Phsicochemical Treatment of Hazardous Wastes - Chapter 4 docx

This is because the radical reactionwill have a much lower activation energy barrier than regular reactions do;therefore, oxidants such as oxygen, hydrogen peroxide, and ozone are com-bi

is considered to be thermodynamically favorable The oxidation potentialsfor common oxidants suitable for environmental applications are listed inTable 4.1

As can be seen in the table, the hydroxyl radical has an oxidation tial of 2.80 V The hydroxyl radical is a short-lived and extremely potentoxidizing agent, according to its potential as shown in the table Becausethey are extremely potent oxidizing agents, hydroxyl radicals react withorganic compounds by three mechanisms: hydrogen abstraction, electrontransfer, and hydroxylation (Huang et al., 1993) From a thermodynamicpoint of view, the higher the oxidation potential is, the stronger the oxidantspecies will be

poten-Another factor is how fast the reaction is The fundamental theory lining the mechanisms involved in AOPs is the transition state theory (TST),which provides theoretical guidance for the search of the most efficient AOP.According to the TST, hydroxyl radicals may accelerate the oxidation rates

under-of an organic compound by several orders under-of magnitude compared withoxidation rates for common oxidants This is because the radical reactionwill have a much lower activation energy barrier than regular reactions do;therefore, oxidants such as oxygen, hydrogen peroxide, and ozone are com-bined with catalysts such as transition metals, metal oxides, photons, andultrasound to generate hydroxyl radicals

For each AOP, the degradation rate is investigated to search for the mostefficient process We begin this chapter with basic chemical kinetics followed

by discussion on the TST, oxidants, and catalysts used in AOPs

86 Physicochemical Treatment of Hazardous Wastes

O3 + 2H + + 2e = O2 + H2O 2.07

H2O2 + 2H + + 2e = 2H2O2 (acid) 1.76 MnO4 + 4H + + 3e = MnO2 + 2H2O 1.68 HClO2 + 3H + + 4e = Cl – + 2H2O 1.57 MnO4 + 8H + + 5e = Mn 2+ + 4H2O 1.49 HOCl + H + + 2e = Cl – + H2O 1.49

Advanced Oxidation Processes 87

The following is a typical first-order reaction:

(4.5)The rate law for the first-order reaction is:

(4.6)

Integrating this equation, the time dependence of concentration A becomes:

(4.7)The concentration profile of reactant A is:

−d = ′[ ] =d

A

t = −k A[ ]

−ln[ ]A = − +kt constant

88 Physicochemical Treatment of Hazardous Wastes

(4.11)

Assume the x moles of reactants A and B have been reacted and x moles of

C have been produced, then the production rate of C should be:

A

t = −k A[ ] [ ]B

dd

x

t =k A( 0−x) (1 B −x)

0 1

ln

Advanced Oxidation Processes 89

For the nth order reaction with respect to one reactant, the general solution

for rate expression is:

4.3 Transition State Theory

Chemical reactions are studied in terms of elementary reactions involving

only one step for bond breaking, bond formation, or electron transfer A

characteristic of elementary reactions is the molecularity In other words, if

0 0

11

1 0 1

90 Physicochemical Treatment of Hazardous Wastes

a reaction takes place in a single irreducible act at a molecular level without

any detectable intermediates, the reaction is called as an elementary reaction:

(4.23)

However, when reactants A and B are approaching each other, most of the

motions in a reacting molecular system are ordinary vibrations, rotations,

and translations Only one normal mode corresponding to the reaction

coor-dinates is involved in breaking or forming a chemical bond to form a new

molecule The new chemical bond results in the rearrangement of atoms

The collection of these atoms is defined as the reactive center

As the distance between the two atoms becomes shorter and shorter, the

electron clouds interact with each other due to the rapid motion of electrons

As a result, a multidimensional and continuous potential surface is

devel-oped The potential field becomes stronger and stronger as the two molecules

approach each other On this multidimensional surface exists a most

eco-nomic energy path for reactants A and B to interact with each other The

reaction path is referred as the reaction coordinate Along this coordinate, the

highest energy along the most economic reaction path defines an activated

energy can be seen in Figure 4.1 The energy between reactants A and B and

[AB]‡ is defined as the activation energy barrier, as shown in Figure 4.2

According to the TST, the following reaction will represent the formation

of the activated complex and subsequent production of C:

(4.24)

The AC usually has several characteristics The molecules going over the

barrier are in equilibrium with all the other reactant molecules, and the AC

can be treated as a normal molecular species except that one of its vibrational

modes is missing and must be replaced by translation along the reaction

coordinate Also, the rate of formation of product C through the AC is the

universal frequency ν = ; therefore, the reaction rate k can be expressed

FIGURE 4.1

Potential energy surface with a late transition state (From Boudart, M., Kinetics of Chemical

Processes, Prentice Hall, Englewood Cliffs, NJ, 1968 With permission.)

rBC

rAB

Transition State Complex

Furthermore,

(4.31)The equation takes the final form as follows:

(4.32)

Experimentally, the Arrhenius equation describes the relationship between

the rate constant k and the temperature as follows:

(4.33)

By taking the natural logarithms of both sides, the equation becomes:

(4.34)

FIGURE 4.2

Potential energy profile for elementary reaction (From Boudart, M., Kinetics of Chemical Processes,

Prentice Hall, Englewood Cliffs, NJ, 1968 With permission.)

E O,2 =EO,C+EO,D

Eo Activation Barrier Transition state: Z

Products: C+D Reactants: A+B

Reaction coordinate

Heat of reaction

The above equation suggests that the activation energy barrier can be imentally determined In addition, if the transition state equation is com-pared with the Arrhenius equation, the following will also be true:

exper-(4.35)

Because E a is the potential energy of activation, the difference between ∆H‡

and E a is the kinetic energy of activation, which is usually small compared

to E a For a bimolecular reaction, this difference is RT As a result, the vation energy will be the same as the change of activation enthalpy:

acti-(4.36)

It is important to point out that for exothermic reaction, the activation energy

∆E‡ equals the activation energy:

(4.37)However, for endothermic reaction, the activation energy equals the activa-tion barrier plus the heat of reaction ∆H:

(4.38)

The transition state theory indicates that the rate of a reaction is not a matter

of energy alone, but also requires a favorable configuration by a change ofentropy In addition, the rate of a reaction can be speeded up through thefollowing methods These methods are the guiding principles in the searchfor the most efficient AOPs:

• By increasing the temperature to increase the universal collisionfrequency ν = , where is the Boltzmann constant (1.38*10–23J•K–1) and h is the Planck constant (6.63*10–24 J•s); supercritical wateroxidation uses this method

• By increasing the ground energy of reactants using ultraviolet tons and ultrasound to reduce the activation energy barrier

pho-• By increasing pressure to increase the positive entropy of tion, In supercritical water oxidation, either air or oxygen isused as the oxidant Although the activation energy barrier isextremely high, ∆S is large due to extremely high pressure As aresult, the reaction is still fast enough to oxidize concentrated organicwaste

activa-• By decreasing the enthalpy of the reaction, ∆H‡

e∆S‡ /R

The last approach is the fundamental approach used in AOPs In terms ofthermodynamics, ordinary oxidants such as oxygen, ozone, and hydrogenperoxide will form activated complexes with organic pollutants with largeenthalpy, ∆H‡, and the reactions will be thermodynamically less favorablethan reactions between hydroxyl radicals and organic compounds; therefore,one way to increase reaction rates is to convert these oxidants to hydroxylradicals first As a result, the enthalpy change when the hydroxyl radicalattacks an organic molecule is several orders smaller than when a commonoxidant attacks an organic molecule; thus AOPs can be found by any com-bination of oxidants such as oxygen, ozone, and hydrogen peroxide withcatalysts such as UV photons, transition metals, and ultrasound Based onthese guiding principles in searching for AOPs, Table 4.2 provides possibleAOPs with different combinations of oxidants and catalysts

Because AOPs take advantage of the high reactivity of hydroxyl radicals,initial, propagation, promotion, recombination, and reversible reactions arecommonly involved in the degradation of organic pollutants Table 4.3 liststhese major elementary reactions

Possible AOPs with Different Combinations of Oxidants and Catalysts

Catalyst Metals and Ions

Metal Oxides Oxidants Photon

sound Electron Oxidant Fe 2+ Fe Pt TiO 2 Fe 2 O 3 OH – O 3 H 2 O 2 UV US

Ultra-e – (reductant)

Although the redox potential of oxygen is only 0.4 V, it can form oxygen anion (•O2) in water by a rapid self-redox reaction to O2 and H2O2,which is a two-electron reduction The third electron will further convert

super-H2O2 to •OH, which can be achieved by adding Fe2+ salts The fourth electronand one proton will produce water as the terminal reductive product of theabove reaction The pH effect on the thermodynamics of the reduction ofoxygen is summarized in Figure 4.3

4.4.2 Hydrogen Peroxide

Thenard identified hydrogen peroxide, H2O2, as a chemical compound in 1818(Schumb et al., 1995) It has been commercially available since the middle of19th century, but the scale of manufacture and use has increased rapidly sincearound 1925, when electrolytic processes for its manufacture were introduced

to the United States, and its use in industrial bleaching became increasinglyimportant Commercially, hydrogen peroxide is handled as an aqueous solu-tion in a wide concentration range Hydrogen peroxide has a covalent oxy-gen–oxygen bond and two covalent hydrogen–oxygen bonds All oxidativeproperties of H2O2 originate from these bonds Hydrogen peroxide is a liquid

at normal temperatures and has a melting point of –0.43°C, and a boilingpoint of 15.2°C at 1 atm Its commercial usage is in the form of an aqueoussolution; therefore, the rate constant in aqueous solutions is more importantthan that of the pure compound

FIGURE 4.3

Standard potentials for redox reactions involving oxygen at 25°C (From Sawyer, D.T and

Nanni, E.J., Jr., Redox chemistry of dioxygen species and their chemical reactivity, in Oxygen

and Oxy-Radicals in Chemistry and Biology, Rodgers, M.A.J and Powers, E.L., Eds., Academic

Press, New York, 1981, 15–44 With permission.)

+0.695

+1.229 +0.70

+2.813

O 2 -0.33 O 2 · +0.89 H 2 O 2 +0.30

+0.59 +1.20

+0.281

+0.815 +0.29

+2.40

O 2 -0.33 O 2 · +0.20 HO 2- -0.251

-0.03 +0.65

-0.065

+0.401 -0.13

+1.985

-+1.349

+0.867

4.4.2.1 Molecular Structure

The four atoms in a hydrogen peroxide molecule are structurally joined

by simple covalent bonds, H–O–O–H, in a nonpolar structure The structurecan be defined by four parameters: the O–O distance, the O–H distance, theO–O–H angle, and the angle between two planes, each of which is defined

by two oxygen atoms and one hydrogen atom The best values for theseparameters in the solid state are 1.453 ± 0.007 Å for the O–O distance; 0.988

± 0.005 Å for O–H distance; 102.7 ± 0.3 Å for O–O–H angle; and 90.2 ± 0.6

Å for the dihedral angle between the O–O–H planes The dihedral angleappears sensitive to its environment and thus may be different in the vaporphase or in other crystals containing H2O2 (Othmer, 1991)

Hydrogen peroxide is 100% soluble in water The freezing point of themixture depends on the percentage of hydrogen peroxide by weight Table4.4 lists the freezing point changes with the composition of the mixture The eutectics exist at 45.2 and 61.2 wt% H2O2, and the compound

H2O2•2H2O (48.6 wt% H2O2) exists in the solid state The evidence indicatesthat solid solutions are not formed in this system, although it is extremelydifficult to obtain water-free solid H2O2 The recommended value for theheat of fusion of H2O2 is 87.84 cal/g For the liquid/vapor phase relationshipfor aqueous hydrogen peroxide, partial pressures of the vapors over theliquid are lower than the calculated value for ideal solutions Table 4.5 shows

TABLE 4.4

Freezing Point of Hydrogen Peroxide

H 2 O 2 Concentration

(wt%)

Freezing Point (°C)

H 2 O 2 Concentration (wt%)

Freezing Point (°C)

Source: Othmer, D.F., Hydrogen peroxide, in Encyclopedia of Chemical Technology,

4th ed., Howe-Grant, M Ed., John Wiley & Sons, New York, 1991 With mission.

per-atmospheric boiling points and related liquid and vapor compositions Theheats of vaporization for aqueous hydrogen peroxide at 25 and 60°C areprovided in Table 4.6

The values for a number of thermodynamic properties of H2O2 can beobtained from chemical engineering data handbook The average heat capac-ity from 25 to 60°C for 100% H2O2 is 0.628 cal/g/°C) The mixing heat of

H2O2 and 100% water ranges from –590 cal/mol H2O2 at 0°C to –1110 cal/mol H2O2 at 75°C Table 4.6 presents the heats of vaporization for aqueoushydrogen peroxide (Othmer, 1991)

The formation heat of H2O2 from the reaction 4.40 at 25°C has been lated at –32.52 kcal/mol:

calcu-H2 (gas) + O2 (gas) → H2O2 (gas) (4.40)The free energy of formation of anhydrous H2O2 liquid at 25°C is calculated

as –28.78 kcal/mol The decomposition heat of pure liquid H2O2 to waterand oxygen at 25°C is –23.44 kcal/mol

Boiling Point (°C)

Vapor Composition (wt% H 2 O 2 )

resulting from simple decomposition The various reactions can be simplified

to five general types, as follows:

is very complex in many cases and may depend on the types of catalyst andreaction conditions

TABLE 4.6

Total Heat of Vaporization of Aqueous H2O2

Heat of Vaporization (cal/g solution)

H 2 O 2 Concentration

0 20 40 60 80 100

582.1 534.5 503.1 460.4 414.1 362.7

563.2 526.4 487.4 446.0 401.3 351.3

Source: Othmer, D.F., Hydrogen peroxide, in Encyclopedia of Chemical Technology, 4th ed., Howe-Grant, M Ed., John Wiley & Sons, New York,

1991 With permission.

4.4.2.5 Ionization

Hydrogen peroxide exhibits a weak acidic character, having a dissociationconstant of about 1.5 × 10–12; thus, pure aqueous solutions of hydrogenperoxide have pH values below 7 The dissociation constant of hydrogenperoxide is 10–6

(4.47)Depending on the pH of the medium, ionic species will have either nucleo-philic character (pH > 7) or electrophilic character (pH < 7) In an alkalinemedium, hydrogen peroxide reacts with hydroxide anions to give perhy-droxyl anions HOO– according to the equilibrium

(4.48)This equilibrium is shifted to the right when the basicity of the mediumincreases This shows the nucleophilic character of the peroxidic linkage O–Odue to formation of the HOO– anion This perhydroxyl anion is considered

to be a supernucleophile, as its reactivity is about 200 times higher than that

of the OH– anion The pH effect on the kinetic rate constant for the tation of superoxide is shown in Figure 4.4

Depending on whether the electron pair of the broken bond is shared or not

by the two new entities, the reaction sequence will involve either an ionic

or a free-radical pathway, as shown in Equation (4.49) and Equation (4.50)

H

OOH

HO

HO

O.

+ H.

H

OO

+