Cellulose 7: 263–286, 2000.c© 2000Kluwer Academic Publishers. Printed in the Netherlands.

Alkaline hydrogen peroxide bleaching of cellulose

ROBERT E. BROOKS and SAMUEL B. MOORE∗Burlington Chemical Company, Inc., P.O. Box 111, 615 Huffman Mill Road, Burlington,NC 27215

Received 16 February 2000; accepted 20 June 2000

Abstract. A closed system bleaching apparatus was designed to determine the kinetics andeffects of various factors on alkaline hydrogen peroxide bleaching of textile cellulose fabrics.It was confirmed that perhydroxyl anion is the primary bleaching moiety in alkaline hydrogenperoxide systems. The use of the apparatus in the measurement of fabric color, waste oxygen,and the subsequent calculation of hydroxyl ion, and molecular hydrogen peroxide confirmedthat pH and titration of ‘free’ hydrogen peroxide in alkaline bleaching systems are not goodindicators of bleaching mechanism. The role of the cellulose itself in the chemical bleach-ing system was determined. The rate of bleaching on cotton fabric was shown to be a firstorder reaction in concentration of perhydroxyl anion at 60 and 90◦C. An activation energyof 17 kcal/mole was estimated. Decomposition of H2O2 into waste oxygen was found to besecond order kinetics.

Key words: bleaching, cellulose, hydrogen peroxide

Introduction

The bleaching of cellulose by alkaline hydrogen peroxide is an establishedindustrial process. This does not mean however, that the mechanism of al-kaline peroxide bleaching of cellulose is well elucidated. A review of thepertinent literature indicates continued disagreement on both the mechan-ism of bleaching and the appropriate techniques to measure the chemicalchanges in aqueous alkaline peroxide solutions in contact with bleachable cel-lulose (Simon and Drelich, 1946; Evans and Uri, 1949; Duke and Haas, 1961;Goodman and Wilson, 1962; Koubeket al., 1963; Taher and Cates,1975;Steinmiller and Cates, 1976; Isbellet al., 1981; Galbacs and Gangi, 1983;Evans and Upton, 1985; Dannacher and Schlenker, 1996; Spiro and Griffith,1997; Thompsonet al., 1993, 1994, 1999; Milne, 1998).

Textile, paper, and the domestic and commerical laundry industries areall interested in producing the highest bleaching effects on cellulose withoutdamaging the cellulose polymer. To achieve this ‘holy grail’ of bleaching

∗ Author for correspondence.

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requires the ability to precisely measure many chemical parameters of al-kaline peroxide solutions in contact with cellulose.

Goals

In order to clearly define the chemistry and action of alkaline bleaching, thefollowing measurements are necessary:

• Amount of oxygen (O2) released during bleaching,• Bleach bath pH,• Hydrogen peroxide and sodium hydroxide concentrations during bleach-

ing,• Fabric ‘whiteness’ i.e., bleaching effectiveness,• Temperature.

These measurements required the fabrication of a gas tight bleaching sys-tem and ‘real time’ measurements during the bleaching cycle. From thesemeasurements, calculation of the following factors is necessary to model thebleaching process:

• kinetics of fabric bleaching (whiteness over time)• kinetics of the decomposition of alkaline H2O2 into oxygen• Perhydroxyl anion, hydroxylanion, and undissociated H2O2 concentra-

tions

The research described in this paper discusses some relevant literatureon bleaching with alkaline H2O2. It describes the apparatus and techniquesdeveloped to measure the factors discussed above, and the calculations ne-cessary to determine the species involved in bleaching performance. In theconclusions it lays the groundwork for further investigations into the effect-iveness of auxiliary chemicals in textile bleaching.

Literature review

Bleaching of cellulose with H2O2 has been extensively studied, but problemsremain in quantification of the bleaching species versus whiteness of fabrics.As a general consensus, the perhydroxyl anion is proposed as the primaryspecies responsible for bleaching and is a product of the reaction betweenH2O2 and NaOH in aqueous solution. However, debate continues on the rolesplayed in bleaching by free radicals, nascent oxygen, or some other form ofoxygen.

Taher and Cates (1975a) proposed a mechanism where the hydroxyl radi-cal OH· and the perhydroxyl radical OH2· forms a chain reaction, which

265

results in the decomposition of H2O2 into oxygen and water. Bleaching oc-curs when the fabric reacts with these radicals and thereby acts in a chaintermination role. Steinmiller and Cates (1976a) later examined the effect ofdifferent fabric treatments and were able to demonstrate the presence of freeradicals by initiating the polymerization of methyl methacrylate in a bleachbath, but the existence of ‘free radicals’ alone does not confirm their role inbleaching.

Dannacher and Schlenker (1996a) rejected the role of the perhydroxylanion in bleaching because of a belief that there is an optimum pH beyondwhich bleaching decreases. But these authors also point out that this has notbeen substantiated experimentally. This pH dependence of H2O2 bleachingsuggests that the perhydroxyl anion is not the active bleaching agent. Theypurposed instead that the active oxygen is the superoxide radicalO

−2 , formed

in an alkaline medium from the perhydroxyl radical. The relationship theyuse to estimate the concentration of this radical is not clearly defined but hasa lot of similarity to that used by Duke and Haas (1961a) to measure thedecomposition of H2O2.

Spiro and Griffith (1997a) claim that there is no need to involve any otheroxidizing species other than HO−2 and H2O2. The extensive investigationsby Thompsonet al., (1993), (1994) and (1999a) were concerned with thebleaching of certain dyes by H2O2 in aqueous solution and did not includefabric. The need to include fabric in any study of textile bleaching was clearlydemonstrated by Simon and Drelich, (1946a).

In many of the references, measurements of H2O2 concentration are made,but many without cotton fabric being present. In actual textile and paperbleaching production situations, the perhydroxyl anion is not measured but in-stead total H2O2 available is determined by titration under acidic conditions.pH and total alkalinity are also measured during the course of bleaching, butthis data is not used to calculate perhydroxyl anion concentrations, whichmust be determined to define bleaching efficiency.

Our belief is that the perhydroxyl anion formation is the key to textilebleaching and that the degree of whiteness obtained is directly related tothe concentration of perhydroxyl anion in the bleaching solutions. There aretwo competing reactions of H2O2 in the bleaching bath: first, there is thereaction with colored bodies on the fabric to give bleached fabric and second,decomposition to form oxygen. An enclosed apparatus was developed thatcan measure the release of gaseous oxygen and the whiteness of bleachedfabrics. The design of the apparatus and its operation are described in theexperimental section.

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Experimental

Calculations of perhydroxyl anion in alkaline H2O2 bleach solutions wereperformed and measurements were made of the effects on bleaching perform-ance on cotton, cellulose based textile fabrics.

Hydrogen peroxide was provided by Solvay Interox as Ultra Pure 31% andwas diluted with deionized water for use as a stock solution of about 0.7 M.This stock solution was then standardized against potassium permanganatejust before each use and diluted accordingly.

Sodium hydroxide was reagent grade, certified 0.5 or 1.0 N and was di-luted and used as needed. Aliquots of each of the above stock solutions weretaken by use of a Rainin2 automatic pipette, which could be adjusted, from 0to 10 ml in 0.1 ml increments.

De-ionized water was used in all dilutions and great care was used toprevent contamination by trace metal ions.

To insure complete wetting and to aid in the removal of entrapped air, anonionic wetting agent, Neodol 25–20, (Shell Chemical, USA) was addedfrom a stock solution at a level of 0.10% in each bleach run. This surfactantis a primary alcohol (C12–C18) with 20 moles of ethylene oxide per mole ofalcohol.

Fabric

Desized but unbleached, 100% cotton fabric was used to measure the rate ofbleaching. This fabric was obtained from Test Fabrics, Inc., West Pittson, PA,as Style 428-U, Army Carded Sateen. The fabric as received had a uniformlight tan color. Each bleaching experiment used a piece cut 2 inches by 4inches and weighed on average 1.21–1.22 g. Analysis of this fabric gave 0.2%total ash content and indicated the following concentration of metals in mg/kgas seen in Table 1.

Apparatus

All bleaching experiments were conducted in an 125 ml Erlenmeyer flaskplaced in a constant temperature bath controlled to±0.5◦C. Stirring duringbleaching was provided by a magnetic stir bar in the flask placed over a sub-mersible variable speed magnetic stirrer. An overhead glass column providedan outside addition funnel connected though a stopcock to a tube dipping tobelow the surface of the liquid in the bleach bath for peroxide addition afterassembly and reaching temperature. This column also had a stopcock to a vent

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Table 1. Fabric analysis metal concentrations

Element mg/kg

Ca 528.0

Mg 166.0

Al 24.0

Fe 12.0

Mn 1.6

Zn 1.3

Cu 0.5

tube, which could be opened to allow the peroxide addition at atmosphericpressure. At the top of the column was a pressure transducer to detect theoxygen pressure generated. The output voltage was converted directly to mmHg by a DP25-S Strain Gage Meter from Omega Engineering, Inc., Stamford,CT. The total overhead volume was kept to a minimum and was calibrated foreach flask for the purpose of calculating the moles of oxygen produced. Theoverhead temperature was monitored during each run. Figure 1 illustrates thisapparatus.

Conditions

The conditions for each experiment were set as follows unless otherwisestated: total volume of bleach bath 120 ml; one piece of fabric 2 inches by4 inches; temperature of the bath 90◦C; time of measurement of oxygenpressure every 2 min for 12 min; time of fabric bleach 12 min; time of equi-libration at temperature before the addition of H2O2 30 min.

Procedure

Each flask was charged with de-ionized water, nonionic surfactant and therequired amount of NaOH. A sample of loosely rolled fabric was then addedto the flask. The flask was placed in the water bath and an overhead condenserwas attached to the flask. The flask was brought to temperature and allowed toequilibrate to remove any air trapped in the fabric. After 30 min the flask wasmoved and connected to the overhead column through a spherical glass joint(without silicone grease). A Teflon O-Ring sealed this joint and was held inplace by a locking clamp. The required amount of H2O2 was added with the

268

Figure 1. Experimental bleaching apparatus.

vent tube open. Immediately all stopcocks were closed and the timer was star-ted. Oxygen pressure was measured every 2 min for a total of 12 min. After12 min, the fabric was removed and rinsed thoroughly in running cold water.The fabric was air dried over night and the K/S (whiteness) value measured.

pH measurement

The pH was not measured continuously throughout the procedure because ofthe constraints of the procedure and equipment. The bleach bath was cooledin an ice water bath immediately after termination of the 12 min bleach cyclesand the pH measured at 25◦C. This final pH is slightly different than the start-ing pH because a small amount of H2O2, a weak acid, was used for bleachingand oxygen production.

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K/S measurement

All cotton samples were dried at ambient conditions for 12 h prior to reflect-ance measurements. The reflectance of the fabric was read on a SpectraflashSF-600 reflectometer by Datacolor International with the readings convertedby the instrument to the Kubelka and Munk K/S value. An ‘unbleached’and a ‘complete bleach’ fabric sample were included in each set of experi-mental bleached samples as standards. The ‘unbleached’ sample was fabrictreated 120 min at bleaching temperature in a bath of 0.1% Neodol and 0.10 MNaOH. The ‘complete bleach’ sample was treated the same as above with0.10 M H2O2. For a quantitative measure of ‘color matter’ the followingequations were used:

Ao =[K

S

]o

= ‘Unbleached’ Fabric

A∞ =[K

S

]∞= ‘Complete Bleach’ Fabric

As =[K

S

]s

= ‘Experimental Sample’ Fabric

Ct , the ‘color’ remaining in a bleached sample after time,t , is equal to:

Ct =[As − A∞Ao − A∞

](1)

Oxygen measurement

Immediately after the H2O2 was added to the bleaching apparatus at temperat-ure, the pressure was recorded every 2 min for a total of 12 min. This pressurerepresents the total pressure in the apparatus, which is composed of air, watervapor, and oxygen generated in the decomposition of the peroxide. A plotof the logarithm of the pressure versus time is linear after the initial up-setconditions. Extrapolation back to zero time gives the logarithm of the pres-sure correction, log(Po), which is converted, back toPo. This is backgroundpressure and corrects for air, water vapor, etc. and is subtracted from everytotal pressure reading to give an estimate of the partial pressure of oxygenproduced by the decomposition of H2O2 to form gaseous oxygen. While anestimate of the oxygen partial pressure is used, it is the rate of change that we

270

use to calculate the rate constant(k) and would be independent of a constanterror in the absolute partial pressure.

The partial pressure of oxygen is converted to moles of oxygen by use ofthe ideal gas laws and apparatus over-head volume and temperature. Molesof oxygen are multiplied by two to give moles of H2O2 used for oxygenproduction by the stoichometry of the Equation (2)

2H2O2 −→ 2H2O+O2. (2)

This defines the quantity of H2O2 used to produce oxygen and is subtractedfrom the initial concentration of H2O2, Ao, to giveAt , the concentration ofhydrogen peroxide left at each time,t . Finally the reciprocal of this H2O2 con-centration is tabulated for each measurement time. The integrated expressionof the second order kinetic rate equation is:

1

At− 1

Ao= kot. (3)

Plotting 1/At versus time gives straight line with a slope ofKo, the observedrate constant. The overall rate constant,k1, is calculated using the followingequation (see derivation in the oxygen kinetic section)

k1 = ko

Keq[OH−]. (4)

As a matter of experimental practice, reaction times are kept as short as pos-sible to produce an initial rate where reactants and pH remain effectivelyconstant.

Bleach kinetics

If bleaching were dependent upon the concentration of the perhydroxyl anion,then the equations for bleaching cotton fabrics would be represented by:

OH− + H2O2� HO−2 + H2O. (5)

Followed by the rate controlling step

C+ HO−2 � Bleached fabric. (6)

Where the symbol C is intended to represent the reaction component of thecoloring matter in the fabric. The mechanism of bleaching as depicted inEquation (6) is not well understood. Some of the chromophoric bodies oxid-ized during the bleaching process are undoubtedly conjugated double bonds.

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The mechanism of oxidation proceeds to alcohol formation via an inter-mediate epoxide as suggested by Milne (1998a).

The rate equation for the bleaching process in Equation (6) may bedescribed as follows:

−dC

dt= koverall[HO−2 ][C] . (7)

If the bleach experiment is conducted under conditions where a large excessof peroxide and caustic are maintained over the amount of ‘color’, then thekinetics will be pseudo first-order in terms of ‘color’ removal. Let

kobs= koverall[HO−2 ], (8)

Then

dC

dt= −kobs[C], (9)

which integrates to

ln

[Ct

Co

]= kobst, (10)

wherekobs is the specific rate constant for the particular concentration ofperoxide and caustic being used. Then a plot of the experimental observedkobs versus the perhydroxyl anion concentrate should be linear with the slopeequal to the overall rate constant,koverall, as given in Equation (8).

Oxygen kinetics

The decomposition of hydrogen peroxide may proceed by both a chain anda non-chain process as observed by Isbellet al., (1981a). The chain reactionis thought to be started by radicals mainly generated by metal ions such asiron and copper (Koubeket al., 1963a; Galbacs and Gangi 1983a; Thompsonet al., 1993, 1994, 1999b). We used every precaution to avoid the introductionof such ions, and no chelating agents were used.

In the absence of metal ions, a non-chain process occurs when one moleof H2O2 and one mole of perhydroxyl anion form an intermediate complex,which decomposes to molecular oxygen. The formation of this intermediateis the rate-controlling step. This is the mechanism suggested by Duke andHaas (1961b) and later by Goodmanet al. (1962a).

In alkaline solutions this can be represented by a series of reactions:

OH− + H2O2

keq

�HO−2 + H2O, (11)

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H2O2+ HO−2k1

�K2

HO− + H2O+O2, (12)

H2O2kion H+ + HO−2 . (13)

From reaction Equation (11), the equilibrium constant can be calculated:

Keq= [HO−2 ]

HO−[H2O2]= Kion

Kw, (14)

where

Kw = [OH−][H+] and, (15)

Kion = [H+][HO−2 ]

[H2O2], (16)

for the rate controlling reaction Equation (12), the rate equation is calculatedby

dO2

dt= +K1[H2O2][HO−2 ]. (17)

The stoichiometry of the overall equation is obtained by adding Equation (11)and (12)

2H2O2→ 2H2O+O2 (18)

and thus

dH2O2

dt= −2

dO2

dt. (19)

The factor of minus two(−2) is accounted experimentally when we calculatethe moles of H2O2 used, from the oxygen pressure,(

dH2O2

dt

)cal. = dO2

dt. (20)

Substitution of Equation (20) into Equation (17) gives the rate equation interms of hydrogen peroxide.

dH2O2

dt= k1[H2O2][HO−2 ]. (21)

Substituting perhydroxyl anion (HO−2 ) from Equation (14), this becomes

dH2O2

dt= k1[H2O2]2[HO−]Keq. (22)

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At a fixed starting concentration of NaOH and H2O2, the concentration of[OH−] is determined completely by the equilibrium constant and is thereforefixed. Let

ko = k1Keq[OH−].

Substituting and separating variables Equation (22) becomes

dH2O2

[H2O2]2= ko dt. (23)

This is now a second-order equation in H2O2 and may be integrated to give

1

[H2O2]t− 1

[H2O2]o= kot. (24)

A plot of 1/[H2O2]t versust is linear with a slope equal toko. Rearrangementof the equationko = k1Keq[OH−], from above, gives the overall rate constantk1.

k1 = ko

keq[OH−]. (25)

To determine the change in rate as a function of pH, [OH−] = Kw/[H+] issubstituted into Equation (25) and recall that,

Keq=[

HO−2[HO−][H2O2]

]= Kion

Kw

from Equation (14). This gives

ko = k1× Kion

[H+]. (26)

Taking the logarithm of both sides of the equation and rearranging gives

logK1 = logk0− logKion− pH. (27)

Therefore, the rate of oxygen produced should be a linear function of the pH.

Perhydroxyl anion calculation

It is necessary that the ionization constant,Kion, for H2O2 be calculated at thetemperature at which the bleaching is to be carried out. The ionization rateconstantKion is available in the literature for 25◦C but must be calculated forthe experimental temperature.

H2O2kion H+ + HO−2 (28)

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Kion = [H+][HO−2 ]

[H2O2]. (29)

Using the values given by Evans and Uri (1949a):

Kion = 2.24× 10−12@25◦C,

with a heat of reaction equal to+8.2 kcal/mol. This allows for the utilizationof the van’t Hoff equation for the calculation of theKion at temperatures otherthan 25◦C. Assuming the heat of reaction does not change much over thetemperature range.

LogK2− LogK1 = 1H

4.576

[T2− T1

T2× T1

], (30)

where

1H = Heat of reaction (cal),

T = absolute temperature(◦K).

Using this relationship the following ionizations constants were calculated:

Temperature Ionization constant

35◦C 3.51× 10−12

60◦C 9.49× 10−12

90◦C 2.66× 10−11

Dividing this ionization constant by the ionization constant for water,Kw =[H+][OH−], we obtain the equilibrium constant for the reaction:

OH− + H2O2

keq

HO−2 + H2O (31)

Keq=[

HO−2[HO−][H2O2]

]= Kion

Kw. (32)

This resulted in the following equilibrium constants:

Temperature Keq Kw

35◦C 168 2.09× 10−14

60◦C 99 9.61× 10−14

90◦C 62 4.30× 10−13

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whereKw is value as given in the Handbook of Chemistry and Physics, 66thEdn, page D-164.

Evans and Uri (1949a) also expected that the second dissociation of H2O2

to be extremely small due to the high endothermicity that they estimated,∼100 kcal for this step. Therefore, we did not include the reaction:

HO−2 � H+ +O−2 , (33)

with the above equilibrium constants, the concentration of perhydroxyl anionmay now be calculated from the starting NaOH and H2O2concentrations.

Let

N = Starting NaOH concentration in moles/lH = Starting H2O2 concentration in moles/lX = [HO−2 ] formed and is also equal to [OH−] reacted.

Then by charge balance

[HO−2 ] + [OH−] = [Na+] + [H+]. (34)

However, at pH above 10, [H+] approaches zero, and upon rearrangementEquation (34) becomes

[OH−] = N− X. (35)

By H2O2 balance

H = [HO−2 ] + [H2O2], (36)

which gives

[H2O2] = H− X. (37)

Where [H2O2] is undissociated hydrogen peroxide left at equilibrium. Then

Keq= [HO−2 ]

[OH−][H2O2]= X

(N− X)(H− X), (38)

solving forX gives

X = −R−√R2− 4NH

2, (39)

where

R

(−N −H − 1

Keq

), (40)

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from the value of X= [HO−2 ]

[H2O2] = H −X by Equation (37)

and

[OH−] = N− X by Equation (35).

From the value of [OH−]

[H+] = Kw

[OH−]

and

pH= log[OH−] − logKw,

whereKw must be for the same temperature as that at which the pH is meas-ured. Because of the large difference inKw at various temperatures, this canbe significant. For this reason and because of the importance of pH in thediscussion of the kinetic of bleaching, we calculated the pH at both ambientand at bleach temperature.

The importance of calculating the perhydroxyl anion concentration canbe illustrated by the following examples. Because of the interaction of thehydroxyl ion and H2O2, a weak acid, different solutions where either OHor H2O2 may be in excess might have the same pH and yet have a 10-folddifference in perhydroxyl anion concentration. For example:

Starting Calculated

NaOH H2O2 [HO−2 ] [H2O2] [OH−] pH@60◦C pH@25◦C

0.075 0.010 0.0087 0.0013 0.0663 11.84 12.83

0.150 0.100 0.0864 0.0136 0.0636 11.82 12.80

Conversely, two solutions may give the same approximate concentrationof perhydroxyl anion and have a different pH.

Starting Calculated

NaOH H2O2 [HO−2 ] [H2O2] [OH−] pH@60◦C pH@25◦C

0.050 0.080 0.0399 0.0401 0.0101 11.02 12.00

0.075 0.050 0.0391 0.0109 0.0359 11.57 12.56

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This data clearly demonstrates why plotting reaction rates versus pH can-not describe bleaching results. The more preferred correlation is between thespecific rate constant and perhydroxyl anion concentration. pH alone doesnot adequately define the concentration of various ionic forms of hydrogenperoxide and hydroxide ions present. This is addressed in the performedexperiments.

Some past investigations have used buffers to maintain a constant pH toavoid this issue. If the purpose is to study the effects of various species onbleaching, then these reactants must be analyzed or calculated by the use ofequilibrium constants. The use of a buffer will prevent this calculation of theperhydroxyl anion.

Fabric bleaching results

Bleaching experiments were made varying the concentration of NaOH andH2O2 [one at a time]. Table 2 demonstrates no correlation between bleachrate and the concentration of these reactants.The specific rates constant forbleaching was plotted against the pH to try to verify the work of (Dannacherand Schlenker, 1996b). This is shown in Figure 2. There seems to be no directcorrelation between pH and the reaction rate, particularly at higher pH wherewe are most interested.

The results from the experiments listed in Table 2 demonstrates the needto know the concentrations of the three species, hydroxyl, perhydroxyl, andundissociated H2O2. Titration for H2O2 under acidic conditions, as is usuallyperformed, in bleaching is of no help since this does not reveal the concen-tration of the species present under alkaline conditions of bleaching. Usingthe procedure given earlier, the concentrations of each of these species wascalculated from the starting concentrations of NaOH and H2O2. Also calcu-lated was the pH the solution would have at 90◦C. These results are given inTable 3. The calculated concentration of perhydroxyl anion [HO−

2 ] is plottedagainst the specific rate constants, in Figure 3. The data and calculationsfrom the bleach experiments show the linear relationship expected from themechanism proposed in the kinetics section of this paper. The slope of thegraph is equal to the overall rate constant,koverall, as given in Equation (10).

koverall= 27.2× 10−3 l/mol s@90◦C.

To get an estimate of the effect of temperature, a limited number of bleachruns were made at 60◦C. These results are in Table 4.

Again, the specific rate constant for bleaching was plotted against thepH to look for direct correlation. As shown in Figure 4, there is a general

278

Table 2. rates at 90◦C; rates in l/mol s usingKeq= 61.78@90◦C

Run # Start NaOH Start H2O2 25◦C pH 25◦C pH Bleach -kxmeasured calculated 1000

15 0.025 0.037 11.34 12.02 1.093

16 0.025 0.044 11.30 11.96 0.815

17 0.025 0.052 11.22 11.90 0.907

18 0.025 0.059 11.13 11.85 0.870

19 0.006 0.044 10.40 11.26 0.491

20 0.013 0.044 10.83 11.60 0.691

21 0.019 0.044 11.10 11.80 0.927

22 0.025 0.044 11.32 11.96 0.918

23 0.031 0.044 11.49 12.09 0.966

24 0.038 0.044 11.64 12.19 1.046

25 0.044 0.044 11.80 11.93 0.957

26 0.050 0.044 11.94 12.37 1.041

27 0.063 0.044 11.94 12.52 1.104

28 0.075 0.044 12.10 12.63 1.095

29 0.088 0.044 12.20 12.73 1.082

30 0.100 0.044 12.33 12.81 1.255

31 0.050 0.044 11.80 12.37 1.071

32 0.025 0.059 11.07 11.85 1.043

41 0.100 0.059 12.16 12.74 1.856

42 0.100 0.074 12.12 12.66 1.895

43 0.100 0.089 12.06 12.58 2.098

44 0.100 0.103 12.00 12.50 2.273

45 0.075 0.059 11.96 12.54 1.625

46 0.075 0.074 11.88 12.45 1.733

47 0.075 0.089 11.76 12.36 1.975

48 0.075 0.103 11.65 12.28 2.060

Table 2 shows the starting concentration, together with the measured final pHand the specific rate constantskobs. Also in Table 2 is shown the calculatedvalue the pH of the starting mix would have if cooled to 25◦C.

increase in the rate as the pH increases. However, at pH values above 11.5,data is scattered and it would be difficult to draw any conclusions. When thesesame rates are plotted against the corresponding calculated concentrations ofperhydroxyl anion as in Figure 5, a linear relationship is revealed. The slopeof this line gives the overall rate constant for bleaching at 60◦C.

koverall= 5.5× 10−3 l/mol s@60◦C.

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Figure 2. Bleach rate (90◦C).

Finally the Arrhenius equation was used to estimate the activation energy,Ea,from the overall rate constant at 60◦C and 90◦C.

Ea = 17 kcal/mole.

This is close to the values of 16.4 and 16.7 kcal/mole given by (Thompsonet al., 1993, 1994, 1999c) for the bleaching of Alizarin and Crocetin by H2O2.These authors suggest that the mechanism of the bleach reaction is similar andinvolves formation of an epoxide.

Oxygen results

It is difficult to study the rate of alkaline decomposition of H2O2 to oxygenbecause the reaction is so sensitive to adventitious catalysis. Many materialscan act as catalysts, both homogeneous and heterogeneous. Even the natureof the reaction vessel and atmospheric carbon dioxide are reported to increasethe ‘spontaneous’ decomposition of peroxide. Perhaps the most unavoidableis the purity of the alkali used. Literature values for the uncatalysed decom-position of at H2O2 35◦C vary with the degree of purification from 7.4×10−4

l/mole s to 3.7× 10−8 l/mole s (1985a).The object of this investigation was not to determine the absolute rate for

the spontaneous decomposition H2O2 but to find the order of reaction withrespect to H2O2 and perhydroxyl anion and the dependency of the rate uponpH under the same conditions as the bleach experiments.

280

Table 3. Rates at 90◦C; rates are in liter/moles s usingKeq= 61.78@90◦C

Run # Bleachkx Calculated Calculated Calculated Calculated pH

1000 [HO−2 ] [H2O2] [OH−] @ 90◦C

15 1.093 0.145 0.022 0.011 10.390

16 0.815 0.159 0.028 0.009 10.330

17 0.907 0.017 0.035 0.008 10.270

18 0.870 0.018 0.041 0.007 10.220

19 0.491 0.005 0.040 0.002 9.630

20 0.691 0.009 0.036 0.004 9.970

21 0.927 0.013 0.032 0.006 10.170

22 0.918 0.159 0.029 0.009 10.320

23 0.966 0.019 0.025 0.012 10.450

24 1.046 0.022 0.023 0.016 10.560

25 0.957 0.024 0.020 0.015 10.660

26 1.041 0.026 0.018 0.024 10.740

27 1.104 0.030 0.014 0.033 10.890

28 1.095 0.032 0.012 0.043 11.000

29 1.082 0.034 0.010 0.054 11.100

30 1.255 0.035 0.009 0.065 11.180

31 1.071 0.026 0.180 0.024 10.740

32 1.043 0.018 0.041 0.007 10.220

41 1.856 0.046 0.014 0.054 11.100

42 1.895 0.055 0.019 0.046 11.020

43 2.098 0.062 0.027 0.038 10.950

44 2.273 0.068 0.035 0.032 10.870

45 1.625 0.040 0.019 0.035 10.910

46 1.733 0.047 0.027 0.028 10.810

47 1.975 0.052 0.037 0.023 10.730

48 2.060 0.056 0.047 0.019 10.650

Table 2 shows the specific rate contant,kobs, and calculated values.

The experimental oxygen data generated during bleaching could bemodeled on the basis of the equation for the second order rate constant,ko, asdeveloped in the section on oxygen kinetics. Table 5 gives this rate constant,ko, together with the overall rate constant,k1, and the measured final pH at25◦C.

The derived Equation (27) predicts that the logarithm of the overall rateconstant,k1, should be a linear function of the pH. Figure 6 shows a plot ofLog k1 versus pH. This graph gives a reasonable linear fit considering theexperimental difficulties in keeping out spurious catalytic decomposition.

281

Figure 3. K/S bleach rate (90◦C).

Table 4. Results for bleach study at 60◦C; all concentrations are in moles/l rate constant isin l/mole s

Second Start Start Calculated Calculated Calculated Measured Bleachkx

set run NaOH H2O2 [HO−2 ] [H2O2] [OH−] pH @ 60◦C 1000

1 0.015 0.050 0.012 0.038 0.003 10.770 0.583

2 0.020 0.050 0.016 0.035 0.005 11.030 0.664

3 0.030 0.050 0.022 0.028 0.008 10.330 0.653

4 0.050 0.050 0.032 0.018 0.018 11.820 0.687

5 0.080 0.050 0.040 0.010 0.040 12.220 0.846

6 0.080 0.080 0.056 0.024 0.024 12.020 1.154

7 0.100 0.080 0.063 0.017 0.037 12.220 0.955

8 0.100 0.100 0.730 0.027 0.027 12.090 1.000

9 0.120 0.100 0.080 0.020 0.040 12.280 1.035

10 0.150 0.100 0.086 0.014 0.064 12.460 0.886

11 0.170 0.100 0.089 0.011 0.081 12.580 0.942

12 0.200 0.110 0.100 0.010 0.100 12.640 1.191

Note that in general the specific rate constant,ko shows some increasewith pH, but does not pass through a maximum and decrease as has beenreported by some (Duke and Haas, 1961c). Our findings are more like thoseof Thompsonet al. (1993, 1994, 1999d) who found the rate continued toincrease even at pH above 12.

282

Figure 4. K/S bleach rate (60◦C).

Figure 5. K/S bleach rate (60◦C).

To understand the contribution made by the fabric, a series of experimentswere made at 90◦C. with no fabric in the peroxide bath. This data is presentedin Table 6 and a plot of the logarithm of the overall rate constant, Logk1,versus pH is shown in Figure 7. Again a linear relationship is shown to sup-port the mechanism purposed in the section on oxygen kinetics. Comparingthe rate constants at 90◦C with and without fabric is made difficult by thefact that while the starting range of alkali concentration was the same, the

283

Table 5. Rates at 90◦C; rates are in l/moles s usingKeq= 61.78 @ 90◦c sorted by pH

Run # Start Start Measured ko × 1000 k1× 1000 Logk1NaOH H2O2 pH

19 0.0063 0.0444 10.40 1.57 14.12 −1.85020

20 0.0128 0.0444 10.83 1.86 7.53 −2.12340

18 0.0250 0.0589 11.13 2.12 0.84 −2.31580

17 0.0250 0.0517 11.22 1.61 3.26 −2.48710

16 0.0250 0.0444 11.30 2.27 4.04 −2.39390

22 0.0250 0.0444 11.32 4.04 7.19 −2.14350

15 0.0250 0.0367 11.34 2.85 4.39 −2.35720

24 0.0375 0.0444 11.64 3.14 3.24 −2.48980

48 0.0750 0.1032 11.65 3.13 2.65 −2.57630

47 0.0750 0.0885 11.76 1.96 1.38 −2.86030

25 0.0438 0.0444 11.80 2.56 2.86 −2.88000

31 0.0500 0.0443 11.80 1.93 1.32 −2.54400

45 0.0750 0.0588 11.88 2.39 1.11 −2.95400

46 0.0750 0.074 11.88 2.51 1.45 −2.83990

26 0.0500 0.0444 11.94 3.37 2.31 −2.87400

27 0.0630 0.0441 11.94 2.75 1.34 −2.63610

44 0.1000 0.1032 12.00 3.07 1.57 −2.80480

43 0.1000 0.0885 12.06 2.86 1.22 −2.91430

28 0.0750 0.0441 12.10 2.77 1.04 −2.98180

42 0.1000 0.0739 12.12 2.76 0.98 −3.00800

41 0.1000 0.0592 12.16 3.42 1.02 −2.99240

29 0.0880 0.0441 12.20 3.43 1.03 −2.98880

pH is higher with no fabric in the bath. We do not have an explanation forthis fabric effect. However, there are two experiments that have the same pHof 12.2, (No. 29 with fabric and No. 58 without fabric). Comparing thesetwo runs, we find an oxygen decomposition rate of 1.03 with fabric and 6.81without. This indicates that the cotton fabric used in these bleaching trials actas a stabilizer for the peroxide. This supports the findings of Taher and Cates(1975b).

Conclusion

We have demonstrated an apparatus that is capable of independently measur-ing the oxygen produced and the color removed from a piece of fabric in a

284

Figure 6. Oxygen rate (90◦C).

Table 6. Rate constants 90◦C – No fabric sorted by pH

Run # pH k1× 1000 Log (k1× 1000)

58 12.22 6.81 −2.1671

57 12.28 7.23 −2.1408

56 12.31 6.4 −2.1940

51 12.57 3.35 −2.4743

59 12.61 3.53 −2.4522

50 12.66 2.32 −2.6344

55 12.73 2.18 −2.6607

54 12.75 2.18 −2.6614

53 12.78 2.56 −2.5925

52 12.83 1.98 −2.7042

bleach bath. In addition, procedures were developed that allow the calculationof the concentrations of the perhydroxyl ion, hydroxyl ion and undissociatedH2O2 from the starting levels of NaOH and H2O2.

The rate of bleaching of cotton fabric was shown to be a first-order reac-tion in concentration of perhydroxyl ion as calculated by this procedure. First-order kinetics was shown at both 60◦C and 90◦C with estimated activationenergy of 17 kcal/mole.

285

Figure 7. Oxygen rate; no fabric (90◦C).

The rate at which H2O2 decomposes into oxygen was found to be second-order kinetics. A mechanism involving the reaction of undissociated H2O2

and perhydroxyl ion was developed. The kinetics of this reaction was suppor-ted by the linear relationship of the overall rate constant and the pH of thebath.

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