pH BUFFERING ANOXIC

OF PORE WATER OF RECENT IMARINE SEDIMENTS’ Sam Ben-Yaakov2

Department

of Geology,

University

of California,

Los Angeles

90024

ABSTRACT A model is proposed to explain the relative pH stability in pore water of recent anoxic marinc scdimcnts. Oceanic data arc not inconsistent with a model which assumes that the pH of the port waters is controllecl by the byproducts of organic decomposition, sulfate reduction, and precipitation of sulfide and carbonate. The model predicts that the pH of pore waters should remain in the range 6.9 to 8.3, which is in agreement with mcnsured values,

Anoxic

marinc

scdimcnts

gcnerdly

alkalinity of 2.2 mmole kg-l (Fig. 1). Maximunr buffer capacity is reached at low pH values below the first apparent dissociation constant of carbonic acid ( pK1) which is about 6.0 (Lyman 1957), but the buffer capacity around the pH of ocean and pore waters is very low and dots not exceed 0.5 mmolc kg-l pH-1. That is, an addition of 0.5 mmolc kg-l of CO2 will reduce the pH of scawatcr by more than a pH unit, It is thcrcforc surprising that the concentration of CO2 in port water of marine sediments may reach values of 60 mmole kg-l, and yet show pH shifts less than one unit. It is evident that the increase in CO2 must be counter balanced by other processes that tend to increase the pJ1 of seawater (such as the production of ammonia and incrcasc in total alkalinity) so that the net balance results in a fairly constant pH. The purpose of this study is to examine the chemistry of port water of anoxic marine sediments and to propose a model that many explain mechanisms con trolling its pII. The derived relationships will then be compared to field data. The pH buffering model presented hcrc differs from the model of Thorstenson (1970) in that it considers the roles of diffusion and chemical equilibria with solid phases and dots not assume thermodynamic equilibria bctwcen all the spccics.

con-

tain active sulfate-reducing bacteria ( ZoBell and Rittcnberg 1948) which oxidize organic matter through the reduction of sulfate to sulfide. This process consumes sulfate from the port water and produces a variety of protolytic spccics, such as H2S and NH3, whose amount is large in comparison to their amount in the original seawater. For example, Prcslcy ( 1969) has found that the concentration of dissolved CO2 in poro water of rcccnt scdimcnts (c.g, Saanich Inlet, B.C. ) may reach the value of 60 mmolc kg-‘, about 30 times the total CO2 conccn tration in the overlying waters. Despite the vast amount of protolytic spccics added to the pore waters during dccomposition of organic matter, the pH of pore water of rcccnt marinc scdimcnts is fairly constant and rarely cxcceds the limits pH 7.0 to 8.2. This stability can better be exemplified, perhaps, if one considers the buffer capacity of seawater with rcspcct to (,KO,) dissolved COZ. This paramctcr is defined hcrc as the incremental change in total dissolved inorganic CO2 rcquircd to shift the pI1 of a solution by one pH unit. The buffer capacity of average seawater was calculated using the method of BenYaakov (1970n) for scawatcr with total -___ l This study was supported by U.S. Atomic Energy Commission Contract AT( O4-3)-34 F.A. 134. 2 I wish to thank Professor I. R. Kaplan for critically reading the manuscript nncl M. Goldhaber and J, Cline for fruitful discussion related to the research. LIMNOLOGY

ANT)

OCEANOGRAPHY

A METEIOIJ

FOR PII

CALCULATION

It is convenient to divide the ions in a complex system, such as seawater, into protolytic and nonprotolytic species. Protolytic 86

JANUARY

1973, V. 18(l)

pH

BUFFERING

rcspcct to addition (or removal) of weak acids, weak bases, or gases to the solution. This is so bccausc the addition of a weak acid, for cxamplc, does not add any net charge to QI> since the charges of the added anions arc balanced against the charge of the added hydronium ions. The total charge held by a weak monobasic acid ( Qnl ) can bc calculated from :

6.00~

1.00. \ OS.0

c--z+

--h-m-7.0

6.0

6.0

.

87

OF PORE WATER

10.0

PH

Fig. 1. Buffer capacity of seawater with respect to total dissolved COs (broken line). Solid line represents the concentration of total dissolved CO, at any given pII. Curves arc for TA = 2.2 mcq kg-’ (based on data in Ben-Yaakov 1970b ) ,

ions arc dcfincd as ions of weak acids or weak bases; a “weak acid’ (“weak base”) is dcfincd hcrc as one whose total anionic ( cationic) charge changes appreciably over the pI1 range of interest. The principle of charge neutrality requires that the net total charge (i.e. the diffcrcncc bctwccn the total charge of the cation and the total charge of the anions) held by the protolytic ions ( Q1>) bc balanced against the net total charge held by the nonprotolytic ions ( QAT1>).Namely:

Qr+Qm=O

(1) The value QIa (or QNP) is closely rclatcd to the parameter “alkalinity” or “titration alkalinity” ( Anderson and Robinson 1946) which is often used in chemical occanography. IGwcver, this equivalency breaks down when the sample water contains weak bases such as ammonia. Total alkalinity (TA) of a given solution is equal to the charge held by ions of weak acids ( IICOs-, CO:i2-, etc. ) plus the concentration of the undissociatcd bases such as Nl&OII, but dots not include the concentration of NI&+. Thcsc differcnccs may result in some confusion and the term alkalinity will not bc used here in conjunction with port waters. In open ocean waters : TA = IQiwJ= IQPI. (2) The total charge held by each part of is invariant with the solution

Qdl=

-T (acid)

“’ K’1 + all”

(3)

whcrc T( acid) is the total concentration ( dissociated plus undissociated species) of the acid in the solution, K’1 is the apparent dissociation constant of the acid and &I+ is the activity od the hydronium ion. Apparent dissociation constants rather than thermodynamic constants arc usccl to pcrmit analysis in terms of concentration rather than activities. The error in applying the apparent constant (which can bc defined Only for solutions of constant composition) to port water is probably insignificantly small, bccausc the ionic strength of intcrstitial water of unlithified scdimcnt is similar to that of the overlying scawatcr. The total charge held by the ions of a dibasic acid can bc calculated from: K’I aI I.+ 2 K1 Kt2 en'

=-T(acid)[aII+]2

+ KlaHt.+

K/1.,2'(4)

whcrc K’l and K12 arc the first and second apparent dissociation cons tams of the acid. Similarly, the total charge held by the cation of a monoacidic base can bc calculatcd from the cxprcssion:

QH = T(base)K’

K’ + ( aHa0

. K,&zII+)

4%

whcrc Klv is the thermodynamic dissociation constant of water and a1120 is the activity of water at the given ionic strength. The apparent dissociation constant of a weak acid is rclatcd to the thermodynamic constant by:

88

SAM BEN-YAAKOV

whcrc Y&N and yB+ arc the activity cocfficicnts of the base and its cation, rcspcctively. The pII of a solution containing both weak acids and weak bases can now bc calculated from the expression: QP=ZQ,