PHYSICAL - CHEMICAL TREATMENT OF WATER AND WASTEWATER - CHAPTER 13 pdf

Removal of Iron and Manganese by Chemical Precipitation Iron concentrations as low as 0.3 mg/L and manganese concentrations as low as 0.05 mg/L can cause dirty water complaints. At these concentrations, the water may appear clear but imparts brownish colors to laundered goods. Iron also affects the taste of beverages such as tea and coffee. Manganese flavors tea and coffee with medicinal tastes. Some types of bacteria derive their energy by utilizing soluble forms of iron and manganese. These organisms are usually found in waters that have high levels of iron and manganese in solution. The reaction changes the species from soluble forms into less soluble forms, thus causing precipitation and accumulation of black or reddish brown gelatinous slimes. Masses of mucous iron and manganese can clog plumbing and water treatment equipment. They also slough away in globs that become iron or manganese stains on laundry. Standards for iron and manganese are based on levels that cause taste and staining problems and are set under the Environmental Protection Agency Secondary Drinking Water Standards (EPA SDWA). They are, respectively, 0.3 mg/L for iron and 0.05 mg/L for manganese. Iron and manganese are normally found in concentrations not exceed- ing 10 mg/L and 3 mg/L, respectively, in natural waters. Iron and manganese can be found at higher concentrations; however, these conditions are rare. Iron concen- trations can go as high as 50 mg/L. Iron and manganese may be removed by reverse osmosis and ion exchange . The unit operation of reverse osmosis was discussed in a previous chapter; the unit process of ion exchange is discussed in a later chapter. This chapter discusses the removal of iron and manganese by the unit process of chemical precipitation. One manufacturer claimed that these elements could also be removed by bio- logical processes. It claimed that a process called the bioferro process encourages the growth of naturally occurring iron assimilating bacteria, such as Gallionella ferruginea , thus reducing iron concentration. An experimental result shows a reduc- tion of iron from 6.0 mg/L to less than 0.1 mg/L. They also claimed that a companion process called bioman could remove manganese down to 0.08 mg/L. This uses naturally occurring manganese bacteria to consume manganese. Figure 13.1 is a photograph showing growths of “bioferro” and “bioman” bacteria. 13 TX249_frame_C13.fm Page 593 Friday, June 14, 2002 2:32 PM © 2003 by A. P. Sincero and G. A. Sincero 594 13.1 NATURAL OCCURRENCES OF IRON AND MANGANESE Iron and manganese have the electronic configurations of [Ar]3 d 6 4 s 2 and [Ar]3 d 5 4 s 2 , and are located in Groups VIIIB and VIIB of the Periodic Table, respectively. They are both located in the fourth period. [Ar] means that these elements have the electronic configuration of the noble gas argon. The letters d and s refer to the d and s orbitals; the superscripts indicate the number of electrons that the orbitals contain. Thus, the d orbital of iron contains 6 electrons and that of manganese contains 5 electrons. Both elements contain 2 electrons in their s orbitals. This means that in their most reduced positive state, they acquire oxidation states of 2 + . Also, because of the d orbitals, they can form a number of oxidation states. The multiplicity of oxidation states give iron and manganese the property of imparting colors such as the imparting of brownish colors to laundered goods. Surface waters always contain dissolved oxygen in it. Thus, iron and manganese would not exist in their most reduced positive state of 2 + in these waters. The reason is that they will simply be oxidized to higher states of oxidation by the dissolved oxygen forming hydroxides and precipitate out. Groundwater is a source where these elements could come from. Groundwaters occurring deep down in the earth can become devoid of oxygen, thus, any iron or manganese present would have to be reduced. Therefore, the waters where removal of iron and manganese could be undertaken are groundwaters and the form of the elements are in the 2 + oxidation states, Fe(II) and Mn(II), respectively FIGURE 13.1 A photograph of “bioferro” and “bioman” bacteria. TX249_frame_C13.fm Page 594 Friday, June 14, 2002 2:32 PM © 2003 by A. P. Sincero and G. A. Sincero 595 13.2 MODES OF REMOVAL OF IRON AND MANGANESE The best place to investigate for determining the mode by which the elements can be removed is the table of solubility products constants as shown in Table 13.1. In general, a precipitation product that has the lowest K sp means that the substance is the most insoluble. As shown in the table, for iron, the lowest K sp is that of Fe(OH) 3 , an Fe(III) iron, and has the value of 1.1(10 − 38 ). For manganese, the lowest K sp is that of MnS, an Mn(II) manganese, and has the value of 1.1(10 − 22 ). These K sp ’s indicate that the elements must be removed in the form of ferric hydroxide and manganese sulfide, respectively; however, from the table, manganese can also be removed as Mn(OH) 2 at a K sp = 4.5(10 − 14 ). Of course, lime has many uses, while sulfide has only few. Sodium sulfide is used in photographic film devel- opment; however, lime is used in water and wastewater treatment, as an industrial chemical, as well as being used in agriculture. Thus, because of its varied use, lime is much cheaper. In addition, using a sulfide to remove iron and manganese would be a new method. Its health effect when found in drinking water is not documented. On the other hand, lime has been used for years. We will therefore use lime as the precipitant for the removal of iron and manganese. The probable use of sulfide in removing iron and manganese could be a topic for investigation in applied research. 13.3 CHEMICAL REACTIONS OF THE FERROUS AND THE FERRIC IONS The chemical reactions of the ferrous and the ferric ions were already discussed in a previous chapter. From the topic in the preceding section, iron is more efficiently removed as ferric hydroxide. The natural iron is in the form of Fe(II), so this ferrous must therefore oxidize to the ferric form in order to precipitate as the ferric hydroxide, if, in fact, the iron is to be removed in the ferric form. In Chapter 12, this was done using the dissolved oxygen that is relatively abundant in natural waters. It must be TABLE 13.1 Solubility Product Constant of Iron and Manganese Precipitation Products Precipitation Product Solubility Product, K sp Fe(OH) 2 3.16(10 −15 ) FeCO 2 2.11(10 −11 ) FeS 8(10 −26 ) Fe(OH) 3 1.1(10 −38 ) Mn(OH) 2 4.5(10 −14 ) MnCO 3 8.8(10 −11 ) MnS 4.3(10 −22 ) TX249_frame_C13.fm Page 595 Friday, June 14, 2002 2:32 PM © 2003 by A. P. Sincero and G. A. Sincero 596 noted, however, that based on the K sp values, iron can also be removed in the ferrous form as Fe(OH) 2 . For convenience, the reactions are reproduced next. Ferrous: (13.1) (13.2) (13.3) (13.4) Ferric: (13.5) (13.6) (13.7) (13.8) (13.9) Remember that the values of the solubility product constants, and and all the other equilibrium constants for the complex ions apply only at 25 ° C. The presence of the complex ions increase the solubility of the iron species and therefore increase the concentration of these species in solution. For the ferrous and the ferric species, they are sp FeII and sp FeIII and were derived in Chapter 12, respec- tively, as (13.10) (13.11) Fe(OH) 2 s()  Fe 2+ 2OH − K sp,Fe OH() 2 + 10 −14.5 = Fe(OH) 2 s()  FeOH + OH − K FeOHc + 10 −9.4 = Fe(OH) 2 s() OH −  Fe OH() 3 − K Fe OH() 3 c + 10 −5.1 = Fe(OH) 2 1 4 O 2 1 2 H 2 O → Fe(OH) 3 ↓ (conversion from ferrous to ferric)++ Fe(OH) 3 s()  Fe 3+ 3OH − K sp,Fe OH() 3 + 10 −38 = Fe(OH) 3 s()  FeOH 2+ 2OH − K FeOHc + 10 −26.16 = Fe(OH) 3 s()  Fe(OH) 2 + OH − K Fe OH() 2 c + 10 −16.74 = Fe(OH) 3 s() OH −  Fe OH() − 4 K Fe OH() 4 c + 10 −5 = 2Fe(OH) 3 s()  Fe 2 OH() 2 4+ 4OH − K Fe 2 OH() 2 c + 10 −50.8 = K sp,Fe OH() 2 K sp,Fe OH() 3 , sp FeII []Fe 2+ []FeOH + []Fe OH() 3 − []++= K sp Fe OH() 2 , γ H 2 H + [] 2 γ FeII K w 2 K FeOHc γ H H + [] γ FeOHc K w K Fe OH() 3 c K w γ Fe OH() 3 c γ H H + [] ++= sp FeIII []Fe 3+ []FeOH 2+ []Fe OH() 2 + []Fe OH() 4 − []2Fe 2 OH() 2 4+ []++++= K sp Fe OH() 3 , γ H 3 H + [] 3 γ FeIII K w 3 K FeOHc γ H 2 H + [] 2 γ FeOHc K w 2 K Fe OH() 2 c γ H H + [] γ Fe OH() 2 c K w K Fe OH() 4 c K w γ Fe OH() 4 c H + [] ++ += 2K Fe 2 OH() 2 c γ H 4 H + [] 4 γ Fe 2 OH() 2 c K w 4 + TX249_frame_C13.fm Page 596 Friday, June 14, 2002 2:32 PM © 2003 by A. P. Sincero and G. A. Sincero Example 13.1 From the respective optimum pH’s of 11.95 and 8.2 for sp FeII and sp FeIII , calculate the concentrations [sp FeII ] and [sp FeIII ], respectively. Assume the water contains 140 mg/L of dissolved solids. Solution: Therefore, sp FeII [] K sp,Fe OH() 2 γ H 2 H + [] 2 γ FeII K w 2 K FeOHc γ H H + [] γ FeOHc K w K Fe OH() 3 c K w γ Fe OH() 3 c γ H H + [] ++= K sp Fe OH() 2 , 10 14.5– = µ 2.5 10 −5 ()TDS γ 10 0.5z i 2 µ () 1+1.14 µ () – == µ 2.5 10 −5 ()140()3.5 10 3– ()== γ H γ FeOHc γ Fe OH() 3 c 10 0.5 1() 2 3.5 10 3– ()[] 1+1.14 3.5 10 3– ()[] – 0.94== = = γ FeII 10 0.5 2() 2 3.5 10 3– ()[] 1+1.14 3.5 10 3– ()[] – 0.77K w 10 14– ()=== K FeOHc 10 9.4– K Fe(OH ) 3 c 10 5.1– == sp FeII [] 10 14.5– 0.94() 2 10 11.95– [] 2 0.77()10 14– () 2 −−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−− 10 9.4– ( ) 0.94()10 11.95– [] 0.94() 10 14– () −−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−− 10 5.1– ()10 14– () 0.94( ) 0.94()10 11.95– [] −−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−++ = 4.57 10 11– ()4.47 10 8– ()8.01 10 8– ()++ 1.0 10 7– () gmol/L Ans== 0.0056 mg/L= sp FeIII [] K sp,Fe OH() 3 γ H 3 H + [] 3 γ FeIII K w 3 K FeOHc γ H 2 H + [] 2 γ FeOHc K w 2 K Fe OH() 2 c γ H H + [] γ Fe OH() 2 c K w K Fe OH() 4 c K w γ Fe OH() 4 c H + [] ++ += 2K Fe 2 OH() 2 c γ H 4 H + [] 4 γ Fe 2 OH() 2 c K w 4 + K sp,Fe OH() 3 10 38– γ FeIII 10 0.5 3() 2 3.5 10 3– ()[] 1+1.14 3.5 10 3– ()[] – 0.56== =K FeOHc 10 26.16– = γ FeOHc 10 0.5 2() 2 3.5 10 3– ()[] 1+1.14 3.5 10 3– ()[] – 0.77 K Fe OH() 2 c 10 16.74– == = γ Fe OH() 2 c 10 0.5 1() 2 3.5 10 3– ()[] 1+1.14 3.5 10 3– ()[] – 0.94== TX249_frame_C13.fm Page 597 Friday, June 14, 2002 2:32 PM © 2003 by A. P. Sincero and G. A. Sincero Therefore, 13.3.1 PRACTICAL OPTIMUM pH RANGE FOR THE REMOVAL OF FERROUS AND FERRIC As shown in Chapter 12, at 25°C and at a solids concentration of 140 mg/L, the optimum pH’s correspond to 11.95 and 8.2 (or around 12 and 8), respectively, for ferrous and ferric. The respective concentrations for sp FeII and sp FeII at these condi- tions as obtained in the previous example are [sp FeII ] = 0.0056 mg/L and [sp FeIII ] = 0.0000016 mg/L. A pH range exists, however, at which units used for the removal of the elements can be operated and effect good results. This range is called the practical optimum pH range. Tables 13.2 through 13.5 show the respective concentrations of sp FeII and sp FeIII at other conditions of pH and total solids. The values for [sp FeII ] were obtained using Equation (13.10) and the values for [sp FeIII ] were obtained using Equation (13.11). Note that these equations require the values of the activity coefficients of the ions. The activity coefficients are needed by the equations and, since activity coefficients are functions of the dissolved solids, dissolved solids are used as parameters in the tables, in addition to pH. The previous tables indicate that the total solids (or equivalently, the activities of the ions) do not have a significant effect on the optimum pH values, which for sp FeII remain at about 12.0 and for sp FeIII remain at about 8.0. For practical purposes, however, the practical optimum pH for sp FeII ranges from 11 to 13 and for sp FeIII , it ranges from 5.0 to 13.0. Note that for the range of pH for sp FeII , it is assumed the element is to be removed as Fe(OH) 2 . If it is to be removed as Fe(OH) 3 , the pH of the solution during its oxidation by dissolved oxygen or any oxidizer need not be adjusted since the practical optimum pH for the precipitation of ferric hydroxide varies over a wide range from 5.0 to 13.0 and already includes the range for the ferrous removal. K Fe(OH) 4 c 10 5– γ Fe(OH) 4 c γ Fe(OH) 2 c 0.94 K Fe 2 (OH) 2 c 10 50.8– === = γ Fe 2 (OH) 2 c 10 0.5 4() 2 3.5 10 3– ()[] 1+1.14 3.5 10 3– ()[] – 0.36== sp FeIII [] 10 38– ()0.94() 3 10 8.2– [] 3 0.56()10 14– () 3 −−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−− 10 26.16– ( ) 0.94() 2 10 8.2– [] 2 0.77()10 14– () 2 −−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−− 10 16.74– ( ) 0.94()10 8.2– [] 0.94()10 14– () −−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−++ = 10 5– ()10 14– () 0.94()10 8.2– [] 2 10 50.8– ()0.94() 4 10 8.2– [] 4 0.36()10 14– () 4 ++ 2.09 10 63– () 5.6 10 43– () 2.43 10 43– () 7.7 10 29– () 1.08 10 25– () 9.4 10 15– () 1.0 10 19– () 5.93 10 9– () 3.92 10 84– () 3.6 10 57– () ++++= 3.73 10 21– ()3.16 10 15– ()1.15 10 11– ()1.69 10 11– ()1.08 10 27– ()++++= 2.84 10 11– () gmol/L 0.0000016 mg/L Ans== TX249_frame_C13.fm Page 598 Friday, June 14, 2002 2:32 PM © 2003 by A. P. Sincero and G. A. Sincero TABLE 13.2 Concentration of sp FeII as a Function of pH at 25°° °° C pH, Dissolved Solids = 140 mg/L [sp FeII ], mg/L 0 5.0(10 +18 ) 1 2.0(10 +16 ) 2 2.0(10 +14 ) 3 2.0(10 +12 ) 4 2.0(10 +10 ) 5 2.0(10 +8 ) 6 2.0(10 +6 ) 7 2.0(10 +4 ) 8 2.2(10 +2 ) 9 4.2(10 +0 ) 10 2.4(10 −1 ) 11 2.3(10 −2 ) 12 7.3(10 −3 ) 13 5.1(10 −2 ) 14 5.0(10 −1 ) 15 5.0(10 +0 ) TABLE 13.3 Concentration of sp FeII as a Function of pH at 25°° °° C pH, Dissolved Solids = 35,000 mg/L [sp FeII ], mg/L 0 5.0(10 +18 ) 1 5.0(10 +16 ) 2 5.0(10 +14 ) 3 5.0(10 +12 ) 4 5.0(10 +10 ) 5 5.0(10 +8 ) 6 5.0(10 +6 ) 7 5.0(10 +4 ) 8 5.2(10 +2 ) 9 7.2(10 +0 ) 10 2.7(10 −1 ) 11 2.4(10 −2 ) 12 1.5(10 −2 ) 13 1.3(10 −1 ) 14 1.3(10 +0 ) 15 1.3(10 +1 ) TX249_frame_C13.fm Page 599 Friday, June 14, 2002 2:32 PM © 2003 by A. P. Sincero and G. A. Sincero TABLE 13.4 Concentration of sp FeIII as a Function of pH at 25°° °° C pH, Dissolved Solids = 140 mg/L [sp FeII ], mg/L 0 1.2(10 +10 ) 1 8.6(10 +5 ) 2 1.3(10 +3 ) 3 5.3(10 +0 ) 4 5.5(10 −2 ) 5 1.4(10 −3 ) 6 1.1(10 −4 ) 7 1.0(10 −5 ) 8 1.6(10 −6 ) 9 6.0(10 −6 ) 10 6.0(10 −5 ) 11 6.0(10 −4 ) 12 6.0(10 −3 ) 13 6.0(10 −2 ) 14 6.0(10 −1 ) 15 6.0(10 +0 ) TABLE 13.5 Concentration of sp FeIII as a Function of pH at 25°° °° C pH, Dissolved Solids = 35,000 mg/L [sp FeIII ], mg/L 0 1.2(10 +10 ) 1 1.2(10 +7 ) 2 1.3(10 +4 ) 3 2.3(10 +1 ) 4 1.2(10 −1 ) 5 1.1(10 −3 ) 6 1.1(10 −5 ) 7 2.0(10 −7 ) 8 9.4(10 −7 ) 9 9.4(10 −6 ) 10 9.4(10 −5 ) 11 9.4(10 −4 ) 12 9.4(10 −3 ) 13 9.4(10 −2 ) 14 9.4(10 −1 ) 15 9.4(10 +0 ) TX249_frame_C13.fm Page 600 Friday, June 14, 2002 2:32 PM © 2003 by A. P. Sincero and G. A. Sincero The previous optimum values of pH apply only at 25°C. As indicated in the formulas, equilibrium constants are being used to compute these values. The values of the equilibrium constants vary with temperature, so the optimum range at other temperatures would be different. Equilibrium constants at other temperatures can be calculated using the Van’t Hoff equation which, however, requires the values of the standard enthalpy changes. At present, these values are unavailable making values of optimum pH range at other temperatures impossible to calculate. For this reason, the pH range found above must be modified to a conservative range. Hence, adopt the following: for ferrous removal as Fe(OH) 2 , 11.5 ≤ optimum pH ≤ 12.5 and for ferrous removal as Fe(OH) 3 , 5.5 ≤ optimum pH ≤ 12.5. 13.4 CHEMICAL REACTIONS OF THE MANGANOUS ION [Mn(II)] Manganese can be removed as Mn(OH) 2 using a suitable source. Upon intro- duction of the hydroxide source, however, it is not only this solid that is produced. Manganese forms complex ions with the hydroxide. The complex equilibrium reac- tions are as follows (Snoeyink and Jenkins, 1980): (13.12) (13.13) (13.14) (13.15) The values of the equilibrium constants given above are at 25°C. The complexes are Mn(OH) + , Mn and Mn . Also note that the ion is a participant in these reactions. This means that the concentrations of each of these complex ions are determined by the pH of the solution. In the application of the previous equations in an actual treatment of water, conditions must be adjusted to allow maximum precip- itation of the solid represented by Mn(OH) 2(s) . To allow for this maximum precipita- tion, the concentrations of the complex ions must be held to the minimum. The pH corresponding to this condition is the optimum pH. From the previous reactions, the equivalent mass of Mn 2+ is Mn/2 = 27.45. 13.4.1 DETERMINATION OF THE OPTIMUM pH Let sp Mn represent the collection of species standing in solution containing the Mn(II) species. Thus, (13.16) All the concentrations in the right-hand side of the above equation will now be expressed in terms of the hydrogen ion concentration. This will result in expressing OH − Mn(OH) +  Mn 2+ OH − K MnOHc + 10 3.4– = Mn(OH) 2 s()  Mn 2+ 2OH − K sp,Mn(OH) 2 + 4.5 10 14– ()= Mn(OH) 2 0  Mn 2+ 2OH − K Mn(OH) 2 c + 10 6.8– = Mn(OH) 3 −  Mn 2+ 3OH − K Mn(OH) 3 c + 10 7.8– = (OH) 2 0 , (OH) 3 − OH − sp Mn []Mn 2+ []Mn(OH) + []Mn(OH) 2 0 []Mn(OH) 3 − []+++= TX249_frame_C13.fm Page 601 Friday, June 14, 2002 2:32 PM © 2003 by A. P. Sincero and G. A. Sincero [sp Mn ] in terms of the hydrogen ion. Differentiating the resulting equation of [sp Mn ] with respect to [H + ] and equating the result to zero will produce the minimum concentration of sp Mn and, thus, the optimum pH determined. Using the equations and equilibrium constants of Eqs. (13.12) through (13.15), along with the ion product of water, we proceed as follows: (13.17) (13.18) (13.19) (13.20) γ Mn , γ MnOHc , γ Mn(OH)2c , and γ Mn(OH)3c are, respectively, the activity coefficients of Mn(II) and the complexes Mn(OH) + , Mn , and Mn . is the solubility product constant of the solid Mn(OH) 2(s) . K MnOHc , and are, respectively, the equilibrium constants of the complexes Mn(OH) + , Mn and Mn . Equations (13.17) through (13.20) may now be substituted into Equation (13.16) to produce (13.21) Differentiating with respect to [H + ], equating to zero, rearranging, and changing H + to , the concentration of the hydrogen ion at optimum conditions, the following expression is produced: (13.22) The value of [H opt ] may be solved by trial error. Mn 2+ [] Mn 2+ {} γ Mn K sp Mn, OH() 2 γ Mn OH − {} 2 K sp Mn, OH() 2 H + {} 2 γ Mn K w 2 K sp Mn, OH() 2 γ H 2 H + [] 2 γ Mn K w 2 == = = Mn OH() + [] Mn OH() + {} γ Mn OH()c Mn 2+ {}OH − {} γ Mn OH()c K MnOHc K sp Mn, OH() 2 γ H H + [] γ Mn OH()c K MnOHc K w == = Mn(OH) 2 0 [] Mn(OH) 2 0 {} γ Mn(OH) 2 c=1 Mn(OH) 2 0 {} Mn 2+ {}OH − {} 2 K Mn OH() 2 c K sp Mn, OH() 2 K Mn OH() 2 c === = Mn(OH) 3 − [] Mn(OH) 3 − {} γ Mn OH() 3 c Mn 2+ {}OH − {} 3 γ Mn OH() 3 c K Mn OH() 3 c K sp Mn, OH() 2 K w γ Mn OH() 3 c K Mn OH() 3 c γ H H + [] == = (OH) 2 0 OH() 3 − K sp,Mn(OH) 2 K Mn(OH) 2 c , K Mn(OH) 3 c (OH) 2 0 ,OH() 3 − sp Mn [] K sp Mn, OH() 2 γ H 2 H + [] 2 γ Mn K w 2 K sp Mn, OH() 2 γ H H + [] γ Mn OH()c K MnOHc K w K sp Mn, OH() 2 K Mn OH() 2 c ++= K sp Mn, OH() 2 K w γ Mn OH() 3 c K Mn OH() 3 c γ H H + {} + H opt + 2 γ H 2 H opt + [] 3 γ Mn K w 2 γ H H opt + [] 2 γ Mn OH()c K MnOHc K w + K w γ Mn OH() 3 c K Mn OH() 3 c γ H = TX249_frame_C13.fm Page 602 Friday, June 14, 2002 2:32 PM © 2003 by A. P. Sincero and G. A. Sincero [...]... calculate the activity coefficient of Mn(OH) At a pH of 11.97 and dissolved solids of 140 mg/L, the concentration of − spMn is 0.0179 mg/L Calculate the equilibrium constant of Mn ( OH ) 3 © 2003 by A P Sincero and G A Sincero TX249_frame_C13.fm Page 623 Friday, June 14, 2002 2:32 PM 13. 19 13. 20 13. 21 13. 22 13. 23 13. 24 13. 25 13. 26 13. 27 13. 28 13. 29 13. 30 13. 31 13. 32 13. 33 13. 34 The temperature is 25°C... the hydrogen ion The optimum pH for precipitating Mn(OH)2 is 11.97 and the dissolved solids content of the water is 140 mg/L The temperature is 25°C Calculate © 2003 by A P Sincero and G A Sincero TX249_frame_C13.fm Page 622 Friday, June 14, 2002 2:32 PM 13. 9 13. 10 13. 11 13. 12 13. 13 13. 14 13. 15 13. 16 13. 17 13. 18 the ion product of water The given temperature may be used to calculate other parameters... - V φa   – pH (13. 62) – pH  10 to – 10 cur  M HNO3 pH = 63.01  [ A cadd ] geq = [ A cur ] geq + - V φa   (13. 63) And, also, gleaning from Equation (13. 55) and the respective equivalent masses of NaOH and Na2CO3 and the cubic meters of water treated, V , – pH – pH  10 cur – 10 to  = 40.0  [ A ccur ] geq + - V φb   M NaOHpH – pH (13. 64) – pH... Removal of Soluble Iron and Manganese Am Water Works Assoc., Denver McNeill, L and M Edwards (1997) Predicting as removal during metal hydroxide precipitation J Am Water Works Assoc., 89, 1, 75–86 Perry, R H and C H Chilton (1973) Chemical Engineers’ Handbook p (3 138 ) McGrawHill, New York, 3 138 Ranganathan, K and C Namasivayam (1998) Utilization of Waste Fe(III)/Cr(III) hydroxide for removal of Cr(VI) and. .. Hedberg, T and T A Wahlberg (1997) Upgrading of waterworks with a new biooxidation process for removal of manganese and iron Water Science Technol., Proc 1997 Int Conf on Upgrading of Water and Wastewater Syst., May 25–28, Kalmir, Sweden, 37, 9, 121–126 Elsevier Science Ltd., Exeter, England Holtzclaw Jr., H F and W R Robinson (1988) General Chemistry D C Heath and Company, Lexington, MA, A7, 809 Knocke,... Cubic meters of water treated Activity coefficient of the hydrogen ion Activity coefficient of the ferrous ion Activity coefficient of the ferric ion + 2+ Activity coefficient of FeOH and FeOH + Activity coefficient of Fe ( OH ) 2 − Activity coefficient of Fe ( OH ) 3 − Activity coefficient of Fe ( OH ) 4 4+ Activity coefficient of Fe2 ( OH ) 2 Activity coefficient of Mn(II) ion + Activity coefficient of Mn(OH)... FeOH 2+ FeOH + Fe ( OH ) 2 − Fe ( OH ) 3 − Fe ( OH ) 4 Alkalinity to be added to water Gram equivalents per liter of alkalinity to be added to water Acidity to be added to water Gram equivalents per liter of acidity to be added to water Current alkalinity of water Gram equivalents per liter of current alkalinity of water mg/L of ferrous to be removed Monohydroxo Fe(II) complex ion Monohydroxo Fe(III) complex... the water 13. 7 CHEMICAL REQUIREMENTS The chemical requirements are those for the hydroxide source, chlorine, permanganate, ozone, and oxygen The discussion will be subdivided into requirements in the ferrous reactions and into requirements in the manganous reactions The treatment on chemical requirements, in effect, is reduced to the determination of the equivalent masses of the pertinent chemicals 13. 7.1... place and the filter are always necessary The unit operations of mixing, flocculation, settling, and filtration were already discussed in the earlier chapters of this book As can be deduced from the addition of lime, the unit operations above are the same as those that would be applied in water softening Thus, for reasons of economics, iron and manganese removal are normally incorporated into water softening... constant of Mn ( OH ) 3 At a pH of 11.97 and dissolved solids of 140 mg/L, the concentration of − spMn is 0.0179 mg/L Calculate the activity coefficient of Mn ( OH ) 3 The temperature is 25°C The data on dissolved solids may be used to calculate − other parameters but not to calculate the activity coefficient of Mn ( OH ) 3 At a pH of 11.97 and dissolved solids of 140 mg/L, the concentration of spMn . Fe[] mg V == TX249_frame_C13.fm Page 611 Friday, June 14, 2002 2:32 PM © 2003 by A. P. Sincero and G. A. Sincero (13. 45) (13. 46) (13. 47) (13. 48) (13. 49) (13. 50) (13. 51) (13. 52) (13. 53) Example 13. 4 A raw water contains. reproduced next. Ferrous: (13. 1) (13. 2) (13. 3) (13. 4) Ferric: (13. 5) (13. 6) (13. 7) (13. 8) (13. 9) Remember that the values of the solubility product constants, and and all the other equilibrium. reaction = mg/L of ferrous to be removed = mg/L of manganous removed = cubic meters of water treated Therefore, (13. 38) (13. 39) (13. 40) (13. 41) (13. 42) (13. 43) (13. 44) M Cl 2 Fe M Ca ClO() 2 Fe M KMnO 4 Fe M O 3 Fe M O 2 Fe M Cl 2 Mn M Ca