Standard Methods for Examination of Water & Wastewater_12 potx

Removal of Iron and Manganese by Chemical PrecipitationIron concentrations as low as 0.3 mg/L and manganese concentrations as low as0.05 mg/L can cause dirty water complaints.. If it is

Removal of Iron and Manganese by Chemical Precipitation

Iron concentrations as low as 0.3 mg/L and manganese concentrations as low as0.05 mg/L can cause dirty water complaints At these concentrations, the water mayappear clear but imparts brownish colors to laundered goods Iron also affects thetaste of beverages such as tea and coffee Manganese flavors tea and coffee withmedicinal tastes

Some types of bacteria derive their energy by utilizing soluble forms of ironand manganese These organisms are usually found in waters that have high levels

of iron and manganese in solution The reaction changes the species from solubleforms into less soluble forms, thus causing precipitation and accumulation of black

or reddish brown gelatinous slimes Masses of mucous iron and manganese can clogplumbing and water treatment equipment They also slough away in globs thatbecome iron or manganese stains on laundry

Standards for iron and manganese are based on levels that cause taste and stainingproblems and are set under the Environmental Protection Agency Secondary DrinkingWater Standards (EPA SDWA) They are, respectively, 0.3 mg/L for iron and 0.05 mg/Lfor 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 trations can go as high as 50 mg/L

concen-Iron and manganese may be removed by reverse osmosis and ion exchange Theunit operation of reverse osmosis was discussed in a previous chapter; the unitprocess of ion exchange is discussed in a later chapter This chapter discusses theremoval 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 encouragesthe 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 companionprocess called bioman could remove manganese down to 0.08 mg/L This usesnaturally occurring manganese bacteria to consume manganese Figure 13.1 is aphotograph showing growths of “bioferro” and “bioman” bacteria

13

and s orbitals; the superscripts indicate the number of electrons that the orbitalscontain 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 intheir 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 ofoxidation states give iron and manganese the property of imparting colors such asthe imparting of brownish colors to laundered goods

Surface waters always contain dissolved oxygen in it Thus, iron and manganesewould 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 dissolvedoxygen forming hydroxides and precipitate out Groundwater is a source where theseelements could come from Groundwaters occurring deep down in the earth can becomedevoid 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 aregroundwaters and the form of the elements are in the 2+ oxidation states, Fe(II)andMn(II), respectively

FIGURE 13.1 A photograph of “bioferro” and “bioman” bacteria

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 Ingeneral, a precipitation product that has the lowest K sp means that the substance isthe 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, anMn(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 ferrichydroxide and manganese sulfide, respectively; however, from the table, manganesecan also be removed asMn(OH)2 at a K sp = 4.5(10−14) Of course, lime has manyuses, 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 industrialchemical, 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 theprecipitant for the removal of iron and manganese The probable use of sulfide inremoving 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 efficientlyremoved as ferric hydroxide The natural iron is in the form of Fe(II), so this ferrousmust 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 doneusing 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)

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 andtherefore increase the concentration of these species in solution For the ferrous andthe ferric species, they are spFeIIand spFeIII and were derived in Chapter 12, respec-tively, as

Fe(OH)2+14O2+12H2O → Fe(OH)3↓ (conversion from ferrous to ferric)

Fe(OH)3 s( )
Fe3++3OH− K sp,Fe OH( )3 = 10−38Fe(OH)3 s( )
FeOH2++2OH− K FeOHc = 10−26.16Fe(OH)3 s( )
Fe(OH)2++OH− KFe OH( )2c = 10−16.74Fe(OH)3 s( ) OH−
Fe OH( )−

γFe2( OH )2c K w

4

+

-Example 13.1 From the respective optimum pH’s of 11.95 and 8.2 for spFeIIand spFeIII, calculate the concentrations [spFeII] and [spFeIII], respectively Assume thewater contains 140 mg/L of dissolved solids.

=

0.5z i2( µ ) 1+1.14 ( µ ) - –

-0.94

γFeII 10

0.5 2 ( ) 2 [ 3.5 10 ( 3) ] 1+1.14 3.5 10 3

–

−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−− 10

9.4 –

( ) 0.94 ( ) 10 – 11.95

0.94 ( ) (10–14)

−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−−− 10

5.1 –

( ) (10–14) 0.94

γFe2( OH )2c K w

4

+

0.56

0.5 2 ( ) 2 [ 3.5 10 ( 3) ] 1+1.14 3.5 10 3

–

γFe OH ( )2c 10

0.5 1 ( ) 2 [ 3.5 10 ( 3) ] 1+1.14 3.5 10 3

–

-0.94

13.3.1 P RACTICAL O PTIMUM pH R ANGE FOR THE R EMOVAL

OF F ERROUS AND F ERRIC

As shown in Chapter 12, at 25°C and at a solids concentration of 140 mg/L, theoptimum pH’s correspond to 11.95 and 8.2 (or around 12 and 8), respectively, for

ferrous and ferric The respective concentrations for spFeII and spFeII at these

condi-tions as obtained in the previous example are [spFeII] = 0.0056 mg/L and [spFeIII] =

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 spFeII and spFeIII

at other conditions of pH and total solids The values for [spFeII] were obtained using

Equation (13.10) and the values for [spFeIII] 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 coefficientsare functions of the dissolved solids, dissolved solids are used as parameters in thetables, 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

spFeII remain at about 12.0 and for spFeIII remain at about 8.0 For practical purposes,

however, the practical optimum pH for spFeII ranges from 11 to 13 and for spFeIII, it

ranges from 5.0 to 13.0 Note that for the range of pH for spFeII, it is assumed theelement is to be removed as Fe(OH)2 If it is to be removed as Fe(OH)3, the pH ofthe solution during its oxidation by dissolved oxygen or any oxidizer need not beadjusted since the practical optimum pH for the precipitation of ferric hydroxidevaries over a wide range from 5.0 to 13.0 and already includes the range for theferrous removal

–

50.8 –

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

TABLE 13.2

Concentration of spFeII as a Function of pH at 25 °°°°C

pH, Dissolved Solids = 140 mg/L [spFeII ], mg/L

pH, Dissolved Solids = 35,000 mg/L [spFeII ], mg/L

TABLE 13.4

Concentration of spFeIII as a Function of pH at 25 °°°°C

pH, Dissolved Solids = 140 mg/L [spFeII ], mg/L

pH, Dissolved Solids = 35,000 mg/L [spFeIII ], mg/L

The previous optimum values of pH apply only at 25°C As indicated in theformulas, equilibrium constants are being used to compute these values The values

of the equilibrium constants vary with temperature, so the optimum range at othertemperatures 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, adoptthe following: for ferrous removal as Fe(OH)2, 11.5 ≤ optimum pH ≤ 12.5 and forferrous 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 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):

intro-(13.12)(13.13)(13.14)(13.15)The values of the equilibrium constants given above are at 25°C The complexesare 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 ionsare 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 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 pHcorresponding to this condition is the optimum pH From the previous reactions, theequivalent mass of Mn2+ is Mn/2 = 27.45

precip-13.4.1 D ETERMINATION OF THE O PTIMUM pH

Let spMn 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 beexpressed in terms of the hydrogen ion concentration This will result in expressing

OH−

Mn(OH)+
Mn2++OH− K MnOH c = 10–3.4Mn(OH)2 s( )
Mn2++2OH− K sp,Mn(OH)2 = 4.5 10( –14)Mn(OH)20
Mn2++2OH− KMn(OH)2c = 10–6.8Mn(OH)3

[spMn] in terms of the hydrogen ion Differentiating the resulting equation of [spMn]with respect to [H+] and equating the result to zero will produce the minimum

concentration of spMn 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)

γMn, γMnOHc, γMn(OH)2c, and γMn(OH)3c are, respectively, the activity coefficients

of Mn(II) and the complexes Mn(OH)+, Mn , and Mn isthe solubility product constant of the solid Mn(OH)2(s) K MnOHc, and

are, respectively, the equilibrium constants of the complexes Mn(OH)+,

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 followingexpression is produced:

-Hopt+

2γH 2

Hopt+[ ]3

Example 13.2 Calculate the optimum pH for precipitating Mn(OH)2 Assumethe water contains 140 mg/L of dissolved solids.

Solution:

Therefore,

Solving by trial and error, let Y = 2.30(1028)[ ]3 + 2.51(1017)[ ]2

spMn in mg/L Assume the water contains 140 mg/L of dissolved solids

–

-0.94

γMn 10

0.5 2 ( ) 2 [ 3.5 10 ( 3) ] 1+1.14 3.5 10 3

–

=2.30 10( 28) Hopt[ + ]3

7

( ) 4.81 10 5

–4.81 10( 5) 2.74 10 7

– -

13.4.2 P RACTICAL O PTIMUM pH R ANGE

From Example 13.2, at 25°C and at a solids concentration of 140 mg/L, the optimum

pH for the removal of manganese is 11.97 The corresponding concentration for

spMn is 0.0179 mg/L As in the case for the removal of ion, there is a practical pHrange at which units used for the removal of manganese can be operated and effectgood results Tables 13.6 and 13.7 show the respective concentrations of spMn at

other conditions of pH and total solids The values for [spMn] were obtained usingEquation (13.21)

Again, as in the case of the removal of iron, the tables show that total solids (orequivalently, the activities of the ions) do not have a significant effect on the optimum

pH value for the removal of manganese as Mn(OH)2 The optimum pH remains ataround 12.0 For the practical optimum pH, we adopt the following:

manganese removal as Mn(OH)2: 11.5 ≤ optimum pH ≤ 12.5,

which is the same as that for ferrous removed as Fe(OH)2

13.5 OXIDATION OF IRON AND MANGANESE

TO REDUCE PRECIPITATION pH

To summarize, the optimum pH’s of precipitation are as follows: ferrous removed

as Fe(OH)2, 11.5 ≤ optimum pH ≤ 12.5; ferrous removed as Fe(OH)3, 5.5 ≤ optimum

pH ≤ 12.5; and Mn removed as Mn(OH)2, 11.5 ≤ optimum pH ≤ 12.5 The reasonfor the high pH values for the removal of iron as Fe(OH)2 and for the removal ofmanganese as Mn(OH)2 is the formation of the complex ions These ions are FeOH+and Fe for the ferrous ion and Mn(OH)+, Mn(OH)0, and Mn for the

( ) 10– 14

+

-=4.5( ) 10– 14

10–6.8 - (4.5) 10– 14

( ) 10– 14

0.94( ) 10– 7.8

26 –

3.74 10( –18)

- 2.84 10( 7) 4.5 10( –28)

1.5 10( –20) -+

=3.27 10( 7) geq/L 0.0179 mg/L Ans

TABLE 13.6

Concentration of spMn as a Function of pH at 25 °°°°C

pH, Dissolved Solids = 140 mg/L [spMn ], mg/L

pH, Dissolved Solids = 35,000 mg/L [spMn ], mg/L

manganous ion The values of the equilibrium constants of these complex ions aresuch that their concentrations diminish the hydroxide ions needed to precipitate the solidhydroxides Fe(OH)2 and Mn(OH)2 For these solids to precipitate out, more hydrox-ide ions must be added, resulting in a high pH If these complex ions were destroyed,then the metal ions released would precipitate as the respective hydroxides, prevent-ing the complex equilibria to occur.

To destroy the complex ions requires the use of oxidants Chlorine, potassiumpermanganate, and ozone are normally used for this purpose In the case of iron,the ferrous state is oxidized to the ferric state This oxidation includes the oxidation

of the ferrous complexes Because Fe(OH)3 precipitates at a wider optimum pHrange, the precipitation pH can therefore be reduced to a much lower value, i.e., toeven a pH of 5.5

Let us tackle the situation of precipitating manganese at a lower pH range Oncethe complex ions have been destroyed, the reactions left would be those for a pre-cipitation product of a higher oxidation state than 2+ The reactions for the destructionusing chlorine, potassium permanganate, and ozone are, respectively, as follows:

(13.23)(13.24) (13.25)The previous reactions show that manganese is oxidized from Mn(II) to Mn(IV).This oxidation poses the possibility of Mn(IV) forming a complex with the hydrox-ides The review of the literature, however, did not uncover any evidence for this to

be so It did reveal that Mn(IV) forms a complex with fluorine and potassium(Holtzclaw and Robinson, 1988) This complex is K2[MnF6] The review also didnot reveal any solubility product constant for MnO2 It could very well be that itdoes not have any If, in fact, MnO2 does not have any solubility product constant,then the previous reactions can be construed to be possible at any pH Until a study

is done to show the complex formation of Mn(IV) and the accompanying solubilityconstant and to be compatible with the removal of ferrous as Fe(OH)3, we will adoptthe practical pH for removal of manganese through its oxidation by the abovereactions as 5.5 ≤ optimum pH ≤ 12.5 This range is, however, arbitrary and applieswhen the complexes have been destroyed For more accurate results, a pilot plantinvestigation for a given raw water should be undertaken

13.6 UNIT OPERATIONS FOR IRON AND

MANGANESE REMOVAL

The unit operations involved can be divided into two general categories: removal atthe pH range of 11.5 to 12.5 and removal at the pH range of 5.5 to 12.5 The former

range is called the high pH range and the latter is called the low pH range The

latter is arbitrarily categorized as low, because it is possible to adjust the operation

at a low pH value In these ranges, both iron and manganese can be removed at the

Mn2++Cl2+2H2O→ MnO2 s( )↓ 2Cl+ −+4H+3Mn2+ 2KMnO4 2H2O→ 5MnO2 s( ) ↓ 2K+

4H+++

3Mn2++O3+3H2O→ 3MnO2( ) ↓ 6Hs + +

same time and at the same unit The differences and similarities of these categories

of operations will be explored

13.6.1 H IGH pH R ANGE

Under this category, iron is removed as Fe(OH)2 and manganese is removed asMn(OH)2.The unit operations will then involve a unit (tank) for containing theaddition of the hydroxide source (lime) to raise the pH and for mixing the chemicalsand allowing completion of the precipitation of Fe(OH)2 and Mn(OH)2 The precip-itates formed will still be in suspension; thus, it is necessary for them to flocculate

by adding a flocculation tank After the flocculation tank will be a settling tank Thesettling tanks removes the flocculated solids A filter then follows to filter out anysolids that were not removed by the settling tank

The sequence of unit operations discussed above may not all be needed in agiven application What will actually be used depends upon the concentrations ofiron and manganese in the raw water If the concentrations are small such as lessthan 10 mg/L, the resulting floc concentration will be small; the chemically treatedwater, in this instance, may simply be introduced directly into the filter, withoutusing any flocculation or settling basin In all cases, however, the mixing chamberfor the reactions to take place and the filter are always necessary The unit operations

of mixing, flocculation, settling, and filtration were already discussed in the earlierchapters of this book

As can be deduced from the addition of lime, the unit operations above are thesame as those that would be applied in water softening Thus, for reasons of economics,iron and manganese removal are normally incorporated into water softening and allthe unit operations involved are identical In fact, iron and manganese are hardness ions

13.6.2 L OW pH R ANGE

Under this category, iron is removed as Fe(OH)3 and manganese is removed asMnO2, which involves oxidation of the metal ions to their higher oxidation states.The unit operations under the previous category are applicable under the presentcategory, only that under the present scheme, oxidation of the metals is involved.The unit process involves one of Eqs (13.23) through (13.25) for the case ofmanganese and similar equations for iron The equation designed into the unitoperations depends upon the unit process contemplated The unit process part,however, will only be at the mixing tank where the reactions should occur

As in the previous category, the unit operations of flocculation and settling maynot be necessary Again, in all cases, the mixing chamber for the reaction to takeplace and filtration is always necessary The reaction does not immediately produceparticles capable of being filtered, therefore, sufficient detention time should beprovided in the reaction chamber to allow the particles to grow into filtrable sizes.This normally ranges from 20 to 30 minutes In any case, a pilot plant may benecessary to determine the exact detention time

When oxidizing iron and manganese using dissolved oxygen, the process isusually carried out under catalytic reaction on some contact surfaces To accomplish

the reaction, the water is generally made to trickle over small rock surfaces such aslimestone, coke, or pyrolusite (MnO2) Pyrolusite possesses high catalytic power,springing from the iron and manganese oxides precipitated from the raw water thatcoat over and around the rocks These coatings act as catalysts between the reaction

of the ferrous and manganous ions in the water and the oxygen from the air as theycome into intimate contact over the contact surfaces of the rocks

The contact medium sizes vary from 3.5 to 5 cm The accumulated flocs producedfrom the deposition of the hydroxides or oxides are periodically flushed out by rapiddrainage This is done by filling the tank containing the bed of rocks to capacityand quickly releasing the water

13.7 CHEMICAL REQUIREMENTS

The chemical requirements are those for the hydroxide source, chlorine, nate, ozone, and oxygen The discussion will be subdivided into requirements in theferrous reactions and into requirements in the manganous reactions The treatment

permanga-on chemical requirements, in effect, is reduced to the determinatipermanga-on of the equivalentmasses of the pertinent chemicals

13.7.1 R EQUIREMENTS IN THE F ERROUS R EACTIONS

Under these requirements are the reactions at the high and low pH ranges At thehigh range, the hydroxide reaction is

From the previous equivalence, the equivalent of mass of NaOH is 2NaOH/2 = 40and the equivalent mass of CaO is CaO/2 = 28.05

Chlorine, permanganate, and ozone, will only be used at the low pH range There

is no need for them to be used at the high pH range, because under these conditions,the ferrous ion will be forced to precipitate as the ferrous hydroxide It is true thatafter precipitating, chlorine, permanganate, or ozone could then be used to oxidizeferrous hydroxide to the ferric state, but no practical reason exists for designing theunit process this way Thus, the unit process design will be to oxidize the ferrous

Fe2++2OH−→ Fe OH( )2 s( ) ↓

2NaOH ⇒ 2OH−

Ca OH( )2 ⇒ 2OH−