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The corrosion of silver by atmospheric sulfurous gases
CorrosionScience,Vol. 25, No. 2, pp. 133-143,1985 Printedin Great Britain
0010-938X/85$3.00+ 0.(~J © 1985PergamonPressLtd.
T H E C O R R O S I O N OF SILVER BY A T M O S P H E R I C S U L F U R O U S GASES J. P. FRANEY, G. W. KAMMLOTTa n d T. E. GRAEDEL AT&T Bell Laboratories, Murray Hill, NJ 07974, USA Abstraet--Polycrystalline silver has been exposed to the atmospheric gases H2S , O C S , CS 2 and SO2 in humidified air under carefully controlled laboratory conditions. OCS is shown to be an active corrodant while CS2 is quite inactive. At room temperature, the rates of sulfidation by H2S and OCS are comparable, and are more than an order of magnitude greater than those of CS2 and SO2. It appears that OCS is the principal cause of atmospheric sulfidation of silver except near sources of H2S where high concentrations may render the latter gas important. At constant absolute humidity, the sulfidation rate of silver by both HzS and OCS decreases from 20 to 40°C and then increases to 40 to 80°C. This behavior is interpreted as indicating a decrease in water-enhanced sulfidation as the relative humidity is reduced, followed by dominance of sulfidation by a water-independent process with strong positive temperature dependence. The initial sulfidation of silver by 3.9 + 0.2 ppm H2S in humidified air at 22°C has been studied in detail. The data are consistent with an initial stage of sulfidation involving rapid attack by HeS at surface defect sites. As these corrosive products spread and merge, diffusion of silver to the surface will be impeded. In agreement with this picture, the results show that the fraction of H2S molecules striking the surface that become incorporated into the sulfide film drops sharply from ~3 × 10-~ (at t = 5 min) to ~1 × 10-8 (at t = 50Oh). INTRODUCTION THE ATMOSPHERIC sulfidation of silver has b e e n recognized for a long time a n d r e m a i n s a p r o b l e m today, b e i n g d e m o n s t r a t e d to l a y m a n a n d scientist alike by the t a r n i s h i n g of electrical e q u i p m e n t , j e w e l r y a n d h o u s e h o l d silverware. T h e sulfide film that forms o n silver m a y have a sufficiently low c o n d u c t i v i t y that use as an electrical c o n t a c t is unsatisfactory. 1-3 T h u s there is wide i n t e r e s t in u n d e r s t a n d i n g the principal agents of silver sulfidation, a n d the m e c h a n i s m s by which they o p e r a t e . W e b e g i n by s u m m a r i z i n g the c u r r e n t state of k n o w l e d g e . Silver sulfidizes u p o n e x p o s u r e to a wide variety of gaseous s u l f u r - c o n t a i n i n g c o m p o u n d s . P e r h a p s the most extensive investigations of sulfidation have utilized h y d r o g e n sulfide (HES) as the reactive agent. 1,3-15 It has b e e n d e m o n s t r a t e d that HES c o n c e n t r a t i o n s as low as 0.2 parts per billion (ppb) are sufficient to cause sulfidation,11 a n d that high relative h u m i d i t i e s m a r k e d l y accelerate the process. 1'4'5A°'11'15E x t e n sive e x p e r i m e n t s have also b e e n p e r f o r m e d with flowers of sulfur (Ss). 3,s'16-1s C o n c e n t r a t i o n s as low as 30 p p b have b e e n s h o w n to sulfidize silver.17 Silver also reacts readily with a variety of m e r c a p t a n s a n d o t h e r o r g a n i c s u l f u r - c o n t a i n i n g species. 19 I n c o n t r a s t to its reactivity with r e d u c e d sulfur c o m p o u n d s , the sulfidation of silver by sulfur dioxide (SO2) is very slow, l's either dry or in the p r e s e n c e of water vapor. 8 Manuscript received 28 July 1984. 133
134
J.P. FRANEY,G. W. KAMMLOTTand T. E. GRAEDEL
Unlike some other metals, silver does not naturally form a surface oxide.11'20'21 Exposure to reduced sulfur compounds or to flowers of sulfur produces a film of silver sulfide (Ag2S). 8'12 Long exposures produce small amounts of a second component, which may be either silver oxide (Ag20) or silver sulfate (AgESO4).3'8'12 These compounds are also seen on silver that has been exposed in field environments. 22 Several researchers have examined the kinetics of silver sulfidation. For HES the rate at moderate exposure is parabolic, 13 becoming linear for severe exposures. 3'6 For flowers of sulfur the rate is linear throughout a very wide range of exposures. 3'12'17'18These data have been interpreted to mean that gas phase diffusion is the controlling process in sulfidation except for moderate H2S exposures, where the rate of the surface reaction controls the sulfidation rate. 13 In the light of the sulfidation research discussed above, it is of interest to consider the atmospheric concentrations of sulfur compounds. On a global basis, the most abundant is carbonyl sulfide (OCS),23'24 a compound whose effects on silver have not yet been assessed, although it is known to corrode copper readily. 24 The next most abundant gas is SO2, which is produced principally by the combustion of sulfurcontaining fossil fuels. SO2 thus tends to be concentrated downwind of power plants and industrial operations. H2S is less abundant, being emitted by widely dispersed sources of lower intensity such as marshes, volcanoes, crude oil processing, etc. Mercaptans and sulfides are only occasionally present and then only in very low concentrations. 25'26Free sulfur is rarely present in the atmosphereY As these remarks illustrate, silver sulfidation is a significant and ubiquitous problem. Sulfidation is known to be caused by a variety of reduced sulfur compounds and by free sulfur. None of these species are abundant in the atmosphere. Sulfur dioxide, a more common atmospheric constituent, has a negligible effect on silver sulfidation. No investigations have thus far been performed to study the effects on silver of carbonyl sulfide, the atmosphere's most abundant sulfur species. The kinetics of sulfide film growth have been determined only at rather severe exposures, except for free sulfur (known to be rare in the atmosphere). The controlling processes at mild exposures are virtually uninvestigated. In the work reported in this paper, we attempt to remedy some of these uncertainties. We have conducted the first laboratory exposures of silver to carbonyl sulfide, deriving kinetic data in the process. We have exposed silver to hydrogen sulfide over a wide range of conditions, including those typical of benign and severe field environments. We have confirmed earlier work on the relative inertness of silver to sulfur dioxide. Finally, we have exposed silver to carbon disulfide (CS2), an atmospheric gas of low abundance but moderate reactivity and a possible precursor to carbonyl sulfide. The results of these experiments are described below. EXPERIMENTAL
METHOD
Samplepreparation The silver samples used in this study were 99.999% pure. They were prepared by machining 1 cm z coupons from flat stock. The coupons were wet polished with 600-grit A1203 paper, followed by Linde C and Linde A aluminum oxide polish, using the method of Davis and Louzon.27 The resulting coupons had a mean surface roughness of ~<0.3 ttm, as measured from scanning electron microscope (SEM) micrographs, and undetectable (<0.1 nm; see below) amounts of sulfur contamination.
Corrosion of silver by atmospheric sulfurous gases
135
Generation, calibration and measurementof corrosive atmospheres The exposures were carried out in a system consisting of three modules: a gas generation system, an exposure facility and a monitoring and data reduction system. The system and its calibration have been described in detail elsewherefl 8'29For H2S, OCS and SO~, all of which are gases at room temperature and are available in pressurized form, the corrosive gas is generated by pressurizing a length of teflon tubing with the gas and allowing it to permeate through the tubing into a stream of carrier gas. Concentrations are adjusted by varying the length of tubing, the pressure of corrosive gas in the tubing and the flow rate of carrier gas around the tubing. The mixed gas is humidified as desired by passing a portion of the gas stream through a water bubbler. The mixed gas then flows into the exposure chamber, within which temperature and relative humidity are monitored continuously. As the mixed gas leaves the exposure chamber, it enters a sulfur gas detection system which continuously monitors the sulfur gas concentration. As the result of establishing and maintaining a detailed system of calibrations, the measured H2S, OCS and SO 2 concentrations are accurate to _+5% throughout the experiments. The fourth atmospheric sulfur-containing species included in the experiments, carbon disulfide (CS2), is a liquid at room temperature. Gas dilution techniques are therefore not suitable for generating a mixture of CSz in humidified air. We performed a simple exposure experiment as follows: a mixture of water and CS2 was prepared in a beaker. The odor of CS2 at the beaker rim was detectable but not overpowering. Since the odor threshold for CS 2 is 0.9 ppm, 3° we estimate the CS2 concentration within the beaker as - 1 0 ppm. Silver samples prepared for exposure were then placed within a small Petri dish which was floated on the mixed liquid, and the beaker was capped. After the designated exposure period, the samples were removed for study. While crude, this procedure for CS: proved satisfactory since the effects of CS 2 on silver were shown to be very slight (see below). Had they been severe, our calibration of the CS_~ concentration would have had to have been made with a precision similar to that employed for the other sulfurous gases.
Sulfidefilm analysis The morphology of the silver sulfide films was studied with a Kent-Cambridge 2A SEM equipped with a solid state X-ray detector and a multichannel analyzer. The thicknesses of the corrosion films were determined by energy-dispersive X-ray analysis (EDXA). Following the technique developed for cuprous sulfide films, 3J thickness standards were prepared by vapor deposition of silver films on silicon substrates. The samples were cleaved perpendicular to the surface and their thicknesses measured. They were then exposed to HzS in humidified air for a period long enough to completely sulfidize the films. The films were floated off the substrates, mounted on bulk silver substrates, and their EDXA sulfur and silver X-ray intensities measured. If sulfur concentration gradients in the exposed silver samples are disregarded and the sulfur assumed to be present solely as AgzS, 8'1z'22the intensity ratios may be converted to equivalent average sulfide film thicknesses. The result is a calibration curve of film thickness as a function of X-ray intensity ratio accurate to 20% or _+10nm, whichever is smaller.
EXPERIMENTAL
RESULTS
Relative sulfidation of silver by atmospheric sulfurous gases Silver coupons were exposed to humidified air containing trace concentrations of o n e o f t h e a t m o s p h e r i c s u l f u r g a s e s H2S, O C S , CS2 o r S O z. E x p o s u r e t i m e s v a r i e d f r o m as s h o r t as 5 r a i n t o as l o n g as 5 w e e k s . U p o n r e m o v a l f r o m t h e m u l t i p o r t chamber the silver sulfide on each of the samples was determined by EDXA. At least three individual measurements of each sample were made. T h r e e s e t s o f s a m p l e s w e r e e x p o s e d t o H2S a t r o o m t e m p e r a t u r e a n d h i g h h u m i d i t y , t h e s e t s d i f f e r i n g i n H2S c o n c e n t r a t i o n b y as m u c h as 4 0 % . T h e r e s u l t s a r e s h o w n in Fig. 1, w h e r e t h e m e a s u r e d s u l f i d e film t h i c k n e s s e s a r e p l o t t e d a g a i n s t t o t a l exposure (i.e., the product of exposure time and corrosive gas concentration). The c o r r e l a t i o n c o e f f i c i e n t f o r t h e c o m b i n e d d a t a is 0 . 8 7 .
136
J.P. FRANEY,G. W. KAMMLOTTand T. E. GRAEDEL io 2
H2S
suLfidation of silver
c
¢0 .c i0 ~ E
g <
I
I
1
I
I0 0
I01
102
103
Total
exposure ( p p m h - I ]
Fl6. 1. The thickness of the Ag2S corrosion film as a function of total exposure. All exposures were made with RH = 92 + 4%, t = 21 _+2 °C. H2S concentrations were stable within 2% for each exposure; the concentration range for the entire sample set was 2.66-3.64 ppm. The line is a least-squares fit to the logarithmsof the data. The error bars on the film thickness measurements are indicated on this and followingfigures,
The thicknesses of sulfide films formed on silver at room temperature by OCS are shown in Fig. 2. These data are the first to be presented for the OCS-silver system. The most important information communicated by the figure is that silver is sulfidized relatively rapidly by OCS. This consequence of the presence of OCS in the atmosphere has not previously been appreciated. The figure shows data for experiments in which the OCS concentrations differed by a factor of ten; within the limits of our measurements, the sulfidation appears to be a function of total OCS exposure rather than exposure time. The sulfidation of silver by SO2 and CS2 proved to be much less rapid than was the case for either H2S or OCS at equivalent concentrations. In Fig. 3 we display our SO2 data along with the sulfidation domains of H2S and OCS. These domains represent the areas bounded by the 90% confidence limits for the least-squares fits of Figs 1 and 2. 32 It can be seen that within these confidence limits there is no difference between the silver sulfidation rates of H2S and OCS. No sulfur was detected on the sample for the CS2 exposure; we therefore have a single upper limit for CS2 sulfidation. Notwithstanding the limitations of the data, it is clear that the sulfidation rate for silver exposed to either SO 2 or CS2 is at least one order of magnitude less than that for either H2S or OCS. The silver sulfidation produced by the four gases may be compared if we utilize typical atmospheric concentrations of the sulfurous gases and recognize that sulfidation of materials often scales with total exposure, at least for H2 S33 and O C S . 24 In Table 1 we compute the time required for the formation of 5 nm
Corrosion of silver by atmospheric sulfurous gases OCS
sulfidation
137
of silver
io 2
c
i0 I
== .c 4J
I0 o
Io-' i0 -I
I
1
1
I
i0 0
i01
10 2
i0 3
Total exposure (ppm h -=)
FIG. 2. The thickness of the Ag2S corrosion film as a function of exposure time for exposures at two different OCS concentrations. All exposures were made with R H = 92 + 4%, t = 21 °C. The line is a least-squares fit to all the data. IS],2.5 _+ 0.05 ppm OCS. ©, 0.26 _+ 0.02 ppm OCS.
films of Ag2Sin near-source and background field environments. (The values for CS2 assume that its silver sulfidation rate is proportional to that of SO2, the upper limit rate at a total-exposure of 640 ppm h-] serving to establish that proportionality.) The tabulated results show that H2S is the most potent of the gases in the near-source regime, while OCS fills that role in areas well removed from sources. The relative effects of SO2 are not large, and those of CS2 are negligible.
Temperature dependence of silver sulfidation Silver sulfidation experiments over the temperature range 20-80°C were conducted for H2S and OCS, the two most active sulfidizing gases. This temperature TABLE 1.
TIME REQUIRED TO FORM 5 n m A g 2 S FILMS IN THE AMBIENT ATMOSPHERE
Corrosive gas
formfilm(ppbh -])
H~S OCS CS 2
102 103 >107
SO~
7 × I 0 ~'
Exposure to
Ambient concentration (ppb)* Nearsource Background 5 0.5 0.3 I00
Time to form 5 nm film (days)
Nearsource
Background
0.03 0.5 0.03
I 80 >1.4 x l0 n
140 80 >1.4 × 107
0. I
3 x 103
3 x 107
* Sources for these typical concentrations are given in Reference 24.
138
J.P. FRANEY,G. W. KAMMLOTTand T. E. GRAEDEL Sukfidation
of
silver
102
i01
== o
E
i0 0 -
/ io-I
10-1
SO2J
+CS2
I
I
I
I
I
I00
I01
102
i0 5
104
Total
exposure ( p p m
h -t)
FIG. 3. The thickness of the Ag2Scorrosion film as a function of total exposure to H2S, OCS, SO2 and CS2, All exposures were made with RH = 92 _+4%, t = 21 °C. The H2S and OCS domains are discussed in the text.
range encompasses that likely to be encountered by silver exposed to the atmosphere in typical field uses. The results are presented in Fig. 4 for HzS and Fig. 5 for OCS. In each case the sulfidation rate is relatively high at room temperature, decreases as temperature is increased, and then increases as temperature is further increased. This behavior can be seen more readily in Fig. 6, where the thickness of the sulfide film is plotted as a function of temperature for a total exposure of 100 ppm hr- ~.
The initial stage of silver sulfidation under atmospheric conditions Since it is known that H2S is a significant factor in the atmospheric sulfidation of silver, we have performed short-term exposures of silver to H2S in order to examine the initial sulfidation process in more detail. These experiments used the capability of rapid sample insertion and removal in the multiport chamber 28to study sulfide film growth in the 0-30 nm region. We then use the measured film thicknesses to derive values of the incorporation coefficient y, i.e. the fraction of H2S molecules striking the surface which result in the incorporation of sulfur into the film, using the relationship 29
1.2 y-
[H2S]t,
(1)
Corrosion of silver by atmospheric sulfurous gases H2S on Ag (21°C)
139
HzS on Ag (42°C)
102
IO'
i0 °
Ec c ~c
I
IO-r
_u
I
I
I
I
I
I
I
+,
_ HzS on Ag (80°C)
E
~>
Fitted Lines
80°C
iO 2
i01
.J" i00 -
iO-~ i0 -j
I
I
I
I
i0 0
I0 r
10 2
i0 3
Total
i0 -I
I
J
I
I
i0 0
i01
10 2
10 3
exposure (ppm h -r)
FIG. 4. The thickness of the Ag2S corrosion films as a function of total exposure to H2S at different temperatures. All experiments were conducted at an absolute humidity of 18 4- 1 torr (94% RH at 21 °C).
where • is the equivalent Ag2S film thickness (nm), [ H 2 S ] is the average concentration of the corrosive gas, t is the exposure time (s) and the constant (1.2) applies to a film of Ag2S. The results are shown in Fig. 7. From an initial relatively stable value of 3, - 3 x 10 - 6 at t = 5 min, the coefficient decreases to a value of 3' - 1 x 10 -s at t = 500 h. The figure also indicates the exposure at which 10 equivalent monolayers of Ag2S were formed. 3' begins a rather sharp decline after the formation of the first few equivalent monolayers. O t h e r researchers °11 have observed corrosive 'mounds' at early stages of sulfidation. We infer that an average coverage of ten monolayers of sulfide is approximately the point at which the mounds of Ag2S coalesce into a relatively cohesive sulfide film, impeding the diffusion of silver ions and thus retarding further film formation.
140
J.P. F'RANEY,G. W. KAMMLOTTand T. E. GRAEDEL OCS on Ag (21"C)
OCS on Ag ( 3 8 ° C )
102
p
/
IO t / p
I0 o
,9."
E
.X u .c 4-,
s
s
i
I
i0 -[
OCS on Ag (80"C)
E
Fitted
I
I
I
Lines
g o
i 0 2 --
<
iO r -•
80"C/
/
//
/~,/'380C
lO o -- ~ / ¢
/ /
tO - I i0 o
I I0 j
I
[
10 2
10 3
Total
i0 °
exposure
(ppm
I
I
I
iO t
i0 2
I0 3
h -I)
FIG. 5. The thickness of the Ag 2S corrosion films as a function of total exposure to OCS at different temperatures. All experiments were conducted at an absolute humidity of 18 + 1 torr (94% RH at 21 °C).
CONCLUSIONS The results presented in the previous section provide information on the sulfidation of silver by different gases, the temperature dependence of sulfidation and its initial rate. In this section, we use the results to draw conclusions about silver sulfidation in the atmosphere. The data presented in Fig. 3 provide, to the best of our knowledge, the only directly comparable information on the relative susceptibility of silver to sulfidation by different sulfur-containing atmospheric gases. By combining these data with typical atmospheric concentrations of the gases, comparative ambient sulfidation information can be derived• To accomplish this, we read from Fig. 3 the total exposure necessary to produce a silver sulfide film of 5 nm. (Such a film visibly degrades the appearance of the silver and is thick enough to cause continuity
Corrosion of silver by atmospheric sulfurous gases
141
Temperature dependence 50
40
H2S
E v
3o u
E
g
2O ~10
18 -r I0
4 I 0
20
I 40
Surface H20 -[ . . . . . I 60
80
-
g
I00
Temoerature (*C) FIG. 6.
Ag2S film thickness at total exposures of 100 ppm h i H2 S or OCS, as functions of temperature. The values are derived from the fitted lines of Figs 4 and 5.
problems when silver is used as an electrical contact material.) Combining the exposure with typical atmospheric concentrations of the sulfurous gases gives the required times for film formation. As seen in Table 1, H2S is expected to be the principal sulfidizer near its sources, while OCS will be dominant elsewhere. Sulfidation due to SO2 and CS2 is unimportant. We have not investigated possible synergistic effects on sulfidation of the combined presence of two or more of the sulfurous gases or of other atmospheric trace gases. Such studies will be necessary if a complete picture of the atmospheric sulfidation of silver is to be assured. The temperature dependence shown in Fig. 6 indicates clearly that sulfidation is occurring by one mechanism in the lower temperature range and a second mechanism in the higher temperature range. An inverse temperature dependence on sulfidation has been seen with copper under similar exposure conditions, 29 where it was shown to be consistent with the amount of water adsorbed on the surface of the material at different relative humidities. The same process is doubtless also operating for silver, since silver sulfidation is relatively humidity-dependent. 1'4'5']°'uJ5 A film of three monolayers of water retains some of the properties of bulk water, 34 but as seen from Fig. 6, the average water film on silver becomes thinner than this above 30°C. 35 As the temperature is increased from this point, the positive temperature dependence of the dry sulfidation reaction becomes the dominant feature, as has been seen in higher temperature studies of silver sulfidation.36 The variation in incorporation coefficient with total exposure reflects the response of the silver surface to the hydrogen sulfide. When first exposed, the silver rapidly
142
J.P. FRANEY,G. W. KAMMLOTTand T. E. GRAEDEL 10-5
10-6
L~ 10-7
_
Ten equi aL monotoyers of Ag z S
I
I
I
I
I
i0 -I
i0 0
I01
102
103
I0 4
10-8
10-9
~
Exposure (ppm h-i)
FIG. 7. The H2S-incorporationcoefficienton silver as a function of total exposure. The exposure at which 10 monolayers (on average) of sulfide film is present on the surface is indicated. The lines are least squares fits to the data points.
adsorbs the sulfurous molecule, probably at grain boundaries or defects in the surface structure. 37 We show elsewhere 38 by experiments at constant temperature and varying relative humidity that both HaS and OCS sulfidation of silver are enhanced by the presence of water (as suggested by Fig. 6); the initial sulfidation stage must therefore involve water in some way. The initial regions of sulfidation serve as nucleating centers for the sulfide film, which can evolve into the corrosive 'mounds' seen in sulfidation experiments by other workers. Once the surface sulfide layer has coalesced, it can grow only through the reaction of hydrogen sulfide with silver ions that have diffused through the sulfide. The result is the very slow increase in sulfidation with time seen in silver sulfidation experiments at high exposure. 6 In summary, we have performed the first unified group of experiments on the sulfidation of silver by different atmospheric sulfur gases: H2S, SO2, OCS and CS2. These are the first experiments with the latter two compounds. OCS is shown to actively sulfidize silver, while the sulfidation by CS: is negligibly small. Temperaturedependent studies of silver sulfidation by H2S and OCS at constant absolute humidity demonstrate that in both cases an initial water-enhanced process is subsumed at higher temperatures by a water-independent sulfidation reaction with strong positive temperature dependence. H2S sulfidation at very low exposures demonstrates an initial rapid incorporation of H2S onto the surface of the silver, followed by a gradual transition into a regime in which sulfide film growth is controlled by the rate of silver
Corrosion of silver by atmospheric sulfurous gases
143
ion diffusion through the sulfide layer. Finally, an assessment of the sulfidation rates and typical atmospheric concentrations of the sulfurous gases demonstrates that OCS is probably the principal sulfidizing agent, except near sources of H2S and SO2. REFERENCES 1. J. A. LORENZEN,Proc. Inst. Environ. Sci. 110 (1971). 2. G.J. Russ, IEEE Trans. Parts, Materials and Packaging PMP-6, 129 (1970). 3. R. V. CHIARENZELLL Proc. 3rd Int. Syrup. Elect. Contact Phenom., p. 83-93. University of Maine, Orono (1966). 4. S. LILIENFELDand C. E. WHITE,J. Am. Chem. Soc. 52. 885 (1930). 5. J. DROTr, Ark. Kemi 15,181 (1959). 6. H. FISCHMEISTERand J. DROTr, Acta Metall. 7,777 (1959). 7. P. BACKLUND,B. FJELLSTROM,S. HAMMARB,~CKand B. MAIJGREN,Ark. Kemi 26,267 (1966). 8. W. H. ABBOTt, Electrical Contacts 1968, p. 53-54, IIT Research Inst., Chicago (1968). 9. W. E. CAMPBELLand U. B. THOMAS,Electrical Contacts 1968, p. 233-265. IIT Research Institute, Chicago (1968). 10. D. POPE, H. R. GIBBENSand R. L. Moss, Corros. Sci. 8,883 (1968). l 1. H. E. BENNETT, R. L. PECK,D. K. BURGEand J. M. BENNETr,J. appl. Phys. 40, 3351 (1969). 12. W.A. CROSSLANDand E. KNIGHT,Proc. 5th Int. Symp. Elect. Contact Phenom., Berlin, p. 324 (1970). 13. W. H. ABBott, IEEE Trans. Parts, Hybrids, Packaging PHP-10, 24 (1974). 14. D. SIMON, J. BARDOLLEand M. BUJOR, IEEE Trans. Components, Hybrids, and Manuf. Technol. CHMT-3, 13 (1980). 15. D. W. RICE, P. PETERSON,E. B. RIGBY,P. B. P. PHIPPS, R. J. CAPPELLand R. TREMOUREUX,J. electrochem. Soc. 128,275 (1981). 16. W. H. ABBOTt, Electrical Contacts 1969, p. 1-5. liT Research Institute, Chicago (1969). 17. B. T. REAGORand J. D. SINCLAIR,J. electrochem. Soc. 128,701 (1981). 18. W. H. ABBOt, Electrical Contacts 1970, p. 21-25. liT Research Institute, Chicago (1970). 19. J. D. SINCLAIR,J. electrochem. Soc. 129, 33 (1982). 20. W. E. CAMPBELLand U. B. THOMAS,Trans. electrochem. Soc. 76,303 (1939). 21. A. W. CZANDERNA,J. phys. Chem. 68, 2765 (1964). 22. D.W. R,CE, R. J. CAPPELL,W. KINSOLVINGand J. J. LASKOWSKI,J. electrochem. Soc. 127,891 (1980). 23. R. P. TURCO, R. C. WHITFEN, O. B. TOON, J. B. POLLACKand P. HAMILL,Nature283, 283 (1980). 24. T. E. GRAEDEL,G. W. KAMMLOTFand J. P. FRANEY,Science 212,663 (1981). 25. T. E. GRAEDEL,Chemical Compounds in the Atmosphere. Academic Press, New York (1978). 26. P. J. MAROULISand A. R. BANDY,Geophys. Res. Lett. 7,681 (1980). 27. M. E. DAvis and T. A. LouzoN, Metallography 13,195 (1980). 28. J. P. FRANEY,Corros. Sci. 23, 1 (1983). 29. T. E. GRAEDEL,J. P. FRANEYand G. W. KAMMLOTr,Corros. Sci. 23, 1141 (1983). 30. W. H. STAHL(Ed.), Compilation of Odorand Taste Threshold Values Data, ASTM DS-48. American Society for Testing and Materials, Philadelphia (1973). 31. G. W. KAMMLOVr,Appl. Spectroscopy 35,324 (1981). 32. M. KENDALLand A. STUART,Advanced Theory o f Statistics, Vol. 1, 4th edn. Macmillan, New York (1977). 33. J. P. FRANEY,T. E. GRAEDELand G. W. KaMMLOrr, in Atmospheric Corrosion (ed. W. H. A1LOR), pp. 383-392. McGraw-Hill, New York (1982). 34. P. B. P. PHIPPSand D. W. RlCE, in Corrosion Chemistry (eds. G. R. BRUBAKERand P. B. P. PnIpps) ACS Symp. 89, p. 235-261. American Chemical Society, Washington, D.C. (1979). 35. D. W. RICE, R. J. CAPPELL,P. B. P. PHIPPSand P. PETERSON,in Atmospheric Corrosion (ed. W. H. AILOR), p. 651--666. McGraw-Hill, New York (1982). 36. C.J. WARDE,J. CORrSn and C. D. O'BRIAIN,J. electrochem. Soc. 122, 1421 (1975). 37. V. A. PH,LLIPS,J. appl. Phys. 33,712 (1962). 38. T. E. GRAEDEL,J. P. FRANEY, G. J. GUALTIERI, G. W. KAMMLO'Iq"and D. L. MALM, submitted for publication.
0010-938X/85$3.00+ 0.(~J © 1985PergamonPressLtd.
T H E C O R R O S I O N OF SILVER BY A T M O S P H E R I C S U L F U R O U S GASES J. P. FRANEY, G. W. KAMMLOTTa n d T. E. GRAEDEL AT&T Bell Laboratories, Murray Hill, NJ 07974, USA Abstraet--Polycrystalline silver has been exposed to the atmospheric gases H2S , O C S , CS 2 and SO2 in humidified air under carefully controlled laboratory conditions. OCS is shown to be an active corrodant while CS2 is quite inactive. At room temperature, the rates of sulfidation by H2S and OCS are comparable, and are more than an order of magnitude greater than those of CS2 and SO2. It appears that OCS is the principal cause of atmospheric sulfidation of silver except near sources of H2S where high concentrations may render the latter gas important. At constant absolute humidity, the sulfidation rate of silver by both HzS and OCS decreases from 20 to 40°C and then increases to 40 to 80°C. This behavior is interpreted as indicating a decrease in water-enhanced sulfidation as the relative humidity is reduced, followed by dominance of sulfidation by a water-independent process with strong positive temperature dependence. The initial sulfidation of silver by 3.9 + 0.2 ppm H2S in humidified air at 22°C has been studied in detail. The data are consistent with an initial stage of sulfidation involving rapid attack by HeS at surface defect sites. As these corrosive products spread and merge, diffusion of silver to the surface will be impeded. In agreement with this picture, the results show that the fraction of H2S molecules striking the surface that become incorporated into the sulfide film drops sharply from ~3 × 10-~ (at t = 5 min) to ~1 × 10-8 (at t = 50Oh). INTRODUCTION THE ATMOSPHERIC sulfidation of silver has b e e n recognized for a long time a n d r e m a i n s a p r o b l e m today, b e i n g d e m o n s t r a t e d to l a y m a n a n d scientist alike by the t a r n i s h i n g of electrical e q u i p m e n t , j e w e l r y a n d h o u s e h o l d silverware. T h e sulfide film that forms o n silver m a y have a sufficiently low c o n d u c t i v i t y that use as an electrical c o n t a c t is unsatisfactory. 1-3 T h u s there is wide i n t e r e s t in u n d e r s t a n d i n g the principal agents of silver sulfidation, a n d the m e c h a n i s m s by which they o p e r a t e . W e b e g i n by s u m m a r i z i n g the c u r r e n t state of k n o w l e d g e . Silver sulfidizes u p o n e x p o s u r e to a wide variety of gaseous s u l f u r - c o n t a i n i n g c o m p o u n d s . P e r h a p s the most extensive investigations of sulfidation have utilized h y d r o g e n sulfide (HES) as the reactive agent. 1,3-15 It has b e e n d e m o n s t r a t e d that HES c o n c e n t r a t i o n s as low as 0.2 parts per billion (ppb) are sufficient to cause sulfidation,11 a n d that high relative h u m i d i t i e s m a r k e d l y accelerate the process. 1'4'5A°'11'15E x t e n sive e x p e r i m e n t s have also b e e n p e r f o r m e d with flowers of sulfur (Ss). 3,s'16-1s C o n c e n t r a t i o n s as low as 30 p p b have b e e n s h o w n to sulfidize silver.17 Silver also reacts readily with a variety of m e r c a p t a n s a n d o t h e r o r g a n i c s u l f u r - c o n t a i n i n g species. 19 I n c o n t r a s t to its reactivity with r e d u c e d sulfur c o m p o u n d s , the sulfidation of silver by sulfur dioxide (SO2) is very slow, l's either dry or in the p r e s e n c e of water vapor. 8 Manuscript received 28 July 1984. 133
134
J.P. FRANEY,G. W. KAMMLOTTand T. E. GRAEDEL
Unlike some other metals, silver does not naturally form a surface oxide.11'20'21 Exposure to reduced sulfur compounds or to flowers of sulfur produces a film of silver sulfide (Ag2S). 8'12 Long exposures produce small amounts of a second component, which may be either silver oxide (Ag20) or silver sulfate (AgESO4).3'8'12 These compounds are also seen on silver that has been exposed in field environments. 22 Several researchers have examined the kinetics of silver sulfidation. For HES the rate at moderate exposure is parabolic, 13 becoming linear for severe exposures. 3'6 For flowers of sulfur the rate is linear throughout a very wide range of exposures. 3'12'17'18These data have been interpreted to mean that gas phase diffusion is the controlling process in sulfidation except for moderate H2S exposures, where the rate of the surface reaction controls the sulfidation rate. 13 In the light of the sulfidation research discussed above, it is of interest to consider the atmospheric concentrations of sulfur compounds. On a global basis, the most abundant is carbonyl sulfide (OCS),23'24 a compound whose effects on silver have not yet been assessed, although it is known to corrode copper readily. 24 The next most abundant gas is SO2, which is produced principally by the combustion of sulfurcontaining fossil fuels. SO2 thus tends to be concentrated downwind of power plants and industrial operations. H2S is less abundant, being emitted by widely dispersed sources of lower intensity such as marshes, volcanoes, crude oil processing, etc. Mercaptans and sulfides are only occasionally present and then only in very low concentrations. 25'26Free sulfur is rarely present in the atmosphereY As these remarks illustrate, silver sulfidation is a significant and ubiquitous problem. Sulfidation is known to be caused by a variety of reduced sulfur compounds and by free sulfur. None of these species are abundant in the atmosphere. Sulfur dioxide, a more common atmospheric constituent, has a negligible effect on silver sulfidation. No investigations have thus far been performed to study the effects on silver of carbonyl sulfide, the atmosphere's most abundant sulfur species. The kinetics of sulfide film growth have been determined only at rather severe exposures, except for free sulfur (known to be rare in the atmosphere). The controlling processes at mild exposures are virtually uninvestigated. In the work reported in this paper, we attempt to remedy some of these uncertainties. We have conducted the first laboratory exposures of silver to carbonyl sulfide, deriving kinetic data in the process. We have exposed silver to hydrogen sulfide over a wide range of conditions, including those typical of benign and severe field environments. We have confirmed earlier work on the relative inertness of silver to sulfur dioxide. Finally, we have exposed silver to carbon disulfide (CS2), an atmospheric gas of low abundance but moderate reactivity and a possible precursor to carbonyl sulfide. The results of these experiments are described below. EXPERIMENTAL
METHOD
Samplepreparation The silver samples used in this study were 99.999% pure. They were prepared by machining 1 cm z coupons from flat stock. The coupons were wet polished with 600-grit A1203 paper, followed by Linde C and Linde A aluminum oxide polish, using the method of Davis and Louzon.27 The resulting coupons had a mean surface roughness of ~<0.3 ttm, as measured from scanning electron microscope (SEM) micrographs, and undetectable (<0.1 nm; see below) amounts of sulfur contamination.
Corrosion of silver by atmospheric sulfurous gases
135
Generation, calibration and measurementof corrosive atmospheres The exposures were carried out in a system consisting of three modules: a gas generation system, an exposure facility and a monitoring and data reduction system. The system and its calibration have been described in detail elsewherefl 8'29For H2S, OCS and SO~, all of which are gases at room temperature and are available in pressurized form, the corrosive gas is generated by pressurizing a length of teflon tubing with the gas and allowing it to permeate through the tubing into a stream of carrier gas. Concentrations are adjusted by varying the length of tubing, the pressure of corrosive gas in the tubing and the flow rate of carrier gas around the tubing. The mixed gas is humidified as desired by passing a portion of the gas stream through a water bubbler. The mixed gas then flows into the exposure chamber, within which temperature and relative humidity are monitored continuously. As the mixed gas leaves the exposure chamber, it enters a sulfur gas detection system which continuously monitors the sulfur gas concentration. As the result of establishing and maintaining a detailed system of calibrations, the measured H2S, OCS and SO 2 concentrations are accurate to _+5% throughout the experiments. The fourth atmospheric sulfur-containing species included in the experiments, carbon disulfide (CS2), is a liquid at room temperature. Gas dilution techniques are therefore not suitable for generating a mixture of CSz in humidified air. We performed a simple exposure experiment as follows: a mixture of water and CS2 was prepared in a beaker. The odor of CS2 at the beaker rim was detectable but not overpowering. Since the odor threshold for CS 2 is 0.9 ppm, 3° we estimate the CS2 concentration within the beaker as - 1 0 ppm. Silver samples prepared for exposure were then placed within a small Petri dish which was floated on the mixed liquid, and the beaker was capped. After the designated exposure period, the samples were removed for study. While crude, this procedure for CS: proved satisfactory since the effects of CS 2 on silver were shown to be very slight (see below). Had they been severe, our calibration of the CS_~ concentration would have had to have been made with a precision similar to that employed for the other sulfurous gases.
Sulfidefilm analysis The morphology of the silver sulfide films was studied with a Kent-Cambridge 2A SEM equipped with a solid state X-ray detector and a multichannel analyzer. The thicknesses of the corrosion films were determined by energy-dispersive X-ray analysis (EDXA). Following the technique developed for cuprous sulfide films, 3J thickness standards were prepared by vapor deposition of silver films on silicon substrates. The samples were cleaved perpendicular to the surface and their thicknesses measured. They were then exposed to HzS in humidified air for a period long enough to completely sulfidize the films. The films were floated off the substrates, mounted on bulk silver substrates, and their EDXA sulfur and silver X-ray intensities measured. If sulfur concentration gradients in the exposed silver samples are disregarded and the sulfur assumed to be present solely as AgzS, 8'1z'22the intensity ratios may be converted to equivalent average sulfide film thicknesses. The result is a calibration curve of film thickness as a function of X-ray intensity ratio accurate to 20% or _+10nm, whichever is smaller.
EXPERIMENTAL
RESULTS
Relative sulfidation of silver by atmospheric sulfurous gases Silver coupons were exposed to humidified air containing trace concentrations of o n e o f t h e a t m o s p h e r i c s u l f u r g a s e s H2S, O C S , CS2 o r S O z. E x p o s u r e t i m e s v a r i e d f r o m as s h o r t as 5 r a i n t o as l o n g as 5 w e e k s . U p o n r e m o v a l f r o m t h e m u l t i p o r t chamber the silver sulfide on each of the samples was determined by EDXA. At least three individual measurements of each sample were made. T h r e e s e t s o f s a m p l e s w e r e e x p o s e d t o H2S a t r o o m t e m p e r a t u r e a n d h i g h h u m i d i t y , t h e s e t s d i f f e r i n g i n H2S c o n c e n t r a t i o n b y as m u c h as 4 0 % . T h e r e s u l t s a r e s h o w n in Fig. 1, w h e r e t h e m e a s u r e d s u l f i d e film t h i c k n e s s e s a r e p l o t t e d a g a i n s t t o t a l exposure (i.e., the product of exposure time and corrosive gas concentration). The c o r r e l a t i o n c o e f f i c i e n t f o r t h e c o m b i n e d d a t a is 0 . 8 7 .
136
J.P. FRANEY,G. W. KAMMLOTTand T. E. GRAEDEL io 2
H2S
suLfidation of silver
c
¢0 .c i0 ~ E
g <
I
I
1
I
I0 0
I01
102
103
Total
exposure ( p p m h - I ]
Fl6. 1. The thickness of the Ag2S corrosion film as a function of total exposure. All exposures were made with RH = 92 + 4%, t = 21 _+2 °C. H2S concentrations were stable within 2% for each exposure; the concentration range for the entire sample set was 2.66-3.64 ppm. The line is a least-squares fit to the logarithmsof the data. The error bars on the film thickness measurements are indicated on this and followingfigures,
The thicknesses of sulfide films formed on silver at room temperature by OCS are shown in Fig. 2. These data are the first to be presented for the OCS-silver system. The most important information communicated by the figure is that silver is sulfidized relatively rapidly by OCS. This consequence of the presence of OCS in the atmosphere has not previously been appreciated. The figure shows data for experiments in which the OCS concentrations differed by a factor of ten; within the limits of our measurements, the sulfidation appears to be a function of total OCS exposure rather than exposure time. The sulfidation of silver by SO2 and CS2 proved to be much less rapid than was the case for either H2S or OCS at equivalent concentrations. In Fig. 3 we display our SO2 data along with the sulfidation domains of H2S and OCS. These domains represent the areas bounded by the 90% confidence limits for the least-squares fits of Figs 1 and 2. 32 It can be seen that within these confidence limits there is no difference between the silver sulfidation rates of H2S and OCS. No sulfur was detected on the sample for the CS2 exposure; we therefore have a single upper limit for CS2 sulfidation. Notwithstanding the limitations of the data, it is clear that the sulfidation rate for silver exposed to either SO 2 or CS2 is at least one order of magnitude less than that for either H2S or OCS. The silver sulfidation produced by the four gases may be compared if we utilize typical atmospheric concentrations of the sulfurous gases and recognize that sulfidation of materials often scales with total exposure, at least for H2 S33 and O C S . 24 In Table 1 we compute the time required for the formation of 5 nm
Corrosion of silver by atmospheric sulfurous gases OCS
sulfidation
137
of silver
io 2
c
i0 I
== .c 4J
I0 o
Io-' i0 -I
I
1
1
I
i0 0
i01
10 2
i0 3
Total exposure (ppm h -=)
FIG. 2. The thickness of the Ag2S corrosion film as a function of exposure time for exposures at two different OCS concentrations. All exposures were made with R H = 92 + 4%, t = 21 °C. The line is a least-squares fit to all the data. IS],2.5 _+ 0.05 ppm OCS. ©, 0.26 _+ 0.02 ppm OCS.
films of Ag2Sin near-source and background field environments. (The values for CS2 assume that its silver sulfidation rate is proportional to that of SO2, the upper limit rate at a total-exposure of 640 ppm h-] serving to establish that proportionality.) The tabulated results show that H2S is the most potent of the gases in the near-source regime, while OCS fills that role in areas well removed from sources. The relative effects of SO2 are not large, and those of CS2 are negligible.
Temperature dependence of silver sulfidation Silver sulfidation experiments over the temperature range 20-80°C were conducted for H2S and OCS, the two most active sulfidizing gases. This temperature TABLE 1.
TIME REQUIRED TO FORM 5 n m A g 2 S FILMS IN THE AMBIENT ATMOSPHERE
Corrosive gas
formfilm(ppbh -])
H~S OCS CS 2
102 103 >107
SO~
7 × I 0 ~'
Exposure to
Ambient concentration (ppb)* Nearsource Background 5 0.5 0.3 I00
Time to form 5 nm film (days)
Nearsource
Background
0.03 0.5 0.03
I 80 >1.4 x l0 n
140 80 >1.4 × 107
0. I
3 x 103
3 x 107
* Sources for these typical concentrations are given in Reference 24.
138
J.P. FRANEY,G. W. KAMMLOTTand T. E. GRAEDEL Sukfidation
of
silver
102
i01
== o
E
i0 0 -
/ io-I
10-1
SO2J
+CS2
I
I
I
I
I
I00
I01
102
i0 5
104
Total
exposure ( p p m
h -t)
FIG. 3. The thickness of the Ag2Scorrosion film as a function of total exposure to H2S, OCS, SO2 and CS2, All exposures were made with RH = 92 _+4%, t = 21 °C. The H2S and OCS domains are discussed in the text.
range encompasses that likely to be encountered by silver exposed to the atmosphere in typical field uses. The results are presented in Fig. 4 for HzS and Fig. 5 for OCS. In each case the sulfidation rate is relatively high at room temperature, decreases as temperature is increased, and then increases as temperature is further increased. This behavior can be seen more readily in Fig. 6, where the thickness of the sulfide film is plotted as a function of temperature for a total exposure of 100 ppm hr- ~.
The initial stage of silver sulfidation under atmospheric conditions Since it is known that H2S is a significant factor in the atmospheric sulfidation of silver, we have performed short-term exposures of silver to H2S in order to examine the initial sulfidation process in more detail. These experiments used the capability of rapid sample insertion and removal in the multiport chamber 28to study sulfide film growth in the 0-30 nm region. We then use the measured film thicknesses to derive values of the incorporation coefficient y, i.e. the fraction of H2S molecules striking the surface which result in the incorporation of sulfur into the film, using the relationship 29
1.2 y-
[H2S]t,
(1)
Corrosion of silver by atmospheric sulfurous gases H2S on Ag (21°C)
139
HzS on Ag (42°C)
102
IO'
i0 °
Ec c ~c
I
IO-r
_u
I
I
I
I
I
I
I
+,
_ HzS on Ag (80°C)
E
~>
Fitted Lines
80°C
iO 2
i01
.J" i00 -
iO-~ i0 -j
I
I
I
I
i0 0
I0 r
10 2
i0 3
Total
i0 -I
I
J
I
I
i0 0
i01
10 2
10 3
exposure (ppm h -r)
FIG. 4. The thickness of the Ag2S corrosion films as a function of total exposure to H2S at different temperatures. All experiments were conducted at an absolute humidity of 18 4- 1 torr (94% RH at 21 °C).
where • is the equivalent Ag2S film thickness (nm), [ H 2 S ] is the average concentration of the corrosive gas, t is the exposure time (s) and the constant (1.2) applies to a film of Ag2S. The results are shown in Fig. 7. From an initial relatively stable value of 3, - 3 x 10 - 6 at t = 5 min, the coefficient decreases to a value of 3' - 1 x 10 -s at t = 500 h. The figure also indicates the exposure at which 10 equivalent monolayers of Ag2S were formed. 3' begins a rather sharp decline after the formation of the first few equivalent monolayers. O t h e r researchers °11 have observed corrosive 'mounds' at early stages of sulfidation. We infer that an average coverage of ten monolayers of sulfide is approximately the point at which the mounds of Ag2S coalesce into a relatively cohesive sulfide film, impeding the diffusion of silver ions and thus retarding further film formation.
140
J.P. F'RANEY,G. W. KAMMLOTTand T. E. GRAEDEL OCS on Ag (21"C)
OCS on Ag ( 3 8 ° C )
102
p
/
IO t / p
I0 o
,9."
E
.X u .c 4-,
s
s
i
I
i0 -[
OCS on Ag (80"C)
E
Fitted
I
I
I
Lines
g o
i 0 2 --
<
iO r -•
80"C/
/
//
/~,/'380C
lO o -- ~ / ¢
/ /
tO - I i0 o
I I0 j
I
[
10 2
10 3
Total
i0 °
exposure
(ppm
I
I
I
iO t
i0 2
I0 3
h -I)
FIG. 5. The thickness of the Ag 2S corrosion films as a function of total exposure to OCS at different temperatures. All experiments were conducted at an absolute humidity of 18 + 1 torr (94% RH at 21 °C).
CONCLUSIONS The results presented in the previous section provide information on the sulfidation of silver by different gases, the temperature dependence of sulfidation and its initial rate. In this section, we use the results to draw conclusions about silver sulfidation in the atmosphere. The data presented in Fig. 3 provide, to the best of our knowledge, the only directly comparable information on the relative susceptibility of silver to sulfidation by different sulfur-containing atmospheric gases. By combining these data with typical atmospheric concentrations of the gases, comparative ambient sulfidation information can be derived• To accomplish this, we read from Fig. 3 the total exposure necessary to produce a silver sulfide film of 5 nm. (Such a film visibly degrades the appearance of the silver and is thick enough to cause continuity
Corrosion of silver by atmospheric sulfurous gases
141
Temperature dependence 50
40
H2S
E v
3o u
E
g
2O ~10
18 -r I0
4 I 0
20
I 40
Surface H20 -[ . . . . . I 60
80
-
g
I00
Temoerature (*C) FIG. 6.
Ag2S film thickness at total exposures of 100 ppm h i H2 S or OCS, as functions of temperature. The values are derived from the fitted lines of Figs 4 and 5.
problems when silver is used as an electrical contact material.) Combining the exposure with typical atmospheric concentrations of the sulfurous gases gives the required times for film formation. As seen in Table 1, H2S is expected to be the principal sulfidizer near its sources, while OCS will be dominant elsewhere. Sulfidation due to SO2 and CS2 is unimportant. We have not investigated possible synergistic effects on sulfidation of the combined presence of two or more of the sulfurous gases or of other atmospheric trace gases. Such studies will be necessary if a complete picture of the atmospheric sulfidation of silver is to be assured. The temperature dependence shown in Fig. 6 indicates clearly that sulfidation is occurring by one mechanism in the lower temperature range and a second mechanism in the higher temperature range. An inverse temperature dependence on sulfidation has been seen with copper under similar exposure conditions, 29 where it was shown to be consistent with the amount of water adsorbed on the surface of the material at different relative humidities. The same process is doubtless also operating for silver, since silver sulfidation is relatively humidity-dependent. 1'4'5']°'uJ5 A film of three monolayers of water retains some of the properties of bulk water, 34 but as seen from Fig. 6, the average water film on silver becomes thinner than this above 30°C. 35 As the temperature is increased from this point, the positive temperature dependence of the dry sulfidation reaction becomes the dominant feature, as has been seen in higher temperature studies of silver sulfidation.36 The variation in incorporation coefficient with total exposure reflects the response of the silver surface to the hydrogen sulfide. When first exposed, the silver rapidly
142
J.P. FRANEY,G. W. KAMMLOTTand T. E. GRAEDEL 10-5
10-6
L~ 10-7
_
Ten equi aL monotoyers of Ag z S
I
I
I
I
I
i0 -I
i0 0
I01
102
103
I0 4
10-8
10-9
~
Exposure (ppm h-i)
FIG. 7. The H2S-incorporationcoefficienton silver as a function of total exposure. The exposure at which 10 monolayers (on average) of sulfide film is present on the surface is indicated. The lines are least squares fits to the data points.
adsorbs the sulfurous molecule, probably at grain boundaries or defects in the surface structure. 37 We show elsewhere 38 by experiments at constant temperature and varying relative humidity that both HaS and OCS sulfidation of silver are enhanced by the presence of water (as suggested by Fig. 6); the initial sulfidation stage must therefore involve water in some way. The initial regions of sulfidation serve as nucleating centers for the sulfide film, which can evolve into the corrosive 'mounds' seen in sulfidation experiments by other workers. Once the surface sulfide layer has coalesced, it can grow only through the reaction of hydrogen sulfide with silver ions that have diffused through the sulfide. The result is the very slow increase in sulfidation with time seen in silver sulfidation experiments at high exposure. 6 In summary, we have performed the first unified group of experiments on the sulfidation of silver by different atmospheric sulfur gases: H2S, SO2, OCS and CS2. These are the first experiments with the latter two compounds. OCS is shown to actively sulfidize silver, while the sulfidation by CS: is negligibly small. Temperaturedependent studies of silver sulfidation by H2S and OCS at constant absolute humidity demonstrate that in both cases an initial water-enhanced process is subsumed at higher temperatures by a water-independent sulfidation reaction with strong positive temperature dependence. H2S sulfidation at very low exposures demonstrates an initial rapid incorporation of H2S onto the surface of the silver, followed by a gradual transition into a regime in which sulfide film growth is controlled by the rate of silver
Corrosion of silver by atmospheric sulfurous gases
143
ion diffusion through the sulfide layer. Finally, an assessment of the sulfidation rates and typical atmospheric concentrations of the sulfurous gases demonstrates that OCS is probably the principal sulfidizing agent, except near sources of H2S and SO2. REFERENCES 1. J. A. LORENZEN,Proc. Inst. Environ. Sci. 110 (1971). 2. G.J. Russ, IEEE Trans. Parts, Materials and Packaging PMP-6, 129 (1970). 3. R. V. CHIARENZELLL Proc. 3rd Int. Syrup. Elect. Contact Phenom., p. 83-93. University of Maine, Orono (1966). 4. S. LILIENFELDand C. E. WHITE,J. Am. Chem. Soc. 52. 885 (1930). 5. J. DROTr, Ark. Kemi 15,181 (1959). 6. H. FISCHMEISTERand J. DROTr, Acta Metall. 7,777 (1959). 7. P. BACKLUND,B. FJELLSTROM,S. HAMMARB,~CKand B. MAIJGREN,Ark. Kemi 26,267 (1966). 8. W. H. ABBOTt, Electrical Contacts 1968, p. 53-54, IIT Research Inst., Chicago (1968). 9. W. E. CAMPBELLand U. B. THOMAS,Electrical Contacts 1968, p. 233-265. IIT Research Institute, Chicago (1968). 10. D. POPE, H. R. GIBBENSand R. L. Moss, Corros. Sci. 8,883 (1968). l 1. H. E. BENNETT, R. L. PECK,D. K. BURGEand J. M. BENNETr,J. appl. Phys. 40, 3351 (1969). 12. W.A. CROSSLANDand E. KNIGHT,Proc. 5th Int. Symp. Elect. Contact Phenom., Berlin, p. 324 (1970). 13. W. H. ABBott, IEEE Trans. Parts, Hybrids, Packaging PHP-10, 24 (1974). 14. D. SIMON, J. BARDOLLEand M. BUJOR, IEEE Trans. Components, Hybrids, and Manuf. Technol. CHMT-3, 13 (1980). 15. D. W. RICE, P. PETERSON,E. B. RIGBY,P. B. P. PHIPPS, R. J. CAPPELLand R. TREMOUREUX,J. electrochem. Soc. 128,275 (1981). 16. W. H. ABBOTt, Electrical Contacts 1969, p. 1-5. liT Research Institute, Chicago (1969). 17. B. T. REAGORand J. D. SINCLAIR,J. electrochem. Soc. 128,701 (1981). 18. W. H. ABBOt, Electrical Contacts 1970, p. 21-25. liT Research Institute, Chicago (1970). 19. J. D. SINCLAIR,J. electrochem. Soc. 129, 33 (1982). 20. W. E. CAMPBELLand U. B. THOMAS,Trans. electrochem. Soc. 76,303 (1939). 21. A. W. CZANDERNA,J. phys. Chem. 68, 2765 (1964). 22. D.W. R,CE, R. J. CAPPELL,W. KINSOLVINGand J. J. LASKOWSKI,J. electrochem. Soc. 127,891 (1980). 23. R. P. TURCO, R. C. WHITFEN, O. B. TOON, J. B. POLLACKand P. HAMILL,Nature283, 283 (1980). 24. T. E. GRAEDEL,G. W. KAMMLOTFand J. P. FRANEY,Science 212,663 (1981). 25. T. E. GRAEDEL,Chemical Compounds in the Atmosphere. Academic Press, New York (1978). 26. P. J. MAROULISand A. R. BANDY,Geophys. Res. Lett. 7,681 (1980). 27. M. E. DAvis and T. A. LouzoN, Metallography 13,195 (1980). 28. J. P. FRANEY,Corros. Sci. 23, 1 (1983). 29. T. E. GRAEDEL,J. P. FRANEYand G. W. KAMMLOTr,Corros. Sci. 23, 1141 (1983). 30. W. H. STAHL(Ed.), Compilation of Odorand Taste Threshold Values Data, ASTM DS-48. American Society for Testing and Materials, Philadelphia (1973). 31. G. W. KAMMLOVr,Appl. Spectroscopy 35,324 (1981). 32. M. KENDALLand A. STUART,Advanced Theory o f Statistics, Vol. 1, 4th edn. Macmillan, New York (1977). 33. J. P. FRANEY,T. E. GRAEDELand G. W. KaMMLOrr, in Atmospheric Corrosion (ed. W. H. A1LOR), pp. 383-392. McGraw-Hill, New York (1982). 34. P. B. P. PHIPPSand D. W. RlCE, in Corrosion Chemistry (eds. G. R. BRUBAKERand P. B. P. PnIpps) ACS Symp. 89, p. 235-261. American Chemical Society, Washington, D.C. (1979). 35. D. W. RICE, R. J. CAPPELL,P. B. P. PHIPPSand P. PETERSON,in Atmospheric Corrosion (ed. W. H. AILOR), p. 651--666. McGraw-Hill, New York (1982). 36. C.J. WARDE,J. CORrSn and C. D. O'BRIAIN,J. electrochem. Soc. 122, 1421 (1975). 37. V. A. PH,LLIPS,J. appl. Phys. 33,712 (1962). 38. T. E. GRAEDEL,J. P. FRANEY, G. J. GUALTIERI, G. W. KAMMLO'Iq"and D. L. MALM, submitted for publication.