© Division of Chemical Education  • www.JCE.DivCHED.org  •  Vol. 85 No. 2 February 2008  •  Journal of Chemical Education 213

Chemistry for Everyone

Textbooks in environmental chemistry apply chemical and physical principles to numerous topics of current environmental concern such as ozone depletion, acid rain, chemical wastes, urban smog, and global warming. When the impact of rising levels of CO2 in the atmosphere is treated in these texts it is usually emphasized that CO2 is a greenhouse gas that can lead to global warming (1–3). In that context, the role of the ocean in dissolving atmospheric CO2 is considered a “good thing” since it removes a greenhouse gas from the atmosphere. What is often neglected is the likelihood that increasing levels of CO2 in the ocean will significantly lower the pH of surface water and lead to the dissolution of calcium carbonate in coral and other calcifying organisms. Thus, while acidification of terrestrial waters (and the related carbonate chemistry) is typically treated in the context of “acid rain” (1–3) the acidification of surface seawater is often not emphasized. Over the last few decades, however, this topic has risen to the fore as a serious environmental concern and has been reviewed recently in widely read scientific journals such as Nature (4) and Scientific American (5).

It should be emphasized here that acid–base carbonate chemistry in seawater is not entirely the same as in freshwater. The presence of electrolytes in the ocean affects the activities of solution species significantly, and the temperature of seawater also varies over a wide range. Thus “apparent” acid ionization constants (Ka) and the autoionization constant of water (Kw) in seawater must be written as functions of both salinity (S) and temperature (T ). In addition a complete treatment of the topic may necessitate the inclusion of multiple weak acids—such as phosphoric, silicic, and boric acids—although these can be ne-glected to first-order approximation. It is also important to note that the ocean is not a system at equilibrium. In particular, we may not assume that calcium carbonate (in the form of aragonite or calcite) is in equilibrium with calcium and carbonate ions in the surrounding water. In fact, the Earth’s oceans are currently supersaturated with respect to these ions and may become un-saturated in the future.

In a previous article published in this Journal (6), Weston treats the topic of ocean acidification and its effects on coral reefs, in the following manner. First, he outlines a detailed treatment of “carbonic acid equilibria” in pure water. Then, after establishing the basic concepts, he refers the reader to more ad-

vanced treatments in seawater without elaborating many details. Readers are referred to a publicly available program (CO2SYS; ref 7) to carry out more advanced calculations.

In this article, we attempt to fill in the pedagogical gap between Weston’s elementary and advanced treatments. The seawater carbonate calculations we present lie between those two approaches. We treat the effects of seawater salinity and temperature, and we drop the assumption of solubility equilib-rium for CaCO3, in contrast to Weston (6, eq 8). The equations we present are amenable to simple spreadsheet applications and are thus less likely to be treated as “black box” calculations by students.

The essential chemistry is as follows: Dissolved CO2 lowers the pH of water by dissociating into protons and bicarbonate. Protons can then react with dissolved carbonate to form more bicarbonate ion:

The net effect of CO2 acidification is the sum of the above two reactions (4):

2 HCO3H2O CO2 CO32

Thus the level of carbonate ion decreases when dissolved carbon dioxide increases. At a sufficiently low pH (or high CO2 con-centration) the carbonate concentration will drop below the saturation level of aragonite (CaCO3) and thus cause dissolving of coral structures. This scenario is most likely to be realized first in cold arctic surface waters where CO2 is most soluble (4, 5).

As noted by Weston (6), the chemistry of ocean acidi-fication can be used to illustrate the practical utility of many chemical topics, such as polyprotic acid ionization, buffers (the Henderson–Hasselbalch equation), effect of pressure on gas solubility (Henry’s law), and the effect of temperature on equilibrium constants (Le Châtelier’s principle).

(acidi�cation due to CO2) H HCO3H2O CO2

HCO3H CO32

(1)

A Simplified Model To Predict the Effect of Increasing Atmospheric CO2 on Carbonate Chemistry in the Ocean

Brian J. Bozlee* and Maria JaneboDepartment of Chemistry, Hawaii Pacific University, Kaneohe, HI 96744-5297; *bbozlee@hpu.edu

Ginger JahnDepartment of Chemistry, University of Maryland, College Park, MD 20742

214 Journal of Chemical Education  •  Vol. 85 No. 2 February 2008  • www.JCE.DivCHED.org  • © Division of Chemical Education

Chemistry for Everyone

Overview of Relevant Ocean Chemistry

Table 1 lists the concentrations of some of the major chemi-cal species in seawater and is modified from Dickson and Goyet (8). It should be noted that the pKa values are given for 25 °C and a salinity of 35‰ (parts per thousand). The pH of ocean water is near 8.1, and CO2* represents the sum of aqueous CO2 and H2CO3, which are difficult to distinguish experimentally (6, 8). There are many other chemical species in the ocean as well, including Na+ and Cl− ions, that do not influence the pH except through the effects of ionic strength.

The surface of the ocean is expected to drop in pH by at least 0.1 units in the next 100 years owing to acidification by CO2 (1). We will assume here that the concentrations of metal ions will not change appreciably over that pH range. Similarly, the concentrations of anions from strong acids (Cl−, Br−, etc.) will be unaffected by pH. What may change are the ratios of weak acids (HA) to their conjugate bases (A−) according to the Henderson–Hasselbalch equation,

which can be rearranged to

pH p aK10HAA

pH p a logKHA

A

(2)

Combining eq 2 with [HA] + [A−] = total concentration of A species results in

The predicted fraction of the acid form from eq 3 can be plotted, for different pH, as a function of pKa (Figure 1).

This graph shows that the fraction of HA (and thus the fraction of A−) will not change much for a 0.1 unit drop in pH, except for species with pKa values between about 7 and 9. Referring to Table 1, it therefore appears that only carbonate and borate species will be affected by a pH lowering of 0.1. Since the total concentration of borate is lower than the total concentration of carbonate, we will treat only the latter in what follows. Similarly, since the concentrations of phosphate and silicate are low in most of the ocean we will neglect these as well, although they could be significant in the near-shore environment.

A Simplified Model To Treat Carbonate Chemistry in the Ocean

We will assume that surface waters reach equilibrium with atmospheric CO2 (CO2* represents the sum of aqueous CO2 and H2CO3).

CO2*(aq)CO2(g)

The equilibrium constant for the dissolution of CO2 is called the Henry’s law constant and is given by

where PCO2 is the pressure of CO2 in atmospheres.

The dissolved carbon dioxide in turn undergoes a first and second acid ionization,

HCO3 HCO2*

CO32 HHCO3

leading to

Autoionization of water

frmole

action HA A

pH p a

1

1 10 K

pH p a10 K

HA

HA

HA

HA HA(3)

CO2

CO

*0

2

KP

(4)

1KH HCO3

CO2*(5)

CO3K2

2H

HCO3

(6)

H OHH2O

pH = 8.1pH = 8.0

change infractionof HA

Fra

ctio

n of

Con

juga

te A

cid

pKa of Conjugate Acid

1.0

4 6 8 10 12

0.8

0.6

0.4

0.2

0.0

Figure 1. Effect of pKa value on conjugate acid–base ratios at pH 8.1 and 8.0.

Table 1. Concentration of Conjugate Acid–Base Pairs in the Ocean

Acid Form (HA)

Conc/ M

Base Form (A–)

Conc/ M

Ratio [A–]/[HA]

pKa of HA

HSO4– ~0 SO4

2– 0.02824 ∞ 0.9989

HF ~0 F– 0.00007 ∞ 2.519

B(OH)3 0.00032 B(OH)4– 0.00010 0.31 8.597

CO2* 0.00001 HCO3– 0.00177 177 5.856

HCO3– 0.00177 CO3

2– 0.00026 0.15 8.925

© Division of Chemical Education  • www.JCE.DivCHED.org  •  Vol. 85 No. 2 February 2008  •  Journal of Chemical Education 215

Chemistry for Everyone

gives

And charge balance (i.e., molarity of positive charge equals molarity of negative charge) may be written as

Here, [cations] is the sum of all positive ions whose concentra-tions do not change with pH. Similarly [anions] is the sum of all negative ions whose concentrations do not change with pH. Subtracting [anions] from both sides results in

where [net cations] is the residual positive charge concentration after subtracting [anions]. Since both [H+] and [OH−] are small, we may make the approximation

HCO3 CO32net caations 2

Using the concentrations in Table 1, the value of [net cations] is therefore approximately 0.00229 M. This concentration is assumed to be pH independent so that charge balance may be finally written as

Equations 4–8 represent a system of five equations with six unknowns. From these it should be possible to write [CO3

2−] as a function of PCO2

. However, since this proves to be math-ematically difficult, we have chosen instead to express PCO2

as a function of [CO3

2−] according to the equations

COP x22

where

and

Kw OHH (7)

H HCO 23 CO32

OH

cationns

anions

H HCO3

CO32 OH

net c

2

aations

0 00229

2

mol

L.H

HCO3 CO32 OH

(8)

(9)

xb b ac

a4

2

2

(10)

aK K K0 1 2 K K0 1

CO32 (11)

0 00229 mol

L.b

K K K

K K K

0 1 2

0 1 22

12

12

CO32

CO32

(12)

wc K (13)

The effective values of the equilibrium constants (K0, K1, K2 and Kw) are given as functions of temperature (T ) and salinity (S) in the online supplement and are taken from Dickson and Goyet (8). By using these, it is possible to predict the pressure of CO2 that corresponds to any chosen carbonate concentration in seawater at any temperature and salinity.

It is worth noting that other concentrations of interest may also be expressed as functions of x, as defined in eq 10. Although these are not directly related to the saturation of calcium carbon-ate, we list them here for the sake of completeness.

Relationship to Oceanographic Parameters

Oceanographers express inorganic carbon chemistry in terms of four measurable parameters. These are the pressure of carbon dioxide (PCO2

), pH, total inorganic carbon (CT), and alkalinity (A). The measurement of any two can be used to calculate the remaining two (6, 7). In more careful work the pressure of CO2 is replaced by the fugacity of CO2, but this is only a small correction that we disregard (6, 7). Three of the four oceanographic parameters may also be expressed as functions of x and [CO3

2−]:

The alkalinity (A) may be defined 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 one kilogram of sample”(8). In the context of this model, the alkalinity is essentially the carbon alkalinity:

According to eq 1 when 1 mole of CO32− is lost, 2 moles of

HCO3− are gained. This results in no net gain or loss of alkalinity

as a result of acidification of seawater by CO2. Thus we expect A to remain nearly constant at the current value of 2290 μM.

02K xCO2* (14)

0 1x

K K

2KHCO3

CO32

(15)

0 1 2xK K K

HCO3

2 (16)

COP x22 (9)

pH log llog xK K K0 1 2HCO3

2(17)

(18)

CT

K x xK K

K02 0 1

2

CO2* HCO3 CO32

CO32

CO32

A 2HCO3 CO32

216 Journal of Chemical Education  •  Vol. 85 No. 2 February 2008  • www.JCE.DivCHED.org  • © Division of Chemical Education

Chemistry for Everyone

298 K

273 K

Concentration CO2 / ppm

[CO

32] /

(μm

ol/L

)

1000

800

600

400

200

00 200 400 600 800 1000

Figure 2. Effect of temperature (or latitude) on carbonate levels in seawater.

Latitude / deg

[CO

32] /

(μm

ol/L

)

300

200

100

350

250

150

50

080 60 40 20 0 20 40 60

pre-industrial(270 ppm)

current(370 ppm)

2100 C.E.(563 ppm)*

2100 C.E.(788 ppm)*

Figure 3. Calculated seawater carbonate concentrations as a function of latitude for pre-industrial, current, and future scenarios. (The CO2 concentrations are shown in parentheses; the data marked with an asterisk are from ref 4. Reference 9 was used to associate latitudes with approximate surface temperatures and salinities: 60o = 278 K, 32.5‰; 40o = 285 K, 34.3‰; 20o = 298 K, 35.3‰; 0o = 300 K, 35.0‰; –20o = 296 K, 35.8‰; –40o = 287 K, 34.8‰; –60o = 274 K, 34.0‰.

Predictions of Carbonate Levels in Seawater and Other Calculated Results

Carbonate concentrations ranging from zero to 1000 μM were entered into a spreadsheet column as the independent variable and the corresponding values of x were determined, using eqs 10–13. The values of x, in turn, were squared to find the values of PCO2

(eq 9). Plots were made with concentration of CO2 in ppm on the horizontal axis and [CO3

2‒] on the vertical axis. Figure 2 shows such plots for two different temperatures, assuming a salinity of 35‰. It is interesting to note that the concentration of carbonate drops off dramatically in cold water as the pressure of atmospheric CO2 increases, owing to the enhanced solubility of CO2 at low temperature (in accord with Le Châtelier’s principle for an exothermic process). At low temperatures and high CO2 levels the concentration of carbon-ate in surface seawater is predicted to fall below the saturation limit of aragonite (approximately 70 μM carbonate; ref 4) resulting in the destruction of certain marine organisms (4–6). Even in tropical regions, where the calcium carbonate remains supersaturated, the health of coral may still be compromised by lowered carbonate levels (6).

List 1. Input Options Used in CO2SYS

Type of Input Single Input

Choice of constants Mehrbach et al.

Choice of KSO4[sic] Dickson

Choice of pH scale Total scale

Total carbon 2030 μM (from Table 1)

Pressure of CO2 360 ppm

Input temperature 20 oC

Salinity 35‰

Input pressure (i.e., depth) 0 (for surface waters)

[Phosphate] zero

[Silicate] zero

Table 2. Comparison of Results CO2SYS and This Work

Parameter CO2SYS This Work

[HCO3–]/mM 1819 1836

[CO32–]/mM 199 224

[CO2*]/mM 12 12

Alkalinity (‰) 2310 2290

pH (total scale) 8.085 8.100

Further calculations are displayed in Figure 3. The pre-dicted concentrations of carbonate in seawater are plotted versus latitude for CO2 levels that correspond to pre-industrial (270 ppm), current (370 ppm), and modeled levels for the year 2100 C.E. (563 ppm from S650 model; 788 ppm from IS92a model; ref 4). Reference 9 was used to estimate the average temperatures and salinities at latitudes ranging from 60° S to 60° N. Our cal-culated carbonate concentrations are lower than those predicted by Orr et al. (4, Figure 1) at polar latitudes (as much as 33%), and higher near the equator (as much as 23%). Nevertheless we show the same trends and are in qualitative agreement (average percent difference of ‒3%). Similar to Orr et al. we find that ara-gonite (CaCO3) may become unsaturated in the surface waters of polar regions by the year 2100 AD.

We have also compared our calculations to the Oak Ridge National Laboratory’s program CO2SYS (7) using the input parameters summarized in List 1. The output results compare favorably with our own calculations (eqs 14–18; Table 2). The Oak Ridge calculations include the effect of borate, while our calculations do not. In spite of this, the agreement is quite heart-ening. This helps to justify our original assumption that borate may be neglected, to a first degree of approximation.

© Division of Chemical Education  • www.JCE.DivCHED.org  •  Vol. 85 No. 2 February 2008  •  Journal of Chemical Education 217

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Summary

We have presented a simplified model for dissolved inor-ganic carbon chemistry in seawater that is adaptable to the class-room and to spreadsheet calculations. The primary application of this model has been the acidification of surface seawater by rising CO2 levels. In agreement with oceanographic researchers we predict that CaCO3 will become less oversaturated or even unsaturated in the future, resulting in the probable destruction of calcifying organisms, especially in the Earth’s polar regions.

Literature Cited

1. Baird, C. Environmental Chemistry, 2nd ed.; W. H. Freeman and Company: New York, 1999.

2. Langmuir, D. Aqueous Environmental Chemistry; Prentice Hall: Upper Saddle River, NJ, 1997.

3. Eby, G. N. Principles of Environmental Geochemistry; Thomson Brooks/Cole: Pacific Grove, CA, 2004.

4. Orr, J. C.; Fabry, V. J.; Aumont, O.; Bopp, L.; Doney, S. C.; Feely, R. A.; Gnanadesikan, A.; Gruber, N.; Ishida, A.; Joos, F.; Key, R. M.; Lindsay, K.; Maier-Reimer, E.; Matear, R.; Monfray, P.; Mouchet, A.; Najjar, R. G.; Plattner, G.-K.; Rodgers, K. B.; Sa-bine, C. L.; Sarmiento, J. L.; Schlitzer, R.; Slater, R. D.; Totterdell, I. J.; Weirig, M.-F.; Yamanaka, Y.; Yool, A. Nature 2005, 437, 681–686.

5. Doney, Scott C. Scientific American 2006, 58–65. 6. Weston, Ralph E. J. Chem. Educ. 2000, 77, 1574–1577. 7. Lewis, E.; Wallace, D. Program Developed for CO2 Calculations,

ORNL/CDIAC-105. Carbon Dioxide Information Analysis Center, Oak Ridge National Laboratory, U.S. Department of Energy: Oak Ridge, TN, 1998. http://cdiac.ESD.ORNL.GOV/oceans/co2rprt.html (accessed Nov 2007).

8. Handbook of Methods for the Analysis of the Various Parameters of the Carbon Dioxide System in Sea Water, version 2; Dickson, A. G., Goyet, C., Eds.; ORNL/CDIAC-74; Carbon Dioxide Information Analysis Ceter, Oak Ridge National Laboratory, U.S. Department of Energy: Oak Ridge, TN, 1994.

9. Tomczak, Matthias; Godfrey, J. Stuart. Regional Oceanography: An Introduction, 2nd ed.; Daya Publishing House: Delhi, 2003; p 19, Figure 2.5a,b.

Supporting JCE Online Materialhttp://www.jce.divched.org/Journal/Issues/2008/Feb/abs213.html

Abstract and keywords

Full text (PDF) with links to cited URLs and JCE articles

SupplementThe effective values of the equilibrium constants (K0, K1, K2, and Kw), given as functions of temperature (T ) and salinity (S)