Ozone chemistry

LITERATURE

2.3 Ozone chemistry

Chapler 2 Lilernture

For an on-line sensor, the ozone loss due to decay in the sampling tine must be reduced in order to obtain near real-time response (Rakness et al., 2000). The sampling detention time should be short enough to ensure that the measured residual is at least 900;. of the acrual residual inside the contactor (Figure 2-4).

, ..

•• " _'\ ^r--:: _r"::. ~ ^-- --

•• ~ ~ -... _f---

^{HL 4mi}

.. ^, ,.

:2

_• ••

~

• ,.

0

2

0

..

• •

~

"

•

~

0

"- "---

^{I uL~ 2rlli}

" ^~

^{Omi "-}

^{' I:!!-}

^p

^"-

^lmn

^{h --}

Ozone H~f·Life

_ _ ~l _ 0.' .. i.

~ "---

, -

... ~l _ I ... i.

t--..

'" ^" ..

^__ -~L ^~L _ ^{• .I} 4 m^mill ill

^-...

• • , . .. .. .. .. ..

^{100 110 110}

Detention time in sample line (sec)

Figure 2-4: Reductions of measured ozone residual due to ozone decay and sample line HDT (Rakncss et al., 2000, p242)

On a full-scale ozonation plant, commercial sensors are currently available to measure the ozone residual.

Additional equipment may be necessary if a point within the contactor is required.

ChRptcr 2 L!I~nlrure

2.3.1 Chemical and physical properties

At room temperature ozone is a blue gas, diamagnetic and unstable. It is readily detectable at concentrations as low as 0.02 to 0.05 ppm (vlv) and appears colourless in dilute fonn (US EPA, 1999).

When the concentration reaches 240 J-tg.kL" in air, it is considered a nasaltoxicant at the lowest observable level (Hoigne, 1998).

Ozone has a high standard reduction potential in both acidic and basic media (Brady and Holum, 1993).

Scheme 2·2 shows the tendency of ozone to transfer an oxygen atom during reaction, producing oxygen gas.

EO - + 2.08V EO-+I.24 V Scheme 2-2: Redox potential of ozone (Brady and Holum, 1993)

This high reactivity makes ozone a health hazard, for example, as a pollutant in smog. Exposure to ozone at a level of 1.0 ~g.kL·1 for over la min is dangerous for all people (Brady and Holum, 1993). When applying this high reactivity appropriately, ozone is a powerful oxidising agent. Amongst the commonly used chemicals in water treatment, ozone is only second to the hydroxyl free radical, OH·, in tenns of the oxidation ability (US EPA, 1999). Therefore any unutilised ozone in the gas stream requires further treatment (such as a destruction facility) before releasing the spent gas to atmosphere.

Ozone is generated in gaseous fonn. The mechanism and equipment of ozone generation was discussed in Section 2.2.1 (page 2·6). Therefore ozone must be dissolved before aqueous reactions can take place.

The solubility characteristics of ozone have been studied extensively (Roth and Sullivan, 1981; Bablon et al., 1991a; Rischbieter et al., 2000). In literature, the solubility data are commonly reponed either in terms of the solubility, S, or as an absorption coefficient,

p.

Both are related to the dissolved concentration of ozone in water at saturation, C ...

S is often expressed as the so--called solubility ratio as milligrams per litre of ozone in water to milligrams per litre of ozone in gas (Bablon et al., 1991a).

p,

also known as Bunsen absorption coefficient, is the volume of gas expressed at NTP, dissolved at equilibrium by a unit volume of liquid at a given temperature when the panial pressure of the gas is the unit atmosphere (Bablon et al., 199Ia). In general, ozone dissolution is accepted to follow Henry's law, shown in Eqn. (2.1):

(2-1) where P_j is the partial pressure of the solute,}, in the gas phase, expressed in atmosphere (atm); H is called the Henry's law constant, in units of atm,mole fraction"; and Xj' is the mole fraction of the solute in the liquid phase.

The dissolved ozone concentration is related to the ozone panial pressure at thennodynamic equilibrium {Lee el al., 1999). Thus CS<l1 can be related 10 the panial pressure of ozone Pj at a given temperature.

Under thermodynamic ideality, the expression for the C_ in kg.kL,l can be written as:

2·14

Chapter 2 l.ltennu~

C~", "" pMP} (2-2)

where M is the density of gas at standard temperarure and pressure (2.14 g.e¹ for ozone).

Hence the solubility data of gases are also found to be reported in the Henry's law constant. Data of thc common water treatment oxidants using air. oxygen (in aeration) or chlorine (Cl_{l ),} chlorine dioxide (CIo,) (in chlorination) arc compared to the values of ozone in Table 2-3.

Table 2-3: Solubility data ohar-ious gaseous oxidants (LUey et aI., 1997; Bablon et aI., 1991a)

GUCOU5 olltidants p

c.

Hit 20"C

,N mg.L·' for P_J - I (11m.mol fraction")

Air 66400

Chlorine 4.54

, ....

^{1.773 i} CI,.L·' at 100 mmHg

Chlorine diolltide %6' ISOOOO 4.3a ClO:.L·' at 5% by yol.

Olltygcn 0.049 70 400'

Ozone 0.64 1400 J760

Ozone is sparingly soluble in wilter (jJ"" 0.64) but more soluble than oxygen (jJ - 0.049). At 100% ozone, only 1400 mg.L·! is achieved at saturation; whereas chlorine, under the same condition, is about 10 times more soluble.

Roth and Sullivan (1981) and Rischbieter et al. (2000) have critically reviewed the experimental data on ozone's solubility in water and both found the data considerably scattered. This can be attributed to the decomposition of ozone, pH or temperature ofthe water. Table 2-4 listed values of the Henry's constant of ozone reported in literature at various temperatures. Although some incongruence of the data may have been contributed by the earlier, less accurate analytical methods (Roth and Sullivan et al., 1981), the dependence of solubility on the temperature can still be marked.

Table 2-4: Effect of temperature on ozone solubility (Bablon et aI., 1991a; Roth and Sulllvan, 1981)

Klwlmur. (1932) Stumm (1958)

o

Kirk-Othmer(1967) 2590 Bablon et al. (l99la) 1940

"

''''

3150 3330

319.

H (alm.mole rractloo")

" "

³⁰

"

4610 6910 9520

3930 4BO

4310 6210 8500

.. ^"

^{655S 8302}

The scattering of data can also be attributed to experimental difficulties related to the decomposition of ozone which also depends on a number of variables such as pH (Rischbieter et aI., 2000). However the effect of pH is often suppressed by using a buffered aqueous system in order to measure the effect of temperature. In literature, the combination of temperature and pH or the temperature alone constitute the input variables co the prediction of Henry's constant of ozone.

2.3.2 Reactions

Ozone reactions in water are generally known to be complex. Figure 2-6 depicts that dissolved ozone undergoes reactions with various dissolved compounds (M) in two ways:

2-15

Chaptcr2 Lltcrature

;t undissolved Ol (g) 0, (g)

-++ . ... . . :~ ^.. f. .. - .... - ... - .,

^{--1:.+" , -}

^{M ---1~M}

^•• direct reactions decomposition

l< ... . .

· · ·· ···· ··· l· ···" ··· ~~ai n re~:~i·~·n·

...

secondary __ _ c+",M,,-_ _

+ • •

^M'~

oxidants radical.type reactions

Figure 2-t): Aquatic ozone reacCions (adapted from Hoigne and Badcr. 1983a)

• Direct reactions: with molecular ozone; selecljve reactions to give Mox as the oxidation product.

• Radical-type reactions: with secondary oxidants; extremely fast (in JlS scale) but less selective reactions to generate M' ox as the oxidation product.

The dashed arrow signifies the chain reaction induced by some oxidation products which catalyse the ozone decomposition.

The kinetics of the ozone reactions are typically first·order with respect to ozone and first-order with respect to the dissolved compound. Therefore the rate equation can be wrinen as (Hoigne and Bader.

1983a; Bablon et al .. 199Ia):

0) + M ---+ Products

d[O,]

- - -

~'o,

[0, )[M ]

dt

(2-3)

(2-4)

Ozone decomposes during the ozonation process and transfonns into secondary oxidants such as·OH . the hydroxyl radical. These arc highly reactive oxidants but are generally consumed in a fast and non-selective manner. For examples, almost all dissolved organic compounds present act as an OH sink.

This renders the radicaly·type reactions to be relatively inefficient for oxidising target compounds which may be present at very low concentration.

In natural waters, these aquatic reactions are affected by various parameters such as pH. temperature and dissolved compounds in the water matrix. In the treatment process, ozone is generated in the gaseous form which requires dissolution methods prior to its aquatic reactions (as depicted in Figure 2-6). Therefore the complexity of reactions is compounded due to the influence of mass transfer of ozone to water. The classification of ozone reactions in literature often incorporates parameters to account for the effect ofmass transfer (Beitran, 1995). The rate of ozone transfer may become a limiting factor when ozone is consumed quickly. However. in the present investigation, the transfer of ozone is assumed to occur mainly in the static mixer before the combined flow of water and the ozone·containing gas reaches the contactor (see Section t.2.2 ). The present section will thus focus only on the intrinsic kinetics relating ozone to other constituents in the water matrix.. The effects of mass transfer will be briefly discussed in Section 2.4.2 .

2-16

Ch:apler 2 Literature

The kinetics of ozone decomposition is oudined in Section 2.3.2.1 . An extensive list of ozone reactions with prescribed compounds can be found in literature. For the purpose of this thesis, only the pertinent categories of compounds in the respective ozonation functions (discussed in Section 2.1.2 ) will be discussed.

2.3.2.1 Decomposition

Ozone decomposition occurs through a series of reactions, resulting in the formation of short-lived radical species (Bablon et aI., 1991). While ozone reacts more selectively with certain constituents in water, other radicals are responsible for the less restricted reactions. Various reaction mechanisms have been postulated for the decomposition of ozone.

The Hoigne. Staehelin and Bader Mechanism (Bablon et aI., 1991 a) presents a consistent set of chain reactions. which accounts for initiation, propagation and break-in chain reactions for ozone decomposition al near neutral pH values.

Initiation step Propagation steps

0) + OH-~ .HO} + '02 .H0₂

r====

'02 + H~

0] + '02 ~ ^{.0]' +02} O• H+~ O· H- ' )+ ....--:---. 2+

•. >

'HO)~ ·OH+Oz Dj'

+

.OH ~ .HO~

·HO~ ----=.!....t

•

.H0₂

+

02

Termination steps .HO~+ .HO~ ~ H₂0₂+ 20]

.HO~+ 'HO]~ H₂0} +0]+ 02

where M is the molar concentration (mol.L·¹⁾

*1"" 7.0x 10¹

(M",,·')

pK" "" 4.8

*2 =1.6x109

(M",,')

k) =5.2xI0^1O

(M".")

k.

3

^=2.3xI02

(S·I)

*4 =l.Ix10$

(,' )

k5 = 2.0x 109

(M",S')

*6

= 2.8x 10

4 (s' )

*7

=5.0x109

(M" ",)

ks =5.0xI09

(M".")

Scheme 2-3: HSB mechanism of ozone decomposition (Sablon et aI., 199111).

[I]

[I') (2) (3 )

(4) [s) (6) (7) (8)

[n pure wafer, these radical reactions are initiated by the reaction of ozone with the hydroxide ions, 0..- (Staehelin et al., 1984). From the rate constant values, it can be seen thatlhe free-radical initiating step is the rate-determining step in the mechanism. Moreover the generation of hydroperoxide radical .H02 and superoxide radical '02 in Scheme 2-3[1] from the hydroxyl ion OH- indicates the stoichiometry is unity and o(her species capable of consuming (he hydroxyl ion without generating the two radicals bring stability on the ozone molecule in water (Bablon et aI., 1991 a).

However, the HSB mechanism shown in Scheme 2-3 introduces some transient compounds, -HO) and .H04 as chain intennediates. Although Staehelin and co-workers demonstrated the existence of

·H0₄ with the aid of computer simulation (Staehelin et al.. 1984), this posrulated species has not been

2·\7

Chapter 2 Literature

measured directly (Bablon et al.. 1991a; Westerhoff et al., 1997). The validity of the model also depends on a controversial equilibrium value for the ozonide radical,

·03"

(WesterhofTet al., 1997).

Scheme 2-4: Initiation step of the TGF mechanism

Tomiyasu, Gordon and Fukutomi, have proposed another reaction scheme (the TGF mechanism) that involves a two-election transfer process or an oxygen atom transfer from ozone 10 OH- (Bablon et al., 1991a) (shown in Scheme 2-4), as opposed to the one-electron transfer in the HSB model.

The TGF model does not include the posrulaled intennediates as in the HSB model. However, the mechanism proposed in the TGF model was only verified experimentally in basic mediwn (Tomiyasu et al., 1985; Westerhoffet al., 1997) and further experimemai work may be necessary to demonstrate its validity.

Both mechanisms indicate the decomposition increases with alkalinity as ozone is initiated only by OH""

(Staehelin et al., 1984; Tomiyasu et al., 1985). Staehelin and Hoigne (1982) found that OJ decomposition was slow due to the accumulation ofHzOz at pH below 8 and the rates were measurable if pH was below 10.

At pH above 10, decomposition was extremely fast.

The decomposition is also dependent on temperature. At pH 7, the rate increases with increasing temperature from 10 tQ 40 GC and becomes exceedingly fast when the temperature is above 30"C and at a pH above 8 (Sotelo et al., 1987).

In non-pure water, ozone decomposition may also be affected by the lype of solute present in water (Staehelin and Hoigne, 1985; Bablon et al., 199Ia). Inorganic compounds (such asOH-, .HOz and some cations) and organic compounds (such as glyoxylic acid, fonnic acid and humic substances) and ultraviolet radiation at 253.7 nm are all capable of inducing the foonation of ·Oi (Bablon et al., 199Ia), which assist in the initialion step in Scheme 2-3, reaction [I].

Some inorganic (e.g.: phosphate species) and organic (e.g.: primary alcohols, humic acids) act as the promoters oflhe free-radical reaction. They regenerate the superoxide anion,.Oi. from OH-. On the other hand, those compounds which scavenge OH- without generating the chain carrier .Oi are called the inhibitors to the chain reactions. Some common inhibitors include bicarbonate and carbonate ions, teniary alcohols (e.g.: ten-butyl alcohol), alkyl groups and humic substances (denoted as HS in Figure 2-7) (Staehelin and Hoigne, 1985) which consume OH"" but do not produce the chain intennediates. Such solutes quench the radical-type chain reaction and in some way stabilise the aqueous ozone (Hoigne, 1998).

The overall mechanism in non-pure water is shown in Figure 2-7.

2-18

Chapter 2 Luer.lture

Initiation:

OH-, H0₂^-, Fe²*.

CHOO-, UV, HS ...

H, O

Promotion:

0 ,

^{DJ. HS, ...}

,_

Inhibition:

.r-

^{COl-, HCO}-,}

I-BuOH, HS ...

0, ·OH

j--(

0 ,

Figure 2·7: Mechanism diagram of ozone decomposition Induced by radical· type reaction (Bablon Ct al.,199Ia)

Various rate equations have been suggested for the decomposition of ozone. Since the strong influence of 01"-was observed in the mechanisms discussed, the reaction order is often related to OH-as well as OJ.

Values of the reaction order with respect to ozone typically ranges from 1.5 - 2 in acidic environment to 1.5 - I in basic medium; with respect to hydrox:ide ion from 0 to I (Sotelo, 1987).

Most natural waters that undergo ozonation contain other solutes which may compete with ozone to react with Oa-(Hoigne, 1998). The reaction products also ox:idise further by ox:ygen to fonn perox:y and ox:y radicals, and eventually release ·0 2 or .H0₂ or H₂0_{2 •}

In summary, the decomposition of aqueous ozone occurs more rapidly at higher pH or in presence of elevated concentrations of compounds which behave as the initiators or the promOlers in the radical-type reaction scheme. The rate is impeded when pH is decreased or compounds which scavenge Oa- but do not produce chain-carrier radicals, such as bicarbonates.

2.3.2.2 Inorganic compounds

In general. the reaction of ozone with inorganic compounds present in water follows a first-order kinetic law with respect to ozone and the compound itself as described by Eqn. (2-4). The rate constants are affected by pH as many inorganic solutes are only reactive when not protonated (Hoigne et al., 1985). A selected range of inorganic compounds that pose problems in the water treatment effort will now be discussed in more detail.

Reactions with Iron and mangancse. Iron and manganese are two of the most abundant elements present in nature. Both iron and manganese exist in water in the reduced fonns (Fe²• and Mn²) which are soluble and cause unpleasant effects of staining and colouring (Lin. 1993).

Cha.pt~r 2 l.ucnlture

The removal of iron and manganese are usually achieved by oxidising the ions into the respective insoluble forms as indicated in Scheme 2-5.

2Fe²+ ° l - °2

~

^{2Feh H20 I Fe(OH))}

J.

^0.43 mg O)/mg Fe

Mn^l'" 0, - °2

~

^{Mn4... H20}

~

^Mn0₂

J.

0.88 mg Ol/mg Mn

Scheme 2-5: The oxidation reaction of Iron and manganese with ozone (Bablon et al., 19918; Bablon et al., 1991b)

Iron can be easily oxidised by ozone directly as shown in Scheme 2-6 [I]; by the hydroxyl radical from the ozone decomposition product [2]; or possibly react with the ozonide, ·0)" (3).

Fe^h

+

0) ----t Fe³+

+

.O~

Fe² ... +·OH

--+

OH- + Fe)'"

Fe²+ +.0)" + H₂0 ----t Fe)'" + 02

+

20H-

[I]

[2]

[l]

Scheme 2-6: Mechanism responsible for the oxidation of iron by ozone and its related products The removal of manganese is more difficult than iron as indicated from the stoichiometry in Scheme 2-5.

Per unit manganese requires approximately double the amount of ozone (0.88 mg) as iron (0.43 mg) (refer to Scheme 2-5). The oxidation of manganese can be severely impeded when other substances in water exhibit faster rates of ozone consumption.

Reactions with reduced non-metal species. Although it is thennodynamically possible for the halides to be oxidised by ozone, the reaction rate varies greatly from very fast for iodides to negligible in the case of chlorides (Hoigne, 1998; Bablon et al.. 1991). The rapid reaction of ozone with iodide shown in Eqn. (2-5) is often applied for the analysis ofozonated solution when the concenttation is high.

(2-5)

Free bromine (Br2). hypobromite (BrOl. free iodine (h). hypoiodite ion (10 ) and iodate ion (101) are some of the halogen-containing oxidants that can be produced when bromide or iodide ions are present in ozonated waters (Bablon et al.. 199Ia). Apart from these oxidants, the presence of bromide ion in the raw water or the use of chlorine as a secondary disinfectant can lead to the fonnation of halogenated disinfection by-products (OBPs). The halogenated OBPs such as trihalomethanes and haloacetic acids are suspected to be carcinogens and arc therefore a health concern (Pryor and Freese, 2000). In particular, the brominated OBPs appear 10 pose a greater health risk than non-brominated OBPs (US EPA. 1991). Ozone has been found to be effective in reducing the OBPs precursors (Richardson et al.. 1999; US EP A. 1991).

Ozone reacts with the inorganic fonns of sulphur fairly fast. The reaction rates decrease with the degree of anion protonation (Bablon et al., 1991). Sulphide is oxidized to sulfate by the general reaction shown in Eqn. (2-6). This reaction is of importance when treating sulphide-containing groundwater.

S2-+ 40) - - + SO!-+ 402 Ammonia can be oxidised by the following reaction in Eqn. (2-7»:

NH} +40) ----+NOj" + 402 + H30~

(2-6)

(2-7)

2-20

Chapler 2 LitcruNre

The rate of reaction is very slow and it decreases with pH due to the protonation of ammonia as ammonium ion (Hoigne. 1998). The ozonation of surface waters may enhance the nitrification of ammonia as the organic compounds are converted into more biologically degradable types.

2.3.2.3 Organic compounds

Ozone reacts with dissolved organic pollutants. narural or synthetic. In general. ozone reactions with the substrate M.,.., are first·order with respect to M .... and ozone:

d[O ,] [ ]

-~=ko ^•. o~[O,] MM, ^(2-8)

The reactivity ofa target pollutant must be high in order for degradation to occur. For organic pollutants.

the functional groups in their molecular structures influence their reactivity towards ozone. The reaction sites are either multiple bonds or atoms carrying a negative charge, attracting electrophilic oxidants such as ozone. Radical reactions are therefore not selective.

Natural water may be polluted by narural substances (e.g.: amino acids and carbohydrates), products or wastes related to human activity (e.g.: solvents. hydrocarbons and pesticides) or compounds produced due to high concentration of organic matter. The presence of organic solutes can influence the ozone consumption and reactions pathways in several ways: the direct consumption of ozone (e.g.: phenol); the initiation of ozone decomposition by producing ozonide and ·OH; the inhibition of ozone decomposition by ozone-scavenging species; and the promotion of Ol decomposition by ozone-scavenging species with production ofsuperoxide radical ions (e.g.: formate).

Many specific pollutants have been studied and data may be found in literature. Bablon et al. (1991) and Hoigne and Bader (19838 & 1983b) have compiled extensive lists of the organic waler pollutants and their reaction with ozone. In natural waters, analysis of organic pollutants is usually limited by their diversity and the analytical methods required for quantification of the pollutants, especially at low concentrations.

Therefore raw water is generally analysed for the priority pollutants. Some of the more commonly encountered problems with organic pollutants are described in details.

Reactions with odour and taste causing compounds. Amongst the micro-pollutants. control of tastes and odours is a common problem for water treatment works (Ho et al., 2002; Morioka et al., 1993; Glaze et aI., 1990). The occurrences of the odour- and tasle-causing compounds are as a resuh of the microbial activities in rivers. reservoirs or the treatment plant. Compounds like geosmin and 2-MIB are detectable at concentrations of nanograms per litre. The removal of the taste and odour is accomplished by altering or breaking the organic structure of the source compounds through oxidation. Because of the bulky chemical structure of geosmin and 2-MtB (shown in Figure 2-8), only strong oxidants such as ozone can oxidise these compounds at reasonably low dosages.

2·21