Notes 857

References from improper logging, on the bottom fauna

ANDERSON, N. H., AND D. M. LEHXIKUHL. 1967. of a small trout stream in the southern Ap-

Catastrophic drift of insects in a woodland palachians. Prog. Fish-Cult. 17: 64-70.

stream. Ecology 49 : 198-206. WATERS, T. B. 1964. Recolonization of denuded

LAGLEFI, K. F. 1970. Freshwater fishery biology, stream bottom areas by drift. Trans. Am.

2nd ed. Brown. Fish. Sot. 93: 311-315.

MENHINICK, E. F. 1964. A comparison of some -* 1965. Interpretation of invertebrate

species-individuals diversity indices applied drift in streams. Ecology 46 : 327-333.

to samples of field insects. Ecology 45 : 859-861. Submitted: 2 April 1973

TEBO, L. B. 1955. Effects of siltation, resulting Accepted: 11 March 1974

Thermodynamic limitations on the use of the platinum electrode in Eh measurements

Abstract-The thermodynamic stability of surface oxide and sulfide coatings on the platinum electrode is considered. In well mixed environments thermodynamic consider- ations and experimental evidence suggest that the platinum surface acts as an oxide electrode that responds to pH rather than to oxygen partial pressure. In stagnant condi- tions in the marine or estuarine environment the slow formation of platinum sulfide may push the measured Eh values to more negative potentials if permanently emplaced electrodes are used. These reactions may restrict the operational use of Eh measurements.

Since the work of Gillespie (1920) the potential of the platinum electrode (Eh) has been widely used to characterize nat- ural environments and on occasion to draw conclusions about the precise oxidation-re- duction conditions within a particular sys- tem. Although it is now clear that such de- ductions are only valid in rather exceptional circumstances (Berner 1963; Bricker 1965; Hem 1961) there is still a considerable interest among biologists and geologists in the use of Eh as an operational parameter for mapping variations in the level of oxi- dative degradation in the field. This em- pirical application, originally suggested by ZoBell ( 1946)) h as received considerable circumstantial support from a variety of sources (Hewitt 1950; Sato 1960; Mortimer 1971; and others).

The main physicochemical objections to the use of Eh as an operational parameter can be summarized as follows:

Experimental difficulties associated with the disturbance of the sample (release or absorption of gases, e.g. 02, H$ ) and reactions at the liquid junction of the reference electrode (precipitation of heavy metal sulfides, effect of suspended matter).

Low exchange current densities at the platinum surface and the predominance of mixed potentials should give rise to in- stability and poor reproducibility in the measurements.

The response of the platinum electrode to changes in the environment will depend to a large measure on the properties of the platinum surface and the presence of ad- herent surface coatings.

It must be stressed that these are prob- lems associated with the operational use of Eh. Additional problems arise if the Eh values are to be interpreted as oxidation- reduction potentials (Morris and Stumm 1967; Stumm 1967; Whitfield 1969). How- ever, if the objections raised above can be overcome-or at least understood in more detail-then the practical use of Eh values can proceed with greater confidence.

In previous communications I have con- sidered procedures for improving the quality of the experimental measurements (Whitfield 1969, 1971~) and for estimating the equilibrium exchange current densities at the platinum surface (Whitfield 1972). This paper will consider some facets of the final obstacle affecting Eh measurements -the properties of the platinum surface.

858 Notes

Table 1. Stability constants for reactions involving platinum metal in natural aqueous solutions at 25°C and 1 -atm pressure.

Reac tlon No.

Equation

(1)

(2)

(3)

(4)

(5)

(6)

(7)

(8)

(9)

(10)

01)

02)

(13)

(14)

(15)

(16)

pt" + H20 ' Pt-0-k + 2H+ + 2e- --

pt" + H20 ,- PtO -F 2H+ + 2e-

pt" + 2H20 a Pt02 + 4H+ + 4e-

pt" + 3H20 - PtO. 2H20 + 2H+ + 2e-

pt" + 4H20

pt" + 5H20 - pto2. 3H20 + @I+ + 4e-

pt" + 6~~0 \ pto2. @I20 + @I+ + 4e-

pt" + 3H20 - Pto3 + 6H+ + 6e-

pt" + @.I20 ----A Pto + 8~+ + 8e- -4

so + @I20 - Pt304 + 8~+ + 8e-

pt" + 2H20 - Pt(OH)2 + 2H+ + 2e-

pt" + 6H20 - Pt(OH)62- + 6H+ + 4e-

pt" + so 2- 4

+ 8~+ + 6e- - \- PtS + 4H*o

pt" + 401- \ Ptcl 2- 4 + 2e-

pt" + 601- y PtC162- + 4e-

pt" - pt2+ + 2e-

-0.88

-0.9

-2.08

-1.00

-0.96

-0.98

-1.06

-1.5

-1.6

-1.11

-0.98

-29.73'

-30.4

-141

-33.8

-64.9

-66.3

-71.6

-152

-216

-150

-33.1s

-97.555

52.0511

-24.7s

-47.7s

-qo.65

*Calculated only for the formation of adherent oxide phases.

tPt-0 represents a layer of adsorbed oxygen atoms on the platinum surface rather than a layer of oxide, PtO (Hoare 1968: pa 34). Solid phases are underlined.

$Data from Hoare (1968) unless otherwise stated.

§De Bethune and Swendeman-Loud (1964).

II See text.

When Eh measurements are made, the consequences of direct attack of the plati- platinum electrode is assumed to be an num surface. It will then be possible to inert sensor providing a site for electron decide how far these fundamental processes exchange (Whitfield 1971b). If the metal support or militate against the use of Eh actually reacts with substances dissolved as an environmental parameter. in the water then its properties will alter The stability of various coatings on the and it will display potentials that are char- platinum surface will be calculated from acteristic of the electrode rather than of reversible thermodynamics. Since this form the environment. Here I will explore the of analysis assumes that all reactions attain

Notes 859

equilibrium, it results in an oversimplified picture of the state of the platinum surface. In practice kinetics will control the transi- tion from one surface coating to another. At present it is not possible to consider in detail the influence of these rate processes in such complex aqueous media.

Stability constants for the most important equilibria affecting the chemistry of the platinum surface are collected in Table 1. Data for the formation of platinum oxides and hydroxides were taken from the com- pilation of Hoare ( 1968, p. 19). This list- ing is more complete than other tabulations (Pourbaix 1966; Sillen and Martell 1964) and includes data on most of the phases that have been considered in electrochemi- cal investigations of the platinum-oxygen system. Data for the dissolved phases were taken from DeBethune and Swendeman- Loud ( 1964) and adjusted to refer to direct attack of the platinum metal. The relation- ship for the formation of PtS( s) was de- rived from the equations (DeBethune and Swendeman-Loud 1964)

log10 K” Pt” + H,S( aq) zs PtS(s) + 2H++ 2e-= 11.05 SOI'- + 10 H’ + 8e- e H&(aq) + 4Hz0 = 41.0

Pt” + SOI- + 8H’ + 6e- e PtS(s) + 4H,O = 52.05

Since sulfate reduction is actually mediated by bacterial action and rarely runs to com- pletion in the natural environment this equation will set a pessimistic upper limit for the onset of platinum sulfide formation. The stability constants used refer to in- finite dilution. No attempt was made to correct for the effect of ionic strength (e.g. Stumm 1967; Dyrssen et al. 1968) since the corrections involved ( -+ 10%) will not ma- terially affect the deductions drawn from the data; they will definitely be less than the uncertainties in the values of the sta- bility constants used. For the formation of the oxides or hydroxides, the equilibrium relationships (equations l-11 : Table 1) can be written in the form

log (aoX : aFto) - x(pH - pE) = log K”; (17)

pE (= -log,, a,-) corresponds to the oxida- tion power of the system at equilibrium

PE

Fig. 1. Stability for platinum compounds in solutions where pH = 8.3, pC1 = 0.45, and pS0, = 2.47 at 25°C and l-atm pressure. Numbers on the curves identify the equations in Table 1 that describe the formation of the various components, e.g. line 6 refers to the formation of PtO, * 3H,O.

(Sill& 1965). aos is the activity of the oxide or hydroxide phase and apt0 is the activity of the platinum metal. The corre- sponding equations for the formation of the sulfide and of the chloride complexes are

log (arts : al+) + pS04 + 8pH + 6pE = 52.05, (18)

log (aPtCI,?-: apto) + 4pCl- 2pE = -24.7, (19)

and

log (aPtC1,2-: arto) + 6pCl- 4pE = -47.7, (20)

where pCl = -logl,,acl etc.

A plot of the activity ratio of complex to free metal [log (apti : apto) ] versus pE with given values of pH, pC1, and pS04 will indicate which complexes are predominant at different points in the pE range (Sillen 1959; Butler 1964). Since the stoichio- metric numbers (x) for pH and pE are equal for each oxide phase (equation 17), the relative positions of the lines will not be altered by pH shifts and so a single activity ratio diagram can be used to select the oxide phases to be considered as stable coatings on the platinum surface over the entire pH range.

860 Notes

+8

PE

-8

4 8

P”

Fig. 2. Predominance area diagram for plati- num compounds on the pH-pE plane for solutions where pC1 = 0.45 and pSOI = 2.47. The vertical line corresponds to the conditions shown in Fig. 1. Broken line indicates the corresponding boundary in freshwater with pSOa = 3.92. The shaded area indicates conditions where water itself is unstable.

The diagram for seawater at pH 8.3 (Fig. 1) indicates that the solid phases PtS, Pt”, Pt-0, and PtO, l 2H20 are stable over different spans of the pE range. Using this information we can construct a diagram which shows, on a pH/pE plane, the areas of predominance of the various phases (Fig. 2). The boundaries separating the stability regions can be obtained from equations 1, 5, and I3 by setting the ratios uos : uPtO and arts : aPta equal to one in equations 17 and 18 respectively. The stability boundaries for the formation of the chloride complexes (PtC1b2- and PtClG2-) can be obtained by setting the ratio aPtO: aPtCiX2- equal to unity in equations 19 and 20. These boundaries will be independent of pH (Fig. 2) but their position will depend on the chloride ion activity, shifting to higher pE values as pC1 increases. The anodic dissolution of platinum electrodes in acid chloride solu- tions is well documented (Frankenthal and Pickering 1965).

In well aerated environments oxygen

is the dominant electron acceptor and the formation of an oxide phase on the surface of the platinum electrode has a significant influence on its behavior. Oxygen rapidly becomes adsorbed onto a fresh platinum surface at moderate oxygen pressures and the “derma-sorbed” layer, once formed, is difficult to remove (Hoare 1968). Further- more there is evidence that oxygen actually dissolves in the bulk platinum and if the surface oxide layer is removed by cathodic stripping in a nitrogen-saturated solution then a new oxide coating is formed by dif- fusion from the body of the metal. Several monolayer equivalents of oxygen may be dissolved in the platinum so that even under conditions where the free platinum surface is thermodynamically stable, oxide coatings may still persist.

Despite detailed investigations (e.g. Hoare 1968; Ives and Janz 1961) the nature of the oxide coat remains unresolved. Ac- cording to Hoare (1968, p. 41) “the defini- tive experiment is yet to be done to de- termine conclusively in what form the adsorbed oxygen on platinum exists.” Pesh- chevitsky et al. (1967) have suggested that the formation of Pt ( OH) 2 *xHzO phases might be predominant in controlling the potential of the platinum electrode in the natural environment. Stumm and Morgan (1970) suggest that the formation or de- composition of Pt ( OH) may give rise to spurious potentials in natural media. These suggestions would appear unlikely on the basis of the available thermodynamic data ( Fig. 1). Hoare ( 1968) noted specifically that Pt(OH) 2 was not considered in the interpretation of electrode potentials be- cause it is very unstable under normal conditions. Most electrochemical investiga- tions have suggested (as is borne out by Fig. 1) that the stable phases are probably Pt-0, PtO,, or various admixtures of these components. These coatings are electron- ically conducting, and, if the surface layer is intact, the platinum electrode should record the reversible oxygen potential ac- cording to the equation

O2 + 4H+ + 4e- * 2H20 E" = 1.229 V. (21)

Notes 861

Eh (VI

0 4 8

PH Fig. 3. Predominance area diagram for platinum compounds ( Fig. 2 ) superimposed on the collected

observations of Baas-Becking et al. (1960) for E/z/pH data in natural aqueous environments. Lines a, b, and d correspond to the boundaries shown in Fig. 2. Line c follows the “irreversible” oxygen potential (Merkle 1955). Vertical boundaries represent pH barriers as discussed in the text. The posi- tion of the stability boundaries for the chloride complexes are for solutions where pC1 = 0.45.

Region A includes measurements made in well mixed environments where the oxygen partial pres- sure is high. Region B contains data from stagnant environments where oxygen has been replaced by other electron acceptors.

862 Notes

The electrode potential . would then re- spond to changes in the oxygen partial pres- sure. The contribution of this reaction will however be negligible unless impurities are rigorously removed from the solution. When the impurity concentration ap- proaches 10-lOM the exchange current den- sity of the reversible oxygen electrode is already being swamped (Bockris and Reddy 1970). In natural media one would expect the platinum electrode to act pri- marily as an oxide electrode, measuring PH.

This suggestion, which has also been discussed by Peshchevitsky et al. ( 1967)) is most conveniently checked by consider- ing the data of Baas-Becking et al. (1960) who collected over 4,000 Eh/pH readings and used them to delineate chemically dis- tinct environments. Although their detailed boundaries might appear rather speculative in the light of the more recent work of Stumm ( 1967) and Sillen (1965), their paper contains the most comprehensive survey of empirical Eh values to date. The readings they record cover most of the Eh/pH span available, but more than 90% of their data on natural aqueous environ- ments lies within a fairly restricted area (Fig. 3).

At high Eh values (region A, Fig. 3) the platinum electrode exhibits a pH-sensi- tive response that is described quite ac- curately by the equation summarizing the function of the Pt-0 coated electrode as a pH electrode:

E = Eopt,rt.O - 0.06pH [ E”~~t,pt~O = 0.88 V] .

The bulk of the readings fall within k-53 mV of this line ( Fig, 3, line b) . This same equation was attributed by Peshchevitsky et al. (1967) to the pH function of Pt( OH)a l 2H20 on the platinum surface. The calculated E” value is similar to that observed experimentally by Baas-Becking et al. (1960: E” = 0.81 V) and by Cooper (1937: E” = 0.846 V) in oxygenated solu- tions at various pH values. These E” values are also close to that for the peroxide re- action (Ives and Janz 1961).

H202 + e- = OH + OH- E” = 0.835 V.

Fig. 4. Stability diagrams analogous to Fig. 1 for the other platinum group metals with pH = 8.3. Only the stable phases are shown for clarity. Data to *construct the diagram were taken from Hoare ( 1968). The stable phases identified were ( reading from the high pE phase in each case ) OsOa; RuOz; IrOz, Ir203, Ir-0; I&03, Rh-0; PtO, .2H20, Pt-0; PdO.

This agreement must be considered coinci- dental unless it can be shown that the necessary concentrations of peroxide ( > 10~~M) exist in the natural environment. The corresponding equations for the PtO,* 2H20 oxide electrode (E” = 0.96 V: line a, Fig. 3) and for the so-called “irreversible” oxygen electrode (E” = 0.70 V: line c, Fig. 3: Merkle 1955; Garrels and Christ 1965) appear to set upper and lower limits to this pH function but cannot be considered representative of the bulk of the data points.

The other noble metals (Rh, Ir, Pd, Ru, OS) form similar oxide coats and are liable to show pH functions with E” values close to that of platinum under the same condi- tions (Fig. 4). Gold does not form a stable oxide coating but it is characterized by very small exchange currents in the natural environment (Whitfield 1972) and does not give a useful Eh electrode. The data points shown in region A ( Fig. 3) were taken exclusively in environments where the oxygen partial pressure is high and include

Notes 863

normal soils, shallow groundwater, sea- water, and river water. Meteoric waters account for most of the readings on the acid side of the pH 6 boundary.

Under these circumstances the platinum electrode will not respond to changes in the partial pressure of oxygen. The Eh value will therefore not give any indication of the oxidation-reduction status of the system unless other electroactive couples are pres- ent in sufficient concentration ( > 10e5M) to dominate the electrode potential (Pesh- chevitsky et al. 1967; Bohn 1968). There is evidence to suggest that iron (Hem 1961) and manganese ( Bricker 1965) couples may occasionally act in this way. From the operational point of view the consequence of the “oxide electrode” func- tion of the platinum surface is to associate large, positive Eh values with well aerated systems. If we exclude meteoric waters the Eh values for such systems should range from +500 mV to +300 mV.

Readings represented in region B (Fig. 3) reflect the behavior of the platinum electrode in stagnant environments where oxygen has been replaced by other electron acceptors. This region includes measure- ments made in freshwater lakes, evaporite basins, and in the interstitial waters of freshwater, marginal marine, and open-sea sediments. The vertical boundaries of region B in Fig. 3 reflect the importance of the carbon dioxide system in buffering the pH of isolated or stagnant environ- ments. The buffering capacity of the car- bon dioxide system is at a maximum when pH = pKi and pH = p&, where pK1 and pK2 are the negative logarithms of the first and second acidity constants respec- tively of the carbon dioxide system. The relevant values at 25°C are pH 6.4 and 10.4 in freshwaters and pH 6.0 and 9.2 in sea- water (Stumm and Morgan 1970). Pytko- wicz ( 1967) has convincingly demonstrated the role of the carbon dioxide system in buffering the oceans against changes in pH over periods of thousands of years. If this mechanism is accepted then the pH of natural systems will tend to vary around the buffer intensity minimum for the carbon

dioxide system between the peaks at pH = pK1 and pH = pK,. In all but the softest waters the pH will rarely rise above 9.5 because of the onset of precipitation of cal- cium and magnesium carbonates.

In region B (Fig. 3) the platinum elec- trode is no longer responding primarily to pH. Either the formation of the oxide coat- ing is precluded by the low pE values experienced as oxidative degradation pro- ceeds-as is suggested by the thermo- dynamic evidence-or else the platinum surface is responding to other redox systems -either as mixed potentials, reaction po- tentials, or redox potentials. There is evi- dence to suggest that the platinum elec- trode is responsive to changes in the status of the nitrogen (Bailey and Beauchamp 1971)) iron ( Doyle 1968), and sulfur (Berner 1963) systems as oxidative degra- dation proceeds. The incomplete oxide coating which probably still persists, de- spite the increasing stability of the clean platinum surface, will not hinder these other reactions since it is electronically conducting and able to transmit the effects of surface reactions to the bulk metal. However, at low pE values platinum sul- fide becomes the stable phase (see also Najdeker and Bishop 1973) and this might present a lower limit to the usefulness of the platinum electrode. In reduced systems containing sulfur it has been shown over a wide geographical range (Berner 1963; Skopintsev et al. 1966; Whitfield 1969) that the platinum electrode responds to the po- tentials established by the system

HS-(aq) + SO(rhomb.) + H+ + 2em.

This is equivalent to the half cell

S2- * So (rhomb.) + 2 e-,

for which

Eh = -0.485 + 0.0295 pS.

We can rewrite the equation for the forma- tion of platinum sulfide in a similar form so that

Pt” + S2- * PtS + 2e- log K = 32.05

for which

Eh = -0.96 + 0.0295~s.

864 Notes

Therefore one would expect the attack of the platinum surface by sulfide to be ac- companied by a relatively sharp drop in Eh. Apparently the rate of attack is slow and the experience of Hayes et al. (1958) is typical. They left a platinum electrode overnight in sulfide-bearing mud and found that the potential had drifted downward by 0.36 V. This result is of the same order as the predicted drop of 0.47 V. The elec- trode, when removed from the mud, had a distinct amber tarnish that was difficult to remove even under extremely oxidizing conditions ( coned nitric acid). The slow- ness of the tarnishing reaction and the di- rection of drift in potential were also con- firmed by Baas-Becking et al. (1960).

A thermodynamic analysis of the be- havior of the platinum electrode over the pE/pH range of the natural environment indicates that the metal is susceptible to attack with the formation of oxides in the high pE range and platinum sulfide in the low pE range. The predictions of this an- alysis are borne out by a consideration of the trends in the Eh/pH data accumulated by Baas-Becking et al. ( 1960). The plati- num electrode appears to be acting as an oxide electrode in well aerated systems and it responds to pH in a manner analogous to that predicted for the Pt-0 electrode. It is, however, insensitive to changes in the oxygen partial pressure. In the intermediate region where the thermodynamic analysis predicts that free platinum is stable, the surface of the electrode is still likely to be coated with oxide since the rate of release of oxygen by the platinum is very slow compared to the reading times normally used in fieldwork. The behavior of the electrode in this region has been inade- quately studied, although it has been shown that under restricted circumstances the electrode may respond to redox reactions in the manganese, iron, nitrogen, and sulfur systems. These reactions, where they are established, continue the hierarchy of de- creasing Eh values with the progressive oxidative degradation of organic matter. The attack of the platinum by bisulfide

ions has been confirmed by observations in sulfur-bearing muds and a correlation in sign and magnitude has been confirmed between the potentials observed and those predicted by the thermodynamic equations. These points taken together indicate that the operational use of Eh measurements may be restricted to region B of Fig. 3.

M. Whitfield

The Laboratory Citadel Hill Plymouth PLl 2PB England

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Notes 865

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Submitted: 25 October 1973 Accepted: 28 February 1974