photolysis rates of organomercurials at various depths in natural
waters by measurements of sunlight intensities.

Light absorption by an organomercury compound may be expressed by’
the equation

hy
(RHgx)° = (RHgx)* (38)

where (RHgX)° and (RHgX)* represent the compound in its ground and
electronically excited states, respectively. The primary quantum
yield for any photochemical process is simply the fraction of
electronically excited molecules that undergo the process. If the
electronically excited molecule undergoes two or more competitive
primary photoprocesses, the sum of the quantum yields for these
processes is unity, but the quantum yield for a specific process is
less than unity. In liquid-phase photochemistry, the measured
quantum yield for photoreaction of a compound is often affected by
secondary chemical reactions. For example, the primary quantum
yield for photocleavage of an organomercury compound (eq 39) may be
high, but the measured quantum yield may be lower because of the
secondary reaction in equation 40.

(RHEX)® > Re + *HeX (39)
Rt + ‘HgX 7 (RHgx)° (40)

Photodecomposition of weakly absorbing compounds can often be
accelerated by addition of other compounds that absorb light more
strongly. Such acceleration, or photosensitization, can result
from electronic energy transfer from the strong light-absorber or
sensitizer (8) to the photoreactive compound (eq 42).°* ®

hy
(sy = (S)* (1)
(S)* + (RHgx)° = (S)° + (RHgx)* (42)

Since numerous substances that absorb sunlight more strongly than
organomercury compounds are present in natural waters, one goal of

this study was to define those types of compounds that can photo-
sensitize degradation of organomercurials,

Quantum yields for photoreactions are sometimes lowered by energy
transfer from the electronically excited, photoreactive molecule to
a quencher molecule (Q).°°

(RHgX)* + (@° > (RHgx)° + (Q)* (43)

52


SECTION V

MATERIALS AND METHODS

MATERIALS

Reagent grade dimethylmercury and methylmercuric salts were purchased
from various commercial sources, Dimethylmercury was purified by
distillation: bp 94-95° (760 mm). Diphenylmercury and phenylmercuric
salts were obtained from the Perrine Primate Laboratory, U. S.
Environmental Protection Agency, Perrine, Florida, Diphenylmercury
was chromatographed on Woelm basic alumina (activity 1), and
recrystallized from 95% ethyl alcohol. Phenylmercuric acetate was
recrystallized from 5% acetic acid in water and phenylmercuric
nitrate was recrystallized from 95% ethyl alcohol. Reagent grade
mercuric acetate was purified by recrystallization from glacial
acetic acid.

Water that was distilled, passed through ion exchange columns, and
redistilled was used in all experiments except those carried out
with natural waters, The natural water samples were obtained from
a pond near Athens, Georgia, and a western North Carolina stream
located at the Coweeta Hydrologic Laboratory.

Benzene was acid-washed, dried, and distilled. Other reagent grade
solvents were used as received. Cis-1,3-pentadiene was distilled
(bp 42-43°, 760 mm), and stored under nitrogen at -20°C. Reagent
grade thioglycolic acid and 2-mercaptoethanol were used as

received and stored at -20°C. Cysteine hydrochloride and NagS-9H20
were used without purification. Reagent grade mercuric oxide, acids,
and bases were used as received.

ANALYTICAL PROCEDURES

Acidolysis of dimethylmercury was followed by two analytical
techniques: (1) gas liquid chromatography (gle) of dimethylmercury
on a Porapak QS column using tetrahydrofuran as an internal

standard, and (2) a conductometric technique, described in the
literature.°® Aqueous solutions of methylmercuric and phenylmercuric
salts were analyzed by the dithizone method of Gran .*° Phenyimercuric
salts were also analyzed by ultraviolet spectroscopy.°° Diphenyl-
mercury acidolysis in water was followed by ultraviolet spectroscopy,
as described by Kaufman and Corwin,®? Kinetics of the mercuric-salt
cleavages of aqueous dimethylmercury and diphenylmercury were
determined by measuring the disappearance of mercuric salt by
flameless atomic absorption (AA) spectroscopy.°”

18


RESEARCH REPORTING SERIES

Research reports of the Office of Research and
Monitoring, Environmental Protection Agency, have
been grouped into five series. These five broad
categories were established to facilitate further
development and application of environmental
technology. Elimination of traditional grouping
was consciously planned to foster technology
transfer and a maximum interface in related
fields. The five series are:

1. Environmental Health Effects Research
2. Environmental Protection Technology
3. Ecological Research

4. Environmental Monitoring

5. Socioeconomic Environmental Studies

This report has been assigned to the ECOLOGICAL
RESEARCH series. This series describes research
on the effects of pollution on humans, plant = and
animal species, and materials. Problems are
assessed for their long- and short-term
influences. Investigations include formaticn,
transport, and pathway studies to determine the
fate of pollutants and theix effects. This work
provides the technical basis for setting standards
to minimize undesirable changes in living
organisms in the aquatic, terrestrial and
atmospheric environments.


Table 12.

EFFECT OF QUENCHERS UPON PHOTODECOMPOSITION OF
SULFUR-BONDED METHYLMERCURY COMPLEXES IN WATER®

10° [9]
Complex Quencher (Q) M tug.
CH, HgS CH, CO2H None 0 0.16
2,4-hexadien-1-ol 86 0.063
2,4-hexadien-1-ol 57 0.075
2,4-hexadien-1-ol 38 0.082
2,4-hexadien-1-ol 19 0.093
sodium trans-cinnamate 20 0.085
oxygen 1.4° 0.075
CH HgSCH2 CH, OH None 0 0.14
2,4-hexadien-i-o1 39 0.043
2,4-hexadien-1-ol 4.9 0.046
2,4-hexadien-1l-ol 0.49 0.087
CHgHgSTM None 0 0.65
Oxygen 1.4° 0.08
“Quenching studies carried out at 313 nm under conditions in which

no light absorbed by quencher.
° Quantum yield for formation of inorganic mercury,
°Oxygen-saturated water.

61


10.

SECTION VIL

REFERENCES

T. Takeuchi, presented in part at the International Conference
on Environmental Mercury Contamination, Ann Arbor, Michigan,
September 1970.

K. Borg, H. Wanntorp, K. Erne, and E, Hanko, J. Appl. Ecol., 3
(Suppl.), 171 (1966).

L. Friberg, et al., Nord. Hyg. Tidskr,, 52, Suppl. 4 (1971).

A, Jernelév, "Chemical Fallout," M, Miller and G. Berg, Eds.,
Charles C, Thomas, Publisher, Springfield, Illinois, 1969,
Chapter 4.

(a) F. M, D'Itri, Technical Report No. 12 to the Michigan
(USA) House of Representatives (1971); (b) N. Nelson, et al.,
Environ. Res., 4, Chapter 1 (1971); (c) P. A. Krenkel, R. S.
Reimers, E. B, Shin, and W. B. Burrows, "Mechanisms of Mercury
Transformations in Bottom Sediments," Technical Report No. 27,
Environmental and Water Resources Engineering, Vanderbilt
University, Nashville, Tennessee, 1971.

(a) F. R. Jensen and B. Rickborn, "Electrophilic Substitution

of Organomercurials," McGraw-Hill Book Co., New York, 1968;

(b) 0. A, Reutov and I. P. Beletskaya, "Reaction Mechanisms of
Organometallic Compounds," North-Holland Publishing Co.,
Amsterdam, 1968; (c) L. G. Makarova and A. N. Nesmeyanov in
"Methods of Elemento-Organic Chemistry," Vol, 4, A, N. Nesmeyanov
and K, A. Kocheshkov, Ed., North-Holland Publishing Co.,
Amsterdam, 1967; (d) L. G. Makarova in "Organometallic Reactions,"
Vol. I, E. I. Becker and M, Tsutsui, Ed., Wiley-Interscience,

New York, 1970; (e) T. G. Brilkina and V. A. Shushunov, "Reactions
of Organometallic Compounds with Oxygen and Peroxides," CRC Press,
Cleveland, Ohio, 1969, Chapter 2.

D. Seyferth, Organometal., Chem. Rev. B, 5, 288 (1969).

D. Seyferth, Organometal. Chem. Rev. B, 6, 239 (1970).

D. Seyferth, Organometal, Chem, Rev. B, 8, 425 (1971).

(a) K. C,. Bass, Organometal. Chem. Rev., 1, 391 (1966); (b) L. G.
Makarova in "Organometallic Reactions," Vol. II, E. I. Becker

and M, Tsutsui, Ed,, Wiley-Interscience, New York, 1971.

72


SECTION IV

BACKGROUND

Chemists have long been intrigued by reactions of mercury compounds,
and a vast body of literature on this subject has accumulated during
the past century. Since a review of that work is well beyond the
scope of this report, the interested reader is referred to several
excellent books that have recently been published.® Also, annual
reviews of mercury chemistry have been compiled since 1968 by
Professor Dietmar Seyferth and his coworkers.’~° Bass and Makarova
have recently reviewed the photochemistry of organomercurials.?°

Since most past studies of mercury chemistry were not prompted by
environmental considerations, relatively few studies were carried
out under conditions that could be extrapolated to the environment.
A review of the environmentally pertinent literature is presented
with particular emphasis on publications related to kinetics of
organomercury degradation,

EQUILIBRIA OF MERCURY REACTIONS IN WATER

Early studies of aqueous solutions of mercuric, >> alkylmercuric,}°
and phenylmercuric’® compounds indicated that many of these com-
pounds react with water to form acidic solutions. Subsequent
investigations’*~?® showed that reactions 1-5 account for the acid
formation, The symbol X represents any electron-withdrawing ligand
that forms an ionic bond with mercury, and R stands for an organic
group such as methyl or phenyl. Mercury compounds with ionic Hg-X
bonds are often designated as mercury "complexes" or "salts" in the
literature.

HeXe 2 Het + 2x7 qd)
Hg?* + 4,0 = Hgow +H (2)
HgOH’ + HO = Hg(OH)2 + HY (3)
RHgX = RHg’ + xX (4)
RHg’ + H,0 = RHgOH + Hf G)


Table 14.

QUANTUM YIELDS FOR DIRECT PHOTOLYSIS
(313 nm) OF PHENYLMERCURY COMPOUNDS

Disappearance
Compound Solvent Quantum Yield
(CgHs ) oHg Benzene 0.33 + 0.01
(Cele )oHg Acetonitrile 0.40 + 0.02
(CeHs )oHg Water, pH 7.0 0.27 + 0.02
Ce Hs HgC10, Water, pH 2.3 0.25 + 0.02
Cy Hs HgC1Q; Water, pH 10.3 0.24 + 0.02

“Extrapolated using viscosity data.

67


SECTION VI

RESULTS AND DISCUSSION

CALCULATED EQUILIBRIUM CONCENTRATIONS OF METHYLMERCURIC SPECIES IN
AQUATIC SYSTEMS

Although the literature is devoid of quantitative information about
formation of organomercury complexes in the environment, previous
studies have provided abundant information concerning equilibrium
constants for methylmercury complexes in water.1®1+° We have devised
a technique utilizing these equilibrium constants to calculate
relative concentrations of methylmercury complexes in aqueous systems
containing chemical species found in natural waters, i.e., hydroxide,
chloride, hydrogen sulfide and its dissociated forms, thiols (RSH),
amines (RNHz), phenols (humic acid), ammonia, and orthophosphate ,°*
Bicarbonate and organic carboxylic acids were not considered because
they form very weak complexes with the methylmercury moiety. To
render these calculations relevant to the aquatic environment, we
used recently measured concentrations of chemical species in Lake
Erie®® where high concentrations of mercury have been found in fish.

66

The Project Hypo study of the Lake Erie central basin hypolimnion
showed that high concentrations of hydrogen sulfide are present in
the water column near the bottom sediments.°” Formation of locally
high concentrations of hydrogen sulfide is a phenomenon observed in
many natural waters.°* Complexation of the methylmercury group by
reduced sulfur species (HpS, SH, and S°~) results in the formation
of two species, methylmercuric sulfide ion (CHgHgS ) and bis(methyl-
mercury) sulfide ((CHgHg)2S).1®° The two species are in equilibrium
as shown in equation 18.

2CH,HgS’ + HY = (CHjHg)2S + SH (18)

Qur calculations indicate that these two species should account for
very high fractions (> 95%) of the methylmercury complexes in
natural waters that contain reduced sulfur species (Figures 6 and 7).
Relative concentrations of CHgHgS and (CH3;Hg)2S depend upon the pH
and total concentration of reduced sulfur species, At high concen-
trations of reduced sulfur species (i0o-® to 10°* Mw), CHgHgS” accounts
for virtually all of the complexed methylmercury in the pH 5-9 range
(Figure 6). As the reduced sulfur concentration drops, (CHgHg)2S
becomes the predominant methylmercury complex in acidic waters
(Figure 7) and CH,HgS remains the predominant complex in basic
waters (Figure 6).

25


Table BA,

DEGREES OF DISSOCIATION OF 10 ORGANOMERCURY COMPOUNDS

IN AQUEOUS SOLUTION AT 25°C

Degree of dissociation (%) at
Compound 10* M 10-7 M 10°?° M
CgHgHg0COCHs 32 > 99 > 99
CgHsHgSR < 0.1 < 0.1 < 0.1
CHg HgC1 21 98 > 99
CHgHgBr 4.8 75 > 99
CH, HgL 0.5 3.3 91
CH, HgOCOCH, 76 > 99 > 99
CH; HgHPO,~ 26 99 > 99
CH HgNH2R* 1.6 39 > 99
CH3 HgSR < O11 < OL < 0.1
CH3HgSTM < 0.1 < 0.1 < 0.1

96


_ 1
4” [Ake

t: (24)

where [A] = [B]

Another special case results when the effective concentration of one
reactant does not change with time (pseudo-first-order reaction).

This occurs when one reactant is present in large excess or its
effective concentration does not change with time because of buffering.
Half-life of reactant A is dependent only on the concentration of the
reactant in excess [B] and the magnitude of the rate constant (kz)

(eq 25).

_ 0.693

(25)
[BTke

%

when [B| >>> [A]

A graphical representation of this relationship is given in Figure 10
in which t, is plotted as a function of concentration for specific
rate constants,

The application of kinetics in terms of the reactivity or transforma-
tion of a pollutant requires an understanding of the reaction process
and reaction conditions, This is particularly true in extrapolation
of laboratory to the aquatic environment.

Acidolysis of Dimethyl- and Diphenylmercury

Cleavage of the carbon-mercury bond by protic acid in dialkyl- or
diarylmercury compounds is referred to as acidolysis (eq 26).°%?
This reaction has been proposed by several investigators? °* 72 tbo
be a pathway for the chemical transformation of dimethylmercury to
methylmercury derivatives (Figure 1) in the aquatic environment.

H20
CHjHgCH, + Ho +X -— CHjHg” + X + Ch (26)

X= Cl, Bro, I”, Clo,~, No.”

However, as discussed earlier (Section IV), the rate constant for
acidolysis of dialkyl- or diarylmercury compounds in water has not
been determined. Even with compounds for which the kinetics have
been determined in organic solvents, the reaction mechanisms are
not clearly defined,

Preliminary work was done to evaluate a gic method for following
the acidolysis kinetics, Although the organomercuric salts were

32


100

80

60

20

1 J a r — ——

300 320 340 360 380 400 420
WAVELENGTH, nm

Figure 16, UV absorption spectra of methylmercuric
sulfide ion (A) and bis(methylmercuric)
sulfide (B).

57


In solutions containing the chemical species at concentrations
indicated in Figure 6, but no reduced sulfur, nearly all the methyl-
mercury would be complexed by organic thiols, e.g., sulfhydryl-
containing proteins, in the pH 5-9 range (Figure 8). Thus, when
either reduced sulfur species or organic thiols are present in
natural waters, methylmercury is quantitatively complexed by
sulfur-bonded ligands.

Exclusion of reduced sulfur species and organic thiols from the
system permits complexation by more weakly binding chemical species.
The plot of calculations shown in Figure 9 indicates that in the
absence of sulfides and thiols, methylmercuric hydroxide is the
major complex in basic waters (pH 7-10) and methylmercuric chloride
predominates in acidic waters.

One noteworthy aspect of the calculations is their prediction that
commonly occurring chemical species such as orthophosphate, ammonia,
phenolic groups in humic acid, and amino groups in protein have
relatively little impact upon complexation of methylmercuric ion
under environmental conditions, Moreover, the calculations indicate
that methylmercury can be freed from sulfur-bonded complexes by
strong acidification. Experiments have also shown that methylmercury
can be freed from sediments or biological samples by treating the
samples with strong acid.°*

Equilibrium calculations are strictly applicable only to closed
systems, and equilibrium is closely approximated in natural waters
only if the rate of approach to equilibrium is more rapid than the
rate of change of environmental conditions.’° Kinetic studies of
ligand exchange reactions discussed previously have shown that
achievement of equilibrium is very rapid for such reactions. More-
over, the Lake Erie Time Study of Kramer and coworkers indicated
that the rate of change in concentration of the chemical species in
Lake Erie considered in the above calculations is relatively slow
even during periods of high biological activity.°® Thus, near
approach to the calculated equilibrium concentrations of methyl-
mercury complexes is probable in local regions of lakes and slower
moving rivers and streams.

Although the literature does not contain a great deal of information
concerning equilibrium constants for phenylmercury complex formation,?®> 7
sufficient data are available to indicate that relative stabilities

of phenylmercury complexes parallel those of methylmercury complexes.
KINETICS OF ORGANOMERCURY REACTIONS IN WATER

Theoretical Considerations

The reaction mechanism, the rate expression, and the rate constant
are required to evaluate the significance of a reaction in the

28


laboratory occurred over a period of a few days on the desk top. At
high concentrations (0.10 M) with excess sulfide present, the complex
was degraded by a dark reaction to form mercuric sulfide. This
reaction was not investigated in detail because it proceeded at a
much slower rate at concentrations < 0.01 M. Assuming second-order
kinetics for the reaction, its rate constant was estimated to be

~ 10° t/mole sec.

Photosensitization of the methylmercury-thioglycolic acid complex
(CHgHgSCH2C0,H) by acetone proceeded with about the same quantum
yield (0,13) measured for the direct photolysis (0.15). Presumably,
the other thiol complexes would undergo such sensitized photodecompo-
sition also, since they all have the same structure about the mercury-
sulfur chromophore. On the other hand, humic acid, a substance likely
to be found in natural waters, did not photosensitize the decomposi-
tion of CHgHgSCH2COjH.

Quantum yields for photodecomposition of the methylmercury-thiol
complexes were lowered by addition of quenchers (Table 12). Stern-
Volmer plots (see eq 44) of the data were non-linear (Figure 18),
indicating that two or more excited states are involved in the
photodecomposition of the complexes.°* Whatever the nature of the
quenching processes, concentrations of quenchers such as oxygen are
sufficiently high under certain environmental conditions to lower
quantum yields for photodecomposition of the complexes. For example,
the concentration of oxygen in air-saturated water (3 x 10°* M)
would lower the quantum yield for CH,;HgSCH2,CH20H from 0.14 to 0.10.

Maximum rates of sunlight photodecomposition were calculated as
described previously (Table 13) using quantum yields from Table 11
and spectral data from Figures 16 and 17. The rapid photolysis rate
of the sulfide complex is due to the combination of its high quantum
yield and large sunlight absorption rate constant, Since methyl-
mercury ion is often complexed with sulfur-containing ligands in
the environment, these data suggest that sunlight photodecomposition
of sulfur-bonded methylmercury complexes plays an important role in
the conversion of methylmercury to inorganic mercury compounds,

Photocleavage of Phenylmercury Compounds

Spectroscopic studies showed that phenylmercuric hydroxide,
phenylmercury ion, and diphenylmercury (Figure 19) absorb at wave-
lengths > 290 nm, The absorption in this spectral region is due to
singlet-triplet electronic transitions that have enhanced intensity
because of perturbation by the heavy atom, mercury.°* °* ‘The
singlet-triplet spectra of phenylmercuric ion and phenylmercuric
hydroxide are identical, Thus, although changes in pH affect the
composition of phenylmercuric species in aqueous solution (eq 47),
the rate of sunlight absorption is pH-independent.

60


Table B5. EQUILIBRIUM CONSTANTS FOR HYDROLYSIS OF MERCURIC,
ALKYLMERCURIC, AND PHENYLMERCURIC IONS AT 25°C

Reaction pk (log Kly Reference
He’* +H,0 = Hgont +H 3.70 + 0.07 14
He0H* + H,0 2 Hg(OH)2-+ H* 2.60 + 0.09 14
Hg** + 2H,0 = Hg(OH)o + 2H" 6.30 + 0.05 14
CHjHg* + H20 2 CHgHg0H + Ht 4.50° 18
CoHgHg* + H,0 2 C,HsHg0H + Ht 4.90 16a
CgHsHe* + H20 + CgH,Hg0H + H” 4.12 16

*Tonic strength was 0.5 for all constants listed.
’Standard deviations not given.

97

U.S, GOVERNMENT PRINTING OFFICE: 1973 54/-312/228 1-3.


In the case of protic acid cleavage of DMM in water, where the strong
mineral acids are completely dissociated, the reaction probably takes
place by a pure Sp 2 mechanism (eq 28).

R

4
ReHg-R ReHg R R
+ =f > | + 4 (28)
a i i He

Because a myriad of chemical species exist in the aquatic environment,
several reactants that might be expected to affect the mechanism were
investigated to determine their effect on the rate constant (Table 4).
None of the salts tested appreciably affected the rate constant.

The small increase noted in some cases may be attributed to a salt
effect, Iodide ion, a better nucleophile, complicated the reaction
because CH,HgI precipitated, However, based on the first 20% of the
reaction, a rate constant of 7.6 x 10°? 2/mole sec at 85° indicated
no apparent rate alteration.

In the presence of S~, the reaction could not be followed because
NazS addition resulted in interfering gle peaks. However, gle
analysis 4id reveal that DMM was stable to Na,S (10°* M) at 85° for
24 hours in water,

Since proteins are common to the aquatic environment, an experiment
was designed to check for any effect they might have on acidolysis.
Cysteine, a thiol-containing amino acid was chosen as a representa-
tive compound. The rate constant for the acidolysis of DMM by cysteine
hydrochloride is given in Table 4, The kinetics were complicated by
the inability to accurately determine the actual H’ concentration.

The acid concentration was obtained by measuring the pH at 65°

(pH = 2.32). The rate constants determined were about 50% lower

than those obtained with HC1; the possibility of a large rate
alteration by proteins was therefore ruled out.

The rate constant for DPM acidolysis by HC10, was determined in both
aqueous ethanol and in water (Table 5), Aqueous ethanol was used
because the solubility of DPM in water (o-® M) was too low for
convenient determination of the rate constant under second-order
conditions, Product studies indicated that the reaction proceeded
according to equation 29,

H20
ColigHeCgH + HY + C10,” + CoHoHe” + C10,7 + CsHs (29)

The reaction obeyed second-order kinetics through 50% to 75%
completion in both 30% and 40% ethanol-water solutions.

38


APPENDIX A

METHYLMERCURIC ION EQUILIBRIUM CALCULATIONS

The total concentration of methylmercury species present in a system
was represented as the sum of the concentrations. of all complexes
incorporating the methylmercury moiety, plus the concentration of
the free ion itself, The formation constant expressions for the
various complexes considered were solved for the concentration of
the complex. Substitution of the concentration expression so
derived into the methylmercury concentration equation yielded an
equation for Z[CH,Hg] as a function of any one methylmercury complex.
Dividing both sides by the concentration of the complex and rearrang-
ing terms yields an expression for the fraction of the total methyl-
mercury that exists as a given complex. Because the formation
constant for bis(methylmercury) sulfide contains a squared term, the
resulting expression is second-order and requires an iterative
technique for solution.

An example has been worked out in detail for the complexing agents,
OH, Cl, S=, thiol (RSH), amino nitrogen (RNHg), ammonia, humic
acids (Ar0H), and orthophosphate. The formation constants (Appendix
B) indicate that few other complexes are likely to be important in
the environment, so other complexing agents were excluded from the
sample calculation.

The total concentration of methylmercury in the system is expressed
as:

E[CHsHg| = [CHsHgS” ] + [CHsHgOH] + 2[ (CHyHg) 2S]
+ [CHeHgC1] + [CHgHgNH,R"] + [CHgHgSR] + [CHsHgHPO,” |

+ (CHoHgNHs*] + [CHeHgoAr] + [CHaHe"] (58)
The following formation equilibrium expressions and the constants
(Appendix B) were then solved for the concentration of the respective
complexes,

Ky
CHjHg’ +S" 2 CH,HgSTM

(cHaHes” ] 2
%. = Touue’) (sj = 10°*

80


SECTION I

CONCLUSIONS

1. The composition of dissolved methylmercuric species in aquatic
systems is dependent on the type and concentration of complexing
agents present and the pH, An analytical expression is derived that
gives the relative percent of the methylmercuric species as a function
of concentration and pH in the presence of eight complexing agents.

2, At pH's and concentrations expected in the environment, methy1-
mercuric ion is not the predominant methylmercuric species, Because
natural systems contain much organic and inorganic sulfur, methyl-
mercury will exist predominantly as the sulfide or thiol complex. In
the absence of sulfide or thiols, the hydroxide or chloride complex
would predominate,

3. Reaction of hydrogen ion with dimethylmercury and diphenylmercury
gives methane and benzene, respectively, and the corresponding organo-
mercuric salt. The reaction is first-order with respect to both
hydrogen ion and organomercury concentrations. The second-order rate
constant extrapolated to 25° for dimethylmercury is 7.33 X 107° 2/mole
see and 9.67 x 10°° 4/mole sec for diphenylmercury, With the known
rate constants and kinetic expression, the acidolysis half-lives for
these compounds can be calculated at various pH's.

4. Dimethylmercury and diphenylmercury react with mercuric salts in
aqueous solution to give methylmercuric and phenylmercuric salts,
respectively. The reaction rate shows a strong pH dependence, increas-
ing dramatically as the pH is decreased. At alkaline pH's and mercuric
species concentrations common to the aquatic environment, this would
not be a significant degradative pathway. At acidic pH's the observed
rate constant is higher and the reaction may be significant, The half-
lives can be calculated from mercuric salt concentrations and pH data,

5. Formation of dimethylmercury by symmetrization reactions of methyl-
mercuric ion or methylmercuric hydroxide occurs at a rate too slow to
be significant under environmental conditions.

6. Dimethylmercury in aqueous solution does not react with dissolved
oxygen, hydroxide, sulfide, iodide, or albumin at a rate fast enough
to be a significant pathway for degradation.

7. Evaporative loss of dimethylmercury and elemental mercury from
aqueous solution may be significant for turbulent systems, Half-lives
for evaporative loss can be estimated from literature data and


ABSTRACT ,

LIST OF FIGURES

CONTENTS

LIST OF TABLES... . . ee ee ee oe
ACKNOWLEDGMENTS . . oe wee . .
Sections

L CONCLUSIONS , , 2... 1 ee eee

IL RECOMMENDATIONS 2. 2 6 eee ee eee
TIL INTRODUCTION. 2. we ee ee ee ee ee
Iv BACKGROUND. . 2. 2 eee eee ee ee

EQUILIBRIA OF MERCURY REACTIONS IN WATER ,
CHEMICAL DEGRADATION OF ORGANOMERCURIALS .

PHOTODECOMPOSITION OF ORGANOMERCURIALS . .

MATERIALS AND METHODS ..........
MATERIALS 66 6 ee ee ee ee ee
ANALYTICAL PROCEDURES .... 2... eee
APPARATUS 2. 6 6 ee ee et ee we we

ACIDOLYSIS OF DIMETHYLMERCURY, METHOD 1 .
ACIDOLYSIS OF DIMETHYLMERCURY, METHOD 2 ,
MERCURIC SALT CLEAVAGE OF DIMETHYLMERCURY
ACIDOLYSIS OF DIPHENYLMERCURY .....

MERCURIC SALT CLEAVAGE OF DIPHENYLMERCURY

iii

vii

ix

13

16

18

18

18

19

21

21

21

22

22


40 b of
cas 85°
Kx

t
eae
a|°
Cc
Sb
65°
e)
8 = mm
J
4 it 1 l l l it

40 80 120 160 200 240 280
TIME, minutes

Figure 11. Acidolysis of dimethylmercury.

35


Previous studies by Russell and Bridger have shown that reaction 57
occurs much more rapidly than reaction 56 in cyclohexane and carbon
tetrachloride,°® Products that resulted from reactions of the phenyl.
peroxy radical were not identified because of their complexity.

Disappearance quantum yields for photodecomposition of the phenyl-
mercury compounds with 313 nm light are summarized in Table 14.
Phenylmercuric perchlorate was used in the study because it is
completely dissociated in water at the high concentrations used in
the experiments, At pH 2.3 and pH 10.3, the phenylmercuric species
existed as phenylmercuric ion and phenylmercuric hydroxide, respec-
tively (eq 47). However, results in Table 14 show that the quantum
yield for photodegradation of phenylmercuric salts is the same at
both pH's. Because the broad-band (> 290 nm) quantum yield for
disappearance of phenylmercuric perchlorate (0.23 + 0.03) was about
the same as the 313-nm quantum yield, the quantum yield was assumed
to be wavelength-independent, Benzene (0.80 moles per mole of
phenylmercury compound decomposed) and inorganic mercury containing
precipitates were formed at both pH values in the acetone-sensitized
studies.

Photodecomposition of phenylmercurials was sensitized efficiently by
acetone (triplet energy 80 kcal/mole), but not by sensitizers with
triplet energies < 74 kcal/mole (Table 15). The triplet-state
energy of phenylmercurials must therefore be between 74 and 80 kcal/
mole, because efficient photosensitization occurs only when the
sensitizer triplet energy is equal to or greater than that of the
photoreactive compound.*° Ultraviolet absorption spectra of the
phenylmercury compounds (Figure 19) show that their triplet energies
are ~ 80 kcal/mole. The similarity of quantum yields in the direct
and sensitized photolyses is not surprising since both involve
excitation of the phenylmercury compounds to their triplet states,

Photosensitized conversion of phenylmercury to inorganic mercury
compounds is unlikely to be generally important in the aquatic
environment for the following reasons:

® Most compounds with triplet energies = 80 kcal moleTM+ do not
absorb strongly at wavelengths > 290 nm,

e Observed concentrations of mercury compounds in natural waters
are generally very low (< 10°° M). Because much higher concentra-
tions of competing energy acceptors, such as oxygen, are usually
present in the environment, photosensitization would be negligible.

Stern-Volmer plots for the quenching of diphenylmercury photo-
decomposition were linear (Figure 20). In acetonitrile, the slope
of the plot, k, +, was only 0.20, Assuming a value of 1.0 x 10°° 4/
mole sec for k, in acetonitrile,°” the indicated triplet lifetime,

66


46.

47.

48.

49,

50.

Sl.

52,

53.

54,

55.

56.

57.

58.

59.
60.

61.

62.

63.

0. P. Strausz, T. DoMinh, and J. Font, J. Amer. Chem. Soc., 90,
1930 (1968).

H. Takehara, M. Kotakemori, and T. Kajimura, Nippon Nogei Kagaku
Kaishi, 39, 448 (1965); Chem. Abstr., 64, 10342g (1966).

H. Shiina, R, Nishiyama, M. Ichihashi, and K. Fiyikama, Nippon

Nogei Kagaku Kaishi, 38, 481 (1964); Chem, Abstr., 63, 10594h
(1965).

G. Gran, Svensk Paperstidn, 53, 234 (1950).

A. Eldridge and T, R, Sweet, Anal, Chem., 28, 1268 (1956).

F. Kaufman and A, H. Corwin, J. Amer. Chem, Soc., 77, 6280 (1955).
W. R, Hatch and W. L. Ott, Anal, Chem., 40, 2085 (1968).

J. E, Longbottom, R. C. Dressman, and J, J, Lichtenberg, private
communication.

P. J. Wagner, P. A. Kelso, and R. G. Zepp, J. Amer. Chem. Soc.,
94, 7480 (1972).

(a) W. H. Thompson et al., Trans, Faraday Soc., 32, 681 (1936);
(b) G, Davis, Dissertation, Kansas State University, 1971; (c)

R. N. McDonald and G, E, Davis, J. Org. Chem., 38, 138 (1973).

F. G. Moses, R. S. H. Liu, and B, M. Monroe, Mol, Photochem., 1,
245 (1969).

R. Livingston in "Technique of Organic Chemistry,” Vol. VIII,
Part I, A. Weissberger, Ed., Interscience Publishers, Inc., New
York, 1961, Chapter V.

P. J, Wagner and A. E, Kemppainen, J. Amer. Chem. Soc., 94,
7495 (1972).

A, A, Lamola and G, S. Hammond, J. Chem. Phys., 41, 2129 (1965).
C. S. Foote, Accounts Chem. Res., 1, 104 (1968).

R. C, Weast, Ed., “Handbook of Chemistry and Physics,” 45th Ed.,
The Chemical Rubber Co., Cleveland, Ohio, 1964.

Pp. A, Leighton, "Photochemistry of Air Pollution," Academic
Press, New York, N. Y¥., 1961, pp. 26-30.

E. R. Allen, R. D. McQuigg, and R. D. Cadle, Chemosphere, 1, 25
(1972).

75


660/3-73-012
1973

Ecological Research Series



ABSTRACT

Kinetics in water of some chemical and photochemical reactions
postulated as key transformations in the environmental mercury cycle
were investigated. Decomposition of dimethylmercury (DMM) and
diphenylmercury (DPM) by acids and mercuric salts was shown to be
pH dependent and too slow to be significant under most environmental
conditions. Degradation of organomercuric salts by acid is even
slower, Theoretical evidence indicates that loss of elemental

mercury or DMM at the air-water interface can be important in turbu-
lent systems.

Dimethylmercury, methylmercuric chloride, methylmercuric hydroxide,
and methylmercuric ion were not decomposed by sunlight, but phenyl-
mercury and sulfur-bonded methylmercuric species were readily
decomposed to inorganic mercury. Detailed equilibrium calculations
indicate that the sulfur-bonded methylmercuric species are the
predominant species in natural waters. Quantum yields for these
reactions are presented along with a technique for calculating
sunlight photolysis rates from laboratory data.

The report also includes a review of the chemical literature con-
cerning the kinetics of chemical and photochemical decomposition of
organomercurials.


PHOTOLYSIS OF METHYLMERCURIC SULFIDE ION

Aqueous solutions of methyimercuric sulfide ion (0,010 M) were pre-
pared by reacting methylmercuric hydroxide with a two-fold molar
excess of sodium sulfide, Photolysis by > 290 nm light caused
formation of gas and a black precipitate. The gas was characterized
by gic and the precipitate was identified by X-ray diffraction.

Yields of the mercuric sulfide precipitate were determined as

follows: (1) The photolyzed solutions were diluted 1:100. Mercuric
sulfide is soluble in concentrated sodium sulfide and therefore did
not completely precipitate prior to dilution.®* The diluted solutions
were allowed to stand in the dark for 12 hours to ensure complete
precipitation. (2) The supernatant liquid was partially decanted

and the remaining mixture was centrifuged. (3) After centrifugation,
the precipitate was washed with water, recentrifuged, and dissolved
in aqua regia. (4) Mercuric ion concentration was determined as
usual by atomic absorption spectrophotometry.

SUNLIGHT PHOTOLYSIS OF METHYLMERCURY COMPOUNDS

Aqueous solutions (1.00 x 10°* M) of methylmercuric hydroxide,
methylmercuric chloride, methylmercuric bromide, and methylmercuric
iodide were degassed as usual in 10.0 + 0.1 mm quartz tubes, The
tubes were sealed under vacuum and irradiated by September sunlight
between 9 AM and 3 PM on the roof of the Southeast Environmental
Research Laboratory (SERL), Athens, Georgia. Methylmercuric iodide
was irradiated for 3.7 hours and the other methylmercury compounds
were irradiated for 17.1 hours,

ACETONE-SENSITIZED PHOTODECOMPOSITION OF PHENYLMERCURIC ION AND
PHENYIMERCURILC HYDROXIDE

Degassed aqueous solutions of phenyimercuric perchlorate (1.00 x 107° M)
and acetone (0,60 M) adjusted to pH 2.3 and pH 10,2 were irradiated
by Pyrex-filtered light. Benzene was determined by gle and precipi-
tates were analyzed for mercury content by flameless AA spectropho-
tometry. Nearly quantitative yields of inorganic mercury precipitated
under basic conditions (0.95 mole Hg per mole CgH,Hg0H decomposed)

and lower yields (0.55 mole Hg per mole CgH;Hg* decomposed) precipi-
tated under acidic conditions, Analysis of the supernatant of the
acidic solution showed that it contained an additional 0.4 mole

Hg per mole Ce Hs He” decomposed, presumably in the form of mercurous
and/or mercuric ions.

QUENCHING STUDIES OF PHENYLMERCURY COMPOUNDS

Solutions containing the phenylmercury compound and various con-
centrations of quencher (0-4 M) were prepared, degassed (unless

23


Ke
CHjHg’ + NH; = CH,HgNH,*

+
ec ee
° ~ (cHgHe* | INH 1
Kg
CHgHg’ + ArO” = CH3HgOAr
(cH, Hg0Ar }
Ke = 10°

~ TorgHe® ] [aro]

Substitution of the above in equation 58 and division of both sides
by [CH3HgS” ] gives the following equation:

E[CHgHe | _ 14 ke Lor] , 2K. [cogs
(cH, HgSTM J x, [s*] Ky (s=]

ak {cer 4k LRNH2 | 4 Ke {sr7] a kr (HPoz7]
kK, fs] x, [s7] x, [s*) & [8]

Ko [We | Ke [Aro] 1

+— st = = (59)
Kk, [s*] x, [s=) RK [877
Simplifying and taking the reciprocal of both sides gives the
fraction of methylmercury existing as CH,HgS.
(CHgHgsTM 1]
© [cusug]
K,(s"]
K,[S"] + Ke[OW ] + 2k, (CHgHgS”] + K& [CI] + K,[RNH2] + Kg[SR]
+ Ky HPO?” ] + Ke (NH, ] + Ko [Aro] +1 (60)

The concentrations of [STM], [RW], [RS], [Hpog7], [nwo], and [Aro-]
are difficult to measure and are also pH-dependent; they must there-
fore be defined in terms of more easily measurable variables.
Equations 67-72 redefine the variables in terms of quantities that
can be more readily estimated.

82


80.

81.

82.

83.
84.

85.

86.

87.

88,

89,

90,

91.

92,

93,

94,

95.

96.

97.

J. G. Calvert and J. N. Pitts, Jr., "Photochemistry," John Wiley
& Sons, Inc., New York, N. Y., 1966, p. 742.

Reference 62, Chapters II and III.

E. R. Allen, R. D. McQuigg, and R. D, Cadle, Chemosphere, 1, 25
(1972).

S. Q. Duntley, J. Opt. Soc, Amer., 53, 214 (1963).

A. A, Lamola, Photochem, Photobiol., 8, 601 (1968).

N. J, Tarro, "Molecular Photochemistry," W. A. Benjamin, Inc.,
New York, 1967, Chapter 5.

(a) T. Frankiewicz and R. S, Berry, Environ. Sci, Technol., 6,
365 (1972); (b) J. N. Pitts, Jr., in "Chemical Reactions in
Urban Atmospheres," L, S$, Tuesday, Ed., American Elsevier
Publishing Co,, Inc., New York, N. Y., 1971, pp. 3-31.

J. F. Uthe, J. Solomon, and B. Grift, J. Assoc, Offic. Anal.
Chem,, 55, 583 (1972).

C, Marcker, Ann., 136, 81 (1865).

R. J. Kern, J. Amer. Chem. Soc., 75, 1865 (1953).

W. L. Hughes, Ann. N. Y. Acad. Sci., 65, 454 (1957).

B. G. Gowenlock and J. Trotman, J. Chem, Soc., 1454 (1955).
Pp. J. Wagner in "Creation and Detection of the Excited State,"
Vol. 1, Part A, A. Lamola, Ed,, Marcel Dekker, New York, 1971,
Chapter 4,

A, P, Marchetti and D. R. Kearns, J. Amer, Chem, Soc., 89, 768
(1967).

W. J, Turro, "Molecular Photochemistry,” W, A. Benjamin, Inc.,
New York, 1967, pp. 48-52.

G. A, Russell and R. F. Bridger, J. Ar
(1963).

Chem, Soc., 85, 5765

K. Sandros and H. L. J. Backstrom, Acta Chem. Scand., 16, 958
(1962).

P. J. Wagner and D, J. Bucheck, J. Amer, Chem, Soc., 91, 5090
(1969).

77


SECTION IIL

INTRODUCTION

Although the occupational hazards of mercury have been known for
many years, its environmental impact only recently became apparent.
The recent interest in the environmental implications of mercury
stems primarily from two incidents in widely separated countries.
First, a massive case of poisoning in Japan--the Minamata Bay
incident --was attributed to industrial discharge of organomercurials.*
Secondly, declining bird populations in Sweden was blamed on the use
of phenyl and methylmercurial pesticides as seed dressings.” Pio-
neering work, particularly in Sweden, has since shown that mercury
accumulates in fish to concentrations much higher than that in
surrounding waters and that it exists in fish predominantly in a
methylated form.” The widespread occurrence of methylmercury
compounds remained a mystery until it was shown by Jernelév that
inorganic forms could be biologically methylated in aquatic systems.*
No direct proof exists, however, that the mercury in fish is
methylated prior to uptake.°® Several extensive reviews of the
overall mercury problem are available and should be consulted for
more detail,°>®

Tentative descriptive models similar to Figure 1 were offered in
explanation of the transformations and cycling of mercury in
aquatic systems.* The phenylmercurials were included primarily
because of their widespread release into waterways as a result of
their use as fungicides.

While prior research has clearly demonstrated the importance of
organomercurials, few data are available concerning the mechanisms
or rates of their reactions, And although some work has been done
on the non-aqueous chemistry of organomercurials, very few studies
were carried out in water or, in the case of photochemical reactions,
in light of wavelengths characteristic of solar radiation. The lack
of relevant information may partially explain the strong reliance
of environmental chemists on thermodynamics to predict the aquatic
chemistry of mercury and mercury compounds. The environmental
literature, as a result, contains many erroneous conclusions about
the forms and reactions of organomercurials as pollutants.

The present report includes an extensive background section with a
discussion of pertinent equilibria and a review of the chemical and
photochemical degradation reactions of mercury and mercury compounds,
Because the chemical and photochemical reactions occur by different
mechanisms and require different experimental techniques, they are
discussed separately,


t i a al i

10 20 30 40 50
TIME, minutes

Figure 13, Desymmetrization of dimethylmercury
by mercuric perchlorate.

AT

60


CONDITIONS:
LS =10°M

[CH3HgS~]
> [CHsHg]

ERS = 10M

ERS = 108M

ERS = 107M

ERS = 10°°M

Figure A4, Relative concentrations of CH3HgS in systems
containing a low concentration of reduced

sulfur species.

89


89

Table 15, QUANTUM YIELDS FOR PHOTOSENSITIZATION OF PHENYLMERCURIALS

Triplet

Energy Disappearance
Compound Sensitizer (keal moleTM*) Solvent Quantum Yield
(Cg Hs )oHg Acetone 80 Benzene 0.33
(CeHs Jog Acetophenone 74 Benzene < 0.003
(CgHs )eHg Benzophenone 68 Benzene < 0.006
(Cg Hs Jog Naphthalene 62 Benzene 0.027
(Cg Hs Jog Anthracene 42 Benzene < 0.003
Cg Hg Hg02 CCHg Acetone 80 Water, pH 4.0 0.23
CgHsHgC10, Acetone 80 Water, pH 2.3 0.23



PERCENT [CH3HQ] ror

r —

? (CHsHgC1] 10 [CH3HgSR]
100 [CH3Hg}] ,o, [CH3Hg] ror
go b+ SPECIES CONC. M
i (CHsHg],,, <0"!
60 [H2PO47] + [HPO.—] 10°
(CI7] 1074
— (NH4*] + [NH3] 10°?
40 + [RNH3*] + [RNH, ] 10°
_ [RSH] + [RS7] 10°°
oo (ArOH] + [Ar07} 7 10°?
[H2S] + [SH>] + [S=] 0
g _: _! | | l | j j
0 2 4 6 8 10 12

Figure 8. Relative concentrations of CHgHgCi and CH,HgSR in
systems containing multiple chemical species,
excluding reduced sulfur

29


When the pseudo-first-order rate constant was determined in water
using a large excess of HC1Q,, the reaction obeyed first-order
kinetics through one half-life (see Figure 12),’

The rate constant increased as the ethanol concentration was reduced
from 30% to 0%. Although the rate increase was small for such a
large solvent change, qualitatively it agrees with that expected
from the increase in acidity with decreased ethanol concentration.”°
The small decrease in rate constant in going from 40% to 30% ethanol
was probably due to solvation effects that could not be evaluated
without further study.

Extrapolation of the rate data obtained at higher temperatures to 25°
gives a constant of 9.67 x 10°° t/mole sec for water.

The acidolysis half-lives for DMM and DPM can be calculated at 25°
from these data (Tables 2 and 5). Assuming that the acid concentra-
tion remains constant because of buffering, pseudo-first-order
conditions would prevail (see eq 25 where [B) = [H*]). At pH's
commonly found in the environment, DMM would have a relatively long
half-life. For example, at pH 5, tz would be 33 years at 25°, For
DPM under the same reaction conditions, it would be 0.25 years.
Thus, acidolysis may be important for DPM under certain conditions,
but for DMM it would be important only at low pH's.

The literature contains limited quantitative data on the relative
rates of acidolysis for dialkyl- and diarylmercurials. However, a
qualitative order of reactivity is given by Kharasch and Grafflin’®
(Table 6). They did not measure actual rates, but rather determined
relative rates using unsymmetrical organomercurials in competition
studies. They also demonstrated that the relative order of reactivity
did not change for a variety of acids and solvent systems. Our data
for DMM and DPM are in agreement with this order.

Since the phenyl group is the most reactive moiety listed in Table 6
and acidolysis of DPM is low under environmental conditions, other
organomercurials containing the alkyl groups listed would react even
slower.

Desymmetrization of Dimethyl- and Diphenylmercur,

Carbon-mercury bond cleavage of dialkyl- or diarylmercury compounds
by mercuric salts is termed desymmetrization’**” (eq 30). The
equilibrium for this reaction generally lies far to the right as
shown in Table 7.

desymmetrization
ReHg + Hexe = 2RHex (30)
symmetrization

41


9”

Table 9, pH DEPENDENCE FOR DESYMMETRIZATION OF DIPHENYLMERCURY BY MERCURIC

PERCHLORATE AT 27°

Conc. Hg (C10, )2 Conc. DPM k
pH M M 2/mole sec
2.0 2.0 x 1072° 1.0 x 107° >2x 10%
4.8 5.0 x 10°? 4.1 x 107° 3 x 10°"
5.8 7.0 x 10°? 4.5 x 107° 3 x 108°
6.9 5.2 x love 4.6 x 10° 4 x 10°
8.1 6.57 x 107° 4.56 x 10°° 3 x 10°°
9.4 1.06 x 10-7 2.6 x 10° 3 x 10%

“Using atomic absorption,

* oH determined at the start and end of the reaction.

“Lower Limit.

“Calculated from three single point determinations,

°Calculated by computer program.


(Ag — Aco)
(A — Koo)

Ln

0.5

0.4

0.3

0.2

0.1

4

6

10

TIME, sec x 107?

Figure 12, Acidolysis of diphenylmercury.

42


TL

Table 16, COMPARISON OF CALCULATED AND EMPIRICAL RATES FOR
PHOTODEGRADATION OF PHENYLMERCURY COMPOUNDS IN SUNLIGHT

Calculated “va Empirical Empirical ty
Compound 10° (ea )ints @» bre? 10" (ke inte O° hrs
(Cols )2 He 8.14 1.7 8.5 41.8
CgHsHgOCocH, ‘ 3.3 4,340.3 16.0 + 2
CgHs HgNOg 3.3 3.5 + 0.2 20.041
Ce Hs HeBO, 3.3 5.0 4+ 0.3 4.042
Cols Hg OH 3.3 4,340.3 16.0 +2

“Calculated from solar radiation data, extinction coefficients, and quantum yields in Table 14,
* average of three replicates irradiated by sunlight.
°Gaiculated assuming first-order kinetics for photodecomposition,


Ke
CH,Hg* + OH = CH,HgOH

= (CH, HgoH] 2.37
“o * Taigua® 1 Cor] = 7°

Ke
CHgHg” + CHysHgS” = (CHaHg)2S

( (Cig Hg) 23]

= 1016?
(Cig He” | [CHsHgs” }

Ky =

K
CHgHg’ + Cl” = CHgHgCl

(cH, Hgc1}

& = Toagne’} tery = 1"

Ks
CHgHe” + RNH, = RNH,HgCH,”

(RNB HgCH,* J

_ 7 p8.28
Ks = Tanz) CougneTM) ~ 2°
Ke

CHgHg’ + SR 2 CHgHgSR

_ (CHsHgSR} 102822

Due) ber

(CHgHe | (Sk)
Ky

CHgHg* + HPG;” 2 CHyHgHPQ,~

_ _ [CHa heliro,” } 10828
Ky = Tongue’ | [HEG:) ~

81


9€

Table 2, DIMETHYLMERCURY ACIDOLYSIS RATE DATA MEASURED BY GLC
Conc, Acid Conc, DMM Temp.
Acid M M °¢ k 2/mole sec k sec?
b
HCL 8.27 x 1o-* 3.46 x 10°° 65 (7.19 + 0.32) x 107°
HCL 8.27 x lo 5.43 x 10° 65 (7.43 + 0.19) x 10°°
b
HCL 8.27 x Lo-* 4,90 x 107° 85 (4.89 + 0,07) x 10°?
HCL 8.27 x 10° 4.60 x 107° 85 (4.87 + 0.39) x 10°?
e
HCL 40 5.08 x 10°*
°
HCL 25 7.33 x 107°
a
HCL 1.0 x 10+ 6.16 x 10°° 40 5.12 x 10° (5.12 + 0.12) x 107
HCL 1.9 x 1073 5.28 x 10% 40 4,87 x 10°* (9.16 + 0.09) x 107°
HBr 8.24 x 10°° 3.47 x 10°° 85 (4.46 + 0.12) x 107°
HBr 8.24 x 10° 6.82 x 10°? 85 (4,69 + 0.12) x 10%
HC1O, 8.20 x 10-* 4.79 x 107° 65 (8.06 + 0.33) x 10°°
HC10, 8.20 x 10°° 6.08 x 107° 85 (5.02 + 0.15) x 107-7

“See Section V,
’ The activation parameter based on the average of the
= -3,1 + 0.6 eu,

and AS#
°Extrapolated.

“Calculated from pseudo-first-order rate constant.

two values are AH+ = 22,1 + 0.1 kcal/mole


EPA-660/3-73-012
September 1973

CHEMISTRY OF ORGANOMERCURIALS IN AQUATIC SYSTEMS

by

George L. Baughman
John A. Gordon
N. Lee Wolfe
Richard G. Zepp
Southeast Environmental Research Laboratory
National Environmental Research Center-Corvallis
, Athens, Georgia 30601

Project 310301 QQc
Program Element 1BA023

NATIONAL ENVIRONMENTAL RESEARCH CENTER
OFFICE OF RESEARCH AND DEVELOPMENT
U. §. ENVIRONMENTAL PROTECTION AGENCY
CORVALLIS, OREGON 97330

For sale by the Superintendent of Documents, U.S. Government Printing Oflce, Washington, D.C, 20402 - Price $1.30


Table 11. QUANTUM YIELDS FOR PHOTODEGRADATION OF SULFUR~BONDED
METHYLMERCURY COMPLEXES AT 313 nm IN DISTILLED WATER

Complex S
CH3Hg-egg albumin 0.12 + 0.02
CHgHgSCH2COsH 0.15 + 0.01" °
CH, HgSCH, CH, OH 0.14 + 0,01 °
CH, HeSCH2 CH (NH,* )CO2H 0.16 + 0.02
CH HgSTM 0.65 + 0.05"

* Disappearance quantum yield in degassed solutions.
°Quantum yield shown to be concentration independent.
°Quantum yield shown to be pH independent.


sy

Table 8. pH DEPENDENCE FOR DESYMMETRIZATION OF DIMETHYLMERCURY AT 27°

Hig (C10, Jo DNM k
pit M M 2/mole sec
1,2 7x 10° 7 x lo® 5 x 10*

3.1 7x 107° 5 x 1o7® 2 x 10%
5.8 3x 1077 3x 10-7 1 x 10°°
9.0 2 x 10°8 1x 10% 2x 102°

“Using atomic absorption.

pH determined at start and end of reaction,

“Calculated by computer program,


Table

B2, LOGARITHMS OF FORMATION CONSTANTS FOR
PHENYLMERCURIC COMPLEXES (C,H,HgX) IN

H20, 25°C
(Ce Hs Hex |
Ligand X Log TesHsHe | Ox) Reference
- OH 9.89 16
-OCOCHs” 4.82 16
-OCOCHs CH,” 4,51 16
-SCg Hg > 16 7

94


Actual yields for sensitized photolysis of diphenylmercury in benzene
were determined by gle on a 3% OV-1 column. Yields for photodecompo-
sition of methylmercury-thiol complexes were measured by nuclear
magnetic resonance (nmr) spectroscopy and gas liquid chromatography.°°
Mercury yields in the precipitates resulting from photolysis of
phenylmercuric acetate and methylmercury-thiol complexes were
measured by digesting the precipitates with aqua regia (3:1 HNO, -HC1)
and determining the mercury by flameless atomic absorption spectros-
copy. Yields, of benzene from the photolysis of phenylmercury
compounds were measured by gle on a Porapak PS column. Product
yields in the chemical actinometer were determined by gas liquid
chromatography as described elsewhere,°*

APPARATUS

Kinetic studies of methylmercury and phenylmercury compounds were
conducted in a thermostated oil bath that regulated temperature
within + 0.05°C, One of the major problems in working with DMM was
its volatility (vapor pressure 50 mm Hg at 20.5°C) °°" which made it
difficult to accurately weigh and transfer small quantities. In
addition, DMM dissolved slowly (several hours) in water. Special
bombs were constructed for the studies of dimethylmercury chemical
reactions and photolysis (Figure 5). The closed bombs were almost
completely filled with solutions to minimize loss of gaseous dimethyl-
mercury. When the stopcock of the bomb was opened, aliquots of the
reaction solution were removed by inserting a syringe needle through
the septum. The aliquots were analyzed by glc as described above.
Control experiments showed that dimethylmercury did not volatilize
from the bomb at the elevated temperatures used in the kinetic runs.
Dimethylmercury acidolysis was also followed with procedures and
apparatus described by Davis and McDonald.°°?° Solutions for the
conductance studies were prepared and added to the cell under a
nitrogen atmosphere in a glove box.

Quantum yield studies were carried out on a photochemical apparatus
consisting of a rotating turntable assembly contained in a water bath.
A Hanovia 450-W medium pressure mercury lamp, positioned in the
center of the turntable, was employed as the light source, The
apparatus has been described in detail elsewhere.°° Samples for
photochemical studies were degassed by several freeze-thaw cycles
under vacuum.

Glc analyses were performed on a Tracor MT-220 gas chromatograph,
equipped with flame detectors and a Ni-63 electron capture detector,
Glc peaks were integrated by a Varian Model 477 Digital Integrator.
Mass spectra were obtained on a Hitachi-Perkin Elmer RMU-7 Mass
Spectrometer. Nmr spectra of the methylmercury-thiol complexes were
measured on a Varian HA~100 NMR Spectrometer. Flameless AA analyses
were carried out using a Perkin-Elmer 403 Atomic Absorption

19


Acidolysis reactions of non-ionic organomercury compounds (eq 11) and
organomercuric salts (eq 12) have been studied by several investiga-
tors.

ReHg + HX - RHgX + RH (il)

RHgX + HX + HgX, + RH (12)

Zimmer and Makower reported that HBr reacted with diphenylmercur

(R = CgHs, eq 11) in 5% water-methanol twice as fast as did HC1."*
Under these conditions, sulfuric, perchloric, nitric, acetic, and
trichloroacetic acids reportedly did not react. In aqueous tetra-
hydrofuran and dioxane solvents, second-order kinetics were obeyed
when HCl was the acid.*® Reutov and coworkers reported that
acidolysis of dibenzylmercury (R = CgH,CHz, eq 11) was first-order
in dibenzylmercury and first-order overall,°° but Jensen and Rickborn
have strongly criticized Reutov's experimental technique and inter-
pretation.’ Kinetic data for dimethylmercury acidolysis (R = CH,
eq 11) in water are not available in the literature. Dessy and
coworkers reported that the acidolysis of dimethylmercury by HCl in
dimethylsulfoxide-dioxane (10:1) was too slow to measure at 40°C,°°
This limited and sometimes confusing information emphasized the need
for kinetic studies of acidolysis in water by dissociated acids.

Cleavage of organomercuric salts in water by acids has been inves-
tigated, Kreevoy reported the second-order rate constant for
acidolysis_ of methylmercuric iodide (R = CHg, X =I), eq 12, by

1 M HeSQ ee Extrapolation of these data from higher temperatures
to 25°C gave the very low rate constant, 3 x 10°° t/mole sec. Under
Kreevoy's conditions the calculated half-life was 3 x 10° days at
25°C, Acidolysis of phenylmercuric chloride (R = Css, X = Cl),
eq 12, in water containing 10% ethyl alcohol was investigated by
Brown,*° Extrapolating Brown's results to 25° gave a second-order
rate constant of 1.7 x 10°° 2/mole sec. Comparison of Kreevoy's
and Brown's data indicates that methylmercuric salts undergo
acidolysis at much slower rates than phenylmercuric salts.

Another possible route for chemical degradation of organomercuric
salts is the so-called "demercuration reaction" shown below.

H20
RHgX — ROH + He? + xX (13)

Quellette** has shown that reaction 13 is kinetically first-order
and that the reaction rate is very rapid for branched aikylmercuric
salts, such as t-butylmercuric and cyclohexylmercuric halides.
However, demercuration of methylmercuric salts is very slow, Extra-
polation of Quellette's data for CH,HgX to 25° gives the first-order

is


SELECTED WATER jh Revere Wo] 2 " d. Accession No.
RESOURCES ABSTRACTS | '

INPUT TRANSACTION FORM ! Ww
4. Title 1 5. Report Date
CHEMISTRY OF ORGANOMERCURIALS IN AQUATIC SYSTEMS 1¢
1@. FP rtormy.g Orgai. zation
7. Author(s) Baughman, George L., Gordon, John A., Wolfe, N. Report No.
Lee, and Zepp, Richard G, 10. Project No.

_— 310301QQ6
9. Organization .
Southeast Environmental Research Laboratory 11. Contract/Grant No.
National Environmental Research Center-Corvallis
U.S. Environmental Protection gency ‘lr Type FRepo.and
Athens, Georgia 30601 pitied Covered.

12, S: nsorin Organ’ ion J. 8. Environmental Protection Agency — Final Report

15. Supplementary Notes

Environmental Protection Agency report number,

EPA-660/3-73-012, September 1973.
16. Abstract
Kinetics in water of some chemical and photochemical reactions postulated as key
transformations in the environmental mercury cycle were investigated. Decomposition
of dimethylmercury (DMM) and diphenylmercury (DPM) by acids and mercuric salts was
shown to be, pH dependent and too slow to be significant under most environmental con-
ditions, Degradation of organomercuric salts by acid is even slower, Theoretical
evidence indicates that loss of elemental mercury or DMM at the air-water interface can
be important in turbulent systems.
Dimethylmercury,methylmercuric chloride, methylmercuric hydroxide, and methylmercuric
ion were not decomposed by sunlight, but phenylmercury and sulfur-bonded methylmercuric|
species were readily decomposed to inorganic mercury, Detailed equilibrium calcula-
tions indicate that the sulfur-bonded methylmercuric species are the predominant species
in natural waters. Quantum yields for these reactions are presented along with a
technique for calculating sunlight photolysis rates from laboratory data,
The report also includes a review of the chemical literature concerning the kinetics
of chemical and photochemical decomposition of organomercurials. (Baughman -
Southeast Environmental Research Laboratory)

17a. Descriptors
*Heavy metals, *Hydrolysis, *Kinetics, Metal organic pesticides, Chemical degradation,
Aqueous solutions, Water chemistry, Evaporation, Air-water interfaces

17b, Identifiers

*Photolysis, *Organomercury, Photodegradation, Complex formation, Mercury,
Methylmercury, Phenylmercury, Diphenylmercury, Dimethylmercury

17. COWRR Field & Group 5B

18. Availability 19, Security Cless, | 21. No.of | send To:
“Repoi.) Pages
lon See rityeiss. | 2 Poe ATER aran aces peur ygzormaron cenren
1 (Page) WASHINGTON. D.C. 202
Abstractor George L. Baughman I Institution Southeast Environmental Research Laboratory}

WRSIC 102 (REY, JUNE 1971)


100
90
80
70
ml _ 0
2\¢
=)
Slz
42,
~
s
-” 49
SPECIES CONC, M
{CHyHo) <107'?
0 Fr (hiP0."] + [HPO] 10"
(cI-} 10"
2 L {NHa"] + [NH5} 10"?
[RNH;*) + (RNH2) 10°
(RSH) + [RST] 0°
10 {ArOH] + [ArO-] 10°?
(HeS] + [SH] + (S=] xs
/ —
1 1 n
1 2 3 4 5 6 7 8 9
pH
Figure 7, Relative concentrations of (CHgHg)25 in systems

containing multiple chemical species,

27


Quantitative calculations of sunlight photolysis rates can be made
from specific absorption rates (k,) and quantum yields (@) both
determined in the laboratory.® Concentrations of pollutants such
as organomercury compounds are generally so low that absorption at
wavelengths greater than 290 nm is very weak, For weakly absorbing
systems, rates of sunlight photodecomposition (-d{RHgX]/dt) can be
expressed as

- SEH) - 1, ofRnex] (35)

where the term [RHgX] represents concentration of organomercury
compound. This equation assumes that sunlight photodecompositions
follow first-order kinetics.

Specific absorption rates for the organomercurials were calculated
using solar radiation data published by Leighton®® and extinction
coefficients at wavelengths > 290 nm, Leighton's data are most
appropriate for Los Angeles during August through November, the
period we used for our experiments in sunlight. Since Athens,
Georgia, has about the same latitude and elevation as Los Angeles,
the spectral flux distribution of sunlight at Athens should closely
parallel that at Los Angeles. The value of k, changes during the
day because the solar spectrum is a function of the solar zenith
angle, z. For this reason, it is convenient to use an integrated
specific absorption rate constant, (k,):nt, for predictions of sun-
light photolysis rates ,°*

N
"

m/2
kez d2/t (36)
z= -7/2

(ka dinte =

In the above expression, k,, is the absorption rate constant at
solar zenith angle, z. The half-life (tx) for a photoreaction can
then be expressed as

ne 0.693 (7)
2 (ka ent od

Specific absorption rate constants in natural waters are also a
function of competitive light absorption by water and other sunlight:
absorbing species, light scattering, and other factors.°* In this
study, the primary purpose was to determine minimum half-lives for
photodegradation of organomercury compounds by sunlight, i.e.,
half-lives for sunlight photolysis near the surface of a body of
water. Using the results of this study, it is possible to predict

51


Dy/ fo

1.0 K~

L \ 1 1

0 ] 2 3 4
[Pentadiene], moles/|

Quenching of diphenylmercury photolysis by

Figure 20,
cis-1,3-pentadiene in acetonitrile.

69


Spectrophotometer, Ultraviolet spectra were measured by a Perkin-
Elmer 352 Ultraviolet Spectrophotometer and uv analyses were
carried out on Bausch and Lomb Spectronic 505 and Beckman DU Spec-
trophotometers. X-ray diffraction studies were carried out at J. F.
Kennedy Space Center, Florida,

ACIDOLYSIS OF DIMETHYLMERCURY, METHOD 1

A weighed portion of dimethylmercury was added to a volumetric
flask containing 0.0040 M tetrahydrofuran (the gle internal standard)
in water and a magnetic stirring bar, The resulting mixture was
stirred for 12 hours to dissolve the dimethylmercury, which was
found to have a solubility limit of 0.02 M in water at room tempera-
ture, An aliquot of standardized acid solution was added to the
aqueous dimethylmercury solution, and the resulting solution was
transferred by a syringe to the kinetic bomb (Figure 5), which was
then totally immersed in the thermostated oil bath. Periodically,
the bomb was removed from the bath and cooled by immersing in water
at room temperature. After a 1.0-microliter aliquot was removed
from the bomb and analyzed by gic, the bomb was returned to the
bath. Second-order rate constants for acidolysis were calculated
by computer using a least-square analysis of data.>”

ACIDOLYSIS OF DIMETHYLMERCURY, METHOD 2

Dimethylmercury was added to water in a volumetric flask containing
a magnetic stirring bar. The flask was placed inside the glove box
under nitrogen and the dimethylmercury was dissolved by stirring for
12 hours, An aliquot of standardized acid solution was added to the
dimethylmercury solution under nitrogen in the glove box, and the
resulting solution was transferred to the conductance cell. The cell
was immersed in the thermostated oil bath, and the conductance of
the solution was recorded at appropriate time intervals, Reactions
were carried out with a 100-fold excess of dimethylmercury, and
pseudo-first-order kinetics were observed, Pseudo-first-order rate
constants were calculated by computer employing a least-squares fit
of the data,°”

MERCURIC SALT CLEAVAGE OF DIMETHYLMERCURY

Aqueous solutions of mercuric perchlorate (10°° to 10°° M) were
prepared by the addition of perchloric acid to mercuric acetate
solutions, The pH was measured with a pH meter and the initial
mercuric salt concentration was measured by flameless AA spectro-
photometry. The mercuric perchlorate solution was then transferred
to a reaction bomb and the bomb was immersed in the oil bath
thermostated at 27.0°. An aliquot of aqueous dimethylmercury
solution was added and, at appropriate time intervals, aliquots of
the reaction solution were removed and the mercuric salt concentration

21


[CH3HgSR]
DY [CH3Hg]

ERS =10°°M

CONDITIONS: 2S =10°M

RS = 10°M

ERS = 10°M

0 2 4 6 8 10 12

Figure A2, Relative concentrations of methylmercuric-thiol

complexes in systems containing a low concentra-
tion of reduced sulfur species,

87


calculated transfer coefficients. Based on these calculations
elemental mercury is lost from solution about twice as fast as
dimethylmercury,.

8. The low sunlight absorption rate constants for dimethylmercury,
methylmercuric ion, and methylmercuric hydroxide preclude photodecom-
position as a significant pathway for degradation. Neither photo-
sensitization nor singlet oxygen effects their decomposition.

9. Diphenylmercury absorbs sunlight in aqueous solution and photolyzes
with a disappearance quantum yield of 0.27. The photolysis results in
carbon-mercury bond cleavage to give elemental mercury and phenyl
radicals, The minimum photochemical half-life determined in sunlight
experiments is 8.5 hours.

10. Dissolved phenylmercuric salts undergo photochemical decomposi-
tion with pH independent disappearance quantum yields of 0.24. The
photochemical reaction results in carbon-mercury bond cleavage and
formation of mercurous salts and phenyl free radicals. The experi-
mentally determined minimum half-lives @ 17 hours) show that photo-
decomposition of these compounds may be environmentally significant
under certain conditions,

li. Methylmercuric thiol and methylmercuric sulfide ion complexes
undergo photodecomposition in sunlight. The methylmercuric thiol
complexes have quantum yields from 0.12 to 0.16 with minimum photo-
chemical half-lives of 46 to 120 hours, Methylmercuric sulfide ion
has a quantum yield of 0.65 and a minimum half-life of 0.43 hours.
The major products are methane and inorganic mercury species.
Although oxygen lowers the quantum yields, photochemical reaction
may still be a significant degradative pathway.


7, for diphenylmercury is 2.0 x 10°17 sec, Disappearance quantum
yields in water for phenylmercury ion and phenylmercuric hydroxide
were not decreased by quencher concentrations less than 0.10 M,
indicating that + for these compounds is < 10°*° sec, Because
concentrations of potential quenchers in natural waters are much
lower than 0.1 M, quantum yields in Table 14 should be unaffected
by quenching processes in the environment.

To determine empirically the effect of materials dissolved in natural
waters on the quantum yield for photodecomposition of a phenylmercuric
salt, air-saturated solutions of phenylmercuric acetate (1.0 x 1o-? w)
in two different natural waters and in distilled water were subjected
to equal exposures of Pyrex-filtered mercury-lamp light ( 290 nm).
The phenylmercuric salt photodecomposed at the same rate in all three
solutions, in agreement with the predictions derived from the quench-
ing studies, Dark controls showed no decomposition.

Sunlight photolysis rates calculated from laboratory data and
photolysis rates measured in sunlight are compared in Table 16.
Other experiments showed that the same salts were completely decom-
posed after 18 days in sunlight while dark controls showed no change.
With the exception of diphenylmercury, the compounds were largely
dissociated and hydrolyzed to form 25% phenylmercuric ion and 75%
phenylmercuric hydroxide under the experimental conditions, Results
in Table 16 show that the calculated and empirical photolysis rates
are in reasonable agreement.

70


64.

65.

66.

67.

68.

69.

70.

71.

72.

73.

74.

75.

76.

77,

78.

79,

See Appendix A of this report for more detail.

J. R. Kramer, H. E. Allen, G. W. Baulne, and N. M. Burns,
"Lake Erie Time Study," Canada Centre for Inland Waters Paper
No. 4, Burlington, Ontario, 1972.

R. L. Carr, C. E. Finsterwalder, and M. J. Schibi, Pestic.
Monit. J., 6, 23 (1972).

N. M. Burns and C, Ross, "Project Hypo," U. S, Environmental
Protection Agency, Technical Report TS-05-71-208-24, 1972,
Chapter 7.

G. E. Hutchinson, "A Treatise on Limnology," John Wiley & Sons,
Inc., New York, 1957, Chapter 13.

L, Friberg, et al., Nord. Hyg. Tidkr., 52, Suppl. 4, Chapter 3,
(1971).

J. J. Morgan, “Equilibrium Concepts in Natural Water Systems,"
Advances in Chemistry Series, No. 67, American Chemical
Society, Washington, D. C., pp. 1-29.

A, A. Frost and R. G. Pearson, "Kinetics and Mechanisms," 2nd Ed.,
Wiley and Sons, Inc., New York, N. Y., 1961.

S. I. Shibko and N. Nelson, Environmental Research, 4, 23 (1971).

G. L. Baughman, M, H. Carter, N. L. Wolfe, and R. G. Zepp,
J. Chromatogr., 76, 471 (1973).

G. M. Barrow, "Physical Chemistry," 2nd Ed., McGraw-Hill,
New York, N. Y., 1966, p. 651.

Roger G. Bates, "Determination of pH," John Wiley and Sons, Inc.,
New York, N. ¥., 1954, p. 212.

M. S. Kharasch and M, W. Grafflin, J. Amer. Chem. Soc., 47, 1948
(1925). ~

(a) K. Hartley, H. 0, Pritchard, and H. A. Skinner, Trans, Faraday
Soc., 46, 1019 (1950); (b) H. B, Charman, E. D. Hughes, and
C. K. Ingold, J. Chem. Soc., 2530 (1959).

E. C, Tsivoglou, "Tracer Measurement of Stream Reaeration,"”
USDI, Federal Water Pollution Control Administration, Washington,
D. C., 1967.

M. P,. Thekaekara, R. Kruger, and C, H. Duncan, Applied Optics,
8, 1713 (1969).

76


10

11

12

13

14

TABLES

KINETIC DATA FOR ORGANOMERCURY LIGAND EXCHANGE

DIMETHYLMERCURY ACIDOLYSIS RATE DATA MEASURED
BY GLC 6 ww ee ee ee ee eee ee ee

DIMETHYLMERCURY ACIDOLYSIS RATE DATA DETERMINED

CONDUCTOMETRICALLY . . 1 1 6 ee ee ee ee ee
DIMETHYLMERCURY ACIDOLYSIS IN THE PRESENCE OF
NUCLEOPHILIC SPECIES ... 2. eee ee eee
DIPHENYLMERCURY ACIDOLYSIS RATE DATA ......

INCREASING EASE OF ACID CLEAVAGE OF ARYL AND
ALKYL GROUPS IN ORGANOMERCURIALS ........

CALCULATED EQUILIBRIUM CONSTANTS FOR THE
REACTION OF DIMETHYLMERCURY WITH MERCURIC
HALIDES . 6 6 ee ew ee ee eee we ee te

pH DEPENDENCE FOR DESYMMETRIZATION OF DIMETHYL-
MERCURY AT 27° 2. ww ee ee ee

pH DEPENDENCE FOR DESYMMETRIZATION OF DIPHENYL-
MERCURY BY MERCURIC PERCHLORATE AT 27° .....

EXPERIMENTAL PHOTOLYSIS RATES FOR FOUR METHYL-
MERCURY COMPOUNDS IN SUNLIGHT .........

QUANTUM YIELDS FOR PHOTODEGRADATION OF SULFUR-
BONDED METHYLMERCURY COMPLEXES AT 313 nm IN
DISTILLED WATER... «2 1 ee ee eee

EFFECT OF QUENCHERS UPON PHOTODECOMPOSITION OF
SULFUR-BONDED METHYLMERCURY COMPLEXES IN WATER

CALCULATED SUNLIGHT PHOTOLYSIS RATES FOR SULFUR-
BONDED METHYLMERCURY COMPLEXES AT 25° IN WATER

QUANTUM YIELDS FOR DIRECT PHOTOLYSIS (313 nm) OF
PHENYLMERCURY COMPOUNDS . .. . ns

vii

Page

14

36

37

39
40

43

43

45

46

55

59

61

63

67


27.

28.
29,

30.

31.
32.
33.
34,
35.

36.

37.

38.

39.

40.

41.

42.

43.

45.

D. L. Currell and C. Toppolo, Biochim. Biophys, Acta, 263, 82
(1972).

Y. Lin, et al., Biochem. Biophys. Res, Commun., 47, 1209 (1972).

L. Levine and R. K. Brown, Biochim. Biophys. Acta, 25, 329 (1957).
M. Eigen, G. Geier, and W. Kruse, "Essays in Coordination
Chemistry," Experientia (Birkhauser Verlag, Basel, 1964), Suppl.
9, p. 164.

R. B. Simpson, J. Chem, Phys., 46, 4775 (1967).

H. S. Gutowsky and C. H. Holm, J, Chem. Phys., 25, 1233 (1956).

J. D. Hem, U. S. Geol. Survey Prof, Paper 713, 1970.

H. Zimmer and S, Makower, Naturwissenschaften, 41, 551 (1954).
F. Nerdel and S. Makower, Naturwissenschaften, 45, 491 (1958).

I. P. Beletskaya, L. A. Fedorov, and 0. A. Reutov, Proc. Acad.
Sci., USSR Chem, Sect., 163, 794 (1965).

See reference 6a,

R, E. Dessy, G. F. Reynolds, and J, Kim, J. Amer, Chem. Soc., 81,
2683 (1959).

M. M. Kreevoy, J. Amer, Chem. Soc,, 79, 5927 (1957).

R. D. Brown, A. S. Buchanon, and A. A. Humffray, Aust. J. Chem.,
18, 1507 (1965).

R. J. Ouellette, Ph.D. Thesis, University of California, Berkeley,
1962.

E. G. Janzen and R. J. Blackburn, J. Amer. Chem. Soc., 91, 4481
(1969). ~

(a) A, Terenin, J. Chem, Phys., 2, 441 (1934); (b) J. W. Linnett
and H. W, Thompson, Trans. Faraday Soc., 33, 501 (1937); (c) R.
Gomer and W, A, Noyes, Jr., J. Amer. Chem. Soc., 71, 3390 (1949);
(d) R. A, Holroyd and W. A. Noyes, J. Amer, Chem, Soc., 76, 1583
(1954), ~

R. E. Rebbert and P. Ausloos, J, Amer, Chem. Soc., 85, 3086 (1963).

TI. Fagerstrom and A. Jernelév, Water Research, 6, 1195 (1972).

74


UcHsHgSR] _ K, (F3) ERS

(79)
© [chug] 1+Kk, [a]
(CHyHg0H] _ Ky (F3) & (80)
= [CHsHg) ty]
(CHaHeHPO”] _ K, (3) EPQ @1)

= [CHgHg | 1+k, {t]

(CHjHeNH,*] Ks (F3) EN «s2)
= (cHyHe] 1+, [at]
(CHsHgOAr'| _ ky (F3) ZArO (83)

E (cHyHg) 1+k, (HI

The computer program used for the calculations is shown in Figure 1A.
Originally developed as a Conversational Programming System (CPS)
program, it was later converted to Fortran IV for plotting of the
calculated fractions, Calculations were carried out on an IBM 360/67
computer. Calculations from the latter were plotted on a Tektronix
4010-1 terminal. Typical computer plots are shown on the following
pages (Figures 2A-7A). Concentrations of all chemical species used
in the calculations, except sulfur compounds, are given below,

= (CH,Hg] = 107° M
ERN= 10° M

DPQ = 10° M

EN=10°5TMM
= Aro = 10° M
{cr ] = 10* mM

85


The effect of added quencher upon the quantum yield is defined by
the following Stern-Volmer expression,

Po

= 1+k, 7 [Q] (44)
%Q

where $, and ¢g are quantum yields without and with added quencher,
respectively, k, is the bimolecular rate constant for the quenching
process, and t is the lifetime of the excited, photoreactive molecule
in the absence of added quencher. Since several excellent quenchers,
such as oxygen, are present in the aquatic environment, the effect of
known quenchers upon photodecomposition rates of organomercury com-
pounds was also examined.

Photodecomposition of Dimethylmercury, Methylmercuric Ion,
Methylmercuric Hydroxide, and Methylmercuric Halides

The ultraviolet absorption spectra of (CHg )oHg, CH, Hg" , and CHgHg0H
in water revealed that these species absorb virtually no light at
wavelengths greater than 290 nm (Figure 15). Specific absorption
rate constants for these species are extremely low, so sunlight
photolysis rates are very slow. Spectral predictions of low sunlight
photoreactivity were corroborated by irradiating aqueous solutions of
(CHy)gHg, CHsHg*, and CH,Hg0H with Pyrex-filtered ultraviolet light
from a mercury lamp, Prolonged irradiations resulted in no photo-
decomposition of these species.

Other laboratory experiments showed that the decomposition of these
methylmercury species was not photosensitized by acetone, a high
energy sensitizer (triplet energy, 80 kcal moleTM+), Moreover,
dimethylmercury was not degraded by singlet oxygen, a chemical species
important in environmental chemistry.°°

Sunlight irradiation of aqueous 10°* M solutions of methylmercuric
hydroxide, methylmercuric chloride, methylmercuric bromide, and
methylmercuric iodide resulted in rapid photodecomposition of CH,HgI,
slow photodecomposition of CHgHgBr, and negligible degradation of
CHgHgC1 and CH3HgOH (Table 10). Dark controls at the same tempera-
ture (30°) showed no change. These and following results demonstrate
that changes in the ligand bonded to methylmercury markedly affect
photoreactivity.

Facile photodecomposition of CH3HgI could cause significant errors in
the analysis for methylmercury content of environmental samples since
several widely used procedures call for gas chromatographic analysis
of methylmercuric iodide in organic solvents.°* 78,87 Because methyl-
mercuric iodide is photodecomposed by fluorescent lights, precautions
should be taken to shield the CH,HgI solutions from light during the
analysis.

53


[ugtl} ((ugIT1 = [ne** 1 + [ugon* ] + [Hg(0H)o]) is in excess. Using
a value of 0.03 ppb for [Hg!f] and k = 1 x 10°, at pH 5.8 (25°), the
half-life for reaction with DMM is about 50 days, For DPM under the
same conditions, the half-life would be four hours.

Dimethylmercury Stability to Oxygen and Base

DMM was shown to be stable to 1M KOH, After heating for 24 hours
at 85°, gle analysis showed that DMM had not reacted, Likewise DMM
neither reacted with KI (five hours at 85°) nor did it react with a
saturated solution of egg albumin (23 hours at 30°).

Secondary and tertiary dialkylmercury compounds are reported to
undergo slow oxidation by oxygen in organic solvents.® Primary
dialkylmercury compounds are less susceptible to oxidation. DMM
reaction could not be detected by gle analysis after standing for 20
hours at 85° in an aqueous solution saturated with oxygen, Lack of
reaction under these stringent conditions precludes any significant
contribution to degradation in the environment.

Evaporative Loss

Evaporative loss of organomercurials from the aquatic environment is
a physical process, but because of its potential importance as a
means of transport, it is discussed here. Many organomercurials have
a substantial vapor pressure. Although the ionic or polar covalent
mercurial salts when dissolved in water will not be readily lost to
the atmosphere because of solvation effects, mercury and dimethyl-
mercury are non-ionic compounds and evaporative loss may be important.

An estimate of the magnitude of loss may be obtained by employing
the method and data of Tsivoglou. 78 Using the relationship between
transfer coefficients and molecular diameters (K,/Kg = dz/d,) and
van der Waals radii®! to calculate molecular diameters, the ratio
of transfer constants for DMM to oxygen is calculated to be 2.4,
Utilization of Tsivoglou's’© experimental data reveals that for a
moderately turbulent river section, DMM would have an evaporative
half-life of about 12 hours, This relationship also predicts that
dissolved elemental mercury would be lost from the river at a rate
2.3 times faster than DMM.

KINETICS OF ORGANOMERCURY PHOTODECOMPOSITION

Since tropospheric solar radiation has negligible intensity at
wavelengths less than about 290 nm (Figure 14),°% 7° systems must
have appreciable absorptivity at wavelengths greater than 290 nm if
significant photoreaction is to occur in sunlight. Predictions of
sunlight photoreactivity can be derived from laboratory experiments
employing Pyrex-glass filtration of light from a mercury lamp since
this filter transmits only wavelengths greater than 290 nm.°°

49


100

fe}
So
|

PERCENT CH3HgOH

20

ron
ro)
T

&
T

Figure 3,

pH dependence of the hydrolysis of methylmercuric ion.

11


0”

Table 5, DIPHENYLMERCURY ACIDOLYSIS RATE DATA

Temp. | % Ethanol Conc. DPM | Conc, HC1O, Ant AS*
°¢ in Water M M k secTM* x 10*] k 2/mole sec x 10! | kcal/mole | eu
70 40 8.80 x 107° | 2.40 x 1074 1.88 + 0,07
70 40 8.80 x 10°] 2.40 x 107% 1.69 + 0,07
70 30 8.80 x 10°] 2.40 x 1o* 1.44 + 0,04
70 30 8.80 x 10° | 2.40 x 10% 1.31 + 0,05
70 0 1.02 x 10°© | 1.37 x 1o-® | (7,80 + 0.3) 5.69 + 0,19

18.1 -7.0
50 0 1,02 x 107° | 1.37 x 10°° | (1.53 + 0,02) 1.12 + 0,02
25 9.67 x 10%"

*Extrapolated value.


Kinetic data for the desymmetrizations of DMM and DPM by Hg(C10; )o
are summarized in Tables 8 and 9, respectively. Rate constants were
obtained by use of the integrated second-order rate expression, 7
which was obeyed through two half-lives as shown in Figure 13.

In the kinetic studies, Hg(C1Q;)z2 formed by the reaction of Hg(OAc)2
with perchloric acid gave results identical to those obtained with
Hg(C10,)2 formed by reacting HgO with HC1Q,. The perchlorate was
chosen because it is completely dissociated in aqueous solution,
eliminating interference by any associated salt, which simplifies
the reaction kinetics (see Section IV).

The desymmetrization rate constants for DMM and DPM show a strong

pH dependence varying six orders of magnitude over the pH range of

1 to 9. The reaction velocity for DPM at pH 2 was so fast (near
diffusion controlled) that the rate constant could only be estimated
by single-point determinations. The rate constant obtained (2 x 10° t/
mole sec)therefore represents a lower limit. A calculation of
acidolysis half-lives at the pH's employed indicates that acidolysis
was not a competing reaction under the conditions of the desymmetri-~
zation reactions.

The strong pH dependency of the desymmetrization reaction may be due
to the dependence of the concentration of the various mercuric
Species on pH, The hydrolysis of mercuric ion (eqs 31 and 32)
involves three different mercuric species (ag'* , HgOH', and Hg(OH)2)
whose relative concentrations are a function of hydronium ion con-
centration,

at

Hg** +H,0 2 Heton + HY (31)

Hg" OH + H20 = Hg(OH)2 + H (32)

Figure 2 (Section IV) presents a plot of relative mercuric ion con-
centration as a function of pH. Below pH 1.5 the predominant mercuric
species is Hg"* and above pH 4.5 the predominant species is Hg(OH)2.
Between pH 1.5 and 4.5, the third species, ust OH, never exceeds about
10% of the total mercuric species.

At low p's desymmetrization is primarily due to electrophilic attack
by Hg** on the organomercurial, represented in equation 33 as a pure
Sp 2 mechanism, The reaction at pH 9 is attributed to electrophilic
attack by Hg(OH)2, represented in equation 34 as an Sp i mechanism.
A pure Sp 2 mechanism is probably also operative for Hg*oH.®*

44


11.

12,

13.

14,

15.

16.

17.

18.

19,

20.

21.

22,

23.

24,

25.

26.

(a) H. Ley, Z. Phys. Chem., 30, 248 (1899); (b) A. B. Garrett and
A. E, Hinchler, J, Amer. Chem. Soc., 60, 299 (1938); (c) J. Bjerrum,

Dissertation, Copenhagen, 170 (1941); (d) A, B. Garrett and W. W.
Howell, J. Amer. Chem, Soc., 61, 1730 (1939).

(a) J. L. Maynard and H. C, Howard, J. Chem. Soc., 123, 960
(1923); (b) I. B. Johns, W. D., Peterson, and R. M. Hixon, J. Amer.
Chem. Soc., 34, 2218 (1930),

(a) R. M. Schramm, J. Amer, Chem. Soc., 69, 1831 (1947); (b) R.
Benesch and R. E, Benesch, J. Amer, Ch Soc., 73, 3391 (1951).

S, Hietanen and L. D, Sillén, Acta Chem, Scand., 6, 747 (1952).
R. B. Simpson, J, Amer, Chem. Soc., 83, 4711 (1961).

(a) T. D, Waugh, H. E. Walton, and J. A. Laswick, J. Phys. Chem.,
59, 395 (1955); (b) S. S. Parikh and T, R. Sweet, J. Phys. Chem. ,
65, 1909 (1961); (c) E., A. Smith, Diss. Abstr, Int. B., 32, 5075
71932) ; Dissertation, Rutgers University, 1971.

L. F. Sytsma, Diss, Abstr, Int. B, 32, 6311 (1972).

G. Schwarzenbach and M. Schellenberg, Helv. Chim, Acta, 48, 28
(1965).

J. D, Calvert and J. N. Pitts, "Photochemistry," John Wiley and
Sons, Inc., New York, 1966, p. 270.

(a) P. L. Goggin and L. A. Woodward, Trans. Faraday Soc., 56, 159
(1960); (b) J. H. R. Clarke and L. A. Woodward, Trans. Faraday
Soc., 644, 1041 (1968).

F. A, Cotton and G, Wilkinson, "Advanced Inorganic Chemistry,"
Interscience Publishers, New York, 1966, p. 616.

See Appendices for Tables of equilibrium constants and additional
discussion.

H, Edelhoch, E, Katchalski, R. H. Mayberry, W. L. Hughes, Jr.,
and J, T. Edsall, J. Amer. Chem. Soc., 75, 5058 (1953).

P. D. Boyer in "The Enzymes," P. D. Boyer, H. Lardy, and
K. Myrbach, Ed., Academic Press, New York, 1959, Chapter 11.

Pp. D. Boyer, J. Amer. Chem, Soc., 76, 4331 (1954).

E, Chiancone, D. L. Currell, P. Vecchine, E. Antonini, and J.
Wyman, J. Biol, Chem., 245, 4105 (1970).

73


(CeHs)oHg

| os
wwe |

Hg® ———» Hgt* ——» (CH3)eHg

CH30CH»CHeHg*

Figure 1. Model of the environmental mercury cycle.


R-Hg-R Reg’ R R
+ > toe > booty
He** Hig” Hg Hg (33)
Sp 2
OR
R-Hg-R R--Hg R R
+ | \ (34)
Hg-OH : H Hg + ug
’ Hg---01 i t
OH t ie) to}
’ t
OH H
Sp i

The high rate of electrophilic substitution by Hg’* as compared to
Hg(0H)2 represents about five orders of magnitude in increased
electrophilicity towards the carbon-mercury bond. Although extensive
background data are not available to evaluate the desymmetrization
reaction for other mercuric species, the more electropositive the
mercury atom, the larger the rate constant. For example, in ethanol
the order of reactivity is Hg(NOs)p> Hg(OAc)z > Hg(Br)2, which is
also the order of decreasing ionic character of the mercury bond.””?
The reverse reaction, symmetrization (eq 30) has been suggested as
a possible chemical pathway for the conversion of methylmercuric ion
to DMM.*° It was proposed that the equilibrium would be shifted to
the left by removal of mercuric ion as insoluble mercuric sulfide.

A maximum value for the symmetrization rate constant can be obtained
from equilibrium data (Kg = kiesyn/Ksyq). Experiments showed that
the equilibrium constant (K,,) was > 10*; thus the rate constant for
symmetrization must be four orders of magnitude smaller than that

for desymmetrization, Using a methylmercuric hydroxide concentration
of 1 ppb (4 x 10° M) and the second-order rate constants in Table 8,
calculations (eq 24) show this reaction to have a half-life of
several years,

The effect of sulfide on the symmetrization reaction was examined
employing 0.4 M sodium sulfide and 0,1 M methylmercuric hydroxide in
water at 25°. At the end of 24 hours, the organomercurial was 25%
reacted, The second-order rate constant derived from these data
was ~ 10° Z/mole sec, too low to permit the reaction to be environ-
mentally significant.

Based on the rate constants (Tables 8 and 9), half-lives can be
calculated, under pseudo-first-order reaction conditions, where

48


VI

VIL

VIIL

IX

PHOTOCHEMICAL PROCEDURES

PHOTOLYSIS OF METHYLMERCURIC SULFIDE ION

SUNLIGHT PHOTOLYSIS OF METHYLMERCURY COMPOUNDS

ACETONE-SENSITIZED PHOTODECOMPOSITION OF
PHENYLMERCURIC ION AND PHENYLMERCURLC HYDROXIDE .

QUENCHING STUDIES OF PHENYLMERCURY COMPOUNDS

SUNLIGHT PHOTOLYSIS OF PHENYLMERCURY COMPOUNDS

RESULTS AND DISCUSSION

CALCULATED EQUILIBRIUM CONCENTRATIONS OF METHYL-
MERCURIC SPECIES IN AQUATIC SYSTEMS .

KINETICS OF ORGANOMERCURY REACTIONS IN WATER

Theoretical Considerations

Acidolysis of Dimethyl- and Diphenylmercury ...

Desymmetrization of Dimethyl- and Diphenylmercury

Dimethylmercur:

Stability to Oxygen and Base

Evaporative Loss

KINETICS OF ORGANOMERCURY PHOTODECOMPOSITION

Photodecomposition of Dimethylmercury, Methyl-

mercuric Ion

Methylmercuric Hydroxide, and

Methylmercuric Halides

Photoreactivit

Complexes .

Photocleavage of Phenylmercury Compounds

REFERENCES

PUBLICATIONS

APPENDICES

iv

of Sulfur-bonded Methylmercur:

23

23
23
24

25

25
28
28
32
41
49
49

49

53

56
60
72
78

79


K
CeHsHg* + H,0 = CeHsHgoH + Ht (47)

K = 6.8 x 10° (Reference 16)

Irradiation of diphenylmercury, phenylmercuric hydroxide, and phenyl-
mercury ion in degassed solutions caused the following reactions:

hy
(Ces )2Hg ~ CoH yHg’ + CeHs° (48)
CoHsHg + CgHy: + He} (49)
hy
CegHjHe* - CsHs* + *He* (50)
2°Hg* + Heg3* (51)
hy
CgHsHg0H > CoH,* + +HgOH G2)
2-HgOH - Hge (OH)p 63)
Hg2(OH)2 ~ Hg0! + He®l + H20 (54)
Cols’ + RH - CoHs +R (55)
2CgHs* > (Ces Jat (56)

Quantitative yields of metallic mercury and phenyl free radicals
resulted from irradiation of diphenylmercury. Phenylmercuric ion
photolyzed to give phenyl radicals and mercurous ions, and phenyl-
mercuric hydroxide yielded phenyl radicals and nearly quantitative
yields of metallic mercury and yellow mercuric oxide. The fate of
the phenyl radicals depended upon the composition of the reaction
media, With organic materials (RH) present in the aqueous media,
phenyl radicals reacted to form nearly quantitative yields of
benzene (eq 55). In oxygen free distilled water containing no
additional organic substances, the most important phenyl radical
reaction was coupling to form biphenyl (eq 56). With oxygen present,
no biphenyl was formed, presumably because the phenyl cadicals were
scavenged by the following reaction.

Cog’ + 02 7 CoH. O2° (57)

65


SEPTUM

i, TEFLON STOPCOCK
oN =
—

ay,

DIMETHY LMERCURY
7 SOLUTION

4

GLASS TUBE

NN”

Figure 5. Enclosed bomb for kinetic studies.

20


Table 13, CALCULATED SUNLIGHT PHOTOLYSIS RATES FOR SULFUR-BONDED
METHYLMERCURY COMPLEXES AT 25° IN WATER

10? (ke )inte 10? Oka inte ty
Complex hr? hyo? hr
CH,Hg-egg albumin 10 12 58
CHsHgSCH, CO2H 10 15 46
CH, HgSCHz CH2 OH il 15 46
CHg HgSCH2 CH (NHg* )CO2H 3.6 5.8 120
Cig HgSTM 240 1600 0.43

63



Photoreactivity of Sulfur-bonded Methylmercury Complexes

As discussed earlier, under certain environmental conditions, a
large fraction of the methylmercury species exists as sulfur-bonded
methylmercury complexes. Earlier studies®® °° showed that mercuric
mercaptides are readily decomposed by light from a mercury lamp or
even ordinary room light.

Previously recorded ultraviolet spectra showed that some organo-
mercuric mercaptides have significantly large extinction coefficients
at wavelengths > 290 nm.”° Methylmercuric sulfide ion, bis(methyl-
mercuric) sulfide, and methylmercury-thiol complexes also absorb at
wavelengths > 290 nm (Figures 16 and 17). Preliminary experiments
showed that Pyrex-filtered ultraviolet light decomposed the sulfide
and thiol complexes with cleavage of the methylmercury bond.

hv
CHygHeS' + CH,t + HgSt (45)

hy
CHjHgSR + CHkt + ColHigt

+ Inorganic mercury precipitate | RSSR (46)

Products of the photolyses were methane and ethane (20:1) gases and
inorganic mercury precipitates. The black precipitate from photolysis
of methylmercuric sulfide ion was found to be mercuric sulfide, and
the precipitate from the methylmercury-thiol photolysis was not
identified, but was shown by mass spectrometry to contain no organi-
cally bound mercury. Yields of inorganic mercury in the precipitates
quantitatively accounted for disappearance of the methylmercury
complexes. Under these same conditions, unphotolyzed controls

formed no inorganic mercury during the photolysis period,

Disappearance quantum yields for photodegradation of several
different methylmercury-thiol complexes by 313 nm light were similar
(Table 11). Quantum yields for photodecomposition of the cysteine
and egg albumin complexes are particularly significant, since these
complexes are good models for methylmercury-thiol complexes that
form in biological systems,>® °° In addition to complexation by
sulfhydryl groups, the methylmercury ion was probably also complexed
by other functional groups such as amino and hydroxyl groups in the
albumin protein.°° Such complexes would not be decomposed by 313 nm
light, because methylmercury-nitrogen and methylmercury-oxygen
complexes do not absorb light at wavelengths > 290 nm.**

Photodecomposition of the methylmercuric sulfide ion proceeded with a

high quantum yield (Table 11). The complex was so light-sensitive
that extensive photodegradation by fluorescent lights in the

56


showed that the reaction of p-chloromercuribenzoate (PMB), a substi-
tuted phenylmercury compound, with sulfhydryl groups could be used
to distinguish kinetically different types of -SH groups in proteins.
Subsequent studies by other groups have established that second-
order rate constants for reactions of PMB with protein sulfhydryl
groups range from about 10° t/mole sec ("masked" -SH groups)"°~"* to
10° 2/mole sec ("free -SH groups)®® at temperatures in the 20-30°C
range. Rapid exchange also occurs with complexed divalent inorganic
mercury. For example, the sterically hindered reaction of the mercury
dimer of serum albumin (ASHgSA) shown below proceeds with a rate
constant of 5 x 10° 2/mole sec at 24°C,?°

ASHgSA + Hg(HPQ:)2 — 2ASHg (HPQ, ) (9)

Presumably, unhindered thiol-mercury complexes would exchange much
more rapidly.

Recent work has provided kinetic data on the very rapid exchange
reactions involving non-sterically hindered methylmercury complexes,
Using the temperature-jump method, Eigen, Geier, and Kruse have
determined forward and reverse rate constants for reaction 10.°°

k
£
CHgHgX + OW = CHgHgOH + X (10)
k,
Tr

Simpson®* determined rate constants for ligand exchange using the
nuclear magnetic resonance technique of Gutowsky and Holm.*? Rate
constants for the exchange reactions frequently were found to be
close to diffusion-controlled, with the slowest rate constants about
10 Z/mole sec (Table 1). These kinetic data indicate that achieve-
ment of ligand exchange equilibrium in systems such as natural
waters that contain multiple chemical species and mercury compounds
is quite rapid even at the very low concentrations observed in the
aquatic environment.

Relative equilibrium concentrations of reduced and oxidized inorganic
mercury have been calculated by Hem for a model system containing
chloride and sulfate.*° Hem used these calculations to predict the
most stable forms of inorganic mercury in the aquatic environment.

No calculations of relative equilibrium concentrations of organo-
mercury complexes under environmental conditions have been reported
in the literature. Calculations that we have carried out for
methylmercury complexes are discussed later in this report.

CHEMICAL DEGRADATION OF ORGANOMERCURTALS
Kinetic studies of chemical degradation of organomercurials

generally have not been carried out in water. Nonetheless, a few
relevant studies have been gleaned from the literature,

13


PERCENT [CH3H9] jor

(CH3HgC!] , (CHsHgOH]
[CH3H9] sor [CH3Hg} ror
100 F
30 + SPECIES CONC, M
(CH3Hg].,, <1o7??
i [HePOe7} + [HPO] 106
60 [cl] 10-*
- [NHe*] + [NHs] 10
40 ([RNH3*] + [RNH»] 10-6
[RSH] + [RS7] 0
fs [ArOH] + [ArO7] 10°
20 r [HS] + [SH] + [S=] 0
Q bed LL Ne td
0 2 4 6 8 10 12

Figure 9, Relative concentrations of methylmercury complexes
in systems containing multiple chemical species,
excluding reduced sulfur species and organic thiols.

30


J, x10" PHOTONS, cm”? sec7?10 nm=?

5b
ab
34+
2b
l —
| | | | !

300 400 500 600 700 800
WAVELENGTH, nm

Figure 14, Estimated solar irradiance at
solar zenith angle = 0°,

50


rate constant, 8 x 10°2° sec”? or a half-life of about 3 x 10* years.
No data on demercuration of phenylmercuric salts are available from
the literature.

PHOTODECOMPOSITION OF ORGANOMERCURIALS

Most of the abundant literature on the photochemistry of organo-
mercury compounds is concerned with product studies.?° Upon
irradiation, the carbon-mercury bond of organomercurials is cleaved
to give alkyl or aryl free radicals and inorganic mercury (eqs 14-17).

hy
ReHg + R: + ‘HgR (14)
‘HgR > Hg +R: (15)
hy
RHgX - R* + *HgX (16)
2*HeX 7+ Hgeke (17)

That alkyl and phenyl free radicals are formed from photolysis of
organomercuric salts has been demonstrated by electron spin resonance
studies of Janzen and Blackburn.*®

Quantitative mechanistic studies of organomercury photoreactions are
difficult to find in the literature. However, the photolysis of
dimethylmercury has been intensely studied. In the vapor phase,
dimethylmercury (R = CHg, eqs 14 and 15) photodecomposed at 30°

with a quantum yield of unity to give methyl radicals and elemental
mercury.*® In the liquid phase, a significant fraction of the methyl
radicals underwent cage recombination to form ethane, and the
remainder abstracted hydrogen atoms to form methane. 4 Although
Fagerstrom and Jernelév have suggested that photodegradation of
dimethylmercury is important in the environment,*® the insignificant
absorption of (CHg)2Hg at wavelengths > 280 nm? indicates that its
sunlight degradation must be very slow. Experiments of Strausz,

Do Minh, and Font have shown that photodecomposition of a dialkyl-
mercury compound is extremely slow when wavelengths available from
solar radiation, i.e., Pyrex-filtered light (> 290 nm), are
employed .*&

Although no quantum yield data were found for the photodecomposition
of phenylmercury compounds or alkylmercuric salts, a few publica-
tions on the sunlight stability of these compounds are available.
Takehara and coworkers irradiated several phenylmercuric and
alkylmercuric compounds with several light sources.*” Irradiation

16


ES = [H:S] + [s7] + [si] (61)

ZRN = [RNH,] + [RNH3*] (62)
urs = [Rs”] + (Rsw] (63)
rPQ, = [H,PQj”] + [HPO] (64)
N= [Nay] + (yH.* J (65)
LAro = [ArOH] + [Aro] (66)

Substituting expressions for acid-base equilibria into equations
61-66 and rearranging terms results in the following:

7 6.358 ‘
~ 10°? (Ht + 1o'® Ht] + 6.3 (67)

(s=]

=RN

[RNH2 J = 14k, ti)

K, = 10°° (68)

-, _ —  2RS 499.52
(1 = Tye TI K, = 10° (69)

=
ec se re @o

— =n = 9.42
Oe) = TA K, = 10 one)

DArO
[aro] = aaa K, = 10% (72)

where K,, Ks, K,, Ka, and K, are equilibrium constants for protonation
of amino acids, thiolate ions, hydrogen phosphate, ammonia, and
phenolate ions, respectively.

83


LE

Table 3, DIMETHYLMERCURY ACIDOLYSIS RATE DATA DETERMINED CONDUCTOMETRICALLY*

Pseudo-First Order

Second-order
Rate Constant?

Cone, Acid Conc, DMM Temp. Rate Constant
Acid M x 104 M x 107 °c k sec7? x 10* k 2/mole sec x 10”
HC1 7.13 1.16 65 0,792 0.70
HCL 7,13 1.07 85 4.12 + 0.0005 4.0
HC1OQ, 7.20 0.501 65 0.756 + 0,01 0.78
HC1O, 7.20 1,01 85 1.94 + 0.01 4.1
HNOg 6.90 1.15 65 0.937 + 0.48 0.81
HNO; 6.90 0,52 85 2.70 + 0.01 5.4
“See Section V.
- Obtained from pseudo-first-order rate constants,


aan

SECTION VIII

PUBLICATIONS

G. L. Baughman, M. H. Carter, N. L. Wolfe, and R. G. Zepp,
Gas-liquid Chromatography-Mass Spectrometry of Organomercury
Compounds. J. Chromatogr. (Amsterdam), 76, 471 (1973).

N. L. Wolfe, R. G. Zepp, J. A. Gordon, and G. L. Baughman,
Chemistry of Phenylmercury Compounds in the Aquatic Environ-
ment. Chemosphere (Oxford), 1, 273 (1972).

R. G. Zepp, N. L. Wolfe, and J. A. Gordon, Photodecomposition
of Phenylmercury Compounds in Sunlight. Chemosphere (Oxford),
2 (in press).

N. L. Wolfe, R. G. Zepp, J. A. Gordon, and G. L. Baughman,
Chemistry of Methylmercurials in Aqueous Solution. Chemosphere
(Oxford) (in press).

R. G. Zepp, N. L. Wolfe, and G. L. Baughman, Methylmercury
Complexes in Aquatic Systems. Environ. Lett. (in preparation).

78


Table 10.

EXPERIMENTAL PHOTOLYSIS RATES FOR FOUR
METHYLMERCURY COMPOUNDS IN SUNLIGHT

; 10° Oks Dangg ty
Compound’ hr hrs
CHgHgl > 230 < 3
CH, HgBr 6.2 110
CH; HgC1 < 1.3 > 530
CH; HgOH < 1.3 > 530

* Concentration 1,00 x 10°* M in water.
° Expressed in hours of sunlight, not actual hours.

55


A.

Be

SECTION IX

APPENDICES

Methylmercuric Ion Equilibrium Calculations .. .

Equilibrium Constants for Formation, Dissociation,
and Hydrolysis of Mercury Species . +... +...

79


molecular oxygen was quencher), and irradiated in parallel at 313 nm
on the photochemical apparatus.

SUNLIGHT PHOTOLYSIS OF PHENYLMERCURY COMPOUNDS

Air-saturated solutions of diphenylmercury (1.0 x 10°° M) and
phenylmercuric salts (4.0 x 10°° M) in water were sealed in 13.0

+ 0.1 mm quartz tubes, The tubes were irradiated by sunlight in an
exposed area on the roof of the Southeast Environmental Research
Laboratory, Athens, Georgia. In a preliminary experiment, tubes
were irradiated for 18 days during July (average temperature, 30°C
during daylight hours). In a second experiment, tubes were irra-
diated for two days in August (average temperature 31°C). During
this period the weather was mostly fair, and the tubes received
20.0 hours of sunlight. Unphotolyzed controls showed no decrease
in phenylmercurial concentration during the irradiation periods.
Based on the fraction of phenylmercurial that disappeared, photolysis

rate constants and half-lives were calculated assuming first-order
kinetics ,°* °*

24


ACKNOWLEDGMENTS

We express our appreciation for the support and assistance of

Dr, Walter M. Sanders III, Chief, National Pollutants Fate Research
Program, Southeast Environmental Research Laboratory, Environmental
Protection Agency, Athens, Georgia.

We also wish to thank Mrs, Helein Bennett of NASA's J. F. Kennedy
Space Center for her assistance in X-ray diffraction analysis, and
the staff of the National Water Contaminants Characterization Research
Program, Southeast Environmental Research Laboratory, for their help
and the use of their instrumentation.

ix


6€

Table 4, DIMETHYLMERCURY ACIDOLYSIS IN THE PRESENCE OF NUCLEOPHILIC SPECIES*
Cone, Acid Conc. DMM Temp.
Acid M M °¢ k £/mole sec Nucleophile
HCL 8.27 x 10° | 3.40 x 10° 85 (5.13 + 0.05) x 10-*
1.03 x 10°? M NaCl

HCL 8.27 x 10° 15.72 x 10 85
Cysteine-HCl’ | 7.11 x 10° 16.36 x 1o-® 65 --
Cysteine-HCl” | 7.84 x 10° | 5.60 x 10°? 85 --
HCL 8.27 x 1o8 | 3.50 x 107% 85 1.02 x 10°? M Nal
HCL 1.0 x 10" | 5.05 x 10% 85 Nags
*Gas-liquid chromatography technique (see Section V).

FpH 2.23,


Page
QUANTUM YIELDS FOR PHOTOSENSITIZATION OF
PHENYLMERCURIALS «2 ee ee eee te ee ee ee 68

COMPARISON OF CALCULATED AND EMPIRICAL RATES
FOR PHOTODEGRADATION OF PHENYLMERCURY COMPOUNDS
IN SUNLIGHT. ....

a rs |

viii


41.2

4.12

0.41 a
= 3
oS <4
# Ed
z 3
= =

0.041

0.0041

. : 0.01
10-> 10° 107 1078
(B], moles/1

Figure 10. Half-life of A vs. concentration of B for second-order

reactions when [B] >>> [A].

33


3

80 F
x=
fo)
2
a 60 —_
x=
a CsHsHgOH
a
o 40 -
oc
wi
a
20 F
0 1 l
0 2 4 6 8 10

Figure 4, pH dependence of the hydrolysis of phenylmercuric ion.

12


A comprehensive kinetic study is presented of some chemical and
photochemical reactions expected to result in degradation and cycling
of organomercurials in the aquatic environment, Study was limited to
the simple phenyl and methyl derivatives because they are known to
enter aquatic systems and because higher homologs are less stable.
The following reactions were studied.

(Clg )olig CHgHg* ——y
Inorganic mercury species
—_—
(Cells )allg CsHgHg” ——t

The rate equations are the result of laboratory studies and should
predict the rates of the respective reactions in water. They will
not predict the net rate of the reaction in natural systems unless
the reaction in question is the rate determining step.


10

11
12

13

14

15

FIGURES

Model of the environmental mercury cycle .. 4... «es
pH dependence of the hydrolysis of mercuric ion... .

pH dependence of the hydrolysis of methylmercuric

GON see ww ee ee eee es

pH dependence of the hydrolysis of phenylmercuric

GON se ee ee ee ee
Enclosed bomb for kinetic studies - +--+ +++ sees

Relative concentrations of CH,HgS in systems
containing multiple chemical species .....

Relative concentrations of (CHgHg)2S in systems
containing multiple chemical species ... +... ses

Relative concentrations of CH,HgCl and CH3HgSR
in systems containing multiple chemical species,
excluding reduced sulfur . 2. 6 6 0 se ee ee eee

Relative concentrations of methylmercury complexes
in systems containing multiple chemical species,
excluding reduced sulfur species and organic thiols . -

Half-life of A vs. concentration of B for second-order

reactions when TB] >>> [A]... 0 ee eee eee
Acidolysis of dimethylmercury . .. 1. 6 6 + ee ee ee
Acidolysis of diphenylmercury . . . 1. 1 6 1 ee te ee

Desymmetrization of dimethylmercury by mercuric
perchlorate , . ee ee ee ee ee te

Estimated solar irradiance at solar zenith
angle = 0° ... wee ee ee ee ee

UV absorption spectra of dimethylmercury, methyl-
mercuric ion, and methylmercuric hydroxide in

wateL . 6 we ee ee ee ee ee ee

v

10

11

12

20

26

27

29

30

33
35

42

47

50

54


re — ———y

a
3 CH gHgSCHyCH2OH |
eo i
c ®  CHsHgSCHeCOOH
e2 —
ie a
rE
N j | Jeena j
2 4 6 8 0 86

[Q] x 10?, moles/1

Figure 18. Effect of quencher, 2,4-hexadien-1-ol,
upon quantum yields for photocleavage
of methylmercury-thiol complexes.

62


of several phenylmercuric salts by an intense 2537A low-pressure
mercury lamp or sunlight resulted in extensive photodecomposition.
Photoreaction was more rapid when the compounds were dissolved in
water than when they were irradiated as pure solids or in dust
formulations. Sunlight irradiation of aqueous phenylmercuric
acetate caused 25% decomposition in 10 hours. Major products from
photolysis of phenylmercuric acetate and phenylmercuric chloride in
water were reported to be Hg,0 and Hg,Clz, respectively. Shiina
and coworkers reported results that seemingly conflicted with
Takehara's findings, These workers reported that irradiation of
phenylmercuric acetate, phenylmercuric chloride, methylmercuric
iodide, and several other phenylmercuric salts resulted in little
decomposition after light exposure “equivalent to 7 summer days” of
sunlight.*® The apparent conflict in the two publications was
probably due to differences in analytical procedures, Takehara
analyzed directly for the residual organomercury compounds, whereas
Shiina measured the "residual pesticide effect" of the photolyzed
organomercurials upon a fungus culture. Comparison of the two
publications suggests that the photoproducts also act as effective
fungus growth inhibitors.

Taken together, the published studies indicate that sunlight photo-
decomposition of organomercury compounds may provide an important
pathway for conversion of such compounds to inorganic mercury. The
lack of quantitative data on the rates of these photoreactions
prompted the studies included in this report,

7


SECTION II

RECOMMENDATIONS

1. Studies are needed to elucidate the rates and mechanisms of the
release of dimethylmercury and elemental mercury from sediments to
overlying water. Because these materials are non-ionic and have a
significant vapor pressure, they may play an important role in
mercury cycling through loss at the air-water interface.

2, Calculations presented here indicate that elemental mercury and
dimethylmercury are likely to volatilize from the aquatic environment
Because little is known about the behavior of gaseous mercury species
in the atmosphere, research should be undertaken to determine the
fate of these materials.

3. The chemical and physical behavior of mercury in the environment
can be determined by species with which the mercury complexes.
Characterization of these complexes is essential for evaluation of
the chemical, physical and biological processes in the mercury cycle.
Therefore, water from sediments and from the water column should be
analyzed to identify the mercury species present.

4. Redox properties of natural aquatic systems should be investi-
gated with particular emphasis on the rates and mechanisms of
oxidation-reduction reactions. An understanding of these properties
is required before general statements can be made about the effect
of redox potentials on organomercuric or other pollutants.

5. Ligand exchange rates should be examined in detail for complexes
of mercury and other heavy metals, Our literature survey indicates
that some exchange reactions with ligands common to natural waters
are fast enough to approach equilibrium under environmental condi-
tions. This study should include humic acids as well as proteins
and other biological ligands.

6. Concentrations of reduced sulfur species and sulfhydryl-containing
organics should be measured in natural waters, These measurements
should include samples of interstitial water from sediments as well
as samples from the water column. Sensitive techniques for analysis
of these substances should be developed because very low concentra-
tions can have important effects upon the complexation of organo-
mercuric and other metallic species.

7. Light absorption and light scattering characteristics of
natural waters should be determined for the wavelength region 290-700


was measured by flameless AA spectrophotometry. Control experiments
established that the mercuric salt concentration was not decreased by
adsorption on the glass walls of the reaction vessel during the time
periods of the kinetic runs, Second-order rate constants were calcu-
lated as described above,

ACIDOLYSIS OF DIPHENYLMERCURY

Aqueous solutions of diphenylmercury (~ 10°® M) were prepared by
adding weighed amounts of diphenylmercury to water and stirring for
several days, Exact concentrations were determined by uv analysis,
An aliquot of standardized perchloric acid was added and the rate of
disappearance of diphenylmercury was determined by uv spectroscopy.°*

MERCURIC SALT CLEAVAGE OF DIPHENYLMERCURY

The procedure for the kinetic runs was essentially the same as
described above for dimethylmercury with the exception that the
mercuric perchlorate and diphenylmercury concentrations were lower
(107° to 10°1° M), Attempts to follow the reaction by ultraviolet
spectroscopy with 10°° M reactant concentrations were unsuccessful,
because complete cleavage occurred within a matter of seconds in
neutral and acidic media,

PHOTOCHEMICAL PROCEDURES

Solutions of organomercury compounds were irradiated by broad-band
(> 290 nm) and monochromatic (313 nm) light from the mercury lamp.
Light from the mercury lamp was filtered through a Pyrex sleeve for
the broad-band studies and through a Pyrex sleeve and 1.0 cm of 4
solution of 0,001 M potassium chromate in 2% aqueous potassium
carbonate to isolate the 313 nm line, Procedures for preparing
and degassing sample and actinometer tubes have been previously
described,°*

A valerophenone actinometer°® was used for the studies with 313 nm
light and a benzophenone-cis-1,3-pentadiene actinometer®® was used
for the broad-band irradiations, After irradiating the organo-
mercurial and actinometer solutions in parallel in the quantum
yield apparatus, quantum yields were calculated by comparing actual
yields of photoreaction in the solutions.°*

Sensitized photolyses were carried out with sufficient concentrations
of sensitizers to absorb > 99% of the light. Singlet oxygen was
generated by methylene blue sensitization.°° Quenching studies

were carried out under conditions where no light was absorbed by

the quenchers, Disappearance quantum yields for phenylmercurials
and sulfur-bonded methylmercury complexes were independent of the
extent of reaction to at least 35% of completion. All quantum
yields were determined at 25°C.

22


2[(CH3Hg)2S]

¥ [CH3Hg]

Figure A6.

ERS = 108M

CONDITIONS:

"ERS = 107M *
- ZS = 10M

ERS = 10M

Relative concentrations of (CHgHg)25 in systems
containing a high concentration of reduced
sulfur species.


The final solution is obtained by substituting equations 67-72 into
equation 60, eliminating the [OH ] term, and rearranging to give

(cHsHes} _ x, [s=] (73)
ElcHgHe] X+k, [S7]

where X= 1+ a + 2g [CHoHgs] +k [or] +

K,=RN
T+k GF)

KeERS 4, KyEPQ: 4g KEN 4 _ KeEArO
1+ (#] 1+x, Gt) 14+, I] 14+, fH)

and K, represents the ionization constant of water.

Using similar procedures, the fraction of methylmercury existing as
(CHgHg)2S was derived (eq 74).

2 [(CHsHg)2S} _ 2ko (CHgHes
x[Cislig] [sTM] xlcHsHg] (74)

((ciigHe)2$) [sTM]

107*-° 75
(cHyHgsTM 7° 7)

Kio =

Expressions derived for the relative concentrations of other methyl.
mercury complexes are shown below.

where 3 = HsHe'] _ —_1_ (76)
m (CHsHg] x+x, [s7]
(CHsHgC1] tery
a. = cl F3 77
© ([CHgHg] “ e) oD
+
(cH, HgNHoR"]) _ K, (RN) (F3) (78)

= ([CHgHg) 1+k, [a]

84


nm, In conjunction with solar radiation intensities, these data can
be used to calculate the penetration of sunlight into natural waters.

8. A general mathematical model is needed for photochemical
processes in natural waters. Such a model should be capable of
predicting integrated photolysis rates in natural systems based on
quantum yields, ultraviolet absorption spectra, solar radiation
intensities, and turbulence levels.

9. Applicability of current information on sensitization and quench-
ing of photoreactions is limited by lack of knowledge concerning the
photochemical properties of natural waters. Research is needed to
define the nature of the sensitization process, quantum yields, and
conditions under which sensitization occurs. These studies should
also attempt to identify chemical and physical properties of natural
waters that indicate the presence of potential sensitizers or
quenchers,

10. A mathematical model should be developed describing the aquatic
mercury cycle, The model should have sufficient detail to permit
determination of the relative importance of the various proposed
transport processes,


0.05

0.00

Figure A5,

CONDITIONS:
rs=10°M

ERS = 10M

XRS = 1077M
ERS = 108M

Relative concentrations of methylmercuric-thiol
complexes in systems containing a high concen-
tration of reduced sulfur species.

90


Table B3. LOGARITHMS OF FORMATION CONSTANTS FOR 18
METHYLMERCURIC COMPLEXES (CH,HgX) IN WATER

{CHgHgxX|
Ligand X%* Log, Teague”) OT
-F 9.37 (9.5)
-cl- 5.25 (5.45)
-Br 6.62 (6.7)
-" 8.60 (8.7)
-or 9.37 (9.5)
OCs Hg” @~ 6.5)
-OCOCH3TM @ 3.6)
-HPQ,*7 5.03
-HPO,°7 4.67
-$27 21.2
-SCHp CH OH” 16.12
-SR (Cysteine) (15.7)
-S0327 8.11
Sp 0377 10.90
-SCN" (6.1)
-NHg 7.60 (8.4)
~NH2 CH CHzNHa 8.25
-cN” 14.2

*values in parentheses taken from reference 15 (ionic strength =
0.5, temp. 25°C); other values from reference 18 (ionic strength =
0.1, temp. 20°C).

95


30 en

st A, CH3Hg-EGG ALBUMIN |
\ B. CHsHgSCH,COOH
' C. CH3HgSCHCH,0H

20 D. CH3HgSCH,CH(NH3t)COOH

| | | |
300 320 340 360 380 400
WAVELENGTH, nm

Figure 17. UV absorption spectra of methylmercury-thiol
complexes,

58


0.4 CoHsHg* or

-. -
PON en

Noon
(CgHs)2Hg \

CHsHgOH
0.2 - .
0.0 1 L | 1 |
300 310 320 330 340 350 360
WAVELENGTH, nm
Figure 19.

UV absorption spectra of phenylmercury compounds.

64


Although spectroscopic studies have shown that mercuric*” and organo-
mercuric®° ions, like the hydrogen ion, exist as hydrated forms in
water, these species will be represented for convenience as Hg’,
RHg*, and H*, respectively.

Mercuric®* and organomercuric!*"1° perchlorates, nitrates, and
sulfates are very ionic and are thus completely dissociated in
aqueous solution (eq 1 and 4). Other, less ionic mercury compounds
only partially dissociate at higher concentrations, the degree of
dissociation being concentration dependent.®* Thermodynamic
studies'*-?% 2 have shown that the dissociation tendencies of
mercuric (HgX2,) and organomercuric (RHgX) compounds depend upon the
nature of X as follows:

F > OCOCHg~ > HPQ,®” ~ CI” > Bx > NHg > OH > SR” > S?7

Thiol and sulfide compounds are particularly stable, i,e., they have
very low tendencies to dissociate.
Mercuric’* and organomercuric'®"1° ions react rapidly with water
(hydrolyze) to form corresponding hydroxides (eqs 2, 3, and 5).

These reactions are pH-dependent and will not occur to a significant
extent in acidic water (pH < 3-4). However, within pH ranges usually
found in natural waters (pH 5-9), the ions are almost completely
hydrolyzed (Figures 2-4). Experiments discussed later in this

report indicate that these hydrolysis equilibria have important
effects upon the rates of organomercury reactions in water.

Natural waters contain a variety of chemical species that can under-
go “ligand exchange reactions" with mercuric or organomercuric
complexes, Ligand exchange reactions involve reaction of mercuric
or organomercuric complexes with some chemical species, Y, to form
a new complex (eqs 6-8).

HgX2 + Y = HgXy + X (6)
HgXY +Y 2 HgYo + X 7”)
RHgX + Y = RHgY + X (8)

Ligand exchange reactions have been studied by a variety of techniques.
Early attempts to measure the rate of reaction of mercuric salts

with human serum mercaptalbumin, a protein that contains sulfhydryl
groups, were unsuccessful because of the rapid rate of the reaction.”°
Interest in the function of sulfhydryl groups in enzyme catalysis®*
prompted a number of kinetic studies of the reactions of protein-SH
groups with organomercuric salts, Pioneering studies by Boyer®®>


[CHsHgS~]

Figure A7.

ERS = 109M

[= = 10M

CONDITIONS:
mS =10M

Relative concentrations of CH,HgS in systems
containing a high concentration of reduced
sulfur species.


Table 1.

KINETIC DATA FOR ORGANOMERCURY LIGAND EXCHANGE REACTIONS

CHgHgX + a CHgHgy + X
7
x Y Log ke Log k, Reference
cr oH” 8.18 4.04 30
Bro ot 8.08 5.34 30
vr on 7.61 6.84 30
SCN” or 8.70 5.30 30
$0377 on” 6.70 5.40 30
oN” on” 4.2 8.9 31
cn” cr Ll 9.9 31
cnr $0377 2.8 8.8 31
CN” Sp0,°7 3.5 6.7 31

14


Table 6. INCREASING EASE OF ACID CLEAVAGE OF ARYL
AND ALKYL GROUPS IN ORGANOMERCURIALS

1. Methyl 5. Benzyl

2, Ethyl 6. Cyclohexyl
3. -Butyl 7. Phenyl

4, q-Propyl

Table 7, CALCULATED EQUILIBRIUM CONSTANTS FOR THE REACTION
OF DIMETHYLMERCURY WITH MERCURIC HALIDES
(EQ 30, R = CH,)”*

Salt K

HgCle 3.5 x 107+
HgBra 2.5 x 10°
Hele 3.4 x 10°

43


2[(CH3Hg)2S ]
YX [cHsHg]

Figure A3.

DRS =10°M

CONDITIONS:
rS=10M

IRS = 108M

Relative concentrations of (CHgHg)2S in

systems containing a low concentration
of reduced sulfur species.

88


CH3HgS~
[CH3Hg} tor

10

a SPECIES CONC, M
(CHsH9] ,., <1?
90 L [HrPo.7] + (HPO=] 0
{ci7] 10"
(NHa*] + [NH3] 0"?
80 [| [RNH,*) + [RNH2} 10°
[RSH] + [RST] 10?
[ArOH] + [ArO7] 10s
70 F [hes] + [SH] + (S=] xs
ce
so ob
40
30 fF
20 OF
10
0 : L
1 2 3 4 5 6 7 8 9 10
pH

Figure 6. Relative concentrations of CH,HgS in systems
containing multiple chemical species.

26


oT

UOT OFAIndIew Jo stskjTozphy ay Jo souspuodep yd *z eandty
ray
4
a
x
2
D> 7 7
=x
4
+ =x
x al
oO
D>.
=x ee al
D>
x=
N
it | i
2
Oo & 4 a) @ s

NOI D1YNIYIW LNIDYId


CH HgOH
(pH 10.0)

CH3Hgt (pH 2.0)

(CH3)gHg (pH 7.0)

S

200

Figure 15.

220 240 260 280 300
WAVELENGTH, nm

UV absorption spectra of dimethylmercury,
methylmercuric ion, and methylmercuric
hydroxide in water.

54


No.

16

17

18

19

20

UV absorption spectra of methylmercuric sulfide
ion and bis(methylmercuric) sulfide . .

UV absorption spectra of methylmercury-thiol

complexeS . + 6 ee ee eee ee

Effect of quencher, 2,4-hexadien-1l-ol, upon
quantum yields for photocleavage of methylmercury-

thiol complexes .... s+ ees

UV absorption spectra of phenylmercury compounds

Quenching of diphenylmercury photolysis by

cis-1,3-pentadiene in acetonitrile

58

62

64

69


12:32:55 4/05/7323

ILZ4D (MEHGFR?

WLIST
de Bl: GET LIET(C), C203, 04,0509, 010-C1195
tel GET LISTCHB, LAST. DELTs XI
a. PUT Liste "Os
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Bel PUT LIST(' CHEPG4-)+CHPO4: 209s
ae PUT _LISTC"’(CL-d=",C2)3
hel PUT LISTC'(NH4424CiHG)=", C1093
4ee PUT LISTC'(CAR2H)+¢ et, Cl tre
Se PUT LISTC' CRUH3+)+4CRNH22=",C03)3
6. PUT LIST(' (RSH)+(RS- 2 Cure
Te PUT LISTC'TOTAL MEKG CO@NC=" CS)5
be PUT LISTC' "3
9.
10. St TO LASTS we
Me PUT LISTC('P| LEGIOCIZC6)D3
126 PUT LISTC' "35
1a. Ca=.1E-103
15. CTS 60 G*Cl/(~ LES3+C6e% 246 LELO*CE+6. 35
16. PUT LISTC*CS=)=',C0795
V7s IF C7=0 THEN C8=0;
20. Z2 1+ 10#e- 463/054 10%¥5.254C2F
21. ZI=LO**B. 25" C37 14+10%*10%C6) 4 10*¥ 166 12CA/ C14 10**9. S2*C
2 6
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623 “c
23. Z4=1Oes6. SCL IS C141 O%e9e BXCE DF
25. IT: ZO=Z+Z 142242442" 10416. 3xCBS
306 O#e21. 24CT/ (2341044214 QCTIT
35. 14*C53
366 IF FI=0 THEN GS Te Bes
37 IF ABS((C8-CSP)/CS)<DELT THEN G2 TB C2MPs
40. CB=EXP( (LEG CCED+LEGC CEP) 9/293
45. GO TO 173
50. COMP: FOs2"10%4-4.9"CS#*2/ (C7 *C5 95
SS Ba: SCZAZ14Z 247 4410442). 2HCT+ 241 016+ 34CT IS
S56 OHS. QTHF Se. 13/063
576 OF#5« O3KFSKCO/ (1410464 79"C6I5
56. OF=7e6*F3*CIO/ (1410849. 42%T6)F
59. If Fl=
60. Fa=10845
656 FS=10*4S. B5#COMF3/C 14.1 E11 C603
66. FI1= 105% 6. S#F3"C11/C1+104*9-8*C6 25
70. Fi O%#* 16. 124F3*C4/ 0] 4+10=49- 526063
MMe Fi +P Qt 347 as S+F6+T7+FO*F 1O+FI 1s
75. PUT LISTC'FRACTIGN FLD
BQ. PUT LISTC' FRACTION
es. PUT LISTC'FRACTICN
6. PUT LISTC' FRACTION
87. . PUT LISTC'FRACTIG
8B. PUT LIST('FRACTIA
90. PUT LISTC" cT1a
Si. PUT LISTC(' FRACTION
oS. PUT LISTC*FRACTIAN »
100. PUT LISTC* FRACTION MEHGSR=', F693
101. PUT ‘SUM OF FRACTION
102. PUT LISTC' '93
103. PUT LISTC* "93
104. C62 C64 1G RHXS
1056 END 3
110. co TS Bis
FLeG?

TINE 12:39:523 TIME US:

CPU OO100:05s TER 09:06:18 PAGE vur05:367_

Figure Al. Computer program used for calculations in Appendix A,

86


degradation or transformation of a pollutant. Also determination of
rate constants at different temperatures will allow calculation of
activation parameters, which are indicative of the reaction pathway,
and will allow extrapolation of the rate constant to other tempera-
tures. It is desirable to carry out these studies in water when
possible because of pronounced solvent effects on rates and products.

Since rate constants are sometimes difficult to relate to the life-
time of a reactant, the half-life (ti) expression is often employed.”>
The half-life is defined as the time required for the concentration
of a reactant to be reduced to one-half its initial concentration.

The ti expression is especially convenient for first-order’?
reactions (eq 19). The rate of disappearance of reactant A is given
by differential equation 20.

ky

A> B (19)
d

- TAI og fay (20)
dt

The tz, expression derived from the integrated form of equation 20 is
given in equation 21,

_ 0.693

7 (21)

t
%

The tz is independent of the concentration of reactant A and
dependent only on the magnitude of the rate constant (k,).

The tz expression for second~order’+ reactions requires a statement
of reactant concentrations, The disappearance of reactant A (eq 22),
given by equation 23, includes the concentrations of both reactants.

ke
A+B-> C (22)
- Wd. [atts) (23)
dt

The derivation of the half-life expression by integration of equation
23 gives an equation too complex to be of practical value. However,
in the special case of equal initial concentrations of reactants,

the ty expression may be expressed as

31


APPENDIX B

EQUILIBRIUM CONSTANTS FOR FORMATION, DISSOCIATION, AND HYDROLYSIS
OF MERCURY SPECIES

Table Bl. LOGARITHMS OF FORMATION CONSTANTS
FOR Hg°* COMPLEXES IN WATER AT 25°C

{Hex, |
Ligand x* Log Tae] OCP
“cl” 13,2°
-Bro 16.8°
Tv 23,8°
- OH 21.7°
-NHg 17.4
-NH2R (Histidine) 21.2
-SK (Cysteine) 41.0
-SCNr 17.48
~CN" 34.7

*Sign to right denotes charge of unattached ligand,
®Taken from reference 15,

“Tonic strength =
‘Tonic strength = 2
“Tonic strength = 0.15.
tonic strength = 1
Etonic strength = 0

93


found to decompose during chromatography,’® DMM was found by combined
GC-MS to be stable to gle conditions employed. The kinetics of
acidolysis were therefore determined by following the disappearance
of DMM by this technique. The methyl~ and phenylmercuric salts were
shown to be stable under the acidolysis reaction conditions, as

; 0
reported earlier .*° #7

A plot of the second-order rate expression”) (Figure 11) indicated
that the dimethylmercury acidolysis reactions obeyed second-order
kinetics through 50-75% of the reaction. The variation in slope
with temperature illustrates the temperature dependence of the
reaction,

The kinetic data for HCl, HBr, and HC10, in Table 2 demonstrates
that the second-order rate constants are independent of the nature
of the acid, In water at the low reactant concentrations indicated,
strong mineral acids are completely dissociated and no anion depen-
dence is observed, For HCl the extrapolated rate constant at 25°C
is 7.33 X 10° L/mole sec.

Verification of the kinetics was obtained by employing a 20-fold
excess of HCl compared to DMM and determining the pseudo-first-order
rate constant at 40°, The second-order rate constant was obtained
by dividing the pseudo-first-order rate constant by the HCl concen-
tration. The constant so obtained, 4.99 + 0.13 x 10°* 4/mole sec,
agrees with the extrapolated value of 5.08 x 10°* 2/mole sec obtained
under second-order conditions.

The gle method of determining the reaction velocity was verified by
a conductance method®® * ° employing a large molar excess of DMM compared
to HCl. The reaction was followed through 25% to 50% completion by
monitoring the decrease of conductance with time. This was possible

because of the high specific conductance of H’ compared to DMM or
MMX.

The conductometrically determined rate constants for HCl, HC10,, and
HNO, are given in Table 3. The calculated second-order rate constants
agree with the values reported in Table 2 (gle method) for HCl and
HC1O, .

The above results definitively show the reaction to be first-order
with respect to both DMM and acid (an electrophilic substitution
reaction (Sg 2)). Based on stereo-chemical studies, these reactions
have previously been thought to proceed by a four-centered or
similar type mechanism (Sg i), as shown in equation 27,

R
R---Hg
K-Hg-R ' : R R
+ 7 ' > to + f (27)
HCt ' a Hg
-- CL
cL

34