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<title>Guidelines for Measuring Changes in Seawater pH and Associated Carbonate Chemistry in Coastal Environments of the Eastern United States</title>
<type>single page tiff</type>

United States
Environmental Protection
EPA/600/R-17/483 | March 2018
Guidelines for Measuring
Changes in Seawater pH and
Associated Carbonate Chemistry
in Coastal Environments of the
Eastern United States
Office of Research and Development
National Health and Environmental Effects Research Laboratory


j^CPA	EPA/600/R-17/483 | March 2018
united states
Environmental Protection
Guidelines for Measuring Changes in
Seawater pH and Associated Carbonate
Chemistry in Coastal Environments
of the Eastern United States
Adam R. Pimenta
Jason S. Grear
Atlantic Ecology Division
National Health and Environmental Effects Research Laboratory
Narragansett, RI
National Health and Environmental Effects Research Laboratory
Office of Research and Development
U.S. Environmental Protection Agency
Washington, DC 20460

The U.S. Environmental Protection Agency through its Office of Research and
Development funded and managed the research described herein. This document has been
subjected to the Agency's peer and administrative review and has been approved for publication
as an EPA document. Mention of trade names or commercial products does not constitute
endorsement or recommendation for use.
These guidelines are written for a variety of audiences ranging from shellfish growers
interested in monitoring pH with inexpensive equipment to citizen monitoring groups to
advanced chemistry laboratories interested in expanding existing capabilities. The purpose is to
give an overview of available sampling, analytical and data reporting approaches that will
contribute to the usefulness of coastal acidification measurements for both the needs of those
intending to monitor as well as those of other interested stakeholders along the Atlantic seaboard
of the US. The state of the science, including recommended best practices, is rapidly evolving,
so certain sections may be either too sparse or too detailed. Thus, we encourage users of the
guidelines to begin with a careful review of the detailed contents listing and to take note of
references to other guidelines available in the open literature.
Brian Rappoli and Matthew Liebman provided early and continuous feedback on the
scope and content of this report. External reviews of earlier versions were provided by Dwight
Gledhill, Christopher Hintz, Christopher Hunt, Esperanza Standoff, Elizabeth Turner, and
Zhaohui Aleck Wang.
ii|Measuring Changes in Coastal Carbonate Chemistry

3.1	Eutrophi cation	5
3.2	Nitrification	6
3.3	Denitrification	6
3.4	Stratification	6
3.5	Regional lithology, aquatic geochemistry, carbon import/export, and climate	6
3.6	Calcification and dissolution	7
3.7	Sediment diagenesis	8
4.1	Overview	9
4.2	Dissolved inorganic carbon (DIC)	10
4.3	Carbon dioxide (CO2)	11
4.4	Total alkalinity (TA)	11
4.5	pH	11
4.6	Bicarbonate	13
4.7	Carbonate	13
4.8	Calcite and aragonite saturation states	14
5.1	Sample containers	15
5.2	Cleaning	15
5.3	Preferred order of sample collection	15
5.4	Filling sample bottles	16
5.5	Long-term storage bottles	16
5.6	Short-term storage bottles	16
5.7	Sample preservation and shelf life	17
5.8	Sampling without preservation	17
5.9	Transport	17
Contents liii

5.10	Other sampling equipment	17
5.11	Filtration	17
5.12	Analyte-specific sampling techniques	18
6.1	Certified reference materials (CRMs)	19
6.2	Alkalinity	19
6.3	Dissolved inorganic carbon	21
6.4	pH	22
6.5	pCO:	25
9.1	Precision and accuracy	28
9.2	Weather vs. climate data quality benchmarks	30
Table A-l. Comparison of methods and uncertainty for carbonate system analyses	31
Bottle Sampling and Preservation Checklist	32
B.l Essential elements	34
B.2 Example LOP: Adapting a total alkalinity SOP to a specific laboratory setting	34
B.3	Example LOP: Measuring DIC using an NDIR-based instrument	37
C.	1 A water quality research laboratory	43
C.2 A single-instrument setup in a basic water quality laboratory	44
C.3 A monitoring effort with external laboratory support	44
C.4 Shellfish growers and hatcheries	45
iv | M c a s u i" i n. g € h a n g c s i n € oast a 1 € a r b o n a t c € h c m i s t r y

Coastal and estuarine systems hold significant economic and recreational importance for
communities along the Atlantic seaboard. These systems support finfish, bivalve, crustacean and
seabird populations and play vital roles in water quality and the cycling of nitrogen and carbon.
However, seawater pH and other characteristics of coastal carbonate chemistry are changing
through a process known as coastal acidification, which is a fundamentally similar but more
complex version of ocean acidification. Coastal acidification has the potential to disrupt the
species composition and ecological functioning of coastal biological communities and threaten
commercially important aquatic life. As in the open ocean, the carbonate system in coastal
waters consists of the major forms of inorganic carbon present in seawater, which are carbon
dioxide, bicarbonate and carbonate. Although there are numerous groups interested in
monitoring pH or other carbonate parameters in coastal waters, there is little available guidance
on how these groups can best utilize or expand their existing capabilities.
Coastal acidification differs somewhat from ocean acidification, which is a global process
that involves a reduction in the pH of the ocean (see section below on the seawater carbonate
system). It is caused primarily by carbon dioxide from the atmosphere entering the ocean.
Coastal acidification is a more localized, further reduction in pH. It is primarily driven by high
levels of respiration (typically by bacteria involved in decomposition), which releases carbon
dioxide into the water. Coastal acidification is often fueled by nutrients entering the water from
land, stimulating phytoplankton blooms that subsequently decompose on or near the seabed.
Coastal acidification happens in coastal waters because that is where high nutrient levels and
algal blooms occur [].
In the past few decades, only half of the CO2 released by human activity, including fossil
fuel emissions, land use change and cement production, has remained in the atmosphere; of the
remainder, about 30% has been taken up by the ocean and 20% by the terrestrial biosphere
(Khatiwala et al., 2009; Sabine et al., 2004). The evidence for decreasing pH in the open ocean
is unequivocal (Caldeira and Wickett, 2003; Doney et al., 2009), as is the evidence for negative
effects on many marine organisms when these chemical changes are simulated under controlled
laboratory conditions (e.g., Kroeker et al., 2013; Talmage and Gobler, 2009). However,
scientists are just beginning to test the severity of these effects in ocean and coastal ecosystems
where an organism's chemical environment is only one of many ecological factors affecting its
There are many clear cases of extreme biological sensitivity to acidification among
economically important coastal organisms such as shellfish and corals, but the biological
responses of many other species are variable and difficult to predict (Kroeker et al., 2010). For
example, many types of marine plants and algae may be harmed by lower pH (i.e., higher
acidity) but may also benefit from increases in the carbon dioxide they require for photosynthesis
(Riebesell, 2004). Thus, although species composition may change in the future, neither the
details nor the ecosystem level consequences (e.g., food production) are predictable (Grear et al.,
2017). The continued study of these effects needs to be accompanied by a clear understanding of
how coastal carbonate chemistry varies through space and time. A number of methods have
been described for the coastal current and upwelling zones of the US west coast (e.g.,
McLaughlin et al., 2014; McLaughlin et al., 2013). While coastal upwelling occurs on the east
coast, deep water upwelling does not strongly influence acidification in the short term (Wang et
al., 2013). Thus, observations from the mid- and outer-shelf may be less comparable to the
1 .Introduction I 1

inshore environment than on the west coast. Moreover, many coastal organisms have sensitive
estuarine and nearshore life stages that coincide with mid and late summer extremes in dissolved
oxygen, pH, and other characteristics of the carbonate system and are thus expected to be
especially vulnerable (Wallace et al., 2014). These issues raise concern about coverage in the
nearshore environments that fall outside of areas covered by the major federal observing
programs (e.g., ECOMON and GOMECC) and which tend to be too infrequent to capture either
seasonal or more frequent excursions in carbonate chemistry.
The decrease in pH in the open ocean during the industrial age has been on the order of
0.1 to 0.2 pH units per century (Caldeira and Wickett, 2003), which translates to more than a
25% increase in the concentration of hydrogen ions. Relative to coastal environments, pH in the
open ocean is generally less dynamic in terms of diurnal and seasonal variations (Hofmann et al.,
2011),	which has made open ocean trends easier to distinguish from background variability. In
addition, the ocean is extremely important to the global carbon cycle, so scientists have been
taking highly precise measurements of the carbonate system in the open ocean for decades.
However, due to greater variability of pH in the coastal environment, a trend of similar
magnitude would require a larger number, and longer time-series, of samples to detect (Keller et
al., 2014). This creates a unique challenge for coastal monitoring because current best practices
for handling and analyzing samples for carbonate system parameters are expensive, and therefore
possibly not feasible for the high frequency and spatially extensive sampling that would be
necessary to detect decadal and spatial trends in the coastal environment. For example, while pH
is easy to measure with handheld meters or multi-function autonomous sensors that use glass
membrane pH electrodes, chemical oceanographers often question the value of these
measurements for the study of carbonate chemistry, including acidification (Re'rolle et al.,
2012).	Although this criticism is sometimes unwarranted because of differences in study goals
(e.g., see the "climate vs. weather quality" discussion below; Newton et al., 2014), accepted
protocols are unlikely to change without an improved understanding of coastal acidification, and
until issues relating to appropriate pH scales, calibration standards, instrument drift, and indirect
pH estimation are further refined or agreed upon by the research community.
These guidelines are meant to be a resource for learning about and performing
measurements of the seawater carbonate system, especially as they relate to coastal acidification.
The intended audience includes scientists in academic, government, and non-government
organizations including those involved in citizen science and shellfish management. Many such
organizations are already monitoring or beginning to monitor components of the seawater
carbonate system that may be partially or completely sufficient for assessing coastal
acidification. For example, specific organizations in the northeast are examining coastal
acidification as a potential cause for recent declines in shellfish abundance. Other organizations
study coastal carbonate chemistry and acidification as part of a broader interest in coastal carbon
cycles. Clearly, there is a wide diversity of rationales and capabilities for monitoring
acidification in the coastal environment.
Numerous publications exist in the peer-reviewed and online "gray" literature that
describe recommended practices for measuring and calculating the various components of the
seawater carbonate system. Most of these resources are written by and for oceanographic
researchers and place less emphasis on informing, for example, the expansion of an existing
shellfish or nutrient monitoring program to include coastal acidification parameters. In addition,
the available resources (e.g., Dickson et al., 2007; McLaughlin et al., 2014; Riebesell et al.,
2010) tend to be generalized to accommodate a wide variety of instrument and laboratory
2 | Mcas u i" i n. g € h a ngcs i n C oast a 1 € a r b o n a t c € h c m i s t r y

configurations. This creates a challenge for the new investigator and slows the rate at which new
monitoring efforts can be implemented. Thus, this document will attempt to address the unique
issues and some of the available solutions for measuring the seawater carbonate system in coastal
and estuarine environments. This includes alternatives and clarifications of existing best
practices that will make them suitable for typical environments of the US east coast. Because the
study of ocean and coastal acidification is changing rapidly, these guidelines are not meant to be
prescriptive, but are intended to facilitate the development of compatible datasets for sharing
insights and experiences from the present community of investigators interested in coastal
acidification. Although these guidelines are intended to apply throughout the east coast, much of
the nearshore research in the eastern US has been conducted in the northeast. There are likely to
be solutions that we have not covered here, so continued communication through acidification
monitoring networks will be critical.
Before describing specific methods, we provide an overview of coastal acidification, the
seawater carbonate system, and ecological considerations that normally affect study design.
There is a vast literature on these topics for the ocean and a growing one for the coasts. For a
recent overview of the state of the sci ence in the US northeast, see Gledhill et al. (2015) and
other resources on the Northeast Coastal Acidification Network (NECAN) website
( For other eastern US coastal regions, see ongoing developments on the
SOCAN ( and MACAN websites ( Although the
seawater carbonate system is described later in greater detail, our overview begins with the
summary in Figure 1-1.
crt« «L
^ «CQ,
mineral ?jj
(calcifying pla
e.g. foraminifera)
Figure 1-1: Carbonate system equilibrium reactions and relationship to calcification
(Barker and Ridgwell, 2012).
1 .Introduction I 3

Although the fundamental factors driving carbonate chemistry are essentially the same in
ocean and coastal waters, the relative importance of each of these factors differs a great deal.
This results in large differences in spatial and temporal variability, with nutrient dynamics,
primary production, and respiration being especially important in coastal areas. In nutrient-poor
surface waters of the open ocean, for example, equilibration of carbon dioxide concentration
with the atmosphere via air-sea exchange is less significantly perturbed by short-term variations
in biological activity and terrestrial run-off events. Biological fixation and eventual removal
from surface waters of carbon favors further uptake of carbon dioxide from the atmosphere by
the oceans, but the large magnitude of these processes per unit area in coastal environments
result in much more dynamic systems both temporally and spatially (Johnson et al., 2013). This
is why productive estuaries exhibit much greater diurnal and seasonal variation in pH than open-
ocean systems (Hofmann et al., 2011). In addition, although the net effects of atmospheric
nitrogen and sulfur dioxide deposition on ocean pH are thought to be small, they are potentially
more important in coastal ecosystems (Doney et al., 2007).
A pH probe placed in the upper water column of a productive estuary will typically
indicate rising pH (lower acidity) to levels sometimes exceeding >8.4 during the day and then
declining into the <7 range during the night. Similarly, pH tends to be higher during the growing
season than in the fall. These fluctuations occur in estuaries because pH change is determined
primarily by the amount of carbon dioxide that enters and leaves the water column, although
other acid species (e.g., organic acids) can significantly affect water pH in coastal oceans. The
key sources for carbon dioxide in the water column include absorption from the atmosphere,
respiration by marine organisms, and inputs from other systems, such as rivers, wetlands, and
sediments. Conversely, when photosynthetic organisms (e.g., phytoplankton) produce oxygen,
they consume carbon dioxide. These ecosystem processes explain why daily and seasonal
patterns of rising and falling pH and pCOi often mirror trends seen in oxygen measurements.
The sensitivity of the local carbonate system to these processes also depends on alkalinity (the
ability of substances in seawater to react with the addition of a strong acid and convert it to an
uncharged species) and buffer capacity, which in turn is sensitive to nearby ocean characteristics
as well as the limnological and lithological characteristics of the watershed. These various
processes interact to control pH and carbonate ion availability (Cai et al., 2011; Feely et al.,
2010; Mucci et al., 2011; Wallace et al., 2014; Wang et al., 2013).
Although strong diurnal and seasonal variability can be exacerbated by human activity,
there is an expectation that coastal organisms have evolved under highly variable conditions and
may be less sensitive than oceanic biota to long-term changes in carbonate chemistry. Although
differences in acclimation and adaptation potential among subpopulations from environments
with differing pH variability have been observed, neither the specific characteristics of pH
variability that affect organisms - i.e., daily minimum pH, total amplitude, mean pH, breeding
season pH, etc., nor the phylogenetic traits that confer tolerance or adaptive capacity are well
known, so the sensitivity of most coastal biota remains unclear.
4|Measuring Changes in Coastal Carbonate Chemistry

Against a background of increasing carbon di oxide concentrations in the atmosphere due
primarily to fossil fuel combustion, there are several factors contributing to geographical
variation in coastal acidification that are particularly important on the US east coast (Salisbury et
al., 2008; Wang et al., 2013). Consideration of these factors will help readers to identify survey
design issues, areas of possible collaboration with other scientists, and coordination with existing
monitoring efforts. This section gives an overview of these factors.
3.1 Eutrophication
One of the most important factors driving acidification of coastal and estuarine waters is
nutrient loading. Local input of nutrients, especially NFh and NO.f. can be increased due to
urbanization, wastewater input, and agricultural/urban runoff. The northeast US has been shown
to host some of the largest nutrient loading rates in the world (Anderson and Taylor, 2001;
Howarth, 2008). These nutrient inputs can lead to enhanced primary production in estuarine and
coastal systems by phytoplankton (Beman et al., 2005; Carpenter et al., 2008; Feely et al., 2010;
Newton and Van Voorhis, 2002; Simonds et al., 2008), which eventually senesce and sink to the
bottom waters where their microbial decay consumes oxygen and produces carbon dioxide. As
already noted, this microbial respiration and remineralization of organic matter, whether from in
situ phytoplankton production or from watershed loading of organic matter, can increase
localized CO2 concentrations in aquatic ecosystems (Figure 3-1; Cai et al., 2011). While this can
occur in any part of the water column, these conditions can be exacerbated in bottom waters due
to reduced vertical mixing thus enhanced respiration/decomposition during periods of
stratification (Feely et al., 2010; Gobler and Baumann, 2016; Wallace et al,, 2014). As a result,
bottom waters especially, can be affected by lower pH and carbonate ion concentration (Feely et
al., 2010), often expressed as a reduction in the saturation state of the calcium carbonate minerals
aragonite and/or calcite. Through turnover and mixing events, usually in early fall, these
undersaturated waters can also make their way to the surface layer.
co2	co2	co2	co2
Figure 3-1. A conceptual model for a
large river plume eutrophication and
subsurface water hypoxia and
acidification. From Cai et al. (2011) by
permission from Springer Nature:
Macmillan Publishers, Nature Geoscience.
3.Factors Contributing to Acidification |5

3.2	Nitrification
Nitrification is a biochemical process carried out primarily by sediment bacterial
communities and results in the oxidation of ammonium to nitrate. Changes in total alkalinity
may accompany nitrification because carbonate species are consumed in the conversion of
ammonia to nitrite (Wolf-Gladrow et al., 2007). Nitrification can occur in freshwater, brackish
and marine systems. During nitrification, the oxidation of ammonium to nitrate produces two
moles of H+ ions for every mole of ammonium oxidized, with a corresponding decrease in
2 NH4+ + 3 02 = 2 N02" + 4 H+ + 2 H20
2 N02" + 02 = 2 N03"
3.3	Denitrification
Denitrification is a biochemical process, in which nitrate is reduced to molecular
nitrogen. The process results in the consumption of a proton and production of carbon dioxide
(Brenner et al., 2016; Drtil et al., 1995). The overall net effect on pH that can be attributed to
both nitrification and denitrification will be small in a well-buffered system because these
processes tend to co-occur in time and space (Rysgaard et al., 1996).
3.4	Stratification
Due to temperature and salinity differences between bottom and surface waters and
between riverine and oceanic waters, it is common for stratification (i.e., layering) to occur in
estuarine systems. Stratification tends to be strongest in the summer, especially during periods
of reduced wind-driven mixing and tides (i.e., neap tides) and in systems with sufficient surface
water heating to form a thermocline. Seasonal changes in precipitation and runoff are additional
factors affecting stratification because of their effect on water buoyancy via salinity. Sinking of
organic matter during periods of stratification causes these materials to accumulate in bottom
waters and near the sediment layer where they fuel C02 production due to microbial respiration.
Since mixing is reduced under strong stratification, C02 builds up in the bottom waters,
increasing acidification. In some systems, especially in late summer and early fall, bottom water
C02 concentrations can greatly exceed atmospheric concentrations (Frankignoulle, 1988) and
remain high until temperature, tidal, or wind-driven mixing occurs during the fall and winter
months (Gledhill et al., 2015; Wallace et al., 2014; Wang and Cai, 2004).
3.5	Regional lithology, aquatic geochemistry, carbon import/export, and climate
The geologic and land use makeup of the watershed can have a pronounced effect on the
pH, dissolved inorganic carbon (DIC) concentration, alkalinity and buffering capacity of the
freshwater entering an estuary. Rivers that drain upland areas with exposed bedrock (especially
carbonate rocks like limestone or marble) generally have higher carbonate content, pH and
buffering capacity. Rivers originating in coastal plains may have lower pH and buffering
capacity and higher amounts of C02 and also may contain high concentrations of humic, fulvic
and tannic acids (Cai and Wang, 1998; Suchet et al., 2003). Changes in precipitation amounts
and timing in coastal watersheds have the potential to change the alkalinity and thus the
buffering capacity of receiving waters. Such watershed effects were invoked as a potential
explanation of the lower buffering capacity observed in shelf waters off of New England (Wang
et al., 2013).
6 | Mcas u i" i n. g € h a ngcs i n C oast a 1 € a r b o n a t c € h c m i s t r y

Carbonate chemistry can be strongly influenced by oxidation of organi c matter that
originates from the watershed (Cai et al., 2011). In addition, Wang et al. (2016) showed that
coastal marshes are a significant source of DIC and alkalinity and reported estimated extremely
high annual DIC exports to coastal waters from the salt marsh at their Massachusetts study site.
3.6 Calcification and dissolution
Calcifying organisms extract Ca and carbonate system ions from the water surrounding
them in order to build shell. Although the calcification process is complex and variable among
species (Jury et al., 2010; Leung et al., 2017; Riebesell et al., 2010), many marine organisms
directly use carbonate ions through either passive or active uptake. Because this often involves
active transport of protons across the cell membrane to maintain pH balance, dissolution or a
slowing of calcification due to energetic costs occur when pi I either alters the proton gradient
across the membrane or otherwise affects calcification physiology. Calcification and dissolution
reactions can also affect observed carbonate chemistry (e.g., Ailing et al., 2012), especially
through abiotic reactions at the sediment-water interface (Soetaert et al., 2007).
! Air
H2C03 < > H+ + HC03-
—> H20
C032- + H20 «. > OH- + HCO3-
Rock, soil,
or sediments CaC03(s)
Figure 3-2. A conceptual model describing the interactions between carbonate system species and
associated atmospheric and sediment fluxes. (Image from Patricia Sliapley
In the ocean, CaCCb is found in sediments. While CaCO.3 exists in equilibrium with its
constituent ions (see Figure 3-2), the equilibrium is shifted toward dissolution as the
concentration of CO2 increases. Dissolution occurs via the following reaction:
CaC03 + CO2 + H20 = Ca2 2HC03"
While these precipitation and dissolution processes are of interest in studies of ecosystem
processes governing DIC pools (e.g., Ailing et al., 2012), the concentration of carbonate ions
tends remain near 10% of the DIC pool via equilibrium reactions over large spatial and temporal
scales. However, it is important to note that riverine calcium concentration is more variable than
in the ocean, so the solubility coefficients for carbonate minerals may also be less stable in
coastal environments.
3.Factors Contributing to Acidification |7

3.7 Sediment diagenesis
Sediment diagenesis describes a diverse set of mechanisms that transform sediment
following deposition and may significantly affect alkalinity (Jahnke and Jahnke, 2004; Krumins
et al., 2013). These processes are complex and can work over timescales from hours to hundreds
of thousands of years. Some of these processes can be biologically mediated by bottom-dwelling
communities that ingest, burrow and otherwise transform the sediment, altering the sediment-
water interface. Most research on ocean acidification has focused on surface waters and does not
consider sediment processes, so this is seen as a major science gap in the coastal environment
where the effect of sediments on alkalinity should be considered.
8 | Mcas u i" i n. g € h a ngcs i n C oast a 1 € a r b o n a t c € h c m i s t r y

4.1 Overview
Concise overviews of the seawater carbonate system are provided in Dickson et al.
(2007) and Riebesell et al. (2010), both of which are available online. Seawater carbonate
chemistry consists of a system of chemical states and reactions by which CO2 gas, whether
originating from the atmosphere, water column or sediment is taken up and transformed in the
seawater. CO2 and its dissociation products HCO3" (bicarbonate ion) and CO32" (carbonate ion)
are further utilized in other important chemical and biological processes. CO2 entering the water
column via atmospheric exchange or produced during respiration combines with H2O to form
carbonic acid, which immediately dissociates into HCO3" and H+; when HCO3" further
dissociates into CO32" another H+ ion is released (Figure 1-1). In addition to being short lived,
carbonic acid only constitutes 0.1% of dissolved CO2 and thus, is not considered biologically or
chemically significant (Riebesell et al., 2010). With increasing concentrations of dissolved CO2,
dissociation reactions increase the concentration of H+ in the water, decreasing pH. If CO2
concentrations decrease, these reactions go in the opposite direction, decreasing the
concentration of H+ in the water and increasing pH. The equilibrium expressions below in
Figure 4-1 show the major stoichiometric reactions that control CO2 speciation in water.
1.	C02(g) = C02(aq)
2.	CO2 (aq) + H20(1) = H+(aq) + HCO3 (aq)
3.	HCO3 (aq) = H+(aq) + CO32 (aq)	
Figure 4-1: These equilibrium expressions are based on A. Dickson's chapter in Riebesell et al.
(2010), with the letters g, 1, and aq denoting gas, liquid, and aqueous solution, respectively.
DIC Speciation



, \
/ v
/ *

/ *

— co2 \
1 »
/ *
/ *
1 *
1 \







2	4	6	8	10	12
Figure 4-2. Concentration of inorganic carbon species changes as a function of pH. The above
Bjerrum plot was created using the seacarb package in R (Lavigne et al., 2011; R Core Team, 2017).
Dissolved CO2 is the predominant species at low pH but at pH values above 5, the concentration of
dissolved CO2 rapidly decreases. Conversely, the concentration of CO32" is near zero at low pH, and
it becomes the predominant species at pH values above 9.
4.The Seawater Carbonate System |9

Acidity is determined by the concentration of hydrogen ions. In seawater, hydrogen ion
concentrations are buffered (i.e., maintain a nearly constant concentration) by the interconversion
of bicarbonate and carbonate and concomitant release or consumption of a hydrogen ion (Figure
4-1, reaction 3); consequently, hydrogen ion concentration is largely controlled by the ratio of
bicarbonate to carbonate. However, the buffer capacity of seawater is finite and increasing
concentrations of carbon dioxide will ultimately increase the concentration of hydrogen ions.
The relative abundance of bicarbonate and carbonate over a range of pH values is shown in
Figure 4-2. In typical seawater, as hydrogen ion concentration increases (i.e., pH decreases),
carbonate ion concentration decreases and bicarbonate ion concentration increases.
Additional information on the seawater carbonate system can be found in chemical
oceanography textbooks (e.g., Pilson, 2013). Concise summaries can also be found in Dickson
et al. (2007), Riebesell et al. (2010) and Zeebe (2012). In addition, we find Wolf-Gladrow et al.
(2007) discussions of processes affecting alkalinity especially helpful. A simplified visual
representation of ocean acidification is given at, part of which
is also shown in Figure 1-1.
The well-studied seawater carbonate system of the open ocean and, to some extent,
nearshore waters, can be fully described by measuring only two of its four major (directly
measurable) parameters along with temperature and salinity. This often allows effort to be
focused on high quality measurements and equipment for two parameters. The equilibrium
constants used in making these calculations are critically dependent on high quality
measurements of temperature, salinity and pressure. Millero (2010) and Riebesell et al. (2010)
provide detailed descriptions of these constants and where/when they should be used (also see
section below on software packages that include documentation of the constants). Careful
selection and reporting of the units of scale for all parameters is also critical. pH, for example,
can be measured on a variety of scales, each of which has a different meaning and usage in the
carbonate system calculations. DIC, total alkalinity and dissolved CO2 can all be expressed in
gravimetric (micromoles per kilogram) or volumetric (micromoles per liter) units. In order to
minimize the chance for reporting or calculation error, all measured carbonate variables, with the
exception of pH, temperature and salinity, should be reported in micromoles per kilogram. This
straightforward conversion will be possible because sample salinity and temperature will be
4.2 Dissolved inorganic carbon (DIC)
DIC, sometimes referred to as total CO2, /C02, or ECO2, is the sum of the inorganic
carbon species that are dissolved in a solution.
DIC = [C02(aq)] + [H2CO3] + [HC03"] + [C032"]
The majority of DIC in seawater exists as bicarbonate and carbonate ions. At surface,
open-ocean equilibrium conditions, the DIC pool consists of carbon dioxide: C02(aq)* (~1%),
carbonate ion: CO32" (-10%) and bicarbonate ion: HCO3" (-89%). DIC is typically measured
using instruments that incorporate infrared, coulometric, and spectrophotometric detection (see
section on analyzing bottle samples for details).
10 | Mcas u i" i 11 g C h a ngcs i 11 C oast a 1 € a r b o n a t c € h c m i s t r y

4.3	Carbon dioxide (CO2)
Under equilibrium conditions aqueous CO2 (C02(aq)*) may only make up approximately
1% of the total DIC pool but is the most labile inorganic carbon species. The mole fraction of
CO2 (XCO2) can be directly measured using a non-dispersive infrared gas analyzer after
equilibration of gas and liquid phases in the measurement system (Frankignoulle et al., 2001;
Hales et al., 2004; Wanninkhof and Thoning, 1993). pCOi and /CO2 can be calculated from
XCO2 with other related parameters (e.g., temperature and pressure)and are often reported in
micro-atmospheres. [CC>2*aq] can be calculated from /CO2 via Henry's Law (Dickson et al.,
4.4	Total alkalinity (TA)
The measurement of alkalinity quantifies the ability of substances in seawater to react
with the addition of a strong acid and convert it to an uncharged species. For this reason, it is
sometimes informally denoted as "buffering" or "acid buffering capacity". At its simplest, TA is
the excess of proton acceptors over proton donors (Wolf-Gladrow et al., 2007) though it is
defined in many ways2. The TA of a system is dependent on a large number of seawater
constituents and can be expressed as follows:
TA= [HCO3-] + 2[C032"] + [B(OH)4-] + [OH"] + [HPO42"] + 2[P043"] + [SiO(OH)3"] + [NH3] +
[HS"] - [H+]f - [HSO4"] - [HF] - [H3PO4] ... (other undefined acid-base species)
Alkalinity tends to be pseudo-conservative with salinity (i.e., alkalinity does not dilute
linearly with salinity because there are biogeochemical processes that can add or remove
alkalinity). Generally, higher salinity waters (containing a greater concentration of salt and
carbonate ions) will have higher alkalinity (Millero et al., 1998) and a greater ability to neutralize
acidic inputs. In estuarine and near-shore systems, the contribution to total alkalinity by organic
alkalinity (derived from dissolved organic matter) can be significant (Cai and Wang, 1998; Yang
et al., 2015). TA can be measured through the use of acidimetric titration using potentiometric
or colorimetric methods (Dickson et al., 2007; Gran, 1952).
4.5	pH
Water molecules dissociate into a hydrogen ion [H+] and a hydroxyl ion [OH"]. These
hydrogen ions can exist freely in solution or combine with water ions. pH is a negative
logarithmic numeric scale used to specify the acidity of an aqueous solution.
pH = -logio[H+]
1	Aqueous CO2 concentration is proportional to CO2 fugacity (/CO2) via dissolution of CO2 gas into water
(Henry's Law). Values of /CO2 are similar to, but not the same as partial pressure of CO2 (pCOi). that would
be exerted by the CO2 if it existed at the observed temperature in isolation from other gasses (i.e., ideal gas)
in equilibrium with seawater. pC02 is the product of mole fraction of CO2 (XCO2) and atmospheric pressure,
accounting for the interaction with other gases in air. The major difference between pCO-j. and ./CO2 is that the
/CO: definition incorporates the non- ideal nature of CChgas, making /CO2 a more precise measure (Schuster et al.
2009b). In practical terms, differences between pCO-j. and /CO2 are often not discernible. In typical estuarine
conditions the difference is a few tenths of a percent, which is in the range of and often smaller than measurement
uncertainty except on state-of-the art equipment.
2	Dickson et al., 2007 defines seawater TA as "the number of moles of hydrogen ion equivalent to the excess of
proton acceptors (bases formed from weak acids with a dissociation constant K < 10-4.5 at 25°C and zero ionic
strength) over proton donors (acids with K > 10-4.5) in 1 kilogram of sample".
4 . T h c Sea w a t c r C a r fa o 11 a t c S v s t c 111 | 11

Because the pH scales are logarithmic, a one unit change in pH represents a tenfold change in
hydrogen ion concentration. In oceanic, coastal and estuarine systems, surface seawater pH
averages 8.1 ±0.1 (Millero, 2007). pH is a parameter that can be directly measured using
potentiometric or colorimetric methods.
There are four pH scales that are commonly used in seawater pH measurement, and the
lack of a consistently agreed upon scale can make interpretation between studies difficult
(Marion et al., 2011; Orr et al., 2009). Dickson (1993) provides an excellent overview of the
major pH scales and the suitability of each. They are determined by the method of calibration
and are a critical piece of reporting. The four scales are the NBS (National Bureau of Standards)
also known as IUPAC (International Union of Pure and Applied Chemistry) scale (pHnbs),
seawater scale (pHsws), free hydrogen scale (pHf) and total hydrogen scale (pH-r). The NBS
scale measures hydrogen ion activity; the seawater, free hydrogen and total hydrogen scales
measure hydrogen ion concentration but differ from each other by incorporating measurements
of different dissociation protons. These scales can be converted from one to the other, but the
calculations can introduce error. Choosing the correct scale is important and can make a large
difference in the final pH values, e.g. for a seawater sample at a salinity of 35 and a temperature
of 25°C the difference between pH values using the pHsws and pHi scales can be as much as 0.1
units, a sizeable difference on a logarithmic scale (Marion et al., 2011). The difference in values
between pHnbs and pHi can range from hundredths to tenths of a pH unit, but more importantly,
these differences are rarely consistent or predictable. Conversions from pHnbs to pHi are
sometimes possible but require rigorous calculations and values for activity coefficients and
liquid junction potential of the electrode (Pilson, 2013). Moreover, the relationship between the
total pH scale and the NBS scale typically used in handheld meters and sondes is not
conservative or straightforward because the two methods measure differing chemical properties.
For carbonate chemistry measurements/calculations in oceanic and estuarine systems
using established stoichiometric models, the total hydrogen scale is the most suitable partly
because of the high sulfate concentration in seawater. The total hydrogen scale (pHr) includes
the concentration of the free hydrogen proton as well as the proton that dissociates from
hydrogen bisulfate (HSO/f)
The NBS (or IUPAC) scale is optimized for glass membrane electrodes and uses NBS or
similar buffers. This method measures the free hydrogen ion activity, but the low ionic strength
buffers may not be suitable for non-freshwater systems (Dickson, 1984). In any case,
preparation of calibration buffers with ionic strengths that are close to those of the expected
sample is an important consideration in pH measurement (see ANALYZING BOTTLE
SAMPLES). The free hydrogen scale accounts only for the free hydrogen ion concentration.
The chemical composition of seawater makes calibration when using the free hydrogen scale
difficult by neglecting to account for the hydrogen ion from hydrogen bisulfate, and is not
recommended for seawater. Both the seawater scale and total hydrogen scale measure the H+
and HSO/fconcentration. The seawater scale additionally incorporates the concentration of HF
in the solution. Due to the relatively small differences between pHr and pHsws, it is clearer to
treat fluoride explicitly as a minor acid-base species when needed rather than to incorporate it
implicitly into the definition of pH (Dickson, 1993). While this document recommends using the
total hydrogen scale whenever possible and especially when high-quality data are the goal, one
must be mindful that the constants used in the calculation match the intended scale. Using
incorrect constants can contribute systemic error in calculations and reporting (Riebesell et al.,
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4.6	Bicarbonate
The bicarbonate ion is the result of the first dissociation reaction of carbonic acid into
seawater. At a pH ~8 bicarbonate makes up 89% of the inorganic carbon species constituting the
DIC pool. Bicarbonate is not currently directly measured, but is calculated using measurements
of at least two other parameters of the carbonate system along with salinity and temperature.
4.7	Carbonate
The carbonate ion is the result of the second dissociation reaction of carbonic acid into
seawater. At a pH ~8 carbonate makes up -10% of the inorganic carbon species constituting the
DIC pool. Carbonate is not typically measured directly, but is calculated from measurements of
at least two other parameters of the carbonate system along with salinity and temperature.
O 0.9
0.8 -
0.7 -
0.6 ¦


¦ co32-

1 co2( ,

Constant Pressure(1 atm)
TA (2143), pH (8.034),
Salinity (31.7)

0 5 10 15 20 25 30
Temperature (°C)
26 28 30 32 34 36 1400 1600 1800 2000 2200
Salinity	TC02
Figure 4-3. Partitioning of the carbonic acid system as a function of temperature, salinity, and
REPORTING")- Note that as DIC is increased either by atmospheric exchange (i.e. ocean acidification)
or by respiration, the relative proportion of carbonate ion decreases. Temperature and salinity also
impart important controls on carbonate ion availability (Dwight Gledhill, personal communication).
4.The Seawater Carbonate System | 13

4.8 Calcite and aragonite saturation states
This metric refers to the actual concentrations of calcium and carbonate ions relative to
their expected concentrations at equilibrium with their respective allotropic form of calcium
carbonate. Solubility product constants are used to describe saturated solutions of ionic
compounds of relatively low solubility. A saturated solution is in a state of dynamic equilibrium
between its constituent ions and the undissolved solid. For calcium carbonate, the solubility
product is expressed as follows:
Ksp = [Ca2+][C032-]
We refer to the degree to which seawater is saturated with respect to allotropic form of calcium
carbonate as saturation state, which is represented by Q (omega).
Q = [Ca2+][C032-]/Ksp
Saturation states less than 1.0 indicate "undersaturation" while saturation states greater
than 1.0 indicate "supersaturation" and are thus of interest for evaluating potential biological and
ecological effects, especially for calcifying organisms. Calcium carbonate dissolution is
thermodynamically favorable in undersaturated waters and calcium carbonate formation is
favored in supersaturated waters (note that effects are sometimes seen at saturation states well
above 1.0). Saturation states are calculated from measurements of at least two other parameters
of the carbonate system along with salinity and temperature measurements (see "CARBONATE
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Sample collection and preservation is often an overlooked part of water quality chemistry
but is an area where standardization is important and achievable. General standard operating
procedures (SOPs) are available and commonly used, but there remains considerable flexibility
within those SOPs that could affect sample quality (e.g., maximum time between sample
collection and sample preservation). This section discusses recommended sample containers,
collection techniques, and the pros and cons of varying preservation techniques. We recommend
use of the standard operating procedures described in Dickson et al. (2007) including those for
sample collection and preservation. Here, we provide only brief summaries of those methods
and focus on additional considerations for coastal sampling.
5.1	Sample containers
Sample containers should be chosen based upon the intended analysis, volume of sample
required for analysis, length of anticipated storage and collection method. Important
considerations in bottle choice include volume, leaching of bottle material, gas permeability,
opening size, neck size, and sealing (see sample storage sections below).
Whenever possible, volume should be collected in a sufficient quantity to perform the
expected analysis with additional volume for rinsing pipettes, instruments etc. Headspace should
be kept to a minimum while allowing for sample expansion. See the section below on storage
methods for additional information.
5.2	Cleaning
To minimize contamination of carbonate samples, the analyst must ensure that everything
that comes in contact with the sample (sample collection devise, sample bottles, pipettes) is
rigorously cleaned. The cleaning, rinsing and filling of sampling devices and sample containers
should be done in a standard and reproducible way in order to minimize inconsistencies that can
reduce precision. Ideally, they should be thoroughly cleaned with a 5-10% (v/v) solution of 36%
hydrochloric acid and then thoroughly rinsed (at least 3 rinse cycles) with deionized water to
remove acid residue (i.e., "acid washed"). Careful, thorough rinsing is an important step as
residual acid can greatly effect carbonate chemistry of the sample. Detergents are not preferred
but may be an acceptable alternative cleaning agent; however, their effects on sample quality
should be assessed beforehand (e.g., via analysis of sample blanks). Containers/caps should be
inverted and allowed to air-dry and capped or covered to minimize contamination. As with most
water quality sampling, nitrile or latex gloves should be worn to further minimize contamination.
For large pieces of equipment such as large rosette units, acid washing of equipment between
sampling stations is not feasible, so rinsing with sample water from the sample location becomes
especially important.
5.3	Preferred order of sample collection
The dissolved CO2 concentration in the water sample will rapidly equilibrate with the
atmosphere once a sample is collected; consequently, samples will change over time. The rate of
change is dependent on factors such as the sample volume, surface area, surface area to volume
ratio, CO2 gradient between the seawater and atmosphere, and biological activity. For this
reason, the minimization of atmospheric exchange during sample collection and storage is
paramount in obtaining accurate measurements of carbonate parameters. Some carbonate
parameters are more sensitive to atmospheric gas exchange compared to other measured
5.Collecting Bottle Samples | 15

parameters. Due to the low concentration of CO2 found in seawater (typically about 2xl0"5 mol
kg"1), dissolved CO2 measurements are extremely sensitive to atmospheric exchange. For this
reason, it is useful to collect samples in decreasing order of their sensitivity to atmospheric
exchange, which ispCOi> pH > DIC > TA. In order to minimize the level of atmospheric
exchange, Niskin, Go-Flow and Van Dorn sampling devices are common choices. After
retrieval, the device should be vertically oriented with its outlet lying below the vented top
surface, to prevent bubbles from entering the device. Flexible silicon, FEP or Teflon© tubing is
typically used to fill and transfer samples from the bottom outlet of the sampling device to the
sample bottles. The tubing should be rinsed with sample seawater prior to introducing it into a
sample bottle. Skin contact with the portion of the tube that enters the bottle should be avoided,
even in cases where it will be rinsed.
5.4	Filling sample bottles
Prior to filling, sample bottles cleaned as above should then be rinsed and emptied 3
times with the sample seawater. The flexible tube is then placed with its opening at the bottom
of the sample container and allowed to fill at a rate that minimizes visible turbulence in the
bottle. The sample bottle should be filled and allowed to overflow for approximately twice the
time needed to fill the container to the top (e.g., if it requires 5 seconds to fill the bottle, the
overflow process should continue for 10 seconds). If bubbles enter the sample, the entire
overflow process should be repeated.
The tube can be slowly withdrawn from the sample bottle and the sample can be closed
with a screw cap, ground glass stopper or conical displacement cap. Normally, a headspace of
about 1% of the total sample volume is left to allow for sample expansion. This can be achieved
by pinching the tube before it is completely withdrawn.
Samples should not be taken for pCOi, pH and DIC when the water volume in the
sampling device falls to -25% of the total device volume, due to the potential of atmospheric
exchange significantly altering the sample. In some cases, and for certain purposes, bucket
collection of seawater will be the only option. However, the same principles of rinsing and
reducing surface exposure and turbulence apply.
5.5	Long-term storage bottles
The preferred method for long-term storage (>1 months) of samples uses a narrow mouth
borosilicate glass necked bottle with a ground borosilicate glass stopper (Dickson et al., 2007).
Apiezon high vacuum grease (or similar) is applied in a thin layer on the stopper which is
inserted into the bottle and turned to completely coat the surfaces where the stopper contacts the
bottle. To further ensure there is no gas exchange nor stopper displacement the stopper can be
secured with positive pressure (e.g., using an elastic band or plastic strap).
5.6	Short-term storage bottles
For short-term storage of samples (<1 months) borosilicate glass screw thread sample
vials are a suitable storage option. Huang et al. (2012) demonstrated that DIC concentration of
samples preserved and stored in a specific type of screw cap vial vs. borosilicate stoppered
bottles were not statistically different over the 148-day study period. These vials are available in
a variety of volumes and can be ordered with caps that are tailored for the expected analysis.
Solid caps with Teflon © liners retard gas exchange with the atmosphere.
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5.7	Sample preservation and shelf life
A saturated mercuric chloride solution (in DI water) is the standard preservative used for
seawater carbonate system sampling. Preserving a sample is sometimes also referred to as
"fixing" a sample. Immediate preservation is advised in SOPs, but there is no clear guidance on
how quickly samples degrade after collection and what exactly constitutes immediate
preservation (i.e., delays less than seconds, minutes, or hours). This will be addressed in future
versions or additions to these guidelines but is largely related to the relative productivity of the
waters. Water samples with high levels of nutrients and organic matter probably require more
immediate fixing. Freezing is not an acceptable method of preserving samples. In any case, it
has been suggested that a well preserved sample in a high quality sealed bottle has a shelf life of
at least one year (e.g., Tris samples:
5.8	Sampling without preservation
In some cases, if the samples are to be analyzed quickly (i.e., within a few hours),
preservation of properly handled samples may be unnecessary. Proper handling of samples can
include keeping them cool and dark after collection to slow down biological activity. We
recommend storing samples in a refrigerator (Dickson et al., 2007) or a cooler with ice or cold
5.9	Transport
Sample containers should be stored in a cool, dark environment (e.g., a cooler). Shipping
of samples preserved with mercuric chloride or other toxic chemicals requires special handling
and labeling. Mercuric chloride preserved samples fall under Department of Transportation
Hazard Class 6.1; poisonous materials must be transported in compliance with 49 CFR 173.132.
Refer to these Department of Transportation guidelines for further information.
5.10	Other sampling equipment
As noted above, latex or nitrile gloves should be worn when sampling to prevent
contamination of the water sample. If the sample is to be fixed with a hazardous preservative,
personal protective equipment (PPE) that is appropriate for working with that preservative
should be worn. This typically includes nitrile gloves, eye protection and a lab coat.
5.11	Filtration
Instruments with small diameter syringes and injection ports or combustion columns or
optically-based measurements may provide improved results if samples are pre-filtered.
Filtering, however, can have a deleterious effect on carbonate measurements for pCOi, pH and
DIC. Samples are subjected to pressure changes and increased turbulence in most filtering
apparatus. Benchtop and syringe filtering using glass fiber filters should not be used because
their use will lead to atmospheric exchanges and, thus, alter the sample. Bockman and Dickson
(2014) described a method using a peristaltic pump in conjunction with a membrane filter to
remove phytoplankton and CaCCb particles from a sample without altering the carbonate
chemistry of the sample. This method may be useful when trying to characterize systems high in
biological and inorganic particulate matter but, if used, details need to be carefully reported.
In-line barrel filters, such as those used in groundwater sampling (0.45 |im), properly
filled with the water being sampled are sometimes used in the bottle-filling process when
filtering large volumes of water. Effects on sample quality are not known so we do not
5 .	:i pics | 17

recommend this method. In cases where water is so turbid that such filtering is necessary, other
issues emerge that would only be addressed by advanced chemistry labs equipped for "over
determination" of the carbonate system.
5.12 Analyte-specific sampling techniques
pCCh: Borosilicate glass should be used along with a cap that is impermeable to gas
exchange. As already noted, pCOi samples should be collected with a headspace approximately
1% of the bottle volume. This is to prevent possible breakage due to the expansion of the sample
if its temperature increases. Required sample volume is instrument and analysis dependent.
Systems for measuring discrete pCOi samples are not common and are often custom made in
research laboratories (Hales et al., 2004). Recent evaluations of a commercially available CO2
probe can be found in Moran et al. (2010) and Pfeiffer et al. (2011). More commonly, pCOi
measurements are taken using a flow-through system. Flow-through systems present a separate
set of protocols to be followed to ensure that a representative environmental sample is measured.
pH: In general, there is no accepted method for preservation of pH samples, so many
investigators opt for direct in situ measurement or immediate analysis of collected samples (i.e.,
within seconds or minutes). A recent study (Chou et al., 2016) examined the viability of
preserving pH samples in the field for later lab analysis, so this method deserves further
consideration. pH samples should be collected with minimal headspace and should be stored in a
borosilicate glass container with a cap that is impermeable to gas exchange. The amount of
sample volume that is necessary to collect is instrument and analysis dependent. Samples
collected for analysis using the potentiometric method must be collected with minimal headspace
and the volume should be sufficient to allow the electrode to be immersed. Due to gas exchange
driven pH changes and ion consumption by the electrode, a larger volume of sample collected
will potentially give more stable and precise readings. When performing a measurement in an
open container, the surface area to volume ratio of the sample should be minimized (to reduce
gas exchange) by choosing smaller diameter/tall vessels vs. larger diameter/short vessels.
Electron starvation and gas exchange become more of an issue when there is large variation
between samples and thus longer times for stabilization of the reading.
DIC: Each DIC sample bottle should be filled as described above and stored in a
borosilicate glass container with a cap that is impermeable to gas exchange; samples with high
pCOi are especially susceptible to gas exchange concerns. The amount of sample volume that is
necessary to collect is instrument and analysis dependent although most commercially available
systems require a sample volume of 5-25 mL. The need for splitting samples or repeat analysis
is also a factor.
Total Alkalinity: Samples for total alkalinity can be collected in borosilicate glass or
HDPE sample containers. Borosilicate glass is preferable to "softer" glasses for sample storage.
Storing samples in soda lime glass can cause significant increases in alkalinity concentrations as
seawater can leach sodium and other compounds from the glass over time (Huang et al., 2012).
Alkalinity samples are less sensitive to gas exchange than DIC samples so it is less critical but
still advisable to collect them in gas impermeable containers. The amount of sample volume that
is necessary to collect is instrument and analysis dependent. A minimum volume of 40 mL is
recommended although accuracy will be greater with a larger sample volume.
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All of the measured carbonate parameters discussed below can be analyzed using
different protocols and procedures. There are existing comprehensive resources that go into
great detail regarding carbonate system measurement principles and method SOPs (Dickson et
al., 2007; Martz et al., 2015; Schuster et al., 2009a). For most of these parameters, there exists
an array of instruments ranging from highly customized "research instruments" to off-the-shelf
commercial instruments. Martz et al. (2015) provides a good overview regarding the availability
and intended usage of instruments measuring the seawater carbon system. For all analysis,
standard or certified reference materials (CRMs) are recommended to quantify sample precision
and accuracy.
6.1	Certified reference materials (CRMs)
The importance of CRMs for carbonate system analysis cannot be overstated. Careful
use of certified reference materials can dramatically improve measurement precision and
accuracy. Martz et al. (2015) noted that establishment of the CRM program reduced residual
errors in DIC from -14 jamol kg-1 during GEOSECS (Bradshaw et al., 1981) to ~3 jamol kg-1
during WOCE/JGOFS/OACES (Lamb et al., 2002). Andrew Dickson's Marine Physical
Laboratory at the University of California at San Diego provides CRMs that currently serve as
the industry standard for the traceable QA in inorganic carbon system analyses. These CRMs are
provided for DIC and alkalinity in sterilized natural seawater; uncertified reference materials for
pHi are made in artificial seawater. These CRMs have been widely used in oceanographic
research beginning around 1990. DIC and alkalinity concentration values and salinity are
provided along with the concentration of phosphate, silicate, nitrite and nitrate (all values are in
micromoles per kilogram). Salinity is usually close to 33 PSU so care must be taken when using
these oceanic reference materials in low ionic strength estuarine systems (i.e., < 20 PSU).
Detailed information on each batch and storage instructions can be found at
( Some investigators prepare
inorganic carbon or borate solutions in seawater and then analyze their DIC and alkalinity
against the CRMs. These can then be used as secondary standards for monitoring drift during an
analysis session, thereby reducing the use of CRMs and their associated waste stream. However,
we recommend that all DIC and alkalinity sessions include a minimum of at least two CRM
determinations (beginning and end). If the electrode system is affected by switching salinity
between samples, it may be beneficial to try measuring a diluted (by weight) CRM as well as an
off-the-shelf one to confirm that such a problem does/or does not exist (Dickson, personal
6.2	Alkalinity
Relative to the inorganic carbon parameters, total alkalinity tends to remain stable during
short term exposure to the atmosphere during sample collection and analysis. However,
alkalinity measurements require careful determination of salinity and temperature. Care must be
taken to use an instrument that incorporates these parameters into the measurement or one must
record these parameters using calibrated sensors and incorporate them into final calculations.
Commercial alkalinity measurement equipment is available and can be purchased for
approximately $10,000 to $20,000, depending on extra features (Riebesell et al., 2010; also see
6.Analyzing Bottle Samples | 19

IOCC instrument list). See Table A-l for typical precision goals. Two methods are described
Potentiometric method: A seawater sample is placed in a thermally jacketed, open container.
A known amount of a hydrochloric acid titrant solution is added to the sample. During titration,
this acid titrant reacts with the alkaline species present and drives the pH down below 3. Once
pH decreases past the second equivalence point, the alkalinity is back calculated based on the
sample weight, amount of hydrochloric acid titrant solution added, corrections for preservative
effects and the molarity of the titrant solution (typically requiring a Tris calibration; see Dickson
et al., 2007; Gran, 1952). The molarity of the hydrochloric acid titrant solution can be
determined using Tris [tris (hydroxymethyl) aminomethane, an organic compound with the
formula (HOCEh^CNEh)] buffers (Marion et al., 2011) prepared in an NaCl solution of similar
salinity (approximately) to that of the samples being analyzed. During the alkalinity titration, the
temperature of the sample should be held constant and near to 25°C. This can be done using a
jacketed beaker plumbed to a temperature-controlled water circulator. Continuous mixing using
a magnetic stir bar or similar is necessary to ensure that the titrant is uniformly dispersed and the
electrode reading is representative.
Potentiometric method considerations
•	Benefits: High precision (±0.1% at 2000 |imol kg"1). Widely used and well-documented
method. Good analytical stability (Riebesell et al., 2010).
•	Disadvantages: In samples with high nutrient and/or dissolved organic matter
concentrations (e.g. eutrophic estuarine systems), interpreting alkalinity titrations can be
difficult, refer to Riebesell et al. (2010).
•	Careful determination of titration reagent molarity is required.
•	Temperature control (±0.1°C) is required for high precision.
•	Samples should be kept in a temperature-controlled bath prior to analysis, and titration
should be carried out in a temperature-controlled jacketed beaker.
•	Acid addition should proceed slowly (0.01 mL/second) so that evolved CO2 gas can
escape solution. Gas exchange can be facilitated through the gentle bubbling of air (or
•	CRM: Dickson seawater reference material.
Colorimetric method: The colorimetric method of determining total alkalinity was commonly
used for environmental samples in the past and there is ample documentation from USGS and
EPA (USEPA, 1974) on this method. Colorimetric alkalinity measurements work on the
principle of using a reagent that contains a known quantity of acid and a color indicator, which is
typically methyl orange. The color change that occurs when a known amount of sample water is
added to the reagent is used to calculate the alkalinity. The acid-methyl orange indicator works
under the principle that "any addition of alkalinity causes a loss of color directly proportional to
the amount of alkalinity. This method is applicable to drinking, surface, and saline waters,
domestic and industrial wastes. The applicable range is 10 to 200 mg/L as CaCCb" (USEPA,
1974). There is an autonomous system currently under development using bromocresol purple
indicator dye as the color reagent and calculating alkalinity via a spectrophotometer (Spaulding
et al., 2014).
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Colorimetric method considerations
•	Disadvantages: Slow analysis time (individual sample time is > 10 minutes). High
turbidity can interfere with analysis. Interpreting alkalinity can be difficult (Riebesell et
al., 2010).
•	This method is not currently widely used in monitoring environmental samples.
•	High purity indicator dye must be used.
•	CRM: Dickson seawater reference material.
6.3 Dissolved inorganic carbon
DIC measurement requires careful determination of salinity and control of temperature.
DIC can be measured using either infrared or coulometric methods. Because of its shorter
processing time, the infrared CO2 analysis method is more commonly used. As with alkalinity,
quality assurance is much improved through the use of CRMs (see Certified Reference Materials
CRMs). For both methods, samples are preserved, stored and prepared in the same way.
Inorganic carbon standards are sometimes prepared or purchased but this practice is rare
because of the lack of traceability and the unknown sources of error it may introduce.
Commercial dissolved inorganic carbon measurement equipment is available and can be
purchased for roughly $50,000 (Riebesell et al., 2010).
Infrared method: When measuring DIC, a known quantity of a water sample is dispensed by
the instrument into a stripping chamber. The water sample is then acidified, typically with
phosphoric acid, to a pH below 3.0. By decreasing the pH to less than 3.0, all of the inorganic
carbon species in solution (H2CO3, HCO3", CO32") are converted to CO2, which is then removed
from the acidified solution by sparging with a high purity, CO2 free, inert carrier gas (N2 is
commonly used). This C02-containing carrier gas is passed through a drying column, to remove
water vapor, and then is transferred to the built-in non-dispersive infrared gas analyzer for
measurement of CO2 concentration. CRMs are used for calibration of the DIC analyzer.
Infrared method considerations
•	Benefits: Small sample volume needed (>lmL). Short measurement time (-10-15
minutes). High precision (0.05-0.1% CV). One reagent needed, and smaller amount of
hazardous waste produced.
•	Disadvantages: Expensive (but similar cost to coulometric). Calibration is less stable
than coulometric and thus more frequently required. Instrument drift must be evaluated
using periodic check standards (i.e., every five or ten samples).
•	Temperature control (±0.1°C) is required for high precision.
•	Samples should be kept in a temperature-controlled bath prior to analysis and
temperature-controlled jacketed beakers should be used for the reagents and samples.
•	CRM: Dickson seawater reference material.
•	CRM stability is subject to headspace equilibration. Portion large CRM bottle into many
smaller, gas-tight containers.
Coulometric method: The preliminary sample manipulation steps for determining DIC
coulometrically are similar to the infrared DIC method (i.e. acidifying, sparging, drying). The
difference lies in how the CO2 entrained in the carrier gas is measured. The amount of CO2 in
the carrier gas is measured by trapping the CO2 in an absorbent containing ethanolamine and
6 . A 11 a 1 y z i 11 g Bo 111 c S a 111 pics | 21

coulometrically titrating the hydroxyethylcarbamic acid that is formed. Further information on
coulometric methods can be found in Lindberg (1978) and Johnson et al. (1985).
Coulometric method considerations
•	Benefits: High Precision (± 1 |imol kg"1). Longer instrument stability following
•	Disadvantages: Expensive but similar to infrared; greater number of chemical reagents
needed and larger volume of waste produced. Sample run time is longer than IR method.
•	Preferred CRM: Dickson seawater reference material.
•	CRM stability is subject to headspace equilibration. Portion larger CRM bottles into
many smaller, gas-tight containers.
6.4 pH
Water naturally dissociates into its constituent positive (H+) and negative (OH") ions. pH
measurements are an attempt to quantify hydrogen ion (H+) activity or concentration. However,
as discussed above, there are varying scales on which pH is measured and expressed. Seawater
pH measurement using the potentiometric method can be accomplished using a number of
different electrode configurations along with newer technologies such as solid state sensors.
Potentiometric pH measurement requires periodic calibration with a carefully selected calibrant
buffers suited to the measurement method, pH scale to be used and expected environmental pH
and ionic strength. Unlike the potentiometric method, spectrophotometric pH measurement
(Clayton and Byrne, 1993) employs a pH-dependent dye and does not require calibration with a
buffer. However, we recommend periodic checks using a standardized Tris buffer solution. For
both the potentiometric and spectrophotometric methods, temperature needs to be carefully
recorded or controlled and salinity needs to be recorded. Because pH is affected by CO2 flux
between the sample and atmosphere, laboratory pH measurements should be carried out in closed
containers. When measuring pH directly in a water body using a portable probe, this is not
Colorimetric/spectrophotometric method: In spectrophotometric pH measurements, which
measure total pH, a known quantity of purified m-Cresol purple, a sulfonphthalein acid-base
indicator, is added to a seawater sample (Clayton and Byrne, 1993; Dickson et al., 2007; Easley
and Byrne, 2012; Liu et al., 2011). Mosley et al. (2004) discusses applications to a wide salinity
range. Absorbances are measured with a spectrophotometer at the maximum absorption
wavelengths corresponding to the different forms of the indicator dye, which dissociates into
acid and base forms with distinct light absorbance spectra. This dissociation, and hence the
absorption ratio, depends on the pH of the sample (including the effect of the sulfate ion). The
calculation to determine total pH is made from the measured absorbance ratios, blank correction,
pH perturbation (due to dye addition), temperature, salinity, pressure and dissociation constant of
the indicator dye. Temperature must be tightly controlled (±0.1°C) because of its effect on pH,
so jacketed cells are often used. Dye impurities are known to affect the accuracy and precision
of measurements, so highly purified dye (e.g., from R. Byrne's lab at the University of South
Florida) are typically required for high quality measurements. Mosley et al. (2004) found that
precision using this method was better at salinities greater than 30 (±0.0005) as opposed to
salinities under 5 (±0.002). Additionally, correcting the dye pH to match expected sample pH
will be more difficult in highly variable estuarine systems.
22 | M c a s 111" i 11 g € h a ngcs i 11 C oast a 1 € a r b o n a t c € h c m i s t r y

Colorimetric/spectrophotometric method considerations
•	Benefits: High precision (±0.002 to ±0.0005) can be one or two orders of magnitude
better than glass electrode method. Measures pH-r, the preferred scale for seawater.
•	Disadvantages: Required equipment is more expensive and less portable than pH
electrodes. Sample analysis time greater than with handheld electrode. Limited
availability of purified m-Cresol dye, no batch tracking or certification. Making Tris
buffers can be time consuming and difficult.
•	Temperature control (±0.1°C) is required for high precision.
•	Samples should be kept in a temperature-controlled bath before analysis and a jacketed,
temperature-controlled optical cell should be used in the spectrophotometer.
•	Dye pH is subject to change, prior to use it should be adjusted to 8.1 ±.05 (e.g., using HC1
or NaOH). Some investigators argue that the dye should be close in pH to the samples,
so that the perturbation effect of the dye on the sample is small.
•	Dye perturbation analysis (measuring the pH change caused by the addition of dye alone)
should be carried out at least once per week (R. Byrne, personal communication).
•	Preferred CRM: Tris buffers are available as reference materials from the Dickson lab at
Scripps (not certified) or can be prepared in artificial seawater with considerable effort
and potential for error (Nemzer and Dickson, 2005). Tris buffer salinity should be
adjusted to match sample salinity.
Potentiometric method: In general, potentiometric pH measurement utilizes electrodes to
measure electromotive force, i.e., the difference in voltage measured between two points with
different electrical potential values. Electrodes for pH measurement must be ion selective for the
H+ ion and will consist of a reference and a glass electrode. These electrodes can be combined
(typical in handheld pH meters) or separate. The three main types of pH electrodes in use are
hydrogen, metal, and glass. Of these, glass electrodes are the most frequently used (the term
glass electrode is a description of the membrane material, i.e., glass membrane, rather than
of the material used to construct the electrode). Combination glass/reference pH electrodes are
commonly used but measurement precision can be improved by employing separate glass and
reference electrodes (Dickson et al., 2007). The reference electrode is contained in a medium
that provides a constant electrical potential value that is not affected by the pH of the sample,
while the glass electrode returns an electrical potential value that is dependent on the hydrogen
ion activity in the sample. These two electrical potential values determine the pH based upon
the premise that the system provides a Nernstian response. Because the electrochemical
potential measured by a glass electrode system is dependent upon an external liquid junction
potential, which is sensitive to changes in the seawater environment (e.g., salinity, temperature),
frequent re-calibration (weekly or more frequently; Easley and Byrne, 2012) is necessary to
ensure that the electrodes are not exhibiting non-Nernstian behavior. Dickson et al. (2007)
describes the methodology for obtaining accurate total hydrogen scale pH measurements from
glass electrodes. This method is more complex and requires more involved and frequent
calibration compared to the spectrophotometric method, but the equipment cost is smaller.
We generally do not recommend measurement of pH on the NBS scale, except in very
limited circumstances (e.g., educational settings, monitoring of qualitative patterns to
compliment other carbonate measurements). Within stable water conditions, some investigators
will calculate a regression between a pHnbs and a pHi determination to allow conversion, but
this is only reliable if the biological and chemical composition of the sample source is similar to
6 . A na 1 y z i n.g Bo 111 c S am pics | 23

the samples used for the regression. There are several different pH buffers that are either
commercially available or can be prepared in-house. NBS (or IUPAC) buffers are available
from a variety of manufacturers. These buffers are suited for calibrating electrodes on the NBS
scale only; if using NBS buffers it is recommended to use those in individual foil packets. pH
buffers will start to degrade when opened, whereas unexpired buffers in foil packets ensures
fresh buffers for each calibration (after dispensing into suitable beakers). Researchers measuring
ocean pH are more likely to use Tris buffers (Marion et al., 2011). Tris has a pKa of 8.07, which
makes it an effective buffer in the pH range of 7 to 9 and can be prepared in artificial seawater at
a variety of ionic strengths (matching the ionic strength of the intended sample is necessary to
minimize residual liquid junction errors [Dickson et al., 2007]). Calibrating with Tris buffers is
acceptable for the total hydrogen pH and free hydrogen pH scales (Millero, 1986), but low
quality Tris preparations may introduce as much error as the methods such as the regression
approach described above. A detailed description of Tris buffer use, including information on
stability preparation can be found in (Nemzer and Dickson, 2005). Glass electrodes can also
be calibrated to total pH using spectrophotometric pH measurements (Easley and Byrne, 2012).
More discussion of the varying pH scales in included in the section on the seawater carbonate
Potentiometric method considerations
•	Benefits: Equipment is easy to use in field conditions. Fast response time -30 seconds.
Comparatively inexpensive. Many commercial systems can be set up to continuously
monitor and record pH measurements.
•	Disadvantages: Precision is not as good as spectrophotometric method. Requires
frequent calibration. Making Tris buffers can be time consuming and difficult.
The availability of inexpensive meters contributes to insufficient calibration practices
and misinterpretation of data.
•	Electrodes must be periodically monitored for stability/degradation and non-Nernstian
•	Preferred CRM: Tris buffer in artificial seawater.
•	If the only available buffers are NBS, opt for those sold in single use foil packets and
try to calibrate at the temperature specified in the buffer label.
Solid state sensors: All solid state pH sensors are constructed using metal oxide
semiconductor field effect transistors and are typically referred to as ion-sensitive field-effect
transistor (ISFET) instruments. When the metal oxide semiconductor is exposed to an aqueous
solution, the exchange of protons produces an interfacial potential on the ion selective electrode.
The voltage difference between the ion selective electrode and either the internal or external
reference electrode is converted to a pH measurement (Bresnahan et al., 2014). These sensors
are used in a wide variety of chemical, industrial and environmental uses. Martz et al. (2010)
and Bresnahan et al. (2014) describe best practices of sensor conditioning and calibration.
ISFET instruments designed for handheld use and autonomous application are commercially
available. As of this writing, Honeywell Durafet probe and meter cost is $2500-$3000.
24 | M c a s u i" i n. g € h a ngcs i n C oast a 1 € a r b o n a t c € h c m i s t r y

Solid state method considerations
•	Benefits: Easy to use and requires minimal training. Fast response time -30 seconds.
Many commercial systems can be set up in a continuous monitor and record pH
measurements. Can remain stable for multiple months when continuously deployed
in seawater (Bresnahan et al., 2014). Precision of well calibrated instrument comparable
to spectrophotometric method.
•	Disadvantages: Expensive compared to potentiometric method. Lengthy time period
required for sensor conditioning and calibration (Bresnahan et al., 2014).
•	Not intended for use with bottle samples.
•	Deployed systems subject to biofouling and/or sedimentation.
•	Preferred CRM: paired, lab analyzed samples.
6.5 pCOi
Membrane/infrared method: For instruments that utilize infrared detection, the first step in
measuring pCOi is to extract the CO2 dissolved in the water sample into the gas phase. This can
be accomplished using headspace equilibration (Frankignoulle, 1988; Frankignoulle and Borges,
2001; Hales et al., 2004). A CO2 free gas (usually CO2 free air) is introduced into a gastight
chamber with the water sample. The water and air is agitated together for a period of time
(~1 min) until the CO2 concentration in the water and the headspace are equivalent.
Alternatively, many commercial instruments use a gas permeable membrane that is immersed in
seawater and CO2 can passively diffuse across this membrane. Once the CChin the gas phase,
the gas is transferred to an optical cell for measurement via non-dispersive infrared analysis
(NDIR). The infrared absorbance measurement of the CO2 gas sample is compared to the
reference measurement, which does not contain CO2, that allows calculation of a pCOi value
because sensor absorbance response complies with the Beer-Lambert law. Alternatively, a set of
calibrated reference gases can be used to generate a standard curve. Most commercial pCOi
measurement systems are in situ or flow through systems. The measurement of pCOi is
extremely sensitive to changes in temperature and pressure; consequently, this analysis is well
suited to an in situ application.
Membrane/infrared method considerations
•	Benefits: Continuous measurement of pCOi over weeks/months. Does not require
frequent calibration.
•	Disadvantages: Accurate sensors are expensive. Many of these systems are of the
homemade variety, extensive documentation regarding calibration and usage may not be
•	Not intended for use with bottle samples.
•	Deployed systems subject to biofouling and/or sedimentation.
•	Preferred CRM: Calibrated gas cylinder.
6 . A 11 a 1 y z i 11 g Bo 111 c S a 111 pics | 25

In situ and autonomous instrumentation is a rapidly developing area of research and will
be covered in greater detail in future versions or additions to these guidelines. Although
commercial options are available, many customized instrument configurations are being
developed and used in autonomous applications.
There is commercial instrumentation suitable for extended periods of data collection for
both pH and pCOi. Martz et al. (2015) and Byrne et al. (2010) provide good overviews of the
commercially available autonomous sensors and principles of measurement. Autonomous DIC
analysis is a developing field and as a result, instrumentation tends to be custom built by
individual research groups. Liu et al. (2013) describes an autonomous instrument that uses
spectroscopy to measure ambient DIC. Bell et al. (2011) describes a deployable membrane inlet
mass spectrometry instrument that can be used for DIC analysis. For alkalinity measurement,
Spaulding et al. (2014) describes a SAMI-ALK system that is in development and is suitable for
autonomous determination of alkalinity for a period of up to one month. Those readers
interested in autonomous and in situ instrumentation will have to consider the rapid
developments in this field and monitor new developments. Jones et al. (2016) recently reported
good results using autonomous measurements of pH and an alkalinity approximation from sea
surface temperature and salinity for their California coastal site.
Getting high quality, precise carbonate system measurements from in situ and
autonomous instrumentation has several challenges. The instrumentation used will need to be
carefully calibrated either in the lab prior to deployment or autonomously, to provide the best
data. Several of the commercially available sensors are factory calibrated but the majority of
users surveyed in Martz et al. (2015) conducted their own calibrations using CRM's or paired,
lab-analyzed samples. Biofouling is a persistent issue with any deployed instrumentation. The
severity of biofouling can vary between environments and is affected by fouling species present
and seasonal differences in settling rates. Biofouling resistance is understood to limit the
performance of the instrument, especially as deployment time increases (Byrne et al., 2010). For
carbonate chemistry the biofouling effect includes the changes in pCOi (and thus pH and DIC)
that can be caused by respiring organisms.
26|Measuring Changes in Coastal Carbonate Chemistry

Unmeasured carbonate system parameters can often be calculated from measured
parameters. For the open ocean, it is widely accepted that the measurement of any two carbonate
parameters can be used with salinity, temperature and pressure to "constrain" the carbonate
system and calculate the remaining parameters. This depends on a number of tightly constrained
thermodynamic constants that suit only specific environmental conditions and which govern the
dissociation reactions in the seawater carbonate system. The computations have been
implemented in a number of open source software tools, each with good documentation on how
to select the appropriate constants before performing calculations (ten such packages were
reviewed in Orr et al. 2015). In some cases, these constants are less appropriate for coastal
monitoring. Currently, the ki and k2 constants developed by Millero (2010) are often considered
suitable for carbonate system calculations at estuarine salinity. The Dickson and Riley (1979)
option for kf, and the Dickson (1990) option for ks, are available for most calculation packages
and are shown to be suitable over a large salinity range (Millero, 2010; Orr et al., 2015). In
addition, CRMs used for quality control are prepared in full strength seawater and are considered
inappropriate at low salinities. Because of this, and since there appears to be no consensus on
the best pair of carbonate parameters to monitor in the coastal environment, many investigators
"overdetermine" the carbonate system by measuring at least three parameters. Over-
determination allows identification of cases where one pair of parameters produces a different
picture of the carbonate system than a different pair. If measurement error or instrument
interferences (e.g., biofouling; colored dissolved organic materials effects on optical
measurements and non-carbonate alkalinity) can be ruled out, such an outcome suggests that the
open ocean stoichiometry may be poorly suited to the study site and that higher level calculations
(e.g., aragonite saturate state) should be performed with caution.
When reporting these higher level calculations (e.g., aragonite saturation state), it is
possible and useful to incorporate the uncertainty in the underlying parameters. As noted in
Martz et al. (2015), error propagation from measurement to final calculations would be a useful
addition to available software (see "Carbonate System Calculations"). This allows attribution of
the dominant sources of uncertainty, which informs selection of the set of carbonate parameters
that will most effectively minimize uncertainty in monitoring programs (see Table A-l).
While some investigators will have the opportunity to overdetermine the carbonate
system by measuring three or more parameters, others may be limited to only one parameter. As
monitoring data accumulate, there is the possibility that proxies will be identified that can make
such data useful, so these efforts should not be discouraged. For example, alkalinity can
sometimes be approximated from pre-existing alkalinity-salinity regressions (Carter et al., 2016;
Gledhill et al., 2015; Lee et al., 2006; Millero et al., 1998), andpCC>2 or JCO2 can sometimes be
crudely predicted using production/respiration calculations or proxies thereof, such as changes in
dissolved oxygen concentrations (e.g., Sunda and Cai, 2012).
9.Sampling Concepts and Uncertainty | 27

9.1 Precision and accuracy
Uncertainty is a critical component of all environmental data reporting and is usually
expressed in terms of accuracy and precision. These terms are sometimes incorrectly used as
though they are interchangeable. However, they have very specific and different meanings that
are important to keep in mind when planning, collecting, or interpreting observations
(Figure 9-1).
Figure 9-1: Precision vs. accuracy.
Accuracy refers to how close a given obseivation or measurement is to the true value,
whereas precision refers to how close multiple measurements are to each other (i.e.,
repeatability). An instrument that is always biased away from the true value would be
considered inaccurate but if this bias is extremely consistent, the instrument would also be
considered highly precise. An instrument that rarely returns the true value but which, after many
randomly varying measurements, produces an average value that is close to the true value, would
be considered imprecise but accurate when a sufficient number of measurements is averaged.
Imprecise methods are sometimes useful, accessible and/or inexpensive, but are problematic if
replication is too low. These same principles of accuracy and precision apply to environmental
sampling methods. For example, repeated measurements of pFI at a small number of fixed
locations in a large water body may be extremely consistent (i.e., high precision), but accuracy of
the mean and variance estimates depends on whether the sample times and locations are
representative of the point, water body, or parcel about which generalizations are being made.
Calculation of final sample precision and accuracy involves detailed methods for combining
error sources into a single estimate (Ellison and Williams, 2012) but is beyond the scope of these
The sources of uncertainty affecting a set of samples can be lumped into observation
error and process variability. Process variability is the result of deterministic processes (e.g.,
biogeochemical reactions, spatial heterogeneity). Since these processes can operate on tiny
scales of space and time, it is theoretically impossible to sample a process more than once before
it changes state. Nonetheless, duplicate sampling is an important component of quality
assurance because it helps to identify sources of error. Samples intentionally taken over wider
28|Measuring Changes in Coastal Carbonate Chemistry

intervals or spaces serve to describe process variability over the sampling domain or water parcel
from which the final reported values are averaged.
Sample handling, analyst skill, preparation and storage of reagents used in instrument
calibration, and instrument error are some of the obvious contributors to observation error, but as
described above, the total variance of a set of samples also includes the effect of process error.
Thus, a well-designed sampling program will include quality control measures to assess and
constrain these sources individually. As already noted, sample duplication is one example,
where multiple samples are collected simultaneously (to minimize the effect of process error)
and compared to assess observational precision (i.e., repeatability). Samples may also be split
and repeatedly analyzed to assess instrument precision. Many instruments perform this
repetition internally. Analysis of "blanks" (e.g., deionized water) are often used to assess
contamination during sampling, handling and analysis. And as discussed below, analyses of the
carbonate system should be referenced to certified materials to assess accuracy.
Although the dissection of total variance into observation error and process error is rarely
reported, it can be extremely useful for improving and fine tuning sampling programs. For
example, if observation error can be estimated through methods like the ones discussed above,
the remaining error is conventionally attributed to process error. The investigator can then focus
sampling design on capturing or constraining the effect of this process noise on inferences and
calculated endpoints. Although these are well established concepts in sampling design, our sense
is that the study of coastal acidification could benefit from more systematic application of these
As already noted, pHr typically varies much more in the coastal environment than in the
open ocean. When designing a sampling program, there is usually a trade-off between accuracy
and precision that depends on this variability and the purpose of the study. For example, the
number of samples required to characterize a well-mixed or homogeneous water body is smaller
than for a highly dynamic system. Conversely, in a heterogeneous system, a larger number of
samples enabled by cheaper (i.e., lower accuracy) measurements may produce a more
representative estimate than a single high accuracy measurement. In other cases, such as long
term moored time series, high precision makes subtle trends more detectable if key assumptions
about background variability are met. A key point is that these trade-offs should always be
considered and, in some cases, can be quantitatively optimized when goals of monitoring are
clear (e.g., using power analyses, etc.).
Many measurements are involved in the determination of individual carbonate
parameters. Determination of total alkalinity by titration, for example, involves measurements of
temperature, salinity, sample mass, titrant concentration, titrant volume, and voltage.
Uncertainty in each of these measurements propagates to uncertainty in the final alkalinity
determination. When known, these uncertainties can be combined into a single estimate of
measurement uncertainty using standard methods (Ellison and Williams, 2012). This allows the
relative contribution of each source to be assessed so that areas needing improvement can be
identified. However, for reporting purposes, the repeatability (precision) and accuracy
(determined using CRMs) are usually sufficient. In all cases, ID numbers of reference materials
should be included. Often, publications will report"+/-" bounds for a measured quantity,
without defining how these bounds were determined and whether they are standard error,
standard deviation, or confidence limit. Without such information, the uncertainty estimate is
virtually useless. Additional information and reporting templates for ocean acidification data are
9 . S a in p 1 i n. g Co 11 c c p t s a n. d U 11 ccr t a i n. t y | 29

described in Jiang et al. (2015) and Ellison and Williams (2012) and provide methods for both
spreadsheet calculation and Monte Carlo simulations of combined standard uncertainty for
chemical measurements.
9.2 Weather vs. climate data quality benchmarks
There are many criteria and study goals that might be considered in designing a
monitoring effort, two of which are captured by the Global Ocean Acidification Observing
Network ( as "climate" and "weather" data quality objectives. The following
definitions are from Newton et al. (2014).
•	Defined as data of quality sufficient to assess long term trends with a defined
level of confidence.
•	With respect to OA, this is to support detection of the long-term
anthropogenically-driven changes in carbon chemistry over multi-decadal
•	Defined as data of sufficient and defined quality used to identify relative spatial
patterns and short-term variation.
•	With respect to OA, this is to support mechanistic interpretation of the ecosystem
response to and impact on local, and immediate OA dynamics.
30 | Mcas u i" i n. g € h a ngcs i n C oast a 1 € a r b o n a t c € h c m i s t r y

Table A-l. Comparison of methods and uncertainty for carbonate system analyses.
Preferred Method of
Reference Materials3
Relative Uncertainties (%) in
Calculating [C032~| when Pairet
with: b
Spectrophotometric /colorimetric
(m-Cresol Purple dye)
Tris buffer in synthetic seawater
Solid-state pH sensor (Durafet)
Factory calibration, or calibrated to
a single independent measurement
Commercially available systems
(in situ or flowthrough)
Factory calibration, calibrated gas
http ://www. esrl .noaa. sov/ smd
~2 |iatm
Custom headspace equilibration
system / infrared detection
Calibrated gas cylinders
httD ://www. esrl .noaa. sov/smd
Acidification / gas stripping /
infrared detection
Sterilized natural seawater
2-3 |imol kg-1
Acidification / gas stripping /
coulometric detection
Sterilized natural seawater
2-3 |imol kg-1
Potentiometric determination
Sterilized natural seawater
2-3 |imol kg-1
a Information on reference materials from Riebesell et al. (2010) except for solid state pH, which is from Bresnahan et al. (2014).
b Measurement uncertainty and relative uncertainties adapted from table found in Riebesell et al. (2010), except solid state pH, for which uncertainty is reported as accuracy
in Bresnahan et al. (2014). Assumed analytical procedures used are those found in Dickson et al. (2007), performed by an experienced laboratory with quality assurance
programs and reference materials.
Appendix A | 31

Bottle Sampling and Preservation Checklist
D Narrow mouth borosilicate glass bottle
D Gas impermeable cap (ground glass stopper or Teflon © lined screw cap)
D No preservation necessary if analysis occurs within 4 hours of collection
D Mercuric chloride for long term preservation (10 |iL saturated mercuric chloride
solution per 50 mL of sample volume)
Preparation for Sampling
D Rinse three times all sample bottles and sampling devices with a 5-10%
hydrochloric acid solution
D Allow sample bottles and sampling devices to air dry and cap/cover to keep clean
D Collect personal protective equipment, preservatives and coolers/ice
Ready to Sample?
D Rinse three times all sample bottles and sampling devices with sample water
D Use a sampling device that minimizes exchange with atmosphere
D Fill sample container from the bottom in a controlled manner using flexible tubing
D Allow to overflow by at least twice the time needed to fill the container to the top
(e.g., if it requires 5 seconds to fill the bottle, the overflow process should
continue for 10 seconds)
D Allow 1% of bottle volume for headspace
D Add and gently disperse pickling agent if desired
D Tightly cap sample bottle
D Place in a cool, dark location (cooler) for transport back to lab
(continued on next page)
32 I Mcas u r i n. e C h a n e c s i n € oast a 1 € a r b o n a t c € h c m i s t r v

(continued from previous page)
Sample Storage
D Samples must not be frozen
D Samples should be stored tightly capped and in a dark, refrigerated space
Carbonate Chemistry Sampling Tips
D Order of collection is important. Carbonate system parameters are unequally
affected by atmospheric gas exchange
D Order of collection: pCOi> pH > dissolved inorganic carbon > total alkalinity
D Samples should not be taken forpCO2, pH and dissolved inorganic carbon when
the water volume in the sampling device falls to -25% of the total device volume
D Always collect environmental metadata (location, time of sample, depth of
sample, salinity, water temperature, atmospheric pressure)
Appendix A | 33

B.l Essential elements
Many of the available SOPs, as well as the narrative descriptions in these guidelines lack
many of the details needed for practical application. However, documentation of these details is
a critical piece of quality assurance, reporting, and sharing of expertise within and between
laboratories. To ensure consistency, such documentation takes the form of "Laboratory
Operating Procedures" (LOPs) which, ideally, are adapted from SOPs (e.g., Dickson et al., 2007)
to suit the equipment and process details of a specific laboratory. LOPs version identifiers can
then be referenced in data sheets or lab notebooks to document the specific procedures in effect
on the date of the analysis. This section provides examples of LOPs to demonstrate typical
B.2 Example LOP: Adapting a total alkalinity SOP to a specific laboratory setting
DRAFT Determination of Total Alkalinity in Seawater by Titration Laboratory Operating
Point of Contact:
US Environmental Protection Agency
27 Tarzwell Drive
Narragansett, RI 02882
Researcher:	QA Officer:	
SHEMP Manager:	
DISCLAIMER: This procedure was written to meet the needs of the research program at the
U.S. EPA Atlantic Ecology Division. It is not a U.S. EPA standard method and must not be
referred to as such. Mention of trade names or commercial products does not constitute
endorsement or recommendation for use.
This procedure is intended for use with a Metrohm Titrino 877 Plus automated titration
unit for the determination of total alkalinity in seawater and estuarine waters. This
titration occurs quickly and can be considered a closed cell titration. This is a lab-
specific procedure based on the general approach in Dickson et al. (2007).
Metrohm Titrino 877 Plus and accessories
Stir plate
Combined pH electrode
Temperature-controlled circulating water bath
Jacketed sample beaker or holder plumbed to circulating water bath
1000 mL volumetric flask
Concentrated hydrochloric acid (reagent grade or better)
Sodium chloride (reagent grade or better)
Sodium carbonate (reagent grade or better)
Indicating soda lime
34|Measuring Changes in Coastal Carbonate Chemistry

Tris [tris (hydroxymethyl) aminomethane]
pH calibration solutions (pH 4.01, 7.00, 10.01)
250 mL beakers
50 mL beakers
125 mL sample bottles
Type 1 ultrapure water
Stir bars
Solvent wash bottle
Metrohm storage solution (p/n 6.2323.000)
3 M potassium chloride solution
Silicone tubing
USB storage device
Analytical balance(s)
Graduated pipette
Pipette bulb
3.1. Instrumental Setup and Reagent Preparation
3.1.1.	Titrino 877 Plus Startup
a)	[...details on instrument startup... ]
b)	Flush the reagent syringe. Navigate to Menu»Manual Control»Dosing, choose prep
from the lower menu and press OK. Repeat this step two more times.
3.1.2.	Calibration of Combination pH electrode
Calibrate the pH electrode daily by following this procedure:
a)	Navigate to Method»Cal_pH and choose load from the bottom menu using the arrow
keys. Press OK.
b)	Press start. The unit will proceed through a 5 point calibration (4.01, 7.00, 10.01, 7.00,
and 4.01). Follow the on-screen instructions. Remember to always rinse the pH
electrode (collect rinse water in a beaker) between calibration solutions and wick away
droplets with a Kimwipe. NEVER dry the electrode's glass membrane.
c)	Place the pH electrode back in storage solution after completion of the calibration and
between uses. NEVER leave the electrode exposed to the air for any period of time.
3.1.3.	Preparation of Titrant Solution
a)	Weigh an empty 1000 mL volumetric flask on a calibrated analytical balance. Record
weight, air temperature, humidity and ambient air pressure.
b)	Add about 700 mL of Type 1 Water to volumetric flask.
c)	Draw 8.4 mL of concentrated hydrochloric acid (-37% w/w) from a small aliquot of
the main container with a graduated pipette, dispense to volumetric flask and mix
thoroughly. Dispose excess acid immediately and safely.
d)	Add 35 g of sodium chloride, or appropriate amount to match ionic strength of media
to be titrated, and mix until dissolved completely.
e)	Fill volumetric flask to 1000 mL and mix thoroughly.
f)	Weigh filled flask and record weight. Be sure outside of flask is completely dry.
g)	Determine solution density in kg L"1.
Appendix B | 35

h)	Transfer to titration vessel and label accordingly.
i)	Prep syringe as described in startup procedure.
3.1.4. Calibration of Titrant Solution with Tris and Determination of Titer
a)	In an oven dry about 5 g Tris for 2 h at 105°C and cool and store in a desiccator.
b)	Dissolve precisely 0.940 g of Tris in 250 mL of Type 1 water (with NaCl added as
above to achieve the same ionic strength as the titrant). Record exactly the weight
c)	Multiply the exact weight measured by 0.16. This is the sample weight.
d)	Using a graduated pipette, measure 40 mL of Tris solution and dispense into a 50 mL
beaker. Add a small stir bar and insert combination pH electrode and dosing tip into
e)	Load "Tris" method from the Titrino menu, enter sample weight from step C, and
press start. Titration of buffer will proceed.
f)	When endpoint is reached, remove electrode and dosing tip.
g)	Repeat procedure 5 times total.
h)	Statistics will be generated by the Titrino unit and Titer value displayed. The Titer
value is the result of the following calculation which provides a correction factor by
which the concentration of the titrant solution must be adjusted to determine the actual
Titer = COO / (C01 * C02 * EP1)
COO = sample weight in mg
C01 = 121.4 g mol"1 Tris
C02 = 0.1 mol L"1 titrant concentration
EP1 = titrant consumption in mL
Titrant (mol L"1) *Titer = Actual Titrant (mol L"1)
3.2. Sample Analysis
a)	In a tared clean beaker, add the sample and record the sample weight in grams.
b)	Place the sample in the jacketed beaker or holder.
c)	Load the SW TotalALK method (this is the instrument file containing titration
parameters such as initial dose, mL per subsequent dose, etc.). Enter the sample name.
d)	Press start. The unit will perform the titration and determine the equivalence points by
the Fortuin method and export the results to the USB drive with the sample name as
the filename.
3.3. Total Alkalinity Calculation
a) At the endpoint of the titration (pH = 3.0) the solution Alkalinity can be described by.
mC — mni4r
	= [H+]f + [HSO4] + [HF] + [H3P04] + [H2P04~]
m0 + m
m is the mass of acid consumed during titration
mo is the starting mass of sample
C is the concentration of acid in mol kg"1
^4-ris the total alkalinity in mol kg"1 Where:
36 | Mcas u r i 11 g C h a ngcs i 11 C oast a 1 € a r b o n a t c € h c m i s t r y

This equation can be solved with a customized computer program or an existing software
function that accepts titration volumes, millivolt readings, titrant molarity, temperature,
and salinity as input arguments ((e.g., the AT function in the R package 'seacarb',
Lavigne et al., 2011).
4.	QA/QC
Replicate samples should be run to ensure data quality. At least one set of duplicates
should be run every ten samples, and the absolute difference recorded. After an initial
dataset of at least 12 duplicates is assembled a control chart should be produced. See
SOP 22 in Dickson et al. (2007) for specific recommendations on preparation and
interpretation of control charts.
Instruments should only be evaluated by experienced personnel. All other instrumental
troubleshooting should be performed under the guidance of the manufacturer or qualified
service personnel.
Dickson, A. G., et al., Eds. (2007). Guide to best practices for ocean CO2 measurements
PICES Special Publication 3.
Fortuin, J., 1961. Method for determination of the equivalence point in potentiometric
titrations. Anal Chim Acta, 24: 175-191.
7.1. Alkalinity titration parameters
[This section should document specific instrument settings such as start conditions, end
condition, initial titrant dose, titrant dosage increments, stir bar speed, etc.]
B.3 Example LOP: Measuring DIC using an NDIR-based instrument
DRAFT Determination of Dissolved Inorganic Carbon Laboratory Operating Procedure
Point of Contact:
US Environmental Protection Agency
27 Tarzwell Drive
Narragansett, RI 02882
Researcher:	QA Officer:	
SHEMP Manager:	
DISCLAIMER: This procedure was written to meet the needs of the research program at the
U.S. EPA Atlantic Ecology Division. It is not a U.S. EPA standard method and must not be
referred to as such. Mention of trade names or commercial products does not constitute
endorsement or recommendation for use.
Appendix B | 37

This procedure is intended for use with Apollo SciTech AS-C3 dissolved inorganic
carbon analyzer for the determination of total dissolved inorganic carbon (DIC) in
seawater and estuarine waters. This analysis takes approximately 15 minutes per sample
(if using 5 injections per sample). For greatest precision, temperature must be carefully
controlled at 25°C for reagents and samples. This is a lab-specific procedure based on
the general approach in Schumacher and Smucker (1983) and on specific instrument
instructions provided in the AS-C3 User Manual.
Apollo SciTech Model AS-C3 and accessories
Temperature control recirculating water system
40-60 mm diameter silicon or tygon tubing
4-6 mm diameter silicon or tygon tubing
Jacketed beaker for acid (9.5 cm diameter)
Jacketed beaker for sample (9.5 cm diameter)
N2 regulator suitable for precise operation from 5-15 mL/min
Titanium diffuser
Silicone stopper with two 5 mm ports
5 mL syringe with luer lock fittings
Replacement rubber o-rings
Replacement 0.20 |im, hydrophobic PTFE filters with luer lock fittings
250 mL beakers
50 mL beakers
24-40 mL sample bottles
100 mL graduated cylinder
1000 mL volumetric flask
1000 mL glass stoppered bottle
N2 gas supply
Andrew Dickson Scripps certified reference material
Concentrated phosphoric acid (reagent grade or better)
Sodium chloride (reagent grade or better)
Indicating soda lime
Magnesium perchlorate drying agent
DI wash bottle
USB storage device
Analytical balance(s)
38 | Mcas u r i 11 g C h a ngcs i 11 C oast a 1 € a r b o n a t c € h c m i s t r y

3.1. Instrumental Setup and Reagent Preparation
3.1.1.	Preparation of 5% Phosphoric Acid Reagent
a)	Fill a 1000 mL volumetric flask with 500 mL of milliQ water.
b)	Measure out 59 mL of 85% phosphoric acid into a graduated cylinder.
c)	Slowly add phosphoric acid to the volumetric flask with milliQ water and gently stir to
d)	Add 100 grams of sodium chloride and stir to dissolve.
e)	Finally fill to the 1000 mL line of the volumetric flask with milliQ water.
f)	Transfer to a stoppered or screwcap bottle.
3.1.2.	Apollo AS-C3 setup
a)	[...details on instrument startup... ]
b)	Allow the instrument to warm up for at least 60 minutes.
c)	Set the temperature-controlled water recirculator and a temperature-controlled water
bath to a set temperature, we recommend 25°C for both.
d)	Place samples to be run in the water bath and place 2" of DI water into the jacketed
water beakers attached to the water recirculator system.
e)	Bring phosphoric acid reagent to fume hood and gently bubble for at least 15 minutes
with pure N2 gas using a titanium diffuser.
f)	After 15 minutes, cap the bottle and transfer to the jacketed beaker to allow the reagent
to come up to temperature.
g)	Place samples into water bath and allow to come up to temperature.
h)	After instrument has warmed up for at least 60 minutes, load DIC analysis program
from computer.
i)	Insert the silicone cap into acid reagent bottle and attach tubing from port C of the DIC
j) Place sample into the second jacketed beaker and place tubing from port B into the
sample bottle, leave tubing 2 cm off of bottom.
3.1.3.	Calibration of instrument and drift correction
a)	Dickson CRM is used in calibration. This can be portioned off from the 500 mL
sample bottle to smaller sample vials in order to minimize the amount of change in DIC
concentration due to headspace equilibration, as follows:
Using acid washed/DI rinsed, small diameter tygon tubing, create a syphon and slowly
gravity fill from a freshly opened CRM bottle into 24 mL borosilicate glass scintillation
vials with Teflon© displacement caps (this method will produce -20 headspace free,
daily running standards; not suitable for long term storage).
b)	DIC concentration of Dickson CRMs is -2000 |imol L"1. Use the specific batch DIC
concentration to bracket the expected DIC concentration of your samples. Run the DIC
standard as three separate samples. For example, changing the volume for each sample
to solution volumes of 0.8, 1.0, and 1.2 mL will produce a standard curve bracketing
samples from 1600 to 2400 |imol L"1.
c)	Drift will occur during a full day of running the instrument. To account for this run the
DIC standard curve at the beginning of the run and end of the run. To further adjust the
drift correction, run one volume (1.0 mL) of the standard as an unknown every 5-8
Appendix B | 39

d)	After standardization volumes have been completed, use a linear regression software
program to determine the relationship between peak area and DIC. For example, if the
standard DIC concentration is 2000 jamol L"1 a 1 mL injection is equal to 2 |imol, a
.8 mL injection is equal to 1.6 jamol. This can then be converted to DIC concentration
per liter by multiplying by 1000.
e)	Calculate the DIC concentration of samples by applying the linear model obtained
during standardization to peak area of samples.
f)	Convert sample concentrations from jamol L 1 to |imol kg-1 by using a density equation
that incorporates temperature, salinity and pressure of the sample.
3.2. Sample Analysis
a)	It is possible to perform single sample runs but it is recommended that the analyst uses
the batch process function.
b)	Open up batch process window and select a measurement "scheme" the software
contains 3 schemes to choose from or the analyst can create their own.
c)	We recommend that the analyst uses a scheme with 0.1% error allowed within 3 repeats
out of a maximum of 5 measurements.
d)	Create sample inventory under the sample list tab by adding samples by clicking the
"+" button to add lines.
e)	Name samples and set injection volume (0.8 - 1.2 mL).
f)	Press the "Sample Measurement" button to start the batch measurement.
g)	Instruments syringe will automatically begin drawing correct volumes of sample and
reagent before sending to the reactor chamber.
h)	After -180 seconds sample will be expelled and instrument will automatically begin
next injection (i.e. 2 of 5).
i)	Upon sample completion, there will be an audible beeping notification and you will be
greeted with a dialogue box asking to run the next sample.
j) If the sample fails the replicate precision parameters|i|i specified in the scheme, you will
be presented with a dialogue box asking if you want to re-run the sample.
k) Results will be saved within the program and can be viewed by clicking the "Test
Result" button.
1) Results will need to be exported prior to closing the program, this can be done through
the test result screen.
m) At the end of the run, add a new sample called "DI rinse", set volume to 1.5 mL.
n) Place sample tube and acid reagent tube into DI water and run the sample.
o) Take the tubes out of the DI water and keep in air.
p) Click the "Connect/Disconnect" button, When the prompt appears, press enter.
q) Close down the DIC analysis program.
r) Turn off the power switch for the LI-7000 bottom unit, Apollo AS-C3 upper unit and
gas supply.
3.3. Dissolved Inorganic Carbon Calculation Example
Dickson CRM batch #151 standard concentration: 2033.83 ± 0.62 |imol kg"1
Density at salinity (33.345), Temperature (25°C): 1022.091 (kg m3)
Dickson CRM batch #151 standard concentration: 2033.83 *(1022.091*10"3) =
2078.76 |imol L"1
40 | M c a s u i" i n. g € h a ngcs i n C oast a 1 € a r b o n a t c € h c m i s t r y

Standard curve:
2078.76 |imol L"1 * 0.0008 L"1 = 1.663 |imol
1.663 |imol = 11381 raw CO2 peak area
2078.76 |imol L"1 * 0.0010 L"1 = 2.079 |imol
2.079 |imol = 14246 raw CO2 peak area
2078.76 |imol L"1 * 0.0012 L"1 = 2.495 |imol
2.495 |imol = 17072 raw CO2 peak area
$ 15000
CO, Peak Area ~ DIC Content
1.8	2.0	2.2	2.4
Micromoles DIC
Figure B-l. Dissolved inorganic carbon peak area as a function of dissolved inorganic carbon
Example calculation:
Sample Volume: 1 mL 1 (0.001 L"1)
Measured Peak Area: 13300
Regression Equation: y(peak area) = 6840.1 * x(DIC) + 12.007
Re-Arranged Regression Equation: x(DIC) = (13300 - 12.007)/6840.1
Calculation: DIC = (13300 - 12.007)/6840.1 = 1.942 |imol mL"1
Volume Correction Factor: (Scale to 1 L): 1.942 |imol mL"1 / .001 1 = 1942 |imol L"1
4. QA/QC
As Dickson CRMs are used in the creation of standardization curves, it is important to
utilize an independent check standard. This check standard should be made up in
artificial seawater with a salinity ±20% expected sample salinity. The check standard
DIC concentration should also be ±20% of expected DIC concentration. Replicate
samples should be run to ensure data quality. At least one set of duplicates should be run
every ten samples, and the absolute difference recorded. After an initial dataset of at
least 12 duplicates is assembled a control chart should be produced.
Appendix B | 41

Instruments should only be evaluated by experienced personnel. All other instrumental
troubleshooting should be performed under the guidance of the manufacturer or qualified
service personnel.
Schumacher, T. E., and A. J. M. Smucker. 1983. Measurement of C02 dissolved in
aqueous solutions using a modified infrared gas analyzer system. Plant Physiology 72:
42 | M c a s u i" i n. g € h a ngcs i n C oast a 1 € a r b o n a t c € h c m i s t r y

The following scenarios illustrate the range of equipment that can support differing
C.l A water quality research laboratory
EPA's Atlantic Ecology Division (AED) in Narragansett conducts coastal acidification
research focusing primarily on land-based drivers. In order to separate these from oceanic
drivers and to determine when and where costly methods are necessary, high quality
measurements are necessary. The following is a description of equipment used in AED's studies
of coastal carbonate chemistry. This work is closely coordinated with nutrient sampling and
Research vessels equipped with winches and GPS
Go Flo Bottles and van Dome samplers with silicon transfer tubes
CTD rosette, from which in situ temperature and salinity is determined (Hydrolab)
Handheld YSI for environmental measurements in shallow waters
400 hundred sample bottles (borosilicate glass-stoppered; borosilicate septum vials, and HDPE)
Auto-pipettors for dispensing preservatives (dedicated sole-purpose to prevent cross-
Apiezon grease in dispensing syringes
Personal protective equipment
Spill prevention and response kit
Coolers for sample transport
General support facilities
Walk-in refrigerator for storage of samples and reference materials
Laboratory glassware cleaning facility
DI and ultra-pure Milli-Q water supply
Waste handling facilities and protocols
Seawater wetlab for testing field instruments
General laboratory equipment
General glassware and calibrated cylinders, flask and pipettors
Ventilated storage area (i.e., fume hood) for hazardous reagents (phosphoric acid, hydrochloric
acid, mercuric chloride)
Calibrated NIST-traceable digital thermometer for checking water bath controls, instrument
sensors, etc.
Drying oven for drying powder/crystal reagents
Personal protective equipment (gloves, eye protection, lab coats)
Eyewash station
Appendix C | 43

Total alkalinity analysis
Metrohm Titrino 877 autotitrator with magnetic stir plate, ANOVA cooling/heating water bath,
thermometer, combination electrode, jacketed beakers connected to water bath, glassware for
acid titrant and Tris preparation, high quality calibrated analytical balance
DIC analysis
Apollo-Scitech AS-C3 Dissolved Inorganic Carbon Analyzer; jacketed beaker holders connected
to water bath (temperature-controlled), continuous supply of research grade N2 gas, computer for
instrument control, moisture and CO2 scrubber columns
pH determinations
Perkin-Elmer Lambda 35 dual beam UV-VIS Spectrophotometer with jacketed quartz cell (10
cm path), solid calibration standards, dedicated temperature-controlled water bath, computer,
micro-pipettor for dispensing w-cresol dye
Other pH equipment
Satlantic SeaFET autonomous pH sensor
WTW 3310 pH meters with Sentix probes
Reagents and reference materials
Tris (for acid titrant calibration), hydrochloric acid, phosphoric acid, sodium carbonate,
bicarbonate, sodium chloride, sodium hydroxide, mercuric chloride, NBS buffers, m-cresol dye
Certified reference materials for DIC, TA
Reference materials for pHi(Tris buffers)
C.2 A single-instrument setup in a basic water quality laboratory
As one example, we were recently encouraged by a presentation at the ASLO 2017
meeting by Erin Guyler and Robert Byrne that described a single instrument scenario for full
constraint of the seawater carbonate system. Their method used both beams in a single
spectrophotometer to obtain pHr and alkalinity. At low carbonate saturation states, they
obtained "climate quality" measurements. Although the TA + pHr pair has higher uncertainty
due the sensitivity of carbonate system determination to error in pH (see Table A-l), the results
are likely to be far superior to the use of parameter pairs involving either the salinity-TA
regression or low quality pH measurements. When grounded to Tris reference materials for pH
and certified reference materials for alkalinity, these measurements could conceivably be used to
ground-truth autonomous sensors. Careful consideration of sample holding times and
preservation would be required.
C.3 A monitoring effort with external laboratory support
A growing number of monitoring groups are partnering with water quality research
laboratories to obtain ground-truthing of their field measurements. This allows the monitoring
group to focus on the sampling design issues (i.e., trade-offs between sample
frequency/resolution and sample quality) and operation of autonomous sensors which, at present,
are gaps in the application of existing open ocean methods to the coastal zone. See Jones et al.
(2016) and Bresnahan et al. (2014).
44 | M c a s 111" i 11 g € h a nges i 11 C oast a 1 € a r b o n a t c € h c m i s t r y

C.4 Shellfish growers and hatcheries
A high quality handheld pHnbs meter and probe can be purchased for under $1000.
These are useful for monitoring diurnal and spatial patterns with a given growing area, but
require calibration and occasional probe replacement. With extensive additional effort and
purchase of non-commercial Tris buffers, or purchase of DuraFET sensors ($2500-$3000, with
proper power supply), handheld devices can be used for measuring pH-r. In any case, our
opinion is that growers concerned about low pH zones or seasonal excursions should not be
discouraged from taking their own measurements. Additional guidelines regarding the
importance of consistent methodology and sample timing would likely be beneficial.
Appendix C | 45

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