McIlroy, R. A,, “Absorption of Sulfur Dioxide by Molten Carbonates,” MS Thesis, The University of Akron, Ohio, December 1969. Oldenkamp, R. D., “Development of the Molten Carbonate Process For Control of Sulfur, Oxide Emissions,” presented at AIChE National Meeting, Cleveland, Ohio, May 5-7, 1969. Rohrmann, F. A., Ludwig, J . H., Chem. Eng. Progr., 61 (9), 59 (1965). Rosen, G., “Data and Calculations for Gasification of Spent Cooking Liquors from the Pu,l,p Industry, Part 1: Fundamental Thermodynamic Quantities, Trans. Roy. Insti. Technol., Stockholm, Sweden (159). 1960.
Rossini, F. D., Wagman, D. D., Evans, W. H., Levine, S., Jaffe I., Nat. Bur. Std. Circ. 500, p 496 (1952). Ruthven, D. M., Kenney, C. N., Chem. Eng. Sci., 22,1561 (1967). Squires, A. M., Chem. Eng., 74 (24) 260 (November 6, 1967). Stull, R. D., Ed., JANAF Thermochemical Tables, the Dow Chemical Co., Midland, Mich., August 1965; First Addendum August 1966; Second Addendum August 1967. Thoen, G. N., DeHass, G. C., Austin, R., TAPPI, 51, 6, 246 ( 1968). Received for review January 24, 1973. Accepted July 27, 1973
Adsorption of Hg(ll) by Hydrous Manganese Oxides R. Addis Lockwood and Kenneth Y. Chen’ Environmental Engineering Programs, University of Southern California, Los Angeles, Calif. 90007 The adsorption of Hg(I1) by hydrous manganese oxides suspended in solutions with wide ranges of p H , chloride concentration, and ionic strength was studied t o evaluate the possible importance of manganese oxides in mercury scavenging in natural waters. Manganese dioxide and possibly other oxides were precipitated with the addition of M n S 0 4 solution t o K M n 0 4 solution. Low crystallinity was verified by X-ray. Hg(I1) was adsorbed rapidly when added to aged suspensions of MnOz with low ionic strength. The equilibrium values in the p H range 6-8 fit a Freundlich isotherm:
0.6M NaCl repressed adsorption below p H 9, but not above p H 10. Apparently 0.6M NaC104 changed equilibriu m values only slightly, b u t adsorption rates were many orders of magnitude lower, particularly above p H 8. The uncharged Hg(OH)2 is the adsorbed species. HgC12 was also adsorbed, but not so strongly as H g ( 0 H ) z . It is concluded t h a t MnOz may be important as a mercury scavenger in fresh or brackish water, either in the natural environment or in the treatment of water and waste water.
The enrichment of manganese nodules with rare metals, including mercury, has been reported by various authors. Krauskopf (19561, after observing the results of laboratory adsorption studies and calculating solubility product limitations, stated t h a t a t least 13 of the metal concentrations in the sea are controlled by adsorption. Furthermore, of the various adsorbents, manganese dioxide appears t o be the most effective, with iron oxide t o a lesser degree. Though mercury was not included in his manganese dioxide experiments, the appearance of mercury in manganese nodules suggests t h a t manganese oxide may be effective for scavenging mercury (Goldberg, 1963). According to Krauskopf‘s (1956) estimate, between 0.01 and 0.07% of the amount of mercury entering qceans has remained in the water. The role of manganese in removing mercury from ocean water seems to be important (Harris, 1968). Manganese, though not as abundant as iron, is nevertheless nearly ubiquitous in its appearance in natural waters. Jenne
T o whom correspondence should be addressed. 1028
Environmental Science 8 Technology
(1968) suggested t h a t the activities of earth materials in adsorbing various metals might be localized on manganese and iron oxide sites. Early work in this laboratory indicated t h a t clay materials of nearly pure aluminum silicate are poor adsorbing agents for mercury, but material containing iron showed greater adsorptive activity. The adsorption of unhydrolyzed cations on manganese dioxide has been attributed t o ion exchange in the surface of the precipitate (Posselt et al., 1968). However, mercury is different from the divalent metals used in many adsorption studies because above p H 3 the uncharged metal oxide hydrate H g ( 0 H ) z is t h e dominant species. Also, mercury is easily reduced to the elemental state, which makes the problem more complicated if the redox potential is not maintained a t a high value to prevent the formation of the reduced forms. The main purpose of this work was to evaluate the role of hydrous manganese oxide in t h e scavenging and transport of mercury in water. The degree t o which these solids remove Hg(I1) from water solution was studied using pH, ionic strength, and sodium chloride concentration as variables. The 25°C isotherm function for the adsorption of aqueous mercury by MnO2 solids was also developed.
Experimental Equipment. Temperature was maintained a t 25 f 0.5”C in a constant temperature and humidity environmental chamber. The Perkin Elmer Model 124 doublebeam spectrophotometer was used for manganese analysis. The Unicam Model SP90 atomic absorption spectrophotometer was used in the flameless mode with a 15-cm path quartz absorption cell for the mercury analysis. pH was determined with a Corning Model 110 p H meter using a Corning glass electrode and an Orion double-junction standard calomel electrode. Suspensions were kept agitated with a n Eberbach shaker operating a t about 120 strokes per minute. Reagents. All reagents were analytical reagent grade or primary standard grade except for sodium perchlorate, which was Fischer Scientific purified grade. The analytical reagent red mercuric oxide used as the mercury standard was dried a t 110°C and stored over calcium chloride. The hydroxylamine hydrochloride contained mercury. Solutions were sparged overnight with nitrogen t o remove the mercury before use in the analytical procedures. All other reagents were checked for mercury and were free of interfering quantities. Sodium hydroxide was rendered carbonate-free by centrifuging the saturated solution.
Preparations. A mercury solution 5 x lO-3M (0.1%) was prepared from mercuric oxide dissolved in perchloric acid to give a final acid concentration of 0.1M. This solution was diluted lo4 times to 0.1 ppm for use in preparing the standard analytical working curve. T o allow adsorption experiments with low ionic strength, a second stock mercury solution was prepared with the minimum of perchloric acid required to keep t h e p H below 3. Solutions stronger than 40 ppm, the solubility of HgO (Dyrssen and Tyrrel 1961), start to precipitate HgO in t h e range of p H between 3 and 4 as hydrolysis of the Hg2f ion takes place. This stock was diluted to 20 p p m and neutralized with sodium hydroxide for use in mercury additions to manganese dioxide suspensions. Preparation of manganese dioxide suspension was similar to t h a t of Morgan a n d S t u m m (1964) except t h a t the suspension was prepared a t low p H to minimize the ionic strength of the solution. T h e m a n ganese oxide solids were produced by precipitation in situ a t the approximate concentration to be used in the mercury adsorption phase of the experiments t o avoid possible side effects from separations and resuspension. T h e concentrations were too low in most cases to separate enough solids for a B E T surface area determination. However, X-ray examination of several precipitates was made to determine the identity of the crystalline structure; the results using a Norelco X-ray machine with a copper target a t 30 kV and a goniometer scanning a t l"/min showed t h a t the precipitates were amorphous. Since characterization of the manganese oxide solids was difficult or impossible by other means, the precipitation technique was carefully standardized as the best means to characterize the solids and provide reproducible results. Procedure. Neutral mercury solution was always added to preformed manganese dioxide precipitates, after which the course of the interactions a t 25°C between the solute Hg(I1) and the MnOz solids was followed by filtering out the solids and analyzing the filtrates for mercury. T h e precipitates of MnOz were prepared in pretared 250-ml glass-stoppered conical flasks, starting with weighed quantities of triple-distilled water calculated to make final suspension volumes of 200 ml in each flask. An aliquot of M n S 0 4 stock solution was pipeted into a flask and mixed with the distilled water; then a n equal aliquot of K M n 0 4 stock solution was pipeted while swirling the flask vigorously. Each newly precipitated suspension was placed on an Eberbach shaker operating a t 120 strokes per minute to age for 1 hr a t 25"C, then precomputed amounts of COz-free standard FJaOH solution were added as rapidly as possible to the flasks. NaCl or NaC104 salt was also added at this juncture in most of the tests where they were used. After all additions except for Hg(II), the suspensions were returned to the shaker Ct 25°C in an environmental chamber to complete the procedure for MnOz solids preparation by aging for 12 to 16 hr. The aging procedure allowed the solids to equilibrate with the solution and stabilize the p H before the addition of the Hg(I1) solution. The p H drift during the aging made the measurement of p H in the preparation stages of little value, in addition to which, such a procedure would expose the solutions to contamination from air and possibly from the p H electrodes as well. The approximate relationship of the p H of the suspension to the quantity of standard caustic used was determined beforehand, using a beaker titration of similar MnO2 precipitates to obtain the titration curve for each MnOz concentration used. The p H was measured on separate samples of 20-25 ml after the mercury addition. During the exploratory phase of the work, t h e addition of the neutral Hg(OH)* solutions had no measurable effect on the pH-even when 99% of the mercury was adsorbed from 10 - 6 M Hg(1I) solution.
After the precipitates were aged on t h e shaker, the flasks were opened on a preset schedule; mercury solution (1-5 ml) was pipeted into a flask being shaken by hand; then the flask was returned to the shaker. Samples were withdrawn by syringe a t prescheduled times and pressure filtered through new, dry 25-mm Sartorius cellulose nitrate membranes with 50-nm pore size. Filtration pressure was produced by hand with 25-ml plastic syringes. T h e filtrates were caught in tared test tubes or flasks, weighed, --and analyzed for mercury using a flameless atomic absorption method (Chlorine Institute, 1970). Samples of unfiltered solution were also analyzed. The difference in mercury concentration between the filtered and unfiltered samples was taken to be the amount of mercury adsorbed by the solids. In the final series using sodium chloride as a desorption agent, the test proceeded in the same pattern with the exception that sodium chloride was added after the addition of the mercury. Subsequent to the addition of the mercury, the flasks were shaken for 1 hr before a sample was taken for filtration and mercury analysis; after the sample was withdrawn the preweighed amount of sodium chloride was added in the remaining suspension containing mercury and manganese dioxide. T h e p H adjustment by addition of NaOH increased sodium ion and ionic strength in the high p H range. To test sodium ion or ionic strength as possible factors in the millimolar range, a n experiment with NaC104 additions over the low range of concentrations was perbrmed. No sodium hydroxide was added so that the results with N a f ion concentrations in the millimolar range could be compared with results a t the same p H with near zero N a + concentration.
Results Adsorption Rates. In all cases, an attempt was made to measure the reaction rates and equilibrium values of adsorption. In many r u n s , ' t h e rate was too fast to measure even by the most rapid solids separation methods available. T h e values reported in the range of 1-2 min after the mercury addition are difficult to interpret. First of all, there is a n uncertainty in fixing the time of contact, which stems from the difficulty in determining a starting time in the 20- to 30-sec period for draining the mercury solution from the pipet; also. the sampling time during the 30 sec or more required for sampling and filtering 5-20 ml of the suspension could not be precisely fixed. In many cases it will be seen that the values recorded for the shortest times are, in fact, very close to the final equilibrium values. This is exemplified by experiments with low mercury-to-manganese ratios. When the mercury-manganese ratio was about 10-3.5, changes in adsorption could be only marginally detected after the first set of samples were taken. Results a t somewhat higher ratios (10-2.2) with no added salts, as shown a t the top of Figure 1, are typical of the results observed when the ratios were
10- 3.5. Figure 2 shows the results of pursuing the mercurymanganese ratio to the range l O - l . l , still at low ionic strength. The results clearly show lower adsorption rates a t the higher mercury-manganese ratio. The two experiments with fl lower than 7 show rates that are many times faster in attaining the apparent equilibrium values than those with higher p H . Kote that the values for q , the mole ratio of Hg(I1) to MnOz, are five to 10 times higher in Figure 2 than in Figure 1. The coverage of manganese oxide by mercury apparently has a strong effect on the a d sorption rate; the poor fit of the data to the Langmuir equilibrium isotherm may also be related to the coverage factor. Volume 7 , Number 11, November 1973
1029
pH Effect. Figure 3 is a plot of equilibrium values of q / ( ' (averages of results with 24 hr or more contact time)
from a variety of experiments over the range from p H 1.5- 10.8. It is necessarily somewhat rough, because there is a large range of mercury-manganese ratios included. However, the general flat shape between p H 4 and p H 10 is evident. with decreases in adsorption a t both ends of the p H scale. The value q / C was plotted here because it is
Figure 1 . Time function of salt-induced changes in Hg(ll) adsorption with pH 6 6 and 10.0
Comparison of 0 6 M NaCl and NaC104 with I < 1 0 - 3 M in 1 16 X 10 4M M n 0 2 suspensions containing - 0 8 X 1 0 - 6 M t o t a l H g ( l l )
I
0
eo
40
so
20
240
200
280
TIME (HRS)
Figure 2. Time plot of H g ( l l ) adsorption d t h high H g ( l l ) coverage over p H range 4-9.6 1.16 X 1 0 . - 5 M M n 0 2 suspensions with I tionof 1.00 x 1 0 - 6 M
