of anthropogenic atmospheric CO2 by the ocean (Clayton and Byrne,

1993; Martz et al., 2010; Rérolle et al., 2012). High precision pH can be measured using UV-

visible spectrophotometer and indicator dyes (Bellerby et al., 2002, Clayton and Byrne, 1993,

Tapp et al., 2000). Spectrophometric indicator-based methods, in theory, are simple, fast, and

precise. Spectrophotometric pH typically has a precision of 0.004−0.001 pH units

(DeGrandpre et al., 2014, Easley and Byrne, 2012); although the accuracy of

spectrophotometric pH measurements can be adversely affected by impurities in the indicator

dyes (Lai et al., 2016, Yao et al., 2007).

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Recent work has demonstrated that indicator purification is required to eliminate systematic

pH measurement errors (Liu et al., 2011, Yao et al., 2007); errors as large as 0.018 units have

been measured for impurified batches of meta-cresol purple (mCP) (Liu et al., 2011, Yao et

al., 2007). Meta-cresol purple is commonly used for measurements of oceanic seawater pH

over a range of 7.2−8.2 (Easley and Byrne, 2012, Patsavas et al., 2013). The work by Yao et

al., (2007) describes the effect of indicator impurities on spectrophotometric pH

measurements, and Liu et al., (2011) demonstrates that purified mCP could be used to

determine seawater pH on the total hydrogen ion concentration scale (pHT) accurately and

precisely over a wide range of temperatures (278.15 ≤ T ≤ 308.15) and salinities (20 ≤ S

≤40).

Work by Breland II and Byrne (1993) shows that the pH-sensitive dye bromocresol green

(BCG) is ideal for determining the pH of solutions in the range between 3.4 and 4.6. Thus the

dye is a useful indicator for determining the pH of seawater samples following the addition of

acid for alkalinity determination (Breland II and Byrne, 1993, Martz et al., 2006). Using

BCG as a dye, Breland II and Byrne (1993) were able to reproducibly determine seawater

alkalinity to a precision of ± 1.5 mol kg-1. More recently, Liu et al. (2015), developed an

automated system to measure alkalinity using the dye bromocresol purple (BCP). This

method obtains a precision ±1 mol kg-1 utilising volumetric addition of acid to the sample.

Like BCG and BCP, Bromophenol blue (BPB) can be used for spectrophotometric pH

measurements for the determination of total alkalinity (Haraldsson et al., 1997, Okamura et

al., 2010). Bromophenol blue is a diprotic sulfonephthalein indicator dye with a pKa value

~3.7 (King and Kester, 1989, Okamura et al., 2010), thus making it suitable for monitoring

pH for one-point alkalinity titrations. Here we build on the earlier work of Okamura et al.

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(2010) to investigate the impact impurities have on the use BPB as an end-point indicator for

alkalinity measurements and determining extinction coefficients and its temperature and

salinity dependence.

Theory

In seawater, alkalinity is defined as the number of moles of hydrogen ion required to convert

bicarbonate to carbonic acid (Dickson and Goyet, 1994, Dickson et al., 2007) and can be

expressed as follows:

HC O3−¿+H +¿→ H2 CO 3¿ ¿ (1)

Seawater also contains additional acid-base compounds thus total alkalinity (TA) is expressed

as:

TA=¿

(2)

where [H+]F represents the hydrogen ion concentration

The addition of H+ changes alkalinity by one equivalent unit (i.e. [HCO3-]), while two protons

are required to titrate CO32- to H2CO3. In equation (2), the addition of [H+] and [HSO4

−] will

decrease alkalinity as they are the sources of protons.

If equation (2) is inverted, the total analytical concentration of hydrogen ions can be obtained

relative to the reference proton condition for pure H2CO3. Following this conversion, the

total hydrogen ion concentration (CH) can be defined as:

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CH=¿ (3)

The addition of a mass (m) of hydrogen ions to the solution reduces TA, thus after the

addition of a mass (m) of acid to an initial mass of seawater(m¿¿0)¿, the following equation

can be expressed:

CH=mC−m0 TA

m0+m (4)

where C in the concentration of the acid added to the solution.

If equation (3) is substituted into equation (4) the following equation can be written:

mC−m0 TAm0+m

=¿

(5)

At the end of most open cell titrations, the pH is usually between 3.0 – 4.5 depending on the

initial alkalinity of the sample and the mass of acid added to the sample. At pH values below

~ 4.5, virtually all the carbonate species are converted to H2CO3 and CO2 and lost to the

atmosphere following purging. Thus for CH evaluation, the majority of the species can be

neglected apart from total sulfate and fluoride and equation (5) can be reduced to:

mC−m0 TA

m0+m=¿ (6)

where ¿ is the free hydrogen ion concentration.

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The TA of the sample can be determined by measuring [H+]T, if m0, mC and m are all known.

TA+¿) +¿) + m0+m+mI

m0 ¿) C =0 (7)

The ST and FT are the total sulfate and total fluoride and the Ks and Kf are the dissociation

constant of [HSO4-] and [HF]. Thus are dependent on the salinity and temperature (Dickson et

al. 2003).

The measurement of pH (expressed on the total proton scale) can either be done with a well-

calibrated pH electrode or with pH-sensitive indicator dyes. Two such indicator dyes useful

for determining pH in the range between 3 and 4.5 are BCG and BPB (Breland II and Byrne,

1993, Okamura et al., 2010). The relative colour of the dye is dependent on the proportion of

the dye acid and dye base present in solution (Figure 1), which can be expressed as:

H 2 I ⇋H+¿+H I−¿ ¿¿ (8)

H I−¿⇋H+¿+ I2−¿¿¿ ¿ (9)

Thus, the dissociation constant (Ka) of indicator species can be expressed as:

K a=¿¿ (10)

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The pH of the solution containing the indicator dye can be determined by measuring the

relative amounts of the acid and base present in solution using the following equation:

pH=pK a+ log [ (R−e1)(e2−R .e3) ] (11)

where R is the absorbance ratio of acid (HI-) and base (I2-) forms of the indicator (R=λ2 Aλ1 A ),

at wavelengths λ1 (436 nm) and λ2 (590 nm), respectively for BPB. The symbols e1, e2 and e3

are the molar absorbance ratios for each indicator species (Liu et al., 2011, Yao and Byme,

2001, Yao et al., 2007):

e1=λ2∈HI

λ1∈HI

(12)

e2=λ2∈I

λ1∈HI

(13)

e3=λ 1∈I

λ 1∈HI

(14)

where, λ1∈I and λ2∈I denote the molar absorbance of I2−, and λ1∈HI and λ2∈HI refers to the

molar absorbance of HI− at wavelengths λ1 and λ2 respectively.

In order to determine pH, the addition of the dye to the sample can perturb the pH of the

system, accounted for in the calculation of alkalinity. Thus equation (6) can be extended to

include the HI-:

mC−m0 TAm0+m+mI

=¿ (15)

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where mI is the mass of the indicator added to the sample.

Equation (15) shows that once all the relevant bases present in solution have been neutralised,

the number of moles of the protons in solution will be proportional to the concentration of

acid added to the solution. This expression is akin to a “Gran Function” whereby the end

point of the titration can be obtained from the number of moles of the protons in solution

versus the volume of acid added to the solution (Gran, 1952).

Methods

Dye purification and calibration

The solvents used in the purification of the indicator dyes BPB, BCG and BCP consisted of

water, acetonitrile (MeCN) and trifluoroacetic acid (TFA). The initial separations involved

using thin-layered chromatography (TLC) silica gel 60 F254 (Lot: HX248024) 500 glass

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plates (2.5 x 7.5 cm) plates. Using TLC, the mobile phase composition and concentration

were determined by trial and error, starting at 30% MeCN with 0.05% TFA and 0.1 % TFA

and then increasing the MeCN by 10% increments until the organic phase concentration was

80% MeCN. The optimum solvent composition dye-impurity separation consisted of 80%

MeCN plus 20% MQ-water (H2O) and 0.1% TFA. Once the composition of mobile phase

was optimized, preparative column chromatography using the silica gel 60 (0.040-0.063 mm)

was prepared for purification of the three dyes.

Stock solutions of unrefined BPB, BCG and BCP were prepared by dissolving the sodium

salt (Acros Organics, Lot: A0276342) in 80% MeCN, 20% MQ-water, and with 0.1% TFA.

To separate impurities, a 10 mmol L-1 solution of BPB in the mobile phase solution was

prepared and used for chromatographic separation. Pure BPB was collected at its

characteristic band once impurities had passed through the column. The excess solvent was

removed by rotary evaporation at 35oC under partial evacuation. Portions of the purified BPB

stock was re-dissolved in mobile phase for further High–Performance Liquid

Chromatography (HPLC) characterization. This procedure was repeated for BCG and BCP.

The purity of the BPB, BCG and BCP stock solution was confirmed by HPLC (Thermo

Finnigan Surveyor with a PDA detector) using an anion exchange column (PhenoSphere 5µ

SAX 80A, Column 150 x 4.6 mm) under isocratic conditions using a flow rate of 1.0 mL min-

1 and a 20 µL injection volume.

Characterisation of bromophenol blue extension coefficients

Standard solutions of BPB at pH values of 0.5 and pH 7.5 were prepared for characterisation

with an ionic strength of 0.7 mol kg-1 using sodium chloride (NaCl). pH adjustments were

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made by addition of either sodium hydroxide (NaOH) or hydrochloric acid (HCl). The pH of

each buffer solution was checked using Ross combination electrode (ThermoFisher). Dye

absorbance measurements were made using a UV-Visible spectrophotometer (Varian Cary

300). All BPB characterisation measurements were performed within a water jacketed 10 cm

path-length cell at 25 ± 0.1°C. Measurements were made at wavelengths 436 nm and 590 nm

and corrected at 690 nm. The absorbance measurements were corrected to produce values for

e1, e2 and e3.

Determination seawater alkalinity

The initial HCl solution was prepared to a concentration of ≈0.05 mol kg-1 with the BPB

indicator dye (and using BCG and BCP for comparison experiments) present at a

concentration of approximately 2 mmol L-1. The exact concentration of the acid was

determined by titrating against a standard sodium carbonate solution prepared in 0.7 mol kg-1

NaCl and against a certified reference material (CRM) (Batch 140; alkalinity = 2232 ± 0.8

µmol kg-1). The equivalence point for the titration was determined using the Gran function

with BPB as the pH indicator (Breland II and Byrne, 1993, Gran, 1988).

For laboratory-based alkalinity determination, 100 g (± 0.0002 g) of an in-house seawater

standard (ANU Std) was accurately weighed out into a wide-neck volumetric flask.

Typically, 5 g (± 0.0002 g) of HCl (0.05 mol L-1) containing BPB was added, and the sample

was accurately re-weighed. The amount of acid added is such that the final pH of the solution

was ~3.5. The sample was capped and well mixed before being purged with air for 5 minutes

to drive-off evolved CO2. The absorbance spectrum of each sample was collected at 25 ± 0.2

°C. Similar conditions were used for the BCG and BCP dye experiments except that the final

alkalinity pH was calculated using the pKa for the two respective dyes.

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For field-based alkalinity determination, 50 g (± 0.0002 g) of seawater was accurately

weighed out into a wide-neck volumetric flask. Typically, 2 g (± 0.0002 g) of HCl (0.05 mol

L-1) containing BPB was added to the sample so that final pH was around ~3.5 and was also

purged with air for 5 minutes. A USB UV-VIS spectrophotometer (USB4000, OceanOptics,

USA) coupled to a multi-LED light source (Ocean Optics BluLoop, Ocean Optics, USA) was

used for collecting absorbance spectra. The absorbance measurements were collected at 436

nm, 590 nm and 690 nm. The spectrophotometer, the sample cell holder and light source

were coupled via fibre optic cables. The temperature for each field sample was determined at

the time of measurement using a thermocouple (FEP coated thermocouple, TC Direct,

Australia coupled to a NI USB-9211, National Instruments, Australia) and used for

temperature correction of measured R values. The spectrophotometer and termeprature

sensor were controlled using custom software written in LabVIEW (National Instruments).

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Results and Discussion

Characterisation of bromophenol dye

The purification and characterisation experiments reveal that the off-the-shelf batches of the

BPB, BCG and BCP indicator dyes all contain impurities (Figure 2). We also tested various

brands of off-the-shelve-batches for BPB. All batches contained impurities to varying

degrees. Purification of these off-the-shelve-batches dyes removed >95% of the impurities

based on HPLC analysis (Figure 2). Work by Liu et al. (2011) showed that for the natural

seawater pH indicator dye mCP, impurities could lead to pH errors up to 0.018 pH units for

seawater pH values of ~8.1. We tested whether impurities can also influence the use of BPB

as an end-point pH indicator dye. By using impure BPB across multiple alkalinity

determinations using the one-point titration method, an average alkalinity of 2134 ± 2.7 µmol

kg-1 was obtained for an in-house seawater standard (Table 2). In contrast, using the same in-

house seawater standard and approach, but with the purified dye we obtained a value of 2138

± 1.2 µmol kg-1 (n = 10) (Table 2, Figure 3A). Consequently, there appears to be a small

offset in the “true” value obtained using the impure dye and a decrease in the precision of the

method. For titration of the alkalinity CRM (2232 ± 0.8 µmol kg-1), we obtained a value of

2232 ± 1.3 µmol kg-1 (n = 5) for the purified dye and 2233 ± 3.9 µmol kg-1 for the unpurified

dye (n = 5) (Figure 3B). The standard deviation for the unpurifed BPB dye is larger that for

both the in-house standard and the CRM standard thus it appears that dye impurities can

influence the absorbance spectrum resulting in subtle differences in measured R values,

which will translate into subtle pH variations and hence alkalinity. Depending on the nature

of the impurity, it may disproportionately influence relative size of the absorbance peaks for

acid and based forms of the dye thereby leading to an offset in the calculated solution pH.

This has been observed for impure mCP when used to determine natural seawater pH (Liu et

al., 2011). Our results would suggest that purification of the BPB indicator dye is required

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before it is used for high precision (<3 µmol kg-1) alkalinity determination. However, for

alkalinity measurements that do not require such high precision (>4 µmol kg-1) then the

impure BPB dye is suitable.

Initial investigations aimed at characterising BPB molar absorption coefficients were carried

out by Okamura et al. (2010) at 25°C and a salinity of 35 using unpurified BPB. The molar

absorption coefficients obtained for this dye were from absorbance spectra measured at pH

0.5 and 7.5. The molar absorption coefficients obtained for the purified BPB (this study,

Table 1) are slightly lower than values obtained by Okamura et al. (2010). These differences

for molar absorption coefficients may be because Okamura et al. (2010) did not purify their

dye.

An in-house seawater standard was titrated across a range of acid additions and the alkalinity

at each point computed and compared to its known alkalinity. Note that the alkalinity of the

in-house seawater standard should be constant and should not correlate to the mass of acid

added as long as all of the carbonate and bicarbonate ions present in solution have been fully

titrated (e.g. Figure 4). Using this approach, the in-house seawater standard was titrated

multiple times across a range of pH values between 3.2 and 3.7 (Figure 5A and 5B). The

alkalinity for each titration was computed and then plotted versus the residual H+

concentration (Figure 5A). The pKa of the dye was then adjusted so that the alkalinity across

the pH range was constant (Figure 5B). Using this pKa adjustment procedure, we obtained a

pKa of 3.513 for the in-house seawater standard with a mean alkalinity 2138 ± 1.2 μmol kg−1

(Figure 4). The pKa obtained for BPB in this study is significantly lower than the values of

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3.7037 and 3.695 obtained by Okamura et al. (2010) and King and Kester (1989),

respectively.

As a check, we also used the “Gran Function” approach to calculate the total alkalinity of the

in-house seawater standard (Gran, 1952, Gran, 1988) (Figure 5A). The alkalinity obtained

using this approach was 2138 ± 10 μmol kg−1 and compares favourably to our earlier value of

2138 ± 1.2 μmol kg−1 although the error associated with the Gran Function approach is

approximately eight times larger.

This approach also revealed no systematic variation in alkalinity values with increasing

hydrogen ion concentration, and hence the volume of acid added to the sample. Similarly, the

titration of the unpurified dye did not show any systematic variation with varying hydrogen

ion concentration (Figure 3) although the measured alkalinity concentration was slight offset

from the concentration for the in-house seawater standard determined using the purified dye.

The precision of the method was also slightly reduced when the unpurified dye was used

compared to use of the purified dye (Table 3). Various batches of bromophenol blue dye from

three different vendors (Acros Organic, Sigma-Aldrich and Fisher Chemicals) were also

tested and had no significant difference between the BPB brands (Table 2).

Tests on purified versions of the BCG and BCP dyes were also undertaken to determine

whether they produced results similar to that of BPB. The results show that these two dyes

were comparable to that of BPB indicating that all three dyes are suitable for determining

seawater alkalinity to a precision of ± 4 µmol kg-1 or better (Table 3). Note that the pKa values

for these dyes were slightly adjusted so that they produced consistent alkalinity values across

and the range of hydrogen ion concentrations, as done for BPB in figures 4 and 5 (Table 3).

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Errors associated with weighing, temperature and salinity.

Breland II and Byrne (1993) outlined the main errors associated with the colorimetric

determination of seawater alkalinity with the main errors associated with the weighing of the

sample and the titrant acid. The main errors in our experiments are associated with the weight

of the acid required to titrate the sample and with dye measurement by UV-visible

spectrophotometry. The weighing error was approximately two parts in ten thousand (±

0.0002 g for a 2 g acid addition), which results in an alkalinity error of approximately ± 0.4

mol kg-1. Repeated absorption analysis of fully titrated and purged samples resulted in an

alkalinity error in the range of ± 0.4-1.0 mol kg-1 using the OceanOptics USB4000

spectrophotometer. This range is larger than the error purely associated with weighing the

acid and sample and is likely to represent errors associated with the temperature and salinity

calibration of the dye, and subtle differences in alignment with replacing the sample cell in

and out of the sample cell holder.

Two factors that can influence the accuracy of colourimetric alkalinity determination is that

of temperature and salinity (Breland II and Byrne, 1993, Okamura et al., 2010). To test the

temperature dependence of absorbance ratio R for the acid (HI-) and base (I2-) forms of BPB,

we measured R across a temperature range for a sample with known alkalinity acidified to a

pH of ~3.5. These ratios were then referenced back to 25C (Figure 6). The absorbance ratio

of the acid (HI-) and base (I2-) forms showed strong temperature dependence whereby R

increases as the temperature of the sample decreases. For instance, we observed a 10%

increase in the R when the solution temperature was decreased from 25C to 10C (Figure 6).

To correct for this temperature effect we defined the following equation:

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R(25)=R(t ) [1+(0.006774 ± 0.000009)(25−t)] (16)

where R(25) is the absorbance ratio at 25C and R(t) is the absorbance ratio at a given

temperature (t) between 10C and 35C. The slope of this relationship is slightly less than for

BCG (Breland II and Byrne, 1993).

To determine the influence of salinity on pKa for BPB, we measured the pH of an in-house

standard seawater sample across a salinity range from 36.4 to 21.8. The initial pH of the

solution was measured at 3.30. After each subsequent dilution using Milli-Q water, the pH of

the sample was re-determined (Figure 7). The effect of changing salinity on sample pH can

be described using the following equation:

p Ka=p Ka (35)+[(0.00174 ± 0.00008) (35−S )] (17)

where pKa is the subsequent pH for each of the diluted samples and pKa(35) is the calculated

pH of our acidified and purged seawater sample at salinity 35. In equation (17) S represents

salinity. The change in the pKa for BPB associated with changing salinity was found to be

relatively small (Figure 7), which is consistent with the work of Breland II and Byrne (1993)

where they also noted that salinity had a minor influence on the pKa for BCG.

Field application

As a test, we utilised BPB alkalinity method, in conjunction with natural seawater pH

measurements utilising purified mCP dye (Liu et al., 2011), to track changes in carbonate

system parameters on the algal reef flat at One Tree Island, Queensland Australia. At low-tide

seawater samples were collected from the algal reef flat (Latitude 23°30.695 S, Longitude

152°05.271 E) across a 4-5 hour sampling window. For field-based alkalinity determination,

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60 mL seawater samples were collected periodically (~30 minutes) from isolated pools

during low tide and filtered through a 0.2 µm syringe filters (Pall, Australia). Two dome

experiments were also conducted in parallel during the low-tide period to quantify carbonate

production and dissolution within the algal pools when isolated from the atmosphere and

surrounding waters. During these experiments, the domes were placed on the algal turf, thus

isolating a section of benthos and allowing changes in water chemistry that attribute to the

inorganic and biological processes within. The domes were fully filled with ambient

seawater and any remaining air was removed from the domes to prevent CO2 gas exchange.

Samples were periodically removed from the domes for pH (Liu et al. 2011), alkalinity and

dissolved oxygen determination. The alkalinity samples were filtered upon collection and

then stored on ice and transported back to the One Tree Island research station where they

were analyzed. Samples were measured within 2 to 8 hours of collection. Samples were not

poisoned before analysis. Before analysis, alkalinity samples were accurately weighed into 50

mL wide-neck volumetric flasks and then acidified with HCl containing BPB and re-

weighed. The final pH of each sample was around ~3.5. Samples were warmed to room

temperature (approximately 25C), purged and then analyzed using the USB-

spectrophotometer system noting that the temperature at the time of measurement was

carefully determined and accounted for in the final alkalinity results. The results from these

field measurements are presented in Figures 8A, 8B and 8C. The alkalinity and natural

seawater pH measurements made on samples collected from the isolated algal reef flat pools

and dome experiments showed significant variations. For samples collected during the

daytime period (4 January 2015), seawater alkalinity decreased while natural seawater pH

increased, which is consistent with calcium carbonate production and photosynthesis by

crustose coralline algae. During the nighttime period (5 January 2015), seawater alkalinity

increased while natural seawater pH decreased, which is consistent with calcium carbonate

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dissolution and organic matter respiration. The diurnal variations in seawater alkalinity are

large (up to ± 300 µmol kg-1) and were well outside procedural errors for both alkalinity and

natural pH measurement. The beauty alkalinity of the method presented here is that alkalinity

measurements can be made quickly, approximately 10-15 samples per hour, with a high

precision thus allowing the study of ecosystem changes with high temporal resolution, such

as at One Tree Island.

Marine ecosystems have become a significant research focus especially in coastal areas

where chemical and biological variations can be large and dynamic, such as at One Tree

Island (e.g. Silverman et al. (2012)). In this study, a simple but accurate alkalinity method has

been developed to determine seawater alkalinity. This method is well suited to field

experiments and sites where access to high-tech instrumentation and laboratories is limited.

Moreover, when this method is combined with similar dye methods for constraining natural

seawater pH, it provides an easy way to constrain the marine carbonate system with a high

degree of accuracy and precision (Figure 8C). The method is applicable especially when there

are dynamic diurnal changes in calcification, dissolution, photosynthesis and respiration.

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Conclusions

In this study, we developed a method utilising BPB as an end-point indicator dye for making

simple one-point alkalinity measurements. HPLC analysis of the BPB dye reveals impurities

that can result in a slight loss of precision (e.g. 4 µmol kg-1) of the procedure compared to

alkalinity measurements made with the purified dye (±1.2 µmol kg-1). The absorbance

characteristics of the BPB dye is strongly influenced by changes in temperature (10°C - 35

°C) but is weakly influenced by changes in salinity (20-36). High precision (± 1.2 µmol kg-1)

alkalinity measurements are obtainable using a low-cost and low-powered USB-

spectrophotometer system. If high precision alkalinity measurements are not required, i.e. ≥

±4 µmol kg-1 then unpurified BPB maybe suitable. The one-point alkalinity method utilising

BPB as an end-point indicator appears to be ideal for use in remote field locations where

quick, but precise alkalinity measurements are required to help constrain the marine

carbonate system.

Acknowledgments

VN was supported through an ANU funded PhD scholarship. We acknowledge managers at

One Tree Island for their support during the undertaking of fieldwork on the island. Funding

Support from the Australian Research Council (DP170103067) is also acknowledged.

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Table 1. Measured and calculated molar absorption coefficient ratios for e1, e2, e3 at 25oC and

salinity 35 for this study compared with Okamura et al., (2010).

Lab experiment Okamura et al. (2010)

e1 0.00533 0.0027

e2 2.232 2.458

e3 0.0319 0.1566

Table 2: Measured alkalinity using the purified and impure bromophenol blue from different

vendors for a sample size of n =10.

Bromophenol Blue Pure dye

Average ± SD

(µmol kg-1)

Impure dye

Average ± SD

(µmol kg-1)

P value$

Acros Organic 2138 ± 1.2 2134 ± 2.7 <0.001

Sigma-Aldrich 2138 ± 1.9 2132 ± 3.6 <0.001

Fisher Chemicals 2138 ± 1.4 2134 ± 3.9 0.0069

$ P values indicate that the alkalinity measurements utilising the impure BPB dye are

significantly different to alkalinity measurements utilising the purified BPB dye.

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Table 3. Average alkalinity measurement for ANU standard seawater using the purified

bromophenol blue, bromocresol green and bromocresol purple dyes.

Indicator pKa Alkalinity (Average ± SD, n = 10)

Experimental Literature Literature pKa Adjusted pKa

Bromophenol blue 3.513 3.7037$ 2325 ± 46.6 2138 ± 1.2†

Bromocresol green 4.4099 4.4160# 2143 ± 4.9 2138 ± 3.8†

Bromocresol purple 5.9718 5.9770& 2139 ± 3.8 2136 ± 4.3†

References: $ Okamura et al. 2010, #Breland II and Byrne, 1993; &Yao and Byrne, 2001.

†P values for these alkalinity measurements utilising purified dyes are not significant (P >

0.2) i.e. different from each other.

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Figure captions

Figure 1. Absorbance spectra of BPB dye in acidic (HI-) and basic (I2-) forms.

Figure 2. Plot showing the 3D UV-Vis spectra (300 to 600 nm) for HPLC profiles of the A.

unpurified BPB dye (10 mmol L-1), B. purified BPB dye (10 mmol L-1), C. non-purified BCG

dye (~10 mmol L-1) and D. purified BCG dye (1~0 mmol L-1). For each chromatogram the

total injection volume was 20µL at a flow rate of 1 mL min-1.

Figure 3. A. Multiple alkalinity determinations using impure and pure BPB (average

alkalinity of 2134 µmol kg-1 ± 2.7 µmol kg-1 and 2138 ± 1.2 µmol kg-1 (n = 10) for an in-

house seawater standard (ANU Std) respectively. B. Multiple alkalinity determinations using

impure and pure BPB (an average alkalinity of 2233 µmol kg-1 ± 3.9 µmol kg-1 and 2232 ±

1.3 µmol kg-1 (n = 5) for alkalinity CRM, respectively. The certified alkalinity for the CRM

(batch 140) was 2232.58 ± 0.80 µmol kg–1.

Figure 4. Plot showing changes in alkalinity versus mass of HCl for various pKa values for

BPB. The pKa for BPB was optimised using an in-house seawater standard to a value of

3.513 where the variation in alkalinity concentration was at a minimum.

Figure 5. A. Amount of hydrogen ion versus mass of HCl added to the sample. B. Total

alkalinity measurements for the CRM and an in-house seawater standard versus sample

hydrogen ion concentration measured using the USB UV-VIS spectrophotometer system. The

mean alkalinity and the standard deviation for the CRM (batch 140) and an in-house seawater

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standard (ANU Std) are 2232 ± 1.3μmol kg−1 (n = 5) and 2138 ± 1.2 μmol kg−1 (n = 10),

respectively.

Figure 6. A. Plots of R(25)/R(t) versus the temperature range from 10-35oC of the CRM and in-

house seawater standard (ANU Std). The best fit linear regression line for the data is given

as: R(25)=R(t ) [1+(0.006774 ± 0.000009)(25−t)]. B. Plot of residuals versus the temperature

for the CRM and the in-house seawater standard.

Figure 7. Plot of subsequent pH of diluted and in-house seawater standard (ANU Std) versus

salinity at 25°C. The best fit linear regression line for the data is given as:

p Ka=p Ka (35)+[(0.00174 ± 0.00008) (35−S )]

Figure 8. Examples of A. alkalinity and B. pH measurements versus time of day for samples

collected on the algal ridge at low tide from One Tree Island, Queensland, Australia. Samples

were collected within a 24-hour window and measured within 2-8 hours of collection. C.

Transformation of data from A and B into alkalinity versus DIC plot overlaid with aragonite

saturation state ().

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