Kinetics of Oxidation of Bromcresol Green Miles Pickering and David Heiler Princeton University, Princeton. NJ 08544 Many different clock reactions have been reported in the student experiment literature (1-141. Clock reactions do not require instruments and are also quirk enough so that many variations (temperature, nmcentration, ionic strength) can be studied in a sinele neriod. The maior disadvantaee is that in almost all cases there is a long period in which nothing . bv a sudden dramatic chanee. annears to h a n ~ e n followed which gives thkitudent the unrealistic idea that thereact,"" beeins suddenlv after an induction neriod. The chemistry is also necessaril; complicated since i t least two reactionsare involved. The system we report can be conveniently studied by clock reaction methods with the added advantage that there is a continuous change in color until the sampie matches a control. This eliminates the pedagogical difficulties of traditional clock reactions. This reaction is the bleaching of bromcresol green (BCG) bv hv~ochlorite.There is afast reaction. which is followed in tiis-eiperiment, and a slow reaction in which a colorless product is produced overnight. The fast reaction involves the change of color from the blue basic form of the dye to a yellow intermediate. Since BCG is an acid-base indicator, i t has both an acid form (yellow) and a basic form (hlue):
-.
Students can also reassure themselves that the color change is not simply a pH effect by using p H paper or a pH meter to see that the nH remains a t 7 even after the reaction is over. They can aiso show that the blue form will not reappear even if a lot of base (enough to overwhelm the buffer) is added. This proves that the intermediate has n o indicating power. The reagents for this experiment are unusually cheap and verv safe. Tbermostated baths are unnecessan for this experiment. Hypochlorite bleach can be purchased a t any supermarket and the indicator solution is very dilute. Buffers can be purchased, or made by following recipes available in The CRC Handbook of Chemistry and Physics. Experimental We have published a lab handout for this experiment (15)and will only summarize the procedure here. Each student should nrenare a solution of bleach (,5 mL of 5.25% sodium hvoochlorite soiutkn diluted with 296 mL of distilled wn.. ter,, and the stmeroom ahwhl yuvidr a stork sdution containing ahout O 4 p of HCC Isod~umsalt] per liter of distilled Water. Esrh student should prepare his or her own dilution of the bleach so that the match times are not identical. The control mixture for matching should contain 20 mL of a pH 4 buffer, 30 mL of distilled water, and a known quantity of dye (between 2 and 4 mL). The solution will be yellow with a trace of green. Its pH can be adjusted slightly as needed to amin this color. For the run under standard conditions, exactly 10 mL of diluted bleach is pipetted into a flaskof the same size as that containingthe control mixture. Then 10 mL of pH 7 buffer and 30 mL of HzO is added with mixing. The same number of milliliters of dye solution is added to this flask as to the control mixture, and time is counted from the instant of mixing. The color change should be immediately perceptible, and the match should be ettained in 25 to 60 s depending on how fresh the bleach is. This time can be shortened by adding more bleach to the diluted bleach solution or can be lenethened bv ddutingthroxidnnt ~crlution.Studentsneed cn hereminded that the matching is one ,of hue, not intrnsiry, olo,lor. The bleaching reactim is going on at t h e samr time, and this lowrrs the mtensity significantly. Students should try variations andshould be cautioned to try one variable at a time. The effects of doubling the amount of dye and doubling the amount of oxidant are worth trying. Acids and bases can be tried, as can salts (added as solids).Temperature can also he is added. studied bv warmine or cooline the mixture hefore the dve ,~ At theind ut our'handout &ral pussible rn~chanismsofreartion arp pr~puscd.The rtudent 1, asked to ronardcr thr consirtenry of these with the experimental data. ~
~~~~~~~
~
~~
~
~
~
~
P
basic
acid
The only structural difference between the acid and basic forms is the protonation of a phenolic oxygen.The two forms are in eauilibrium. and at OH1 about 90% is in theacid form. This experiment is posLible because the acid form of BCG and its oxidized form have the same color. Thus as the oxidation of the blue basic form takes place, the colors change first to green (a mixture of unreacted basic form and yellow oxidized form) and then t o pure yellow (all oxidized form). Any mixture of unreacted basic form and oxidized form will have the same color as a mixture containing a comparable ~ercentageof basic and acid form. This latter mixture ran he madesimply by adjusting the pH. For example, if we make a mixture of HCG and pH 4 buffer, the color will he that of SOYoacid and 1090basic form. Thiscolor will be the same a s t h a t of 90% o x i d i z e d k d 10% basic form. ~ h " ; the time t o attain a matchine will be the time reouired to oxidize 90% of the basic formof the dye. We first define a set of standard reaction conditions, and the student times the reaction under these conditions.' The standard conditions are then varied by changing the reactant concentrations. The student changes one variable a t a time, and, since it takes only a few minutes to do a run, many variations can be tried. Typical match times are 30 s, and the reproducibility of the judgment as to when the match is attained is almost always better than i3 s.
~
Re~Its Students show that the reaction is first order in hwochlorite. The time for matrhing is independent of the a k u n t of dvr added. This is cmsiscent with first-order behavior, since the match time is simply a multiple of the half-life, and for a pseudo-first-order system the half-life is independent of initial concentration. The activation energy is about 5.8 kcallmol (students ohtain values of 2-8 kcallmol). There is no kinetic salt effect; however, if enough spectator ions are added, the p H of the buffer will be affected slightly and the reaction speeded up by this indirect effect. The reaction is first order in chloride ions, is stopped by any source of ammonium ions, and is favored by increasing acidity. Volume 64 Number 1 January 1987
81
We ask students first to calculate the order of reaction and the activation energy. Then the salt effect is brought into plav. By being deliberately vague about how much salt is to be added, wecan create the situation in which the students have very different answers as to whether there is a salt effect. We then present a menu of possible attacking species. Some of these can be ruled out immediately on the basis of the salt effect since i t is a given that the dye is anionic. The remaining possibilities are then individually calculated. For each assumed intermediate. the student derives an equilibrium expression and sees if i t matches the data on concentration de~endence.This is verv hard even for our best students, b i t one could easily stop after determining the order of reaction and activation energy. There is only one major pedagogical difficulty in this experiment. The first explanation for the color change that occurs to most students is that the pH is changing in the reaction flask and that this, not oxidation, is causing the color change. I t is essential to focus the student's attention on the stability of pH throughout the reaction.
ready substituted, and it is this which allows the hv~ochlorite to survive. such behavior has been described foyaniline (20).I t is believed, however, that the rearrangement is inter. molecular (21). The nature of the attacking species is also unclear (16). Since the dye is anionic in both acid and basic forms and there is no kinetic salt effect, the species must be neutral. This argues for HOCI, or Clz. These species would be formed by the equilibria, Ht
and the HOCl would be formed by OCI-
B % crB & :Br
&-
We do not believe that the chlorine is bonded t o the carbon skeleton. First, it is unlikely that such bonding would be easily reversible in strong acid or base. Second, adding chlorine to the carbon skeleton would certainly affect the chromophore, and the oxidized molecule and the acid form would no longer have the same color. To test this hypothesis we have taken the infrared (IR) spectrum of the oxidized form and compared it with that of the model compound t-hutyl hypochlorite. The latter shows two new IR bands at 845 (medium) and 695 (weak), and the OH band a t 3400cm-' in thealcohol is absent. The spectrum of the oxidized dye shows a sharp peak a t 865, to be comoared with a broad ueak a t 900 in the acid form. The basic rorm of BCG does nbt show comparable peaks. Presumably the broadening is somehow the effect of hydrogen bonding. Phenol hypochlorites are a known (17-19) but only marginally stable species. They would normally be expected to rearrange to yield products chlorinated in the ortho position. In our molecule, both the ortho and para positions are al82
Journal of Chemical Education
+ H,O
F.
OH-
+ HOCl
The equilibrium constant for the first reaction is K, -
[CLI[OH-] [OCl-][CI-] [Ht]
and for the second, hydrolysis equilibrium,
Dlscusslon
As of this writing, we have not completely elucidated all of the underlying facts about the nature of this reaction. We susoect that even if we could do so, not all facets of the mechanism would be within the scope of a freshman course. However, if a menu of mechanistic choices is presented, students can usually pick out possibilities and rule out others, a useful exercise in itself. This discussion presents what we know about the mechanism a t present. Some work has been done in this laboratory to identify the intermediate. The intermediate will extract into tetrahydrofuran (THF) if the water layer contains enough ionic material. The yellow color persists in T H F overnight, rather than fading as in water. However, thin layer chromatography shows that the material soon hecomes a complicated mixture. If a T H F solution of the intermediate is allowed to remain in concentrated acid or base for a long period, the indicating urooertv will uartiallv reamear. w e h"elie,vithat tGe yeiiow intermediate is the phenolic hvoochlorite ester. which is verv similar to the yellow acid -. form.
+ OCI- + C1- F? C12+ OH-
Kh -
[OH-] [HOCl] [OCI-]
The rate-determining step will be, in either case: Rate = k[dye][attackingspecies] If the attacking species is HOC1, then [om-1 Rate = k[dye][HOCI]= kKh-[OH-l [dye] This argues that additional basirity slows the reaction and that the reaction shuuld he firit ordrr in hyporhlvrite (ohsewed) but does not explain the chloride iondependence. If the attacking species is CI,, however, the rate will be Rate = k[dye][CI,]
This would argue for a first-order hypochlorite and C1- dependence and a second-order base dependence. In fact, detailed studies of the H+ dependence show it to have a nonintegral order, probably arguing that the mechanism is complicated by pH effects on the dye itself by dissociation eauilibria of the HOCl or that there are competing ~ a t h w a y s ( i l ) . Because this point is unclear, we ask siudeui8 to eliminate possible mechanisms on the basis of experimental data rather than to propose mechanisms. Nevertheless, this experiment is a useful addition to student work hecause the operations required in the laboratory are verv simple, s e t in spite of the apparent simplicity the student is foiced-to engage in the major intellectual task of kinetics, that of testing possible mechanisms to see if they are consistent with exp&imental data. Literature Cited
1. Conway, W.J. J. Chem.Edue. 1940,17,398. 2. McA1pine.R. K. J. Chem.Educ. 1945.22.387-390. 3. Suryarmaman,M. G. J. ChemEduc. 1951,28,386387. 4. Klemm, L.M.J. Chem.Educ. l951,28,587.
5. Ami8.E.S. J. ChemEduc. 1951.28.635. 6. Euana. G. 0.J. Chem Edur. 1952.29.139.
7. K1emrn.L.H. J. Chrm.Edur. 1952.29.318. 8. B1ack.A. H.;Ddson,V. M. J.Chem.Edue. 1950.33.562. 9. Whib,E.H..I Chem.Educ. 1957.34,275. 10. Dutton, F. B. J. Chem. Edue 1951.34, A303. 11. Kauffman.G. B.; Haii,C. R. J. Cham.Educ. 1958.35.577. 12. Rarrett, R.L. J. Chem.Educ. 1955.32.78. 13. Jones, P.; Hsggett, M. L.; Longridge, J. L. J. Chrm. Edur. 1964.41,610. I 4 Moem, P. C.: Petrumi. R. M. J. Chem. Edur. 1964.41.549, 15. Pickering. M. ThoRediscouev Book: A Csnrral Chamistry Monuol; Ginn: 1986. 16. Swain. C. G.:Crist,D.R. J.Am. Chem. Sor. I972.94,3195. 17. Bo1tzo.K. H.;Dell,H. D.JuLvsLiebig's Ann. Chsm. 1967,708,6369. 18 Deli, H. D.: Kamp, R. JuaruaLiebig'a Ann. Cham. 1967.709,70-78. 19. M0ye.C. J.:SternhiB,S.Tef.Lott. 1964.2411. 20. Haberfold, P.; Pad, D. F.J. Am. Chrm. Sac. 1965.87.5502, 21. Paul, D.F.:Haberfeid.PJ.O,g. Chem. 1976,41,3120.
