The Seawater CO2-Carbonate System
http://en.wikipedia.org/wiki/File:CarmelOoids.jpg
Carmel formation, Utah, Jurassic
Close-up of Oolitic
limestone
Some of the major geochemical roles of
the CO2 system in seawater include:
• seawater pH control and buffering
• source of carbon for photosynthesis
• long-term sink for carbon via carbonate precipitation and
subsequent burial and preservation of limestone and
dolomite
• formation of carbonate reefs
• exchange of CO2 with the atmosphere: CO2 is a major
greenhouse gas. Oceans are both a source and sink for atmospheric CO2
depending on location; a net sink overall.
• source of biogenic carbonates that are important
paleoindicators for a variety of parameters
The Seawater CO2-Carbonate System
Use of sign inspired by talk given by Andrew Dickson, a noted
CO2-system chemist from Scripps.
The Seawater CO2-Carbonate System
The carbonate system is one of the most important chemical
and biogeochemical systems on earth.
Air
Sea
Proton #1
Proton #2
CO2 (g)

CO2 (aq) + H2O <=> H2CO3 <=> H+ + HCO3- <=> H+ + CO32Ko
K1
K2
For seawater with a salinity of 35 and a temperature of 25 oC:
pKo = 1.547
(From Millero, Table 7.4)
pK1 = 5.847
pK2 = 8.915
Structures of CO2 gas
and bicarbonate ion
O=C=O
CO2
O=C-OH
HCO3-
O-
H2CO3
The total
CO2 stays
constant
but the
speciation
depends
on pH
H2O + CO2
Equivalence points
pH = pK1
pH = pK2
From, Millero, Chemical Oceanography, 1996. p 246
Terminology related to the CO2 system in
seawater
DIC - Dissolved inorganic carbon (CO2(g)+H2CO3+HCO3-+CO32-)
CO2 (g) : gaseous CO2
CO2 (aq) : gaseous CO2 that is dissolved in water
Carbonate ion: CO32Bicarbonate ion: HCO3Carbonic acid: H2CO3 (non-charged,neutral species)
Total CO2 (CO2): Sum of all dissolved components
of inorganic carbon, including CO2 (aq), H2CO3, HCO3-,
and CO32-. Total CO2 = DIC
PIC – Particulate inorganic carbon (calcite & aragonite minerals)
Distribution of DIC in the ocean- Vertical and
horizontal distributions
Total DIC (=Total CO2)
is lowest (but not
zero!) in surface
waters and is enriched
in deeper water - the
enrichment is greatest
in the deep Pacific due
to water mass age.
Why is this pattern
observed?
pH in Seawater – Complex control by CO2
system and Alkalinity
Dissolution of CO2 in water results in formation of
carbonic acid, which dissociates to yield bicarbonate
and carbonate plus protons.
CO2 (aq) + H2O <=> H2CO3 <=> H+ + HCO3- <=> H+ + CO32-
Thus, the addition of CO2 to water increases the {H+}
and therefore lowers the pH of the solution.
Conversely, removal of CO2 from solution removes
{H+} and increases the pH
Biological uptake of carbon by marine plants is mainly
as CO2(g) or H2CO3 i.e. neutral species. (some
phytoplankton can take up HCO3- and convert it to CO2 via carbonic anhydrase)
The uptake of CO2 (g) by photosynthetic organisms (or
chemosynthetic organisms) will raise the pH of the
system due to shift in the equilibria to the left (in the
direction consuming H+. Conversely, respiration of
organic matter reverses the cycle and liberates CO2
which will dissociate and lower the pH (increase {H+}).
CO2 (aq) + H2O <=> H2CO3 <=> H+ + HCO3- <=> H+ + CO32Remove CO2, consume H+ and raise pH
Add CO2, add H+ and lower pH
Vertical distribution of pH in the ocean
pH
7.5
7.7
7.9
Depth (km)
0
2
Atlantic
4
Pacific
6
Indian
8.1
There is an oceanwide pH minimum
starting just below the
euphotic zone and
extending to 500-1000
m. pH is lower in
deep Pacific than in
Atlantic - due to water
mass age and
accumulation of
respired CO2!
Seawater alkalinity - a measure of the
buffering capacity
Simply put:
“The amount of negative charge in seawater that is able to
accept a proton (hydrogen ions) during the titration of
seawater with strong acid to the point where essentially all
the carbonate species are protonated” (paraphrased from
Pilson, p 114).
Alkalinity is not the total negative charge in solution,
but rather just the concentration of negatively charged
species that will accept H+ above certain pH end-point
- defined by the method of titration - usually around
pH 3.5 - 4.5
Chloride is a negatively charged species, but Cl- will
not accept a proton in aqueous solution, even at pH 0!
For most natural waters, total alkalinity (TA) can be
simplified to:
(TA) = [HCO3-] + 2 [CO32-] + [B(OH)4-] + [OH-] - [H+]
~95% of the alkalinity in
seawater is comprised of
the carbonate alkalinity;
Carbonate Alkalinity (CA) =
[HCO3-] + 2 [CO32-]
Carbonate Alkalinity (in molal units)
is always greater than Total CO2 (in
molal units) because each unit of
CO32- contributes 2 units of
alkalinity (can accept 2 protons)
Borate contributes
about 5% to the
alkalinity and
needs to be taken
into account.
Alkalinity according to Dickson (1992) (a detailed view)
Alkalinity = [HCO3-] + 2[CO32-] + [B(OH)4-] + [OH-] +
[HPO42-] + 2[PO43-] + [SiO(OH)3-] + [HS-] + [NH3] + [all
other unidentified weak bases] - [H+] - [HSO4-] - [HF]
- [H3PO4] - [all other unidentified acids]
Alkalinity is not strictly related to pH. For example:
Deep ocean water has higher alkalinity than the surface - but a lower
pH (higher acidity).
Factors affecting Alkalinity and Total CO2 in seawater
 Alkalinity is not affected by T and P (because it is a charge
balance).
 Alkalinity increases with dissolution of carbonate minerals
(which release HCO3- or CO32-). Dissolution of CaCO3 releases
CO32-, thereby increasing alkalinity. Likewise, precipitation of
carbonate minerals consumes (decreases) alkalinity.
Carbonate precipitation also affects CO2.
 Photosynthesis and respiration consume and add CO2
respectively, but do not affect alkalinity. This is because release of
CO2 and subsequent hydration and dissociation yields HCO3- + H+
(one unit of H+ for every unit of negative charge alkalinity).
Remineralization does however, increase CO2.
 The exception to this rule is respiration with sulfate as the electron acceptor.
Sulfate reduction generates HS- which increases alkalinity. The CO2 generated
in sulfate reduction only increases CO2 and not alkalinity.
Vertical and horizontal distribution of Alkalinity in
the ocean
• Similar to that of total DIC - low in surface
waters, increasing with depth in thermocline.
DIC
Alkalinity
Higher in deep Pacific than in Atlantic - due to water mass age and
inputs of CO32- from CaCO3 dissolution
Fig. 15.10 in Libes
Precipitation and dissolution of carbonate
minerals in the ocean
All seawater contains the ions
Ca2+, CO32- and HCO3-. The
effective concentrations (i.e.
activities) of these species,
together with the pH,
temperature, pressure and ionic
strength determine whether the
solution is saturated or
undersaturated with respect to
CaCO3 minerals.
Cocolithophore
Emiliania huxleyi a
haptophyte
phytoplankter secretes
plates (liths) of calcite
(CaCO3)
Carbonate Minerals
Calcite
CaCO3 <=> Ca2+ + CO32- Ksp = 4.47 x 10-9 @ 25oC and Ionic
strength of 0 (a std condition)
Aragonite
CaCO3 <=> Ca2+ + CO32- Ksp = 6.02 x10-9 @ 25oC and Ionic
strength of 0 (a std condition)
Aragonite has the larger Ksp, therefore it is more soluble.
Aragonite is more amorphous (less
ordered crystal) and is more soluble
than calcite. These two compounds
differ only in their crystalline structure
not their chemical formula which is
CaCO3 in both cases.
Calcite is the predominant
form of CaCO3 in the
ocean and it is more stable
than aragonite
(amorphous CaCO3).
Coccolithophore
Foraminiferan
Pteropod
Dinoflagellate
cyst
The organisms that
precipitate calcite include
(Cocolithophores and
foraminifera).
Organisms that precipitate
aragonite include Corals &
Pteropods.
Precipitation of biogenic carbonates and their
subsequent burial in marine sediments represents
the single largest export of carbon from the
biosphere.
Biogenic carbonates in sedimentary rocks (e.g.
limestones and dolomites) are the single largest
reservoir of carbon on Earth. Most of this carbonate
is derived from planktonic microorganisms
Not all biogenic carbonate is preserved in sediments much dissolves in the deep sea. The precipitation and
dissolution of CaCO3 depends on the physicochemical
conditions in seawater
There is 1400
times more
Carbon tied
up in
carbonate
rocks than
there is in
DIC in the
ocean!
Nearly all surface ocean waters are supersaturated with respect
to calcite and aragonite; deep waters are undersaturated.
Despite surface supersaturation, spontaneous precipitation
of calcite or aragonite in surface waters does not occur
(except at very high pH's) due partly to interaction of Mg2+ with
CaCO3 crystal surfaces.
Only in very warm, saline waters
where CO2 solubility is low (hence
CO2 is low) will CaCO3 ppt out as
aragonite without biocatalysis.
Carbonate Ooids are examples of
spontaneously precipitated
carbonates – currently found on
Bahamas platform – but extensive
geological deposits exist
http://www.iun.edu/~geos/
Calcifying organisms overcome the Mg2+ problem with enzymes and
intracellular compartmentalization of pH etc.
Calcification:
Calcium carbonate precipitation can be written simply
as:
Ca2+ + CO32-   CaCO3 (s)
But biogenic CaCO3 precipitation appears to occur primarily
by the following reaction mechanism:
Ca2+ + 2HCO3-   CaCO3(s) + CO2 + H2O
Thus, per mole of CaCO3 formed, calcification i) consumes 2 mole of
alkalinity, ii) consumes 1 mole of DIC and iii) produces 1 mole of
CO2 (i.e. increases pCO2)
The Keq values for the CO2 system reactions
are a function of temperature & pressure
therefore so is CaCO3 solubility
As Temp goes down, pH goes down; Ksp of CaCO3 goes up
(more soluble) (retrograde solubility)
As Pressure goes up, pH goes down; Ksp of CaCO3 goes up
(more soluble)
These effects are due the fact that CO2 gas and charged vs.
neutral species are involved in the equilibrium. Gases are more
soluble at higher pressures and lower temperatures, favoring
CO2 (g) dissolution, hence more carbonic acid forms. Also, as
pressure increases, formation of charged species is favored
because the ions have a lower partial molal volume than the
solid (or neutral species) due to electrostriction.
The degree of saturation (Omega) can be expressed as:
 
[Ca2 ] seawater x [CO3 2 ] seawater
[Ca2 ] saturated x [CO3 2 ] saturated

Ion product
K SP 
Omega is given as an output in the CO2SYS program.
If IP > Ksp*, then solution is supersaturated ( > 1) . If IP <
Ksp* then solution is undersaturated ( < 1) .
Solubility of CaCO3 depends mostly on variations in CO32rather than Ca2+ because Ca2+ is nearly constant in the ocean.
The rate of dissolution of CaCO3 is an exponential function
of the degree of undersaturation.
CaCO3 is not found in the surficial sediments in the
deepest parts of the sea (> ~5000 m) to any great
extent for at least two reasons.
1) The solubility of CaCO3 increases as Pressure  and as Temp. 
2) pH decreases with depth and more CaCO3 will dissolve.
The lysocline is the depth at which significant dissolution
of calcite begins. This depth is different for different ocean
water masses.
The CCD (Calcite compensation depth) is the depth at
which the dissolution of CaCO3 minerals equals the
supply rate (rain rate). No significant accumulation of
CaCO3 occurs below this depth.
Δ
Emerson and
Hedges Fig 12.12
CaCO3
CaCO3
Places where CaCO3
dominates the sediments
are relatively shallow (<
5000 m)
CaCO3
CaCO3
CaCO3
CaCO3
CaCO3
CaCO3
Source: Open University: Ocean chemistry and deep sea sediments
The CCD for aragonite is much shallower than for
calcite because aragonite is more soluble (larger Ksp).
The CCD for calcite is shallower in the Pacific (3.5
km) than in the Atlantic (5 km) due to the lower pH of
the Pacific deep waters (caused by age and CO2
production from respiration).
Many factors govern the CCD including the rate of supply,
chemical composition, minerology, size and shape, rate of
bioturbation. Larger particles may not dissolve quickly. The
distribution of CaCO3 oozes as they are called (sediments with
> 75% CaCO3) is largely restricted to shallower parts of the
oceans (see next slide figure from Open University text).
Calculation of all the parameters of the CO2
system in seawater using the CO2SYS program
(available for free download at:
http://cdiac.esd.ornl.gov/oceans/co2rprt.html).
CO2SYS will do all the calculations for you provided you have input
data for two of the four main parameters of the CO2 system:
Total Alkalinity
pCO2
pH
Total CO2
With input of two parameters (plus temperature, salinity, pressure,
silicate and phosphate data) the other two parameters of the CO2
system will be predicted as well as the concentration of various
species, the degree of calcite or aragonite saturation, and more. The
program is very easy to use.
pH Scales
(defined in CO2SYS)
pHNBS (National Bureau of Standards; standard lab pH buffers are NBS,
but they are low ionic strength and not great for SW.
pHseawater
pHtotal
pHfree
Dickson
recommends
these
Differences in these scales
have to do with how they
consider the sulfuric acid and
hydrofluoric acid
components of seawater
pH values on the total scale (pHtot) are about:
.09 units lower than those on the free scale,
.01 units higher than those on the seawater scale, and
.13 units lower than those on the NBS scale.
Exchange of CO2 (g) between the
atmosphere and the ocean
• Portions of the ocean surface are super saturated
with CO2(g) while other portions are undersaturated.
• Only the CO2 (g) part of the total CO2 system can
exchange with the atmosphere. Thus, knowledge of
the partial pressure of carbon dioxide (pCO2) is
critical for understanding exchanges of carbon
between the atmosphere and oceans.
• There is about 50 times more Total CO2 dissolved
in the oceans than there is CO2 in the atmosphere.
CO2 (aq) is in equilibrium with the atmosphere
such that:
[CO2 aq] = HCO2 * pCO2
Where HCO2 is the Henry’s Law constant for CO2
and pCO2 is the partial pressure of CO2. The
Henry’s Law constant is essentially the equilibrium
constant for the dissolution of the gas:
CO2 (g) <=> CO2 (aq)
Keq = {CO2 (aq)}/ {CO2 (g)}
The activity of CO2 in the gas is essentially its partial pressure
The current concentration of CO2 in the atmosphere (in
2012) is about 392 ppm, or 0.0392%. This concentration
has already increased 40% from pre-industrial values
and is expected to nearly double in the next century.
Atmospheric CO2 is increasing dramatically
The
implications
for marine
systems are
huge!
Get the latest CO2
concentration at
http://co2now.org/
392 in
Sep
2012
Ron enters
graduate school @
340 ppm
Ron born
at 318
ppm
The increase is 78
ppm in 54 years – a
25% increase
Source: C. D. Keeling http://cdiac.esd.ornl.gov/trends/co2/sio-mlo.htm
Source: Buddemeir et al. Pew Report on
Coral Reefs and Global Climate Change
Consequences of global increase of
CO2 in atmosphere
 Greenhouse warming
 Sea level rise - polar ice decline
 Enhanced terrestrial primary productivity
 Decreased seawater pH (more carbonic acid)
 changes in phytoplankton physiology/ecology
 decreased calcification by corals and other
marine organisms
 decrease of pH in rain/snow - greater
terrestrial weathering
 Indirect effects - many
Hydrogen ion activity {H+} as a function of pH
Hydrogen ion activity (molar)
2.5E-08
2E-08
1.5E-08
1E-08
26%
increase
for pH
8.2  8.1
5E-09
58% increase
for pH 8.2 
8.0
0
7.6
7.7
7.8
7.9
8
pH
8.1
8.2
8.3
8.4
2x CO2
Source: Buddemeir et al. Pew Report on Coral Reefs and Global Climate Change
The oceans do not always achieve
equilibrium with respect to atmospheric
CO2
• This is due to sluggish kinetics of the equilibria of gas
exchange and the fact that the ocean is a layered system with
a relatively long residence- and mixing time.
• The marine biota add or remove CO2 in surface waters on
short time scales, thereby affecting direction of the CO2 flux.
• Only about half of the CO2 input to the atmosphere by Man’s activities since the dawn
of the industrial age has accumulated in the atmosphere. The other half has been
absorbed by either the oceans or the terrestrial biota. The ocean’s response takes
time.
Into the ocean
Out of the ocean
Fish otoliths (ear
stones) are made of
aragonite/protein layers
Checkley et al. found
that growing larval
White Sea Bass at
elevated pCO2 caused
otoliths to be 8% and
16% larger in the 1000
and 2500 µatm
treatments compared to
the 430 µatm controls
Control
Stop !
IPCC-FAR
In calculating the Ksp the activity of the ionic
species should be used. In practice, marine
chemists would measure the concentration of
Ca2+ and CO32- at which precipitation occurs.
The resulting solubility product would be the
apparent Ksp’ or stoichiometric constant. It’s
value would depend on the conditions such as
temp, pressure, and Ionic strength.
In practice alkalinity is measured by
titration.
The amount of H+ in equivalents per kg needed to
titrate 1 kg of seawater to the bicarbonate/H2CO3
equivalence point.
Modern methods involve coulombic titrations to determine
end point.
Depth profiles of carbonate mineral saturation state in the
Atlantic and Pacific Oceans. An Omega value of 1 indicates
saturation; above 1 is supersaturated; below, undersaturated.
(From Millero, Chemical Oceanography, 1996. pp 274 & 275.
Equilibria
HBa H+ + BaH2O H+ + OH-
Since there are no other cations to balance the
[Ba-], H+ will adjust to equal the concentration
of Ba- to satisfy both equilibria. This sets pH at
4 in this case
Equivalence points – where pH = pK
Consider dissociation of the weak acid:
H2CO3  H+ + HCO3Ka = {H+} {HCO3-}/{H2CO3}
Rearrange to isolate {H+}
Ka {H2CO3} / {HCO3-} = {H+}
Take “p” or negative Log of both sides
pKa {H2CO3} / {HCO3-} = pH
or
{H2CO3} / {HCO3-} = pH/ pKa
At pH = pKa, then {H2CO3} / {HCO3-} = 1
Thus, at pH =pKa, the {H2CO3} and {HCO3-} must be equal, hence
the equivalence point
pH control in sea water - seawater pH varies from 7.9-8.4 with
an average between 8.1 and 8.2.
Seawater pH is controlled largely by the reaction:
HCO3- <=> H+ + CO32The equilibrium expression for this
reaction is:
2
{H  }{CO3 }
Ka 

{HCO 3 }
where Ka is the dissociation constant of the bicarbonate. (~10-8.9) i.e. pKa ~8.9
Looking at this equilibrium expression another way, the pH will depend on the

ratio of bicarbonate to carbonate
K
{
HCO

a
3 }
and vice versa.
{H } 
2
{CO3 }
Measurement of seawater pH is not straightforward due to the complexity of the solution
and its ionic strength (I=0.7). Several different pH scales are in use (depending on
buffers used for standards). The NBS pH scale is the most common.
Ken Caldeira
Climate System Modeling Group
Lawrence Livermore National Laboratory
7000 East Avenue, L-103
Livermore, CA 94550, U.S.A.
E-mail: kenc@LLNL.gov
Robert Berner
Department of Geology and Geophysics
Yale University
New Haven, CT 06520-8109, U.S.A.
E-mail: berner@hess.geology.yale.edu
http://www.sciencemag.org/cgi/con
tent/full/286/5447/2043a