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Oxidation and Coagulation of Wastewater Effluent Utilizing Ferrate (VI) ion
407
OXIDATION AND COAGULATION OF WASTEWATER EFFLUENT UTILIZING FERRATE (VI) ION T. D. WAITE and K. A. GRAY Department of Civil Engineering University of Miami, Coral Gables, Florida 331 24
ABSTRACT For the past several years our laboratory has been developing the technology for thc usc of fcrratc (VI), [FeO:-], in water and wastewater treatment schemes. TWOproperties of fcrratc (VI), \vliich arc displayed by its oxidation and coagulation reactions, indicate that ferrate (VI)could be an cffective multi-purpose treatment chemical. Studies evaluating the reaction of fcrratc (VI) with various organic compounds have been reviewed. In addition t o being a selective oxidant of organic compounds, ferrate (VI) has been demonstrated to reduce THM potential and oxidize a model USEPA priority pollutant. Data has been collected under a number of different scenarios to illustrate the effectiveness of ferrate (VI) as a coagulant of turbid water systems, including secondary effluent. Current research is aimed at defining solid phase formation with ferrate ion decay in order to elucidate the mechanism of coagulation in colloidal systems. Comparative studies have been conducted which evaluate the efficiency and effectiveness of ferrous, ferric, and ferrate salts in destabilization of a model colloid.
1. INTRODUCTION
For the past several years our laboratory has been developing the technology for the use of ferrate (VI), [FeO:-], in water and wastewater treatment schemes. Ferrate (VI) has many appealing characteristics including its strong oxidizing potential (Ea = 2.3 V) and concomitant formation of reduced iron species which will act as coagulants of suspended materials. In addition, the solid phase formed, when ferrate (VI) is reduced, is a good scavenger of metals and anions such as phosphate. The above properties indicate that ferrate (VI) could be an effective multi-purpose treatment chemical for water and wastewater. This paper will discuss two properties of ferrate (VI), i.e., oxidation and coagulation reactions, as these properties are the most important for water and wastewater treatment. The disinfection capacity of ferrate (VI) has been described in detail in earlier publications. 2. FERRATE CHEMISTRY
Iron in its familiar forms exists in the +2 and +3 oxidation states; However, in a
strong oxidizing environment it is possible t o obtain higher oxidation states of iron. Compounds o f iron (IV), (V) and (VI) have been isolated as the metal salts of ferric acid, however, ir is the hexalent form of iron t!iat is o f interest for water treatment. This seemingly exotic form of iron has been of interest t o analytical chemists since 1841 when Fremy [ 11 first synthesized potassium ferrate. By 1925 a wide variety of metallic iron (VI) salts had been synthesized. It was not until 1948, however, that procedures were developed whereby a stable, crystalline solid of high purity could be synthesized, and analyzed for its iron (VI) content [ 2 , 3, 4, 51. As a result of work by Sclireyer, physical chemists and kineticists have been able t o establish the structure for iron (VI) ferrate and find evidence t o support its existence [6, 71. Although ferrate chemistry is in a state of infancy, several U.S. patents are currently held that relate to the use of ferrate in aqueous solutions. Three of these patents include: removal of color from industrial electrolytic baths [8]. use in making catalysts for the Fischer-Tropsch process [9, l o ] , and purification of hemicellulose [ 111. Ferrate (VI) ion has the molecular formula FeOi- and is a powerful oxidizing agent through the entire pH range. Wood [7] has reported the redox potential of ferrate to vary from -2.2 V to --0.7 V in acid and hase, respectively. The standard electromotive force for the half reaction is: F e 3 ' + 4 H z O + F e O ~ - + 8 H ' + 3e-.E0=-2.2+0.03V
(1)
Latimer [l?] gives a calculated potential estimate for the reaction Fe(OH)3
+ 5 0 H - - + FeO:-+
4 H z 0 -k 3 e - o f E, =-0.77 L0.03 V
Nearly 30 metallic salts of ferric acid (FeOz-) have beeii prepared, but only a few of these compounds are found t o yield a highly pure and staole product. As a matter of interest, ferrate compounds containing Ag, Al, Zn,Cr, Cu, Co, Pb, Mn. Ni. Hg, or T1 have been synthesized by double decomposition of B a F e 0 4 , and the correspondin6 metal nitrate in aqueous solution [ 131. For example: BaFe04
+ 2A1(N03)
+
Ba(N03)z
+ Alz FeO4
(3)
It is difficult to isolate most of these compounds from solution as they are subject t o decomposition at 3OoC, and react rapidly with COz while being dried in air. Of more practical interest are the metal ferrates which form either stable solutions o r stable crystalling solids. These compounds vary widely in their aqueous solubilities. Lithium, sodium, calcium, and magnesium ferrate are reported to be extremely soluble and can be synthesized by double decomposition with alkali metal perchlorate (MCIO,) and potassium ferrate [14]. Products vary in purity from 15% to 69%. Another procedure developed by Schreyer, et al. [4]. employs wet chemical oxidation of Fe (111) by hypochlorite, follwed by chenlical precipitation of F e O i - with KOH, forming K z FeO, . Recrystallization results in a high purity crystalline solid. This method was utilized in generating the potassium ferrate for experiments in this study. Aqueous solutions of ferrate ion have a characteristic violet color much like that of permanganate. Spectroscopic analysis of visible spectra of aqueous ferrate solutions show
409 one maximum peak at 505 nm and two minima at 390 nm and 670 nm. The molar extinction coefficient as determined by Wood [7] is 1070 *30 in lo-.' M NaOH. Potassium ferrate decomposes in aqueous solution generating hydroxide ion and molecular oxygen. The overall decomposition of ferrate (VI) ion in aqueous medium is described by Equation (4): 2Fe0,2-+ 3 H 2 0 + 2FeO(OH)
+ 3/2 O2 + 4 0 H -
(4)
The decomposition rate is strongly dependent on pH, initial ferrate concentration, temperature, and to some extent on the surface character of the hydrous iron oxide formed upon decomposition. Ferrate is most stable in strong base with two regions of maximum stability, one at pH 10-1 1 and the other in solutions greater than 3 M in base [7], although this is highly dependent on the initial Fe(II1) [IS]. Studies on the stability of ferrate in aqueous solution have shown that dilute solutions of ferrate are more stable than concentrated solutions [16]. Wagner, et al. [17], found 1.9 . M ferrate solutions to be only 37.4% decomposed after three hours and 50 minutes at 25OC. Ferrate decomposition rate has also been found to decline markedly in the presence of phosphate, and at low temperatures [7, 16, 171.
-
20
LO [minute]
30 Time
50
60
Fig. 1. Effect of various salts on ferrate (VI) stability (from Ref. 2).
Figure 1 illustrates the effects of several salts, and ferric hydroxide on ferrate (VI) stability. Other solute domains probably exist in the presence of SO:-, F - and dissolved or colloidal organic matter, and these groups can form stable complexes with Fe(II1) which may also alter the decomposition rate of K2 F e 0 4 . Ferrate reacts rapidly with most inorganic reducing agents under both acid and basic conditions. Reactions involving inorganic ammonia have been studied in detail, and oxida-
410
tion o f ammonia appears t o have a n optimum conversion in the pH range 9.5 t o 11.2, although losses t o the gaseous phase might be suspected at the higher pH. Strong [ 181 reported the degree o f conversion of ammonia t o increase as the molar ratio of ferrate to ammonia became greater, and as temperature was increased. Murmann [ 141 reports a sec-' at pseudo first order rate constant for ferrate oxidation of NH3 t o be 7 .0 . pH 10.6 and 2.5 . lo-* sec-' at pH 9.0. 3. OXIDATION OF ORGANICS
Iron (11) and iron (HI)enter into a wide variety of reactions with organic compounds which can include complexation, chelation, precipitation and oxidation-reduction. Although some work has been completed which evaluates these reactions, and provides some insight into mechanisms, the extent of organo-iron (VI) interactions is still largely unknown. The degree of oxidation of amino acids by FeO: - varies with initial ferrate concentration [ 141. Cystine and glycine react completely with excess ferrate forming C 0 2 and N 2 . When the amino acid is in excess, a variety of oxidation products are generated. Most sugars, and glycol are slowly oxidized t o organic acids. Certain organo-ferrate (VI) reactions have been studied, and in one study oxidation of organics by ferrate (VI) was evaluated by monitoring the disappearance of substrate using gas chromatohraphy [19]. Tests were conducted over a pH range of 2 t o 10.5 at 2OoC. A wide variety of substrate t o ferrate molar ratios were examined, utilizing the following substrates: benzene, allybenzene, chlorobenzene and 1-hexene-4-01. The data were evaluated in terms o f pH dependency, effect of substrate-ferrate molar ratio, and synergistic effects in two substrate systems. Ferrate was found t o significantly reduce the concentrations of allybenzene and chlorobenzene, while benzene and 1-1iexene-4-ol were converted by about 50% t o products. The range of maxima for per cent oxidation of substrates for reactions occuring below pH 8 are shown in Table 1. The oxidations are deTab. 1, Oxiddtion of Organic Substrates by Ferrate (VI) at pH values < 8 (from Ref. 19) Coinpound
% oxidation
Benzene Chlorobenzene Allylbenzene 1-hexene-4-oi
18-47 23-76 85-100 32-55
pendent on S : Fe(V1) molar ratios, where an excess of ferrate is shown to be most effective in reducing substrate concentrations. Molar ratios of s:Fe(V1) greater than 1 : 3 did not significantly enhance conversions. This points to the formation of products such as organic acids, rather than complete oxidation t o C 0 2 . It is possible that more complete oxidation of substrates would be obtained with multi-stage additions of ferrate. Studies evaluating the reaction of ferrate (VI) with phenol have also been undertaken. Variable ratios of ferrate (VI) t o phenol were investigated at variable pH; then, per cent
41 1
103
Secondary effluent FeOZ-added in phosphate buffer
*--TOC -609
90
(initial 12 ng/L) [initial 12 8 mg/L)
80 70 0
c
60
al
a
$ 50 40
30
20
Ot
I
L -t
8
I
:2 16 20 24 FeO?dose [mg./L a s Fe:
-
Fig. 2. Oxidation of TOC and BOD in sccondary cfflueiit wit11 ferrate (VI).
removal o f phenol, and COD were determined. The following conclusions were made from tlie study: 1) There is a general increase in ferrate reactivity with increasing pH. 9) Initial ferrate (VI) concentration is an important factor at high ferrate t o phenol ratios. 3) The relative ratio of phenol to phenolate species (C, H, OH/C, H5 0-) regulates reactivity as the ionized species reacts more readily with ferrate (VI). 4) Efficient phenol oxidation occurs when tlie ferrate (VI) to phenol molar ratio is > 10.
412
1
-Unfiltered (TOC = 17 porn) &-Filtered ( T O C = l L ppm) Chiorme 30 ppm PH
'
8.5
Contact Time 4Hrs 25
/
6 20-
/*
c
0
3
D Q)
CK
< O
1510-
/r
//
Fig. 3. Oxidation of THM precursors with ferrate (VI).
Ferrate (VI) in concentrations of less than 10 mg/l as Fe is also able to oxidize biodegradable organics (BOD) in domestic secondary effluent. Figure 2 shows data from an experiment where ferrate (VI) was added at different concentrations t o a secondary effluent with a total carbon content of 13 mg/l and a BOD o f 12.8 mg/l. It can be seen that all of tlie biodegradable carbon was oxidized by ferrate (VI), and approximately 35% of the TOC was removed. It should be noted here that no filtration o f the effluent after ferrate (VI) addition was attempted; therefore TOC and BOD removals are due t o oxidation only. It is anticipated that even greater removals could be achieved if the coagulation capacity of ferrate (VI) were taken into account. Preliminary studies have also been undertaken t o evaluate the ability of ferrate (VI) t o oxidize organic precursors o f trihaloniethane (THM). Water samples were collected from the Fox River wliich is located in Northern Illinois (U.S.A.). The water samples which were analyzed averaged 48 turbidity units (TU), pH = 8.5, and TOC = 17 nig/l. The effect of ferrate (VI) o n trihalonietlime potential was examined at several Fe0:- doses. Ferrate was applied 30 nlinutes prior t o chlorination. In all tests a dose of 30 ppm chlorine was used which represented the approximate demand of the raw water. Trihaloniethane concentrations were measured four hours following chlorination. Figure 3 shows a summary of the experimental data. Ferrate (VI) was able to reduce THM potential up t o approximately 25% in this system. The data indicate that no optimum ferrate (VI) did exist, so it is somewhat difficult to interpret the results. It should also be pointed out that the reduction in formation of THM shown here is due entirely to oxidation of precursors. There was no filtration of the samples, thus, no removal due
413
to coagulation was measured, In addition, the amount o f chlorine added was held constant at 30 ppm even though ferrate (VI) was added. Because the ferrate (VI) would normally perform most o f the disinfection, very little chlorine would have t o be added after ferrate (VI) treatment. This would further reduce THM formation in treated secondary effluents. The above data indicate that ferrate (VI) is a selective oxidant of organic compounds, and may have use as an oxidant of toxic organics in waste streams. One recent study has evaluated ferrate (VI) oxidation of one of the 129 priority pollutants listed by the USEPA [20]. Naphthalene was selected in this study as a model compound, and was reacted with ferrate (Vl) at different molar ratios. Table 2 below shows a summary of the data, and it can be seen that ferrate (VI) was an efficient oxidizing agent for the model priority pollutant. Tab. 2. Oxidation of Naphthalene by Ferrate (VI) (from Ref. 20) Perccntage Removal Molar Ratio Perratc/Naphthalenc
N a p h t h a l e n e (nig/L) 100 320 1000
10 20 30 40 60
22 43.5 63.2 15.5
82.2
40 66 16 83.5 95
46 90 100 100 100
4. COAGULATION
The removal of turbidity is also a major objective in water and wastewater treatment. In natural waters turbidity is largely the result of discrete, negatively charged particles and macromolecules which are stabilized by charge repulsion. Coagulation is the process where the surface chemistry of colloids is modified to pernlit aggregation and subsequent removal by gravity settling. The mechanism of coagulation can involve a number of reactions, but in general, two distinct phenomena probably occur: (a) the potential energy of repulsion is reduced; and (b) particles become enmeshed in a precipitate as it is formed [21]. A great deal of research has been conducted o n iron (11) and iron (111) coagulation which has promoted a widely accepted theory describing the probable n~echanismof colloid destabilization. It has been clearly established that salts of iron and alunlinium undergo hydrolysis in aqueous solution. In turn, the resultant aquometal complex will undergo polynierization by successive elimination of coordinated water molecules by Iiydroxide groups. The extent of polymerization is pH dependent. Under conditions uhich exceed the solubility limit of the metal hydroxide, these various polynuclear hydrolysis products may be considered soluble kinetic intermediates in the gradual precipitation of the metal hydroxide [??I. It has been shown that polynuclear hydroxo-metal complexes are readily adsorbed at the liquid-solid interface, and are more effective than non-hydro-
414
lyzed ions in destabilizing colloids. Although these intermediate metal species could be regarded as indifferent electrolytes producing coagulation by double layer compression, this is thought to be less significant than the reduction of zeta potential [21, 231. If similar reasoning is applied to the iron (VI) system, it is plausible that coagulation may involve a greater variety of intermediate hydrolytic species, and possibly, species of greater net, positive charge. It is known that the ferrate (VI) ion rapidly decomposes in acid solution, and its stability increases with increasing pH above 7 [16]. Although the kinetics of Fe5+ and Fe4' formation have never been specifically defined, differences in both the mechanism of precipitation of solid iron and the coagulation behavior between ferrate (VI), iron (IlI), and iron (11) systems would provide a basis for inferring the existence of such intermediate species. Preliminary studies have demonstrated that ferrate (VI) will effectively coagulate turbid water systems. Coagulation jar tests have been carried out on lake water systems to which bentonite clay was added to increase initial turbidity. The results of these tests are presented in Figure 4 which illustrates final turbidity values for a range of ferrate (VI) dose. These data demonstrate a general trend that turbidity removal increases with ferrate (VI) dose to an optimum value of 5 mg/l Fe0:- as Fe. Beyond this minimum, turbidity removal decreases with increasing ferrate dose. When ferrate was added in phosphate buffer the turbidity was reduced by 95%, while for the carbonate buffer and distilled water applications, the turbidity was reduced 79% and 84% respectively. It appears that the presence of phosphate has a positive effect on
0
-b
FeO:-odded
in carbonate bLffer
4
-.-
Fee:-added
:n distil,ed w a t e r
-
!,
1
Lake W a t e r Bentonite CLay 2Fe04 added in phosphate Duffer
06.1
In
2 05aJ
1
z. 0" 4 -\ L
0
\
ul
-n 0 3 - \ 0
L
~
T__
6
.r-
8
.
7
-
12 16 20 2L FeOZ- dose I r n g I L as F e l
Fig. 4. Coagulation bentonite augmented lake water by ferrate (VI).
r
41 5 ferrate (VI) ability to destabilize colloids. I n order to further investigate this phenomenon, coagulation tests were performed on the lake water and bentonite clay system using the optimum ferrate dose (5 mg/l Fe0;- as Fe) with varying amounts of orthophosphate (0-10 mg/l as P). In the presence of increasing amounts of orthophosphate, turbidity removal improved. The enhanced efficacy of ferrate (VI) in the presence of phosphate may be due to a combination of phenomena; e.g. ( 1 ) the stabilization of the ferrate (VI) ion in the presence of phosphate, due to its chelation of Fe3+,which otherwise accelerates the decomposition of Fe6+; and (2) the formation of a mixed hydroxo-phosphatometal precipitate, which enhances colloid enmeshment or coprecipitation. Comparative coagulation jar tests were carried out on the lake water and bentonite clay system using iron (11) and iron (111) salts. At low coagulant dose, i.e., up to 10 mg/1 as Fe, ferrate and ferrous iron remove turbidity more efficiently than ferric iron. At the optimum ferrate dose of 5 mg/l as Fe, greater turbidity removal is achieved with ferrate than with either ferrous or ferric iron. In order to accomplish the same degree of coagulation, a coagulant dosage greater than 15 mg/l as Fe must be used with both ferrous and ferric iron. Furthermore, at ferrate doses greater than 8 mg/l as Fe, the final turbidities of the coagulated system increase with increasing ferrate dose. This behavior contrasts the trend of increasing turbidity removal with increasing dose at low ferrate concentrations and at all tested doses of iron (11) and iron (111). The ferrate system exhibits either colloid restabilization or the formation of a stable hydrous iron oxide colloid with increasing ferrate concentration. Data collected from a small (10 l/h) bench-scale pilot plant also indicated that ferrate can be an effective chemical for suspended solids removal in tertiary treatment of secondary effluent [24]. The pilot facility included a flash nux reactor, where potassium ferrate, of 90% or greater purity, was added, a flocculation unit, and a sedimentation clarifier. The residence time in the flash mix unit was approximately two minutes, and in the three chambered flocculation unit, approximately 45 minutes. The total sedimentation time in the final clarifier wzs four hours. The plant was operated at steady state
--
t
Ferrate concentratlon
Orng/L
2Q/L
o-.- Cmg/L
6mg/~ 8 mg/L c.....1.0 mg/L c..-
'\.
v)
U
L C
-H!%lVG
Fig. 5. Coagulation of secondary effluent in a pilot treatment plant.
416 throughout testing. Figure 5 shows suspended solids removal through the bench-scale pilot plant at various ferrate doses. Despite the variability of secondary effluent quality, the data demonstrate that the pilot system operated optimally at a ferrate dose of 8 mg/l with approximately 86% suspended solids removal. Better than 80% solids removal was observed in all tests utilizing ferrate in concentrations greater than 6 iiigil. However, higher ferrate dosages seemed t o contribute to the turbidity of the effluent stream. This turbidity was not measurable in the solids deternunation, and was obviously due to the formation of colloidal iron from ferrate addition. The behavior of ferrate in water systems with a wide array of organic and inorganic constituents is extremely complex, as a combination of oxidation, coagulation, and precipitation reactions occur. In order t o focus singly on ferrate’s coagulative capabilities. current research is investigating the behavior of ferrate in a defined colloidal system, buffered at various pH and maintained at constant ionic strength. The objective of this new research is t o consider the operative behavior of ferrate (Vl) decay in coagulation in contrast to t h e behavior of iron (11) and iron (111) salts. Since, in practice, the solubility limit o f metal hydroxide is normally exceeded by the dosage of metal salts required t o destabilize colloids, careful consideration has been given to the kinetics of hydrous iron oxide generation for iron (VI), iron (HI), and iron (11) salts. Therefore, an approach has been developed to consider the decay of the ferrate ion, and the precipitation of the resultant metal hydroxide in comparison t o the iron precipitation reactions of iron (11) and iron (111) salts. F o r equivalent dosages of iron, the rate of turbidity generation is believed to reflect the duration of soluble iron species, intermediate t o the formation of iron hydroxide. The formation of hydrous iron oxide with time, as measured by light scattering, has been studied at a variety of iron salt doses and pH. The appropriate dose of iron was added t o one liter of biocarbonate buffer adjusted t o the specified pH and ionic strength. After disper;ing the dose throughout the volume, a small aliquot was transferred to a sealed vial where iron precipitation and sedimentation was followed under quiescent conditions. Data from experiments conducted for potassium ferrate. ferrous sulfate, and ferric nitrate at doses of 15 mg/l as Fe are presented in Figure 6. These results reflect the rate o f generation of an insoluble phase in aqueous solution, buffered at pH 7. If the turbidity profile of iron (111) is regarded as a standard illustration of the hydrolysis, olation, and precipitation of the ferric ion, then comparison between this and profiles similarly developed for iron (11) and iron (111) will demonstrate any differences in the reaction kinetics of iron hydrolysis. Although turbidity measured by light scattering represents the net result of solid iron formation, coagulation, and sedimentation, the formation of solid iron is considered to be the predominant mechanism in iron (11) and iron (Vl) systems within the first few minutes o f turbidity monitoring. Figure 6 shows that a ferric salt generated a maximum m u n t of turbidity within 30 seconds o f the salt addition. Flocs were visible immediately, and throughout the rest of the analysis, turbidity was seen to decrease, due to the flocculation and settling o f the hydrous iron oxide. The turbidity profile of the same dose of ferrous salt in the buffered aqueous system describes behavior different from that o f the iron (111) system. The Fe” must undergo oxidation to Fe3+ prior to hydrolysis, polymerization, and precipitation of an insoluble iron. Under atmospheric conditions at pH = 7 significant oxidation o f iron (11) takes approximately 20 minutes [25]. However, t h e turbidity profile in Figure 6 may reflect t h e competitive for-
41 7
15mg/L a s Fe pH 17
50 -
0-
x--
0-.-
Fe(NO3I3 Fe S O 4 . 7H20 K 2Fe O4
-z 3 5 x
c
+? 3
I-
i
30-
/
25. 20
-
0
i
/
5
10
15
rime [ m i n u t e ]
Fig. 6. Formation of insoluble phase by different iron salts.
mation of a ferrous carbonate complex due to the use of a sodium bicarbonate buffer. Visible flocs were not observed for the first 60 minutes of monitoring, although the color of the system deepended with time. The iron (11) system generated an iron colloid which was stable for approximately two hours, at which time fine flocs became visible and rapidly settled. The ferrate (VI) system demonstrated turbidity formation distinct from both the iron (11) and iron (111) systems. Decomposition of the ferrate ion occurred within the first 3.5 minutes, and fine flocs were observed after the first minute of monitoring. Initially, turbidity increases at a rate slightly greater than that of the ferrous system, but this rate slows after two minutes and an iron colloid is formed. This colloid was observed to be stable for approximately 90 minutes before significant settling occurred. Figure 7 reports data collected from colloid destabilization experiments. A colloidal silica suspension was developed, buffered at pH 7, and maintained at constant ionic
418
\
\
\
/
-s
100
- 90 a,
c
5 80
E"
f
\
\
70
1 ,
c
{
60
?
50
f
\
40 30 20 10
t
10
20
LO 30 Time [ m l n u t e l
50
60
Fig. 7. Destabilization of a silica colloid by different iron salts.
strength. Colloidal destabilization was determined for each iron salt at 15 mg/l as F e and measured as the amount of turbidity remaining in the systems at various times. After the addition of each salt, t o one liter of colloid suspension stabilized over 24 hours, the system was rapidly mixed for one minute, slowly flocculated for 30 minutes, and allowed t o settle for 30 minutes. Ferrate (VI) achieved the greatest amount of turbidity removal, 89% after 30 minutes of settling, while ferrous and ferric salts produced 76% and 58% removal, respectively. It is thought that ferrate accomplished the largest degree of colloid destabilization because o f the greater coagulating efficiency o f soluble species intermediate to its decay t o Fe3'. The 20% decrease in turbidity within one minute of ferrate addition may be due t o the surface activity of these species. After the initial flash mixing, turbidity increased t o a constant level throughout flocculation. This increase reflected the growth o f metal colloid flocs. The iron (11) system, which generated a great deal of turbidity in the ab-
419 sence of the silica colloid, showed a very small, initial decrease,in turbidity and a large increase in turbidity with flocculation. Throughout flocculation very large flocs were visible. The iron (111) system, which had demonstrated immediate hydrous iron oxide formation and gradual settling, behaved similarly under colloidal destabilizing conditions. There was no increase in turbidity with flocculation, and the least amount of silica turbidity was removed. In fact, only in the system destabilized with iron (111) was the white silica colloid still visible after 30 minutes of settling. In conclusion, both ferrate (VI) and iron (11) were better coagulants than iron (111) for this test system. This may be due to the duration of soluble species and the lesser rate of solid iron generation. Also, ferrate (VI) was observed to remove more turbidity from this model colloid system than iron (11). There may have been some diminishment in the coagulation ability of the ferrous salt due to the interference of the bicarbonate buffer. However, most of the iron (11) is believed to have been oxidized to iron (111) under the conditions of this experiment. It one considers the behavior of iron (11) and iron (VI) in buffered aqueous systems, it may be that the stability of the solid iron generated by iron (11) and iron (VI) is mediated by the variety of soluble iron species intermediate to the transition to ferric hydroxide. It may be those differences in the nature of the soluble species generated by ferrate and ferrous salts that also mediates coagulation efficiency. It is also interesting to note that even though a large amount of turbidity was generated by the iron (11) system it did not effectively coagulate the silica colloid. These observations support the theory that the most efficient mechanism of coagulation proceeds via adsorption of soluble aquometal species at the colloid surface.
REFERENCES 1 E. F. Freniy, Coiiipt. Rend. 12, 23, 1841. 2 J. M. Schreyer, L. T. Ockernian and G. W. Thompson, Anal. Chcm., 22, 1950, 691. 3 Inorganic Synthesis, 4, 1953, 164. 4 J. M. Schreyer, et. al., Anal. Cheni., 22, 1951, 1426. 5 J. M. Schreyer, et. al., Jour. Anier. Cheni. SOC., 73, 1951, 1379. 6 H. J. Hrostowski and A. B. Scott, Mourn. Chcm. Phys., 18, 1950, 105. 7 R. H. Wood, The Heat, Free Energy and Entrophy of the Ferrate (V1) Ion, Journ. Amcr. Chcm. SOC.,80, 1958, 2038. 8 J. M. Schreyer, USP, 2,536,703, 1951. 9 M. A. Mosesnian, USP 2,470,784, 1949. 10 M. A. Mosesnian, USP 2,455,696, 1948. 11 J. B. Harrison, USP 2,728,695, 1965. 12 W. M. Latinier, Oxidation Potentials, Prentice Hall, N.Y., 1952. 13 L. Lozana, Acido ferric0 e ferrati (VI), Gazz. Chim. Ital., 55, 1925, 468. 14 R. K. Murmann, The Preparation and Oxidation Propcrtics of Fcrrate (FeO;.), NTIS Publication PB-2 3 8-057, 1974. 15 R. G . Hairc, A Study of thc Deconiposition of Potassium Ferratc (Vl) in Aqueous Solution, Doctoral Abstracts, 1965. 16 J. M. Schreyer and L. T. Ockernim, Stability of Ferratc (Vl) Ion in Aqueous Solution, Anal. Cliem., 24, 1950, 1498. 17 W. F. Wagner, J . K. Gump and E. N. Hart, Factors Affecting the Stability of Aqueous Potassium Ferrate (V1) Solutions, Anal. Chcm., 24, 1952, 1397.
420 18 A. W. Strong, An Exploratory Work on the Oxidation of Ammonia by Potassium Ferrate (VI), NTIS Publication PB 231873,1973. 19 T. D. Waite and M. Gilbert, Oxidative Destruction of Phenol and Other Organic Residuals by Iron (VI) Ferrate, J. Water Poll. Contrl. Fed., 1978, p. 543. 20 S. Deluca, A. C. Chao and C. Smallwood, Removal of Selected Pollutants with Potassium Ferrate, Proc. of 13th Mid-Atlantic Conf. on Indust. Waste, ed. by C. P. Huang, Ann Arbor Sci. Pub., 1981. 21 C. O’Melia and W. Stumm, Aggregation of Silica Dispersions by Iron (111), J. Colloid and Interface Sci., 23, 1967, 437. 22 C. R. O’Melia, “Coagulation and Flocculation”, in Physicochemical Processes, W. J. Weber, Jr., New York: Wiley-Interscience, 1972. 23 G. R. Wiese and T. W. Healy, Adsorption of AI(II1) at the TiO, -H,O Interface, Jour. Colloid and Interface Sci., 51, 1974, 434. 24 T. D. Waite, Feasibility of Wastewater Treatment with Ferrate, ASCE J. of Environmental Engg. Div., 105, 1979, 1023. 25 J. T. O’Connor, Iron and Manganese, in Water Quality Treatment, AWWA, New York, McGrawHill, 1971.
OXIDATION AND COAGULATION OF WASTEWATER EFFLUENT UTILIZING FERRATE (VI) ION T. D. WAITE and K. A. GRAY Department of Civil Engineering University of Miami, Coral Gables, Florida 331 24
ABSTRACT For the past several years our laboratory has been developing the technology for thc usc of fcrratc (VI), [FeO:-], in water and wastewater treatment schemes. TWOproperties of fcrratc (VI), \vliich arc displayed by its oxidation and coagulation reactions, indicate that ferrate (VI)could be an cffective multi-purpose treatment chemical. Studies evaluating the reaction of fcrratc (VI) with various organic compounds have been reviewed. In addition t o being a selective oxidant of organic compounds, ferrate (VI) has been demonstrated to reduce THM potential and oxidize a model USEPA priority pollutant. Data has been collected under a number of different scenarios to illustrate the effectiveness of ferrate (VI) as a coagulant of turbid water systems, including secondary effluent. Current research is aimed at defining solid phase formation with ferrate ion decay in order to elucidate the mechanism of coagulation in colloidal systems. Comparative studies have been conducted which evaluate the efficiency and effectiveness of ferrous, ferric, and ferrate salts in destabilization of a model colloid.
1. INTRODUCTION
For the past several years our laboratory has been developing the technology for the use of ferrate (VI), [FeO:-], in water and wastewater treatment schemes. Ferrate (VI) has many appealing characteristics including its strong oxidizing potential (Ea = 2.3 V) and concomitant formation of reduced iron species which will act as coagulants of suspended materials. In addition, the solid phase formed, when ferrate (VI) is reduced, is a good scavenger of metals and anions such as phosphate. The above properties indicate that ferrate (VI) could be an effective multi-purpose treatment chemical for water and wastewater. This paper will discuss two properties of ferrate (VI), i.e., oxidation and coagulation reactions, as these properties are the most important for water and wastewater treatment. The disinfection capacity of ferrate (VI) has been described in detail in earlier publications. 2. FERRATE CHEMISTRY
Iron in its familiar forms exists in the +2 and +3 oxidation states; However, in a
strong oxidizing environment it is possible t o obtain higher oxidation states of iron. Compounds o f iron (IV), (V) and (VI) have been isolated as the metal salts of ferric acid, however, ir is the hexalent form of iron t!iat is o f interest for water treatment. This seemingly exotic form of iron has been of interest t o analytical chemists since 1841 when Fremy [ 11 first synthesized potassium ferrate. By 1925 a wide variety of metallic iron (VI) salts had been synthesized. It was not until 1948, however, that procedures were developed whereby a stable, crystalline solid of high purity could be synthesized, and analyzed for its iron (VI) content [ 2 , 3, 4, 51. As a result of work by Sclireyer, physical chemists and kineticists have been able t o establish the structure for iron (VI) ferrate and find evidence t o support its existence [6, 71. Although ferrate chemistry is in a state of infancy, several U.S. patents are currently held that relate to the use of ferrate in aqueous solutions. Three of these patents include: removal of color from industrial electrolytic baths [8]. use in making catalysts for the Fischer-Tropsch process [9, l o ] , and purification of hemicellulose [ 111. Ferrate (VI) ion has the molecular formula FeOi- and is a powerful oxidizing agent through the entire pH range. Wood [7] has reported the redox potential of ferrate to vary from -2.2 V to --0.7 V in acid and hase, respectively. The standard electromotive force for the half reaction is: F e 3 ' + 4 H z O + F e O ~ - + 8 H ' + 3e-.E0=-2.2+0.03V
(1)
Latimer [l?] gives a calculated potential estimate for the reaction Fe(OH)3
+ 5 0 H - - + FeO:-+
4 H z 0 -k 3 e - o f E, =-0.77 L0.03 V
Nearly 30 metallic salts of ferric acid (FeOz-) have beeii prepared, but only a few of these compounds are found t o yield a highly pure and staole product. As a matter of interest, ferrate compounds containing Ag, Al, Zn,Cr, Cu, Co, Pb, Mn. Ni. Hg, or T1 have been synthesized by double decomposition of B a F e 0 4 , and the correspondin6 metal nitrate in aqueous solution [ 131. For example: BaFe04
+ 2A1(N03)
+
Ba(N03)z
+ Alz FeO4
(3)
It is difficult to isolate most of these compounds from solution as they are subject t o decomposition at 3OoC, and react rapidly with COz while being dried in air. Of more practical interest are the metal ferrates which form either stable solutions o r stable crystalling solids. These compounds vary widely in their aqueous solubilities. Lithium, sodium, calcium, and magnesium ferrate are reported to be extremely soluble and can be synthesized by double decomposition with alkali metal perchlorate (MCIO,) and potassium ferrate [14]. Products vary in purity from 15% to 69%. Another procedure developed by Schreyer, et al. [4]. employs wet chemical oxidation of Fe (111) by hypochlorite, follwed by chenlical precipitation of F e O i - with KOH, forming K z FeO, . Recrystallization results in a high purity crystalline solid. This method was utilized in generating the potassium ferrate for experiments in this study. Aqueous solutions of ferrate ion have a characteristic violet color much like that of permanganate. Spectroscopic analysis of visible spectra of aqueous ferrate solutions show
409 one maximum peak at 505 nm and two minima at 390 nm and 670 nm. The molar extinction coefficient as determined by Wood [7] is 1070 *30 in lo-.' M NaOH. Potassium ferrate decomposes in aqueous solution generating hydroxide ion and molecular oxygen. The overall decomposition of ferrate (VI) ion in aqueous medium is described by Equation (4): 2Fe0,2-+ 3 H 2 0 + 2FeO(OH)
+ 3/2 O2 + 4 0 H -
(4)
The decomposition rate is strongly dependent on pH, initial ferrate concentration, temperature, and to some extent on the surface character of the hydrous iron oxide formed upon decomposition. Ferrate is most stable in strong base with two regions of maximum stability, one at pH 10-1 1 and the other in solutions greater than 3 M in base [7], although this is highly dependent on the initial Fe(II1) [IS]. Studies on the stability of ferrate in aqueous solution have shown that dilute solutions of ferrate are more stable than concentrated solutions [16]. Wagner, et al. [17], found 1.9 . M ferrate solutions to be only 37.4% decomposed after three hours and 50 minutes at 25OC. Ferrate decomposition rate has also been found to decline markedly in the presence of phosphate, and at low temperatures [7, 16, 171.
-
20
LO [minute]
30 Time
50
60
Fig. 1. Effect of various salts on ferrate (VI) stability (from Ref. 2).
Figure 1 illustrates the effects of several salts, and ferric hydroxide on ferrate (VI) stability. Other solute domains probably exist in the presence of SO:-, F - and dissolved or colloidal organic matter, and these groups can form stable complexes with Fe(II1) which may also alter the decomposition rate of K2 F e 0 4 . Ferrate reacts rapidly with most inorganic reducing agents under both acid and basic conditions. Reactions involving inorganic ammonia have been studied in detail, and oxida-
410
tion o f ammonia appears t o have a n optimum conversion in the pH range 9.5 t o 11.2, although losses t o the gaseous phase might be suspected at the higher pH. Strong [ 181 reported the degree o f conversion of ammonia t o increase as the molar ratio of ferrate to ammonia became greater, and as temperature was increased. Murmann [ 141 reports a sec-' at pseudo first order rate constant for ferrate oxidation of NH3 t o be 7 .0 . pH 10.6 and 2.5 . lo-* sec-' at pH 9.0. 3. OXIDATION OF ORGANICS
Iron (11) and iron (HI)enter into a wide variety of reactions with organic compounds which can include complexation, chelation, precipitation and oxidation-reduction. Although some work has been completed which evaluates these reactions, and provides some insight into mechanisms, the extent of organo-iron (VI) interactions is still largely unknown. The degree of oxidation of amino acids by FeO: - varies with initial ferrate concentration [ 141. Cystine and glycine react completely with excess ferrate forming C 0 2 and N 2 . When the amino acid is in excess, a variety of oxidation products are generated. Most sugars, and glycol are slowly oxidized t o organic acids. Certain organo-ferrate (VI) reactions have been studied, and in one study oxidation of organics by ferrate (VI) was evaluated by monitoring the disappearance of substrate using gas chromatohraphy [19]. Tests were conducted over a pH range of 2 t o 10.5 at 2OoC. A wide variety of substrate t o ferrate molar ratios were examined, utilizing the following substrates: benzene, allybenzene, chlorobenzene and 1-hexene-4-01. The data were evaluated in terms o f pH dependency, effect of substrate-ferrate molar ratio, and synergistic effects in two substrate systems. Ferrate was found t o significantly reduce the concentrations of allybenzene and chlorobenzene, while benzene and 1-1iexene-4-ol were converted by about 50% t o products. The range of maxima for per cent oxidation of substrates for reactions occuring below pH 8 are shown in Table 1. The oxidations are deTab. 1, Oxiddtion of Organic Substrates by Ferrate (VI) at pH values < 8 (from Ref. 19) Coinpound
% oxidation
Benzene Chlorobenzene Allylbenzene 1-hexene-4-oi
18-47 23-76 85-100 32-55
pendent on S : Fe(V1) molar ratios, where an excess of ferrate is shown to be most effective in reducing substrate concentrations. Molar ratios of s:Fe(V1) greater than 1 : 3 did not significantly enhance conversions. This points to the formation of products such as organic acids, rather than complete oxidation t o C 0 2 . It is possible that more complete oxidation of substrates would be obtained with multi-stage additions of ferrate. Studies evaluating the reaction of ferrate (VI) with phenol have also been undertaken. Variable ratios of ferrate (VI) t o phenol were investigated at variable pH; then, per cent
41 1
103
Secondary effluent FeOZ-added in phosphate buffer
*--TOC -609
90
(initial 12 ng/L) [initial 12 8 mg/L)
80 70 0
c
60
al
a
$ 50 40
30
20
Ot
I
L -t
8
I
:2 16 20 24 FeO?dose [mg./L a s Fe:
-
Fig. 2. Oxidation of TOC and BOD in sccondary cfflueiit wit11 ferrate (VI).
removal o f phenol, and COD were determined. The following conclusions were made from tlie study: 1) There is a general increase in ferrate reactivity with increasing pH. 9) Initial ferrate (VI) concentration is an important factor at high ferrate t o phenol ratios. 3) The relative ratio of phenol to phenolate species (C, H, OH/C, H5 0-) regulates reactivity as the ionized species reacts more readily with ferrate (VI). 4) Efficient phenol oxidation occurs when tlie ferrate (VI) to phenol molar ratio is > 10.
412
1
-Unfiltered (TOC = 17 porn) &-Filtered ( T O C = l L ppm) Chiorme 30 ppm PH
'
8.5
Contact Time 4Hrs 25
/
6 20-
/*
c
0
3
D Q)
CK
< O
1510-
/r
//
Fig. 3. Oxidation of THM precursors with ferrate (VI).
Ferrate (VI) in concentrations of less than 10 mg/l as Fe is also able to oxidize biodegradable organics (BOD) in domestic secondary effluent. Figure 2 shows data from an experiment where ferrate (VI) was added at different concentrations t o a secondary effluent with a total carbon content of 13 mg/l and a BOD o f 12.8 mg/l. It can be seen that all of tlie biodegradable carbon was oxidized by ferrate (VI), and approximately 35% of the TOC was removed. It should be noted here that no filtration o f the effluent after ferrate (VI) addition was attempted; therefore TOC and BOD removals are due t o oxidation only. It is anticipated that even greater removals could be achieved if the coagulation capacity of ferrate (VI) were taken into account. Preliminary studies have also been undertaken t o evaluate the ability of ferrate (VI) t o oxidize organic precursors o f trihaloniethane (THM). Water samples were collected from the Fox River wliich is located in Northern Illinois (U.S.A.). The water samples which were analyzed averaged 48 turbidity units (TU), pH = 8.5, and TOC = 17 nig/l. The effect of ferrate (VI) o n trihalonietlime potential was examined at several Fe0:- doses. Ferrate was applied 30 nlinutes prior t o chlorination. In all tests a dose of 30 ppm chlorine was used which represented the approximate demand of the raw water. Trihaloniethane concentrations were measured four hours following chlorination. Figure 3 shows a summary of the experimental data. Ferrate (VI) was able to reduce THM potential up t o approximately 25% in this system. The data indicate that no optimum ferrate (VI) did exist, so it is somewhat difficult to interpret the results. It should also be pointed out that the reduction in formation of THM shown here is due entirely to oxidation of precursors. There was no filtration of the samples, thus, no removal due
413
to coagulation was measured, In addition, the amount o f chlorine added was held constant at 30 ppm even though ferrate (VI) was added. Because the ferrate (VI) would normally perform most o f the disinfection, very little chlorine would have t o be added after ferrate (VI) treatment. This would further reduce THM formation in treated secondary effluents. The above data indicate that ferrate (VI) is a selective oxidant of organic compounds, and may have use as an oxidant of toxic organics in waste streams. One recent study has evaluated ferrate (VI) oxidation of one of the 129 priority pollutants listed by the USEPA [20]. Naphthalene was selected in this study as a model compound, and was reacted with ferrate (Vl) at different molar ratios. Table 2 below shows a summary of the data, and it can be seen that ferrate (VI) was an efficient oxidizing agent for the model priority pollutant. Tab. 2. Oxidation of Naphthalene by Ferrate (VI) (from Ref. 20) Perccntage Removal Molar Ratio Perratc/Naphthalenc
N a p h t h a l e n e (nig/L) 100 320 1000
10 20 30 40 60
22 43.5 63.2 15.5
82.2
40 66 16 83.5 95
46 90 100 100 100
4. COAGULATION
The removal of turbidity is also a major objective in water and wastewater treatment. In natural waters turbidity is largely the result of discrete, negatively charged particles and macromolecules which are stabilized by charge repulsion. Coagulation is the process where the surface chemistry of colloids is modified to pernlit aggregation and subsequent removal by gravity settling. The mechanism of coagulation can involve a number of reactions, but in general, two distinct phenomena probably occur: (a) the potential energy of repulsion is reduced; and (b) particles become enmeshed in a precipitate as it is formed [21]. A great deal of research has been conducted o n iron (11) and iron (111) coagulation which has promoted a widely accepted theory describing the probable n~echanismof colloid destabilization. It has been clearly established that salts of iron and alunlinium undergo hydrolysis in aqueous solution. In turn, the resultant aquometal complex will undergo polynierization by successive elimination of coordinated water molecules by Iiydroxide groups. The extent of polymerization is pH dependent. Under conditions uhich exceed the solubility limit of the metal hydroxide, these various polynuclear hydrolysis products may be considered soluble kinetic intermediates in the gradual precipitation of the metal hydroxide [??I. It has been shown that polynuclear hydroxo-metal complexes are readily adsorbed at the liquid-solid interface, and are more effective than non-hydro-
414
lyzed ions in destabilizing colloids. Although these intermediate metal species could be regarded as indifferent electrolytes producing coagulation by double layer compression, this is thought to be less significant than the reduction of zeta potential [21, 231. If similar reasoning is applied to the iron (VI) system, it is plausible that coagulation may involve a greater variety of intermediate hydrolytic species, and possibly, species of greater net, positive charge. It is known that the ferrate (VI) ion rapidly decomposes in acid solution, and its stability increases with increasing pH above 7 [16]. Although the kinetics of Fe5+ and Fe4' formation have never been specifically defined, differences in both the mechanism of precipitation of solid iron and the coagulation behavior between ferrate (VI), iron (IlI), and iron (11) systems would provide a basis for inferring the existence of such intermediate species. Preliminary studies have demonstrated that ferrate (VI) will effectively coagulate turbid water systems. Coagulation jar tests have been carried out on lake water systems to which bentonite clay was added to increase initial turbidity. The results of these tests are presented in Figure 4 which illustrates final turbidity values for a range of ferrate (VI) dose. These data demonstrate a general trend that turbidity removal increases with ferrate (VI) dose to an optimum value of 5 mg/l Fe0:- as Fe. Beyond this minimum, turbidity removal decreases with increasing ferrate dose. When ferrate was added in phosphate buffer the turbidity was reduced by 95%, while for the carbonate buffer and distilled water applications, the turbidity was reduced 79% and 84% respectively. It appears that the presence of phosphate has a positive effect on
0
-b
FeO:-odded
in carbonate bLffer
4
-.-
Fee:-added
:n distil,ed w a t e r
-
!,
1
Lake W a t e r Bentonite CLay 2Fe04 added in phosphate Duffer
06.1
In
2 05aJ
1
z. 0" 4 -\ L
0
\
ul
-n 0 3 - \ 0
L
~
T__
6
.r-
8
.
7
-
12 16 20 2L FeOZ- dose I r n g I L as F e l
Fig. 4. Coagulation bentonite augmented lake water by ferrate (VI).
r
41 5 ferrate (VI) ability to destabilize colloids. I n order to further investigate this phenomenon, coagulation tests were performed on the lake water and bentonite clay system using the optimum ferrate dose (5 mg/l Fe0;- as Fe) with varying amounts of orthophosphate (0-10 mg/l as P). In the presence of increasing amounts of orthophosphate, turbidity removal improved. The enhanced efficacy of ferrate (VI) in the presence of phosphate may be due to a combination of phenomena; e.g. ( 1 ) the stabilization of the ferrate (VI) ion in the presence of phosphate, due to its chelation of Fe3+,which otherwise accelerates the decomposition of Fe6+; and (2) the formation of a mixed hydroxo-phosphatometal precipitate, which enhances colloid enmeshment or coprecipitation. Comparative coagulation jar tests were carried out on the lake water and bentonite clay system using iron (11) and iron (111) salts. At low coagulant dose, i.e., up to 10 mg/1 as Fe, ferrate and ferrous iron remove turbidity more efficiently than ferric iron. At the optimum ferrate dose of 5 mg/l as Fe, greater turbidity removal is achieved with ferrate than with either ferrous or ferric iron. In order to accomplish the same degree of coagulation, a coagulant dosage greater than 15 mg/l as Fe must be used with both ferrous and ferric iron. Furthermore, at ferrate doses greater than 8 mg/l as Fe, the final turbidities of the coagulated system increase with increasing ferrate dose. This behavior contrasts the trend of increasing turbidity removal with increasing dose at low ferrate concentrations and at all tested doses of iron (11) and iron (111). The ferrate system exhibits either colloid restabilization or the formation of a stable hydrous iron oxide colloid with increasing ferrate concentration. Data collected from a small (10 l/h) bench-scale pilot plant also indicated that ferrate can be an effective chemical for suspended solids removal in tertiary treatment of secondary effluent [24]. The pilot facility included a flash nux reactor, where potassium ferrate, of 90% or greater purity, was added, a flocculation unit, and a sedimentation clarifier. The residence time in the flash mix unit was approximately two minutes, and in the three chambered flocculation unit, approximately 45 minutes. The total sedimentation time in the final clarifier wzs four hours. The plant was operated at steady state
--
t
Ferrate concentratlon
Orng/L
2Q/L
o-.- Cmg/L
6mg/~ 8 mg/L c.....1.0 mg/L c..-
'\.
v)
U
L C
-H!%lVG
Fig. 5. Coagulation of secondary effluent in a pilot treatment plant.
416 throughout testing. Figure 5 shows suspended solids removal through the bench-scale pilot plant at various ferrate doses. Despite the variability of secondary effluent quality, the data demonstrate that the pilot system operated optimally at a ferrate dose of 8 mg/l with approximately 86% suspended solids removal. Better than 80% solids removal was observed in all tests utilizing ferrate in concentrations greater than 6 iiigil. However, higher ferrate dosages seemed t o contribute to the turbidity of the effluent stream. This turbidity was not measurable in the solids deternunation, and was obviously due to the formation of colloidal iron from ferrate addition. The behavior of ferrate in water systems with a wide array of organic and inorganic constituents is extremely complex, as a combination of oxidation, coagulation, and precipitation reactions occur. In order t o focus singly on ferrate’s coagulative capabilities. current research is investigating the behavior of ferrate in a defined colloidal system, buffered at various pH and maintained at constant ionic strength. The objective of this new research is t o consider the operative behavior of ferrate (Vl) decay in coagulation in contrast to t h e behavior of iron (11) and iron (111) salts. Since, in practice, the solubility limit o f metal hydroxide is normally exceeded by the dosage of metal salts required t o destabilize colloids, careful consideration has been given to the kinetics of hydrous iron oxide generation for iron (VI), iron (HI), and iron (11) salts. Therefore, an approach has been developed to consider the decay of the ferrate ion, and the precipitation of the resultant metal hydroxide in comparison t o the iron precipitation reactions of iron (11) and iron (111) salts. F o r equivalent dosages of iron, the rate of turbidity generation is believed to reflect the duration of soluble iron species, intermediate t o the formation of iron hydroxide. The formation of hydrous iron oxide with time, as measured by light scattering, has been studied at a variety of iron salt doses and pH. The appropriate dose of iron was added t o one liter of biocarbonate buffer adjusted t o the specified pH and ionic strength. After disper;ing the dose throughout the volume, a small aliquot was transferred to a sealed vial where iron precipitation and sedimentation was followed under quiescent conditions. Data from experiments conducted for potassium ferrate. ferrous sulfate, and ferric nitrate at doses of 15 mg/l as Fe are presented in Figure 6. These results reflect the rate o f generation of an insoluble phase in aqueous solution, buffered at pH 7. If the turbidity profile of iron (111) is regarded as a standard illustration of the hydrolysis, olation, and precipitation of the ferric ion, then comparison between this and profiles similarly developed for iron (11) and iron (111) will demonstrate any differences in the reaction kinetics of iron hydrolysis. Although turbidity measured by light scattering represents the net result of solid iron formation, coagulation, and sedimentation, the formation of solid iron is considered to be the predominant mechanism in iron (11) and iron (Vl) systems within the first few minutes o f turbidity monitoring. Figure 6 shows that a ferric salt generated a maximum m u n t of turbidity within 30 seconds o f the salt addition. Flocs were visible immediately, and throughout the rest of the analysis, turbidity was seen to decrease, due to the flocculation and settling o f the hydrous iron oxide. The turbidity profile of the same dose of ferrous salt in the buffered aqueous system describes behavior different from that o f the iron (111) system. The Fe” must undergo oxidation to Fe3+ prior to hydrolysis, polymerization, and precipitation of an insoluble iron. Under atmospheric conditions at pH = 7 significant oxidation o f iron (11) takes approximately 20 minutes [25]. However, t h e turbidity profile in Figure 6 may reflect t h e competitive for-
41 7
15mg/L a s Fe pH 17
50 -
0-
x--
0-.-
Fe(NO3I3 Fe S O 4 . 7H20 K 2Fe O4
-z 3 5 x
c
+? 3
I-
i
30-
/
25. 20
-
0
i
/
5
10
15
rime [ m i n u t e ]
Fig. 6. Formation of insoluble phase by different iron salts.
mation of a ferrous carbonate complex due to the use of a sodium bicarbonate buffer. Visible flocs were not observed for the first 60 minutes of monitoring, although the color of the system deepended with time. The iron (11) system generated an iron colloid which was stable for approximately two hours, at which time fine flocs became visible and rapidly settled. The ferrate (VI) system demonstrated turbidity formation distinct from both the iron (11) and iron (111) systems. Decomposition of the ferrate ion occurred within the first 3.5 minutes, and fine flocs were observed after the first minute of monitoring. Initially, turbidity increases at a rate slightly greater than that of the ferrous system, but this rate slows after two minutes and an iron colloid is formed. This colloid was observed to be stable for approximately 90 minutes before significant settling occurred. Figure 7 reports data collected from colloid destabilization experiments. A colloidal silica suspension was developed, buffered at pH 7, and maintained at constant ionic
418
\
\
\
/
-s
100
- 90 a,
c
5 80
E"
f
\
\
70
1 ,
c
{
60
?
50
f
\
40 30 20 10
t
10
20
LO 30 Time [ m l n u t e l
50
60
Fig. 7. Destabilization of a silica colloid by different iron salts.
strength. Colloidal destabilization was determined for each iron salt at 15 mg/l as F e and measured as the amount of turbidity remaining in the systems at various times. After the addition of each salt, t o one liter of colloid suspension stabilized over 24 hours, the system was rapidly mixed for one minute, slowly flocculated for 30 minutes, and allowed t o settle for 30 minutes. Ferrate (VI) achieved the greatest amount of turbidity removal, 89% after 30 minutes of settling, while ferrous and ferric salts produced 76% and 58% removal, respectively. It is thought that ferrate accomplished the largest degree of colloid destabilization because o f the greater coagulating efficiency o f soluble species intermediate to its decay t o Fe3'. The 20% decrease in turbidity within one minute of ferrate addition may be due t o the surface activity of these species. After the initial flash mixing, turbidity increased t o a constant level throughout flocculation. This increase reflected the growth o f metal colloid flocs. The iron (11) system, which generated a great deal of turbidity in the ab-
419 sence of the silica colloid, showed a very small, initial decrease,in turbidity and a large increase in turbidity with flocculation. Throughout flocculation very large flocs were visible. The iron (111) system, which had demonstrated immediate hydrous iron oxide formation and gradual settling, behaved similarly under colloidal destabilizing conditions. There was no increase in turbidity with flocculation, and the least amount of silica turbidity was removed. In fact, only in the system destabilized with iron (111) was the white silica colloid still visible after 30 minutes of settling. In conclusion, both ferrate (VI) and iron (11) were better coagulants than iron (111) for this test system. This may be due to the duration of soluble species and the lesser rate of solid iron generation. Also, ferrate (VI) was observed to remove more turbidity from this model colloid system than iron (11). There may have been some diminishment in the coagulation ability of the ferrous salt due to the interference of the bicarbonate buffer. However, most of the iron (11) is believed to have been oxidized to iron (111) under the conditions of this experiment. It one considers the behavior of iron (11) and iron (VI) in buffered aqueous systems, it may be that the stability of the solid iron generated by iron (11) and iron (VI) is mediated by the variety of soluble iron species intermediate to the transition to ferric hydroxide. It may be those differences in the nature of the soluble species generated by ferrate and ferrous salts that also mediates coagulation efficiency. It is also interesting to note that even though a large amount of turbidity was generated by the iron (11) system it did not effectively coagulate the silica colloid. These observations support the theory that the most efficient mechanism of coagulation proceeds via adsorption of soluble aquometal species at the colloid surface.
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